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F4 - CHEMISTRY (DLP) Flipbook PDF
Text Book Chemistry Subject for Form 4
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DUAL LANGUAGE PROGRAMME
CHEMISTRY Form 4
RM10.50 ISBN 978-967-2375-43-2
FT714001
CHEMISTRY Form 4
RUKUN NEGARA Bahawasanya Negara Kita Malaysia mendukung cita-cita hendak; Mencapai perpaduan yang lebih erat dalam kalangan seluruh masyarakatnya; Memelihara satu cara hidup demokrasi; Mencipta satu masyarakat yang adil di mana kemakmuran negara akan dapat dinikmati bersama secara adil dan saksama; Menjamin satu cara yang liberal terhadap tradisi-tradisi kebudayaannya yang kaya dan pelbagai corak; Membina satu masyarakat progresif yang akan menggunakan sains dan teknologi moden; MAKA KAMI, rakyat Malaysia, berikrar akan menumpukan seluruh tenaga dan usaha kami untuk mencapai cita-cita tersebut berdasarkan prinsip-prinsip yang berikut:
KEPERCAYAAN KEPADA TUHAN KESETIAAN KEPADA RAJA DAN NEGARA KELUHURAN PERLEMBAGAAN KEDAULATAN UNDANG-UNDANG KESOPANAN DAN KESUSILAAN (Sumber: Jabatan Penerangan, Kementerian Komunikasi dan Multimedia Malaysia)
KURIKULUM STANDARD SEKOLAH MENENGAH
CHEMISTRY FORM Writers Lim Kuok Chen Dr. Nur Jahan Ahmad Dr. Chua Kah Heng Wong Choy Wan Lee Sze Yien
Translators Dr. Chan Siok Gim Lim You Sie Lee Chwee Neo
Editors Amalina binti Ahmad Khairul Maisarah binti Kahar
Designer Loh Hong Seng
Illustrator Ai Khen bin Wong
PENERBITAN PELANGI SDN. BHD. 2019
4
KEMENTERIAN PENDIDIKAN MALAYSIA Book’s Serial No: 0176 KPM2019 ISBN 978-967-2375-43-2 First Published 2019 © Ministry of Education Malaysia All Rights Reserved. No part of this book may be reproduced, stored in any retrieval system or transmitted in any form or by any means, either electronic, mechanical, photocopying, recording or otherwise, without the prior permission of the Director General of Education Malaysia, Ministry of Education Malaysia. Negotiation is subject to the calculation of royalty or honorarium. Published for Ministry of Education Malaysia by: Penerbitan Pelangi Sdn. Bhd. 66, Jalan Pingai, Taman Pelangi, 80400 Johor Bahru, Johor Darul Takzim. Layout and typesetting: Penerbitan Pelangi Sdn. Bhd. Font type: Minion Pro Font size: 11 pt. Printed by: The Commercial Press Sdn. Bhd. Lot 8, Jalan P10/10, Kawasan Perusahaan Bangi, Bandar Baru Bangi, 43650 Bangi, Selangor Darul Ehsan.
ACKNOWLEDGEMENTS The publication of this textbook involves cooperation from various parties. Our wholehearted appreciation and gratitude goes out to all involving parties.
• Committee members of Penyemakan Naskhah Sedia Kamera, Educational Resources and Technology Division, Ministry of Education Malaysia
• Officers in Educational Resources and
Technology Division and the Curriculum Development Division, Ministry of Education Malaysia
• R. Weller/Cochise College
Contents Introduction v Laboratory Rules and Safety Measures viii Themes Chapter
1
Chapter
2
Chapter
3
1
Introduction to Chemistry
2 1.1 Development in Chemistry Field and Its Importance in Daily Life 4 1.2 Scientific Investigation in Chemistry 8 1.3 Usage, Management and Handling of Apparatus and Materials 11 Achievement Test 1 20
Matter and the Atomic Structure
22 2.1 Basic Concepts of Matter 24 2.2 The Development of the Atomic Model 29 2.3 Atomic Structure 32 2.4 Isotopes and Its Uses 37 Achievement Test 2 40 The Mole Concept, Chemical Formula and Equation
42 3.1 Relative Atomic Mass and Relative Molecular Mass 44 3.2 Mole Concept 49 3.3 Chemical Formula 59 3.4 Chemical Equation 69 Achievement Test 3 76
Chapter
4
The Periodic Table of Elements
78 4.1 The Development of the Periodic Table of Elements 80 4.2 The Arrangement in the Periodic Table of Elements 82 4.3 Elements in Group 18 84 4.4 Elements in Group 1 87 4.5 Elements in Group 17 92 4.6 Elements in Period 3 96 4.7 Transition Elements 101 Achievement Test 4 106
iii
Chapter
5
Chapter
6
5.1 5.2 5.3 5.4 5.5 5.6 5.7
Basics of Compound Formation Ionic Bond Covalent Bond Hydrogen Bond Dative Bond Metallic Bond Properties of Ionic Compounds and Covalent Compounds 122 Achievement Test 5 132
7
Rate of Reaction
7.1 7.2 7.3
Manufactured Substances in Industry 252 8.1 Alloy and Its Importance 254 8.2 Composition of Glass and Its Uses 260 8.3 Composition of Ceramics and Its Uses 262 8.4 Composite Materials and Its Importance 265 Achievement Test 8 273
The Data Table of Elements
276
Glossary
277
References Index
279 280
218 220 230
Determining Rate of Reaction Factors Affecting Rate of Reactions Application of Factors that Affect the Rate of Reaction in Daily Life 240 7.4 Collision Theory 243 Achievement Test 7 250
The Periodic Table of Elements 275
iv
108 110 111 114 117 120 121
134 6.1 The Role of Water in Showing Acidic and Alkaline Properties 136 6.2 pH Value 143 6.3 Strength of Acids and Alkalis 149 6.4 Chemical Properties of Acids and Alkalis 152 6.5 Concentration of Aqueous Solution 158 6.6 Standard Solution 162 6.7 Neutralisation 167 6.8 Salts, Crystals and Their Uses in Daily Life 174 6.9 Preparation of Salts 178 6.10 Effect of Heat on Salts 190 6.11 Qualitative Analysis 197 Achievement Test 6 216
8
Acid, Base and Salt
Chapter
Chapter
Chemical Bond
Introduction This Kurikulum Standard Sekolah Menengah (KSSM) Form 4 Chemistry textbook is written based on Form 4 Chemistry Dokumen Standard Kurikulum dan Pentaksiran (DSKP) that is prepared by the Ministry of Education Malaysia. In order to successfully implement and fulfil the needs of DSKP, this book is written based on three domains: Knowledge, Skills and Values through methods of inquiry. There are four themes to be discussed, namely the Importance of Chemistry, Fundamentals of Chemistry, Interaction between Matter and Industrial Chemistry. This book is equipped with various special features that focus on instilling Science, Technology, Engineering and Mathematics (STEM), thinking skills, scientific skills and computational thinking so that students can master skills that are needed for the 21st century and to be individuals who are science-oriented. Special features in this book are as follows:
Introduction to
Activity 1.2
21st Century Skills
CHAPTER 1
Chemistry
Activities that involve 21st century skills: • Thinking & problem-solving skills • Interpersonal and intrapersonal skills • Information and communication skills
Century
21 Skills chemistry and origin of the word interpreting the Collecting and information its meaning d media to gather Share activity. printe -Pairs Think variou the from 1. Carry out g materials et or refer to readin 2. Surf the Intern ing issues and discuss the follow chemistry word (a) Origin of the in front of your class. chemistry (b) Meaning of s in multimedia form discussion result 3. Present your group st
Drugs contain chemicals
y Life
Chemicals in Dail
activities of chemicals. The d us are made up Figure 1.2 All substances aroun involve chemical reactions as well. agriculture daily that we carry out in foods, medicine, commonly used icals chem the shows and industries.
Food • Preservatives • Colouring • Flavouring • Antioxidant • Stabilizers
Ch
ca emi
http://bit. ly/2Mx4M7B
Medicine • Antibiotic • Antiseptic • Vitamin • Chemotheraphy • Analgesic
ls in daily life
Acid, Base and Salt
Applications of Neutralisation in Daily Life
Figure 6.31 shows the application of neutralisation for a variety of uses Hair health
Medicine
Herbicide
Century
21st Skills
CHAPTER 6
in daily life. Agriculture
Dental health
Industry Shampoo
• Paint • Polymer • Glass • Ceramic • Detergent • Colouring • Alloy
E ST PA
H OT TO
Agriculture • Herbicide • Pesticide • Fungicide Milk of ia MagnesFertiliser • • Hormone
kan
gar
nye
Me
5
Figure 6.31 Applications of neutralisation in daily life
Activity 6.11 Solving problems on soil fertility using suitable fertilisers 1. Carry out this activity in groups. 2. Study the following problem statement:
STEM 21st Century Skills
Various activities that emphasise student-centered learning based on Higher Order Thinking Skills (HOTS).
Computational Thinking
als
used chemic Slaked lime, Ca(OH)2 a contains Toothpaste Commonly Weak alkali in the Figure 1.2 which is alkaline, is used base that neutralises shampoo neutralises to treat acidic soil. lactic acid produced by acid on hair. bacteria in our mouth.
Milk of magnesia Mg(OH)2 relieves gastric pain by neutralising the excessive hydrochloric acid in the stomach.
21st Century Learning Activities
CT CT
Activities that involve computational thinking: • Decomposition • Abstraction • Pattern recognition • Algorithm • Logical reasoning • Evaluation
replace nutrients Apart from treating acidic soil, fertilisers need to be added to soil to plants. There is a such as nitrogen, potassium and phosphorus that have been absorbed by variety of fertilisers in the market. Which fertilisers are suitable for plants?
3. Gather information concerning the problem given above. (a) What type of crops were planted? (b) What are the type of elements required by the crops? percentage of (c) Identify the fertiliser that is suitable for the crops by considering the the quantity elements such as nitrogen, phosphorus, and other needs, fertiliser cost and needed for the area. 4. Present your group findings in a multimedia presentation. as urea, potassium Neutralisation reaction is also applied in the production of fertilisers such can be produced from sulphate, K2SO4, ammonium nitrate, NH4NO3 and others. For example, urea about other fertilisers? the neutralisation reaction between ammonia and carbon dioxide. How Try to give the acids and alkalis involved in the production of that fertiliser. 169
STEM (Science, Technology, Engineering and Mathematics) STEM
Activities that involve project-based learning through STEM approach. STEM approach is based on teaching and learning that applies knowledge, skills and STEM values through inquiry, daily problem solving, environment, as well as local and global communities.
Information in the QR code on the front cover: • Explanations of book themes • Authors biodata • Updated information and facts (if any)
v
Components at the end of each chapter: Self-reflection
Chain Concept
Reflection to evaluate students’ learning on the chapter. Students can download Self-reflection by scanning the given QR code.
Brief summary at the end of each chapter in the form of a concept map. Matter and Atomic Structure
Manufactured Substances in Industry
CHAPTER 2
1. Define isotopes.
3. Atoms of oxygen-16, oxygen-17 and oxygen-18 are isotopes. Compare and contrast these three isotopes. 4. Magnesium exists naturally as three isotopes, which are 79.0% of 24Mg, 10.0% of 25Mg and 11.0% of 26Mg. Calculate the relative atomic mass of magnesium.
Element
Proton number
Nucleon number
W
6
12
X
6
13
Y
11
23
Z
12
24
Achievement
Molecule
exists as
Isotopes
Atom
Gas
differ from the aspects of
Attraction forces between particles
Neutron
Nucleon number
Proton
Proton number
http://bit.ly/ 2N7Sa5R
Electron
4.0
High-carbon steel
0.8
A Z
Enrichmen Corner 1. Silicon carbide, SiC is a hard and strong substance that melts at 2700 °C. Silicon carbide, SiC is suitable to be used as an abrasive. Explain why this substance is hard and has a high melting point. Carbon, C
• Low density and easily moulded • Absorb UV rays and is very transparent • High impact resistance
Check Answers
You need a pair of new spectacles. Will you choose lenses made from lead crystal glass or polycarbonate? Explain your answer.
39
Scan the QR code to take the interactive quiz at the end of each chapter.
(c) How do scientists create a very cold condition to investigate the superconductivity phenomena?
(c) Nowadays, spectacle lenses are made from polycarbonate polymer. The properties of polycarbonate are as follows:
Electron arrangement
Quick Quiz
Silicon, Si
273
Scan the QR code to get the complete answer on that chapter.
Questions to test students’ understanding at the end of each chapter. HOTS questions at the application, analysis, evaluation and creation level are marked with the HOTS icon .
Enrichment Corner
Enrichment exercises with HOTS questions at the evaluation and creation level.
Stations Experiment
Discussion Problem solving
Activity
Role-play
Project Technology usage
Figure 1
Check Answers
Achievement Test
Laboratory activities
https://bit.ly/ 2VV6Qct
274
Activities in this book:
vi
Temperature
Figure 1
2. Lead crystal glass can be used to make spectacle lenses. (a) What is the composition of lead crystal glass? (b) Explain the advantages and disadvantages of using lead crystal glass to make spectacle lenses.
X
Superconductor X
4K
(a) Give an example of a ceramic that shows superconductivity properties. (b) Explain the difference in the electrical conductivity properties of the two conductors in Figure 1.
(b) Stainless steel is produced from a mixture of chromium, nickel and carbon. (i) State the percentage of chromium, nickel and carbon in stainless steel. (ii) Stainless steel is suitable to be used to make high quality knife blades. Explain.
Uses
represented as
consists of
Carbon %
Cast iron
Non-metal superconductors
0K
(a) Cast iron is brittle whereas high-carbon steel is hard and strong. Based on Table 1, calculate the percentage of carbon that must be removed from cast iron to produce high-carbon steel.
can form
Quick
8
Steel
phases
Kinetic energy
5. Metals can conduct electricity. Ceramic materials can also be processed to conduct electricity and be made superconductors. Figure 1 shows the change in the electrical resistance value against temperature of two conductors:
Table 1 Ion
Arrangement of particles
4. The various unique properties of ceramics are modified in its use in various fields. State the property of ceramic involved in the manufacture of the following objects: (a) Car engine (b) Spark plug
1. The addition of coke (carbon) in the extraction process of iron is to remove oxygen from iron ore. The iron and carbon mixture will form steel. Table 1 shows two types of steel with different percentage of carbon.
Concept
Liquid
Industrial Chemistry
3. Traditional ceramics are made from clay such as kaolin. (a) Name two oxide compounds found in kaolin. (b) Give the formula of the ion that produces the brown colour in clay. (c) State two uses of traditional ceramics.
What new knowledge have you learned in Manufactured Substances in Industry? Which is the most interesting subtopic in Manufactured Substances in Industry? Why? Give several examples of application of Manufactured Substances in Industry in daily life. Rate your performance in Manufactured Substances in Industry on a scale of 1 to 10; 1 being the lowest and 10 the highest. Why would you rate yourself at that level? https://bit.ly/ 5. What can you do to improve your mastery in Manufactured 2RoRBJi Substances in Industry?
5. Madam Maimunah was diagnosed with bone cancer. (a) What isotope is used to treat Madam Maimunah? (b) Explain the positive and negative effects of using the isotope in (a).
Solid
THEME 4
1. 2. 3. 4.
Table 2.6
2. Based on Table 2.6, which atoms are isotopes? Explain your answer.
Matter
CHAPTER 8
Self Reflection
2.4
Resistance
TestYourself
Simulation
Icons in this book:
Career Ki o s k
Chemistry & Us GivesBrain information about Teaser patriotic elements, cultural or Malaysians’ achievements.
Relates chemical concepts with our daily life.
Brain Teaser
Chemistry Gives extra information regarding the topic.
Questions that hone students’ thinking skills.
Safety Precaution Steps that need to be taken to ensure the results of the experiment or lab activity are accurate.
A reminder to students about the dangers while carrying out the experiment or laboratory activity.
Gives information about careers related to the chemistry field.
Literacy Tips Learning or problem solving tips.
e-Portal Provides the link to the website or the QR code to get additional information.
Cross-curriculum element Gives cross-curriculum information related to the topic.
Guide to Scan AR (Augmented Reality) for Interactive 3-Dimensional Animation In this book, there are several pages that contain AR (Augmented Reality) to explain the contents with interactive three dimensional animation. The pages involved are 25, 29, 113, 115 and 259. The following guides on how to scan the AR:
1
Scan
2
or
to download the app.
Scan the whole page that contains
vii
Laboratory Rules and Safety Measures Before entering the laboratory
• Do not enter the laboratory without the teacher’s permission. • Do not bring bags or any food and drinks into the laboratory.
While in the laboratory
• • • • • • • • •
Do not run or play in the laboratory. Do not eat or drink in the laboratory. Do not taste or smell any chemical substances. Avoid touching chemical substances with your hand. Do not carry out any experiments without the teacher’s permission. Understand all instructions before starting the experiment. Read all labels and safety symbols on reagent bottles before using them. Wear safety equipment such as safety goggles, gloves and laboratory coats while carrying out experiments. Do not direct the opening of the boiling tube to one’s self or to others while heating chemical substances with a Bunsen burner. • Keep all chemical substances away from the fire. Tie your hair and tidy your clothes so that it does not come into contact with fire. • Use the required amount of chemical substances only.
If there is an emergency
• Know the location of safety equipment such as a fire extinguisher and fire hose and ways to use them. • If a chemical substance comes into contact with your eyes, skin or clothing, wash immediately with plenty of water. • If you accidentally swallowed a chemical substance, remove it from your mouth and gargle with plenty of water. • Report to the teacher as soon as possible to get treatment. • Report any accident or injury to the teacher as soon as possible.
Before exiting the laboratory
• • • •
Switch off all water, gas and electrical supplies. Return all used equipments to their original places. Clean and tidy up all used equipments. Throw all used chemical substances as instructed by the teacher. Do not throw chemical substances into the sink. • Wash your hands before leaving the laboratory.
The List of Chemical Symbols Explosive
Flammable
Oxidant
Corrosive
Toxic
Health hazard
Irritant
Gas under pressure
Environmental hazard
(Source: Department of Occupational Safety and Health (DOSH), Malaysia)
viii
EM TH E
1
The Importance of Chemistry
This theme introduces students to the meaning of chemistry, importance of chemistry and careers in the chemical field as well as chemicals in daily life. Besides, the application of scientific skills and problem-solving methods in chemistry are also strengthened from chemical perspective.
EM TH E
Fundamentals of Chemistry
2
This theme aims to introduce chemistry from the microscopic aspect which includes particles, the mole concept, chemical formulae and equations. The Periodic Table of Elements and chemical bonds are also emphasised for a better understanding of the fundamentals of chemistry.
EM TH E
3
Interaction between Matter This theme aims to introduce acids, bases and salts as well as the rate of reaction.
EM TH E
4
Industrial Chemistry This theme introduces properties of materials that are widely used in the development and growth of current technology. 1
CHAPTER
1
Introduction to Chemistry
Keywords
Chemistry Chemical technology Scientific method Personal protective equipment Safety in the laboratory Management of laboratory accidents
What will you learn? 1.1 Development in Chemistry Field and Its Importance in Daily Life 1.2 Scientific Investigation in Chemistry 1.3 Usage, Management and Handling of Apparatus and Materials
2
Bulletin Chemistry helps us to understand matter and the reactions that they go through. Lately, the field of chemistry, especially nanochemistry has been developing rapidly. Nanochemistry focuses on the learning and knowledge of synthesis and properties of particles in nanoscale (as tiny as 10–9 m). In this field, chemists study the properties and uniqueness of the atoms and molecules in nanoscale. Nanochemistry combines nanotechnology, biotechnology, chemistry, biology, physic and mathematic into one single field. The discovery of nanochemistry has benefited humans tremendously, including in the field of medicine, health, agriculture, electronics, sources and energy, manufacturing industry and others. Knowledge in chemistry is the basis for mastering nanochemistry. All applications in the world of nanochemistry begin with the knowledge of chemistry learned at school. Therefore, mastering the fundamental concepts of chemistry in secondary school is the beginning of the expansion of knowledge related to chemistry.
What is the meaning of chemistry? What are the careers that require the knowledge of chemistry? What are the correct methods to store chemicals in the laboratory?
3
THEME 1
The Importance of Chemistry
1.1
Development in Chemistry Field and Its Importance in Daily Life
Based on Figure 1.1, what do you understand about chemistry?
Wow, the fireworks look beautiful! How does it produce so many attractive colours?
Fireworks look attractive especially at night. The attractive colours are caused by the mixture of chemicals in the fireworks.
What is chemistry?
g Learnin tandard S At the end of the lesson, pupils are able to: 1.1.1 State the meaning of chemistry 1.1.2 State examples of chemicals commonly used in daily life 1.1.3 Generate ideas on the development of chemistry field and the contributions of chemical technology towards mankind 1.1.4 State examples of careers related to chemistry field
Figure 1.1 Fireworks and chemistry
Activity 1.1
Century Discussing the meaning of chemistry based on students’ understanding 21st Skills 1. Carry out the Round Table activity. 2. Take turns to state the meaning of chemistry according to your understanding on a piece of blank paper based on: (a) The science knowledge that you have learned from Form 1 to Form 3 (b) Your experience outside the classroom 3. Present your group findings to your classmates.
Meaning of Chemistry Chemistry is a field of science that studies the structures, properties, compositions and interactions between matters. Learning of chemistry is not limited to chemicals found in the laboratory but also substances commonly found in daily life such as salt and soap. Chemistry helps us to understand matter around us. The word chemistry originated from the Arabic word ‘al-kimiya’. Carry out Activity 1.2 to study the origin of the word chemistry and its meaning.
4
Father of Arabic chemistry http://bit. ly/2qruQIF
Introduction to Chemistry
CHAPTER 1
Activity 1.2 Century Collecting and interpreting the origin of the word chemistry and 21st Skills its meaning 1. Carry out the Think-Pair-Share activity. 2. Surf the Internet or refer to reading materials from various printed media to gather information and discuss the following issues: (a) Origin of the word chemistry (b) Meaning of chemistry 3. Present your group discussion results in multimedia form in front of your class.
Chemicals in Daily Life All substances around us are made up of chemicals. The activities that we carry out daily involve chemical reactions as well. Figure 1.2 shows the chemicals commonly used in foods, medicine, agriculture and industries.
Food • • • • •
Preservative Colouring Flavouring Antioxidant Stabiliser
Ch
als in daily l emic
Drugs contain chemicals http://bit. ly/2Mx4M7B
Medicine
ife
• • • • •
Antibiotic Antiseptic Vitamin Chemotheraphy Analgesic
Herbicide
Agriculture • • • • •
Industry
Herbicide Pesticide Fungicide Fertiliser Hormone
• • • • • • •
Paint Polymer Glass Ceramic Detergent Colouring Alloy
Figure 1.2 Commonly used chemicals
5
THEME 1
The Importance of Chemistry
The Development in Chemistry Field and the Contributions of Chemical Technology Researches in various chemical fields are constantly being carried out covering various disciplines. For examples, biochemistry, botany and forensics which require chemical knowledge to solve problems. The need for chemical technology to solve problems spurs the development of chemical technology. Technologies used in the 60s and 70s might not be suitable to be applied in this era. Based on your knowledge, what are the contributions of chemical technology to mankind? Carry out Activity 1.3.
Activity 1.3 Searching for information and making a poster Century 21st Skills 1. Carry out the Gallery Walk activity. 2. Gather information on the following aspects: (a) Contribution of chemists (b) Development of chemical technology (c) Careers in the field of chemistry (d) Chemicals in daily life 3. Prepare the results of your group in an attractive poster. 4. Display your group’s results in the class. 5. Each group has to move around to look at the other groups’ posters. Write your comments about the results of the other groups on sticky notes and paste them on the posters.
Bread with nanocapsules
Food preservation with salt or sugar
Food preservation by freezing
Food preservation using nanotechnology
Food preservation using nanotechnology http://bit.ly/2OKV8jj
Figure 1.3 Development of technology in food preservation
6
Introduction to Chemistry
CHAPTER 1
Careers Related to Chemistry In the era of rapid industrial development, most careers require knowledge in chemistry. For example, careers in the cosmetics, pharmaceutical, biotechnology, nanotechnology and green technology fields are shown in Figure 1.4.
i
Co sme tics
l
y log hno Nanotec Examples of careers: • Nanotechnology engineer • Food scientist
Examples of careers: • Doctor • Pharmacist
ca
e ot
Pha rm ac eu t
gy olo n ch
Bi
Examples of careers: • Biotechnology researcher • Biomedical engineer
FIELDS
Green technology
Examples of careers: • Cosmetic chemist • Cosmetic consultant
Examples of careers: • Green technology chemist • Engineer
Figure 1.4 Several careers related to chemistry
Activity 1.4
Century CT Role-playing activity on careers in the field of chemistry 21st Skills 1. Carry out the Role-Play activity. 2. Gather information from suitable reading resources or websites on careers in the field of chemistry. 3. Assign the members of the groups to various careers. 4. Prepare the acting script and suitable props. 5. Present your group’s act in front of your class.
7
THEME 1
The Importance of Chemistry
TestYourself
1.1
1. What is meant by chemistry? 2. List out five types of chemicals used in daily life. 3. Give one example of development of chemistry in industries. 4. List out at least three careers related to chemistry in the following situations: (a) Searching for an antidote for dengue fever (b) Producing palm trees with a high content of oil
1.2
Scientific Investigation in Chemistry
A scientific investigation is a scientific method used in solving problems in science. Generally, a scientific investigation begins with the observation of a problem. Look at Figure 1.5 and identify the problem that occurs. We can carry out an investigation to solve the problem by using a scientific method.
Mom, I have been stirring this salt for a long time. Why is there still salt in the solution?
Why isn’t the salt dissolving? Boys, try adding some hot water into the salt solution.
Figure 1.5 The solubility of salt in hot water Try to recall the steps in a scientific method that you have learned in Form 1.
8
g Learnin tandard S At the end of the lesson, pupils are able to: 1.2.1 Design an experiment to test a hypothesis 1.2.2 Investigate through experiment the effect of temperature on the solubility of salt in water using a scientific method
Introduction to Chemistry
CHAPTER 1
Scientific Method
Scientific method is a systematic method used by scientists or researchers to solve problems related to science. This method involves several general steps to solve a problem using the correct methods. Making observations
Using senses of sight, hearing, touch, taste or smell to gather information about an object or a phenomenon taking place.
Making an inference
Using data collection and past experience to make a conclusion and explanation about an event.
Identifying the problem
Asking questions based on the inference made.
Making a hypothesis
Making a general statement about the relationship between a manipulated variable and a responding variable to explain an event or observation. This statement can be tested to prove its validity.
Identifying the variables
Identifying the manipulated variable, responding variable and fixed variable(s) in an experiment to test the hypothesis that is formed.
Controlling the variables
In an investigation, a variable is manipulated to observe its relationship with the responding variable. Simultaneously, other variables are fixed.
Planning an experiment
Determining the materials and apparatus to be used, procedure of the experiment, method of collecting data and the ways to analyse and interpret the data.
Collecting data
Making observations or measurements and recording the data systematically.
Interpreting data
Organising and interpreting the data collected. Data can be interpreted through calculations, graphs or charts to find and determine the relationship between the variables.
Making a conclusion
Making a statement on the results of the experiment on whether the hypothesis made is accepted or rejected.
Preparing a report
Communicating in detail on all aspects of the experiment so that the outcome of the experiment can be shared for the development of chemistry knowledge. Figure 1.6 Steps in a scientific method
9
The Importance of Chemistry
Based on the scientific method that you have learned, carry out Experiment 1.1 to study the effect of temperature on the solubility of salt in water.
1.1
Experiment
Aim: To study the effect of temperature on the solubility of salt in water.
Problem statement: Does the temperature of water affect the solubility of salt in water?
Hypothesis: The increase in temperature of water will increase the solubility of salt in water.
Variables: (a) Manipulated : Temperature of water (b) Responding : Solubility of salt in water (c) Fixed : Volume of water, mass of salt, time
Brain Teaser
Materials: Distilled water and salt
Apparatus: 150 cm3 beaker, 100 cm3 measuring cylinder, thermometer, electronic scale, glass rod, Bunsen burner, stopwatch, wire gauze and tripod stand
Results:
10
30
20 10
80
Distilled water
Bunsen burner
Observation
Interpreting data: At which temperature does all the salt dissolve in water? Conclusion: Is the hypothesis acceptable? What is the conclusion of this experiment?
Prepare a complete report after carrying out this experiment.
10
Thermometer
30
40
50
60
70
80
90
100
110
How can you obtain distilled water at 10 °C?
Salt
Table 1.1
Temperature (oC)
Brain Teaser
0
Procedure: 1. Measure 50 cm3 of distilled water at temperature 10 °C with a measuring cylinder and pour it into a beaker. 2. Add 40 g of salt into the beaker and stir the solution with a glass rod for 2 minutes. 3. Observe the solubility of salt in the beaker. 4. Repeat steps 1 to 3 with distilled water heated at Glass rod 30 °C and 80 °C. 5. Record your observations in Table 1.1.
–10
THEME 1
Figure 1.7
Introduction to Chemistry
TestYourself
CHAPTER 1
1.2
1. What is meant by scientific method? 2. Why is a scientific method important in chemistry? State your opinion. 3. You are given a bucket of ice cubes, a little sugar, a measuring cylinder and a stopwatch. Plan an experiment to determine whether sugar increases the melting rate of the ice.
1.3
Usage, Management and Handling of Apparatus and Materials
The laboratory is an important place for learning chemistry, and it is a dangerous place too. All the rules and safety measures in the laboratory should be obeyed. Do you still remember the rules and safety measures you have learned in Form 1? Figure 1.8 shows a poster of general rules in the laboratory.
g Learnin tandard S At the end of the lesson, pupils are able to: 1.3.1 Explain the types and functions of self protective equipments and safety in the laboratory 1.3.2 Demonstrate methods of managing and handling apparatus and materials 1.3.3 Communicate about emergency management procedure in laboratory
Chemistry CSDS (Chemical Safety Data Sheet) is a data sheet that gives information about chemicals to assist consumers in managing the risks of using dangerous chemicals. This data sheet contains complete information on chemical poisoning besides providing information on management and storage methods, and handling methods during an emergency.
Figure 1.8 A poster of general rules in the laboratory (Source: National Institute of Occupational Safety and Health (NIOSH), Malaysia)
11
THEME 1
The Importance of Chemistry
Personal Protective Equipment Wearing personal protective equipment while carrying out experiments in the laboratory is necessary to protect yourself from accidents. Figure 1.9 shows personal protective equipment in the laboratory and their functions. Other than the personal protective equipment, the chemistry laboratory is also equipped with various safety equipment. Figure 1.10 shows the safety equipment in the laboratory and their functions. Safety goggles Used for preventing dust or splashes of chemicals from getting into the eyes.
Face mask Used for protecting the respiratory organs from chemicals in the form of powder or fumes.
Gloves Used for handling chemicals to protect hands from injuries, chemicals or infections.
Laboratory coat Used as protection for the body and clothing against chemical spills such as acids, alkalis and organic solvents.
Laboratory shoes Used for protecting the feet from injuries caused by chemical spills, sharp objects or toxic substances.
Figure 1.9 Personal protective equipment in the laboratory and their functions
12
Introduction to Chemistry
CHAPTER 1
Safety shower Is used to wash and clean the body when a chemical accident occurs on parts of the body. This equipment is also used to extinguish fire on clothings.
Fume chamber
Eyewash
A specially designed equipment to carry out experiments that release toxic vapours, cause combustions or produce pungent smells.
Used for washing and cleaning the eye when accidents occur on parts of the eye.
Fire extinguisher Used for extinguishing fire in the laboratory.
Hand wash Used for removing chemical substances, oil, dirt and microorganisms from the hands.
Figure 1.10 Safety equipment in the laboratory and their functions
Chemistry Examples of substances that release toxic vapour, cause combustion or produce pungent smell: • Concentrated sulphuric acid • Chlorine gas • Ammonia gas • Bromine gas • Alcohol • Benzene
13
THEME 1
The Importance of Chemistry
Methods of Managing and Handling Apparatus and Chemicals Skills in using and handling apparatus and materials are important and very useful in carrying out a scientific investigation. Carry out Activity 1.5 to find out the correct methods of managing and handling apparatus.
Activity 1.5 Using and handling apparatus correctly 1. Carry out this activity in groups. 2. Each group should take turns in carrying out the four activities according to the stations by using and handling apparatus and chemicals with the correct methods. Visit the website or scan the QR code for each station. (a) Station 1 - Weighing and heating of solid (c) Station 3 - Electrolysis, gas collection and gas test Station 1
Station 3
http://bit.ly/ 2MzE6Do
http://bit.ly/ 2MzEMbU
(b) Station 2 - Filtration and distillation
(d) Station 4 - Titration
Station 2
Station 4
http://bit.ly/ 2pI0scC
http://bit.ly/ 2W3RLpl
Storage and Disposal of Chemicals
What do you understand about the storage and disposal of chemicals based on the conversation in Figure 1.11? Most chemicals in the laboratory are hazardous. Chemicals should be stored properly so that they do not endanger the user nor cause accidents. Therefore, chemicals should be stored in their designated places according to their categories. Disposal of chemicals is equally important as the storage of chemicals. Disposal of chemicals without following the correct disposal procedures not only causes environmental pollution but also destroys the habitats of flora and fauna and endangers human health as well. Hence, the storage and disposal of chemicals should be taken seriously by all parties.
How can the chemicals in the laboratory be stored or disposed of?
Different types of chemicals should be stored and disposed of using different methods.
Figure 1.11 Storage and disposal of chemicals
14
Introduction to Chemistry
CHAPTER 1
Storage of Chemicals Chemicals have specific storage methods according to the types of substances. Look at the following examples: Reactive substances Reactive metals such as lithium, sodium and potassium are stored in paraffin oil to prevent reaction with the moisture in the air. Hydrocarbons and organic solvents Volatile and inflammable liquids like hydrocarbons and organic solvents should be stored in shady areas far from sunlight and heat source.
Photograph 1.1 Sodium in paraffin oil
Substances that decompose easily Substances that decompose easily in the presence of light, for example concentrated nitric acid, hydrogen peroxide solution, silver nitrate solution, liquid bromine and liquid chlorine are stored in dark bottles. Substances with pH9
Photograph 1.2 Hydrogen peroxide stored in a dark bottle
Corrosive chemicals (pH9) are usually stored in special storage cabinets that are kept locked. Heavy metals and toxic substances Toxic substances and heavy metals should be kept in special labelled containers and kept in a locked room which is heat free.
Guidelines on storage of hazardous substances
Photograph 1.3 Cabinet for keeping corrosive chemicals
http://bit.ly/2MfxBoT
15
Brain Teaser THEME 1
The Importance of Chemistry
Disposal of Chemicals Laboratory wastes have specific disposal methods according to the types of substances.
Brain Teaser Can all the chemical wastes in the laboratory be disposed of into the sink or dustbin? Justify.
Hydrogen peroxide Hydrogen peroxide wastes with a low concentration can be poured directly into the laboratory’s sink. However, hydrogen peroxide with a high concentration has to be diluted with water and added with sodium sulphite for the decomposition process to take place before being poured into the sink. Solid wastes Solid wastes like glass and rubber have to be disposed into special containers. Organic solvents and hydrocarbons Most organic solvents and hydrocarbons are toxic, carcinogenic, volatile and inflammable. This type of wastes cannot be disposed directly into the sink or the laboratory drain because it would pollute the water source and the environment. Organic solvent and hydrocarbon wastes should be kept in special containers made of glass or plastic. Substances with pH9 Substances with a pH value of pH9 are strong acids and strong alkalis respectively. Strong acid and alkali wastes can cause damage to the sink and react with water to release high heat and toxic gases. Strong acid and alkali wastes should be kept in closed labelled containers during disposal. Heavy metals and toxic substances Solutions containing heavy metals and toxic substances have to be kept in plastic bags and the solutions be left to evaporate in the fume chamber. Then, the bag of heavy metal residue is tied carefully and is put into the container of heavy metal waste. This type of substances should be discarded and disposed according to standard procedures.
Photograph 1.4 A closed container for disposing hazardous wastes
Chemistry According to Malaysian regulations, chemicals that are classified as listed wastes should be disposed according to standard procedures. The guidelines for listed wastes can be obtained from the Department of Environment of Malaysia.
Volatile substances Substances such as alcohol, ammonia and bromine are volatile, that is easily converted to gas at room temperature. Some of the gases produced from volatile substances are hazardous to humans and can be fatal if inhaled in large amounts. Volatile wastes should be stored in closed containers and kept away from sun and heat.
16
Introduction to Chemistry
Chemical wastes and apparatus contaminated by chemicals should be disposed of into bins or bottles labelled with the types of wastes. Then, the waste will be sent to disposal centres.
CHAPTER 1
Photograph 1.5 Bin and bottle for chemical waste disposal
Emergency Management Procedure in the Laboratory
Waste spills continue to occur in laboratories even with safety measures in place. When these accidents happen, you should act according to the correct procedure as shown below: 1 Inform your teacher or the laboratory assistant about the accident immediately. 2 Prohibit other students from entering the accident site. 3 Stop the spill from spreading to other areas by using sand to border it. 4 Clean the chemical spill. 5 Dispose of the chemical spill by following the correct procedures. Mercury thermometers are often used to carry out experiments in the school laboratory. If the mercury thermometer breaks, a pupil faces the risk of mercury spill. Although the quantity of mercury in the thermometer is very minimal, it is enough to cause mercury poisoning. Mercury poisoning occurs when a person is exposed to mercury in certain amount.
Mercury is volatile at room temperature. Mercury vapour poisons the air. Do not touch mercury spill because it can absorb into your body through the skin.
News on mercury spill http://bit.ly/ 2MxQZh9
Symptoms of mercury poisoning: • Nausea • Coughing • Vomiting • Diarrhoea • Chest pain • Sore throat • Difficulty in breathing • Headache • Eye irritation • Vision problem • Increase in blood pressure
Figure 1.12 Mercury spill
17
THEME 1
The Importance of Chemistry
Steps to be taken the moment mercury spill occurs. 1 Inform your teacher or the laboratory assistant about the accident. 2 Make the spill site as the prohibited area. 3 Sprinkle sulphur powder to cover up the spill. 4 Contact the Fire and Rescue Department for further action.
If you are exposed to mercury poisoning, you should: Stay away from the mercury source to prevent further exposure to it Go to the hospital for treatment
Guidelines on handling mercury spill http://bit.ly/ 2MJRqng
Activity 1.6
Century Discussing the emergency management procedures in the laboratory 21st Skills 1. Carry out the Gallery Walk activity. 2. Surf the Internet to gather information on the types of accidents that commonly occur in the laboratory and emergency management procedures. 3. Prepare a presentation of your information. 4. Display the information in the class. 5. Each group should move to observe the information of the other groups. Write comments on the other groups’ information on sticky notes and paste them.
TestYourself
1.3
1. List out three safety steps while in the laboratory. 2.
Give the functions of the following equipment: (a) Fume chamber (b) Safety shower (c) Laboratory coat
3. How would you manage solid wastes such as glass and rubber in the laboratory? 4. Explain how you would test the presence of oxygen and hydrogen gases. 5. Explain how you can get the most accurate reading in titration. 18
http://bit.ly/ 33V9v9c
Quick
• Making observations • Making an inference • Identifying the problem • Making a hypothesis • Identifying the variables • Controlling the variables • Planning an experiment • Collecting data • Interpreting data • Making a conclusion • Writing a report
steps
meaning
Systematic method to solve scientific problems
in the field of
Industries
Biotechnology Nanotechnology
Agriculture
Pharmaceutical
Foods Medicine
used in
Daily life
importance
Chemistry
Cosmetics
method of study
meaning
• Distillation • Titration • Filtration • Gas collection • Gas test • Electrolysis • Heating of solids • Weighing of solids
methods
Apparatus
types
Materials
• Solid waste • Substances with pH9 • Organic solvents • Hydrocarbons • Hydrogen peroxide • Toxic substances • Heavy metals • Volatile substances • Reactive substances
Mercury spill
Chemical spill
Safety equipment
Personal protective equipment
management and handling of
accidents
safety in the laboratory
Study of: • Structure of matter • Properties of matter • Composition of matter • Interaction between matter
Green technology
Careers
Scientific method
Concept
Introduction to Chemistry CHAPTER 1
19
THEME 1
The Importance of Chemistry
Self Reflection 1. What new knowledge have you learned in Introduction to Chemistry? 2. Which is the most interesting subtopic in Introduction to Chemistry? Why? 3. Why is the learning of Introduction to Chemistry important in the next chemistry lesson? 4. Rate your performance in Introduction to Chemistry on a scale of 1 to 10; 1 being the lowest and 10 the highest. Why would you rate yourself at that level? 5. What else would you like to know about Introduction to Chemistry?
Achievement
http://bit.ly/ 2MEAuyw
1
1. Chemicals are substances that cannot be ignored in our daily lives. (a) State five types of chemicals that are commonly used in daily life. (b) For each chemical stated in (a), state its use. 2. (a) Name three industries in Malaysia that use knowledge of chemistry. (b) How do the industries stated in (a) benefit our country? 3. List out three types of personal protective equipment and state the function of each equipment. 4. Complete the following flow chart on the steps involved in a scientific method. Making observation
Making an inference
Identifying the problem
Writing a report
5. S tate the safety measures that should be taken to overcome mercury spill accidents that occur in the school laboratory. 20
Introduction to Chemistry
CHAPTER 1
6. Three pieces of cloth with a size of 10 cm × 10 cm each are sprinkled with 20 cm3 of distilled water. Then, all the three pieces of cloth are folded in different styles and left to dry at room temperature. The time taken for each piece to dry is recorded. (a) Explain why the three pieces of cloth are folded in different styles.
(b) Suggest a hypothesis for this experiment.
(c) Make an inference for this experiment.
(d) Determine the variables involved in this experiment.
(e) Construct a suitable table to record the readings in this experiment.
Enrichmen Corner 1. Figure 1 shows several examples of waste substances in the school laboratory.
Broken conical flask
Concentrated acid
Ethanol
Used gloves
Used litmus paper
Dilute NaOH
Bromine
Figure 1
(a) Based on Figure 1, list out wastes that cannot be disposed of into the school sink or the rubbish bin. (b) Explain how to manage the waste substances listed in (a) correctly. 2. H alim is a farmer. He plants various types of vegetables to supply the local market. However, lately, Halim realised that the produce was unsatisfactory and the growth of vegetables was stunted, or the vegetables had died off. He suspected the soil to be acidic. As a scientist, you are asked to help Halim to determine the most suitable pH value of the soil for planting vegetables. Suggest a suitable hypothesis and state briefly how you can control the variables in this experiment. STEM
Check Answers
http://bit.ly/ 3442zGX
21
CHAPTER
2
Matter and the Atomic Structure
Keywords
Atom Ion Molecule Proton number Nucleon number Electron arrangement Isotopes Natural abundance
What will you learn? 2.1 Basic Concepts of Matter 2.2 The Development of the Atomic Model 2.3 Atomic Structure 2.4 Isotopes and Its Uses
22
Bulletin Have you heard of patients undergoing PET-CT scanning? PET-CT is the abbreviation for Positron Emission Tomography-Computed Tomography. PET-CT scanning can give accurate information on the position of a disease in the patient’s body, especially to detect and treat diseases such as cancer, inflammation and infection. PET-CT is an imaging technique that combines CT scan and PET scan. CT is able to visualise the image of a tissue or organ through the anatomical cross section of organs. PET is able to show the metabolism level of cells and tissues in the body of a patient using radioisotopes as a tracer. Did you know that radioisotopes are isotopes that exhibit radioactivity? What are isotopes? Do isotopes have the same subatomic particles like other atoms of their elements?
Who is the scientist that proved the existence of neutrons in the nucleus? How many valence electrons are there in 126C? What is the use of isotope cobalt-60?
23
THEME 2
Fundamentals of Chemistry
2.1
Basic Concepts of Matter
g Learnin tandard S
Meaning of Matter
Try to recall what matter is based on the conversation in Figure 2.1. What are the mass and volume of a rainbow?
A rainbow does not have mass nor volume.
I agree with Ahmad. A rainbow is not a matter.
At the end of the lesson, pupils are able to: 2.1.1 Describe matter briefly 2.1.2 Explain the changes in the states of matter 2.1.3 Determine the melting point and freezing point of naphthalene through activity
Chemistry
Figure 2.1 Rainbow is not a matter
Matter is something that has mass and occupies space. Matter consists of particles that are tiny and discrete. Matter can exist in three states, namely solid, liquid and gas. What are other examples of matter that you encounter in your daily life?
Changes in the State of Matter
Plasma is the fourth state of matter besides solid, liquid and gas. A plasma is an ionised gas. Although plasma in its natural state is rarely found on Earth, plasma is the state of matter most found in the universe. Most stars exist as plasma.
The change in the state of matter is caused by heating or cooling.
Ice cream that is left at room temperature absorbs heat energy and changes from solid to liquid state.
Photograph 2.1 Ice cream
At night, glass windows release heat to the surroundings causing the surface of the glass window to become cold. Water vapour in the air that comes in contact with the cold surface loses heat and forms water droplets on the surface of the glass window. Photograph 2.2 Water droplets on the surface of a glass window
24
Matter and the Atomic Structure
Figure 2.2 shows the conversion among states of matter through the process of heat absorbed or heat released. When heat energy is absorbed or released, changes occur in kinetic energy, arrangement of particles and attraction force between particles, causing the state of matter to change.
CHAPTER 2
The difference between deposition and sublimation process http://bit.ly/2IJ8b0X
Solid
Solid
Melting
Freezing
• Particles are closely packed in an orderly manner • Kinetic energy of particles is low • Attraction force between particles is strong
Liquid Sublimation
Deposition Liquid
Boiling/ Evaporation
Condensation
• Particles are closely packed but not in an orderly manner • Kinetic energy of particles is higher than solid state • Attraction force between particles is strong, but less than solid state
Gas
Gas
Key:
Heat energy absorbed Heat energy released
• Particles are far apart • Kinetic energy of particles is very high • Attraction force between particles is weak
Figure 2.2 Conversion between states of matter
Activity 2.1 Drawing the arrangement of particles in 2D form 1. Based on Figure 2.2, draw the arrangement of particles in solid, liquid and gas in 2D form. 2. Display your work on the notice board in your class. 25
THEME 2
Fundamentals of Chemistry
Matter can exist in the form of elements or compounds. Elements consist of particles of atoms or molecules while compounds are made up of molecules or ions. Figure 2.3 shows the classification of matter. Matter Element Atom Carbon, C
Compound Molecule
Molecule
Ion
Oxygen, O2
Water, H2O
Sodium chloride, NaCl
Figure 2.3 Classification of matter
Chemistry Ions are formed by transfer of electrons between atoms.
Na
Donates 1 electron
Sodium atom
+ Na
Cl
Sodium ion formed
Chlorine atom
Melting Point and Freezing Point How do scientists determine the melting point and freezing point of a substance? Carry out Activity 2.2 to determine the melting point and freezing point of naphthalene, C10H8.
Accepts 1 electron
– Cl
Chloride ion formed
Literacy Tips • Melting point is the constant temperature when a substance changes from solid state to become liquid at a specific pressure. • Freezing point is the constant temperature when a substance changes from liquid state to become solid at a specific pressure.
Activity 2.2 Aim: To determine the melting point and freezing point of naphthalene, C10H8. Materials: Naphthalene, C10H8 and water Apparatus: Boiling tube, 250 cm3 beaker, thermometer, tripod stand, retort stand with clamp, Bunsen burner, stopwatch, conical flask, wire gauze and spatula Procedure: 1. Fill up one third of a boiling tube with naphthalene, C10H8. 2. Place a thermometer in the boiling tube. 3. Pour water into a beaker until it is half filled. 4. Immerse the boiling tube into the beaker as shown in Figure 2.4. Ensure the level of naphthalene, C10H8 in the boiling tube is below the level of water in the beaker. 26
Avoid touching the naphthalene, C10H8 or inhaling the naphthalene gas.
CHAPTER 2
Thermometer
110
Boiling tube
Thermometer
70
80
Water
60
–10
0
10
90
20
100
30
40
50
60
70
80
90
100
110
Matter and the Atomic Structure
Boiling tube
–10
0
10
20
30
40
50
Naphthalene, C10H8
Conical flask
Naphthalene, C10H8
Figure 2.4 Heating of naphthalene,
10
H8
Figure 2.5 Cooling of naphthalene,
10
H8
5. Heat water and stir the naphthalene, C10H8 slowly using the thermometer. When the temperature of the napthalene, C10H8 reaches 60 °C, start the stopwatch. 6. Record the temperature and state of matter of naphthalene, C10H8 at 30 seconds intervals until the temperature reaches 90 °C. 7. Remove the boiling tube from the water bath. Dry the outer surface of the boiling tube and put it into a conical flask as Deficiency of glucose-6phosphate dehydrogenase shown in Figure 2.5. (G6PD) is a genetic disease. 8. Stir the naphthalene, C10H8 continuously. Exposure to naphthalene, 9. Record the temperature and state of matter of naphthalene, C10H8 C10H8 to a patient with G6PD will cause haemolysis, that at 30 seconds intervals until the temperature decreases to 60 °C. is destruction of red blood 10. Record your observations. cells. This situation will
cause the patient to feel tired
Interpreting data: and dizzy. 1. Plot a graph of temperature against time for the following: (a) Heating of naphthalene, C10H8 (b) Cooling of naphthalene, C10H8 2. On the graphs, label the states of matter of naphthalene, C10H8 whether solid, liquid or both. 3. Determine the melting point and freezing point of naphthalene, C10H8 from the graphs plotted. Discussion: 1. During the heating of naphthalene, C10H8: (a) Why is naphthalene, C10H8 not heated directly using the Bunsen burner? (b) Why is a water bath used? 2. During the cooling of naphthalene, C10H8: (a) Why is the boiling tube put into a conical flask? (b) Why is naphthalene, C10H8 stirred continuously? (c) Predict what would happen if naphthalene, C10H8 is not stirred continuously. 3. Explain why the temperature becomes constant when melting and freezing of naphthalene, C10H8 take place. Prepare a complete report after carrying out this activity. 27
THEME 2
Fundamentals of Chemistry
The graph of temperature against time for heating of naphthalene, C10H8 is shown in Figure 2.6 and the graph of temperature against time for cooling of naphthalene, C10H8 is shown in Figure 2.7.
• Melting occurs • No increase in temperature from B to C because heat energy absorbed by the particles is used to overcome the Temperature (°C) attraction force between the particles until solid changes to liquid
Melting point
B
Chemistry Lauric acid, C12H24O2 is a type of fatty acid that can be obtained from coconuts. This acid is also suitable to be used as a substitute for naphthalene, C10H8 in Activity 2.2.
D • Temperature increases from C to D • Exists in liquid state • When heated, the particles absorb heat and move faster
C
• Temperature increases from A to B • Exists in solid state • When heated, the particles absorb heat energy and vibrate faster because kinetic energy increases
A
Time (min)
Figure 2.6 Heating curve of naphthalene, C10H8 • Temperature decreases from E to F • Exists in liquid state • When cooled, the particles release heat and move slower due to the lost of kinetic energy Temperature (°C)
E
Freezing point
F
• Freezing occurs • No decrease in temperature from F to G because heat energy that is lost to the surroundings is balanced by the heat energy released when the particles attract each other to form solid • Temperature decreases from G to H • Exists in solid state • When cooled, the particles release heat and vibrate slower
G
H Time (min)
Figure 2.7 Cooling curve of naphthalene, C10H8
28
Matter and the Atomic Structure
TestYourself
CHAPTER 2
2.1
1. State the types of particles that exist in a copper wire. 2. Lily dries her hair with a hair dryer. (a) Name the process involved during hair drying. (b) State the changes in the movement of water particles when hair is dried. 3. Lauric acid, C12H24O2 is heated from room temperature to 50 °C. At 43 °C, lauric acid, C12H24O2 starts to melt. (a) Draw a heating curve for lauric acid, C12H24O2. (b) Why is the temperature constant at 43 °C?
2.2
g Learnin tandard S
The Development of the Atomic Model
Subatomic Particles Figure 2.8 shows the subatomic particles found in an atom which is made up of protons, neutrons and electrons. What are the similarities and differences in these three types of subatomic particles? Proton Symbol: p Relative charge = +1 Relative mass = 1
At the end of the lesson, pupils are able to: 2.2.1 State the subatomic particles in atoms of various elements 2.2.2 Compare and contrast the relative mass and relative charge of proton, electron and neutron 2.2.3 Sequence the atomic structure models based on Atomic Models of Dalton, Thomson, Brain Teaser Rutherford, Bohr and Chadwick
Neutron
Electron
Symbol: n Relative charge = 0 (Neutral) Relative mass = 1
Symbol: e Relative charge = –1 1 Relative mass = 1840
Figure 2.8 Subatomic particles
Brain Teaser How are the relative charge and relative mass of subatomic particles determined?
Activity 2.3 Comparing and contrasting the subatomic particles 1. Watch a video clip on subatomic particles by searching the Internet. 2. Based on the video, compare and contrast the relative masses and charges of protons, electrons and neutrons. 3. Present your findings using a suitable graphic presentation software and upload your work to social media. 29
THEME 2
Fundamentals of Chemistry
Development of the Atomic Structure Model
Atoms can neither be seen with the naked eye nor the microscope. Have you ever thought how the atomic structure model is produced? The atomic structure model that we know now is the product of many scientists’ efforts. Studies on atoms started since the introduction of atomic theory by Democritus, a Greek philosopher, around 500 B.C. Figure 2.9 shows the historical development of the atomic structure model.
John Dalton (1766 – 1844)
J.J. Thomson (1856 – 1940)
• Matter is made up of particles called atoms • An atom is the smallest spherical body that cannot be created, destroyed nor divided further • Same elements have the same atoms
Ernest Rutherford (1871 – 1937)
• Discovered negatively-charged particles called electrons • Atom is a positively-charged sphere with several electrons in it
• Nucleus is the centre of the atom • Discovered positively-charged particles called protons in the nucleus • Almost the whole of atomic mass is concentrated in the nucleus • Electrons move outside the nucleus
Electron
Electrons move outside the nucleus
– –
–
–
–
– –
–
–
–
Thomson’s Atomic Model
–
Nucleus consists of protons
Rutherford’s Atomic Model
Figure 2.9 Historical development of the atomic model
30
–
–
Positively-charged sphere
Dalton’s Atomic Model
– +
Matter and the Atomic Structure
CHAPTER 2
Activity 2.4
Century Role-playing on the development of the atomic structure model 21st Skills 1. Carry out the Role-Play activity in groups. 2. In your group, find information on the atomic structure model explained by one of the following scientists:
John Dalton
Ernest Rutherford
J.J. Thomson
CT
James Chadwick
Niels Bohr
3. Prepare the acting scripts and suitable props. 4. Present the group act in front of the class.
Niels Bohr (1885 – 1962) • Electrons in an atom move in shells around the nucleus Shell +
Nucleus that contains protons
James Chadwick (1891 – 1974) • Discovered neutral particles, that are neutrons in the nucleus • Neutrons contribute almost half of the mass of an atom Shell Nucleus that contains protons and neutrons
Electron Electron
Bohr’s Atomic Model
Chadwick’s Atomic Model
31
THEME 2
Fundamentals of Chemistry
TestYourself
2.2
1. Figure 2.10 shows the atomic structure of nitrogen. (a) Name X. (b) State the subatomic particles found in the nucleus of nitrogen atom. (c) Compare X and subatomic particles mentioned in (b) from the aspect of relative charge and relative mass. 2. • Electrons move around the nucleus in shells. • Nucleus of an atom consists of protons and neutrons.
X
Nucleus
Figure 2.10
The statements above show the information on an atomic structure model. (a) Which scientist identified it? (b) Draw this atomic structure model.
2.3
Atomic Structure
g Learnin tandard S
Proton Number and Nucleon Number
Look at Table 2.1, what is the relationship between the number of protons and proton number, and the relationship between the nucleon number and the proton number? Table 2.1 Proton numbers and nucleon numbers of oxygen atom, sodium atom and chlorine atom Atom
Number of protons
Number of neutrons
Proton number
Nucleon number
Oxygen
8
8
8
16
Sodium
11
12
11
23
Chlorine
17
18
17
35
The number of protons in the nucleus of an atom is known as the proton number. The total number of protons and neutrons in the nucleus of an atom is known as the nucleon number. Nucleon number = number of protons + number of neutrons or Nucleon number = proton number + number of neutrons Atoms of different elements have different proton numbers. For example, sodium atom has a proton number of 11 and chlorine atom has a proton number of 17. 32
At the end of the lesson, pupils are able to: 2.3.1 Define proton number and nucleon number 2.3.2 Determine the nucleon number, proton number and number of electrons in an atom 2.3.3 Write the standard representation of an atom 2.3.4 Construct an atomic structure diagram and electron arrangement
Brain Teaser
Brain Teaser Are there any two elements with the same proton number? Explain.
Matter and the Atomic Structure
CHAPTER 2
An atom is neutral when the number of electrons is the same with the number of protons. For example, an oxygen atom has 8 protons and also 8 electrons. The examples and solutions are shown in Example 1 and 2. Example
1
An aluminium atom has 13 protons and 14 neutrons. What are the proton number and nucleon number of an aluminium atom? Solution
Proton number = number of protons = 13 Nucleon number = proton number + number of neutrons = 13 + 14 = 27 Example
2
The nucleon number of a potassium atom is 39. A potassium atom has 19 protons. How many neutrons and electrons are there in a potassium atom? Solution
Number of electrons = number of protons = 19 Number of neutrons = nucleon number – number of protons = 39 – 19 = 20 Table 2.2 shows the comparison among the number of protons, neutrons and electrons when a chlorine atom accepts an electron to form a chloride ion. What are the changes in the number of protons, neutrons and electrons? Table 2.2 Number of subatomic particles of chlorine atom and chloride ion Type of particle
Chlorine atom, Cl
Chloride ion, Cl-
Number of proton
17
17
Number of neutron
18
18
Number of electron
17
18
A chlorine atom accepts one electron to form a chloride ion, thus a chloride ion has one electron more than a chlorine atom. The number of protons and neutrons in a chlorine atom and chloride ion are the same. Therefore, during the formation of ion from an atom, the number of protons and neutrons in the nucleus remain the same.
Literacy Tips When the number of electron increases, an anion is formed, which is a negatively-charged ion. However, when the number of electron decreases, a cation is formed, which is a positively-charged ion.
33
THEME 2
Fundamentals of Chemistry
Standard Representation of an Atom
An atom can be represented using a standard representation as shown in Figure 2.11. Nucleon number Proton number
A
ZX
Symbol of element
Figure 2.11 Standard representation of an atom
What information can you obtain from
12 6
C
? The symbol of carbon element is C,
the nucleon number of a carbon atom is 12, while the proton number of a carbon atom is 6. A sodium atom contains 12 neutrons and 11 protons in the nucleus. What is the standard representation of a sodium atom?
Activity 2.5 Finding the mystery code CT 1. In groups, answer the following questions: (a) The nucleon number and the proton number of fluorine element are 19 and 9 respectively. Is the following statement true or false? The atom of this element has 9 electrons and 9 neutrons in its nucleus. (b) Atom X has 11 protons and 12 neutrons. Find the nucleon number of this atom. (c) What is the proton number of an atom of nitrogen element that has 7 electrons? (d) The standard representation of an atom of oxygen element is 168O. This atom accepts electrons to form an oxide ion, O2–. How many electrons are accepted by an oxygen atom to form an oxide ion, O2–? (e) The nucleus of atom Y has the charge +4 and contains 5 neutrons. State the nucleon number of element Y. (f) A calcium atom has 20 protons and its nucleon number is 40. A calcium ion, Ca2+ is formed when a calcium atom donates 2 electrons. State the number of neutrons in a calcium ion, Ca2+. A n atom of element W has 3 electrons and 4 neutrons. What should the number in (g) 7 W the box be, to represent the atom of element W? 2. Scan QR Code or visit the website provided to obtain the code guidance. 3. Get the mystery code. 34
Code guidance http://bit.ly/2P8zNQV
Matter and the Atomic Structure
CHAPTER 2
Atomic Structure and Electron Arrangement
Electrons of an atom orbit around the nucleus in their respective shells. Electrons will fill the shell closest to the nucleus first. When the shell closest to the nucleus is full, electrons will fill the next shell. The maximum number of electrons in the first three shells for elements with proton numbers 1 to 20 are shown in Figure 2.12.
Nucleus
First shell: 2 electrons Second shell: 8 electrons Third shell: 8 electrons
Chemistry The third shell can be filled with a maximum of 18 electrons for elements with proton number exceeding 20.
Figure 2.12 Numbers of maximum electrons in the first three shells for elements with proton numbers 1 to 20
For example, the proton number of aluminium is 13. This shows that an aluminium atom has 13 electrons. The electron arrangement of aluminium atom is, 2 electrons filled in the first shell, 8 electrons filled in the second shell and 3 electrons filled in the third shell. The electron arrangement of aluminium atom can be written as follows:
2.8.3
Number of valence electrons = 3
The outermost shell filled with electrons is the valence shell. Electrons in the valence shell are known as valence electrons. The chemical properties of an element depend on the number of valence electrons of the atom. Elements with the same number of valence electrons have similar chemical properties. The electron arrangement shows the nucleus and electron arrangement of an atom. For example, the electron arrangement of aluminium atom is shown in Figure 2.13.
Al
Figure 2.13 Electron arrangement of aluminium atom
Chemistry The valence shell is the outermost shell of an atom.
The atomic structure shows the number of protons and neutrons in the nucleus and electron arrangement of an atom. For example, the atomic structure of aluminium is shown in Figure 2.14.
13 p 14 n
Figure 2.14 Atomic structure of aluminium atom
35
THEME 2
Fundamentals of Chemistry
Activity 2.6 Writing the electron arrangement and drawing the atomic structure 1. Get the standard representation of atoms of the first 20 elements in the Periodic Table of Elements from the QR code or the Standard given website. representation Based on the information: of atoms (a) Write the electron arrangement of the 20 elements http://bit.ly/2qusDfC (b) Draw the atomic structure of the 20 elements 2. Display your work on the notice board in your class.
CT
Activity 2.7 Illustrating the atomic structure using a model 1. Carry out this activity in groups. 2. Choose an element from the elements with proton numbers 1 to 20. Produce a model to illustrate the atomic structure using recycled materials. 3. The model that is produced must include the following: (a) Protons and neutrons in the nucleus (b) Electron arrangement in the shells 4. Present the model in front of the class.
TestYourself
CT
2.3
Table 2.3 shows the number of protons and the number of neutrons for elements X, Y and Z. Table 2.3 Element
Number of protons
Number of neutrons
X
10
10
Y
11
12
Z
19
20
1. What is the nucleon number of atom Y? 2. Write the standard representation of element Z. 3. Atom Y donates one electron to form ion Y+. State the number of protons, neutrons and electrons for ion Y+. 4. (a) Write the electron arrangement of atom X. (b) Draw the electron arrangement for atom X. (c) Draw the atomic structure of atom X. Label all the subatomic particles in the diagram.
36
Matter and the Atomic Structure
2.4
Isotopes and Its Uses
g Learnin tandard S
Meaning of Isotopes
Figure 2.15 shows three atoms of hydrogen element. All these three atoms of hydrogen have the same proton number but different nucleon numbers. These hydrogen atoms are known as isotopes.
CHAPTER 2
At the end of this lesson, pupils are able to: 2.4.1 Deduce the meaning of isotopes 2.4.2 Calculate the relative atomic mass of isotopes 2.4.3 Justify the usage of isotopes in various fields
1
H
1 3
H
2
H
1
1
Figure 2.15
Activity 2.8 Generalising the meaning of isotopes 1. Carry out this activity in groups. 2. Compare and contrast the number of protons, electrons and neutrons in the isotopes of silicon, magnesium and phosphorus. 28
Si
14
29
30
Si
14
Si
14
24
Mg
12
25
Mg
12
26
Mg
12
31
P
15
32
P
15
3. Interpret the information obtained and generalise the meaning of isotopes. Isotopes are atoms of the same element with the same number of protons but different number of neutrons. For example, chlorine has two isotopes, chlorine-35 and chlorine-37. Table 2.4 shows the number of subatomic particles for the isotopes of chlorine. Atoms of chlorine-35 and chlorine-37 have different masses because the number of neutrons in the nucleus are different.
Chemistry Chlorine atom with the nucleon number 35 can be represented by Cl-35, 35
Cl or 35Cl.
17
Table 2.4 Number of subatomic particles for the isotopes of chlorine Isotope
Atomic standard representation
Chlorine-35
35
Chlorine-37
37
Number of protons Number of neutrons Number of electrons
Cl
17
Cl
17
17
18
17
17
20
17
Relative Atomic Mass of Isotopes
Most elements exist naturally as a mixture of two or more isotopes. Relative atomic mass of these elements depend on the natural abundance of isotopes in a sample. Natural abundance is the percentage of isotopes present in a natural sample of element. The relative atomic mass can be calculated from the natural abundance of an element containing isotopes using the following formula: Relative atomic mass =
∑(% isotope × mass of isotope) 100 37
THEME 2
Fundamentals of Chemistry
Example
3 35
37
35
37
Chlorine consists of two isotopes, 17Cl and 17Cl . The natural abundance of 17Cl is 75% and 17Cl is 25%. Calculate the relative atomic mass of chlorine. Solution
35
35
37
37
(% isotope 17Cl × mass 17Cl ) + (% isotope 17Cl × mass 17Cl ) 100 (75 × 35) + (25 × 37) = 100 = 35.5 Relative atomic mass of chlorine =
Uses of Isotopes
Isotopes are used widely in various fields as listed in Table 2.5. Table 2.5 Uses of isotopes in various fields Field Medicine
Isotope
Uses
Cobalt-60
• In radiotherapy to kill cancer cells without surgery • Sterilising surgical tools
Iodine-131
• Treatment of thyroid disorders such as hyperthyroidism and thyroid cancer
Agriculture
Phosphorus-32 • Study of plant metabolism
Nuclear
Uranium-235
Archaeology Carbon-14
• Generating electricity through nuclear power generator • Estimation of artifacts or fossils’ age
Lead-210
• In determining the age of sand and earth layers up to 80 years
Industry
Hydrogen-3
• As a detector to study sewage and liquid wastes
Engineering
Sodium-24
• In detecting leakage in underground pipes
Development in the field of science, specifically chemistry has maximised the use of isotopes in various fields. Isotopes are used for sustainability of life. The use of isotopes causes both positive and negative effects on the environment and society.
Activity 2.9
Uses of isotopes http://bit.ly/ 32zUJUP
Century Holding a forum on the issues of using isotopes 21st Skills 1. Carry out this activity in groups. 2. Each group is given a role as a chemist, medical representative, enforcer and others. Based on the role given, search for information concerning issues involving isotopes. 3. Hold a forum to discuss the positive and negative effects of using isotopes. 4. Record the forum proceedings and upload to social media.
38
Matter and the Atomic Structure
TestYourself
CHAPTER 2
2.4
1. Define isotopes.
Table 2.6
2. Based on Table 2.6, which atoms are isotopes? Explain your answer. 3. Atoms of oxygen-16, oxygen-17 and oxygen-18 are isotopes. Compare and contrast these three isotopes. 4. Magnesium exists naturally as three isotopes, which are 79.0% of 24Mg, 10.0% of 25Mg and 11.0% of 26Mg. Calculate the relative atomic mass of magnesium.
Element
Proton number
Nucleon number
W
6
12
X
6
13
Y
11
23
Z
12
24
5. Madam Maimunah was diagnosed with bone cancer. (a) What isotope is used to treat Madam Maimunah? (b) Explain the positive and negative effects of using the isotope in (a).
Concept Ion
Matter
Molecule
exists as
phases
Solid
Liquid
can form
Isotopes
Atom
Gas
differ from the aspects of
represented as
consists of
Arrangement of particles
Quick
Kinetic energy
Attraction forces between particles
Uses
Neutron
Nucleon number
Proton
Proton number
http://bit.ly/ 2N7Sa5R
Electron
A Z
X
Electron arrangement
39
THEME 2
Fundamentals of Chemistry
Self Reflection 1. What new knowledge have you learned in Matter and Atomic Structure? 2. Which is the most interesting subtopic in Matter and Atomic Structure? Why? 3. Give a few examples on the application of Matter and Atomic Structure in daily life. 4. Rate your performance in Matter and Atomic Structure on a scale of 1 to 10; 1 being the lowest and 10 the highest. http://bit.ly/ Why would you rate yourself at that level? 2Mkz7Xa 5. What else would you like to know about Matter and Atomic Structure?
2
Achievement
1. Table 1 shows the melting point and boiling point of substances A, B, C, D and E. Table 1 Substance
Melting point (°C)
A
–101.0
Boiling point (°C) –35.0
B
–94.0
65.0
C
17.8
290.0
D
97.8
883.0
E
801.0
1413.0
(a) Classify substances A, B, C, D and E according to states of matter at room temperature. (b) State the substance that will change from liquid to solid when placed in the freezer at temperature 2 °C. (c) Describe the changes that take place on the particles of substance B with relation to energy and attraction force between particles when cooled from 80 °C to –2 °C. 2. A group of students carried out an experiment to determine the melting point of lauric acid, C12H24O2. Temperature (°C) Figure 1 shows the heating curve obtained. (a) Copy Figure 1 and label the melting point of lauric acid, C12H24O2 on the diagram. (b) Draw the arrangement of particles in lauric acid, Q R C12H24O2 between R and S. (c) The melting point of lauric acid, C12H24O2 is 43 °C. Suggest a suitable method of heating lauric P acid, C12H24O2. Time (min) (d) Draw a labelled diagram to show the set-up of Figure 1 apparatus for the method suggested in (c). 40
S
Matter and the Atomic Structure
CHAPTER 2
3. Chen Ling cleans her wound using alcohol as shown in Photograph 1. Chen Ling’s skin feels cool when wiped with alcohol. Explain this situation. Cotton swab
4. Figure 2 shows the nucleus charges of the atoms of elements X, Y and Z.
+13
+9
+3
Photograph 1
X Y
Figure 2
Z
(a) State the subatomic particle that provides the charges in the atoms of the elements. (b) State the other subatomic particle found in the nucleus of the atoms. (c) Write the electron arrangement of the atoms for elements X, Y and Z. (d) Atom Z contains 14 neutrons. Calculate the nucleon number of atom Z. 5. Figure 3 shows the information on boron. Boron has two isotopes, namely isotope 11B and isotope yB. Based on the information given, calculate the nucleon number of isotope yB.
Relative atomic mass of boron = 10.81 80% 11B 20% isotope Boron-Y
6. Figure 4 shows the standard representation of a platinum atom. A platinum ion contains 74 electrons and has a nucleon number of 195. (a) What are the number of protons and neutrons in the platinum ion? (b) What is the charge of the platinum ion?
Figure 3
195
78 Pt Figure 4
7. Justify the use of iodine-131 in the treatment for hyperthyroidism.
Enrichmen Corner 1. You lost your way while camping in a jungle. You felt thirsty but could not find a water source. In your bag, there were a transparent plastic bag and a string. Using the available things, describe how you could produce water through the condensation process. STEM
Check Answers http://bit.ly/ 2odL87q
41
CHAPTER
3
Keywords
Chemical formula Molar volume Relative atomic mass Relative formula mass Molar mass Relative molecular mass Mole Chemical equation
The Mole Concept, Chemical Formula and Equation
What will you learn? 3.1 3.2 3.3 3.4 42
Relative Atomic Mass and Relative Molecular Mass Mole Concept Chemical Formula Chemical Equation
Bulletin Satay is a favourite food among Malaysians. Satay is made from pieces of spiced meat skewered on coconut or bamboo skewers and grilled over burning charcoal. Did you know that the burning of charcoal is a form of chemical reaction? This reaction can be represented by the following chemical equation. C(s) + O2(g) → CO2(g) The symbol ‘C’ in the equation shows the chemical formula of carbon element in the charcoal. What is a chemical formula? What are the meanings of the other letters and numbers in the above equation?
How do you write the formula of a chemical substance? What information is found in a chemical equation? How do you measure the quantity of a chemical substance?
43
THEME 2
Fundamentals of Chemistry
3.1
Relative Atomic Mass and Relative Molecular Mass Have you ever tried counting the number of rice grains in a sack of rice? Rice grains cannot be counted because their size is extremely small. Chemists face a similar problem too. As atoms are too small, it is difficult to determine their number and the mass of each atom. How do chemists overcome this problem?
Photograph 3.1 Rice
g Learnin tandard S At the end of the lesson, pupils are able to: 3.1.1 Conceptualise relative atomic mass and relative molecular mass based on the carbon-12 scale 3.1.2 Calculate relative molecular mass and relative formula mass
Relative Atomic Mass, RAM
Chemists use the concept of ‘relative atomic mass’ by comparing the mass of atom of an element to the mass of atom of another element that is chosen as the standard. Therefore, we do not need to know the actual mass of an atom. Initially, the hydrogen atom was used as the standard atom because it is the lightest atom. Masses of atoms of all other elements 12 hydrogen were compared with the hydrogen atom. For example, one carbon One carbon atom atoms atom is as heavy as 12 hydrogen atoms. Hence, the relative atomic mass of carbon is 12 while the relative atomic mass of hydrogen is assigned as one as shown in Figure 3.1. Figure 3.1 Mass of carbon atom However, in 1961, chemists across the world agreed to use compared hydrogen atom carbon-12 atom as the standard atom after finding that the usage of Brain toTeaser hydrogen atom as the standard atom encountered various problems. Brain Teaser The relative atomic mass, RAM of an element is defined as the average mass of an atom of the element compared to 1 of the mass Determining the relative 12 atomic mass of hydrogen of one carbon-12 atom. Relative atomic Average mass of one atom of the element = 1 × Mass of one carbon-12 atom mass of an element 12
Explain why there is a value of 1 in the 12 definition of the relative atomic mass based on the carbon-12 scale.
44
atom as the standard atom has encountered various problems. Try to investigate what the problems are.
One atom of carbon-12 is given a definite mass of 12 units. 1 of the mass of a 12 carbon-12 atom is the same as the mass of one hydrogen atom, that is 1 unit. So,
Brain Teaser The Mole Concept, Chemical Formula and Equation
CHAPTER 3
Brain Teaser The relative atomic mass of helium is 4. This means the average mass of one atom of helium is 4 times the mass of 1 12 of carbon-12 atom. One helium atom
4 units
Figure 3.2 Relative atomic mass of helium
One magnesium atom is twice as heavy as one atom of carbon-12. What is the RAM of magnesium?
Literacy Tips The relative atomic masses of elements are given in the Data Table of Elements on page 276. Since the relative atomic mass is a comparative value, it has no unit.
Activity 3.1 Discussing why carbon-12 is used as the standard to determine RAM 1. Carry out the activity in groups. 2. Gather information from printed reference materials or surf the Internet and discuss why carbon-12 is used as the standard to determine the relative atomic mass. 3. Present your group discussion results in a suitable thinking map. Carbon-12 is chosen as the standard because it is a solid at room temperature and thus can be handled easily. Carbon-12 combines easily with other elements. Therefore, this element is found in most substances. Although carbon has three isotopes, carbon-12 is the major isotope with the abundance of 99%. This makes the relative atomic mass of carbon-12 exactly 12.0.
Relative Molecular Mass, RMM
Similarly, we can compare the molecular mass of a substance with the standard carbon-12 atom. The relative molecular mass, RMM of a molecule is the average mass of the molecule compared to 1 12 of the mass of one carbon-12 atom. Relative molecular Average mass of one molecule = 1 × Mass of one carbon-12 atom mass of a substance 12
One molecule of water
18 units
Figure 3.3 Relative molecular mass of water
Figure 3.3 shows water molecule that has a relative molecular mass of 18. This means the mass of a water molecule is 18 times the mass of 1 of carbon-12 atom. Activity 3.2 can strengthen your 12 understanding on the concepts of relative atomic mass and relative molecular mass based on the carbon-12 scale as an analogy. 45
THEME 2
Fundamentals of Chemistry
Activity 3.2 Studying the concepts of relative atomic mass and relative molecular mass by analogy Materials: 36 washers, one 5 cm bolt, five nuts, one flat magnet and strings Apparatus: Two-pan balance Relative atomic mass based on the carbon-12 scale Washers Procedure: 1. You are given three models of carbon-12 atom. Calculate the number of washers required to form Model of carbon-12 atom Atom of element A each model of carbon-12 atom. 2. Separate the washers in each model and use them for the following steps. Atom of element B Atom of element C 3. Place an atom of element A on a two-pan balance as Photograph 3.2 shown in Figure 3.4. A model of carbon-12 atom and 4. Place the washers on the other pan one by one until three elements A, B and C they are balanced. 5. Count and record the number of washers used in Table 3.1. 6. Repeat steps 3 to 5 using atom of element B and atom of element C. 7. Calculate the relative mass of each washer in the Atom of Washers model by assuming that each atom of carbon-12 is element A Pan given the accurate mass of 12 units. Then, deduce the relative atomic masses of elements A, B and C. Figure 3.4 Studying the relative mass by analogy
Results:
Table 3.1
Atom of element
Number of washers used
Relative atomic mass
A B C
Discussion: 1. How many washers form a model of carbon-12 atom? 1 of the mass of carbon-12 atom in this activity? 2. What is represented by 12 3. Define the relative atomic mass of an element based on the carbon-12 scale. Relative molecular mass based on the carbon-12 scale
Procedure: 1. Prepare models of molecules X, Y and Z as in Photograph 3.3. 2. Place molecule X on one of the pans of the balance. 3. Place washers on the other pan one by one until they are balanced. 46
Atom A
Atom B Molecule X
Atom A
Atom B
Molecule Y
Atom C
Atom A Molecule Z
Photograph 3.3 Models of molecules X, Y and Z
The Mole Concept, Chemical Formula and Equation
CHAPTER 3
4. Count and record the number of washers used in Table 3.2. 5. Repeat steps 2 to 4 using molecule Y and Z. 6. Deduce the relative molecular masses of X, Y and Z. Results:
Table 3.2
Molecule
Composition of molecule
X
1 atom A + 1 atom B
Number of washers used
Relative molecular mass
Y Z
Discussion: 1. Based on Activity B, give the definition of relative molecular mass based on the carbon-12 scale. 2. Calculate the relative atomic masses of all the elements that form molecules X, Y and Z. 3. Compare the answers from question 2 with the relative molecular masses you obtained in Activity B. What inference can you make about the relationship between the relative molecular mass and the relative atomic mass? 4. Molecule W is formed from one atom of element A, one atom of element B and one atom of element C. Predict the relative molecular mass of W. Prepare a complete report after carrying out this activity. The relative molecular mass of a molecule can be calculated by adding up the relative atomic masses of all the atoms that form the molecule, as shown in Figure 3.5 and Example 1.
RMM of oxygen gas, O2 = 2(RAM of O) = 2(16) = 32
Literacy Tips Relative molecular mass of a molecule is similar to the total of relative atomic mass of all atoms in the molecule.
RMM of water, H2O = 2(RAM of H) + RAM of O = 2(1) + 16 = 18
Figure 3.5 Calculation of the relative molecular mass, RMM
Example
1
Calculate the relative molecular mass of glucose, C6H12O6. [Relative atomic mass: H = 1, C = 12, O = 16] Solution
RMM of glucose, C6H12O6
= 6(RAM of C) + 12(RAM of H) + 6(RAM of O) = 6(12) + 12(1) + 6(16) = 72 + 12 + 96 = 180 47
THEME 2
Fundamentals of Chemistry
Relative Formula Mass, RFM
The concept of relative mass is also used for ionic substances. The relative mass of an ionic substance is called the relative formula mass, RFM. The relative formula mass is calculated by summing up the relative atomic masses of all the atoms shown in the formula of the ionic substance. This is because the mass of an ion does not differ much from the mass of its atom that forms the ion. Check out Example 2. Example
2
Calculate the relative formula mass of zinc chloride, ZnCl2 and aluminium sulphate, Al2(SO4)3. [Relative atomic mass: O = 16, Al = 27, S = 32, Cl = 35.5, Zn = 65] Solution
RFM of zinc chloride, ZnCl2 = RAM of Zn + 2(RAM of Cl) = 65 + 2(35.5) = 65 + 71 = 136
RFM of aluminium sulphate, Al2(SO4)3 = 2(RAM of Al) + 3[RAM of S + 4(RAM of O)]
= 2(27) + 3[32 + 4(16)] = 54 + 3[96] = 342
Activity 3.3 CT Calculating the relative molecular mass and relative formula mass Determine the relative molecular mass or the relative formula mass of each of the following substances. Refer to the Data Table of Elements on page 276 to obtain the relative atomic mass. 2. O3 3. CO 4. NH3 5. N2O4 1. H2 6. C4H10 7. CuCl2 8. Zn(OH)2 9. K2Cr2O7 10. Fe(NO3)3
Activity 3.4 Tic-tac-toe with relative masses CT Materials: 10 pieces of formula cards and a tic-tac-toe card 1. Carry out the activity in pairs. Material for Activity 3.4 2. Each pair is given a tic-tac-toe card and 10 pieces of formula cards. http://bit.ly/ 2PeH6a5 Each card has the formula of a specific substance and its relative mass. 3. Shuffle the cards and put them at the centre of the table with the written side of the cards facing down. 4. The first player will take a piece of card. Without showing it to the second player, the first player will read the formula of the substance on the card to the second player. 48
The Mole Concept, Chemical Formula and Equation
CHAPTER 3
5. Referring to the Data Table of Elements on page 276, the second player will calculate the relative mass of the substance and show the answer to the first player. If the answer is correct, the second player is allowed to mark the tic-tac-toe card. If the answer is wrong, the second player will lose the chance to mark the tic-tac-toe card. 6. Repeat steps 3 to 4 with the second player taking the card while the first player calculates the relative mass of the substance. 7. Continue to take turns until one of the players succeeds in marking a complete line vertically, horizontally or diagonally or until all spaces are filled up.
TestYourself
3.1
1. Define relative atomic mass based on the carbon-12 scale. 2. Refer to the Data Table of Elements on page 276 to get the relative atomic masses. (a) How many atoms of lithium are required to equalise the mass of one atom of krypton? (b) How many atoms of helium are required to equalise the mass of a silver atom? 3. Calculate the relative molecular mass or the relative formula mass of each of the following substances: (c) Sulphuric acid, H2SO4 (a) Methane, CH4 (b) Magnesium nitrate, Mg(NO3)2 (d) Formic acid, HCOOH
3.2
Mole Concept
In our daily lives, we use units such as pairs and dozens to represent the quantity or number of objects. Photograph 3.4, shows the objects that can be quantified using the units pair and dozen. The unit pair represents 2 objects while the unit dozen represents 12 objects.
Photograph 3.4 Uses of units in daily life
In the field of chemistry, we use the unit mole to measure the amount of substance. What is the amount of substance represented by the unit mole?
g Learnin tandard S At the end of the lesson, pupils are able to: 3.2.1 Define mole 3.2.2 Interrelate the Avogadro constant, NA, the number of particles and the number of moles 3.2.3 State the meaning of molar mass 3.2.4 Interrelate the molar mass, mass and the number of moles 3.2.5 State the meaning of molar volume 3.2.6 Interrelate the molar volume, volume of gas and the number of moles 3.2.7 Solve numerical problems involving the number of particles, number of moles, mass of the substances and volume of gases
49
THEME 2
Fundamentals of Chemistry
According to the International Union of Pure and Applied Chemistry (IUPAC), the new definition of mole is as follows: The mole, with the symbol mol, is the SI unit of amount of substance. One mole of substance contains 6.02214076 × 1023 elementary entities of the substance. This number is a fixed value known as the Avogadro constant, NA that is expressed in mol–1. The Avogadro constant, NA is also called the Avogadro number. The Avogadro constant, NA is defined as the number of particles contained in one mole of substance, that is 6.02 × 1023 mol–1. In other words, 1 mol of a substance contains 6.02 × 1023 particles that form the substance. The type of particles depends on the type of substance, namely atomic substance, molecular substance or ionic substance as shown in Figure 3.6. Atomic substance
1 mol of copper, Cu = 6.02 × 1023 copper atoms, Cu
Molecular substance
1 mol of water, H2O = 6.02 × 1023 H2O molecules
For calculation at this level, Avogadro constant, NA is taken as 6.02 × 1023 mol–1, to three significant figures. Ionic substance
1 mol of sodium chloride, NaCl = 6.02 × 1023 NaCl units
1 mol of an atomic substance contains 1 mol of a molecular substance contains 1 mol of an ionic substance contains 6.02 × 1023 atoms 6.02 × 1023 formula units 6.02 × 1023 molecules
Figure 3.6 Numbers of particles in 1 mol of substance
Number of Moles and Number of Particles
The way to use mole is similar to the way of using the unit dozen. For example, 2 dozens of pencils represent 2 × 12 or 24 pencils. Similarly, the Avogadro constant, NA is used as the conversion factor between the number of moles and the number of particles. Number of moles, n =
Number of particles NA
Diagrammatically, the relationship between the number of mole and the number of particles by using Avogadro constant as the conversion factor is shown below: × NA
Number of moles 50
Number of particles ÷ NA
The Avogadro constant is named after a famous Italian scientist, Amedeo Avogadro (1776 – 1856).
The Mole Concept, Chemical Formula and Equation
CHAPTER 3
The example of conversion between the number of mole and the number of particles by using Avogadro constant, NA are shown in Examples 3, 4 and 5. [Note: Assume Avogadro constant, NA: 6.02 × 1023 mol–1] Example
3
How many atoms are there in 0.2 mol of magnesium, Mg? Solution
Number of magnesium atoms, Mg = 0.2 mol × 6.02 × 1023 mol–1 = 1.204 × 1023 atoms Example
Use the equation: Number of particles = Number of moles × NA
4
A sample of zinc chloride, ZnCl2 contains 3.01 × 1024 ZnCl2 units. Calculate the number of moles of zinc chloride, ZnCl2 found in the sample. Solution
3.01 × 1024 Number of moles of zinc chloride, ZnCl2 = 6.02 × 1023 mol–1 = 5 mol Example
Use the equation: Number of moles Number of particles = NA
5
A gas jar is filled with 2 mol of oxygen gas, O2. (a) How many molecules of oxygen are there in the gas jar? (b) How many atoms of oxygen are there in the gas jar? Solution
Use the equation:
Number of particles (a) Number of oxygen molecules, O2 = 2 mol × 6.02 × 1023 mol–1 24 = Number of moles × NA = 1.204 × 10 molecules (b) Each oxygen molecule, O2 has 2 oxygen atoms, O. Hence, the number of oxygen atoms, O = Number of O2 molecules × 2 = 1.204 × 1024 × 2 = 2.408 × 1024 atoms
Activity 3.5 Calculating the number of moles and number of particles CT [Avogadro constant, NA: 6.02 × 1023 mol–1] 1. Calculate the number of atoms found in (a) 0.1 mol of carbon, C (b) 3.5 mol of neon gas, Ne 2. Calculate the number of molecules found in (b) 0.8 mol of ammonia, NH3 (a) 1.2 mol of hydrogen gas, H2 3. Calculate the number of formula units found in (a) 3 mol of sodium chloride, NaCl (b) 0.25 mol of potassium nitrate, KNO3 51
THEME 2
Fundamentals of Chemistry
4. Calculate the number of moles of each of the following substances: (c) 9.03 × 1022 bromine molecules, Br2 (a) 6.02 × 1024 lead atoms, Pb (d) 3.612 × 1024 carbon dioxide molecules, CO2 (b) 3.02 × 1023 magnesium oxide units, MgO 25 5. A reagent bottle contains 1.806 × 10 units of copper(II) oxide, CuO. (a) How many moles of copper(II) oxide, CuO are found in the bottle? (b) Calculate the number of ions found in that bottle. 6. A sample contains 0.2 mol of ethene gas, C2H4. (a) How many ethene molecules, C2H4 are found in the sample? (b) How many hydrogen atoms, H are found in the sample? (c) Calculate the total number of atoms found in the sample.
Number of Moles and Mass of Substances
The number of moles of a substance is impossible to be determined by counting the number of particles in the substance. Therefore, to get the number of moles, the mass of a substance must be measured and we also need to know its molar mass. What is molar mass? Molar mass is the mass of one mole of substance. The unit for molar mass is gram/mol or g mol–1. Chemists found that the value of molar mass of a substance is the same as its relative mass. For example, the relative atomic mass of carbon, C is 12. Thus, the molar mass of carbon is 12 g mol–1 because 12 g of carbon, C contains 1 mol of carbon, C, that is 6.02 × 1023 atoms of carbon, C. Look at Figure 3.7 to strengthen your understanding.
• Copper consists of copper atoms. • RAM of copper = 64 • Molar mass of copper = 64 g mol–1
• Water consists of H2O molecules. • RMM of water = 2(1) + 16 = 18 • Molar mass of water = 18 g mol–1
• Sodium chloride consists of NaCl units. • RFM of sodium chloride = 23 + 35.5 = 58.5 • Molar mass of sodium chloride = 58.5 g mol–1
Figure 3.7 Determining the molar mass of substances The mass of any fraction of a mole of a substance can be weighed. For example, 12 g of carbon for 1 mol of carbon, 6 g of carbon powder for 0.5 mol of carbon and so on. The molar mass is used as the conversion factor between the number of moles and the mass of substance. The formula and the relationship between the number of moles and the mass of substance by using molar mass as the conversion factor is as follows: Number of moles, n =
Mass (g) Molar mass (g mol–1)
× Molar mass
Number of moles
Mass (g) ÷ Molar mass
52
Why is g used as the unit for mass in the formula? Mass = Number of moles × Molar mass g = mol × mol =g
The Mole Concept, Chemical Formula and Equation
CHAPTER 3
Examples of the conversion between the number of moles and the mass of the particles using molar mass are shown in Examples 6, 7 and 8. Example
6
What is the mass of 1.5 mol of aluminium, Al? [Relative atomic mass: Al = 27] Solution
Molar mass of aluminium, Al = 27 g mol–1 Mass of aluminium, Al = 1.5 mol × 27 g mol–1 = 40.5 g Example
Value of the molar mass of an atomic substance is equal to RAM. Use the formula: Mass = Number of moles × Molar mass
7
How many moles of molecules are found in 32 g of sulphur dioxide gas, SO2? [Relative atomic mass: O = 16, S = 32] Solution
Relative molecular mass of sulphur dioxide, SO2 = 32 + 2(16) Value of the molar mass of a molecular = 64 substance is equal to RMM. Thus, the molar mass of sulphur dioxide, SO2 = 64 g mol–1 32 g Use the formula: Number of moles of sulphur dioxide molecules, SO2 = Mass 64 g mol–1 Number of moles = Molar mass = 0.5 mol Example
8
How many moles are found in 4.7 g of potassium oxide, K2O? [Relative atomic mass: O = 16, K = 39] Solution
Relative formula mass of potassium oxide, K2O = 2(39) + 16 = 94 Thus, the molar mass of potassium oxide, K2O = 94 g mol–1 4.7 g Number of moles of potassium oxide, K2O = 94 g mol–1 = 0.05 mol
Value of the molar mass of an ionic substance is equal to RFM. Use the formula:
Number of moles =
Mass Molar mass
Activity 3.6 Calculating the number of moles and mass CT [Relative atomic mass: H = l, C = 12, N = 14, O = 16, Mg = 24, S = 32, Fe = 56; Avogadro constant, NA: 6.02 × 1023 mol–1] 1. Calculate the mass of each of the following substances: (a) 0.4 mol of iron fillings, Fe (b) 2.2 mol of carbon monoxide, CO 2. Calculate the number of moles in each of the following substances: (b) 8.88 g of magnesium nitrate, Mg(NO3)2 (a) 49 g of sulphuric acid, H2SO4 53
THEME 2
Fundamentals of Chemistry
3. An experiment requires 0.05 mol of ammonium sulphate crystals, (NH4)2SO4. What is the mass of ammonium sulphate, (NH4)2SO4 that should be used? 4. 0.2 mol of substance Y has the mass of 11 g. What is the molar mass of substance Y?
Number of Moles and Volume of Gases
Measuring the volume of a gas is easier compared to measuring its mass because gas is very light. How are the number of moles and the volume of a gas related? From studies, chemists found that the volume of 1 mol of any gas has similar value under the same conditions of temperature and pressure. Thus, the concept of molar volume was explained.
What is the mass of the gas in this balloon?
Molar volume is the volume occupied by 1 mol of a gas. The molar volume of any gas depends on the condition, that is 22.4 dm3 mol–1 at STP or 24 dm3 mol–1 at room conditions. Figure 3.8 Weighing the mass of a gas
This means at STP, • 1 mol of neon gas, Ne occupies 22.4 dm3 • 1 mol of nitrogen dioxide gas, NO2 occupies 22.4 dm3
Remember, the molar volume is used only for gases and not for solids or liquids.
While at room conditions, • 1 mol of neon gas, Ne occupies 24 dm3 • 1 mol of nitrogen dioxide gas, NO2 occupies 24 dm3 How do we use the molar volume to measure the number of moles of a gas? The formula and relationship between the number of moles and the volume of gas by using molar volume as a conversion factor are as follow: Number of moles, n =
Volume of gas Molar volume
× Molar volume
Number of moles
Volume of gas ÷ Molar volume
Chemistry • STP is the abbreviation for standard temperature and pressure, the condition where temperature is at 0 oC and pressure of 1 atm. • Room conditions refer to the condition where temperature is at 25 oC and pressure of 1 atm.
The conversion between the number of moles and the volume of gas using molar volume are shown in Examples 9, 10 and 11. 54
The Mole Concept, Chemical Formula and Equation
Example
CHAPTER 3
9
Calculate the volume of 2.2 mol of hydrogen gas, H2 in dm3 at STP. [Molar volume = 22.4 dm3 mol–1 at STP] Solution
Volume of gas =N umber of moles × Molar volume dm3 = mol × mol = dm3
Volume of hydrogen gas, H2 = Number of moles × Molar volume at STP = 2.2 mol × 22.4 dm3 mol–1 = 49.28 dm3 Example 10
What is the volume of 0.01 mol of ammonia gas, NH3 in cm3 at room conditions? [Molar volume = 24 dm3 mol–1 at room conditions] Solution
Volume of ammonia gas, NH3 = Number of moles × Molar volume at room conditions = 0.01 mol × 24 dm3 mol–1 = 0.24 dm3 Convert unit of volume: = 0.24 × 1 000 cm3 1 dm3 = 1 000 cm3 3 = 240 cm Example 11 How many moles of oxygen gas, O2 has the volume of 600 cm3 at room conditions? [Molar volume = 24 dm3 mol–1 at room conditions] Solution
Volume of oxygen gas, O2 = 600 cm3 600 dm3 = 1 000 = 0.6 dm3
Convert unit of volume: 1 dm3 = 1 000 cm3
Volume of gas Molar volume at room conditions 0.6 dm3 = 24 dm3 mol–1 = 0.025 mol
Alternative solution http://bit.ly/ 2MyrUmi
Number of moles of oxygen gas, O2 =
Activity 3.7 Calculating the number of moles and volume of gases CT [Molar volume of gas = 22.4 dm3 mol–1 at STP or 24 dm3 mol–1 at room conditions] 1. Calculate the volume of 0.6 mol of chlorine gas, Cl2 at STP and room conditions. 2. Calculate the number of moles of each of the following gases: (a) 48 cm3 of argon gas, Ar at room conditions (b) 39.2 dm3 of carbon dioxide gas, CO2 at STP 3. A sample contains 0.2 mol of methane gas, CH4 and 0.3 mol of ethane gas, C2H6. What is the volume of the sample at room conditions? 55
THEME 2
Fundamentals of Chemistry
Activity 3.8 Building a chart showing the relationship between the number of particles, number of moles, mass of substances and volume of gases at STP and room conditions 1. Carry out the activity in groups. 2. Discuss among the group members and build a chart on a flip chart paper that shows the relationship between the number of moles, number of particles, mass of substances and volume of gases. 3. Each member needs to copy the chart onto a small, pocket-sized card to produce a memory card. 4. Use this memory card to solve all the following numerical problems. The relationship between the number of moles, number of particles, mass of substances and volume of gases is shown in Figure 3.9. Mass (g) × NA
× Molar mass
Number of particles
÷ Molar mass × Molar volume
Volume of gas
Number of moles ÷ NA
÷ Molar volume
Figure 3.9 Relationship between the number of moles, number of particles, mass and volume of gases
Examples 12 and 13 show the function of the number of moles as a medium to convert from one quantity to another. Example 12 What is the volume of 26.4 g of carbon dioxide, CO2 at room conditions? [Relative atomic mass: C = 12, O = 16; Molar volume of gas = 24 dm3 mol–1 at room conditions] Solution Question analysis and solution plan Information from the question: Mass = 26.4 g Solution plan
÷ Molar mass
Volume of gas at room conditions?
Number of moles
Step 1
RMM of carbon dioxide, CO2 = 12 + 2(16) = 44 Thus, the molar mass of carbon dioxide, CO2 = 44 g mol–1 Mass Number of moles of carbon dioxide, CO2 = Molar mass 26.4 g = 44 g mol–1 = 0.6 mol 56
× Molar volume
Step 2 Before carrying out step 1, the molar mass must first be determined. Step 1: Mass → Number of moles
The Mole Concept, Chemical Formula and Equation
CHAPTER 3
Volume of carbon dioxide, CO2 = Number of moles × Molar volume Step 2: Number of moles → Volume = 0.6 mol × 24 dm3 mol–1 = 14.4 dm3 Hence, 26.4 g of carbon dioxide gas, CO2 occupies a volume of 14.4 dm3 at room conditions. Example 13 How many molecules are there in 672 cm3 of hydrogen gas, H2 at STP? [Avogadro constant, NA: 6.02 × 1023 mol–1; Molar volume = 22.4 dm3 mol–1 at STP] Solution Question analysis and solution plan Information from the question: Solution plan
Volume of gas = 672 cm3 ÷ Molar volume
Number of molecules?
Number of moles
Step 1
Volume of gas Molar volume 672 cm3 = 22.4 × 1 000 cm3 mol–1 Number of moles of hydrogen gas, H2 =
= 0.03 mol
Number of hydrogen molecules, H2 = Number of moles × NA = 0.03 mol × 6.02 × 1023 mol–1 = 1.806 × 1022 molecules Hence, 672 cm3 of hydrogen gas, H2 at STP consists of 1.806 × 1022 molecules.
Nandini, you need to determine the number of moles of a substance before determining the number of particles, mass or volume of a gas that is required.
× NA Step 2 Step 1: Volume → Number of moles
Step 2: Number of moles → Number of molecules Further example http://bit.ly/ 2MBDA7Z
Yes, teacher. I always refer to my memory card from Activity 3.8 to solve numerical problems until I can really understand and remember all the relationships.
57
THEME 2
Fundamentals of Chemistry
Activity 3.9 CT Solving problems involving the number of particles, number of moles, mass of substances and volume of gases at STP or room conditions 1. Carry out this activity in groups. 2. Read and answer the following questions. [Relative atomic mass: H = 1, He = 4, C = 12, N = 14, O = 16, Al = 27, S = 32; Avogadro constant, NA: 6.02 × 1023 mol–1; Molar volume = 22.4 dm3 mol–1 at STP or 24 dm3 mol–1 at room conditions] (a) Calculate the number of atoms found in the following substances: (i) 6.75 g of aluminium, Al (ii) 5.1 g of ammonia gas, NH3 (b) Calculate the volume of the following gases at STP. (ii) 1.204 × 1022 helium atoms, He (i) 5.6 g of nitrogen gas, N2 (c) What is the mass of oxygen gas, O2 that has the same number of molecules as in 8 g of sulphur trioxide gas, SO3? (d) A sample of methane gas, CH4 occupies a volume of 9.84 dm3 at room conditions. How many molecules are found in that sample? Calculate the mass of the sample. (e) A reaction releases 120 cm3 of carbon dioxide gas, CO2 per minute at room conditions. Calculate the total mass of carbon dioxide, CO2 released after 10 minutes. 3. Write the calculation steps for questions (a) to (e) clearly and systematically on a piece of flip chart paper. 4. Present your group solution in front of the class.
TestYourself
3.2
[Relative atomic mass: H = 1, C = 12, N = 14, O = 16, Na = 23, Cl = 35.5, K = 39, Fe = 56, Pb = 207; Avogadro constant, NA: 6.02 × 1023 mol–1; Molar volume = 22.4 dm3 mol–1 at STP or 24 dm3 mol–1 at room conditions] 1. Calculate the molar mass of each of the following substances: (a) Lead metal, Pb (c) Sodium nitrate, NaNO3 (b) Chloroform, CHCl3 (d) Iron(III) oxide, Fe2O3 2. Calculate the number of molecules found in 8 mol of water. 3. What is the mass of 0.5 mol of ammonia, NH3? 4. How many moles of K2O units are found in 14.1 g of potassium oxide, K2O? 5. Calculate the volume of 16 g of oxygen gas, O2 at STP. 6. The mass of 4 dm3 of a gas is 12 g at room conditions. Calculate the molar mass of the gas. 7.
4 g of hydrogen gas, H2 has greater number of molecules than 14 g of nitrogen gas, N2
Do you agree with the above statement? Give your reason. 58
The Mole Concept, Chemical Formula and Equation
3.3
Chemical Formula
Photograph 3.5 A chemical formula represents a chemical substance
Chemical formula is a representation of a chemical substance using alphabets to represent the atoms and subscript numbers to show the number of each type of atoms found in the elementary entities of the substance. Examples of chemical formulae of elements and compounds are shown in Figure 3.10. Elements Substance: Magnesium Chemical formula: Mg
Substance: Oxygen gas Chemical formula: O2
The chemical formula shows that magnesium consists of magnesium atoms only. The chemical formula shows that oxygen gas molecule consists of two oxygen atoms.
Compounds Substance: Water Chemical formula: H2O
The subscript number shows that two atoms of hydrogen combine with one atom of oxygen.
Substance: Aluminium oxide Chemical formula: Al2O3
The subscript number shows that two atoms of aluminium combine with three atoms of oxygen.
CHAPTER 3
g Learnin tandard S At the end of the lesson, pupils are able to: 3.3.1 State the meaning of chemical formula, empirical formula and molecular formula 3.3.2 Determine the empirical formula of magnesium oxide, MgO through an activity 3.3.3 Determine the empirical formula of copper(II) oxide, CuO through an activity 3.3.4 Solve numerical problems involving empirical formula and molecular formula 3.3.5 Construct chemical formulae of compounds
Chemistry Elements are substances that consist of only one type of atoms. Elements like metals and inert gases are atomic substances while elements such as oxygen gas are molecular substances.
Literacy Tips The subscript number 1 need not be written in a chemical formula.
Figure 3.10 Chemical formulae of elements and compounds
59
THEME 2
Fundamentals of Chemistry
Empirical Formula and Molecular Formula
In general, compounds can be represented by two types of chemical formulae, namely the empirical formula and the molecular formula. What are the empirical formula and the molecular formula?
Activity 3.10 CT Gathering and interpreting information involving chemical formulae, empirical formulae and molecular formulae 1. Carry out this activity in groups. 2. Gather information on chemical formulae, empirical formulae and molecular formulae by referring to reading materials or surfing the Internet. 3. Based on the information gathered, construct a suitable thinking map to show the difference between the empirical formula and the molecular formula using a suitable computer software. 4. List out the examples of chemical formulae in a table and use this list throughout your lesson.
The empirical formula is the chemical formula that shows the simplest ratio of the number of atoms of each element in a compound. The molecular formula, on the other hand, is the chemical formula that shows the actual number of atoms of each element found in a molecule of a compound. Figure 3.11 shows the difference between the empirical formula and the molecular formula. Empirical formula of glucose:
Molecular formula of glucose:
CH2O
C6H12O6
The formula shows one molecule of glucose consists of 6 carbon atoms, 12 hydrogen atoms and 6 oxygen atoms.
Ratio of number of atoms C atom : H atom : O atom = 6 : 12 : 6 = 1 : 2 : 1
The formula shows that the simplest ratio of the number of carbon atoms to hydrogen atoms and oxygen atoms is 1:2:1.
Figure 3.11 Molecular formula and empirical formula of glucose Table 3.3 Molecular formulae and empirical formulae of several substances
60
Substance
Molecular formula
Empirical formula
Water
H2O
H2O
Ammonia
NH3
NH3
Hydrazine
N2H4
NH2
Propene
C3H6
CH2
Benzene
C6H6
CH
Brain Teaser Some compounds have the same empirical formula and molecular formula. However, there are some other compounds that have different empirical formula and molecular formula. Try to think why it is so.
Brain The Mole Concept, Chemical Formula andTeaser Equation Determination of an Empirical Formula
CHAPTER 3
Brain Teaser
The empirical formula is obtained by the analysis of percentage composition of a substance. This is done by determining the simplest whole number ratio of atoms of each element that combines through an experiment. Example 14 is used as a guide to solve Activity 3.11.
Hexane is an organic solvent that is widely used in the food industry. The molecular formula of hexane is C6H14. What is its empirical formula?
Example 14 1.35 g of aluminium combines with 1.2 g of oxygen to form aluminium oxide. What is the empirical formula of aluminium oxide? [Relative atomic mass: O = 16; Al = 27] Solution Element
Al
O
Mass (g)
1.35
1.2
1.35 = 0.05 27
1.2 = 0.075 16
n=
0.05 = 1 0.05
0.075 = 1.5 0.05
Divide each number with the smallest number, that is 0.05.
2
3
Number of moles of atoms Mole ratio Simplest mole ratio of atom
Determine the mass of each element. Mass Molar mass
Multiply each answer by 2 to get the simplest whole number ratio.
2 mol of aluminium atoms combine with 3 mol of oxygen atoms. Thus, the empirical formula of aluminium oxide is Al2O3.
Activity 3.11 Determining the empirical formulae CT [Relative atomic mass: H = 1, C = 12, O = 16, Cl = 35.5, K = 39, Br = 80, Sn = 119, I = 127] 1. A sample of potassium bromide contains 6.24 g of potassium and 12.8 g of bromine. What is the empirical formula of potassium bromide? 2. A sample of 26.1 g of tin chloride contains 11.9 g of tin. State the empirical formula of the tin chloride. 3. 0.03 mol of element Y combines with 7.62 g of iodine to produce an iodide salt. State the empirical formula of the iodide salt. 4. A chemist analysed the compound that gives smell to fully ripe bananas. He found that the compound contains 64.62% carbon, 10.77% hydrogen and 24.61% oxygen. What is the empirical formula of that compound? Photograph 3.6 Bananas Using the calculation skills learned, the determination of the empirical formulae of magnesium oxide and copper(II) oxide can be carried out through Activity 3.12 and 3.13. 61
THEME 2
Fundamentals of Chemistry
Activity 3.12 Aim: To determine the empirical formula of magnesium oxide. Materials: 10 cm magnesium ribbon and sand paper Apparatus: Crucible with lid, tongs, Bunsen burner, tripod stand, pipeclay triangle and electronic balance Lid Procedure: Magnesium 1. Weigh and record the mass of a crucible together Crucible ribbon with its lid. Pipeclay triangle 2. Rub 10 cm magnesium ribbon with a sand paper Heat until shiny. Coil the magnesium ribbon and put it in the crucible. 3. Weigh and record the mass of the crucible together Figure 3.12 with its lid and the coil of magnesium ribbon. 4. Set up the apparatus as shown in Figure 3.12. Safety Precaution 5. First, heat the crucible without its lid. 6. When magnesium ribbon starts to burn, close the crucible Prevent white fumes in the with its lid. crucible from escaping when 7. Using a pair of tongs, lift the lid slightly from time to time carrying out Step 7. and quickly place it back. 8. When the burning of magnesium ribbon is complete, take off the lid and heat the crucible with high temperature for 1 to 2 minutes. 9. Put back the lid of the crucible and allow it to cool to room temperature. 10. Weigh the mass of crucible together with its lid and its contents again. 11. Repeat the heating, cooling and weighing process until a constant mass is obtained. 12. Record the constant mass in Table 3.4.
Results:
Table 3.4 Description
Mass (g)
Crucible + lid Crucible + lid + magnesium ribbon Crucible + lid + magnesium oxide
Interpreting data: 1. Based on your results, determine the masses of magnesium and oxygen that combine. 2. Determine the empirical formula of magnesium oxide.
Discussion: 1. What is the purpose of rubbing the magnesium ribbon with a sand paper before using it? 2. Name the white fumes that are produced. 3. Why are steps 6, 7 and 11 performed? 4. What will happen if the white fumes are released into the environment? Prepare a complete report after carrying out this activity. 62
The Mole Concept, Chemical Formula and Equation
CHAPTER 3
Activity 3.13 Aim: To determine the empirical formula of copper(II) oxide. Materials: Water, copper(II) oxide powder, zinc granules, 1.0 mol dm–3 hydrochloric acid, wooden splinter and cotton buds Apparatus: Boiling tube, rubber stoppers, rubber tube, 12 cm glass tube, 10 cm glass tube, spirit lamp, retort stand with clamp, wooden block, electronic balance and spatula Procedure: 1. Weigh the mass of 12 cm glass tube using an electronic balance and record its mass. Guideline to determine the empirical formula 2. Put some copper(II) oxide powder into the glass tube. Use the wooden splinter to move copper(II) oxide powder of copper(II) oxide http://bit.ly/2VLQHq6 to the middle of the glass tube. Weigh the mass of the glass tube together with its contents and record the mass. 2 3. Fill of the boiling tube with water. 3 4. Close the boiling tube with a rubber stopper that has a 12 cm glass tube. Clamp the boiling tube onto the retort stand. 5. Insert a few zinc granules into another boiling tube. Add 1.0 mol dm–3 hydrochloric acid into 1 the boiling tube until it is full. 3 6. Close the boiling tube with a rubber stopper that has a 10 cm glass tube. Clamp the boiling tube onto the other retort stand. 7. Connect the glass tube that contains copper(II) oxide powder as shown in Figure 3.13. Glass tube (12 cm) Rubber tube Air opening Glass tube (12 cm)
Copper(II) oxide powder
Ethanol
Glass tube (10 cm)
1.0 mol dm–3 hydrochloric acid
Water
Zinc granules Wooden block
Figure 3.13
8. Let the hydrogen gas flow for 10 seconds by allowing the air bubbles to be released in the water before starting the heating process. 9. Heat copper(II) oxide using a spirit lamp with a continuous flow of hydrogen gas through the glass tube. 10. Stop the heating when the black colour of copper(II) oxide turns brown completely. 11. Keep a continuous flow of hydrogen gas until the glass tube is cooled back to room temperature.
Safety Precaution If necessary, hold the spirit lamp by moving it under the glass tube to heat the remaining powder that is still black so that all black powder turns brown.
63
THEME 2
Fundamentals of Chemistry
12. Remove the glass tube that contains brown powder. Eliminate water drops at the end of the glass tube with a cotton bud. 13. Weigh the mass of the glass tube together with its contents and record its mass. 14. Repeat the heating, cooling and weighing processes from steps 9 to 13 until a constant mass reading is obtained. 15. Record the constant mass in Table 3.5.
Results:
Another method of determining the empirical formula of copper(II) oxide http://bit.ly/ 2BeHBbY
Table 3.5 Description
Mass (g)
Glass tube Glass tube + copper(II) oxide Glass tube + copper Copper Oxygen
Interpreting data: 1. Determine the empirical formula of copper(II) oxide in this activity.
Discussion: 1. What is the purpose of using zinc granules and hydrochloric acid in this activity? 2. Why does the hydrogen gas need to flow continuously for a while before starting the heating process? 3. The hydrogen gas is allowed to flow until the product of heating is at room temperature in step 11. Why? 4. Why do the heating, cooling and weighing processes need to be repeated until a constant mass is obtained? Prepare a complete report after carrying out this activity. For reactive metals like magnesium, the metal needs to be heated only slightly before it can react with the oxygen in the air. Figure 3.14 shows how the mass of magnesium and oxygen that combine are calculated to determine the simplest mole ratio of atom. Can you name another reactive metal oxide which empirical formula can be determined using the same method? Magnesium + Oxygen ˜ Magnesium oxide This mass is determined before the reaction takes place.
This mass is calculated as the difference between the masses of magnesium and magnesium oxide.
This mass is determined after the reaction takes place.
Figure 3.14 Calculation of the mass of magnesium and oxygen in magnesium oxide
64
The Mole Concept, Chemical Formula and Equation
However, this method is not suitable in determining the empirical formula of copper(II) oxide because copper is less reactive towards oxygen. Hence, copper(II) oxide is heated in a stream of hydrogen gas so that hydrogen can remove oxygen from the oxide as shown in Figure 3.15.
CHAPTER 3
Reactivity series of metals http://bit.ly/ 2pFVTQb
Copper(II) oxide + Hydrogen ˜ Copper + Water This mass is determined before the reaction takes place.
The mass of oxygen is calculated as the difference between the masses of copper(II) oxide and copper.
This mass is determined after the reaction takes place.
Figure 3.15 Calculation of the masses of copper and oxygen in copper(II) oxide
Determination of a Molecular Formula
The molecular formula of a compound is its multiplied empirical formula. Molecular formula = (Empirical formula)n The value of n is a positive integer. Table 3.6 shows several examples. Table 3.6 Relationship between the molecular formula and the empirical formula Substance
Water
Hydrazine
Propene
Benzene
Empirical formula
H 2O
NH2
CH2
CH
Molecular formula
(H2O)1= H2O
(NH2)2 = N2H4
(CH2)3 = C3H6
(CH)6 = C6H6
1
2
3
6
n
Therefore, to determine the molecular formula of a compound, we first need to know its empirical formula. Examples 15 and 16 show solutions regarding chemical formula. Example 15 A compound has the empirical formula CH2. Its relative molecular mass is 56. What is the molecular formula of the compound? [Relative atomic mass: H = 1; C = 12] Solution
Assume that the molecular formula of the compound is (CH2)n. Based on its molecular formula, the RMM of compound = n[12 + 2(1)] = 14n Given the RMM of compound, 14n = 56 Equate the calculated RMM 56 with the given one. n= 14 =4 Hence, the molecular formula of the compound is C4H8. 65
THEME 2
Fundamentals of Chemistry
Example 16 1.2 g of element Y reacts with bromine to form 6 g of a compound with the empirical formula YBr2. Determine the relative atomic mass of Y. [Relative atomic mass: Br = 80] Solution
The compound consists of element Y and bromine. Therefore, the mass of element Y + mass of bromine = mass of compound formed 1.2 g + mass of bromine = 6 g Mass of bromine = (6 − 1.2) g = 4.8 g Further Assume that the RAM of element Y is x. example
Element
Y
Br
Mass (g)
1.2
4.8
1.2 = ? x
4.8 = 0.06 80
Number of moles of atoms
http://bit.ly/ 32BiQ5J
Based on the empirical formula YBr2, 2 mol of Br atoms combine with 1 mol of atom Y or 1 mol of Br atoms combine with 0.5 mol of atom Y or Based on the empirical formula, 0.06 mol of Br atoms combine with 0.03 mol of atom Y. calculate using the right ratio. Hence, the number of moles of atom Y that reacts = 0.03 mol 1.2 = 0.03 x x = 1.2 0.03 = 40 The RAM of element Y is 40.
Activity 3.14 Solving numerical problems involving empirical formulae and CT molecular formulae [Relative atomic mass: H = 1, C = 12, N = 14, O = 16, Ca = 40, Zn = 65] 1. Ethanoic acid has a molar mass of 60 g mol–1. If its empirical formula is CH2O, determine the molecular formula of ethanoic acid. 2. Hydrocarbons consist of carbon and hydrogen. 5.7 g of a hydrocarbon contains 4.8 g of carbon. If the relative molecular mass of the hydrocarbon is 114, determine its molecular formula. 3. What is the mass of zinc required to combine with 0.5 mol of chlorine to produce zinc chloride, ZnCl2? 4. Assume you are a farmer. You want to choose a fertiliser with a high nitrogen content for your plants. Three types of commonly used fertilisers are as follows. Ammonium nitrate, NH4NO3
Urea, CO(NH2)2
Nitrosol, Ca(NO3)2
Which fertiliser would you choose? Give reasons for your choice. Show the steps used in the calculation. 66
The Mole Concept, Chemical Formula and Equation
CHAPTER 3
Chemical Formulae of Ionic Compounds
Ionic compounds are made up of cations (positively-charged ions) and anions (negatively-charged ions). In order to write the chemical formulae of ionic compounds, you need to know the formulae of cations and anions. Table 3.7 shows the examples of formulae for cations and anions that are commonly used. Figure 3.16 explains how the chemical formula of an ionic compound is constructed. Table 3.7 Formulae of common cations and anions Cation
Formula of cation
Anion
Formula of anion
Sodium ion
Na
Oxide ion
O2–
Potassium ion
K+
Chloride ion
Cl–
Aluminium ion
Al3+
Bromide ion
Br–
Zinc ion
Zn2+
Iodide ion
Magnesium ion
Mg2+
Hydroxide ion
OH–
Iron(II) ion
Fe2+
Carbonate ion
CO32–
Iron(III) ion
Fe3+
Nitrate ion
NO3–
Copper(II) ion
Cu2+
Sulphate ion
SO42–
Calcium ion
Ca2+
Phosphate ion
PO43–
Silver ion
Ag+
Manganate(VII) ion
MnO4–
Lead(II) ion
Pb2+
Thiosulphate ion
S2O32–
Ammonium ion
NH4+
Dichromate(VI) ion
Cr2O72–
+
I–
Name: Zinc chloride Cation: Zinc ion
Anion: Chloride ion
Zn2+
Cl–
1. Based on the name of the compound, determine the cation and anion.
Zn2+
Cl–
2. Cross-change the cation charge and anion charge to determine the number of cations and anions.
The number of ion: 1
2
Check: Positive charge : 1 × (+2) = +2 Negative charge : 2 × (–1) = –2 Total charge : 0 Formula:
ZnCl2
Figure 3.16 Constructing the chemical formula of zinc chloride via cross-change method
Further example on cross-change method 3. Write the chemical formula of the compound. The formula is neutral. The charges of ions are not written in the formula. The subscript number is used to show the number of ions.
http://bit.ly/ 32DGbUu
The basic concept of constructing a chemical formula of an ionic compound http://bit.ly/ 35WMLam
67
THEME 2
Fundamentals of Chemistry
Activity 3.15 Constructing the chemical formulae of ionic compounds 1. Carry out this activity individually. 2. Scan the QR code and download the diagram of ionic formula cards. 3. Print and cut out the ionic formula cards. 4. Use the ion formula cards to help you determine the chemical formula of each of the following ionic compounds: Potassium oxide Sodium chloride Calcium bromide
Sodium hydroxide Aluminium oxide Zinc sulphate
CT
Diagram of ionic formula cards http://bit.ly/ 2N4JVaG
Magnesium nitrate Potassium carbonate Copper(II) sulphate
Calcium nitrate Aluminium chloride Sodium carbonate
5. Record your answers systematically in a table.
Naming of Chemical Compounds For ionic compounds, the name of the cation is written first followed by the name of the anion as in Table 3.8. Table 3.8 Examples in the naming of ionic compounds Cation
Anion
Name of ionic compound
Sodium ion
Chloride ion
Sodium chloride
Zinc ion
Bromide ion
Zinc bromide
Magnesium ion
Nitrate ion
Magnesium nitrate
Chemistry Chemical compounds are named systematically as recommended by the International Union of Pure and Applied Chemistry (IUPAC).
Some metals form more than one type of ions. In order to distinguish these ions, Roman numerals are used in their naming. For example, iron forms two types of cations, namely Fe2+ and Fe3+. Fe2+ ion is named as iron(II) ion while Fe3+ ion is named as iron(III) ion. Take a look at the names of the following compounds: Shows iron(II) ion, Fe2+
Iron(II) oxide Iron(III) oxide
Shows iron(III) ion, Fe3+
When naming simple molecular compounds, the more electropositive element is named first followed by the name of the more electronegative element. The name of the first element remains the same while the second element ends with ‘ide’. Greek prefixes are used to represent the number of atoms of each element in simple molecular compounds. Look at the examples below. CO – Carbon monoxide NO2 – Nitrogen dioxide SO3 – Sulphur trioxide
68
Greek prefixes like ‘mono’, ‘di’ and ‘tri’ show the numbers one, two and three respectively.
Literacy Tips Other Greek prefixes are as follows: tetra – 4 hex – 6 pent – 5 hept – 7
The Mole Concept, Chemical Formula and Equation
CHAPTER 3
Activity 3.16 Naming compounds 1. Name the ionic compounds with the following formulae: (c) Mg(NO3)2 (e) Na2SO4 (a) CaCl2 (b) KBr (d) ZnCO3 (f) NH4Cl 2. Name the molecular compounds with the following formulae: (e) BF3 (a) NO (c) SO3 (b) CO2 (d) CCl4 (f) CS2 3. The molecule of a compound consists of two nitrogen atoms and three oxygen atoms. Name the compound.
TestYourself
CT
3.3
1. What is meant by empirical formula and molecular formula? 2. Caffeine, C8H10N4O2 is a natural stimulant found in coffee, tea and cocoa. What is the empirical formula of caffeine? 3. Calcium carbonate and sodium fluoride are two compounds found in toothpaste. Write the chemical formulae of both compounds. 4. A sample of 5.04 g of oxide for phosphorus contains 2.48 g of phosphorus. [Relative atomic mass: O = 16, F = 31] (a) If the relative molecular mass of the oxide is 126, determine its empirical formula and molecular formula. (b) Name the oxide of the phosphorus.
3.4
Chemical Equation
Photograph 3.7 Burning of an oil lamp
g Learnin tandard S At the end of the lesson, pupils are able to: 3.4.1 Write balanced chemical equations 3.4.2 Interpret chemical equations quantitatively and qualitatively 3.4.3 Solve stoichiometry numerical problems
Did you know that the burning of fuel and the digestion of food in our bodies are all chemical reactions? Chemists have a simple and accurate way to describe chemical reactions, that is through chemical equations. 69
THEME 2
Fundamentals of Chemistry
How to Write Chemical Equations
Chemical equations can be written in the form of words or using chemical formulae. The starting substances or reactants are written on the left-hand side of the equation while the new substances formed or products are written on the right-hand side of the equation. The arrow ‘→’ means ‘produces’. The physical state of each substance, whether solid(s), liquid(l), gas(g) or aqueous solution(aq) is usually indicated in a chemical equation. Figure 3.17 shows the examples of writing the chemical equation for the reaction between hydrogen and oxygen. Reactants
Product
Hydrogen
+
Oxygen
→
Water
H2
+
O2
→
H2 O
+ O2 H2O H2 → (2 H atoms) (2 O atoms) (2 H atoms, 1 O atom) Equation is not balanced 2H2 + O2 (4 H atoms) (2 O atoms) 2H2(g)
+
O2(g)
→ →
2H2O (4 H atoms, 2 O atoms) 2H2 O(l)
1. Write the equation in words. 2. Write down the chemical formula of each reactant and product. 3. Check whether the equation is balanced. 4. Balance the equation by adjusting the in front of the chemical formula. 5. Write the physical state of each reactant and product.
Figure 3.17 Writing the chemical equation for the reaction between hydrogen and oxygen
Chemical equations need to be balanced. Based on the law of conservation of mass, matter can neither be created nor destroyed. Therefore, the number of atoms of each element on both sides of the equation must be the same.
Simulation on balancing chemical equation http://bit.ly/33vr5QQ
Activity 3.17 Balancing chemical equations CT 1. Write a balanced chemical equation for each of the following reactions: (a) Nitrogen gas + Hydrogen gas → Ammonia gas (b) Sodium metal + Water → Aqueous solution of sodium hydroxide + Hydrogen gas (c) Solid copper(II) carbonate decomposes into solid copper(II) oxide and carbon dioxide gas when heated. (d) Burning of aluminium powder in excess oxygen produces Chemistry white aluminium oxide powder. 2. Balance the following chemical equations: (a) KI(aq) + Br2(aq) → I2(s) + KBr(aq) (b) Zn(s) + AgNO3(aq) → Zn(NO3)2(aq) + Ag(s) (c) C3H8(g) + O2(g) → CO2(g) + H2O(l) (d) AgNO3(s)→Ag(s) + NO2(g) + O2(g) Δ 70
Sometimes, chemical equations also show the condition of the reactions. For example, the Greek letter delta (Δ) below the arrow shows that heating is required in the chemical reaction.
The Mole Concept, Chemical Formula and Equation
CHAPTER 3
Activity 3.18 Aim: To write balanced chemical equations. Materials: Copper(II) carbonate powder, CuCO3, concentrated Concentrated hydrochloric hydrochloric acid, HCl, concentrated ammonia acid and concentrated solution, NH3, lead(II) nitrate solution, Pb(NO3)2, ammonia are corrosive. potassium iodide solution, KI and limewater Handle them with care and carry out Activity 3.18 in the Apparatus: Test tubes, delivery tube and rubber stopper, test tube fume chamber. 3 holder, Bunsen burner, 10 cm measuring cyclinder, test tube stoppers and glass tube Procedure: Heating of copper(II) carbonate, CuCO3 Copper(II) carbonate powder 1. Fill a spatula of copper(II) carbonate powder, CuCO3 into a test tube. Observe the colour of the powder. 2. Set up the apparatus as shown in Figure 3.18. Heat 3. Heat copper(II) carbonate, CuCO3 and let the gas produced flow into the test tube filled with limewater. Observe the changes that take place in both test tubes. Limewater 4. When the reaction is completed, remove the test tube of limewater. Figure 3.18 Then, stop the heating. 5. Record your observations. Formation of ammonium chloride, NH4Cl 1. Using a glass tube, put 3 or 4 drops of concentrated hydrochloric acid, HCl into a test tube. Close the test tube with a stopper and leave it for a few minutes. 2. Repeat step 1 using concentrated ammonia solution, NH3 in another test tube. 3. Remove the stoppers from both test tubes. Quickly bring the mouths of both test tubes together as shown in Figure 3.19. 4. Observe and record the changes that take place.
Precipitation of lead(II) iodide, PbI2 1. Pour 2 cm3 of lead(II) nitrate solution, Pb(NO3)2 into a test tube. 2. Pour 2 cm3 of potassium iodide solution, KI into another test tube. 3. Pour potassium iodide solution, KI into lead(II) nitrate solution, Pb(NO3)2 as shown in Figure 3.20. Shake the mixture. 4. Observe and record the changes that take place. Discussion: 1. For each reaction in experiments A, B and C, state: (a) The reactants and products (b) The physical state of each reactant and product (c) The chemical formula of each reactant and product 2. Write a balanced chemical equation for each of the reactions.
Hydrogen chloride gas
Ammonia gas
Figure 3.19
Potassium iodide solution
Lead(II) nitrate solution
Figure 3.20
Prepare a complete report after carrying out this activity. 71
THEME 2
Fundamentals of Chemistry
Using Chemical Equations Chemical equation can be interpreted qualitatively and quantitatively. From the qualitative aspect, chemical equations enable us to identify the reactants and products as well as their physical states. 2Na(s)
+
Cl2(g) →
Reactants: Sodium metal and chlorine gas
2NaCl(s) Product: Solid sodium chloride
From the quantitative aspect, we can study the stoichiometry of chemical equations. Stoichiometry is the quantitative study of the composition of substances involved in a chemical reaction. Coefficients in chemical equations show the ratio of substances involved, either as the ratio of elementary entities of substance or the mole ratio. Take a look at the following example: → 2NaCl(s) 2Na(s) + Cl2(g) (2 atoms) (1 molecule) (2 formula units)
or
or
or
(2 mol)
(1 mol)
(2 mol)
Ratio of basic entities (particles): Two sodium atoms react with one molecule of chlorine to produce two NaCl units.
Mole ratio: 2 mol of sodium react with 1 mol of chlorine gas to produce 2 mol of sodium chloride.
Activity 3.19 Century Interpreting chemical equations qualitatively and quantitatively 21st Skills 1. Carry out the Think-Pair-Share activity. 2. Based on the chemical equations obtained from Activity 3.18, interpret each equation qualitatively and quantitatively, from the aspects of ratio of elementary entities and mole ratio. 3. Discuss with your partner. 4. Share the results of your discussion with the class.
Based on the mole ratio of substances from a balanced chemical equation, we can solve various numerical problems by calculating the number of moles of substances required in the right ratio. 2Na(s)
+
Cl2(g)
→
2NaCl(s)
(2 mol)
(1 mol)
(2 mol)
Initial mole ratio from the stoichiometry
(1 mol)
(0.5 mol)
(1 mol)
All values are divided by 2 Calculated in the right ratio for other values
72
The Mole Concept, Chemical Formula and Equation
CHAPTER 3
The number of moles determined can be converted to mass, number of particles or volume of gas using the molar mass, Avogadro constant or molar volume like all the relationships you have learned before. Example 17 Burning of aluminium in air is as follows:
4Al(s) + 3O2(g) → 2Al2O3(s) What is the mass of aluminium oxide produced if 5.4 g of aluminium is burnt completely in air? [Relative atomic mass: O = 16, Al = 27] Solution Question analysis and solution plan Equation: 4Al(s) + Information from the equation: (4 mol) Information from the question: (5.4 g)
3O2(g)
Step 1 How many moles?
→ 2Al2O3(s) (2 mol) (? g – question to be answered)
Step 2
Step 3
Mole ratio?
What is the mass?
Mass Step 1: Molar mass Mass of Al → Number of moles of Al 5.4 g = 27 g mol−1 = 0.2 mol Based on the equation, 4 mol of aluminium, Al produces 2 mol Step 2: of aluminium oxide, Al2O3. Therefore, 0.2 mol of aluminium, Calculate the mole ratio of Al2O3. Al produces 0.1 mol of aluminium oxide, Al2O3. Hence, the mass of aluminium oxide, Al2O3 produced = Number of moles × Molar mass Step 3: = 0.1 mol × [2(27) + 3(16)] g mol–1 Number of moles of Al2O3 → Mass of Al2O3 = 0.1 mol × 102 g mol–1 = 10.2 g Number of moles in 5.4 g of aluminium, Al =
Activity 3.20 Solving numerical stoichiometry problems CT [Relative atomic mass: H = 1, C = 12, O = 16, Cl = 35.5, Ca = 40, Fe = 56, Zn = 65; Avogadro constant, NA: 6.02 × 1023 mol–1; Molar volume = 22.4 dm3 mol–1 at STP or 24 dm3 mol–1 at room conditions] 1. Decomposition of calcium carbonate by heating is as follows: CaCO3(s) → CaO(s) + CO2(g) Δ What is the mass of calcium carbonate required to produce 1.2 dm3 of carbon dioxide gas, CO2 at room conditions? 73
THEME 2
Fundamentals of Chemistry
2. Zinc reacts with hydrochloric acid as follows: Zn(s) + 2HCl(aq) → ZnCl2(aq) + H2(g) What is the mass of zinc that should be used to produce 0.5 mol of hydrogen gas, H2? 3. A sample of iron(III) oxide, Fe2O3 is heated in a stream of excess hydrogen gas, H2 to produce 5.6 g of iron metal according to the following equation: Fe2O3(g) + 3H2(g) → 2Fe(s) + 3H2O(l) Calculate the mass of the iron(III) oxide sample. 4. Nitrogen and hydrogen gases react according to the following equation: N2(g) + 3H2(g) → 2NH3(g) How many molecules of ammonia, NH3 are produced if 6.72 dm3 of nitrogen gas at STP reacts completely with hydrogen gas?
Activity 3.21 CT Creating a computer worksheet Decomposition of potassium chlorate(V), KClO3 by heat is often used to produce oxygen gas in the laboratory.
2KClO3(s) → 2KCl(s) + 3O2(g) Assume you are a laboratory assistant. You are required to prepare different amounts of oxygen gas from time to time. Repeated calculations using chemical equations can be simplified using a computer worksheet. Use Microsoft Excel or other suitable programmes to prepare a computer worksheet involving the equation above to solve the following problems: [Relative atomic mass: O = 16, Cl = 35.5, K = 39; Molar volume = 24 dm3 mol–1 at room conditions] 1. What are the masses of potassium chlorate(V), KClO3 needed to produce 1 dm3, 5 dm3, 10 dm3, 20 dm3 and 50 dm3 of oxygen gas? 2. What are the volumes of oxygen gas produced if 0.25 kg, 0.5 kg, 1 kg, 1.5 kg and 2 kg of potassium chlorate(V), KClO3 are used?
TestYourself
3.4
1. Write the chemical equations for the following reactions: (a) Copper + Silver nitrate solution → Copper(II) nitrate solution + Silver (b) Hot zinc metal will react with chlorine gas to produce solid zinc chloride 2. Decomposition of hydrogen peroxide, H2O2 occurs according to the following equation: 2H2O2(l) → 2H2O(l) + O2(g) (a) What are the products of the decomposition of hydrogen peroxide, H2O2? (b) Calculate the volume of oxygen produced at STP from the decomposition of 30.6 g of hydrogen peroxide, H2O2. 74
The Mole Concept, Chemical Formula and Equation
CHAPTER 3
Concept Chemical Equation Volume of gas
importance
÷ molar volume
× molar volume
consists of
Stoichiometry
Chemical Formula
information obtained
types
Number of Moles, n
Molecular formula
Empirical formula
× NA ÷ NA
Number of particles
Quick http://bit.ly/ 31DFXv1
×m olar mass ÷m olar mass
Mass (g)
involved in the calculations of
Relative mass
Relative atomic mass Relative molecular mass Relative formula mass
Self Reflection 1. What is interesting about The Mole Concept, Chemical Formula and Equation? 2. Why is the learning of The Mole Concept, Chemical Formula and Equation important in the next chemistry lesson? 3. Rate your performance in The Mole Concept, Chemical Formula and Equation on a scale of 1 to 10; 1 being the lowest and 10 the highest. Why would you rate yourself at that level? 4. What can you do to improve your mastery in The Mole Concept, Chemical Formula and Equation? http://bit.ly/ 5. What else would you like to know about The Mole Concept, 2MiTOlY Chemical Formula and Equation?
75
THEME 2
Fundamentals of Chemistry
3
Achievement
Refer to the Data Table of Elements on page 276. [Avogadro constant, NA: 6.02 × 1023 mol–1; Molar volume = 22.4 dm3 mol–1 at STP or 24 dm3 mol–1 at room conditions] 1. What is meant by molar mass and molar volume? 2. What is the relationship between Avogadro constant, number of particles and number of moles? 3.
Relative atomic mass of nitrogen is 14
State the meaning of the above statement based on the carbon-12 scale. 4. Vitamin C or ascorbic acid is an important antioxidant required for our health. Vitamin C has the molecular formula C6H8O6. (a) What is the empirical formula of vitamin C? (b) What is the relative molecular mass of vitamin C?
Relief Antacid Ingredients
Aluminium hydroxide 160 mg Magnesium carbonate 105 mg Relieves flatulence Relieves heartburn
100 Tablets
Active substance (per tablet) Function Aluminium hydroxide 160 mg .......... Antacid Magnesium carbonate 105 mg .........Antacid
5. Antacid functions to relieve gastric problems. Figure 1 shows the label on a bottle of antacid.
Figure 1
Give the chemical formulae of the two active ingredients in the antacid. 6. Figure 2 shows the aerobic respiration in our body cells to produce energy from glucose, C6H12O6. Write a balanced chemical equation for the process of aerobic respiration.
Glucose
Oxygen
Glucose + Oxygen Carbon dioxide + Water Carbon dioxide
Figure 2
7. Iron(II) sulphate heptahydrate, FeSO4.7H2O is often used to treat anaemic patients suffering from the lack of iron mineral. (a) What is the molar mass of iron(II) sulphate heptahydrate? (b) Calculate the percentage of iron in iron(II) sulphate heptahydrate. 76
The Mole Concept, Chemical Formula and Equation
CHAPTER 3
8. Figure 3 shows the weighing steps taken in the determination of the empirical formula of the oxide of metal Y. Combustion tube Porcelain boat ON OFF
230.3024 g
Metal Y Metal Y oxide
ON OFF
242.2501 g
ON OFF
240.6444 g
Figure 3
Determine the empirical formula of metal Y oxide. [Relative atomic mass: O = 16, Y = 207] 9. P, Q and R are three samples of chemical substances. P – 0.2 mol of calcium chloride Q – 12 dm3 of nitrogen monoxide gas at room conditions R – 2.408 × 1023 carbon dioxide molecules
Arrange the three samples in ascending order of mass. 10. In your opinion, between the empirical formula and the molecular formula, which formula is more suitable to be used when writing chemical equations? Give your reasons.
Enrichmen Corner 1. When steam is passed over a hot iron metal, hydrogen gas and iron(III) oxide are formed. What is the mass of steam required to react completely with 100 g of iron? [Relative atomic mass: H = 1, O = 16, Fe = 56] 2. Lithium hydroxide, LiOH is used to remove carbon dioxide from the exhaled air in the cabin of a spaceship. [Relative atomic mass: H = 1, Li = 7, C = 12, O = 16]
2LiOH + CO2 → Li2CO3 + H2O
An outer space mission is carried out for a period of 18 days involving five people on board. If each person is expected to exhale on the average of 42 g of carbon dioxide per hour and each absorption tube can contain 750 g of LiOH, calculate the number of absorption tubes that should be loaded into the spaceship.
Check Answers https://bit.ly/ 32OHQGV
77
CHAPTER
4
The Periodic Table of Elements
Keywords
Periodic Table of Elements Groups Periods Noble gases Alkali metals Halogens Metalloids Transition elements
What will you learn? 4.1 The Development of the Periodic Table of Elements 4.2 The Arrangement in the Periodic Table of Elements 4.3 Elements in Group 18 4.4 Elements in Group 1 4.5 Elements in Group 17 4.6 Elements in Period 3 4.7 Transition Elements 78
Bulletin Is the chlorine content in the swimming pool harmful to the health of users? The answer is no. You need not worry because water treated with chlorine only kills bacteria or dangerous organisms. Users of the swimming pool do not get infected with contagious diseases. In fact, you need not worry about the effects of accidentally swallowing 100 cm3 of chlorinated water each day during your swimming activity. The strong smell in the swimming pool results from the reaction of urea (from sweat) with chlorine. This reaction produces a substance called trichloramine (nitrogen trichloride). This substance may be harmful to health. Therefore, before entering the swimming pool, it is advisable for you to clean yourself first so that the urea due to sweat on the surface of the skin is removed. Did you know that the characteristics of chlorine element can be obtained by studying the position of chlorine in the Periodic Table of Elements? This is because all the elements in the Periodic Table of Elements are arranged systematically, therefore we are able to know the characteristics of chlorine element.
Who invented the Periodic Table of Elements? How are elements classified in the Periodic Table of Elements? What are the special characteristics of transition elements?
79
THEME 2
Fundamentals of Chemistry
4.1
The Development of the Periodic Table of Elements g Learnin tandard S
The Periodic Table of Elements classifies known chemical elements in a table according to certain characteristics. Many theories were explained by scientists before the modern Periodic Table of Elements was produced. What were those theories? Do you know the scientists involved in the development of the Periodic Table of Elements?
Antoine Lavoisier (1743 – 1794)
At the end of the lesson, pupils are able to: 4.1.1 Describe the historical development of the Periodic Table of Elements 4.1.2 Deduce the basic principle of arrangement of elements in the Periodic Table of Elements
Lavoisier classified elements according to certain groups such as gases, non-metals, metals and metal oxides. However, his classification was less accurate because he also classified light, heat and a few compounds into the groups as elements. Besides, there were several elements that were classified into the same group but showed different chemical properties.
Dobereiner found that the atomic mass of strontium atom, Sr was similar to the average mass of calcium atom, Ca and barium atom, Ba. These elements had the same chemical properties. A similar condition occurred with chlorine, Cl, bromine, Br and iodine, I. The group consisting of these three elements was named triad. Dobereiner’s classification was limited to several elements only. However, his classification showed the relationship between the chemical properties of elements and atomic mass.
John Newlands (1837 – 1898) 80
Johann W. Dobereiner (1780 – 1849)
Newlands arranged elements according to their increasing atomic masses. He arranged seven elements in a row because he found that the chemical and physical properties of the first element recurred at every eighth element. He named the arrangement as the Law of Octaves. The Law of Octaves had only been conformed by the first 17 elements. However, the recurrence in properties of the eighth element showed the presence of periodic pattern in the properties of elements.
The Periodic Table of Elements
Lothar Meyer (1830 – 1895)
Meyer plotted the graph of atomic volume against atomic mass of elements. He found that the elements at equivalent positions on the curve of the graph had similar chemical properties. For example, referring to the alkali metals such as lithium, sodium, potassium and rubidium that were located at the peaks of the curve. Meyer also proved the recurrence in properties of elements similar to Newlands.
Mendeleev arranged the elements according to their increasing atomic masses. Only elements with similar chemical properties were arranged in the same vertical columns. He had left several empty spaces in his periodic table to be filled by elements, yet to be discovered. He was successful in predicting the properties of undiscovered elements based on the properties of elements located above and below an element in the table.
Henry Moseley (1887 – 1915)
CHAPTER 4
Dmitri Mendeleev (1834 – 1907)
Moseley studied the frequencies of X-ray released by various elements and eventually found a relationship between the X-ray spectrum and proton numbers. He proposed that each element ought to have its own proton number. Then, he arranged the elements in the Periodic Table of Elements according to their increasing proton numbers. Moseley also left empty spaces in his periodic table like Mendeleev and successfully predicted four elements, namely technetium, Tc, promethium, Pm, hafnium, Hf and rhenium, Re which were discovered later.
Brain Teaser Basic Principle of Arrangement of Elements in the Periodic Table of Elements Elements in the Periodic Table of Elements are arranged in ascending order of proton numbers, ranging from 1 to 118. Elements with similar chemical properties are placed in the same vertical columns. Several new elements which were discovered such as nihonium, Nh, moscovium, Mc, tennessine, Ts and oganesson, Og were added into Period 7 of the Periodic Table of Elements.
Figure 4.1 The new elements
Brain Teaser From the historical development of the Periodic Table of Elements learned, predict the basic principle in the arrangement of the elements.
New elements discovered are named after the location or the name of the scientist.
81
THEME 2
Fundamentals of Chemistry
Activity 4.1 Discussing the importance of classifying the elements 1. Carry out the Think-Pair-Share activity. 2. Scan the QR code on the right on the development of the Periodic Table of Elements and think of the importance of classifying the elements. 3. Discuss with your partner. 4. Share your outcomes in front of the class.
TestYourself
Century
21st Skills
CT
Development of the Periodic Table of Elements http://bit.ly/ 35Sgp0A
4.1
1. Name the scientists that made the following discoveries: (a) Classified elements into four groups according to their chemical properties, that is gases, non-metals, metals and metal oxides (b) Proposed the Law of Octaves (c) Constructed the triad groups consisting of three elements with similar chemical properties 2. In the historical development of the Periodic Table of Elements, Moseley arranged the elements in ascending order of proton numbers. However, before the modern Periodic Table of Elements was used, scientists made their own discoveries. Compare how Dobereiner and Newlands arranged the elements in the Periodic Table of Elements before Moseley.
4.2
The Arrangement in the Periodic Table of Elements
The modern Periodic Table of Elements is a form of systematic classification of elements in ascending order of proton numbers from left to right and from top to bottom. The arrangement of elements is discussed from the aspects of groups, periods, proton number and electron arrangement. The vertical columns in the Periodic Table of Elements are called Groups. There are 18 groups in the Periodic Table of Elements. The number of valence electrons will determine the position of the group of an element. Figure 4.2 shows the position of the group of an element based on the number of valence electrons.
82
g Learnin tandard S At the end of the lesson, pupils are able to: 4.2.1 Describe briefly the modern Periodic Table of Elements 4.2.2 Generalise the relationship between the proton number and the position of elements in the Periodic Table of Elements
Literacy Tips You have learned the positions of metals, non-metals and noble gases in the Periodic Table of Elements in Form 1.
The Periodic Table of Elements
with one or two valence electrons
Group = Number of valence electrons
with three until eight valence electrons
Group = Number of valence electrons + 10
CHAPTER 4
Element
Figure 4.2 Position of the group of an element
The horizontal rows in the Periodic Table of Elements are called Periods. There are seven periods in the Periodic Table of Elements. The number of shells filled with electrons will determine the position of an element in a period. Table 4.1 explains the relationship between the proton number and the position of an element in the Periodic Table of Elements based on the aspects of groups and periods.
Table 4.1 Relationship between the proton number and the position of elements in the Periodic Table of Elements. Group
Number of shells filled with electrons
Period
1
1
2
2
2.8.8.2
2
2
4
4
13
2.8.3
3
3 + 10 = 13
3
3
14
2.8.4
4
4 + 10 = 14
3
3
Nitrogen, N
7
2.5
5
5 + 10 = 15
2
2
Oxygen, O
8
2.6
6
6 + 10 = 16
2
2
Bromine, Br
35
2.8.18.7
7
7 + 10 = 17
4
4
Krypton, Kr
36
2.8.18.8
8
8 + 10 = 18
4
4
Element
Proton number
Electron arrangement
Valence electron
Lithium, Li
3
2.1
Calcium, Ca
20
Aluminium, Al Silicon, Si
Activity 4.2
Century Predicting the group and period of an element based on its 21st Skills electron arrangement 1. Carry out the Round Table activity. 2. Choose a representative to speak out the proton number of an element. 3. Group members take turns to note down the electron arrangement, group and period of that element on a piece of paper. 4. Discuss the correct answer. 5. Pin up your outcomes on the class notice board as a reference for other groups.
CT
83
THEME 2
Fundamentals of Chemistry
TestYourself
4.2
1. Write the symbols for magnesium, copper and fluorine. 2. State the electron arrangement and group, for each of the following elements. Refer the Data Table of Elements on page 276 to get the proton number of each element. (a) Potassium, K (c) Chlorine, Cl (b) Carbon, C (d) Argon, Ar 3. Draw the electron arrangements of lithium, Li and carbon, C.
4.3
g Learnin tandard S
Elements in Group 18
Group 18 consists of elements of helium, He, neon, Ne, argon, Ar, krypton, Kr, xenon, Xe, radon, Rn and oganesson, Og. Elements in Group 18 are known as noble gases or inert gases. Activity 4.3 shows the relationship between the inert nature and the stability of the electron arrangement in an element. He Ne Ar Kr
At the end of the lesson, pupils are able to: 4.3.1 Relate the inert nature of Group 18 elements to its stability 4.3.2 Generalise the changes in physical properties of elements when going down Group 18 4.3.3 Describe briefly the uses of Group 18 elements in daily life
Xe Rn
More information on inert properties of neon
Og
http://bit.ly/33zc7cm
Figure 4.3 Positions of Group 18 elements in the Periodic Table of Elements
Activity 4.3 Relating the inert nature with the stability of duplet and octet electron arrangements of Group 18 elements 1. Carry out this activity in groups. Octet
Duplet He
Helium
84
Ne
Neon
Figure 4.4
Ar
Argon
The Periodic Table of Elements
CHAPTER 4
2. Based on Figure 4.4, discuss the relationship between the inert nature of Group 18 elements and the stability of electron arrangement in an element. 3. Present your findings in front of the class. Noble gases are chemically unreactive due to valence shells that are fully filled with electrons. Noble gases have achieved a stable duplet or octet electron arrangement, causing the atoms of noble gases to not donate, accept nor share electrons with the atoms of other elements. The atoms of noble gases exist as monoatoms.
Chemistry Noble gas is also known as inert gas.
Changes in Physical Properties of Elements When Going Down Group 18
Going down Group 18, the size of atomic radius increases due to the increase in the number of electrons and electron filled shells. Table 4.2 Physical properties of Group 18 elements Element
Atomic radius (nm) Melting point (°C) Boiling point (°C)
Density (g cm–3)
Helium, He
0.050
–270
–269
0.00017
Neon, Ne
0.070
–248
–246
0.00080
Argon, Ar
0.094
–189
–186
0.00170
Krypton, Kr
0.109
–156
–152
0.00350
Xenon, Xe
0.130
–122
–108
0.00550
Radon, Rn
–
–71
–62
Going down the group, the melting point and boiling point of the elements increase. Increase in the atomic size of elements will increase the attraction force between the atoms. Therefore, the attraction force becomes stronger and more heat energy is required to overcome this force.
– Brain Teaser
Brain Teaser Try to relate the increase in the density of the elements with the atomic mass and atomic size of each element, when going down the group.
Activity 4.4 Century Constructing a model to compare the physical properties and changes in the 21st Skills physical properties of Group 18 elements 1. Carry out the Three Stray One Stay activity. 2. Construct a 2D or 3D model to compare the physical properties of at least two elements in Group 18. 3. Prepare an exhibition corner in class and display the models from each group. 4. Choose a representative to give explanations on the comparison and changes in the physical properties of the selected Group 18 elements. The rest of the members will move around to seek information from other groups on their selected Group 18 elements.
85
THEME 2
Fundamentals of Chemistry
Uses of Group 18 Elements in Daily Life You have identified the list of elements found in Group 18 and studied the changes in the physical properties of the elements as you go down the group. Did you know that Group 18 elements have many uses in our daily life?
Helium
• Used to fill weather balloons • Used in the oxygen tanks of divers
Neon
• Used in advertising board lights
Argon
• Used to fill in electric bulbs • Used to provide an inert atmosphere for welding in high temperature
Krypton
• Used in flashlight of cameras • Used in lasers for eye retina treatment
Xenon
• Used in lighthouse lamps • Used for anesthesia
Radon
• Used to treat cancer
Figure 4.5 Uses of Group 18 elements
Activity 4.5 Summarising the uses of Group 18 elements in daily life 1. Carry out this activity in groups. 2. Watch a video clip on the uses of Group 18 elements in our daily lives by searching the Internet or visiting the link given. 3. Based on the video, discuss with your group members and summarise the uses of Group 18 elements in graphic form. 4. Present your group work in front of the class. 86
Group 18 elements http://bit.ly/2Be9EIw
The Periodic Table of Elements
TestYourself 1. 2. 3. 4.
CHAPTER 4
4.3
State the electron arrangement of helium, He. What is the type of electron arrangement of argon atom, Ar? Explain why neon, Ne does not react with other elements. Compare the boiling point of helium and argon. Explain.
4.4
Elements in Group 1
Group 1 is made up of lithium, Li, sodium, Na, potassium, K, rubidium, Rb, caesium, Cs and francium, Fr. Group 1 elements are also known as alkali metals. Li Na K Rb Cs Fr
Figure 4.6 Positions of Group 1 elements in the Periodic Table of Elements
g Learnin tandard S At the end of the lesson, pupils are able to: 4.4.1 Generalise the changes in physical properties of elements when going down Group 1 4.4.2 Investigate through experiment the chemical properties of Group 1 elements with: • Water • Oxygen gas • Chlorine 4.4.3 Generalise the changes in the reactivity of elements when going down Group 1 4.4.4 Reason out the physical and chemical properties of the other elements in Group 1
Photograph 4.1 shows the uses of several Group 1 elements. What other uses of these elements that you know?
Lithium, Li is used in battery production.
Sodium, Na is used in sodium vapour lamps.
Potassium, K is used in fertilisers.
Photograph 4.1 Uses of Group 1 elements
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Fundamentals of Chemistry
Changes in Physical Properties of Elements When Going Down Group 1 Going down Group 1, the atomic radius of elements increases as shown in Table 4.3. Table 4.3 Physical properties of several Group 1 elements Element
Atomic radius (nm)
Melting point (°C)
Boiling point (°C)
Lithium, Li
0.133
186
1342
Sodium, Na
0.186
98
880
Potassium, K
0.203
64
760
Group 1 elements have low melting point and boiling point if compared to other metals like iron that has a melting point of 1 540 °C and boiling point of 2 760 °C. Why do the melting point and boiling point of elements decrease when going down the group? The increase in the atomic size down the group will weaken the attraction force between the atoms. Therefore, less heat energy is required to overcome the attraction forces between the metal atoms. Group 1 elements are soft metals, with low density and float on the surface of water. These alkali metals also have a grey shiny surface at room temperature.
Chemical Properties of Group 1 Elements
Group 1 elements have one electron in the valence shell. In a chemical reaction, these atoms donate one electron and form an ion with the +1 charge. M → M+ + e – What will happen when Group 1 elements react with water, oxygen gas or chlorine gas?
Experiment
4.1
Aim: To study the chemical properties of Group 1 elements. Problem statement: What are the chemical properties of Group 1 elements, when they react with water, oxygen gas and chlorine gas? Materials: Lithium, sodium, potassium, distilled water, filter paper, red litmus paper, oxygen and chlorine gas Apparatus: Forceps, white tile, basin, knife, combustion spoon, gas jar with lid, 10 cm3 measuring cylinder and Bunsen burner Reaction of Group 1 elements with water (Demonstration by the teacher) Hypothesis: Going down the group, the reactivity of alkali metals with water will increase. Variables: Safety Precaution (a) Manipulated : Type of alkali metal Be careful when putting the (b) Responding : Reactivity of alkali metal with water alkali metal into the water. (c) Fixed : Size of alkali metal Only small quantities should be used.
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The Periodic Table of Elements
CHAPTER 4
Procedure: 1. Cut lithium into small pieces using a knife and forceps. Dry a piece of the metal on a filter paper. 2. Put the piece of lithium slowly into a basin filled with water as shown in Figure 4.7. Basin
Lithium, Li Distilled water
Figure 4.7
3. When the reaction is complete, test the solution with a red litmus paper. 4. Record your observations in Table 4.4. 5. Repeat steps 1 to 4 with sodium and potassium. Reaction of Group 1 elements with oxygen gas
Make a hypothesis and state all the variables for part B.
Procedure: 1. Cut lithium into small pieces using a knife and forceps. Dry a piece of the metal on a filter paper. 2. Put the piece of lithium onto a combustion spoon. 3. Heat until it starts to burn and immediately insert the spoon into a gas jar filled with oxygen gas as shown in Figure 4.8. 4. When the reaction is complete, add 10 cm3 of water into the gas jar and shake. 5. Test the solution using a red litmus paper. 6. Record your observations in Table 4.4. 7. Repeat steps 1 to 6 with sodium and potassium.
Combustion spoon Gas jar
Oxygen, O2
Lithium, Li
Figure 4.8
Reaction of Group 1 elements with chlorine gas
Make a hypothesis and state all the variables for part C.
Procedure: 1. Cut lithium into small pieces using a knife and forceps. Dry a piece of the metal on a filter paper. 2. Put the piece of lithium onto a combustion spoon. 3. Heat until it starts to burn and immediately insert the spoon into a gas jar filled with chlorine gas as shown in Figure 4.8. 4. Record your observations in Table 4.4. 5. Repeat steps 1 to 4 with sodium and potassium.
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THEME 2
Fundamentals of Chemistry
Results:
Table 4.4 Metal
Observation With water
With oxygen gas
With chlorine gas
Lithium Sodium Potassium
Conclusion: Is the hypothesis acceptable? What is the conclusion of this experiment? Discussion: 1. Write the chemical equation for the reaction of lithium, sodium and potassium with: (a) Water (b) Oxygen gas (c) Chlorine gas 2. Arrange the reactivity of alkali metals lithium, sodium and potassium with water, oxygen gas and chlorine gas in ascending order. Prepare a complete report after carrying out this experiment.
Reaction of Group 1 Elements with Water, Oxygen Gas and Chlorine Gas Lithium, sodium and potassium have the same chemical properties but different reactivity.
When alkali metals react with water, alkaline hydroxide solution and hydrogen gas is formed. For example, reaction of lithium with water will produce lithium hydroxide and hydrogen gas.
EduwebTV: Alkali metal http://bit.ly/ 31h4LZN
2Li(s) + 2H2O(l) → 2LiOH(aq) + H2(g) When alkali metals burn in oxygen gas, a white solid that is metal oxide will be formed. For example, reaction of lithium with oxygen gas will produce lithium oxide. 4Li(s) + O2(g) → 2Li2O(s) Solid metal oxide will form an alkaline solution when dissolved in water. For example, reaction of lithium oxide with water will produce lithium hydroxide. Li2O(s) + H2O(l) → 2LiOH(aq) When alkali metals burn in chlorine gas, a white solid, that is metal chloride will be formed. For example, reaction of sodium with chlorine gas will produce sodium chloride. 2Na(s) + Cl2(g) → 2NaCl(s) 90
The Periodic Table of Elements
CHAPTER 4
Changes in Reactivity of Elements Going Down Group 1
Experiment 4.1 shows the reactivity of elements increases when going down Group 1. Why do the changes occur? The reactivity of alkali metals in Group 1 is due to the tendency of an atom to donate its valence electrons. The number of shells filled with electrons increases when going down Group 1. This causes the increase in atomic size. The position of valence electrons is placed further away from the nucleus of an atom. When the nuclear attraction force towards the valence electrons weakens, the electrons are more easily donated.
Li
Lithium, Li
Na
Reactivity increases
Sodium, Na
K
Potassium, K
Figure 4.9 Reactivity of elements increases when going down Group 1
Physical and Chemical Properties of Other Elements in Group 1 You have studied the properties of lithium, sodium and potassium. How about the properties of other atoms such as rubidium, caesium and francium?
Like other alkali metals, rubidium, caesium and francium elements are soft metals with shiny surfaces and have low melting point and boiling point. Rubidium and caesium are metals that are very reactive and burn easily. Rubidium and caesium are usually combined with other elements. Thus, they are difficult to be isolated chemically. Francium element is an unstable radioactive isotope with a short half-life. All three elements are very reactive with water and oxygen.
Reactions of elements in Group 1 with water http://bit.ly/ 2MhYybu
Chemistry & Us Lithium batteries like those used in smartphones can explode when charged excessively because it accepts current rapidly. So, only original and good quality chargers should be used to charge your devices.
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THEME 2
Fundamentals of Chemistry
TestYourself
4.4
1. Give two examples of Group 1 elements. 2. Table 4.5 shows the electron arrangement for elements X, Y and Z. Table 4.5 Elements
Electron arrangement
X
2.1
Y
2.8.8.1
Z
2.8.18.8.1
(a) Give two differences in physical properties between elements X, Y and Z. (b) Element X reacts with oxygen when heated. Write the chemical equation for this reaction. (c) Arrange the reactivity of elements X, Y and Z in ascending order. Explain the difference in reactivity.
4.5
Elements in Group 17
Group 17 consists of fluorine, F, chlorine, Cl, bromine, Br, iodine, I, astatine, At and tennessine, Ts. Group 17 elements are known as halogens and exist as diatomic molecules. F Cl Br I At Ts
g Learnin tandard S At the end of the lesson, pupils are able to: 4.5.1 Generalise the changes in the physical properties of elements when going down Group 17 4.5.2 Summarise the chemical properties of Group 17 elements 4.5.3 Generalise the changes in the reactivity of elements when going down Group 17 4.5.4 Reason out the physical and chemical properties of other elements in Group 17
Figure 4.10 Position of Group 17 elements in the Periodic Table of Elements
Do you know the uses of Group 17 elements in our daily lives? Photograph 4.2 shows the examples of daily uses for elements chlorine, bromine and iodine. 92
The Periodic Table of Elements
CHAPTER 4
BLEACH
Chlorine in bleach
Bromine as a substance in fire extinguishers
Iodine as a disinfectant
Photograph 4.2 Uses of Group 17 elements
Changes in Physical Properties of Elements When Going Down Group 17
Going down Group 17, the physical state of halogens at room temperature changes from gas to liquid and finally to solid as shown in Table 4.6. Table 4.6 Physical properties of several Group 17 elements Element
Physical state
Melting point (°C) Boiling point (°C)
Density (g cm–3)
Chlorine, Cl
Gas
–101
–34
0.00300
Bromine, Br
Liquid
–7
59
3.11900
Iodine, I
Solid
114
184
4.95000
Going down the group, the increase in molecular size will cause the attraction force between molecules to become stronger. The melting point and boiling point of halogens will increase because more heat energy is required to overcome the intermolecular forces. The density of elements also increases with the increase in mass when going down the group. Group 17 elements have different colours. Chlorine gas is greenish yellow, liquid bromine is reddish brown while solid iodine is purplish black.
Chemical Properties of Group 17 Elements
Group 17 elements have seven electrons in the valence shell. In chemical reactions, these atoms receive one electron and form ions with –1 charge. X + e– → X– What will happen if Group 17 elements react with water, metals or alkalis?
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THEME 2
Fundamentals of Chemistry
Activity 4.6 Century Watching the reactions of Group 17 elements 21st Skills 1. Carry out the activity in groups. 2. Based on Internet search, watch the video clips on the reactions of Group 17 elements with water, metal and alkali.
Reaction of halogen with water
Reaction of halogen with iron, Fe
http://bit.ly/2ILzLdZ
http://bit.ly/2VJzpcR
Reaction of halogen with sodium hydroxide, NaOH http://bit.ly/32kGWS4
3. Based on the videos above, discuss the following questions: (a) Write chemical equations for the reactions of chlorine with water, iron and sodium hydroxide. (b) Arrange the reactivity of chlorine, bromine and iodine with iron in ascending order. (c) Halogens are reactive non-metal elements. Explain.
Reaction of Group 17 Elements with Water, Metal and Alkali Chlorine, bromine and iodine have the same chemical properties but different reactivity. When halogens react with water, an acidic solution is formed. For example, the reaction of chlorine with water will produce hydrochloric acid and hypochlorous acid. Cl2(g) + H2O(l)
HCl(aq) + HOCl(aq)
When halogens react with metal, a metal halide is formed. For example, the reaction of iron with bromine will produce iron(III) bromide.
Antoine Balard discovered hypochlorous acid when he added a dilute suspension of mercury(II) oxide into a flask filled with chlorine gas.
2Fe(s) + 3Br2(l) → 2FeBr3(s) When halogens react with an alkaline solution, metal halide, metal halate and water will be formed. For example, the reaction of iodine with sodium hydroxide will produce sodium iodide, sodium iodate(I) and water. I2(s) + 2NaOH(aq) → NaI(aq) + NaOI(aq) + H2O(l)
94
The Periodic Table of Elements
CHAPTER 4
Changes in Reactivity of Elements Down Group 17 Cl
Did you know that the reactivity of elements decreases when going down Group 17? Increasing atomic size will cause the valence shell to be further from the nucleus. This will cause the nuclear attraction force towards the electrons to become weaker. Thus, the difficulty in attracting electrons to fill the valence shell will increase.
Chlorine, Cl
Br
Bromine, Br
Reactivity decreases
I
Iodine, I
Figure 4.11 Reactivity of elements decreases when going down Group 17
Activity 4.7 Watching the safety precautions in handling Group 17 elements 1. Carry out the activity in groups. Example on safety 2. Watch the video clip on safety precautions on handling measures in Group 17 elements by surfing the Internet. handling halogens 3. Based on the video, carry out a forum titled ‘Safety http://bit.ly/33vshDO Precautions in Handling Group 17 Elements’. Discuss the following questions: (a) Group 17 elements are dangerous. Explain. (b) What are the safety precautions taken when handling Be careful when handling halogens like chlorine and bromine in the laboratory?
Group 17 elements because they are dangerous.
Physical and Chemical Properties of Other Elements in Group 17
Based on what you have learned, can you predict the physical and chemical properties of fluorine and astatine? Generally, all halogens are soluble in organic solvents and do not conduct heat nor electricity. Fluorine is a light-yellow poisonous gas. This gas which is very reactive and corrosive will cause a strong explosion when combined with hydrogen gas. Astatine is a rare radioactive element because it is not chemically stable. 95
THEME 2
Fundamentals of Chemistry
TestYourself
4.5
1. State three physical properties of elements down the Group 17. 2. Fluorine is very reactive compared to iodine. The reaction of fluorine with almost all other elements is very vigorous. Why is fluorine more reactive compared to iodine? 3. Give two examples of substances that contain Group 17 elements. 4. Astatine is not used in the science laboratory. Explain.
4.6
g Learnin tandard S
Elements in Period 3
Period 3 consists of elements sodium, Na, magnesium, Mg, aluminium, Al, silicon, Si, phosphorus, P, sulphur, S, chlorine, Cl, and argon, Ar.
Na
Mg
Al
Si
P
S
Cl
Ar
At the end of the lesson, pupils are able to: 4.6.1 Describe the trends in physical properties of elements across Period 3 4.6.2 Conduct an experiment to observe changes in the properties of the oxides of elements across Period 3 4.6.3 Describe briefly the uses of semi-metals
Figure 4.12 Position of Period 3 elements in the Periodic Table of Elements You have learned about the uses of sodium, Na, chlorine, Cl, and argon, Ar. The uses of several other elements in Period 3 are shown in Photograph 4.3.
Magnesium, Mg as a substance in a lighter
Aluminium, Al as a substance in cans
Phosphorus, P as a substance in fireworks
Sulphur, S as fungicides
Photograph 4.3 Uses of Period 3 elements
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The Periodic Table of Elements
CHAPTER 4
Changes in Physical Properties of Elements Across Period 3
Going across Period 3 from left to right, the atomic size will decrease because the atomic radius decreases. Table 4.7 Physical properties of Period 3 elements
Elements
Sodium, Magnesium, Aluminium, Silicon, Phosphorus, Sulphur, Chlorine, Argon, Na Mg Al Si P S Cl Ar
Atomic radius (nm) Electronegativity
0.186
0.160
0.143
0.118
0.110
0.104
0.100
0.094
0.9
1.2
1.5
1.8
2.1
2.5
3.0
–
Physical state
Solid
Gas
Atomic size decreases
Increase in the number of protons across Period 3 will increase the charge in the atom̕s nucleus. The electronegativity of elements will increase because the nuclear attraction force towards the electrons increases. The physical state of Period 3 elements will change from solid to gas from left to right across the period. The same goes to metal elements, semi-metal elements and non-metal elements. Sodium, Na, magnesium, Mg and aluminium, Al are metal elements, silicon, Si is a semi-metal element or metalloid while phosphorus, P, sulphur, S, chlorine, Cl and argon, Ar are non-metal elements.
Changes in Chemical Properties of Oxides of Elements Across Period 3
You have learned the properties of metal elements, semi-metal elements and non-metal elements in Period 3. How about the chemical properties of oxides of elements across Period 3?
Experiment
4.2
Aim: To study the change of chemical properties of oxides of elements across Period 3. Problem statement: How does the chemical properties of oxides of elements change across Period 3? Materials: Sodium oxide, Na2O, magnesium oxide, MgO, aluminium oxide, Al2O3 sulphur dioxide gas, SO2, silicon(IV) oxide, SiO2, distilled water, 2.0 mol dm–3 sodium hydroxide, NaOH and 2.0 mol dm–3 nitric acid, HNO3 Apparatus: Test tube, stopper, test tube holder, 10 cm3 measuring cylinder, pH meter, Bunsen burner, glass rod and spatula Reaction of oxides of Period 3 elements with water Hypothesis: Across Period 3, oxides of elements will change from basic to acidic. Variables: (a) Manipulated : Type of oxide of Period 3 elements (b) Responding : Change in oxide property (c) Fixed : Volume of water Procedure: 1. Pour 10 cm3 distilled water into a test tube containing half spatula of sodium oxide, Na2O and shake. 2. Measure the pH value of the solution in the test tube using a pH meter.
97
THEME 2
Fundamentals of Chemistry
3. Record your observations. 4. Repeat steps 1 to 3 using magnesium oxide, MgO, aluminium oxide, Al2O3 and sulphur dioxide, SO2. Results: Table 4.8
Oxide
Sodium oxide, Na2O
Magnesium oxide, MgO
Aluminium oxide, Al2O3
Sulphur dioxide, SO2
With water pH value
Reaction of oxides of Period 3 elements with sodium hydroxide and nitric acid Make hypothesis and state all the variables for part B. Procedure: 1 1. Fill spatula of magnesium oxide powder, MgO 4 into two different test tubes. Sodium Nitric acid, Magnesium 2. Add 5 cm3 of 2.0 mol dm–3 sodium hydroxide hydroxide, HNO3 oxide, MgO solution, NaOH into the first test tube. NaOH Heat Heat 3. Add 5 cm3 of 2.0 mol dm–3 nitric acid, HNO3 into Figure 4.13 the second test tube. 4. Heat both test tubes gently and stir using a glass rod as shown in Figure 4.13. 5. Observe the solubility of oxide in both solutions and record your observations. 6. Repeat steps 1 to 5 by using aluminium oxide Al2O3 and silicon(IV) oxide, SiO2. Results: Table 4.9
Oxide
Solubility With sodium hydroxide, NaOH
With nitric acid, HNO3
Magnesium oxide, MgO Aluminium oxide, Al2O3 Silicon(IV) oxide, SiO2
Conclusion: Is the hypothesis acceptable? What is the conclusion of this experiment? Discussion: 1. List the basic oxides, amphoteric oxides and acidic oxides. 2. Write the chemical equation for the reaction between basic oxide and nitric acid, HNO3. 3. Write the chemical equation for the reaction between amphoteric oxide and sodium hydroxide, NaOH. 4. List the elements that have basic oxides and acidic oxides across Period 3. Prepare a complete report after carrying out this experiment. 98
The Periodic Table of Elements
CHAPTER 4
Going across Period 3, the properties of oxide change from basic oxide to amphoteric oxide and then to acidic oxide. Na2O
MgO
Basic oxide
Al2O3
SiO2
Amphoteric oxide
P4O10
SO2
Cl2O7
Acidic oxide
Figure 4.14 Properties of oxides of elements across Period 3
Group 1 and 2 elements form metal oxides that are basic. When dissolved in water, both basic oxides will produce alkaline solutions. Basic oxides also react with acid to form salt and water. Aluminium forms metal oxide that is amphoteric in property. Aluminium oxide reacts with both acid and alkali to form salt and water. Elements from Groups 14, 15, 16 and 17 form non-metal oxides that are acidic in property. When dissolved in water, oxides of elements from these groups will produce acidic solutions. Acidic oxides will also react with alkali to form salt and water. You have learned the change of physical and chemical properties of elements across Period 3. Can you predict the change of properties for elements across Period 2?
Activity 4.8
Century Predicting the change of properties for elements in Period 2 21st Skills 1. Carry out the Round Table activity. 2. Based on the change of properties of elements across Period 3, discuss and predict the change of properties for elements across Period 2. 3. Take turns to record the information on a piece of paper. 4. Pin up the outcome of your group discussion on the class bulletin board as a reference to others.
Uses of Semi-Metallic Elements
Semi-metallic elements or metalloid have both the properties of metals and non-metals. These elements are weak conductors of electricity. However, metalloids are good electrical conductors at high temperatures. Based on that property, metalloids like silicon are used as semiconductors in the manufacture of electronic microchips. Photograph 4.4 shows the uses of electronic microchips in the production of computers and mobile phones.
Computer
Mobile phone
Photograph 4.4 Electronic microchips used in the making of computers and mobile phones
99
THEME 2
Fundamentals of Chemistry
Activity 4.9 Century Discussing the uses of semi-metals in the microelectronic industry 21st Skills 1. Carry out the Gallery Walk activity. 2. Collect information from reading materials or suitable websites on the uses of the following semi-metals in the microelectronic industry.
3. Discuss among your group members and prepare a presentation. 4. Display your group work in the class. Move around in groups to see the work of other groups. 5. Write comments on their work and paste them.
TestYourself
4.6
1. Why are the elements sodium, Na, magnesium, Mg, aluminium, Al, silicon, Si, phosphorus, P, sulphur, S, chlorine, Cl and argon, Ar in the same period? 2. Silicon exists as a solid at room temperature while non-metals like chlorine exists as a gas. Explain. 3. Figure 4.15 shows several elements in the Periodic Table of Elements.
C Al
Na
Figure 4.15
Ne Cl
(a) When aluminium and chlorine are compared, which element has a smaller atomic size? Explain. (b) Which element forms an amphoteric oxide? (c) Arrange all the elements in accending order of atomic size.
100
The Periodic Table of Elements
4.7
Transition Elements
g Learnin tandard S
Position of Transition Elements
Transition elements are placed in Group 3 and 12 in the Periodic Table of Elements. Examples of transition elements include chromium, Cr, manganese, Mn, iron, Fe, and copper, Cu. The yellow portion in Figure 4.16 shows the position of transition elements in the Periodic Table of Elements. 1
18
1
H 3
2
13
2 4
Li
Be
11
12
5
B
Na Mg 19
K 37
Rb 55
20
Ca 38
Sr 56
13
3 21
Sc 39
Y 57 – 71
4 22
Ti 40
Zr 72
5 23
6 24
V
Cr
41
42
Nb Mo 73
74
7 25
Mn 43
Tc 75
8 26
Fe 44
Ru 76
9 27
Co 45
Rh 77
10 28
Ni 46
Pd 78
11 29
Cu 47
Ag 79
12 30
Zn 48
Cd 80
Al 31
Ga 49
In 81
14 6
C 14
Si 32
Ge 50
Sn 82
15 7
N 15
P 33
As 51
Sb 83
16
17
8
9
O 16
S 34
F 17
Cl 35
Se
Br
52
53
Te 84
I
85
He 10
Ne 18
Ar 36
54
Xe 86
Ba
Lanthanides
Hf
Ta
W
Re
Os
Ir
Pt
Au
Hg
Tl
Pb
Bi
Po
At
Rn
87
88
89 – 103
104
105
106
107
108
109
110
111
112
113
114
115
116
117
118
66
67
Ra
Actinides
Rf
Db
Sg
Bh
Hs
Mt
Ds
Rg
57
58
59
60
61
62
63
64
65
92
93
94
96
97
La
Ce
Pr
89
90
91
Ac
Th
Pa
Nd Pm Sm U
Np
Eu 95
Gd Tb
Pu Am Cm Bk
Cn Nh
Dy
Ho
98
99
Cf
Es
Fl
Mc
Lv
Ts
68
69
70
71
100
101
102
Er Tm Yb Fm Md No
At the end of the lesson, pupils are able to: 4.7.1 Determine the position of transition elements in the Periodic Table of Elements 4.7.2 Explain the special characteristics of a few transition elements with examples 4.7.3 List the uses of transition elements in industry
Kr
Cs Fr
CHAPTER 4
Position of transition elements
Og
http://bit.ly/ 2B8OlYG
Lu 103
Lr
Figure 4.16 Position of transition elements in the Periodic Table of Elements
All transition elements are metals with the following properties:
Chemistry Scandium and zinc are not considered as transition elements because both do not show characteristics of transition elements.
• Solids with shiny surfaces • Very hard compared to metals in Groups 1 and 2 • Have high densities • Have high melting points and boiling points
Special Characteristics for Several Transition Elements in the Periodic Table of Elements
Transition elements are metals with high melting point and boiling point, hard, shiny surfaces, malleable and ductile. Transition elements also have special characteristics unknown to other metals. What are the special characteristics possessed by transition elements? 1
Transition elements function as a catalyst to increase the rate of reaction without undergoing chemical change at the end of the reaction. For example, iron filings are used as a catalyst in the Haber Process. N2(g) + 3H2(g)
200 − 300 atm 450 − 550 oC, Fe
2NH3(g) 101
THEME 2
Fundamentals of Chemistry
Table 4.10
2 Transition elements form coloured ions or compounds.
Transition element ions
Colour of solution
Chromium(III) ion, Cr3+(aq) Dichromate(VI) ion, Cr2O72–(aq)
Green Orange
Manganese(II) ion, Mn2+(aq) Manganate(VII) ion, MnO4–(aq)
Pink Purple
Iron(II) ion, Fe2+(aq) Iron(III) ion, Fe3+(aq)
Green Brown
Copper(II) ion, Cu2+(aq)
Blue
EduwebTV: Transition elements
Photograph 4.5 Coloured compounds of transition elements
http://bit.ly/2BcvM5Z
Activity 4.10 Observing the colour of the transition element compounds 1. Observe the colour of the following transition element compounds: (f) Iron(II) sulphate, FeSO4 (a) Chromium(IIl) chloride, CrCl3 (b) Potassium dichromate(Vl), K2Cr2O7 (g) Iron(III) chloride, FeCl3 (c) Manganese(II) chloride, MnCl2 (h) Copper(I) oxide, Cu2O (i) Copper(II) oxide, CuO (d) Manganese(IV) oxide, MnO2 (e) Potassium manganate(VII), KMnO4 2. Present your findings in the form of a suitable mind map to share with your friends. Table 4.11
3 Transition elements have more than one oxidation number.
102
Transition element
Oxidation number
Compound
Chromium, Cr
+3 +6
Chromium(III) chloride, CrCl3 Potassium dichromate(VI), K2Cr2O7
Mangan, Mn
+2 +4 +7
Manganese(II) chloride, MnCl2 Manganese(IV) oxide, MnO2 Potassium manganate(VII), KMnO4
Iron, Fe
+2 +3
Iron(II) sulphate, FeSO4 Iron(III) chloride, FeCl3
Copper, Cu
+1 +2
Copper(I) oxide, Cu2O Copper(II) oxide, CuO
The Periodic Table of Elements
CHAPTER 4
Table 4.12 Ion of transition element
4 Transition elements can form complex ions.
Formula
Tetraaminecopper(II) ion
[Cu(NH3)4]2+
Hexacyanoferrate(III) ion
[Fe(CN)6]3–
Hexacyanoferrate(II) ion
[Fe(CN)6]4–
Hexaaquaferrate(II) ion
[Fe(H2O)6]2+
Activity 4.11 Conducting the learning activity based on the problems related to the special characteristics of transition elements 1. Carry out the activity in groups. 2. Read and understand the following passage:
Century
21st Skills
CT
Scientists found a special characteristic of certain transition elements that can ̒remember̕ their form. For example, an alloy called Nitinol is a mixture of nickel and titanium that can return to its original form after the alloy is bent. This alloy is used in making spectacle frames and treating broken bones. 3. Gather information on problems that can be solved using transition elements and their special characteristics. 4. Justify the use of transition elements and relate them to their special characteristics. 5. Prepare a multimedia presentation based on your findings. 6. Present your group work in front of the class.
Uses of Transition Elements in Industry
One of the special characteristics of transition elements is their use as a catalyst in industries. Do you know which transition elements are involved in that function? Figure 4.17 shows the examples of transition elements that function as a catalyst in industries.
Iron, Fe is used in the Haber Process to produce ammonia, NH3.
Platinum, Pt is used in the Ostwald Process to produce nitric acid, HNO3.
Transition Elements Vanadium(V) oxide, V2O5 is used in the Contact Process to produce sulphuric acid, H2SO4.
Nickel, Ni or platinum, Pt is used in the hydrogenation process on vegetable oils to produce margarine.
Figure 4.17 Transition elements as catalysts in industries
103
THEME 2
Fundamentals of Chemistry
Apart from catalyst, other uses of transition elements are shown in Photograph 4.6.
Manganese is used to make stained glass windows
Iron is used to build bridges
Titanium is used to make paints
Photograph 4.6 Uses of transition elements
Activity 4.12 Making a scrap book/ brochure/ pamphlet/ poster to show the uses of several transition elements in various industries 1. Carry out this activity in groups. 2. Gather information from various reading materials and search relevant websites for the use of several transition elements in various industries. 3. Discuss with your group members and present your findings in a scrap book/ brochure/ pamphlet/ poster. 4. Exhibit the scrap book/ brochure/ pamphlet/ poster in the laboratory or class.
TestYourself
4.7
1. Table 4.13 shows three transition elements that exist in gems. Table 4.13 Gem
Transition element
Ruby
Chromium
Sapphire
Iron, titanium
Amethyst
Manganese
(a) What are the special characteristics of the transition elements shown in Table 4.13? (b) Apart from the characteristics given in 1(a), what are the other characteristics found in transition elements? 2. Give examples of several transition elements used in industry. 104
Non-metal
• Atomic size • Melting point and boiling point • Density
Metal
• Atomic size • Melting point and boiling point
physical properties
chemical properties
Position of element
Periodic Table of Elements
Non-metal
http://bit.ly/ 2Jg9oNz
Quick
Henry Moseley
Dmitri Mendeleev
Lothar Meyer
John Newlands
Johann W. Dobereiner
Antoine Lavoiser
• Reactivity • Distance between valence electrons and nucleus • Nuclear attraction force • Ease in attracting electrons
chemical properties
changes when going down the group
Groups
vertical columns is called
discovered by
• Reactivity • Distance between valence electrons and nucleus • Nuclear attraction force • Ease in donating electrons
Metal
Basic oxide → amphoteric → acidic oxide
change across period
Period
horizontal rows is called
arranged
Metal → semi-metal → non-metal
physical properties
Follows the increase in proton number
Concept
The Periodic Table of Elements
105
CHAPTER 4
THEME 2
Fundamentals of Chemistry
Self Reflection 1. What new knowledge have you learned in The Periodic Table of Elements? 2. Which is the most interesting subtopic in The Periodic Table of Elements? Why? 3. Give a few examples of elements in The Periodic Table of Elements that you use in your daily life. 4. Rate your performance in The Periodic Table of Elements on http://bit.ly/ a scale of 1 to 10; 1 being the lowest and 10 the highest. Why 31kIQk8 would you rate yourself at that level? 5. What else would you like to know about The Periodic Table of Elements?
4
Achievement
1. How did Moseley arrange the elements in the Periodic Table of Elements? 2. Figure 1 shows the chemical symbol for element X. (a) Which group does element X belong to in the Periodic Table of Elements? (b) Which period does element X belong to in the Periodic Table of Elements?
35
X 17 Figure 1
3. A restaurant owner uses colourful electric lights to attract his customers. What is the substance suitable for making such lights? 4. State the physical and chemical properties of the element with electron arrangement 2.8.8.1. 5. State the element in Period 3 that forms an amphoteric oxide. 6. Figure 2 shows several elements in the Periodic Table of Elements that are represented by alphabets X and Y. X Y
Figure 2
(a) Write the electron arrangement for atom X and atom Y. (b) Explain two differences in chemical properties between element X and element Y. (c) Why does the reactivity of elements in the same group as X, increase when going down the group, but the reactivity of elements in the same group as Y decreases?
106
The Periodic Table of Elements
CHAPTER 4
7. Chlorine, Cl2 reacts with sodium, Na to form a compound. Write the chemical equation for the reaction. 8. Figure 3 shows the electron arrangement for the element G. (a) Which group does element G belong to in the Periodic Table of Elements? (b) Which period does element G belong to in the Periodic Table of Elements? (c) State a physical property of element G. 9.
G
Figure 3
Na
Mg Al Si P S Cl Ar (a) State which element is a metal, metalloid or non-metal from the list of elements given above. (b) Explain the change in atomic radius across the Period from left to right. (c) Which element is a noble gas? (d) Write the chemical equation when a metal reacts with water.
10. Give the colour of the following transition element ions: (a) Iron(II) ion (b) Iron(III) ion
Enrichmen Corner 1. Photograph 1 shows a few microchips. Silicon element is widely used in the manufacture industries of microchips.
Photograph 1 Microchips
What is the property of silicon that enables this element to be used in the manufacture of microchips instead of a metal like lithium? Explain.
Check Answers https://bit.ly/ 2pHzJ04
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CHAPTER
5
Chemical Bond
Keywords Chemical bonds Ionic bonds Electrostatic attraction force Covalent bonds Hydrogen bonds Dative bonds Van der Waals attraction force
Teflon
What will you learn? 5.1 5.2 5.3 5.4 5.5 5.6 5.7 108
Basics of Compound Formation Ionic Bond Covalent Bond Hydrogen Bond Dative Bond Metallic Bond Properties of Ionic Compounds and Covalent Compounds
Bulletin Dato' Dr. Sheikh Muszaphar Shukor is the first Malaysian astronaut sent to the outer space. During his stay in the outer space, he needs to wear an astronaut suit. The astronaut suit is created specifically to protect the astronaut’s body from the space environment. Did you know that the astronaut suit is made up of five layers? The layers consist of an inner layer of cotton, followed by a layer of blue nylon, a layer of black nylon, Teflon layer and lastly white nylon on the outer side. All layers of nylon and Teflon are macromolecules made from covalent compounds by covalent bonds, a type of chemical bond that is very strong.
What is meant by chemical bond? Why is ethanol soluble in water?
How are dative bonds formed?
109
THEME 2
Fundamentals of Chemistry
5.1
Basics of Compound Formation
Sodium, Na is reactive towards chlorine gas, Cl2 when heated and forms a white solid. Did you know that this white solid is the table salt that you use in your daily life? However, no compound is formed when sodium, Na is heated with neon gas, Ne. Why?
g Learnin tandard S At the end of the lesson, pupils are able to: 5.1.1 Explain the basics of compound formation
Compounds are formed when two or more elements are combined. Do you know how elements are combined to produce compounds?
Activity 5.1 Watching a video on the formation of compounds Ionic bond 1. Carry out the activity in groups. http://bit.ly/2MfCmit 2. Watch video clips on the formation of compounds by electron transfer (ionic bonds) and by sharing of electrons (covalent bonds) from your Internet search. 3. Based on those videos, discuss the following: Covalent bond (a) Formation of compounds by transfer of electrons to http://bit.ly/2BeDXi1 achieve a stable octet or duplet electron arrangement (b) Formation of compounds by sharing of electrons to achieve a stable octet or duplet electron arrangement Brain Teaser 4. Present the findings of your discussion on a flip chart paper in front of the class. Noble gases exist as monoatomic gases and are not reactive chemically because they have achieved a stable duplet or octet electron arrangement. However, for atoms of other elements, stability of electron arrangement can be achieved by transferring or sharing of electrons. Chemical bonds are formed when electron transfer or electron sharing takes place. There are two types of chemical bonds; ionic bond and covalent bond. Chemical bonds only involve the valence electrons. There! I have one extra electron.
Na
I only have seven electrons in the valence shell. I need one more electron to achieve an octet electron arrangement.
Cl
Are you sure? Thank you.
Na
Cl
Figure 5.1 Formation of ionic bond
110
Brain Teaser Why are electrons in the inner shells not involved in chemical bonds?
Yes, I am very positive about it.
Na
Cl
Chemical Bond
I have only one electron.
I also have one electron.
H
TestYourself
H
Let us share our electrons.
CHAPTER 5
Wow! We have achieved duplet electron arrangement.
H
H
Figure 5.2 Formation of covalent bond
5.1
1. What is chemical bond? 2. State two types of chemical bonds. 3. Why noble gases do not form compounds? 4. Is the electron arrangement of sodium atom, Na stable? If not, explain how the electron arrangement can become stable.
5.2
Ionic Bond
Kana, I have nine eggs but this egg container can only hold eight eggs.
There is still space in my egg container because I have only seven eggs.
In that case, let me give you an egg.
Thank you, Siti. Now both the containers are fully filled.
Figure 5.3 Analogy for the formation of ionic bond
Situation in Figure 5.3 gives the analogy for the formation of ionic bond. Ionic bond is formed by the transfer of electrons between a metal atom and a non-metal atom.
g Learnin tandard S At the end of the lesson, pupils are able to: 5.2.1 Explain with examples the formation of ionic bond
111
THEME 2
Fundamentals of Chemistry
Formation of Ions
Metal atom donates valence electron to form a positively-charged ion or cation. Figure 5.4 shows the formation of sodium ion, Na+. To achieve a stable electron arrangement, sodium atom, Na needs to donate an electron. The process of donating an electron from the valence shell of sodium atom, Na is much easier compared to accepting seven electrons from another atom.
After donating its valence electron, sodium ion, Na+ achieves a stable octet electron arrangement. Sodium ion, Na+ has 11 protons and 10 electrons, thus the charge for a sodium ion, Na+ is +1. +
Donates one electron
Na
2.8.1 Sodium atom, Na
Na
2.8 Sodium ion, Na+
Half-equation for the formation of sodium ion, Na+: Na Na+ + e– Figure 5.4 Formation of sodium ion, Na+
Non-metal atom receives electron from a metal atom to form a negatively-charged ion or anion. Figure 5.5 shows the formation of fluoride ion, F–. To achieve a stable electron arrangement, fluorine atom, F will accept an electron. The process of accepting one electron to the valence shell of fluorine atom, F is easier compared to donating its seven valence electrons to another atom.
After receiving one valence electron, fluoride ion, F– achieves a stable octet electron arrangement. Fluoride ion, F– has 9 protons and 10 electrons, so the charge of fluoride ion, F– is –1.
– Accepts one electron
F
2.7 Fluorine atom, F Half-equation for the formation of fluoride ion, F–: F + e– F– Figure 5.5 Formation of fluoride ion, F–
112
F
2.8 Fluoride ion, F–
Chemical Bond
CHAPTER 5
Formation of Ionic Bond
An ionic compound is formed when ions of opposite charges attract one another to form an ionic bond. How do ions of opposite charges attract one another?
Chemistry & Us Sodium atom, Na will donate one electron to achieve a stable octet electron arrangement. Sodium ion, Na+ is formed.
Fluorine atom, F will accept one electron from sodium atom, Na to achieve a stable octet electron arrangement. Fluoride ion, F– is formed.
Besides calcium carbonate, CaCO3, sodium fluoride, NaF is also added to toothpaste to strengthen the teeth.
+ Na
2.8.1 Sodium atom, Na
+
F
2.7 Fluorine atom, F
Na
2.8 Sodium ion, Na+
– F
2.8 Fluoride ion, F–
Sodium ion, Na+ and fluoride ion, F– of opposite charges will attract one another by strong electrostatic attraction force. Electrostatic attraction forces are known as ionic bonds. A compound sodium fluoride, NaF is formed.
Figure 5.6 Formation of sodium fluoride, NaF
Activity 5.2 Century Discussing the formation of ionic bonds 21st Skills 1. Carry out the Gallery Walk activity. 2. Gather information from various reading resources and websites on the formation of ionic bonds for the following compounds:
CT
Sodium oxide, Na2O Sodium chloride, NaCl Magnesium oxide, MgO 3. Scan the AR code to see the formation of ionic compound of sodium chloride, NaCl. 4. Discuss the formation of ionic bonds with your group members and prepare a presentation. You need to write half-equations for the formation of ions in each compound. 5. Display your group work in class. Move around to see the outcome of other groups' discussion. 6. Write comments on their work on sticky notes and paste them.
113
THEME 2
Fundamentals of Chemistry
TestYourself
5.2
Kendiri Uji
1. Aluminium atom, Al has 13 protons while fluorine atom, F has 9 protons. (a) Write the formulae of ions formed from the two atoms respectively. (b) Write half-equations for the formation of ions in (a). (c) Draw the electron arrangement to show the transfer of electrons in the formation of ionic bonds in aluminium fluoride compound. 2. Muriate of Potash is a type of fertiliser that has a high content of potassium chloride compound. [Proton number: Cl = 17, K = 19] (a) Write the chemical formula for potassium chloride. (b) Describe the formation of ionic bonds in potassium chloride compound.
5.3
Covalent Bond
g Learnin tandard S
Did you know that diamond is one of the hardest substances in the world? The property of diamond is caused by the formation of covalent bonds between carbon atoms.
At the end of the lesson, pupils are able to: 5.3.1 Explain with examples the formation of covalent bond 5.3.2 Compare ionic bond and covalent bond
Covalent bonds are formed when non-metal atoms share their electrons to achieve a stable duplet or octet electron arrangement. There are three types of covalent bonds; single bond, double bond and triple bond.
Single Bond
A single bond is formed when two atoms share a pair of electrons. Chlorine atom, Cl needs one electron to achieve a stable octet electron arrangement.
Cl
+
2.8.7 Chlorine atom, Cl
Cl
2.8.7 Chlorine atom, Cl
Two chlorine atoms, Cl each contributes one electron to share a pair of electrons to form a single bond in a chlorine molecule, Cl2.
Cl
2.8.8 2.8.8 Chlorine molecule, Cl2
Figure 5.7 Formation of single bond in chlorine molecule, Cl2
114
Cl
Chemical Bond
The formation of covalent bond can be visualised using the Lewis structure. Lewis structure only shows the valence electrons of the atoms involved. A pair of electrons shared, is represented with a line between the two atoms.
Cl
+
Cl
Cl Cl
or
Cl Cl
Chemistry
Diamond consists of carbon atoms, C. Each carbon atom, C forms four covalent bonds with another four carbon atoms.
Figure 5.8 Lewis structure for the formation of chlorine molecule, Cl2
Double Bond
A double bond is formed when two atoms share two pairs of electrons. Oxygen atom, O needs two electrons to achieve a stable octet electron arrangement.
O 2.6 Oxygen atom, O
2.6 Oxygen atom, O
O
Two oxygen atoms, O each contributes two electrons to share two pairs of electrons to form a double bond in oxygen molecule, O2.
O
+
+
O
O or
O O
or
O
2.8 2.8 Oxygen molecule, O2
O O
Figure 5.9 Formation of double bond in oxygen molecule, O2
Triple Bond
A triple bond is formed when two atoms share three pairs of electrons. The nitrogen atom, N needs three electrons to achieve a stable octet electron arrangement.
N 2.5 Nitrogen atom, N
2.5 Nitrogen atom, N
N
Two nitrogen atoms, N each contributes three electrons to share three pairs of electrons to form a triple bond in nitrogen molecule, N2.
N
+
+
N
N
N N
N
2.8 2.8 Nitrogen molecule, N2
or
or
CHAPTER 5
N N
Figure 5.10 Formation of triple bond in nitrogen molecule, N2
115
THEME 2
Fundamentals of Chemistry
Activity 5.3 Century Visualising the formation of covalent bonds 21st Skills 1. Carry out the Three Stray One Stay activity. 2. Build a model to visualise the formation of covalent bonds in the following compounds.
Hydrogen, H2
Hydrogen chloride, HCl
CT
Oxygen, O2
Carbon dioxide, CO2 Nitrogen, N2 3. Prepare an exhibition corner in the class and display the models from each group. 4. Choose a representative to give an explanation on the formation of covalent bond in a chosen compound. The rest of the members will move around to seek information from other groups.
Comparison between Ionic Bond and Covalent Bond
Similarities and differences between ionic bond and covalent bond are shown in Figure 5.11.
Between metal atoms and non-metal atoms
Ionic bond
Forms a positivelycharged ion and a negativelycharged ion
Sharing of electrons
Involve valence electrons only
Transfer of electrons
Covalent bond
Atoms achieve stable octet or duplet electron arrangement
Figure 5.11 Comparison between ionic bond and covalent bond
116
Between non-metal atoms
Forms a molecules
Chemical Bond
TestYourself
CHAPTER 5
5.3
1. State three types of covalent bonds. 2. How are covalent bonds formed? 3. Draw the formation of covalent bonds in a water molecule, H2O. 4. Can carbon atoms, C share electrons with four hydrogen atoms, H to form a methane molecule? Explain. [Proton number: H = 1, C = 6] 5. State one similarity and two differences between ionic bond and covalent bond.
5.4
Hydrogen Bond
Have you ever thought why an iceberg weighing thousands of tonnes is able to float on the surface of the sea? This is because the density of ice is lower compared with water. Why is water denser than ice? To answer this question, the concept of hydrogen bonds needs to be understood. Hydrogen bonds are attraction forces between hydrogen atom, H that has bonded with an atom of high electronegativity, such as nitrogen, N, oxygen, O or fluorine, F with nitrogen, N, oxygen, O or fluorine, F in another molecule. For example, water molecule, H2O can form hydrogen bonds among water molecules, H2O.
Oxygen atom, O has high electronegativity. Water molecule, H2O consists of two hydrogen atoms, H and one oxygen atom, O. Hydrogen, H and oxygen, O are bonded by sharing electrons with one another. This type of bond is a covalent bond.
H
g Learnin tandard S At the end of the lesson, pupils should be able to: 5.4.1 Explain with examples the formation of a hydrogen bond 5.4.2 Explain the effect of the hydrogen bond on physical properties of substances
H O H O H
Attraction forces between hydrogen atom, H in water molecule with oxygen atom, O from another water molecule forms a hydrogen bond.
Figure 5.12 Formation of hydrogen bonds between water molecules, H2O
117
THEME 2
Fundamentals of Chemistry
Activity 5.4 Century Discussing the formation of hydrogen bonds in hydrogen CT 21st Skills fluoride, HF and ammonia, NH3 1. Carry out the Think-Pair-Share activity. 2. Based on Figure 5.12, think about how hydrogen bonds are formed in hydrogen fluoride, HF and ammonia, NH3. 3. Discuss with your partner. 4. Share your findings in front of the class.
Role of Hydrogen Bonds in Daily Life
Observe Figure 5.13. There are protein molecules that form hydrogen bonds among one another in the hair structure. Do you know why hair sticks together when wet? Protein molecules H O
H O
H O
H O
H O
Hair
H O
Hydrogen bonds
Figure 5.13 Hydrogen bonds between protein molecules in the hair structure
When hair is wet, protein molecules no longer form hydrogen bonds among themselves. Instead, protein molecules will form hydrogen bonds with water molecules, H2O. Water molecules, H2O will also form hydrogen bonds with other hair protein molecules. This causes hair to stick together. Have you ever come across the problem of turning the pages of a book where the pages stick together? To overcome this problem, you lick your finger before turning the pages. Why does a wet finger help to turn the pages of a book? Explanation on this is given in Figure 5.15.
Cellulose in paper consists of hydrogen atoms, H that are bonded to oxygen atoms, O.
H H O
O
H O
O H
H O
H H
H O
Hydrogen bonds H O
H H
O
H O
Brain Teaser Figure 5.14
Formation of hydrogen bond between protein molecule and water molecule
Brain Teaser Why does wavy hair look straight when wet?
Water molecules, H2O on the wet finger will form hydrogen bonds with the cellulose in paper. Because of that, paper will stick to the finger.
Paper cellulose
Figure 5.15 Hydrogen bonds formed between cellulose in paper and water molecule, H2O on the finger
118
Chemical Bond
CHAPTER 5
Effect of Hydrogen Bonds on The Physical Properties of Substances
Compounds in the form of liquids reach boiling point when the attraction forces between molecules are overcomed. In the covalent compound of ethanol, C2H5OH, there are hydrogen bonds formed between molecules, other than weak Van der Waals attraction forces. Strong hydrogen bonds are difficult to break. More heat energy is required to overcome the weak Van der Waals attraction forces, besides breaking the hydrogen bonds. As a result, the boiling point of ethanol, C2H5OH is high. On the other hand, molecules like chlorine, Cl2 which do not form hydrogen bonds have lower boiling point compared to ethanol. Ethanol, C2H5OH is also soluble in water. The solubility of ethanol, C2H5OH in water is due to the formation of hydrogen bonds between the ethanol molecule, C2H5OH and water molecule, H2O. Ethanol, C2H5OH consists of hydrogen atoms, H that form covalent bonds with oxygen atoms, O. So, the oxygen atom, O in ethanol molecule, C2H5OH can form a hydrogen bond with the hydrogen atom, H from water molecule, H2O.
H
Van der Waals attraction force H
H
H
C
C
O
H
H
H
H
O H
H
H
C
C
H
H
O
H
O
H
H
Ethanol molecule, C2H5OH
The hydrogen atom, H in ethanol molecule, C2H5OH can also form a hydrogen bond with the oxygen atom, O from water molecule, H2O.
Figure 5.16 Solubility of ethanol, C2H5OH in water, H2O
Activity 5.5
Century CT 21st Skills Discussing the solubility in water and boiling point of covalent compounds 1. Carry out the Round Table activity. 2. Gather information on the solubility in water and boiling point for hydrogen fluoride, HF and ammonia, NH3 from various reading materials and websites. 3. Compare the solubility and boiling point for these compounds with molecules that do not form hydrogen bonds. 4. Take turns to record the related information on a piece of paper. 5. Pin up your group work on the classroom bulletin board to share the information and references with other groups.
TestYourself
5.4
1. State the meaning of hydrogen bond. 2. Hydrogen fluoride, HF exists as liquid at room temperature. Explain this phenomenon based on the formation of hydrogen bonds. 3. Can hydrogen bonds form among hydrogen chloride molecules, HCl? Justify your answer. 4. Explain why paper sticks together when wet. 119
THEME 2
Fundamentals of Chemistry
5.5
Dative Bond
I do not have electrons.
H+
I have a pair of electrons to be shared.
NH3
Wow! I have achieved duplet electron arrangement.
I am still in octet electron arrangement.
H+
g Learnin tandard S At the end of the lesson, pupils should be able to: 5.5.1 Explain with examples the formation of dative bond
NH3
Figure 5.17 Formation of dative bond
Dative bond or coordinate bond is a type of covalent bond where the electron pair that is shared comes from one atom only. How does such sharing take place? Figure 5.18 shows the formation of dative bond in hydroxonium ion, H3O+. 2 Hydrogen ion, H+ does not have
4
any electron in the shell.
1
Oxygen atom, O achieves octet electron arrangement and hydrogen atom, H achieves stable duplet electron arrangement in water molecule, H2O.
H H O + H+
H H O H
In hydroxonium ion, H3O+, oxygen atom, O and all hydrogen atoms, H have achieved stable octet and duplet electron arrangements respectively.
+
3 The lone pair of electrons that are not involved in covalent bond in water molecule, H2O will be shared with hydrogen ion, H+ through the formation of dative bond.
Figure 5.18 Formation of dative bond in hydroxonium ion, H3O+
Activity 5.6 Discussing the formation of dative bond in ammonium ion, NH4+ 1. Carry out this activity in groups. 2. Based on the statement below, discuss the formation of dative bonds in ammonium ion, NH4+.
3. Present your discussion results in an attractive presentation in front of the class. 120
m
Am
on
ia
cid
When hydrogen chloride gas, HCl and ammonia gas, NH3 are mixed, white fumes of ammonium chloride, NH4Cl is formed as shown in Photograph 5.1.
r Hyd
l och
c ori
a
Photograph 5.1 Formation of white fumes of ammonium chloride, NH4Cl
Chemical Bond
TestYourself
CHAPTER 5
5.5
1. What is dative bond? 2. Explain the formation of ammonium ion through the formation of dative bond between hydrogen ion, H+ and nitrogen atom, N in ammonia, NH3. 3. Boron atom, B found in the compound boron trifluoride, BF3 has not achieved octet electron arrangement because it has only six valence electrons. Can boron atoms, B form dative bonds with nitrogen atoms, N in the compound ammonia, NH3? Explain your answer.
5.6
g Learnin tandard S
Metallic Bond
Did you know that exposed electrical wires can cause electric shock? Electrical wires made from metal can conduct electricity. Why can metals conduct electricity? Metal atoms are arranged closely packed and orderly in the solid state. Valence electrons of metal atoms can be donated easily and delocalised although in the solid state. Metal ions that are positively-charged are formed when valence electrons are delocalised. All delocalised valence electrons can move freely between the metal structure and form a sea of electrons. Electrostatic attraction force between the sea of electrons and the positively-charged metal ions form the metallic bond. Valence electron
+
+ +
+
+ +
+ +
+ + + +
+
+ + + +
+
+ + + +
+
+ + + +
+
+ + + +
+
+ + + +
+
+ + + +
+
+ +
Chemistry Delocalised electron means electron that moves freely and is not owned by any atom nor ion. A sea of electron is formed when the valence shells of metal atoms overlap, resulting in electron delocalisation.
Positivelycharged metal ions
+ +
+
At the end of the lesson, pupils should be able to: 5.6.1 Explain the formation of a metallic bond 5.6.2 Reason out the electrical conductivity of metal
+ +
Sea of electron
Figure 5.19 Formation of metallic bond
When electrons of metal atoms are delocalised in the sea of electrons, the metal can conduct electricity. Electrons that move freely in the metal structure carry the charges from the negative terminal to the positive terminal when electricity is supplied, as shown in Figure 5.20.
Negative terminal − Electron
+
+ +
+
+ +
+ +
+ +
+ +
+ +
+
Positive terminal
+ + + +
+
+
Metal
Figure 5.20 Electrical conductivity of metals
121
THEME 2
Fundamentals of Chemistry
Activity 5.7 Century CT Comparing and contrasting the formation of bonds 21st Skills 1. Carry out the Think-Pair-Share activity. 2. Using suitable mind maps, compare and contrast the formation of all bonds studied from the following aspects: (a) Sharing or transfer of electrons (b) Attraction forces formed (c) Examples of compounds or elements 3. Pin up your mind maps produced on the classroom bulletin board.
TestYourself
5.6
1. What is meant by a delocalised electron? 2. How a metallic bond is formed in metals? 3. Using aluminium, Al metal as an example, explain how metals can conduct electricity.
5.7
Properties of Ionic Compounds and Covalent Compounds earning tandard L Observe salt (sodium chloride, NaCl) and ice in Photograph 5.2. S Are both substances in the same physical state? Which substance will melt at room temperature?
Salt is an ionic compound
Ice is a covalent compound
At the end of the lesson, pupils should be able to: 5.7.1 Compare the properties of ionic compounds and covalent compounds through experiment 5.7.2 Explain with examples the uses of ionic compounds and covalent compounds in daily life
Photograph 5.2 Example of ionic compound and covalent compound
Different compounds have different properties. The different properties of ionic compounds and covalent compounds can be studied through Experiment 5.1.
122
Chemical Bond
Experiment
CHAPTER 5
5.1
Aim: To study the difference in properties between ionic compounds and covalent compounds. Problem statement: What are the difference in properties between ionic compounds and covalent compounds? Materials: Solid lead(Il) bromide, PbBr2, naphthalene, C10H8, magnesium chloride, MgCl2, cyclohexane, C6H12 and distilled water Apparatus: Test tubes, spatula, evaporating dish, Bunsen burner, pipeclay triangle, wire gauze, beaker, 10 cm3 measuring cylinder, tripod stand, battery, switch, light bulb and carbon electrodes A Electrical conductivity of compounds Hypothesis: Ionic compounds can conduct electricity in molten state but not in the solid state while covalent compounds cannot conduct electricity in both states. Variables: (a) Manipulated : Type of compound (b) Responding : Electrical conductivity (c) Fixed : Carbon electrode Procedure: 1. Put lead(II) bromide, PbBr2 powder into the crucible until half full. 2. Set up the apparatus as shown in Figure 5.21. 3. Switch on and observe whether the bulb lights up. 4. Switch off and heat the lead(II) bromide, PbBr2 powder until all solids have melted. Bulb 5. Switch on once again and observe whether the bulb lights up. Crucible 6. Repeat steps 1 to 5 using naphthalene, C10H8 powder. 7. Record your observations on the condition of the bulb in Table 5.1 Results: Table 5.1
Compound Lead(II) bromide, PbBr2
Naphthalene, C10H8
Physical state Solid Molten Solid Molten
Condition of bulb
Carry out this experiment in the fume chamber.
Safety Precaution • Naphthalene, C10H8 is a flammable substance. • Bromine gas, Br2 produced during the heating of lead(II) bromide, PbBr2 is poisonous.
Switch Carbon electrodes Lead(II) bromide, PbBr2 Heat
Figure 5.21
Chemistry & Us Excessive exposure to naphthalene, C10H8 can cause haemolytic anaemia, liver and nervous system failure, cataract and bleeding of the retina.
123
THEME 2
Fundamentals of Chemistry
B Solubility of compound in water and organic solvents Make hypothesis and state all variables. Procedure: Cyclohexane, 1. Put half spatula of magnesium chloride, MgCl2 powder Distilled Magnesium C6H12 water chloride, into the test tube. 3 MgCl 2. Add 5 cm of distilled water into the test tube and 2 shake gently. Figure 5.22 3. Observe the solubility of magnesium chloride, MgCl2 in water. 4. Repeat steps 1 to 3 using cyclohexane, C6H12 as the solvent. 5. Repeat steps 1 to 4 and substitute magnesium chloride, MgCl2 with naphthalene, C10H8. 6. Record your observations on the solubility of compounds in Table 5.2. Results: Table 5.2
Compounds
Solubility in distilled water
Solubility in cyclohexane, C6H12
Magnesium chloride, MgCl2 Naphthalene, C10H8
C Melting point and boiling point of compound Make hypothesis and state all variables. Procedure: 1. Put half spatula of magnesium chloride, MgCl2 powder and naphthalene, C10H8 into separate test tubes. 2. Heat both test tubes in the water bath as shown Magnesium in Figure 5.23. chloride, 3. Observe and record the change in physical states and make MgCl2 inference of both substances in Table 5.3. Results:
Table 5.3 Compound
Observation
Distilled water
Naphthalene, C10H8 Heat
Figure 5.23 Inference
Magnesium chloride, MgCl2 Naphthalene, C10H8
Conclusion: Is the hypothesis acceptable? What is the conclusion of this experiment? Discussion: 1. What type of compound is lead(II) bromide, PbBr2, magnesium chloride, MgCl2 and naphthalene, C10H8? 2. Predict the electrical conductivity, solubility, melting point and boiling point of sodium chloride, NaCl. Prepare a complete report after carrying out this experiment. 124
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Electrical Conductivity
Based on Experiment 5.1, ionic compounds and covalent compounds have different electrical conductivity. Ionic compounds cannot conduct electricity in the solid state but can conduct electricity in the molten state and aqueous solution while covalent compounds cannot conduct electricity in all states.
+ – + – + – + – + – + – – + – + – + – + – + – + + – + – + – + – + – + –
–
Ions cannot move freely because they are tied by strong electrostatic attraction forces. Therefore, solid ionic compound cannot conduct electricity.
+
Solid ionic compound
Ions can move freely because electrostatic attraction forces have been overcomed. Molten or aqueous solution of ionic compound can conduct electricity.
– – – + + + + + – – – – – + – – – – + + + + + – + – + – – – + + + – +
+
–
+ Molten or aqueous solution of ionic compound
–
+
Molecules in covalent compounds are neutral and do not carry any charge. Therefore, covalent compounds do not conduct electricity in all states.
Solid or molten covalent compound
Figure 5.24 Electrical conductivity of ionic compounds and covalent compounds
125
Fundamentals of Chemistry
Solubility in Water and Organic Solvents
Most ionic compounds are soluble in water but are not soluble in organic solvents. On the other hand, most covalent compounds are not soluble in water but are soluble in organic solvents. When dissolved in water, water molecules help to overcome electrostatic attraction force between ions and break down the lattice structure of the solid compound. As a result, ions can move freely in water.
Chemistry Lattice structure is the orderly arrangement of atoms, ions or molecules of a compound in a solid crystal.
Chemistry The figure below shows the solubility of sodium chloride, NaCI in water. Water is a polar solvent that has partial negative charge at the oxygen atom and partial positive charge at the hydrogen atom. Positive ion, Na+ will be attracted to the oxygen atom of water molecule which is negatively-charged while negative ion, Cl– will be attracted to the hydrogen atom of water molecule which is positively-charged. Attraction force between atom of water molecules with the ions of ionic compound are strong enough to overcome electrostatic attraction force between ions themselves. This enables most solid ionic compounds to be soluble in water.
Literacy Tips + + –+
–+
+ –
+
+ + –
+
+ + –
+ –+
+
–
–
+ – +
–
+
+
+
–
– +
+
+ –+
+
–
+ –+
– +
+ –
+
–
+
+
+
–
+ –
+ –
+
–
+
+ –
+
–+
+ + –
+
–
+ –
–+
+
–
+ –
+ – +
+
–
+
+
+
– + +
–
+
+ –+
+ + –
+ +
–
– +
–
+ + –
–
–+
+ –
+
+
– +
+
+
δ+
+ –+
+
+
+ –
THEME 2
Solid sodium chloride, NaCl
H
δ+
O
H
δ–
Water molecule, H2O Chloride ions, Cl− are attracted to the positive side of the water. Sodium ions, Na+ are attracted to the negative side of the water.
In a water molecule, oxygen atom has a higher electronegativity than hydrogen atom.This causes the electrons shared in the covalent bond to be pulled towards the oxygen atom. The unequal sharing of electrons creates partial negative charge, δ− at the oxygen atom and partial positive charge, δ+ at the hydrogen atom.
Organic solvents cannot overcome electrostatic forces between ions in a solid ionic compound. So, ionic compounds are not soluble in organic solvents. Molecules in a covalent compound are neutral and do not carry any charges. So, molecules in a covalent compound are soluble in an organic solvent but not soluble in water.
Melting Point and Boiling Point
You have learned that ionic compounds and covalent compounds are formed by ionic bonds and covalent bonds respectively. Do you know that both the chemical bonds influence the melting point and boiling point of a compound? Are these chemical bonds overcomed when compounds are melted or boiled? 126
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CHAPTER 5
Ionic compounds have high melting point and boiling point. Therefore, ionic compounds are not easily volatile.
– +– – + – + + Sodium ion, Na Sodium + ion, – +Na+ – ++– + + – + – +– – + – + + – – + + – heated Chloride ion,Chloride Cl– – + –ion, + Cl +– +–
+ –
+ –
– + – – + – + –– + + – – – –
Sodium chloride, NaCl Sodium chloride, NaCl
High heat energy is required to overcome the strong electrostatic attraction forces so that the ionic compound can melt or boil. Thus, sodium chloride, NaCl has a high melting point and boiling point.
An ionic compound like sodium chloride, NaCl consists of positive ions, Na+ and negative ions, Cl– that attract one another by strong electrostatic attraction forces.
Figure 5.25 Ionic compounds have high melting point and boiling point
Covalent compounds with simple molecules have low melting point and boiling point. Hence, covalent compounds with simple molecules are easily volatile.
H Van der Waals attraction forces between simple molecules in covalent compounds like methane, CH4 are very weak.
Van der Waals attraction H force C H H
C
H
H
H H
H H
C
H
H
Methane molecule, CH4 Low heat energy is required to overcome the weak Van der Waals attraction forces so that the covalent compound can melt or boil. Thus, methane, CH4 has a low melting point and boiling point.
Figure 5.26 Covalent compounds with simple molecules have low melting point and boiling point
Chemistry Lizards can stick to the surface of walls. This is due to the reaction between some electrons from the molecules of the hundreds of fine hairs found on the sole of the lizard’s feet and some electrons from the molecules of the wall. This reaction forms the electromagnetic attraction known as Van der Waals attraction forces.
127
THEME 2
Fundamentals of Chemistry
Structure of Covalent Compounds
There are two types of molecular structure for covalent compounds which are simple molecular structure and giant molecular structure. What is the difference between the simple molecular structure and giant molecular structure in covalent compounds? Simple molecule
Covalent compound
Giant molecule
Silicon atom, Si
Water, H2O
Oxygen atom, O
Example Covalent bond
Silicon dioxide, SiO2
Carbon dioxide, CO2
Small and simple structures can be found in the form of solids, liquids or gases.
Structure
Covalent bonds are strong in the molecules and Van der Waals attraction forces between molecules are weak.
Chemical bond
Low because only little heat is required to overcome the weak Van der Waals attraction forces between molecules.
Melting point and boiling point
Very large structure, usually exists as solids.
Strong covalent bonds in the molecules only. No Van der Waals attraction forces between molecules because of its giant structure.
High because a lot of heat is required to break the strong covalent bonds.
Figure 5.27 Difference between simple molecule and giant molecule in covalent compounds
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Uses of Ionic Compounds and Covalent Compounds in Our Daily Lives
Most ionic compounds and covalent compounds used in our daily lives are in the industrial sector, agriculture, medicine and domestic use. Industrial sector
Agricultural sector The ionic compounds of ammonium nitrate, NH4NO3 and potassium chloride, KCl are used in fertilisers.
The ionic compound of lithium iodide, LiI is used in batteries.
Paint contains covalent compounds such as pigment and turpentine solvent.
Pesticides used to kill weeds and insects contain covalent compounds such as bromoethane, C2H5Br and chloropicrin, CCl3NO2.
The uses of ionic compounds and covalent compounds in daily life
Medical sector
Sodium Bicarbonate
650 mg 100 Tablets
The ionic compound of sodium bicarbonate, NaHCO3 is used in antacids to relieve gastric pain.
Paracetamol, C8H9NO2 is a covalent compound used to treat fevers and irritation.
Domestic use Detergents contain the ionic compound, sodium chlorate(V), NaClO3, which is used for domestic cleaning.
Glycerol, C3H5(OH)3 is a covalent compound added to skincare products to moisturise skin and help to prevent dry skin.
GLYCEROL
Figure 5.28 Uses of ionic compounds and covalent compounds in our daily lives
129
THEME 2
Fundamentals of Chemistry
Activity 5.8 Carry out a problem-solving project on the use of ionic compounds and covalent compounds in daily life 1. Carry out this activity in groups. 2. Read and understand the following passage:
Century
21st Skills
CT
Plastic particles in the sea can cause problems to aquatic life ranging from plankton, fish to big animals like turtles, dolphins and whales. The problem of aquatic life is not only the direct intake of plastics but also the chemicals in the plastics that can be absorbed into the tissues of these aquatic life.
3. Apart from the problem above, surf the Internet to find information about the problems of using ionic compounds and covalent compounds in one of the following fields: (a) Industry (b) Agriculture (c) Medicine (d) Domestic 4. Discuss the ways to solve the problems. 5. Present your findings in front of the class and carry out a question and answer session to improve the proposals of each group.
TestYourself
5.7
1. Compare the melting point and boiling point of ionic compounds and covalent compounds. 2. Give one similarity between simple molecule and giant molecule of covalent compounds. 3. Magnesium hydroxide, Mg(OH)2 known as milk of magnesia, is a type of ionic compound used to relieve gastric pain. (a) State the solubility of magnesium hydroxide, Mg(OH)2 in water. (b) Can magnesium hydroxide, Mg(OH)2 conduct electricity in the solid state? (c) Explain your answer in 3(b). 4. Diamond is a giant molecule of covalent compound while methane, CH4 is a simple molecule of covalent compound. (a) Compare the melting point and boiling point of diamond and methane, CH4. Explain. (b) Predict the electrical conductivity of diamond. Explain your prediction.
130
electrostatic attraction force
Hydrogen bond
form
Positive ion
Metallic bond
Sea of electrons
delocalised electrons
Concept
Van der Waals forces
Hydrogen atoms
forms
Ionic bond
electrostatic attraction force
attraction forces between molecules
Anion
Simple molecule
Giant molecule
Non-metal elements accept electrons
attraction forces between
Ionic compound Nitrogen atom, oxygen atom or fluorine atom
Cation
donate electrons
Metal elements transfer electrons
involves valence electrons
Chemical bond
Double bond
Single bond
http://bit.ly/ 31EGBs4
Quick
Triple bond
Three pair of electrons
Covalent compound
forms
Two pair of electrons
electrons contributed by both atoms
Covalent bond
One pair of electrons
share electrons
electrons contributed by one atom only
Dative bond
Chemical Bond
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THEME 2
Fundamentals of Chemistry
Self Reflection 1. What new knowledge have you learned in Chemical Bond? 2. Which is the most interesting subtopic in Chemical Bond? Why? 3. Give a few examples on the application of Chemical Bond in daily life. 4. Rate your performance in Chemical Bond on a scale of 1 to 10; 1 being the lowest and 10 the highest. Why would you rate http://bit.ly/ yourself at that level? 31le1ft 5. What else would you like to know about Chemical Bond?
Achievement
5
1. What is the meaning of covalent bond?
2. Why does silicon dioxide, SiO2 have a high melting point and boiling point?
3. Figure 1 shows several elements in the Periodic Table of Elements that are represented by letters A, D, E, G and H. A D
E
G
H
Figure 1
(a) State the elements that can combine to form ionic compounds. (b) Element D reacts with element E to form a covalent compound. Write the chemical formula of the covalent compound formed. (c) Atoms of element H combine to form diatomic molecules at room temperature. Explain the melting point and boiling point of molecule H. 4. Figure 2 shows the chemical symbols for elements Q and R. (a) Write the electron arrangement for atom Q and atom R. (b) Element Q and element R react to form compound S. (i) State the type of chemical bond formed. (ii) Explain the process of formation of compound S.
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24 12
Q
16
Figure 2
8
R
Chemical Bond
5.
CHAPTER 5
• Atom of element J has 12 neutrons and 23 nucleon number • Atom of element K has 9 protons
(a) Which element is a metal? (b) Explain how element J combines with element K to form white solid T.
6. Element D combines with element E to form a covalent compound with chemical formula ED3. Element D has a proton number of 17. Predict the electron arrangement of the atom of element E with reasonable explanation.
7. Water, H2O exists as liquid while hydrogen chloride, HCl exists as gas at room temperature. Explain this phenomenon based on the formation of hydrogen bonds. 8. Copper, Cu is a metal that is commonly used in the manufacturing of electric wires. Explain briefly how this metal can conduct electricity. 9. Kevin found a beaker filled with white solid left on top of the table in a laboratory. He would like to know what type of compound the white solid is. He carried out several tests to investigate the physical properties of the white solid and obtained the following results:
• Soluble in water • Can conduct electricity in liquid state
Based on your observation and knowledge, predict the type of compound of this white solid. Explain your prediction.
Enrichmen Corner 1. Deoxyribonucleic acid, DNA in organism is a complex macromolecule that stores genetic information. DNA consists of polynucleotides that coil around each other to form the double helix structure as shown in Figure 1(a). Based on the DNA structure as shown in Figure 1(b), explain how polynucleotides coil around each other using the concept of hydrogen bonds.
Hydrogen bond 3' H2C
A C
O
P
A
A C
G A
H2C
I
I
A N HN
A G
A
P
H2C
I
5' 3'
CH2
O
P
NH
O
C N HN
5' H2C
T
O
O
P
Check Answers
O
NH
P
CH2
O
O
NH
O
G
C
C
P
HN
O
G NH N C
C
https://bit.ly/ 2PbIFFq
5' CH 2 O
G C
G
I
P
O HN T NH N A O
OH
5'
3'
O
O
G
HN
O
CH2
3' OH
P
(a)
(b) Figure 1
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CHAPTER
6
Acid, Base and Salt
Keywords
Basicity of acids pH and pOH Strength of acids and alkalis Molarity Standard solutions Neutralisation Titration Insoluble salts Recrystallisation Double decomposition reactions
Limestone cave, Taman Negara Mulu
What will you learn? 6.1 6.2 6.3 6.4 6.5 6.6 6.7 6.8 6.9 6.10 6.11 134
The Role of Water in Showing Acidic and Alkaline Properties pH Value Strength of Acids and Alkalis Chemical Properties of Acids and Alkalis Concentration of Aqueous Solution Standard Solution Neutralisation Salts, Crystals and Their Uses in Daily Life Preparation of Salts Effect of Heat on Salts Qualitative Analysis
Bulletin
Stalactites
How are stalactites and stalagmites formed in limestone caves? Limestone caves consist of calcium carbonate, CaCO3. When rainwater falls on the caves seep through the limestone, the following reaction takes place to produce calcium bicarbonate salt, Ca(HCO3)2. H2O(l) + CO2(g) + CaCO3(s) → Ca(HCO3)2(aq)
Stalagmites
The flowing water will carry the soluble calcium bicarbonate, Ca(HCO3)2 through the crevices at the roof of the caves. When the water comes in contact with air in the caves, a small portion of calcium bicarbonate, Ca(HCO3)2 reverts back to calcium carbonate, CaCO3, due to water and carbon dioxide losses. Calcium carbonate, CaCO3 starts to precipitate on these crevices. Hence, the formation of stalactites begins gradually. Water that drips from the ends of the stalactites will fall on the floor of the cave. Over the time, stalagmites will form in the same way as stalactites. This is why stalactites and stalagmites are found together in the caves. Formation of stalactites and stalagmites http://bit.ly/2ISEfPQ
What is the relationship between pH value and concentration of hydrogen ions, H+? Why are all alkalis bases but not all bases are alkalis? How does a laboratory assistant prepare a standard solution? 135
THEME 3
Interaction between Matter
6.1
The Role of Water in Showing Acidic and Alkaline Properties
Situation in Figure 6.1 shows the uses of acidic and alkaline substances in daily life. Identify which subtances are acidic and which substances are alkaline.
Dad, how can I make this coin shiny again?
Mom, why do we need to brush our teeth every morning?
Try scrubbing it with lime.
Figure 6.1 Acidic and alkaline substances in daily life
Acids
When acid is dissolved in water, the hydrogen atoms in acid molecules are released in the form of hydrogen ions, H+. Therefore, based on the Arrhenius theory, acid is defined as follows: Chemical substances ionise in water to produce hydrogen ions, H+. HCl(aq) → H+(aq) + Cl–(aq) HNO3(aq) → H+(aq) + NO3–(aq) When hydrogen chloride gas is dissolved in water, hydrogen chloride molecules will ionise in water to produce hydrogen ions, H+ and chloride ions, Cl–. However, do the hydrogen ions, H+ remain in the aqueous solution? Literally, hydrogen ions, H+ produced will combine with the water molecules, H2O to form hydroxonium ions, H3O+. H H
C1
+
H O
H
H O+
+
H
Figure 6.2 Formation of hydroxonium ion, H3O+
136
In order to neutralise the acid on our teeth.
C1–
g Learnin tandard S At the end of the lesson, pupils are able to: 6.1.1 Define acid and alkali 6.1.2 State the meaning of basicity of an acid 6.1.3 Investigate the role of water in showing acidic and alkaline properties through experiment
Chemistry Although hydroxonium ions, H3O+ are the actual ions existing in the aqueous solution that gives the acidic properties, to simplify explanation, we often use hydrogen ion, H+ to represent hydroxonium ions, H3O+.
Acid, Base and Salt
CHAPTER 6
Basicity of Acids
Basicity of acids refers to the number of hydrogen ions, H+ that can be produced by an acid molecule that ionises in water. Hydrochloric acid, HCl is monoprotic acid because it can produce one hydrogen ion, H+ per acid molecule. How about diprotic acid and triprotic acid? Acids Monoprotic acid
Diprotic acid
Triprotic acid
Hydrochloric acid, HCl produces one H+ ion per acid molecule
Sulphuric acid, H2SO4 produces two H+ ions per acid molecule
Phosphoric acid, H3PO4 produces three H+ ions per acid molecule
Figure 6.3 The classification of acids based on the basicity of the acids
Formic acid, HCOOH is used in the coagulation of latex. Is formic acid, HCOOH a diprotic acid? Why?
Alkalis
Brain Teaser
Base is a substance that reacts with acids to produce salt and water only. Metal oxides and metal hydroxides are bases. For example, magnesium oxide, MgO and calcium hydroxide, Ca(OH)2 are bases because they react with acids to produce salt and water only. MgO(s) + H2SO4(aq) → MgSO4(aq) + H2O(l) Salt Ca(OH)2(aq) + 2HNO3(aq) → Ca(NO3)2(aq) + 2H2O(l) Salt
Brain Teaser Observe the chemical equation below. Mg(s) + 2HCl(aq) → MgCl2(aq) + H2(g) Is magnesium a base? Why?
A base that is soluble in water is called an alkali. Potassium hydroxide, KOH and sodium hydroxide, NaOH are alkalis because they are soluble in water. When sodium hydroxide pellets, NaOH is dissolved in water, sodium ions, Na+ and hydroxide ions, OH– that can move freely in water are produced. Water
Na+ OH– OH– Na+
Na+ OH–
Na+ OH–
Figure 6.4 Dissociation of sodium hydroxide, NaOH into ions that move freely in water
137
THEME 3
Interaction between Matter
An alkali is defined as follows: Chemical substances that ionise in water to produce hydroxide ions, OH–. What will happen to ammonia molecule when ammonia gas is dissolved in water? Why is aqueous ammonia produced an alkali? H H
N
H
+
H
H
H
N+
H
O
H
H
+
O–
H Water, H2O
Ammonia gas, NH3
Hydroxide ion, OH–
Ammonium ion, NH4+
NH4+(aq) + OH–(aq) NH3(aq) + H2O(l) Figure 6.5 Formation of hydroxide ion, OH– from ammonia molecule
By dissolving ammonia gas in water, aqueous ammonia is produced. Aqueous ammonia is an alkali because the ammonia molecules ionise partially to produce hydroxide ions, OH–.
Uses of Acids, Bases and Alkalis
Acids, bases and alkalis are not just chemical substances in the laboratory but they are also widely found in daily life. Toothpaste which is alkaline, functions to neutralise acid on the teeth, while vinegar is an acidic substance used to make pickled chillies. Photograph 6.1 Uses of acid and alkali in daily life
Activity 6.1
Century CT Discussing the uses of acids and alkalis in daily life using 21st Skills examples of acidic and alkaline substances 1. Carry out the activity in groups. 2. Gather information from reading materials or websites on examples of acidic and alkaline substances as well as their uses in various fields.
TABLET
ANTACID
Agriculture
138
Industries
Medicine
Food industry
Acid, Base and Salt
CHAPTER 6
3. Based on the information gathered, discuss the followings: (a) Identify the acid, base or alkali in each substance that you have found (b) State the use of acid, base or alkali found in the substance 4. Pin up your group work on the bulletin board to share the information and references with the other groups.
The Role of Water to Show Acidity and Alkalinity Mom, why is the new soap taken out of its box not slippery?
Try to add some water and rub the soap. What do you feel now? It's slippery, mom!
Figure 6.6 Role of water to show alkalinity
Based on the conversation in Figure 6.6, why is the water added to the soap? Is water needed to allow acids or alkalis to show acidic or alkaline properties?
Experiment
6.1
Aim: To study the role of water in showing acidic properties. Problem statement: Is water needed to allow an acid to show its acidic properties? Hypothesis: Water is needed for an acid to show its acidic properties. Variables: (a) Manipulated : Presence of water (b) Responding : Colour change on blue litmus paper (c) Fixed : Type of acid Materials: Solid oxalic acid, C2H2O4, distilled water and blue litmus paper Apparatus: Test tubes and test tube rack Procedure: 1. Add a spatula of solid oxalic acid, C2H2O4 in a test tube. 2. Insert a piece of dry blue litmus paper into the test tube. 137
THEME 3
Interaction between Matter
3. Observe any changes to the colour of the blue litmus paper. Record your observations. 4. After that, add 2 cm3 distilled water and shake well. 5. Observe any changes to the colour of the blue litmus paper. Record your observations. Results:
Acid is corrosive. Be careful when handling acid. If in contact with acid, run continuous flow of water on the affected area.
Table 6.1 Contents
Observations
Solid oxalic acid, C2H2O4 Solid oxalic acid, C2H2O4 + water
Interpreting data: 1. State the change in colour of the blue litmus paper that is used to detect acidic properties. 2. Based on the observations, state a suitable inference. 3. What are the conditions needed for an acid to show its acidic properties? Conclusion: Is the hypothesis acceptable? What is the conclusion of this experiment? Discussion: 1. Name the ion that is responsible for showing the acidic properties. 2. Solid oxalic acid, C2H2O4 had differences in observation compared to the solid oxalic acid, C2H2O4 that has been dissolved in water. Give a reason. 3. What is the operational definition for acid in this experiment? Prepare a complete report after carrying out this experiment. Acids only show acidic properties in the presence of water. When an acid is dissolved in water, acid molecules will ionise to produce hydrogen ions, H+. The presence of hydrogen ions H+ allows the acid to show its acidic properties. Therefore, blue litmus paper changes to red. Without water, solid oxalic acid, C2H2O4 only exist as molecules. Hydrogen ions, H+ are not present. Thus, the colour of blue litmus paper remains unchanged.
140
Literacy Tips Reflect on the properties of acid: Sour taste Corrosive Has pH value less than 7 Changes moist blue litmus paper to red
Acid, Base and Salt
Experiment
CHAPTER 6
6.2
Aim: To study the role of water in showing alkaline properties. Problem statement: Is water needed to allow an alkali to show its alkaline properties? Hypothesis: Make a suitable hypothesis for this experiment. Variables: State all the variables involved in this experiment.
Sodium hydroxide pellet, NaOH
Red litmus paper
Sodium hydroxide solution, NaOH produced
Sodium hydroxide, NaOH is corrosive. A pellet of sodium hydroxide, NaOH is sufficient to carry out this experiment. If in contact with the alkaline solution, run water over the area continuously until it no longer feels slippery.
Figure 6.7 Method to test the alkaline properties of sodium hydroxide, NaOH
Procedure: 1. Based on Figure 6.7, list out the apparatus and materials required for this experiment. 2. Plan the procedure for this experiment with your group members. 3. Determine the method used to collect data and prepare a suitable table. 4. Carry out the experiment with your teacher̕s permission. 5. Record your observations in the table provided. Results: Record your observation in a table.
Interpreting data: 1. Based on the observations, state a suitable inference. 2. What is the condition of the litmus paper needed to detect alkaline properties? Conclusion: Is the hypothesis acceptable? What is the conclusion of this experiment?
Discussion: 1. Name the ion responsible to show alkaline properties. 2. Explain the difference in observation between using a pellet of sodium hydroxide, NaOH and sodium hydroxide solution, NaOH. 3. Give the operational definition for alkali in this experiment. Prepare a complete report after carrying out this experiment.
137
THEME 3
Interaction between Matter
Akalis only shows alkaline properties when they are dissolved in water. Without water, hydroxide ions, OH– in the sodium hydroxide pellet, NaOH cannot move freely and are still tied in its lattice structure. So, the pellet of sodium hydroxide, NaOH does not show alkaline properties. The red litmus paper cannot change colour. When a pellet of sodium hydroxide, NaOH is dissolved in water, hydroxide ions, OH– are produced and able to move freely in water. Thus, sodium hydroxide solution, NaOH shows alkaline properties. Hence, the moist red litmus paper turns blue.
Literacy Tips Reflect on the properties of an alkali: Tastes bitter and feels slippery Corrosive Has pH value more than 7 Changes moist red litmus paper to blue
The presence of water also enables ammonia gas, NH3 to ionise and produce hydroxide ions, OH– that are responsible for its alkaline properties. Therefore, the moist red litmus paper turns blue. Without water, ammonia gas, NH3 only exists as molecules. Hydroxide ions, OH– are not present. So, red litmus paper remains unchanged.
TestYourself
6.1
1. State the meaning of the following terms: (a) Acid (b) Alkali 2. Carbonic acid is a mineral acid with the formula, H2CO3. What is the basicity of carbonic acid? Explain why. 3. Figure 6.8 shows a conversation between Khairul and his teacher. Khairul, what is your problem?
Teacher, the cleaning powder is alkaline. Why doesn't the red litmus paper turn blue?
Figure 6.8
(a) What possible mistake was committed by Khairul in his test? (b) How can you help Khairul in his test? Explain why.
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Acid, Base and Salt
6.2
pH Value
The pH Values of Acids and Alkalis
g Learnin tandard S
Zahir, what are you doing?
At the end of the lesson, pupils are able to: 6.2.1 State the meaning of pH and its uses 6.2.2 Calculate pH values of acids and alkalis 6.2.3 Investigate the relationship between pH value and the concentration of hydrogen ions and hydroxide ions through experiment
I am taking a sample of water to test the pH value of water in this aquarium. The clownfish that I am rearing need water with pH values between 8.0 to 8.3 in order for them to be healthy.
Figure 6.9 Clownfish need water with specific pH values
Based on the pH value mentioned by Zahir, do clownfish require acidic or alkaline water? Why do you say so? The pH scale which ranges from 0 to 14 is used to show the acidity and alkalinity of an aqueous solution. Solutions with pH value less than 7 is acidic while solutions with pH value more than 7 is alkaline. Universal indicator, pH meter or pH paper is commonly used to determine the pH value. Referring to Figure 6.10, what is the relationship between pH value and degree of acidity or alkalinity? More acidic
0
1
2
3
More alkaline
Neutral
4
5
6
7
8
9
10
11
12
13
14
Figure 6.10 The pH scale
What is actually ̒pH̕? In chemistry, pH is a logarithmic measure of the concentration of hydrogen ions in an aqueous solution. pH = –log [H+] where log is logarithm of base 10 and [H+] is the concentration of hydrogen ions in mol dm–3 of the solution. Using that formula, we can determine the pH value of an acid by calculation.
137
THEME 3
Interaction between Matter
Example
1
Calculate the pH value of nitric acid, HNO3 with 0.5 mol dm–3 of hydrogen ion, H+. Solution
Given the concentration of H+ = 0.5 mol dm–3 Use the formula pH = –log [H+] pH = –log [0.5] = –(– 0.301) = 0.301 pH value of nitric acid, HNO3 = 0.3 Example
2
Determine the molarity of hydrochloric acid, HCl with pH value of 2.0. Solution
pH = –log [H+] 2.0 = –log [H+] log [H+] = –2.0 [H+] = 10–2 = 0.01 mol dm–3 Molarity of hydrochloric acid, HCl = 0.01 mol dm–3 The concentration of hydroxide ion, OH– is used to calculate the value of pOH of an alkali based on the following formula, where [OH–] represents the concentration of hydroxide ions in mol dm–3 of the alkali solution. pOH = –log [OH–] Given that the sum of pH value and pOH value is 14, the pH value of an alkali can be calculated by using the following relationship: pH + pOH = 14 pH = 14 – pOH Example 3
Calculate the pOH value for sodium hydroxide solution, NaOH with 0.1 mol dm–3 hydroxide ions, OH–. Solution
Given that the concentration of hydroxide ion, OH– = 0.1 mol dm–3 pOH = –log [0.1] Use the formula pOH = –log [OH–] = –(–1) =1 pOH value of sodium hydroxide solution, NaOH = 1.0 144
Acid, Base and Salt
Example
CHAPTER 6
4
Calculate the pH value for potassium hydroxide, KOH that has 0.01 mol dm–3 hydroxide ions, OH–. Solution
Given concentration of hydroxide ion, OH– = 0.01 mol dm–3 Use the formula pOH = –log [OH–] pOH = –log [0.01] = –(–2) =2 pOH value of potassium hydroxide solution, KOH = 2.0 pH value of potassium hydroxide solution, KOH = 14.0 – pOH = 14.0 – 2.0 = 12.0
Consider the relationship: pH + pOH = 14
Example 5
Determine the molarity of lithium hydroxide solution, LiOH with pH value 12.0. Solution
pH + pOH = 14.0 12.0 + pOH = 14.0 pOH = 14.0 – 12.0 = 2.0 pOH = –log [OH–] 2.0 = –log [OH–] log [OH–] = –2.0 [OH–] = 10–2 = 0.01 mol dm–3 Molarity of lithium hydroxide solution, LiOH = 0.01 mol dm–3
Did you know that the decimal place of pH is related to the significant numbers in the concentration of hydrogen ions given? If the value of given concentration has two significant numbers, the pH value should be rounded to two decimal places.
The pH value can be calculated based on the concentration of hydrogen ions, H+ in an acid, or the concentration of hydroxide ions, OH– in an alkali. So, the pH scale allows us to compare the concentration of hydrogen ions, H+ or the hydroxide ions, OH– in an aqueous solution. The relationship between hydrogen ions, H+ or hydroxide ions, OH– with the pH value can be studied in Experiment 6.3. 137
THEME 3
Interaction between Matter
Experiment
6.3
Aim: To study the relationship between the concentration of hydrogen ions, H+ and pH value of acid. Problem statement: Does the concentration of hydrogen ions, H+ of an acid affect its pH value? Hypothesis: The higher the concentration of hydrogen ion, H+, the lower the pH value of the acid. Variables: (a) Manipulated : Concentration of hydrogen ions, H+ (b) Responding : pH value (c) Fixed : Type of acid Materials: 0.1 mol dm–3, 0.01 mol dm–3 and 0.001 mol dm–3 hydrochloric acid, HCl Apparatus: 100 cm3 beaker and pH meter Procedure: 1. Pour 20 cm3 of hydrochloric acid, HCl of different concentrations into three beakers. 2. Measure the pH value of each hydrochloric acid, HCl with the pH meter. 3. Record the pH values in Table 6.2. Results: Table 6.2
Concentration of hydrochloric acid, HCl (mol dm–3)
0.1
0.01
0.001
Concentration of hydrogen ions, H+ (mol dm–3) pH value
Interpreting data: 1. Based on the results obtained, how does the pH value change when the concentration of hydrochloric acid, HCl decreases? 2. State the changes in the concentration of hydrogen ions, H+ when the concentration of hydrochloric acid, HCl decreases. 3. What is the relationship between the concentration of hydrogen ions, H+ and pH value? Conclusion: Is the hypothesis acceptable? What is the conclusion of this experiment? Discussion: 1. When an acidic solution is diluted, what are the changes in the: (a) Concentration of hydrogen ions, H+? (b) pH value? (c) Degree of acidity of the aqueous solution? 2. State the relationship between the concentration of hydrogen ions, H+, pH values and degree of acidity of an acidic aqueous solution.
Prepare a complete report after carrying out this experiment. 146
Acid, Base and Salt
Experiment
CHAPTER 6
6.4
Aim: To study the relationship between the concentration of hydroxide ions, OH– and pH value of an alkali. Problem statement: Does the concentration of hydroxide ions, OH– of an alkali affect its pH value? Hypothesis: Make a suitable hypothesis for this experiment. Variables: State all variables involved in this experiment. Materials: 0.1 mol dm–3, 0.01 mol dm–3 and 0.001 mol dm–3 sodium hydroxide solution, NaOH Apparatus: 100 cm3 beaker and pH meter Procedure: 1. Plan the procedure to measure the pH value of sodium hydroxide solution, NaOH. 2. Your plan should include the pH meter. 3. Carry out the experiment with your teacher̕s permission. 4. Record the pH values obtained in your report book. Results: Record the pH values in a table. Interpreting data: 1. Based on the data obtained, how does the pH value change when the concentration of sodium hydroxide solution, NaOH decreases? 2. When the concentration of sodium hydroxide solution, NaOH decreases, what are the changes that occur to the: (a) Concentration of hydroxide ions, OH–? (b) pH value? (c) Degree of alkalinity of sodium hydroxide solution, NaOH? 3. State the relationship between the concentration of hydroxide ions, OH–, pH value and degree of alkalinity of sodium hydroxide, solution NaOH. Conclusion: Is the hypothesis acceptable? What is the conclusion of this experiment?
Prepare a complete report after carrying out this experiment. When the concentration of acid increases, more acid molecules ionise to produce hydrogen ions, H+. The higher the concentration of hydrogen ions, H+, the lower the pH value. Acidity increases when the pH value of the acid solution decreases. Concentration of hydrogen ions, H+
, pH value
On the other hand, the higher the concentration of hydroxide ions, OH–, the higher the pH value. Alkalinity increases when the pH value of the alkaline solution increases. Concentration of hydroxide ions, OH–
, pH value 137
THEME 3
Interaction between Matter
Most substances found in our daily lives contain acids or alkalis. The determination of pH values for these substances can be done in Activity 6.2.
Chemistry & Us Purple cabbages change colour at different pH values.
Activity 6.2 Determining the pH values of various items in daily life 1. You are supplied with the following items: • Soap water • Lime juice
• Milk tea • Coffee
• Carbonated drink • Tap water
2. In pairs, measure the pH value of each item using the universal indicator. 3. Record the items with similar pH values. 4. Prepare a pH indicator using a purple cabbage. Visit websites to know how to prepare this pH indicator. Use the pH indicator that you have prepared to measure the pH value of each of the above items. 5. Using a suitable graphic management tools, present your findings. 6. Pin up your work in class to share with others.
Sewage pipe Battery Lemon Tomato Milk Blood Antacid Soap cleaner Milk
Baking soda
Ammonia
Gastric juice Vinegar Coffee Water Baking soda Ammonia solution
Bleach
Bleach
Figure 6.11 The pH value of few items in daily life tested with the universal indicator
TestYourself
6.2
1. Write the formula to calculate the pH value of acid. 2. Calculate the pH value for hydrochloric acid, HCl that contains 0.001 mol dm–3 hydrogen ions, H+. 3. Determine the pH value for calcium hydroxide, Ca(OH)2 with concentration of 0.05 mol dm–3. [pH + pOH = 14]. 148
Acid, Base and Salt
6.3
CHAPTER 6
Strength of Acids and Alkalis
Observe Figure 6.12. What is meant by a strong acid and a weak acid? Can you tell the similarities or the differences between these two acids?
Teacher, both acids are monoprotic.
Hydrogen ions, H+ are produced when acid is dissolved in water.
HCl is a strong acid, CH3COOH is a weak acid.
Figure 6.12 Similarities and differences between two acids
Strong Acids and Weak Acids
The strength of an acid depends on the degree of dissociation or ionisation of the acid in water. Strong Acids A strong acid is an acid that ionises completely in water to produce a high concentration of hydrogen ions, H+. Hydrochloric acid, HCl is a strong acid because all molecules of hydrogen chloride, HCl that dissolve in water are ionised completely to hydrogen ions, H+ and chloride ions, Cl–. No molecule of hydrogen chloride, HCl exists in this solution. HCl(aq) → H+(aq) + Cl–(aq)
Hydrogen chloride gas, HCl
Water
Hydrogen chloride molecule, HCl
Hydrogen ion, H+ Chloride ion, Cl–
g Learnin tandard S At the end of the lesson, pupils are able to: 6.3.1 Define strong acid, weak acid, strong alkali and weak alkali 6.3.2 Explain the strength of acid and alkali based on its degree of dissociation in water
Chemistry Hydrogen ions, H+ produced from acid molecules will combine with water molecules to form hydroxonium ions, H3O+. The hydroxonium ions, H3O+ is the product of a dative bond formed between hydrogen ion, H+ with water molecule.
Figure 6.13 Complete ionisation of hydrogen chloride, HCl
137
THEME 3
Interaction between Matter
Weak Acids A weak acid is an acid that ionises partially in water to produce low concentration of hydrogen ions, H+. Ethanoic acid, CH3COOH is a weak acid because the molecules of ethanoic acid, CH3COOH ionise partially in water. The degree of dissociation of ethanoic acid molecules is 1.54%. In other words, from 100 molecules of ethanoic acid, CH3COOH, only one molecule of ethanoic acid, CH3COOH ionises to hydrogen ions, H+ and ethanoate ions, CH3COO–. The rest still exist as molecules of ethanoic acid, CH3COOH. H+(aq) + CH3COO–(aq) CH3COOH(aq)
Ethanoic acid, CH3COOH molecule
Water
The reversible arrow shows that ethanoic acid, CH3COOH molecule can form hydrogen ions, H+ and ethanoate ions, CH3COO–. These ions can also combine again to form the acid molecules.
Glacial ethanoic acid, CH3COOH Ethanoic acid, CH3COOH molecule Hydrogen ion, H+ Ethanoate ion, CH3COO–
Ethanoic acid, CH3COOH
Figure 6.14 Partial ionisation of ethanoic acid, CH3COOH
Strong Alkalis and Weak Alkalis
Alkalis also consist of strong alkalis and weak alkalis depending on their degree of ionisation in water. Strong Alkalis A strong alkali is an alkali that ionises completely in water to produce a high concentration of hydroxide ions, OH–. Sodium hydroxide, NaOH is a strong alkali that undergoes complete dissociation when dissolved in water. Only sodium ions, Na+ and hydroxide ions, OH– are present in the solution. NaOH(aq) → Na+(aq) + OH–(aq) Sodium ion, Na+ Sodium hydroxide solution, NaOH
Hydroxide ion, OH–
Figure 6.15 Complete ionisation of sodium hydroxide solution, NaOH
150
Literacy Tips Dissociation is also known as ionisation.
Acid, Base and Salt
CHAPTER 6
Weak Alkalis A weak alkali is an alkali that ionises partially in water to produce a low concentration of hydroxide ions, OH–. Ammonia solution, NH3 is a weak alkali because ammonia molecules, NH3 ionise partially in water. The degree of dissociation of ammonia, NH3 is 1.3%. In other words, from 100 molecules of ammonia, NH3 only one molecule of ammonia, NH3 will receive hydrogen ion, H+ from water molecule. So, only a small number of hydroxide ions, OH– is present in the solution. NH4+(aq) + OH–(aq) NH3(aq) + H2O(l) Ammonia molecule, NH3 Ammonia gas, NH3 Ammonia molecule, NH3 Hydroxide ion, OH– Water
Ammonium ion, NH4+
Figure 6.16 Partial ionisation of ammonia solution, NH3
Activity 6.3
Century CT Carry out a simulation to explain the strength of acids and alkalis 21st Skills 1. Visit the website given. Simulation on acid 2. Vary the regulator for acid strength and observe the degree and alkali + of dissociation and number of hydrogen ions, H present. http://bit.ly/31cGoMQ 3. Repeat step 2 for alkali and observe the degree of – dissociation and number of hydroxide ions, OH present. 4. Interpret the information on strength of acids and alkalis based on the degree of dissociation. 5. Relate the concentration of hydrogen ions, H+ and hydroxide ions, OH– with the degree of dissociation of the acid and alkali. 6. Display your findings in an interesting presentation.
TestYourself
6.3
1. Give the meaning of the following terms: (a) Strong acid (c) Strong alkali (b) Weak acid (d) Weak alkali 2. Why does ammonia solution, NH3 that has the same concentration as potassium hydroxide, KOH have a lower pH value? 3. The pH value 0.1 mol dm-3 nitric acid, HNO3 is different from the pH value 0.1 mol dm–3 oxalic acid, H2C2O4? Explain. 137
THEME 3
Interaction between Matter
6.4
Chemical Properties of Acids and Alkalis
Balloon filled with baking soda
After 30 seconds
Carbon dioxide gas
Vinegar
Figure 6.17 To inflate the balloon using vinegar and baking soda
Have you ever inflated a balloon using vinegar and baking soda? Is this process related to the chemical properties of acid?
Chemical Properties of Acid
The properties of acid is divided into physical and chemical properties. The properties of acid such as having a sour taste, changing moist blue litmus paper to red and having pH values less than 7 are the physical properties of acid. The chemical properties of acid on the other hand refer to the reactions between acid and other substances. Carry out activity 6.4 to study the chemical properties of acid.
g Learnin tandard S At the end of the lesson, pupils are able to: 6.4.1 Summarise the chemical properties of acids by carrying out the reactions between: • Acid and base • Acid and reactive metal • Acid and metal carbonate 6.4.2 Summarise the chemical properties of alkalis by carrying out the reactions between: • Alkali and acid • Alkali and metal ion • Alkali and ammonium salt
Balloon inflation http://bit.ly/2Mh3M7y
Activity 6.4 Aim: To study the chemical properties of acids. Materials: Copper(II) oxide powder, CuO, zinc powder, Zn, marble chips, CaCO3, 1.0 mol dm–3 sulphuric acid, H2SO4, 2.0 mol dm–3 nitric acid, HNO3, 2.0 mol dm–3 hydrochloric acid, HCl, limewater, wooden splinter and filter papers Apparatus: 100 cm3 beaker, glass rod, filter funnel, retort stand with clamp, evaporating dish, Bunsen burner, pipeclay triangle, delivery tube and rubber stopper, tripod stand, spatula, test tubes and test tube holder Reaction between acid and base Procedure: 1. Pour 20 cm3 of 1.0 mol dm–3 sulphuric acid, H2SO4 into a beaker. Heat the acid by using a small flame. 2. Add copper(II) oxide powder, CuO gradually into the acid by using a spatula. Stir the mixture with a glass rod. 3. Observe the change that takes place on copper(II) oxide, CuO that reacts with acid. Record your observation on the solution produced. 152
Acid, Base and Salt
4. Continue adding copper(II) oxide powder, CuO until it can no longer dissolve. 5. Filter out the excess copper(II) oxide, CuO from the mixture. 6. Pour the filtrate into an evaporating dish and heat the filtrate until one third of its initial volume remains. 7. Allow the saturated solution produced to cool until salt crystals are formed. 8. Filter the contents of the evaporating dish to obtain the salt crystals. Rinse the crystals with distilled water. 9. Dry the salt crystals between two pieces of filter papers. 10. Observe the physical properties of the salt crystals and record your findings. Glass rod Spatula Excess of copper(II) oxide powder, CuO
Filter funnel Filter paper
Residue (excess of copper(II) oxide powder, CuO)
Hot 20 cm3 of 1.0 mol dm–3 sulphuric acid, H2SO4
Filtrate (copper(II) sulphate solution, CuSO4)
Filtrate
Evaporating dish
Salt crystals
Pipeclay triangle Heat
Figure 6.18 Preparation of salt crystals from the reaction between acid and base
Discussion: 1. What happens to the copper(II) oxide powder, CuO when added to sulphuric acid, H2SO4? 2. What is the colour of the solution produced from the reaction between sulphuric acid, H2SO4 and copper(II) oxide, CuO? 3. Write a chemical equation for the reaction between sulphuric acid, H2SO4 and copper(II) oxide, CuO. 4. From the chemical equation written above, complete the following equation in words: Acid + Base →
+
137
CHAPTER 6
THEME 3
Interaction between Matter
Reaction between acid and reactive metal Procedure: 1. Plan a procedure to study the reaction between hydrochloric acid, HCl and zinc powder, Zn as shown in Figure 6.19. Glass rod Spatula Excess of zinc powder, Zn
Filter funnel Filter paper Residue (excess of zinc powder, Zn) Filtrate (zinc chloride solution, ZnCl2)
20 cm3 of 2.0 mol dm–3 hydrochloric acid , HCl
Filtrate
Evaporating dish
Salt crystals Pipeclay triangle Heat
Figure 6.19 Preparation of salt crystals from the reaction between acid and reactive metal
2. Discuss with your teacher if you have encountered any problems when Burning 5 cm3 of planning the procedure. wooden 2.0 mol dm–3 3. Make sure that you carry out the splinter hydrochloric chemical test on the gas released as acid, HCl shown in Figure 6.20. Zinc powder, Zn 4. Carry out this test with your teacher̕s permission. Figure 6.20 5. Record your observations. Discussion: 1. What is the observation that indicates acid has reacted with metal when zinc powder, Zn is added to hydrochloric acid, HCl? 2. Name the gas released in this activity. 3. Write a chemical equation for the reaction between hydrochloric acid, HCl and zinc, Zn. 4. From the chemical equation written above, complete the following equation in words: Acid + Reactive metal → 154
+
Acid, Base and Salt
CHAPTER 6
Reaction between acid and metal carbonate Procedure: 1. Plan a procedure to carry out this activity to study the reaction between nitric acid, HNO3 and marble chips, CaCO3. 2. Include safety measures taken in your procedure. 3. Discuss with your teacher if you have encountered any problems when planning the procedure. 4. Make sure that you carry out the chemical test on the gas released as shown in Figure 6.21. 5 cm3 of 2.0 mol dm–3 nitric acid, HNO3
Limewater
Marble chips, CaCO3
Figure 6.21
5. Carry out this test with your teacher̕s permission. 6. Record your observations. Discussion: 1. What is the reason for using excess marble chips, CaCO3 to react with nitric acid, HNO3? 2. How do you remove the excess marble chips, CaCO3 from the salt solution produced? 3. For the reaction in this activity: (a) Name the salt produced (b) Name the gas released 4. Write a chemical equation for the reaction between nitric acid, HNO3 and marble chips, CaCO3. 5. From the chemical equation written above, complete the following equation in words: Acid + Metal carbonate →
+
+
Prepare a complete report after carrying out this activity. From Activity 6.4 that was carried out, it can be summarised that acids have the following chemical properties: Acids react with bases to produce salt and water Acids react with reactive metals to produce salt and hydrogen gas, H2 Acids react with metal carbonates to produce salt, water and carbon dioxide gas, CO2
137
THEME 3
Interaction between Matter
Chemical Properties of Alkalis
Chemical properties of alkalis can be determined through Activity 6.5.
Activity 6.5 Aim: To study the chemical properties of alkali. Materials: Benzoic acid powder, C6H5COOH, 1.0 mol dm–3 sodium hydroxide solution, NaOH, ammonium chloride powder, NH4Cl, copper(II) sulphate solution, CuSO4, distilled water, filter paper and red litmus paper Apparatus: 100 cm3 beaker, glass rod, filter funnel, retort stand with clamp, evaporating dish, Bunsen burner, pipeclay triangle, tripod stand, dropper, spatula, test tube, boiling tube and test tube holder Figure 6.22, Figure 6.23 and Figure 6.24 show three reactions involving alkalis. Glass rod Spatula
Filter funnel
Benzoic acid powder, C6H5COOH
Filter paper Residue (excess acid powder) Filtrate
20 cm3 of 1.0 mol dm–3 sodium hydroxide solution, NaOH Evaporating dish Filtrate
Salt crystals
Pipeclay triangle Heat
Figure 6.22 Preparation of salt crystals from the reaction between alkali and acid Moist red litmus paper
Ammonium chloride powder, NH4Cl
1.0 mol dm–3 sodium hydroxide solution, NaOH
Heat
Figure 6.23 Heating the mixture of alkali and ammonium salt to produce ammonia gas
156
1.0 mol dm–3 sodium hydroxide solution, NaOH
Copper(II) sulphate solution, CuSO4
Blue precipitate
Figure 6.24 Addition of alkali to metal ions to produce insoluble metal hydroxide precipitate
Acid, Base and Salt
CHAPTER 6
Procedure: 1. Based on Figure 6.22 to Figure 6.24, plan a laboratory activity to study the chemical properties of alkalis. 2. Plan and write out the procedure for the laboratory activity to be discussed with your teacher. 3. Record your observations in a report book. 4. Write an equation in words to summarise the chemical properties of alkalis. Prepare a complete report after carrying out this activity. From Activity 6.5 that was carried out, we can summarise that alkalis have the following chemical properties: Alkalis react with acids to produce salt and water When a mixture of alkali and ammonium salt is heated, ammonia gas, NH3 is released Addition of an alkali to most metal ions, will produce an insoluble metal hydroxide precipitate Table 6.3 summarises the chemical properties of acids and alkalis Table 6.3 Chemical properties of acids and alkalis
Acid + Base → Salt + Water
Example: 2HNO3(aq)
+
Nitric acid
CuO(s)
Copper(II) oxide
→
Cu(NO3)2(aq)
Chemical Acid + Reactive metal → Salt + Hydrogen gas properties + Zn(s) → ZnSO4(aq) of acids Example: H2SO4(aq)
Sulphuric acid
Zinc
+
CaCO3(s)
→
CaCl2(aq)
+
Hydrochloric acid Calcium carbonate Calcium chloride
+
Potassium hydroxide
→
H2SO4(aq)
Sulphuric acid
Potassium hydroxide Ammonium chloride
+
Water
CO2(g)
Carbon dioxide gas
K2SO4(aq)
Potassium sulphate
+
Potassium chloride
2H2O(l) Water
H2O(l) Water
+
NH3(g)
Ammonia gas
Alkali + Metal ion → Insoluble metal hydroxide + Cation from alkali
Example: 2NaOH(aq)
H2(g)
Hydrogen gas
H2O(l)
Chemical Alkali + Ammonium salt → Salt + Water + Ammonia gas properties + NH4Cl(aq) → KCl(aq) + of alkalis Example: KOH(aq)
Water
Alkali + Acid → Salt + Water
Example: 2KOH(aq)
+
Zinc sulphate
H2O(l)
Acid + Metal carbonate → Salt + Water + Carbon dioxide gas
Example: 2HCl(aq)
+
Copper(II) nitrate
Sodium hydroxide
+
Mg2+(aq)
Magnesium ion
→
Mg(OH)2(s)
Magnesium hydroxide
+
2Na+(aq)
Sodium ion
137
THEME 3
Interaction between Matter
TestYourself
6.4
1. Write a chemical equation for the reaction between hydrochloric acid, HCl and: (a) Barium hydroxide, Ba(OH)2 (b) Magnesium, Mg (c) Zinc carbonate, ZnCO3 2. Write an equation in words to summarise the reaction of an alkali solution and the following substances: (a) Dilute acids (b) Ammonium salts (c) Metal ions
6.5
Concentration of Aqueous Solution
Dad, why is the colour of my tea is different from the one that you are drinking?
Because the concentration of tea in our glasses are different.
g Learnin tandard S At the end of the lesson, pupils are able to: 6.5.1 State the meaning of concentration of aqueous solution 6.5.2 Solve numerical problems involving concentration of solution
Figure 6.25 Concentration of tea affects its colour
Concentration of a solution is a measurement that shows the quantity of solute dissolved in a unit volume of solution, normally in 1 dm3 solution. The higher the quantity of solute, the higher the concentration of the solution. The quantity of solute dissolved can be measured in gram or mole, hence the concentration of a solution can be measured in unit g dm–3 or mol dm–3. Concentration
in unit g dm–3, is the mass of solute found in 1 dm3 solution. Concentration (g dm–3) =
Mass of solute (g) Volume of solution (dm3)
in unit mol dm–3, is the number of moles of solute found in 1 dm3 solution. This concentration is called molarity.
Concentration
Molarity (mol dm–3) = 158
Number of moles of solute (mole) Volume of solution (dm3)
Acid, Base and Salt
Literacy Tips
÷ molar mass
Concentration (g dm–3)
Molarity (mol dm–3) × molar mass
Figure 6.26 Relationship between concentration and molarity
Example
CHAPTER 6
The unit for molarity is mol dm–3 or molar (M). You have to remember that mole is not the same as molar. Mole is the unit for measuring matter while molar is the number of moles of solute in a given volume of solution.
6
Calculate the concentration in g dm–3, for each solution produced. (a) 40 g of solid copper(II) sulphate, CuSO4 is dissolved in water to produce 20 dm3 solution. (b) 18 g of sodium hydroxide pellets, NaOH is dissolved in water to produce 750 cm3 solution. Solution
Mass of solute (g) Volume of solution (dm3) 40 g = 20 dm3 = 2.0 g dm–3
(a) Concentration of copper(II) sulphate, CuSO4 = (b) Concentration of sodium hydroxide, NaOH Mass of solute (g) = Volume of solution(dm3) 18 g = 0.75 dm3 = 24.0 g dm–3
Example
750 cm3 is converted to dm3 by dividing the volume with 1000.
750 dm3 1000 = 0.75 dm3
7
Calculate the molarity of each solution prepared. (a) 10 mol of solid zinc chloride, ZnCl2 dissolved in water to produce 5 dm3 of solution. (b) 0.1 mol of solid calcium chloride, CaCl2 is dissolved in 500 cm3 of distilled water. Solution
Number of moles of solute (mol) Volume of solution (dm3) = 10 mol3 5 dm = 2.0 mol dm–3
(a) Molarity of zinc chloride solution, ZnCl2 =
(b) Molarity of calcium chloride solution, CaCl2 Number of moles of solute (mol) = Volume of solution (dm3) 500 cm3 is converted 0.1 mol to dm3 by dividing the = 0.5 dm3 volume with 1000. = 0.2 mol dm–3
500 dm3 1000 = 0.5 dm3
159
THEME 3
Interaction between Matter
Example
8
What is the concentration of nitric acid, HNO3 with a molarity of 0.5 mol dm–3 in unit g dm–3? [Relative atomic mass: H = 1, N = 14, O = 16] Solution
Concentration = Molarity × Molar mass HNO3
RAM H RAM N RAM O
= 0.5 mol dm–3 × [1 + 14 + 3(16)] g mol–1 = 0.5 mol dm–3 × 63 g mol–1 = 31.5 g dm–3
Example
9
Convert the concentration of 3.6 g dm–3 lithium hydroxide solution, LiOH to molarity, mol dm–3. [Relative atomic mass: H = 1, Li = 7, O = 16] Solution
Concentration Molar mass LiOH 3.6 g dm–3 = (7 + 16 + 1) g mol–1
Molarity =
RAM Li RAM O RAM H
= 0.15 mol dm–3
We can calculate number of moles of solute dissolved in the solution if its molarity and the volume of the solution are known. Number of moles of solute (mol) Molarity = Volume of solution (dm3) M = n V Therefore, n = MV
Volume of solution is in dm3.
If the volume of the solution is in cm3, thus, unit of volume needs to be converted to dm3. n = M V 1000 Volume of Therefore, n = MV 1000 solution is in cm3.
(
)
Example 10
Calculate the number of moles of potassium hydroxide, KOH found in 2 dm3 of 0.5 mol dm–3 potassium hydroxide solution, KOH. Solution
Number of moles, n = MV = 0.5 mol dm–3 × 2 dm3 = 1 mol KOH 160
This formula is applied because the volume of solution is in dm3.
Acid, Base and Salt
CHAPTER 6
Example 11
A beaker contains 200 cm3 of 0.2 mol dm–3 lead(II) nitrate solution, Pb(NO3)2. How many moles of lead(II) nitrate, Pb(NO3)2 is in the beaker? Solution
Number of moles, n = MV 1000 –3 3 = 0.2 mol dm × 200 cm 1000 = 0.04 mol Pb(NO3)2
This formula is applied because the volume of solution is in cm3.
Activity 6.6 Solving numerical problems related to concentration of solutions
CT
1. 6 g of solid magnesium sulphate, MgSO4 is added into a beaker containing 200 cm3 of water. Calculate the concentration in g dm–3, for the solution produced. 2. 0.4 mol of zinc chloride, ZnCl2 is dissolved in water to produce 2 dm3 of solution. Calculate the molarity of the solution prepared. 3. What is the concentration of 0.5 mol dm–3 sulphuric acid, H2SO4 in g dm–3? [Relative atomic mass: H = 1, O = 16, S = 32] 4. The concentration of sodium chloride solution, NaCl is 1.989 g dm–3. Calculate the molarity of the solution in mol dm–3. [Relative atomic mass: Na = 23, Cl = 35.5] 5. Calculate the number of moles of sodium hydroxide, NaOH in 2.5 dm3 of 0.2 mol dm–3 sodium hydroxide solution, NaOH. 6. Given the molarity of 250 cm3 of barium hydroxide solution, Ba(OH)2 is 0.1 mol dm–3. How many moles of hydroxide ion, OH– is in the solution?
TestYourself
6.5
1. What is meant by concentration in unit g mol–1? 2. State two units to measure concentration. 3. 0.03 mol of potassium nitrate, KNO3 is dissolved in 1.2 dm3 of distilled water. What is the molarity of the potassium nitrate solution, KNO3 produced? 4. Calculate the concentration of sulphuric acid, H2SO4 that has the molarity of 2.0 mol dm–3 in unit g dm–3. [Relative atomic mass: H = 1, O = 16, S = 32] 5. 1.9 g MgY2 is dissolved in 100 cm3 of water to produce a solution with the molarity of 0.2 mol dm–3. What is the relative atomic mass of Y? [Relative atomic mass: Mg = 24] 161
THEME 3
Interaction between Matter
6.6
Standard Solution
Have you seen the syrup dispenser as shown in Photograph 6.2? Did you know that the dispenser is filled with a standard solution of sugar so that the machine can dispense sugar at an amount requested by customers? What do you know about standard solution? Syrup dispenser http://bit.ly/35Ajhiv
g Learnin tandard S
Photograph 6.2 Syrup dispenser
Meaning of Standard Solution
Most chemical reactions involve reactants in aqueous solution. In that case, the preparation of aqueous solution with specific concentrations is very important. Standard solution is a solution with known concentration. In the preparation of standard solutions, mass of solute and volume of distilled water are two parameters that have to be measured accurately.
At the end of the lesson, pupils are able to: 6.6.1 State the meaning of standard solution 6.6.2 Describe the preparation of a standard solution through activity: • From a solid substance • Through dilution of an aqueous solution 6.6.3 Solve numerical problems involving preparation of standard solution and dilution
Preparation of a Standard Solution from a Solid
Activity 6.7 Aim: To prepare 250 cm3 of standard solution of 1.0 mol dm–3 sodium carbonate, Na2CO3. Materials: Distilled water and solid sodium carbonate, Na2CO3 Apparatus: Electronic balance, filter funnel, 250 cm3 volumetric flask, dropper, wash bottle, 250 cm3 beaker and glass rod Procedure: 1. Determine the mass of sodium carbonate, Na2CO3 needed using the formula n = MV . 1000 2. Weigh the mass calculated using the electronic balance. 3. Add 100 cm3 of distilled water to the solid sodium carbonate, Na2CO3 in a beaker. 4. Stir the mixture with a glass rod until all the solid sodium carbonate, Na2CO3 is completely dissolved in the distilled water. 5. Transfer the sodium carbonate solution, Na2CO3 into a 250 cm3 volumetric flask via a filter funnel. 6. Rinse the beaker with distilled water. Make sure all the remaining solution is transferred into the volumetric flask. 7. Then, rinse the filter funnel with a little distilled water. All the remaining solution is transferred into the volumetric flask. 8. Remove the filter funnel. Add distilled water until it approaches the calibration mark on the volumetric flask. 162
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9. Using a dropper, add distilled water slowly until the meniscus level is aligned exactly on the calibration mark of the volumetric flask. 10. Close the volumetric flask with a stopper. Shake well by inverting the volumetric flask several times until the solution is homogenous. Note: Keep the standard solution of sodium carbonate, Na2CO3 that you have prepared for Activity 6.8. Distilled water
Solid sodium carbonate, Na2CO3
Glass rod
Filter funnel
Stir
Volumetric flask (a) Add distilled water
(b) Dissolve the solid
Calibration mark
(c) Transfer the solution into a volumetric flask
Dropper
Stopper
Distilled water Meniscus level of solution
(f) Close the volumetric flask with a stopper before shaking
Calibration mark
Meniscus level of solution
(e) Add distilled water until calibration mark
Calibration mark
(d) Rinse the filter funnel with distilled water
Figure 6.27 Preparation of 1.0 mol dm–3 sodium carbonate solution, Na2CO3 from a solid
Discussion: 1. Why must the beaker and filter funnel be rinsed with distilled water? 2. Why must all the remaining solution be transferred into the volumetric flask? 3. How can you ensure that the meniscus level aligns with the calibration mark of the volumetric flask? 4. Why does the volumetric flask need to be closed after the standard solution is prepared? Prepare a complete report after carrying out this activity. Sodium hydroxide, NaOH is not suitable to be used for the preparation of a standard solution because sodium hydroxide, NaOH is hygroscopic (absorbs water or moisture in the air). Sodium hydroxide, NaOH also absorbs carbon dioxide gas, CO2 in the air to form sodium carbonate, Na2CO3. 2NaOH(s) + CO2(g) → Na2CO3(s) + H2O(l) This causes difficulty to determine the exact mass of sodium hydroxide, NaOH. Therefore, the preparation of a standard solution of sodium hydroxide, NaOH with a known concentration could not be made. 163
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Solid oxalic acid H2C2O4.2H2O can be used to prepare a standard solution in the laboratory.
Chemistry Sodium carbonate, Na2CO3 that is used to prepare standard solutions is alkaline. When sodium carbonate, Na2CO3 is dissolved in distilled water, carbonate ions, CO32– react with water molecules to produce bicarbonate ions, HCO3– and hydroxide ions, OH–. The presence of hydroxide ions, OH– gives the alkaline properties to the solution. CO32–(aq) + H2O(l) → HCO3–(aq) + OH–(aq)
Preparation of a Standard Solution by Diluting Aqueous Solution Another method of preparing solutions of known concentration is by dilution method. This method involves adding water to a concentrated standard solution, or known as stock solution, to produce a more diluted solution. During dilution, water that is added to the aqueous solution will alter the concentration of the solution but it would not alter the number of moles of solute contained in the solution. Add distilled water Solute Solute (a) Concentrated solution
(b) Dilute solution
Figure 6.28 Quantity of solute remains the same in both solutions of different concentrations
Hence, Number of moles of solute before dilution = Number of moles of solute after dilution n1 = n2 M1V1 MV = 2 2 1000 1000
M1V1 = M2V2
where M1 is the molarity of aqueous solution (stock solution) before dilution. V1 is the volume of aqueous solution (stock solution) before dilution. M2 is the molarity of aqueous solution (prepared solution) after dilution. V2 is the volume of aqueous solution (prepared solution) after dilution. As an example, you wish to prepare 500 cm3 of 0.1 mol dm–3 copper(II) sulphate solution, CuSO4 from the stock solution of 2.0 mol dm–3 copper(II) sulphate, CuSO4. Use the following formula: M1V1 = M2V2 (2.0)(V1) = (0.1)(500) (0.1)(500) V1 = 2.0 = 25 cm3 Hence, 25 cm3 of stock solution of copper(II) sulphate, CuSO4 needs to be diluted using distilled water until 500 cm3 solution of copper(II) sulphate, CuSO4 is obtained. 164
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The preparation of a standard solution by dilution method can be carried out through Activity 6.8 using sodium carbonate solution, Na2CO3 prepared in Activity 6.7.
Activity 6.8 Aim: To prepare 100 cm3 of standard solution of 0.2 mol dm–3 sodium carbonate, Na2CO3. Materials: Distilled water and 1.0 mol dm–3 sodium carbonate solution, Na2CO3 from Activity 6.7 Apparatus: 100 cm3 volumetric flask, dropper, filter funnel, pipette, wash bottle, pipette filler and 100 cm3 beaker Procedure: Pipette Stock solution
Beaker
(a) Pour stock solution from Activity 6.7 into a beaker
Stopper
Beaker
(b) Take out calculated volume of solution, V1 cm3 with a pipette
Volumetric flask
Calibration mark
(c) Transfer V1 cm3 of solution into a volumetric flask
Dropper
Meniscus level of solution (f) Close the volumetric flask and shake well by inverting it several times until the solution is homogenous
Stock solution
Calibration mark
(e) Add distilled water slowly with a dropper
Distilled water
Calibration mark
(d) Add distilled water until the solution level approaches the calibration mark
Brain Teaser
Figure 6.29 Preparation of 0.2 mol dm–3 sodium carbonate solution, Na2CO3 by dilution method
1. Based on Figure 6.29, plan a procedure to prepare a standard solution of 0.2 mol dm–3 sodium carbonate, Na2CO3 by dilution method. 2. Include precautionary steps in the process of solution preparation. 3. Show your procedure to your teacher before carrying out this activity. 4. Carry out the procedure as planned. 5. Clean and keep the apparatus at their proper places after carrying out this activity.
Brain Teaser Why is the pipette not rinsed with distilled water but rinsed with 1.0 mol dm–3 sodium carbonate solution, Na2CO3?
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Discussion: 1. What is the volume of the standard solution of 1.0 mol dm–3 sodium carbonate solution, Na2CO3 needed to prepare 100 cm3 of 0.2 mol dm–3 sodium carbonate solution, Na2CO3? 2. What is the size of pipette needed in this preparation process? 3. Why is the beaker not suitable to be used in preparing a standard solution by the dilution method? 4. Do you need to remove the last drop of the solution in the pipette? Why? Prepare a complete report after carrying out this activity. Examples 12 and 13 show samples of calculations involved in the preparation of a standard solution by dilution. Example 12
Figure 6.30 shows 75 cm3 of 2.0 mol dm–3 nitric acid, HNO3 that is diluted to x mol dm–3 when 25 cm3 distilled water is added. Calculate the value of x.
25 cm3 distilled water Nitric acid, HNO3 x mol dm–3
75 cm3 of 2.0 mol dm–3 nitric acid, HNO3 Before dilution
Solution
After dilution
Figure 6.30
M1 = 2.0 mol dm–3 ; V1 = 75 cm3 Volume of solution M2 = x mol dm–3 ; V2 = (75 + 25) cm3 = Volume of HNO3 + Volume of distilled water = 100 cm3 2.0 mol dm–3 × 75 cm3 = x mol dm–3 × 100 cm3 Use the formula M1V1 = M2V2 –3 3 x mol dm–3 = 2.0 mol dm ×3 75 cm 100 cm = 1.5 mol dm–3 Then, x = 1.5 Example 13
Determine the volume of 2.0 mol dm–3 hydrochloric acid, HCl needed to be pipetted into a volumetric flask 250 cm3 to produce 0.2 mol dm–3 hydrochloric acid, HCl. Solution
M1 = 2.0 mol dm–3 ; V1 = ? M2 = 0.2 mol dm–3 ; V2 = 250 cm3 2.0 mol dm–3 × V1 = 0.2 mol dm–3 × 250 cm3 –3 cm3 V1 = 0.2 mol dm × 250 –3 2.0 mol dm = 25 cm3 166
Use the formula M1V1 = M2V2
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Activity 6.9 CT Solving calculation problems involved in the preparation of a standard solution by dilution 1. Calculate the volume of 2.0 mol dm–3 sodium carbonate solution, Na2CO3 needed to prepare 50 cm3 of 0.1 mol dm–3 sodium carbonate solution, Na2CO3. 2. What is the molarity of sodium hydroxide solution, NaOH when 30 cm3 distilled water is added to 50 cm3 of 0.5 mol dm–3 sodium hydroxide solution, NaOH? 3. Calculate the volume of solution produced when 50 cm3 of 1.2 mol dm–3 sodium nitrate solution, NaNO3 is diluted to 0.5 mol dm–3. 4. When 200 cm3 water is added to 50 cm3 concentrated sulphuric acid, H2SO4, sulphuric acid, H2SO4 with concentration 0.2 mol dm–3 is produced. Calculate the molarity of the initial concentration of sulphuric acid, H2SO4.
TestYourself
6.6
1. What is meant by standard solution? 2. X cm3 of 0.15 mol dm–3 zinc nitrate solution, Zn(NO3)2 is pipetted into a 500 cm3 volumetric flask to produce 500 cm3 of 0.018 mol dm–3 zinc nitrate solution, Zn(NO3)2. Determine the value of X. 3. Calculate the new molarity of hydrochloric acid, HCl produced if 25 cm3 of 1.5 mol dm–3 hydrochloric acid, HCl is diluted to produce 150 cm3 of hydrochloric acid, HCl. 4. Determine the volume of distilled water needed to add to 50 cm3 of 0.2 mol dm–3 sodium thiosulphate solution, Na2S2O3 so that a 0.025 mol dm–3 sodium thiosulphate solution, Na2S2O3 is produced.
6.7
Neutralisation
When stung by a bee, the area that had been stung can be treated with baking soda. Vinegar, on the other hand, is used to treat the area that had been stung by a wasp. Why?
g Learnin tandard S At the end of the lesson, pupils are able to: 6.7.1 State the meaning of neutralisation 6.7.2 Determine the concentration of an unknown solution through titration method 6.7.3 Solve numerical problems involving neutralisation
Photograph 6.3 Bee and wasp
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Definition of Neutralisation
Neutralisation is a reaction between an acid and an alkali (base) to produce salt and water only. In the reaction, the salt and water produced are neutral because the acid lose its acidity and the alkali lose its alkalinity. Acid + Alkali → Salt + Water For example, the neutralisation reaction between nitric acid, HNO3 with potassium hydroxide, KOH to produce potassium nitrate solution, KNO3 and water, H2O. HNO3(aq) + KOH(aq) → KNO3(aq) + H2O(l) In neutralisation, the actual reaction that occurs only involves the combination of hydrogen ions, H+, from the acid and the hydroxide ions, OH– from the alkali to produce water molecules, H2O. Hence, the ionic equation for the reaction is as follows: H+(aq) + OH–(aq) → H2O(l) The following shows how the ionic equations for neutralisation reaction can be obtained. Chemical equation: HNO3(aq)
+
KOH(aq)
→
KNO3(aq)
+
H+(aq) + NO3–(aq) + K+(aq) + OH–(aq) → K+(aq) + NO3–(aq) + Ions in nitric acid
Ions in potassium hydroxide solution
Ions in potassium nitrate solution
H2O(l) H2O(l) Water molecules
K+ and NO3– are considered as spectator ions that do not change in the reaction. Thus, these ions are cancelled out in the equation.
Ionic equation:
H+(aq) + OH–(aq) → H2O(l)
Activity 6.10 CT Write chemical equations and ionic equations for neutralisation reactions 1. Complete and balance the following equations. After that, write the relevant ionic equation. (a) HCl(aq) + Ba(OH)2(aq) → (b) H2SO4(aq) + KOH(aq) → (c) HNO3(aq) + NaOH(aq) → 2. Play the role as a chemistry teacher by explaining your findings in front of your classmates.
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Applications of Neutralisation in Daily Life
Figure 6.31 shows the application of neutralisation for a variety of uses in daily life. Medicine
Hair health
Dental health
Agriculture
H OT TO
Shampoo
E ST PA
f Milk oe ia Magn s
an
ark
eg
ny
Me
Milk of magnesia Mg(OH)2 relieves gastric pain by neutralising the excessive hydrochloric acid in the stomach.
Weak alkali in the shampoo neutralises acid on hair.
Toothpaste contains a base that neutralises lactic acid produced by bacteria in our mouth.
Slaked lime, Ca(OH)2 which is alkaline, is used to treat acidic soil.
Figure 6.31 Applications of neutralisation in daily life
Activity 6.11 Solving problems on soil fertility using suitable fertilisers 1. Carry out this activity in groups. 2. Study the following problem statement:
STEM 21st Century Skills
CT
Apart from treating acidic soil, fertilisers need to be added to soil to replace nutrients such as nitrogen, potassium and phosphorus that have been absorbed by plants. There is a variety of fertilisers in the market. Which fertilisers are suitable for plants? 3. Gather information concerning the problem given above. (a) What type of crops were planted? (b) What are the type of elements required by the crops? (c) Identify the fertiliser that is suitable for the crops by considering the percentage of elements such as nitrogen, phosphorus, and other needs, fertiliser cost and the quantity needed for the area. 4. Present your group findings in a multimedia presentation. Neutralisation reaction is also applied in the production of fertilisers such as urea, potassium sulphate, K2SO4, ammonium nitrate, NH4NO3 and others. For example, urea can be produced from the neutralisation reaction between ammonia, NH3 and carbon dioxide, CO2. How about other fertilisers? Try to list out the acids and alkalis involved in the production of that fertilisers. 169
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Activity 6.12 Gather information on various types of fertilisers CT 1. Carry out this activity in groups. 2. Visit the websites or refer to printed materials in the library and resource centre to gather information about: (a) Ways to produce urea through the reaction between ammonia and carbon dioxide. Include the chemical equations involved (b) Types of ammonium fertilisers available in the market (c) Calculate the percentage by mass of nitrogen for urea and ammonium fertilisers in the market. Then, compare and determine the quality of fertiliser based on the percentage of nitrogen 3. Use suitable graphic organisers to present your group work to your classmates.
Titration Method
Titration method is a quantitative analysis method to determine the volume of acid needed to completely neutralise a given volume of alkali and vice versa. In acid-base titration, a solution of known concentration is slowly added from a burette into a conical flask that contains a volume of alkali of unknown concentration. Titration stops as soon as the acid-base indicator changes colour. The point in the titration at which the acid-base indicator changes colour is known as the end point.
Activity 6.13 Aim: To determine the concentration of potassium hydroxide solution, KOH by acid-base titration. Materials: 1.0 mol dm–3 nitric acid, HNO3, potassium hydroxide, KOH (unknown concentration), phenolphthalein indicator and distilled water Apparatus: Burette, 25 cm3 pipette, pipette filler, 250 cm3 conical flask, white tile and retort stand with clamp Procedure: 1. Rinse a 25 cm3 pipette with a little potassium hydroxide Retort Burette solution, KOH. Remove the solution. stand 3 2. Pipette exactly 25 cm of potassium hydroxide solution, KOH. Transfer it into a conical flask. 3. Add a few drops of phenolphthalein indicator into the 1.0 mol dm–3 potassium hydroxide solution, KOH and swirl the flask. nitric acid, HNO3 4. Rinse a burette with 1.0 mol dm–3 nitric acid, HNO3. Then, remove the whole solution. Potassium 5. Fill the burette with 1.0 mol dm–3 nitric acid, HNO3 hydroxide Conical flask and clamp the burette onto a retort stand. Record the solution, KOH + few drops of initial reading of the burette. White tile phenolphthalein 6. Drip 1.0 mol dm–3 nitric acid, HNO3 slowly into the conical flask while swirling it. Figure 6.32 170
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7. Stop adding 1.0 mol dm–3 nitric acid, HNO3 as soon as the Safety Precaution colour of the solution in the conical flask changes from pink Make sure the eye position to colourless. Record the final burette reading. is parallel to the meniscus 3 –3 8. Calculate the rough volume, V cm of 1.0 mol dm nitric level of the solution while acid, HNO3 that is needed for titration. taking a burette reading. 9. Repeat steps 2 and 3. 10. Flow 1.0 mol dm–3 nitric acid, HNO3 until (V – 5) cm3 into Literacy Ti s the conical flask containing 25 cm3 of potassium hydroxide solution, KOH. Then, stop the flow of 1.0 mol dm–3 nitric Although, the addition of acid, HNO3. distilled water to rinse the –3 inner part of a conical flask 11. Subsequently add 1.0 mol dm nitric acid, HNO3, drop by changes the concentration drop, into the conical flask while swirling the flask. of the mixture solution, the 12. Occasionally, rinse the inner surface of the conical flask with number of moles of acid and alkali reacted remain distilled water to ensure all the 1.0 mol dm–3 nitric acid, HNO3 unchanged. Thus, the volume has been titrated into the potassium hydroxide solution, KOH. of acid needed to neutralise 13. Stop the titration as soon as the colour of the solution alkali is not affected. turns colourless. 14. Record the final burette reading. Titration procedure 15. Repeat steps 9 – 14 twice. http://bit.ly/2Bev0Fq 16. Record your readings in Table 6.4. Results:
p
Table 6.4
Number of titration
Rough
1
2
3
Initial burette reading (cm3) Final burette reading (cm3) Volume of nitric acid, HNO3 needed (cm3)
Interpreting data: 1. What is the average volume of nitric acid, HNO3, that is needed to neutralise 25 cm3 of potassium hydroxide solution, KOH by ignoring the rough volume? 2. Write the ionic equation for the reaction between nitric acid, HNO3 and the potassium hydroxide solution, KOH. 3. Calculate the number of moles of nitric acid, HNO3, needed in this neutralisation reaction. 4. Calculate the number of moles of potassium hydroxide solution, KOH needed to react completely with the number of moles of nitric acid, HNO3 calculated in question 3. 5. Determine the molarity of potassium hydroxide solution, KOH. Discussion: 1. Why is a white tile used in this activity? 2. Why should not we rinse the inside of the conical flask with potassium hydroxide solution, KOH before beginning the titration? 3. What is the operational definition for the end point in this activity? Prepare a complete report after carrying out this activity. 171
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The end point of the neutralisation process can be determined when the acid-base indicator changes colour. When the end point is achieved, all the hydrogen ions, H+ completely neutralise all the hydroxide ions, OH– to produce water molecules. Table 6.5 shows the phenolphthalein indicator and methyl orange in acidic, neutral and alkaline conditions. Table 6.5 Colours of indicators in acidic, neutral and alkaline conditions Colour in medium
Indicator Phenolphthalein Methyl orange
Acidic
Neutral
Alkaline
Colourless
Colourless
Pink
Red
Orange
Yellow
Solving Numerical Problems Involving Neutralisation
MV If a mol acid A is completely neutralised by b mol of alkali B, then the formula a a = a can MbVb b be used to solve the calculation related to the neutralisation reaction. a mol b mol
a Acid A + b Alkali B → c Salt
Molarity of acid A = Ma Volume of acid A = Va
+ d Water
Molarity of alkali B = Mb Volume of alkali B = Vb
Based on the equation above, the mole ratio of acid A to alkali B is a:b. Example 14
20 cm3 of 0.25 mol dm–3 sodium hydroxide solution, NaOH is neutralised with 0.2 mol dm–3 hydrochloric acid, HCl. Calculate the volume of hydrochloric acid, HCl needed for this neutralisation reaction. Solution
a = 1 mol
b = 1 mol
HCl +
NaOH → NaCl + H2O
Ma = 0.2 mol dm–3 ; Va = ? Mb = 0.25 mol dm–3 ; Vb = 20 cm3
Write this chemical equation and determine the value of a and b based on the coefficients of this chemical equation.
0.2(Va) MV Use the formula a a = a = 1 MbVb b 0.25(20) 1 0.2(Va) = 1 × (0.25)(20) 1 0.25(20) Va = 0.2 = 25 cm3 Volume of hydrochloric acid, HCl that is needed = 25 cm3
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Further examples http://bit.ly/ 35UzBuQ
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Example 15
4.05 g of zinc oxide, ZnO is needed to complete the neutralisation of 50 cm3 of nitric acid, HNO3. Calculate the concentration of the acid in mol dm–3. [Relative atomic mass: H = 1, N = 14, O = 16, Zn = 65] Solution
a = 2 mol
2 HNO3 +
b = 1 mol
ZnO → Zn(NO3)2 + H2O 4.05 g Number of moles of ZnO, n = (65 + 16) g mol–1
RAM Zn RAM O
Write a balanced chemical equation. Convert the given quantity (4.05 g) to number of moles.
= 0.05 mol
Based on the chemical equation, 2.0 mol of HNO3 reacts with 1.0 mol of ZnO 0.1 mol of HNO3 reacts with 0.05 mol of ZnO
Number of moles of HNO3, n = MV 1000 (M)(50) 0.1 mol = 1000 M = 2.0 mol dm–3
Based on the mol ratio, determine the number of moles of HNO3. Convert the number of moles of HNO3 to molarity.
Molarity of nitric acid, HNO3 = 2.0 mol dm–3
Activity 6.14 Solving numerical problems involving neutralisation 1. 25 cm3 of 0.2 mol dm–3 sodium hydroxide solution, NaOH is titrated with 0.1 mol dm–3 sulphuric acid, H2SO4. What is the volume of sulphuric acid, H2SO4 needed to neutralise sodium hydroxide solution, NaOH?
CT
2. Sulphuric acid, H2SO4, reacts with ammonia solution, NH3 according to the following chemical equation: H2SO4(aq) + 2NH3(aq) → (NH4)2SO4(aq) It is given that T cm3 of 0.125 mol dm–3 sulphuric acid, H2SO4 exactly neutralises 25 cm3 of 1.0 mol dm–3 ammonia solution, NH3. Determine the total volume of solution in the conical flask at the end point of titration. 3. 50 cm3 of nitric acid, HNO3 completely neutralises 50 cm3 of 0.25 mol dm–3 calcium hydroxide, Ca(OH)2. Calculate the molarity of the nitric acid, HNO3. 4. In a titration, 15 cm3 of 0.5 mol dm–3 sulphuric acid, H2SO4 neutralises 20 cm3 of potassium hydroxide solution, KOH. Calculate the concentration of potassium hydroxide solution, KOH. 173
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TestYourself
6.7
1. What is the meaning of neutralisation? 2. State the changes in the methyl orange indicator inside the conical flask containing potassium hydroxide solution, KOH when it reaches end point. 3. 50 cm3 of 0.75 mol dm–3 ammonia solution, NH3 is titrated with 1.0 mol dm–3 nitric acid, HNO3. What is the volume of 1.0 mol dm–3 nitric acid, HNO3 that is needed to neutralise the ammonia solution, NH3? 4. Calculate the volume of 0.05 mol dm–3 hydrochloric acid, HCl that exactly neutralises 25 cm3 of 0.1 mol dm–3 barium hydroxide solution, Ba(OH)2. 5. Based on the following chemical equation, 20 cm3 monoprotic acid, HX reacts completely with 10 cm3 of 0.1 mol dm–3 potassium hydroxide solution, KOH. HX(aq) + KOH(aq) → KX(aq) + H2O(l) What is the molarity of this acid? 6. A student dissolves hydrogen chloride gas, HCl in water to produce 500 cm3 of acidic solution. Calculate the molarity of the solution if 6 g of copper(ll) oxide, CuO is used for a complete reaction with the solution produced. [Relative atomic mass: O = 16, Cu = 64]
6.8
Salts, Crystals and Their Uses in Daily Life
The common table salt used in cooking is made up of sodium ions, Na+ and chloride ions, Cl–. The egg shell is made up of calcium ions, Ca2+ and carbonate ions, CO32–. Is calcium carbonate a type of salt too?
g Learnin tandard S At the end of the lesson, pupils are able to: 6.8.1 State the meaning of salt 6.8.2 Characterise the physical properties of salt crystals 6.8.3 Give examples of salt and their uses in daily life
Figure 6.33 Salt and egg shell
Definition of Salt
Salt is an ionic compound. Salt can be produced from the neutralisation reaction between acid and alkali (base). 174
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Reaction between acid and alkali
Anion from acid
H Cl (aq) + Na OH(aq) → NaCl(aq) + H2O(l) Sodium chloride salt
Cation from base
Can salt only be produced from acid and base reaction? Look at the following chemical equations to find out further on the salt production concept. Reaction between acid and reactive metals Anion from acid
2H Cl (aq) + Zn (s) → ZnCl2(aq) + H2(g)
Cation from metal
Zinc chloride salt
Reaction between acid and metal carbonate
Anion from acid
2H Cl (aq) + Ca CO3(s) → CaCl2(aq) + CO2(g) + H2O(l) Cation from metal carbonate
Calcium chloride salt
Reaction between acid and aqueous ammonia Anion from acid
H Cl (aq) + NH3 (aq) → NH4Cl(aq) Cation from aqueous ammonia
Ammonium chloride salt
Based on the chemical equations above, salt can be defined as follow: Salt is an ionic compound formed when the hydrogen ion, H+ from the acid is replaced with the metal ion or the ammonium ion, NH4+.
Activity 6.15 Gathering and interpreting information on the existence of salts that exist naturally 1. Carry out this activity in groups. 2. Gather information on salts that exist naturally. Your information should include the following: (a) Name of the salts (b) Source or location of the salts (c) Relevant photograph of the salts 3. Interpret the information gathered with suitable graphic organisers. 4. Present the information in front of your class. 175
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Interaction between Matter
Physical Properties of Salt Crystals
All salt crystals have specific features. Can you state the physical properties of a salt crystal? Activity 6.16 can assist you in showing characteristics of a salt crystal.
Activity 6.16 Carrying out a crystal growth activity 1. Carry out this activity in pairs. Crystal growth 2. Watch the video clip on the steps taken to produce a large http://bit.ly/2IOG7ZY crystal through the growth of crystal. 3. Discuss with your partner on the important procedures in producing copper(II) sulphate crystal, CuSO4. 4. Carry out the crystal growth activity of copper(II)sulphate, CuSO4 in a time frame of two weeks with your teacher̕s permission. 5. Dry the crystal produced and observe the crystal under a microscope. 6. Record the physical properties of the crystal and sketch its shape in your notebook.
has a specific geometrical shape such as cube, cuboid, rhombus and prism
has flat surface, straight sides and sharp vertices
different crystals have different geometrical shapes has a fixed angle between two adjacent surfaces
same crystals of different sizes still have the same geometrical shapes
Figure 6.34 Physical properties of salt crystal A crystal has specific properties because the particles in the crystal are arranged in compact and orderly manner according to a specific design arrangement.
Examples of Salts and Their Uses
Besides sodium chloride salt, NaCl that we use everyday, there are more salts that exist naturally as minerals in the Earth’s crust. These salts have their own uses in various fields. 176
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Activity 6.17 Making a multimedia presentation on the uses of various salts 1. Carry out this activity in groups. 2. Surf the Internet or refer to printed materials at the library to gather information on the various uses of salt in the following fields: Agriculture
Medicine Preservation Food preparation 3. Interpret the information obtained and present your group’s work using multimedia presentations. Ammonium nitrate, NH4NO3 as fertiliser while iron(II) sulphate, FeSO4 is used in pesticides to kill pests and grass.
Calcium sulphate, CaSO4 as plaster of Paris to support broken bones while potassium manganate(VII) is used as antiseptic to treat wounds.
Antiseptic
Agriculture
Medicine
Food preparation
Preservation
Sodium chloride, NaCl is used as flavour. Sodium bicarbonate, NaHCO3 is used for raising dough.
Sodium benzoate, C6H5COONa is used to preserve chilli sauce, tomato sauce and oyster sauce. Sodium nitrate, NaNO3 is used to preserve processed meat such as sausages.
Figure 6.35 Various uses of salt in daily life
Activity 6.18 Debating the effects of salt on humans 1. Read and understand the following extract: Table salt, Himalayan salt and bamboo salt are among the salts found on Earth. Humans need salt to maintain the fluid balance in their body, prevent muscle cramps and others. However, a high salt content will cause high blood pressure, stroke, kidney failure and other diseases. 2. Gather information on the effects of salt on humans. 3. Debate the effects of salt on human health. 177
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Interaction between Matter
TestYourself
6.8
1. Define the meaning of salt. 2. List out the physical properties of a salt crystal. 3. Give examples of salts and their uses in the following fields: (a) Agriculture (b) Medicine
6.9
g Learnin Standard
Preparation of Salts
The solid as shown in Photograph 6.4 is the Himalayan salt. Does the Himalayan salt dissolve in water? Is it true that all types of salt dissolve in water? Photograph 6.4 Himalayan salt
Solubility of Salt in Water
Salt is an ionic compound. The solubility of various salts in water is investigated in Experiment 6.5.
Experiment
At the end of the lesson, pupils are able to: 6.9.1 Test the solubility of salt in water and classify them into soluble and insoluble salts through experiment 6.9.2 Describe the preparation of a soluble salt through activity 6.9.3 Describe the preparation of an insoluble salt through activity 6.9.4 Construct an ionic equation using the continuous variation method through experiment
6.5
Aim: Investigate the solubility of various salts in water. Problem statement: Do all salts dissolve in water? Hypothesis: Some salts dissolve in water, some salts do not. Variables: (a) Manipulated : Types of nitrate, sulphate, chloride, carbonate and ammonium salt (b) Responding : Solubility of salt in water (c) Fixed : Volume and temperature of water, mass of salt Procedure: Distilled water
Stir
Glass rod Salt ON OFF
g
Electronic balance
Salt
Figure 6.36 Apparatus set-up to investigate the solubility of salt
1. Based on Figure 6.36, list out the apparatus and materials/substances used in this experiment. 178
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2. Plan the experiment procedure with the members of your group. 3. Prepare an appropriate table to record your observation. Do not taste the salts. There 4. Carry out the experiment with your teacher̕s permission. 5. Record the observation obtained in the table you have prepared. are salts that are poisonous. Interpreting data: 1. Based on the results of the experiment, list out: (a) Nitrate, sulphate, chloride, carbonate and ammonium salts that dissolve in water (b) Sulphate, chloride and carbonate salts that do not dissolve in water 2. Formulate and classify the types of salt that dissolve or do not dissolve in water in an appropriate table. Conclusion: Is the hypothesis acceptable? What is the conclusion of this experiment? Prepare a complete report after carrying out this experiment. Soluble salts are salts that dissolve in water at room temperature and non-soluble salts are salts that do not dissolve at room temperature. Table 6.6 shows the types of salts and their solubilities in water. Table 6.6 Solubility of salts in water
Types of salts Nitrate salt (NO3–)
Soluble in water All nitrate salts
Sulphate salt (SO4 )
All sulphate salts
Chloride salt (Cl–)
All chloride salts
Carbonate salt (CO32–)
Sodium carbonate, Na2CO3 Potassium carbonate, K2CO3 Ammonium carbonate, (NH4)2CO3
Ammonium, sodium and potassium salts
All ammonium, sodium and potassium salts
2–
Insoluble in water None except Lead(II) sulphate, PbSO4 Barium sulphate, BaSO4 Calcium sulphate, CaSO4 except Mercury(I) chloride, Hg2Cl2 Lead(II) chloride, PbCl2 Silver chloride, AgCl Other carbonate salts
None
Chemistry Ammonium Potassium NO ASaP NO3–
P ← PbSO4 B ← BaSO4 C ← CaSO4
Sodium
H ← Hg2Cl2 P ← PbCl2 A ← AgCl
Literacy Tips Mnemonics: All NO ASaP salts dissolve in water. PBC sulphates and HPA chlorides do not dissolve in water.
The empirical formula for mercury(I) chloride is HgCl. Because mercury atoms, Hg tend to form Hg-Hg bonds, therefore the chemical formula for mercury(I) chloride is Hg2Cl2. Bond between Hg-Hg Cl
Hg
Hg
Cl
179
Brain Teaser THEME 3
Interaction between Matter
Lead(II) chloride, PbCl2 and lead(II) iodide, PbI2 are two types of salts that are special. These salts are initially insoluble in water, but can be dissolved in hot water to produce a colourless solution. Solids are reformed when water is cooled.
Brain Teaser Are lead(II) chloride, PbCl2 and lead(II) iodide, PbI2 classified as soluble or insoluble salts? Why?
White solid dissolves in hot water to form a colourless solution
White solid of lead(II) chloride, PbCl2
White solid is formed again when water is cooled
Heat
Yellow solid dissolves in hot water to form a colourless solution
Yellow solid of lead(II) iodide, PbI2
Yellow solid is formed again when water is cooled
Heat
Figure 6.37 Special properties of lead(II) chloride, PbCl2 and lead(II) iodide, PbI2
Preparation of Soluble Salts
The method to prepare salts depends on the solubility of the salt in water and the type of salt required. Figure 6.38 shows the various methods for the preparation of soluble and insoluble salts. Salt Insoluble salt
Soluble salt All ammonium salts, sodium salts and potassium salts Neutralisation reaction between acid and alkali
Acid
Alkali
Not ammonium salt, sodium salt or potassium salt Reaction between: • Acid + Reactive metal • Acid + Metal oxide • Acid + Metal carbonate
Double decomposition reaction (precipitation reaction)
Salt solution A
Salt solution B
Excess of metal solid/ metal oxide/ metal carbonate Acid
Figure 6.38 Methods for preparing soluble and insoluble salts
180
Soluble salt solution
Precipitate of insoluble salt
Acid, Base and Salt
CHAPTER 6
Preparation of Soluble Ammonium, Sodium and Potassium Salts
Activity 6.19 Aim: To prepare soluble salts through a neutralisation reaction between an acid and an alkali. Materials: 2.0 mol dm–3 hydrochloric acid, HCl, 2.0 mol dm–3 potassium hydroxide solution, KOH, phenolphthalein indicator, filter papers and distilled water Apparatus: 250 cm3 beaker, glass rod, filter funnel, retort stand with clamp, 25 cm3 pipette, pipette filler, burette, evaporating dish, Bunsen burner, pipeclay triangle, conical flask, tripod stand, white tile and wash bottle Procedure:
V cm3 of hydrochloric acid, HCl
Hydrochloric acid, HCl
Potassium chloride solution, KCl
Pipeclay triangle Heat
Potassium hydroxide solution, KOH + phenolphtalein
Potassium chloride solution, KCl
Wash bottle Filter papers
Distilled water
Potassium chloride crystals, KCl
Saturated potassium chloride solution, KCl
Potassium chloride crystal, KCl
Dry potassium chloride crystals, KCl
Figure 6.39 Apparatus set-up to obtain potassium chloride crystals, KCl
1. Rinse a 25 cm3 pipette with a small amount of 2.0 mol dm–3 potassium hydroxide, KOH. Then, discard the solution. 2. Pipette accurately 25 cm3 of 2.0 mol dm–3 potassium hydroxide solution, KOH and transfer into a conical flask. 3. Add a few drops of phenolphthalein indicator and swirl the flask. 4. Rinse a burette with 2.0 mol dm–3 hydrochloric acid, HCl. Then, discard the solution. 5. Fill the burette with 2.0 mol dm–3 hydrochloric acid, HCl and clamp the burette to the retort stand. Record the initial reading of the burette. 181
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Interaction between Matter
6. Add acid into the conical flask slowly while swirling it. Safety Precaution 7. Continue adding acid until the colour of the solution in the To know whether the salt conical flask changes from pink to colourless. is saturated or not, dip the 8. Record the final reading of the burette. Then, determine glass rod into the solution the volume of 2.0 mol dm–3 hydrochloric acid, HCl required and then remove the glass to neutralise 25 cm3 of 2.0 mol dm–3 potassium hydroxide rod. If crystals are produced, a saturated solution has solution, KOH (assuming the acid volume is V cm3). been obtained. 3 –3 9. Refill a new 25 cm of 2.0 mol dm potassium hydroxide solution, KOH into a conical flask without phenolphthalein Glass rod indicator. Salt 10. Add V cm3 of 2.0 mol dm–3 hydrochloric acid, HCl from a crystal burette into the conical flask and swirl the mixture to ensure the mixture is even. 11. Pour the content of the conical flask into an evaporating dish. 12. Heat the solution slowly to evaporate the water so that a saturated solution is obtained. 13. Let the saturated salt solution cool down to allow crystallisation to occur. 14. Filter the contents of the evaporating dish to obtain potassium chloride crystals, KCl. 15. Rinse the crystals with a little amount of distilled water. 16. Dry the salt crystals by pressing them between two pieces of filter papers. Discussion: 1. Why is phenolphthalein indicator needed in titration? 2. Why should not phenolphthalein indicator be added to V cm3 of hydrochloric acid, HCl to 25 cm3 of potassium hydroxide solution, KOH? 3. Explain why the resulting crystals can only be rinsed with a little amount of distilled water. 4. Write a balanced chemical equation for this neutralisation reaction. 5. Give two other types of soluble salts that can be prepared by this method. Prepare a complete report after carrying out this activity. Preparation of Soluble Salts which are not Ammonium, Sodium and Potassium Salts
Activity 6.20 Aim: To prepare a soluble salt based on the reaction between an acid and a metal oxide. Materials: 2.0 mol dm–3 nitric acid, HNO3, copper(II) oxide powder, CuO, filter papers and distilled water Apparatus: 250 cm3 beaker, spatula, glass rod, filter funnel, evaporating dish, Bunsen burner, pipeclay triangle, conical flask, tripod stand, wash bottle, 20 cm3 measuring cylinder and retort stand with clamp
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Acid, Base and Salt
Procedure:
CHAPTER 6
Glass rod
Spatula Filter funnel
Excess of copper(II) oxide, CuO
Filter paper Excess of copper(II) oxide powder, CuO
20 cm3 of nitric acid, HNO3
Filter papers
Evaporating dish
Salt solution
Pipeclay triangle Heat
Copper(II) nitrate, Cu(NO3)2
Wash bottle
Salt crystals
Saturated solution
Salt crystals
Distilled water
Dry salt crystals
Figure 6.40 Apparatus set-up to obtain copper(II) nitrate crystal, Cu(NO3)2
1. Pour 20 cm3 of 2.0 mol dm–3 nitric acid, HNO3 into a beaker. Heat the acid using medium heat. 2. Add copper(II) oxide powder, CuO gradually into the acid using a spatula. Stir the mixture with a glass rod. 3. Continue adding copper(II) oxide, CuO until it is no longer dissolved. 4. Filter the excess copper(II) oxide powder, CuO from the mixture. 5. Pour the filtrate into an evaporating dish and heat the filtrate till a saturated salt solution is obtained. 6. Let the resulting saturated solution cool until salt crystals are formed. 7. Filter the content of the evaporating dish to obtain the salt crystals. Rinse the crystals with a little amount of distilled water. 8. Dry the salt crystals by pressing them between two pieces of filter papers.
Discussion: 1. Why is copper(II) oxide, CuO added in excess to the solution? 2. The filtration is done twice in this activity. Explain why. 3. Write a chemical equation for the reaction between nitric acid, HNO3 and copper(II) oxide, CuO. 4. Is the reaction between nitric acid, HNO3 and copper(II) oxide, CuO also considered a neutralisation reaction? Give a reason. Prepare a complete report after carrying out this activity.
183
THEME 3
Interaction between Matter
Activity 6.21 Aim: To prepare a soluble salt based on the reaction between an acid and a reactive metal. Materials: 2.0 mol dm–3 sulphuric acid, H2SO4, zinc powder, Zn, filter paper and distilled water Apparatus: 250 cm3 beaker, spatula, glass rod, filter funnel, evaporating dish, Bunsen burner, pipeclay triangle, conical flask, tripod stand, wash bottle and retort stand with clamp Procedure: 1. In pair, study Activity 6.20 on pages 182 and 183. Then, plan the procedures for the lab activity to prepare soluble salts of zinc sulphate, ZnSO4 based on the reaction between an acid and a metal. 2. Discuss with your teacher if you encounter any problem when planning the procedures.
Discussion: 1. Does zinc powder, Zn, have to be added in excess to the sulphuric acid, H2SO4? Why? 2. Write a chemical equation for the reaction between sulphuric acid, H2SO4 and the metal zinc, Zn. 3. Copper powder is not suitable to prepare copper(II) sulphate salt, CuSO4 by using the method in this activity. Give the reason why. Prepare a complete report after carrying out this activity.
Activity 6.22 Aim: To prepare a soluble salt based on the reaction between an acid and a metal carbonate.
Procedure: 1. In groups, determine a soluble salt that needs to be prepared. 2. Based on the chosen soluble salt, determine the materials and apparatus needed for this activity. 3. Plan and carry out the activity to prepare the soluble salt based on the reaction between an acid and an insoluble metal carbonate. Discussion: 1. Name the gas that is produced. 2. Describe the chemical test for the gas released. 3. Write the chemical equation involved. Prepare a complete report after carrying out this activity. Purification of Soluble Salts by the Recrystallisation Method
The soluble salt produced might contain impurities during preparation. Therefore, the soluble salt can be purified by the recrystallisation method. Carry out Activity 6.23 to find out how to purify a soluble salt. 184
Acid, Base and Salt
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Activity 6.23 Aim: To prepare a pure soluble salt through the recrystallisation method. Materials: Copper(II) sulphate crystals, CuSO4, filter papers and distilled water Apparatus: 250 cm3 beaker, spatula, glass rod, filter funnel, evaporating dish, Bunsen burner, wire gauze, pipeclay triangle, conical flask, tripod stand, wash bottle and retort stand with clamp Procedure: Distilled water
Glass rod
Glass rod Filter funnel
Salt solution Evaporating dish
Filter paper Insoluble impurities
Copper(II) sulphate crystals, CuSO4
Heat
Filter papers
Pipeclay triangle Heat
Filtrate
Salt crystals
Saturated solution Salt crystals
Distilled water
Dry salt crystals
Figure 6.41 Apparatus set-up to purify copper(II) sulphate salt, CuSO4
1. Put the copper(II) sulphate crystals, CuSO4 into a beaker. Safety Precaution 2. Add distilled water gradually while stirring. Heat the Make sure the distilled water solution to speed up the process of dissolving the salt. added is just enough to 3. Filter the hot salt solution to remove insoluble impurities. dissolve all the crystals. 4. Then, pour the filtrate into an evaporating dish and heat the filtrate until a saturated salt solution is obtained. 5. Let the saturated solution cool until salt crystals are formed. 6. Filter the contents of the evaporating dish to get the salt crystals. Rinse the crystals with a little amount of distilled water. 7. Dry the salt crystals by pressing them between two pieces of filter papers.
185
THEME 3
Interaction between Matter
Discussion: 1. What is the purpose of rinsing the crystals formed with distilled water? 2. State the method used to increase the size of the crystals formed. 3. Can recrystallisation be used to purify insoluble salts? Why? Prepare a complete report after carrying out the activity.
Preparation of Insoluble Salts
Insoluble salts can be prepared through double decomposition reaction. In this process, two salt solutions that contain insoluble salt ions are needed. Preparing Insoluble Salts by Double Decomposition Reaction
Activity 6.24 Aim: To prepare an insoluble salt by double decomposition reaction. Materials: 2.0 mol dm–3 sodium sulphate solution, Na2SO4, 2.0 mol dm–3 barium chloride solution, BaCl2, filter papers and distilled water Apparatus: 250 cm3 beaker, measuring cylinder, glass rod, filter funnel, evaporating dish, conical flask and wash bottle Procedure: Barium chloride solution, BaCl2
Glass rod
Sodium chloride solution, NaCl Filter paper
Sodium sulphate solution, Na2SO4
Filter funnel Precipitate of barium sulphate, BaSO4
Precipitate of barium sulphate, BaSO4 Filter papers
Dry insoluble salts
Sodium chloride solution, NaCl
Precipitate of barium sulphate, BaSO4
Figure 6.42 Apparatus set-up in the preparation of insoluble salt
1. Based on Figure 6.42, plan the activity with your group members to prepare the insoluble salt of barium sulphate, BaSO4. 2. Discuss with your teacher if you encounter any problems while planning the procedure. 3. Carry out the activity with your teacher̕s permission. 186
Acid, Base and Salt
CHAPTER 6
Discussion: 1. Write a balanced chemical equation for the preparation of barium sulphate salt, BaSO4. 2. Write an ionic equation for the preparation of barium sulphate salt, BaSO4. 3. Why must the filtered barium sulphate salt, BaSO4 precipitate rinsed with distilled water? 4. In your opinion, is it suitable to prepare barium sulphate salt, BaSO4 based on the reaction between sulphuric acid, H2SO4 and barium carbonate, BaCO3? Explain your answer. 5. Name two other salts that can be prepared by the double decomposition reaction. Then, suggest suitable aqueous solutions for the preparation of the salts mentioned. Prepare a complete report after carrying out this activity. In the double decomposition reaction, the ions in both aqueous solutions exchange with each other to form a new aqueous solution and a precipitate. The ionic equation for the formation of barium sulphate, BaSO4 can be derived from the balanced chemical equations as shown below:
Chemical equation: BaCl2(aq)
+
Na2SO4(aq)
Ba2+(aq) + 2Cl–(aq) + 2Na+(aq) + SO42–(aq) Ions in barium chloride
Ions in sodium sulphate
→
BaSO4(s)
+
→
BaSO4(s)
+ 2Na+(aq) + 2Cl–(aq)
Precipitate
2NaCl(aq)
Ions in sodium chloride
The Na+ ion and Cl– ion are spectator ions which do not take part in the reaction. Thus, these ions are cancelled out in the equation.
Ionic equation:
Ba2+(aq) + SO42–(aq) → BaSO4(s)
Activity 6.25 Writing the ionic equation for the formation of insoluble salts 1. Write the ionic equation for the following reactions: (a) The reaction between silver nitrate, AgNO3 and magnesium chloride, MgCl2 (b) Mixing potassium chromate(VI), K2CrO4 with lead(II) nitrate, Pb(NO3)2 (c) Copper(II) chloride solution, CuCl2 is added to sodium carbonate solution, Na2CO3
187
THEME 3
Interaction between Matter
Construction of Ionic Equations through the Continuous Variation Method
The continuous variation method is used to construct the ionic equation for the formation of insoluble salts. In this method, the volume of one solution, A is fixed, while solution B is added to the solution A by increasing the volume as shown in Figure 6.43. 1 cm3 2 cm3 3 cm3 4 cm3 5 cm3 6 cm3 7 cm3 8 cm3 Solution B
Solution A with fixed volume
Figure 6.43 Continuous variation method
The height of the precipitate formed increases gradually for the first few test tubes and then becomes constant as shown in Figure 6.44. The first test tube that achieves the maximum height of the Precipitate precipitate indicates that all reactants has completely reacted with one another.
Figure 6.44 Changes of the height of the precipitate
Experiment
6.6
Aim: To construct an ionic equation for the formation of lead(II) iodide. Problem statement: How to construct an ionic equation for the formation of lead(II) iodide? Hypothesis: As the volume of potassium iodide solution, KI added to lead(II) nitrate solution, Pb(NO3)2 increases, the height of the precipitate will increase and then remain constant. Variables: (a) Manipulated : Volume of potassium iodide solution, KI (b) Responding : Height of the precipitate (c) Fixed : Volume and concentration of lead(II) nitrate solution, Pb(NO3)2, concentration of potassium iodide solution, KI Materials: 0.5 mol dm–3 lead(II) nitrate solution, Pb(NO3)2, 0.5 mol dm–3 potassium iodide solution, KI and distilled water Apparatus: Test tubes of the same size, glass rod, test tube rack, burette, retort stand with clamp and ruler 186
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Acid, Base and Salt
Procedure: Safety Precaution 1. Label eight test tubes from 1 to 8 and place all the test tubes in a test tube rack. Make sure all the test tubes 3 used are of the same size. 2. Using a burette, fill each test tube with 5 cm of -3 0.5 mol dm lead(II) nitrate solution, Pb(NO3)2. 3. Using a second burette, add 0.5 mol dm-3 potassium iodide solution, KI into each test tube according to the volume stated in Table 6.7. 4. Place a glass rod into the test tube. Swirl the glass rod with both palms to ensure even mixing of the two solutions. 5. Slowly remove the glass rod. Rinse the precipitate that is stuck to the glass rod and the walls of the test tubes with distilled water. 6. Repeat steps 4 and 5 for the rest of the test tubes. 7. Leave the test tubes to stand for 30 minutes for the precipitate to settle to the bottom. 8. Record the colour of the precipitate formed and the solution on top of the precipitate. 9. Measure and record the height of the precipitate in each test tube. Results: Table 6.7
Test tube
1
2
3
4
5
6
7
8
Volume of lead(II) nitrate solution, Pb(NO3)2 (cm3)
5
5
5
5
5
5
5
5
Volume of potassium iodide solution, KI (cm )
1
2
3
4
5
6
7
8
3
Height of precipitate (cm) Colour of the solution on top of the precipitate
Interpreting data: 1. Plot a graph of the height of the precipitate against the volume of potassium iodide, KI. 2. From the graph, determine the volume of the potassium iodide solution, KI that completely reacts with 5 cm3 of lead(II) nitrate solution, Pb(NO3)2. 3. Calculate the number of moles for: (a) Lead(II) ions, Pb2+ in 5 cm3 of 0.5 mol dm–3 lead(II) nitrate solution, Pb(NO3)2 (b) Iodide ions, I– that reacts with 5 cm3 of 0.5 mol dm–3 lead(II) nitrate solution, Pb(NO3)2 4. Determine the number of moles of iodide ion, I– that reacts completely with 1 mol of lead(II) ion, Pb2+. 5. Based on your answer in questions 3 and 4, construct an ionic equation for the formation of the lead(II) iodide precipitate, PbI2. Conclusion: Is the hypothesis acceptable? What is the conclusion of the experiment? Discussion: 1. Why should the test tubes be of the same size? 2. Explain why the height of the precipitate increases gradually and then remains constant. Prepare a complete report after carrying out this experiment. 189
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Interaction between Matter
Brain Teaser
In the continuous variation method, fixing the volume of lead(II) nitrate solution, Pb(NO3)2 while manipulating the volume of potassium iodide, KI is to determine the mole ratio of lead(II) ions, Pb2+ that will react completely with the iodide ions, I–. If x mol of lead(II) ions, Pb2+ reacts with y mole of iodide ions, I–, then the empirical formula of the insoluble salt is PbxIy.
Brain Teaser If 2 moles of silver ions, Ag+ reacts with 1 mole of carbonate ions, CO32–, can you write the ionic equation for the formation of silver carbonate salt?
x(ion Pb2+) + y(ion I–) → PbxIy
TestYourself
6.9
1. Classify the following into soluble and insoluble salts. NaNO3
BaSO4
CaCO3
NaCl
Pb(NO3)2
MgSO4
K2CO3
AgCl
(NH4)2SO4
PbI2
BaCrO4
ZnCl2
2. Suggest suitable aqueous solutions for the preparation of calcium sulphate salt, CaSO4. Then, write an ionic equation for the formation of the salt. 3. Using a diagram, show how zinc nitrate crystals, Zn(NO3)2 can be prepared. In your diagram, include the reagents that are needed.
6.10
Effect of Heat on Salts
Based on the conversation in Figure 6.45, can you identify the anion based on the gas released? What is the gas released?
Limewater turns cloudy. The salt is a carbonate salt because carbon dioxide is released.
Figure 6.45 Carbon dioxide gas turns limewater cloudy
186
g Learnin tandard S At the end of the lesson, pupils are able to: 6.10.1 Describe briefly chemical tests to identify gases 6.10.2 Investigate the effects of heat on salts through experiment
Acid, Base and Salt
CHAPTER 6
Gas Tests
The process to identify a gas can be carried out by the gas test in Activity 6.26.
Activity 6.26 Aim: Identifying the gases released. Materials: Solid potassium chlorate(V), KClO3, dilute sulphuric acid, H2SO4, zinc powder, Zn, solid zinc carbonate, ZnCO3, dilute sodium hydroxide solution, NaOH, solid ammonium chloride, NH4Cl, solid manganese(IV) oxide, MnO2, concentrated hydrochloric acid, HCl, solid sodium chloride, NaCl, concentrated sulphuric acid, H2SO4, concentrated ammonia solution, NH3, solid sodium sulphite, Na2SO3, dilute hydrochloric acid, HCl, acidified potassium manganate(VII) solution, KMnO4, solid lead(II) nitrate, Pb(NO3)2, red litmus paper, blue litmus paper and limewater Apparatus: Test tubes, test tube holder, wooden splinter, rubber stopper with delivery tube, glass rod, spatula, tongs, Bunsen burner and 10 cm3 measuring cylinder Procedure: Carry out the test for gases and record your observations. Table 6.8 Gas test A: Test for oxygen gas, O2 1. Put two spatulas of solid potassium chlorate(V), KClO3 into a test tube. 2. Heat the solid with high heat. 3. Insert a glowing wooden splinter into the test tube.
Observation
Inference
Glowing wooden splinter
KClO3 Heat
B: Test for hydrogen gas, H2 1. Put some a few pieces of zinc powder, Zn into a test tube. 2. Add 4 cm3 of dilute sulphuric acid, H2SO4 into the test tube. 3. Place a lighted wooden splinter near the mouth of the test tube. C: Test for carbon dioxide gas, CO2 1. Put a spatula of solid zinc carbonate, ZnCO3 into a test tube. 2. Heat the solid with high heat. 3. Flow the gas produced into limewater.
This activity has to be carried out in the fume chamber.
Lighted wooden splinter H2SO4 Zn
ZnCO3 Heat
Limewater
191
THEME 3
Interaction between Matter
Gas test D: Test for ammonia gas, NH3 1. Put a spatula of solid ammonium chloride, NH4Cl into a test tube. 2. Add 4 cm3 of dilute sodium hydroxide solution, NaOH into the test tube. 3. Heat the mixture slowly. 4. Then, place a piece of moist red litmus paper to the mouth of the test tube.
Observation Moist red litmus paper NaOH NH4Cl Heat
E: Test for chlorine gas, Cl2 1. Put a spatula of powdered manganate(IV) oxide, MnO2 into a test tube. 2. Carefully add 2 cm3 of concentrated hydrochloric acid, HCl. 3. Heat the mixture slowly. 4. Then, place a piece of moist blue litmus Heat paper to the mouth of the test tube. F: Test for hydrogen chloride gas, HCl 1. Put a spatula of solid sodium chloride, NaCl into a test tube. 2. Add 2 cm3 of concentrated sulphuric acid, H2SO4 carefully. 3. Heat the mixture slowly. 4. Dip a glass rod into concentrated ammonia solution, NH3. 5. Then, hold the dipped glass rod to the mouth of the test tube. Heat G: Test for sulphur dioxide gas, SO2 1. Put a spatula of solid sodium sulphite, Na2SO3 into a test tube. 2. Add 4 cm3 of dilute Na2SO3 hydrochloric acid, HCl. 3. Heat the mixture slowly. 4. Flow the gas released Heat into acidified potassium manganate(VII) solution, KMnO4.
Moist blue litmus paper HCl MnO2
Glass rod with concentrated ammonia solution, NH3 H2SO4 NaCl
HCl Acidified potassium manganate(VII) solution, KMnO4
Note: Acidified potassium manganate(VII) solution, KMnO4 can be replaced with acidified potassium dichromate(VI) solution, K2Cr2O7.
186
Inference
Acid, Base and Salt
Gas test H: Test for nitrogen dioxide gas, NO2 1. Put a spatula of solid lead(II) nitrate, Pb(NO3)2 into a test tube. 2. Heat the mixture with high heat. 3. Then, place a piece of moist blue litmus paper to the mouth of the test tube.
Observation
CHAPTER 6
Inference
Moist blue litmus paper
Pb(NO3)2 Heat
Interpreting data: 1. Based on the observations, write the corresponding inference. 2. Why should the litmus paper be moistened before testing for the gases released? 3. Copy and complete Table 6.9 to summarise the method used for gas test. Table 6.9
Gas Oxygen gas, O2
Chemical tests Method Place a glowing wooden splinter into a test tube filled with the gas.
Observation The glowing wooden splinter rekindles.
Hydrogen gas, H2 Carbon dioxide gas, CO2 Ammonia gas, NH3 Chlorine gas, Cl2 Hydrogen chloride gas, HCl Sulphur dioxide gas, SO2 Nitrogen dioxide gas, NO2
Discussion: 1. What is the expected observation if a glass rod that is dipped into concentrated hydrochloric acid, HCl, is brought closer to the gas released in test D? 2. Name the white fumes formed in test F. 3. The gas released in test G is acidic. Predict the observation when a moist litmus paper is used. Prepare a complete report after carrying out this activity.
Effect of Heat
Most salts decompose when heated. By comparing the colour of the salt and the residue left behind and the gas released, we can identify the cation and anion that might be present in the salt. Experiment 6.7 investigates the action of heat on carbonate salt and nitrate salt.
193
THEME 3
Interaction between Matter
Experiment
6.7
Aim: To investigate the action of heat on carbonate salts. Problem statement: Do all carbonate salts decompose when heated to produce carbon dioxide gas? Hypothesis: All carbonate salts decompose when heated to produce carbon dioxide gas. Variables: (a) Manipulated : Types of carbonate salts (b) Responding : Products of decomposed carbonate salts (c) Fixed : Two spatulas of carbonate salts Materials: Solid sodium carbonate, Na2CO3, solid calcium carbonate, CaCO3, solid zinc carbonate, ZnCO3, solid lead(II) carbonate, PbCO3, solid copper(II) carbonate, CuCO3 and limewater Apparatus: Test tubes, boiling tubes, test tube holder, Bunsen burner and rubber stopper with delivery tube Procedure: Carbonate Delivery salt tube 1. Place two spatulas of solid sodium carbonate, Na2CO3 into a dry boiling tube. Observe the colour of salt and Heat record the observation. 2. Connect the rubber stopper with the delivery tube to the mouth of the boiling tube. Ensure that the other end of Limewater the delivery tube is placed into the limewater as shown Figure 6.46 in Figure 6.46. 3. Heat the carbonate salt with high heat. 4. Observe the changes that occur in the limewater and the colour of the residue in the boiling tube when it is hot and when it is cool. Record the observation. 5. Repeat steps 1 to 4 using other carbonate salts to replace sodium carbonate salt, Na2CO3. Results: Table 6.10 Carbonate salt Sodium carbonate, Na2CO3 Calcium carbonate, CaCO3 Zinc carbonate, ZnCO3 Lead(II) carbonate, PbCO3 Copper(II) carbonate, CuCO3
186
Colour of the salt before heating
Colour of the residue When hot
When cool
Effect on limewater
Acid, Base and Salt
CHAPTER 6
Interpreting data: 1. What is the function of the limewater in the experiment? 2. What gas is released when carbonate salts are decomposed by heat? 3. Identify the carbonate salts that cannot be decomposed by heat. Conclusion: Is the hypothesis acceptable? What is the conclusion of the experiment? Discussion: 1. Write an equation for the action of heat on carbonate salts that can be decomposed by heat. 2. Generally, the decomposition of carbonate salts by heat can be represented by the word equation below. Complete the word equation.
Heat
+
3. Name another carbonate salt that cannot be decomposed by heat. Prepare a complete report after carrying out this experiment.
Experiment
6.8
Aim: To investigate the action of heat on nitrate salts. Problem statement: Do nitrate salts decompose when heated to release nitrogen dioxide gas and oxygen gas? Hypothesis: Propose a suitable hypothesis for the experiment. Variables: List all the variables involved in the experiment. Materials: Solid sodium nitrate, NaNO3, solid magnesium nitrate, Mg(NO3)2, solid zinc nitrate, Zn(NO3)2, solid lead(II) nitrate, Pb(NO3)2, solid copper(II) nitrate, Cu(NO3)2, wooden splinter and blue litmus paper Apparatus: Boiling tube, test tube holder, Bunsen burner and spatula Procedure: 1. Discuss the experimental procedure with your group members. Your discussion should include chemical tests for the Moist blue Glowing gases released. litmus paper wooden 2. Identify the safety precautions when carrying out splinter the experiment. 3. Carry out the experiment with your teacher̕s permission. 4. Record the salt colour before heating, the colour of the residue Nitrate salt when it is hot and cool, the colour of the gas released and its Heat effect on the moist blue litmus paper and the glowing wooden splinter in a table. Figure 6.47 Apparatus set-up for heating a nitrate salt Interpreting data: 1. Identify the nitrate salts that are not decomposed by heat. 2. Name the brown coloured gas released in the experiment. 3. What are the changes observed on the glowing wooden splinter? State the corresponding inference. 195
THEME 3
Interaction between Matter
Conclusion: Is the hypothesis acceptable? What is the conclusion of the experiment? Discussion: 1. Write equations for the actions of heat on the nitrate salts that decompose other than sodium nitrate. 2. Generally, the thermal decomposition of nitrate is represented by the word equation below. Complete the equation.
Heat
+
+
Prepare a complete report after carrying out this experiment. When heated, most carbonate salts decompose to produce metal oxides and carbon dioxide gas. For example, decomposition of zinc carbonate salt, ZnCO3 produces zinc oxide, ZnO and carbon dioxide gas, CO2. ZnCO3(s) → ZnO(s) + CO2(g) When heated, nitrate salts decompose to produce metal oxides, nitrogen dioxide gas and oxygen gas. The following equation shows the thermal decomposition of lead(II) nitrate salt, Pb(NO3)2. 2Pb(NO3)2(s) → 2PbO(s) + 4NO2(g) + O2(g) Only several sulphate salts and chloride salts can be decomposed when heated. Ammonium chloride : NH4Cl(s) → NH3(g) + HCl(g) Zinc sulphate : ZnSO4(s) → ZnO(s) + SO3(g) Iron(II) sulphate : 2FeSO4(s) → Fe2O3(s) + SO2(g) + SO3(g) The cation or anion of a salt can be identified based on the gases released when the salt undergoes thermal decomposition. Carbon dioxide gas → Carbonate salt Nitrogen dioxide gas + Oxygen gas → Nitrate salt Ammonia gas → Ammonium salt
Chemistry The cation presents in some salts can be identified from the colour of the residue after heating. Colour of residue
196
Metal oxide
Cation present in the salt
Hot
Cold
Yellow
White
Zinc oxide, ZnO
Zinc ion, Zn2+
Brown
Yellow
Lead(II) oxide, PbO
Lead(II) ion, Pb2+
Black
Black
Copper(II) oxide, CuO
Copper(II) ion, Cu2+
Acid, Base and Salt
CHAPTER 6
Example 16
Salt X decomposes when heated. A brown coloured gas is released and turns moist blue litmus paper to red. The colour of the residue is brown when it is hot and yellow when it is cool. Name salt X. Solution:
The brown coloured gas is nitrogen dioxide. Salt X contains nitrate ion, NO3–. The residue is lead(II) oxide, PbO. Salt X contains lead(II) ion, Pb2+. Salt X is lead(II) nitrate, Pb(NO3)2.
TestYourself
6.10
1. Identify the colourless gas that turns the colour of acidified potassium dichromate(VI) solution, K2Cr2O7 from orange to green. 2. Copper(II) nitrate salt, Cu(NO3)2 is heated with high heat in a boiling tube. State the observations on the colour of the residue and the gas released. 3. The following results are obtained in a laboratory activity to investigate the effect of heat on a sample of salt Y. • Colourless gas turns limewater cloudy. • The colour of the residue is yellow when it is hot and white when it is cooled. Identify salt Y.
6.11
Qualitative Analysis
Qualitative Analysis to Identify Cations and Anions in Salts
Qualitative analysis of a salt is a technique used to identify the cation and anion present in a salt by analysing its physical and chemical properties. Figure 6.48 shows the steps that are involved in the qualitative analysis of a salt. Observation on the physical properties of the salt
Effect of heat on the salt
Test for anions and cations
Confirmatory tests for cations and anions
Figure 6.48 Steps in the qualitative analysis of a salt
g Learnin tandard S At the end of the lesson, pupils are able to: 6.11.1 Identify the cation and anion present in a salt through experiment 6.11.2 Describe the confirmatory tests to identify cations and anions
197
THEME 3
Interaction between Matter
Observations of the Physical Properties of Salts Observing the physical properties of salts such as colour and solubility in water is the first step to make inferences on the possibility of the presence of cations and anions in the salt. Even though the solubility test of the salts does not confirm the identities of the ions present, it helps us narrow down the possible identities of the ions present. For example, salt X is soluble in water. Hence, salt X might contain NO ASaP ions, but definitely not HPA chloride or PBC sulphate and most probably does not contain the carbonate ion, CO32–.
The colour of the salt is one of the physical properties that enables us to make inference on the cation present in the salt. How is this possible? Carry out Activity 6.27 below.
Activity 6.27 Aim: To investigate the colour of the salts and their solubility in water. Materials: Solid ammonium nitrate, NH4NO3, solid potassium nitrate, KNO3, solid sodium chloride, NaCl, solid calcium carbonate, CaCO3, solid calcium nitrate, Ca(NO3)2, solid magnesium sulphate, MgSO4, solid magnesium carbonate, MgCO3, solid zinc sulphate, ZnSO4, solid zinc chloride, ZnCl2, solid iron(II) sulphate, FeSO4, solid iron(III) chloride, FeCl3, solid lead(II) nitrate, Pb(NO3)2, solid lead(II) chloride, PbCl2, solid lead(II) sulphate, PbSO4, solid copper(II) sulphate, CuSO4, solid copper(II) chloride, CuCl2, solid copper(II) nitrate, Cu(NO3)2, solid copper(II) carbonate, CuCO3 and distilled water Apparatus: Test tubes, test tube rack, glass rod, spatula and wash bottle Procedure: 1. Observe and record the colour of each solid salt. 2. Put a little amount of the salt into a test tube. Fill the test tube with distilled water and stir the mixture. 3. Observe the solubility of the salt and the colour of the solution formed. 4. Record all observations in Table 6.11. Results: Table 6.11
Type of salt
Colour of the salt
Solubility in water Yes No
Colour of the salt solution
Ammonium nitrate, NH4NO3 Potassium nitrate, KNO3
Interpreting data: 1. Name solid salts that are coloured: (a) Green (b) Brown (c) Blue 2. Name green salts that are: (a) Insoluble in water (b) Soluble in water to form a blue solution (c) Soluble in water to form a light green solution 196
(d) White
Acid, Base and Salt
CHAPTER 6
Discussion: 1. Is the colour of the salt suitable to identify the cation presents in the salt? Explain. 2. What is the colour of the solution formed when the white salt dissolves in water? 3. Classify each salt as soluble salts or insoluble salts. 4. What is the operational definition of insoluble salts? Prepare a complete report after carrying out this activity. All white soluble salts dissolve in water to form colourless salt solutions. Salts form coloured solutions due to the presence of ions in transition elements. For example, Blue solution : possibly contains copper(II) ion, Cu2+ Brown solution : possibly contains iron(III) ion, Fe3+ Green solution : possibly contains iron(II) ion, Fe2+
Table 6.12 shows the colour of some salts in the solid state and in aqueous solution.
Chemistry Other than iron(II) ion, Fe2+, nickel(II) ion, Ni2+ and chromium(III) ion, Cr3+ also give green colour in aqueous solutions.
Table 6.12 Colour of salts in solid state and in aqueous solution Colour
Salt
Solid state
Aqueous solution
Green
Green/light green
Salts containing iron(III), Fe3+
Brown
Brown/yellowish brown
Copper(II) sulphate, CuSO4 Copper(II) nitrate, Cu(NO3)2
Blue
Blue
Copper(II) chloride salt, CuCl2
Green
Blue
Copper(II) carbonate salt, CuCO3
Green
Insoluble in water
Salts containing iron(II) ion, Fe
2+
Figure 6.49 shows the example on the qualitative analysis of a solid X which is green in colour and three possible results. Solid X in green colour Add distilled water Glass rod
Glass rod
Glass rod
Green precipitate
Green colour solution
Blue colour solution
Solid X might be CuCO3 salt or iron(II) salt which is insoluble in water.
Solid X might be iron(II) salt which is soluble in water.
Solid X might be CuCl2 salt which is soluble in water.
Figure 6.49 Qualitative analysis based on the solubility of the salt and its colour
197
THEME 3
Interaction between Matter
Effect of Heat on Salts and Gas Tests The gas released when a salt is heated can be identified through its colour, smell, action on moist litmus paper or best, by conducting a gas test. After heating the salt, we can make an inference on the ions that might be present based on the colour of the residue and the gas identified, as shown in Table 6.13 and Table 6.14.
Table 6.13 Colour of residue
Table 6.14 Inference
Black
Salt contains Cu2+ ion
Brown
Salt contains Fe3+ ion
Yellow when hot, white when cool
Salt contains Zn2+ ion
Brown when hot, yellow when cool
Salt contains Pb2+ ion
Gas produced
Inference
Gas turns limewater cloudy
Carbonate salt
Gas is brown and acidic
Nitrate salt
Gas is pungent and alkaline
Ammonium salt
Gas ignites the glowing wooden splinter
Might be nitrate salt or Ag2CO3
Figure 6.50 and 6.51 show the examples of qualitative analysis based on the effect of heat on salts X and Y and the corresponding gas tests. However, the qualitative analysis carried out could not identify the presence of cation in the salts. Salt X is not Fe2+, Fe3+ or Cu2+
Salt X is definitely not Zn2+, Fe2+, Fe3+, Pb2+ or Cu2+
Salt X is white in colour
The colour of the residue is white when it is hot and cool
Heated with high heat
Add water
White solid insoluble in water Salt X is definitely an insoluble salt
+
Colourless gas Pass through limewater
Limewater turns cloudy Carbon dioxide gas released. Salt X contains carbonate ion, CO32-
Figure 6.50 Qualitative analysis based on the effects of heat on salt X and gas test Salt X is a carbonate salt that possibly contains Ca2+ ion, Mg2+ ion or Al3+ ion and not K+ ion or Na+ ion because potassium carbonate and sodium carbonate are not decomposed by heat.
196
Acid, Base and Salt
Salt Y is not Fe2+, Fe3+ or Cu2+ Salt Y is white in colour Add water
CHAPTER 6
Salt Y is definitely not Zn2+, Fe2+, Fe3+, Pb2+ or Cu2+ Heated with high heat
The colour of the residue is white when it is hot and cool
Brown gas
+
+
Colourless gas
Test with glowing wooden splinter
Colourless solution
The glowing wooden splinter rekindles
Salt Y is a soluble salt
Brown nitrogen dioxide gas and oxygen gas present. Salt Y contains nitrate ion, NO3–
Figure 6.51 Qualitative analysis based on the effects of heat on salt Y and gas tests Salt Y is a nitrate salt that possibly contains Ca2+ ion, Mg2+ ion or Al3+ ion and not K+ ion and Na+ ion because potassium nitrate and sodium nitrate do not produce a brown gas.
Anion Tests Only four anions are needed to be identified at this level, which are: Carbonate ion, CO32– Sulphate ion, SO42–
Chloride ion, Cl– Nitrate ion, NO3–
Certain anions can be identified from testing the gases released when the salt undergoes decomposition by heat. However, the identity of the anion in a salt still needs to be confirmed through anion tests. Figure 6.52 shows the flow chart of the anion tests. Anion tests
CO32–
Cl–
+ Dilute acid
Gas turns limewater cloudy
SO42–
NO3–
+ Dilute HNO3 + Dilute HCl + AgNO3(aq) + BaCl2(aq)
White precipitate
White precipitate
Nitric acid, HNO3 and barium nitrate solution, Ba(NO3)2 can replace hydrochloric acid, HCl and barium chloride solution, BaCl2 to test for the presence of sulphate ion, SO42–.
+ Dilute H2SO4 + FeSO4(aq) + Concentrated H2SO4
Brown ring
Figure 6.52 Summary of the anion tests
197
THEME 3
Interaction between Matter
The test for anions in an aqueous solution of the salt can be performed as in Experiment 6.9.
Experiment
6.9
Aim: To identify the anions present in aqueous salt solutions. Problem statement: How to identify the anions that are present in aqueous solutions? Hypothesis: The anions present can be identified through observations from the chemical tests on anions. Variables: (a) Manipulated : Types of anions present in the solution (b) Responding : Observations made (c) Fixed : Volume of aqueous salt solution Materials: 2.0 mol dm–3 nitric acid, HNO3, 0.1 mol dm–3 silver nitrate solution, AgNO3, 2.0 mol dm–3 hydrochloric acid, HCl, 1.0 mol dm–3 barium chloride solution, BaCl2, 1.0 mol dm–3 sulphuric acid, H2SO4, 1.0 mol dm–3 iron(II) sulphate solution, FeSO4, concentrated sulphuric acid, H2SO4, sample of salt A (solid sodium carbonate, Na2CO3), sample of salt B (solid sodium chloride, NaCl), sample of salt C (solid sodium sulphate, Na2SO4), sample of salt D (solid sodium nitrate, NaNO3), distilled water and limewater Apparatus: Test tubes, test tube holder, test tube rack, glass rod, dropper, rubber stopper with delivery tube, spatula, 100 cm3 beaker, 10 cm3 measuring cylinder Procedure: Preparing aqueous solutions of the salts 1. Put salt sample A provided by the teacher into a beaker. 2. Dissolve salt sample A with distilled water to produce 20 cm3 of salt A solution. 3. Pour 2 cm3 of salt A solution into 4 test tubes. Label the test tubes as A1, A2, A3 and A4. 4. Repeat steps 1 until 3 using salt samples B, C and D. (I) Test for carbonate ion, CO32– 1. Add 2 cm3 of 2.0 mol dm-3 hydrochloric acid, HCl into the test tube labelled A1. If effervescence occurs, flow the gas into limewater as shown in Figure 6.53. 2. Record the observation. 3. Repeat steps 1 and 2 using solutions B1, C1 and D1. (II) Test for chloride ion, Cl– 1. Add excess of 2.0 mol dm–3 nitric acid, HNO3 into the test tube labelled A2, followed by 2 cm3 of 0.1 mol dm–3 silver nitrate solution, AgNO3. 2. Record the observation. 3. Repeat steps 1 and 2 using solutions B2, C2 and D2.
202
A1 solution
Limewater
Figure 6.53
Be careful when using silver nitrate solution, AgNO3. Skin that comes into contact with silver nitrate solution, AgNO3 would turn brown. The brown spots would disappear only after a few days.
Acid, Base and Salt
CHAPTER 6
(III) Test for sulphate ion, SO42– 1. Add 2.0 mol dm–3 hydrochloric acid, HCl into the test tube labelled A3, followed by 2.0 mol dm-3 barium chloride solution, BaCl2. 2. Record the observation. 3. Repeat steps 1 and 2 with solutions B3, C3 and D3.
(IV) Test for nitrate ion, NO3– 1. Add 2 cm3 of 1.0 mol dm–3 sulphuric acid, H2SO4 into the test tube labelled A4, followed by 2 cm3 of 1.0 mol dm–3 iron(II) sulphate solution, FeSO4. 2. Shake the mixture. 3. Carefully, drip a few drops of concentrated sulphuric acid, H2SO4 slowly down the wall of the tilted test tube as shown in Figure 6.54. Concentrated 4. Slowly set the test tube upright. sulphuric acid, 5. Record the observation. H2SO4 A4 solution 6. Repeat steps 1 to 5 using solutions B4, C4 and D4. Figure 6.54 Results: Table 6.15
Observation
Test
Salt A solution
Salt B solution
Salt C solution
Salt D solution
Carbonate ion, CO3
2–
Chloride ion, Cl– Sulphate ion, SO42– Nitrate ion, NO3–
Interpreting data: 1. What was the gas released that causes the effervescence? 2. (a) Name the white precipitate formed in the test for chloride ion, Cl–. (b) Write the ionic equation for the formation of the precipitate in 2(a). 3. (a) What is the name of the white precipitate formed in the test for sulphate ion, SO42–. (b) Write the ionic equation for the formation of the white precipitate in 3(a). 4. Identify the anion in the following samples: (a) Salt A (b) Salt B (c) Salt C (d) Salt D Conclusion: Is the hypothesis accepted? What is the conclusion of the experiment? Discussion: 1. What is the purpose of adding excess acid in the test for chloride ion, Cl– and sulphate ion, SO42– before adding the other reagents? 2. A student performs a test for the chloride ion, Cl– in a sample of salt solution. He added the silver nitrate solution, AgNO3 without first adding an excess of nitric acid, HNO3. He then made the inference of the presence of chloride ions, Cl– in the sample when he observed a white precipitate being formed. Was the inference correct? Why? Prepare a complete report after carrying out this experiment. 203
THEME 3
Interaction between Matter
Based on the test for the carbonate ion, CO32– the reaction between acid and carbonate ions, CO32– produces carbon dioxide gas that turns limewater cloudy.
Chemistry
2H+(aq) + CO32–(aq) → CO2(g) + H2O(l)
Reaction between In the test for chloride ion, Cl–, silver ion, Ag+ is used to detect concentrated sulphuric – – the presence of chloride ions, Cl . If chloride ions, Cl is present, acid, H2SO4 and nitrate ion, hence the white precipitate of silver chloride, AgCl is produced. NO3– produces nitrogen However, the carbonate ion, CO32– also gives the same observation monoxide, NO. When when reacting with silver ion, Ag+ due to the formation of a white nitrogen monoxide, NO combines with iron(II) precipitate of silver carbonate, Ag2CO3. Therefore, an excess of sulphate, FeSO4, a complex nitric acid, HNO3 has to be added before adding silver nitrate compound of nitrosyliron(II) solution, AgNO3. If effervescence occurs when nitric acid, HNO3 sulphate, FeSO4.NO, which was added, then, the presence of carbonate ion, CO32– is is a brown ring, is observed. confirmed. However, if no effervescence occurs, the formation of white precipitate confirms the presence of the chloride ion, Cl–. For the test of sulphate ion, SO42–, barium ion, Ba2+ used to detect the presence of sulphate ion, SO42– because the reaction between the barium ion, Ba2+ and sulphate ion, SO42– produces a white precipitate of barium sulphate, BaSO4. Hydrochloric acid, HCl was added in excess before adding barium chloride solution, BaCl2 for the same reason as in the test for chloride ion, Cl-, that is to detect and eliminate the carbonate ion, CO32– that might be present.
Table 6.16 Qualitative analysis based on anion tests
Anion test
Observation
Test for carbonate ion, CO32– Dilute acid Salt solution
Test for chloride ion, Cl– AgNO3 (aq)
HNO3 (aq)
Effervescence occurs. The gas released turns limewater cloudy.
Carbonate ion, CO32– is present. Ionic equation: 2H+(aq) + CO32–(aq) → CO2(g) + H2O(l)
White precipitate is formed.
Chloride ion, Cl– is present. Ionic equation: Ag+(aq) + Cl–(aq) → AgCl(s)
White precipitate is formed.
Sulphate ion, SO42– is present. Ionic equation: Ba2+(aq) + SO42–(aq) → BaSO4(s)
Salt solution
Test for sulphate ion, SO42– HCl (aq)
BaCl2 (aq)
Salt solution
Test for nitrate ion, NO3–
Salt solution
202
Dilute H2SO4
Concentrated Brown ring is formed. H2SO4
FeSO4 (aq)
Inference
Nitrate ion, NO3– is present.
Acid, Base and Salt
CHAPTER 6
Cation Tests Alkalis such as sodium hydroxide solution, NaOH, and ammonia solution, NH3 are two main reagents used to test the presence of cations. The hydroxide ion, OH– from both solutions combine with most metal ions to form a precipitate of metal hydroxide. For example,
Zn2+(aq) + 2OH–(aq) → Zn(OH)2(s)
Metal ion
Hydroxide ion
Metal hydroxide precipitate
Inference about the cations present can be made based on the observation on the colour of the precipitate and its solubility in an excess of alkali solution.
Experiment
6.10
Aim: To identify the cations present in aqueous solutions. Problem statement: How to identify the cations present in aqueous solutions? Hypothesis: Types of cations present in a solution can be identified through observations of the cation tests. Variables: (a) Manipulated : Types of cations present in aqueous solutions (b) Responding : Observations made (c) Fixed : Volume of aqueous salt solution Materials: 2.0 mol dm–3 sodium hydroxide solution, NaOH, 2.0 mol dm–3 ammonia solution, NH3, 1.0 mol dm–3 calcium nitrate solution, Ca(NO3)2, 1.0 mol dm–3 magnesium nitrate solution, Mg(NO3)2, 1.0 mol dm–3 aluminium nitrate solution, Al(NO3)3, 1.0 mol dm–3 zinc nitrate solution, Zn(NO3)2, 1.0 mol dm–3 iron(II) sulphate solution, FeSO4, 1.0 mol dm–3 iron(III) chloride solution, FeCl3, 1.0 mol dm–3 lead(II) nitrate solution, Pb(NO3)2, 1.0 mol dm-3 copper(II) sulphate solution, CuSO4 and 1.0 mol dm–3 ammonium nitrate solution, NH4NO3 Apparatus: Test tubes, test tube holder, test tube rack, dropper, 100 cm3 beaker, red litmus paper, Bunsen burner and 10 cm3 measuring cylinder Moist red Procedure: litmus paper
Add a few drops of aqueous alkali
No precipitate
Shake the mixture
Heat
Add aqueous alkali until excess
2 cm3 of salt solution White/coloured precipitate
Precipitate dissolves in excess aqueous alkali Shake it evenly
Figure 6.55 Steps in identifying the cations in aqueous solutions
Precipitate does not dissolve in excess aqueous alkali
205
THEME 3
Interaction between Matter
Using sodium hydroxide solution, NaOH 1. Based on Figure 6.55, discuss the experimental procedures with your group members. 2. Determine the safety precautions needed in carrying out the experiment. 3. Carry out the experiment with your teacher̕s permission. 4. Record all the observations in Table 6.17. Using ammonia solution, NH3 1. Repeat the procedure in A by using ammonia solution, NH3 to replace sodium hydroxide solution, NaOH. 2. Record all the observations in Table 6.17. Results: Table 6.17
Salt solution
Cation
Small amount of sodium hydroxide solution, NaOH
Observation Excess amount Small amount Excess amount of sodium of ammonia of ammonia hydroxide solution, NH3 solution, NH3 solution, NaOH
Calcium nitrate, Ca(NO3)2
Magnesium nitrate, Mg(NO3)2 Aluminium nitrate, Al(NO3)3
Zinc nitrate, Zn(NO3)2
Iron(II) sulphate, FeSO4
Iron(III) chloride, FeCl3
Lead(II) nitrate, Pb(NO3)2
Copper(II) sulphate, CuSO4
Ammonium nitrate, NH4NO3
Interpreting data: Based on the experiment in part A: 1. List the cations that produce precipitates which are: (a) Green (b) Brown (c) Blue (d) White 2. Which salt solution does not show any changes on the addition of sodium hydroxide solution, NaOH. What is the gas released upon heating? 3. Name the cations that produce white precipitates which are: (a) Soluble in an excess amount of sodium hydroxide solution, NaOH (b) Insoluble in an excess amount of sodium hydroxide solution, NaOH 202
Acid, Base and Salt
CHAPTER 6
Based on the experiment in part B: 1. List the cations that produce precipitates which are: (a) Green (c) Blue (b) Brown (d) White 2. Which salt solution does not show any changes on the addition of ammonia solution, NH3? 3. Name the cations that produce white precipitates which are: (a) Soluble in excess amount of ammonia solution, NH3 (b) Insoluble in excess amount of ammonia solution, NH3 4. Which cations form a precipitate that is soluble in excess amount of ammonia, NH3 to produce a dark blue solution? Conclusion: Is the hypothesis accepted? What is the conclusion of this experiment? Discussion: 1. Identify the cations that react with both alkalis to produce white precipitates that are soluble in an excess amount of the alkali. 2. Based on the cations identified in question 1, write ionic equations for the formation of the precipitates. Prepare a complete report after carrying out this experiment. Figure 6.56 shows the flow chart that summarises the reactions between cations and sodium hydroxide solution, NaOH.
Add a few drops of NaOH
Moist red litmus paper turns blue No
Does precipitate form?
Salt solution
NH3 gas
Yes Does coloured precipitate form?
No Dissolve in excess of NaOH solution
Yes
Zn2+
Al3+
Pb2+
Soluble in excess amount of NaOH solution to produce colourless solution
Heat Yes
Presence of NH4+ ion
No
Mg2+
Ca2+
Insoluble in excess amount of NaOH solution
Fe2+
Fe3+
Cu2+
Insoluble in excess amount of NaOH solution
Figure 6.56 Reactions between cations and sodium hydroxide solution, NaOH
207
THEME 3
Interaction between Matter
Figure 6.57 shows the flow chart that summarises the reaction between cations and ammonia solution, NH3.
Add a few drops of NH3 solution Does precipitate form?
Salt solution
Ca2+ ion or NH4+ ion might be present
No
Yes Does coloured precipitate form?
No
Yes
Soluble in excess amount of NH3 solution
Zn
No
Mg
2+
Soluble in excess amount of NH3 solution to form colourless solution
Yes
2+
Pb
2+
Fe2+
Al
3+
Insoluble in excess amount of NH3 solution
Fe3+
Insoluble in excess amount of NH3 solution
Cu2+ Soluble in excess amount of NH3 solution to produce dark blue solution
Figure 6.57 Reaction between cations and ammonia solution, NH3
Example 17
Salt solution X is blue. Test for cations with sodium hydroxide solution, NaOH:
Salt solution X
A few drops of NaOH
Add NaOH solution until excess Blue precipitate forms
Blue precipitate does not dissolve in excess amount of NaOH solution
Cu2+ ion is present.
Test for cations with ammonia solution, NH3:
Salt solution X
202
A few drops of NH3
Add NH3 solution until excess Blue precipitate forms
Blue precipitate dissolves to form dark blue solution
Acid, Base and Salt
CHAPTER 6
Example 18
Salt solution X is colourless. Test for cations with sodium hydroxide solution, NaOH: A few drops of NaOH
Salt solution X
Add NaOH solution until excess White precipitate forms Zn2+, Al3+, Pb2+, Mg2+, Ca2+ might be present
White precipitate does not dissolve in excess amount of NaOH solution Mg2+, Ca2+ might be present
Test for cations with ammonia solution, NH3: Ca2+ ion is present.
A few drops of NH3
Salt solution X
No change Ca2+, NH4+ might be present
Example 19
Salt solution X is colourless. Test for cations with sodium hydroxide solution, NaOH:
Salt solution X
A few drops of NaOH
Add NaOH solution until excess White precipitate forms Zn2+, Al3+, Pb2+, Mg2+, Ca2+ might be present
Zn2+, Al3+, Pb2+ might be present
Test for cations with ammonia solution, NH3:
Salt solution X
A few drops of NH3
Al3+ or Pb2+ ions might be present. Add NH3 solution until excess
White precipitate forms Zn2+, Al3+, Pb2+, Mg2+ might be present
White precipitate dissolves in excess amount of NaOH solution
White precipitate does not dissolve in excess amount of NH3 solution Al3+, Pb2+,Mg2+ might be present
209
THEME 3
Interaction between Matter
Confirmatory Tests for Cations Both Al3+ ion and Pb2+ ion give the same observations when tested with sodium hydroxide solution, NaOH and ammonia solution, NH3. Thus, a confirmatory test is required to differentiate between Pb2+ ion and Al3+ ion. Other than that, ions such as Fe2+, Fe3+ and NH4+ too can be confirmed using specific reagents.
Experiment
6.11
Aim: To confirm the presence of cations (NH4+, Fe2+, Fe3+, Pb2+) in aqueous solutions. Problem statement: How to confirm the presence of cations (NH4+, Fe2+, Fe3+, Pb2+) in aqueous solutions? Hypothesis: Construct a hypothesis which is suitable for this experiment. Variables: Name all variables involved in this experiment. Procedure: Method I: Confirmation of ammonium ion, NH4+
Add a few drops of potassium hexacyanoferrate(III) solution, K3Fe(CN)6
Add a few drops of Nessler reagent
2 cm3 of ammonium chloride solution, NH4Cl
Method II: Confirmation of Method III: Confirmation of iron(II) ion, Fe2+ iron(III) ion, Fe3+
2 cm3 of iron(II) sulphate solution, FeSO4
Add a few drops of potassium hexacyanoferrate(III) solution, K4Fe(CN)6 2 cm3 of iron(III) chloride solution, FeCl3
Method IV: Confirmation of lead(II) ion, Pb2+ 1 cm3 of potassium iodide solution, KI
2 cm 3 of lead(II) nitrate solution, Pb(NO3)2
3 cm3 of distilled water
Heat
Cooled
Figure 6.58 Confirmatory tests for ammonium ion, iron(II) ion, iron(III) ion and lead(II) ion
1. Based on Figure 6.58, list the apparatus and reagents that are needed for this experiment. 2. Discuss the experimental procedure for the experiment with your group members. Make sure you repeat Method III by using potassium thiocyanate solution, KSCN instead of potassium hexacyanoferate(II), K4Fe(CN)6. 3. Carry out the experiment with your teacher̕s permission. 4. Record your observations in Table 6.18. 210
Acid, Base and Salt
Results:
CHAPTER 6
Table 6.18 Confirmatory test
Observation
Method I: Confirmatory test for ammonium ion, NH4+ Method II: Confirmatory test for iron(II) ion, Fe2+ Method III: Confirmatory test for iron(III) ion, Fe3+ (a) Using potassium hexacyanoferate(II) solution, K4Fe(CN)6 (b) Using potassium thiocyanate solution, KSCN Method IV: Confirmatory test for lead(II) ion, Pb2+
Conclusion: Is the hypothesis accepted? What is the conclusion of this experiment? Discussion: 1. Write an ionic equation for the reaction between potassium iodide solution, KI and lead(II) nitrate solution, Pb(NO3)2. 2. Other than Nessler reagent, what are other reagents that can be used to confirm the presence of ammonium ion, NH4+? Briefly explain how the chemical test is carried out. 3. Predict the observation obtained if potassium chloride solution, KCl is used instead of potassium iodide solution, KI in Method IV. Prepare a complete report after carrying out this experiment. Confirmatory tests for NH4+, Fe2+ and Fe3+
Ammonium ion, NH4+ + Nessler reagent
Brown precipitate
Iron(II) ion, Fe2+ + K3Fe(CN)6
Dark blue precipitate
Iron(III) ion, Fe3+ + K4Fe(CN)6
Dark blue precipitate
+ KSCN
Blood red solution
Figure 6.59 Confirmatory tests for NH4+ ion, Fe2+ ion and Fe3+ ion
Chemistry • •
Potassium hexacyanoferate(II) solution, K4Fe(CN)6 is used to confirm the presence of iron(III) ion, Fe3+. If this solution is added to an aqueous solution that contains iron(II) ion, Fe2+, a light blue precipitate is formed. Potassium hexacyanoferate(III) solution, K3Fe(CN)6 is used to confirm the presence of iron(II) ion, Fe2+. If this solution is added to an aqueous solution containing iron(III), Fe3+, a greenish brown colour is obtained.
211
THEME 3
Interaction between Matter
Al3+ ion and Pb2+ ion form white precipitates that are insoluble in an excess of alkali solution. Thus, potassium iodide solution, KI is used to differentiate between these two ions. The presence of Pb2+ is confirmed.
Possibility I
A few drops of potassium iodide solution, KI
Yellow precipitate forms
The yellow precipitate forms again after cooling
Heat The yellow precipitate dissolved in water to form colourless solution
Salt solution which may contain Al3+ or Pb2+ ion
The presence of Al3+ is confirmed.
Possibility II
No precipitate forms
Figure 6.60 Confirmatory tests for Al3+ ion and Pb2+ ion
Qualitative Analysis on Unknown Salts
In order to identify the cations and anions in an unknown salt, you need to perform a systematic analysis according to the order of the test. Figure 6.61 shows the steps that are involved in the qualitative analysis of salts. Carry out Activity 6.28 to identify the anions and cations in the given salt. Observe the colour
Salt
Heating
Add distilled water
Soluble? Yes
No
Add dilute nitric acid, HNO3
Ions in aqueous solution
Test for cations
Tests for anions
Figure 6.61 Qualitative analysis of salts
212
Test for gas
Acid, Base and Salt
CHAPTER 6
Activity 6.28 CT Aims: To identify the cation and anion in L2 solid. –3 –3 Materials: L2 solid, 2.0 mol dm sodium hydroxide solution, NaOH, 2.0 mol dm ammonia solution, NH3, 2.0 mol dm–3 nitric acid, HNO3, 0.1 mol dm–3 silver nitrate solution, AgNO3, 2.0 mol dm–3 hydrochloric acid, HCl, 1.0 mol dm–3 barium chloride solution, BaCl2, distilled water, blue litmus paper and wooden splinter Apparatus: Test tubes, dropper, test tube rack, test tube holder, Bunsen burner, 10 cm3 measuring cylinder, 100 cm3 beaker, spatula and glass rod Procedure: 1. Carry out qualitative analysis to identify the cation and anion in L2 solid. 2. Put half a spatula of L2 solid into a test tube. Add the remaining L2 solid into a beaker. Add 25 cm3 of distilled water into the beaker to dissolve the L2 solid. 3. Divide the L2 solution into four portions to be used for the chemical tests below: • Tests to identify the cation
Table 6.19
Chemical test
Observation
Inference
(a) Pour 2 cm of L2 solution into a test tube. Add 2.0 mol dm–3 sodium hydroxide solution, NaOH until excess. 3
(b) Add 2 cm3 of L2 solution into a test tube. Add 2.0 mol dm–3 ammonia solution, NH3 until excess.
• Tests to identify the anion
Table 6.20
Chemical test
Observation
Inference
(a) Heat L2 solid with high heat in a test tube. Test the gas released with a glowing wooden splinter and moist blue litmus paper. (b) Pour 2 cm3 of L2 solution into a test tube. Add 2.0 mol dm–3 hydrochloric acid, HCl followed by 1.0 mol dm–3 barium chloride solution, BaCl2. (c) Pour 2 cm3 of L2 solution into a test tube. Add 2.0 mol dm–3 nitric acid, HNO3 followed by 0.1 mol dm–3 silver nitrate solution, AgNO3.
4. Record the observations for the activity. Interpreting data: 1. State the inference that corresponds to each observation. 2. Based on the chemical tests, name L2.
Teacher's note http://bit.ly/ 2P6Cu5G
Prepare a complete report after carrying out the activity. 213
THEME 3
Interaction between Matter
While conducting qualitative analysis of salts, systematic and conscientious attitudes are very important in order to identify the cation and anion in a salt correctly.
Activity 6.29 To identify the cations and anions in an unknown salt
Century
21st Skills
CT
You are provided with salt Q1 that contains one cation and anion. Plan a series of chemical tests to identify the cation and anion in salt Q1. 1. Build a flow chart to help you plan your qualitative technique of salt analysis. 2. Identify and list the reagents and apparatus needed. 3. Discuss with your teacher before carrying out the experiment. Teacher's note 4. Carry out the qualitative analysis of salts in the correct order. http://bit.ly/ 5. Write a report on the qualitative analysis of the salt as described 2J6jrVl in Activity 6.28. Identify salt Q1 in your report.
TestYourself
6.11
1. The aqueous solution of a salt is pale green. What cation might be present in the salt? 2. When a sulphate salt is decomposed by heat, the gas released decolourises the purple colour of potassium manganate(VII) solution, KMnO4. What inference can be made on the gas released? 3. When an aqueous solution of a salt Q was tested with alkali solutions, the following observations was obtained. • A white precipitate that is insoluble in an excess of sodium hydroxide solution, NaOH • No change on the addition of ammonia solution, NH3 to the salt solution. What cation may be present in the salt? 4. Three salt samples S1, S2 and S3 contain the zinc ion, Zn2+, aluminium ion, Al3+ and lead(II) ion, Pb2+ respectively. Describe the qualitative analysis carried out to confirm the presence of the ions in each solution.
214
Weak acid
Strong acid
Sodium salt, potassium salt and ammonium salt • Acid + Alkali
Not sodium salt, potassium salt and ammonium salt • Acid + Reactive metal • Acid + Metal oxide • Acid + Metal carbonate
Acid
dissolves in water
Ionise to produce H+ ions
Insoluble salt
Soluble salt
÷ Molar mass × Molar mass
Concentration
calculation
Molarity • M1V1 = M2V2 • n = MV
•
• pH = – log [H+] • pOH = – log [OH–] • pH + pOH = 14
is defined as
Salt
qualitative analysis
Ionic compound that is formed when the H+ ion from the acid is replaced with metal ion or the NH4+ ion
preparation
produces
Neutralisation
reaction between acid and base
Acid, Base and Salt
chemical reaction
Acid • Acid + Base → Salt + Water • Acid + Reactive metal → Salt + Hydrogen gas • Acid + Metal carbonate → Salt + Water + Carbon dioxide gas Alkali • Alkali + Acid → Salt + Water • Alkali + Ammonium salt → Salt + Water + Ammonia gas • Alkali + Metal ion → Insoluble metal hydroxide + Cation from the alkali
Through the double decomposition reaction
partial ionisation
complete ionisation
strong acid
• Low concentration of H+ ion • High pH value
• High concentration of H+ ion • Low pH value
Concept
Weak alkali
partial ionisation
• Low concentration of OH– ion • Low pH value
http://bit.ly/ 2W5y54f
Quick
• Observations on the physical properties of the salt • Effect of heat on the salt • Cation and anion tests • Confirmatory test for cation
Uses in daily life
Alkali
dissolves in water
Ionise to produce OH– ion
strength of alkali
Strong alkali
complete ionisation
• High concentration of OH– ion • High pH value
Acid, Base and Salt
215
CHAPTER 6
THEME 3
Interaction between Matter
Self Reflection 1. What new knowledge have you learned in Acid, Base and Salt? 2. Which is the most interesting subtopic in Acid, Base and Salt? Why? 3. Give a few examples on the application of Acid, Base and Salt in daily life. 4. Rate your performance in Acid, Base and Salt on a scale of 1 to 10; 1 being the lowest and 10 the highest. Why would you http://bit.ly/ rate yourself at that level? 2OSLFXq 5. What can you do to improve your mastery in Acid, Base and Salt?
6
Achievement
1. Table 1 shows the properties of two acids, P and Q. Table 1
Type of acid Formula of the acid pH value of 0.1 mol dm solution –3
P
Q
H2SO4
CH3COOH
0.7
2.9
Based on the information above, answer the following questions: (a) What is the basicity of the acid: (i) P? (ii) Q? (b) Explain your answer in (a). (c) Explain why acid P and acid Q have different pH values. (d) When 10 cm3 of acid P is added into a test tube containing zinc, effervescence occurs. (i) Write a chemical equation for the reaction that occurs. (ii) Calculate the volume of gas released at room conditions. [Molar volume: 24 dm3 mol–1 at room conditions] (e) You are required to prepare 100 cm3 of 0.05 mol dm–3 acid P. Explain briefly how you would prepare dilute acid P.
2. An experiment was carried out to determine the concentration of sodium hydroxide solution, NaOH by titrating 0.5 mol dm–3 sulphuric acid, H2SO4 with 25 cm3 of sodium hydroxide solution, NaOH. Table 2 Titration Final burette reading (cm ) 3
Initial burette reading (cm3) Volume of sulphuric acid, H2SO4 used (cm ) 3
216
I
II
III
25.55
48.20
28.50
0.45
23.00
3.20
Acid, Base and Salt
CHAPTER 6
(a) Complete Table 2. Determine the average volume of sulphuric acid, H2SO4 used. (b) Write a chemical equation for the neutralisation reaction between sulphuric acid, H2SO4 and sodium hydroxide solution, NaOH. (c) Determine the concentration of the sodium hydroxide solution, NaOH used in this experiment. 3. Figure 1 shows the flow chart for a series of reactions that occurs on solid X salt. Reaction W
Salt X Dissolves in water
Heat + a little amount sodium hydroxide solution
Blue solution Y + a little amount aqueous ammonia solution
Black solid P
Blue precipitate
Blue precipitate
+
Brown gas Q
+ excess amount of sodium hydroxide solution
+ excess amount of aqueous ammonia solution
+ Colourless gas R Blue precipitate insoluble in excess amount of sodium hydroxide solution Blue precipitate dissolves in excess amount of aqueous ammonia solution turns into dark blue solution
Figure 1
(a) Based on the flow chart above, identify: (i) Black solid P (iii) Colourless gas R (ii) Brown gas Q (iv) Blue solution Y (b) Write a chemical equation to represent the decomposition of salt X by heat. (c) Explain briefly how you would confirm the anion presence in salt X? (d) The black solid P can be changed into salt X through reaction W. Suggest a chemical that is suitable for changing the black solid P to salt X. Then, explain briefly how the black solid P can be changed to salt X through reaction W.
Enrichmen Corner
1. Hooi See’s mother is very fond of capstone because she loves blue crystals. However, the price of the capstone is very expensive. As a student studying chemistry, how can you help Hooi See to prepare a large and beautiful blue crystal in the laboratory as a present to Hooi See’s mother? Prepare the crystal by including the labelled diagram.
Check Answers http://bit.ly/ 2pHAdTW
217
CHAPTER
7
Rate of Reaction
Keywords
Reactant Product Rate of reaction Catalyst Collision theory Activation energy Effective collision Energy profile diagram
What will you learn? 7.1 Determining Rate of Reaction 7.2 Factors Affecting Rate of Reactions 7.3 Application of Factors that Affect the Rate of Reaction in Daily Life 7.4 Collision Theory 218
Bulletin Nowadays, the use of blow torch in food preparation is gaining popularity. The flame from the blow torch is produced from the burning of butane gas at high concentration, and the reaction is fast. The high temperature of the flame increases the rate of reaction in the food. As a result, food such as meat can be cooked in a short time. If the meat is grilled as normal, the rate of reaction would be slower. Other than temperature and concentration, what other factors would affect the rate of reaction?
What is the difference between average rate of reaction and instantaneous rate of reaction? How does temperature affect the rate of reaction? Why does food cook faster when they are in smaller pieces?
219
theme 3
Interaction between Matter
7.1
Determining Rate of Reaction
Classification of Rate of Reactions
There are a variety of chemical reactions that occur in our surroundings. Did you realise that chemical reactions also occur in our bodies? Do such reactions occur at a fast or slow rate? Figure 7.1 shows several chemical reactions taking place.
Electric cells reaction
g Learnin tandard S
Ignition of matches
Fast reactions
Fireworks
At the end of the lesson, pupils are able to: 7.1.1 Classify fast and slow reactions that occur in daily life 7.1.2 Explain the meaning of the rate of reaction 7.1.3 Identify changes which can be observed and measured during chemical reactions through activity 7.1.4 Determine the • average rate of reaction • instantaneous rate of reaction 7.1.5 Solve numerical problems based on the average and instantaneous rate of reaction
Combustion of gases
Slow reactions
Food decay
Photosynthesis
Corrosion of metals
Rock erosion
Figure 7.1 Examples of fast and slow reactions
220
Fermentation
Rate of Reaction
Meaning of the Rate of Reaction
What do you understand by the rate of a reaction? The rate of reaction is the changes in the quantity of the reactant per unit time or the changes in the quantity of product per unit time. Change in the quantity of reactant or product Rate of reaction = Time taken for the change to occur
CHAPTER 7
Chemistry & Us In the chemical manufacturing factory, a chemical engineer needs to know accurately about the rate of reaction that happens or the duration of time that is needed for the reaction to complete. In other words, a chemical engineer needs to be proficient in rate of reaction.
During the reaction, the quantity of reactant used decreases, while the quantity of product formed increases. The unit for mass of solids is measured in g, while the volume of gases in cm3 or dm3. For the quantity of soluble substances, the concentration is measured in mol dm–3. The choice in unit of time depends on the rate of reaction. For fast reactions, the time is measured in seconds, while for slow reactions, minutes is used. Therefore, the units for rate of reaction that are commonly used are: Unit for rate of reaction: • g s–1 or g minute–1 • cm3 s–1 or cm3 minute–1 • mol dm–3 s–1 or mol dm–3 minute–1
Changes that Occur during Reactions
Determining rate of reaction must be made based on the changes that are observable and can be measured in a certain period of time. What are these changes?
Formation of precipitate occurs in reactions that produce insoluble salts. Photograph 7.1 shows before and after reaction between silver nitrate solution, AgNO3 and sodium chloride solution, NaCl. In this reaction, silver chloride, AgCl and sodium nitrate, NaNO3 are formed. The formation of silver chloride, AgCl can be seen and the precipitate causes the ‘’ mark to disappear, and the amount of precipitate can be measured. Photograph 7.1 The reaction between silver nitrate solution, AgNO3 and sodium chloride solution, NaCl
221
theme 3
Interaction between Matter
Decrease in the mass of the reactants also occurs in reactions that produce gases. Photograph 7.2 shows the reaction between nitric acid, HNO3 and limestone, CaCO3 that produces calcium nitrate, Ca(NO3)2, carbon dioxide, CO2 and water, H2O. The loss in the mass of limestone can be measured by using an electronic balance. Photograph 7.2 The reaction between nitric acid, HNO3 and limestone, CaCO3
Increase in volume of gases occurs for reactions that produce gases. Photograph 7.3 shows the reaction between hydrochloric acid, HCl, and magnesium, Mg. In this reaction, magnesium chloride, MgCl2 and hydrogen gas, H2 are produced. The hydrogen gas, H2 is collected and the volume of the gas measured using a gas syringe. Photograph 7.3 The reaction between hydrochloric acid, HCl and magnesium, Mg
Chemistry Other observable and measurable changes: • Changes in pressure in reactions involving gases. The change in pressure is measured using a pressure meter • Changes in the electrical conductivity of electrolytes in reactions involving mobile ions. An ammeter is used to measure the change in the electrical conductivities in the electrolyte • Changes in pH values for reactions involving acids or bases in aqueous solutions. A pH meter is used to measure the changes in the pH values with time
Activity 7.1 Aim: Determining the time for reaction with reference to some observable and measurable changes. Materials: Zinc powder, Zn, 0.1 mol dm–3 sulphuric acid, H2SO4, marble chips, CaCO3, 2.0 mol dm–3 nitric acid, HNO3, potassium iodide powder, KI, lead(II) nitrate powder, Pb(NO3)2 and distilled water Apparatus: Retort stand with clamp, burette, basin, 250 cm3 conical flask, 10 cm3 and 100 cm3 measuring cylinders, rubber stopper, delivery tube, electronic balance, stopwatch, cotton wool, petri dish, weighing bottle, filter funnel, ruler and filter paper 220
Rate of Reaction
CHAPTER 7
Reaction between zinc, Zn and sulphuric acid, H2SO4 Procedure: 1. Add 20 cm3 of 0.1 mol dm-3 sulphuric acid, H2SO4 into a conical flask. 2. Fill a burette with water and invert it into a basin of water. Clamp the burette vertically. 3. Adjust the water level in the burette so that the level of water is at the 50 cm3 mark. 4. Arrange the apparatus as in Figure 7.2.
Hydrogen gas, H2 Delivery tube
Burette
Conical flask Sulphuric acid, H2SO4
Basin Water
Zinc, Zn
Figure 7.2
5. Add 5 g of zinc powder, Zn into the conical flask that contains sulphuric acid, H2SO4. 6. Immediately close the conical flask with the rubber stopper that is connected to the delivery tube. Start the stopwatch. 7. Record the burette reading at 0.5 minute intervals for 5 minutes. 8. Record your results in Table 7.1 given below. Results:
Time (minute)
Safety Precaution Make sure the extension on the apparatus is tight so that the released gas flows to the burette.
Table 7.1 0.0
0.5
1.0
1.5
2.0
2.5
3.0
3.5
4.0
4.5
5.0
Burette reading (cm3) 50.00 Volume of gas (cm3)
0.00
Discussion: 1. State the observable changes and the measurement made in the activity. 2. Name the gas released. 3. Write a chemical equation for the reaction that occurred. 4. How would you know that the reaction is completed?
221
theme 3
Interaction between Matter
Reaction between nitric acid, HNO3 and marble chips, CaCO3 Procedure: 1. Put 100 cm3 of 2.0 mol dm–3 nitric acid, HNO3 into a conical flask. 2. Close the mouth of the conical flask loosely with Marble cotton wool. chips, CaCO3 3. Set up the apparatus as shown in Figure 7.3. g 4. Add 10 g of marble chips, CaCO3 into the conical flask. Figure 7.3 5. Immediately close the conical flask and start the stopwatch. 6. Record the reading of the electronic balance at intervals of 30 seconds. 7. Observe the changes that occur in the conical flask and record all observations. 8. Record your data in a table. ON OFF
Cotton wool Nitric acid, HNO3 Electronic balance
Discussion: 1. State the observable changes that was recorded in the activity. 2. Why does such a change occur? Explain your answer with the aid of a suitable chemical equation. 3. How would you know that the reaction is completed?
Reaction between potassium iodide solution, KI and lead(II) nitrate solution, Pb(NO3)2 Procedure: 1. Using separate weighing bottles, weigh 2 g of Distilled water potassium iodide powder, KI and 2 g of lead(II) Potassium Lead(II) iodide, KI nitrate powder, Pb(NO3)2 by using two nitrate, Pb(NO3)2 different bottles. Petri dish 2. Pour distilled water into the petri dish to a depth Figure 7.4 of 0.5 cm. 3. Add the potassium iodide powder, KI into the water at the edge of the petri dish. 4. Add lead(II) nitrate powder, Pb(NO3)2 diagonally across from the potassium iodide powder, KI as shown in Figure 7.4. 5. Start the stopwatch immediately. 6. Record the time when the reaction is completed, that is, when no more precipitate is formed. 7. Filter the contents in the petri dish and wash the precipitate with distilled water. 8. Dry and weigh the precipitate. 9. Record your data in a table. Discussion: 1. What is the colour of the precipitate formed? 2. Write an equation for the reaction that occurred. Prepare a complete report after carrying out this activity. Determine the time of reaction based on the observable and measurable changes shown through Activity 7.1. How would you determine the rate of reaction? 220
Rate of Reaction
CHAPTER 7
Average Rate of Reaction and Instantaneous Rate of Reaction
There are two types of rate of reaction, the average rate of reaction and instantaneous rate of reaction. The analogy for the average of reaction and instantaneous rate of reaction is shown in Figure 7.5. A car intends to travel for 400 km. Due to the ever changing traffic conditions, the driver is not able to maintain a constant speed and took 4 hours to reach the destination. The average speed of the car is 100 km hour–1. This is equated as the average rate of reaction. A police officer aims the speed camera in the direction of the car because it is travelling at a speed above the speed limit. At that instant, the camera recorded a speed of 140 km hour–1. The speed at that moment is equated to the instantaneous rate. Figure 7.5 Analogy of average rate of reaction and instantaneous rate of reaction
The average rate of reaction is the average value for the rate of reaction that occurs in a particular time interval. The following explains the way to calculate the average rate of reaction for reactions that release gases. The overall average rate of reaction Total volume of gas collected = Time taken =
V cm3 s-1 t
Volume of gas (cm3)
V
0
Reaction stops here
t Time (s)
The average rate of reaction for the first t1 seconds
Total volume of gas collected in the first t1 seconds = Time taken =
V1 − 0 cm3 s-1 t1 − 0
V = 1 cm3 s-1 t1
Volume of gas (cm3)
Figure 7.6 The overall average rate of reaction
V1
0
t1 Time (s)
Figure 7.7 The average rate of reaction for the first t1 seconds
221
theme 3
Interaction between Matter
The average rate from t1 to t2 Total volume of gas collected from t1 to t2 = Time taken =
Volume of gas (cm3)
V2 − V1 cm3 s-1 t2 − t1
V2 V1
0
t1
t2 Time (s)
Figure 7.8 The average rate of reaction from t1 to t2
The instantaneous rate of reaction is the rate of reaction at a particular point of time. It is determined from the experimental data by plotting a graph of changes in the quantity of the reactants or products against time, and measuring the tangent gradient to the curve at that point of time. Figure 7.9 shows the way to calculate the instantaneous rate of reaction. Method 2
Volume of gas (cm3)
Draw a tangent to the curve at time t. Tangent
0
t
Use the tangent to complete a vertically-angled triangle. The triangle can be drawn in various sizes. The bigger the triangle, the more accurate it can be to determine the tangent gradient.
Volume of gas (cm3)
Method 1
0
Time (s)
t
Time (s)
Method 3
3
Volume of gas (cm )
Calculate the gradient of the tangent to the curve. V2
Rate of reaction at time t = Gradient of the tangent at time t
0 t 1
ΔV Δt V2 − V1 = cm3 s-1 t2 − t1
=
V1 t t2 Time (s)
Figure 7.9 Methods to calculate the instantaneous rate of reaction
220
Rate of Reaction Method to draw the tangent gradient
Total mass of reactants (g)
For reactions that involve a decrease in the total mass of the reactants, a graph as shown in Figure 7.10 is obtained.
CHAPTER 7
0
t
Time (s)
Figure 7.10 The instantaneous rate of reaction that involves the decreasing of reactants
Activity 7.2
Determining the average rate of reaction and instantaneous rate of reaction Answer the following questions based on the data obtained from Activity 7.1. 1. For the reaction between zinc, Zn and sulphuric acid, H2SO4: (a) Plot a graph of gas volume against time. (b) Calculate the following average rates of reaction: (i) For the first minute (ii) For the fifth minute (iii) For the overall reaction (c) Based on the graph plotted, calculate the rate of reaction at the following time: (i) At the first minute (ii) At the third minute (iii) At the end of the reaction 2. For the reaction between potassium iodide solution, KI and lead(II) nitrate solution, Pb(NO3)2: (a) Calculate the average rate of reaction. (b) Can you determine the rate of reaction at 30 seconds? Explain.
Solving Numerical Problems Based on Rate of Reactions Example
1
A student adds magnesium carbonate crystals, MgCO3 until excess into sulphuric acid, H2SO4. The volume of carbon dioxide, CO2 released is collected in a gas syringe and the volume of gas recorded in Table 7.2 for 1 minute intervals for 10 minutes. Table 7.2 Time (minute)
0
1
2
3
4
5
6
7
8
9
10
Volume of gas (cm3)
0
25
40
51
58
63
68
70
70
70
70
(a) Based on Table 7.2, plot a graph of the volume of gas against time. (b) Calculate the following average rate of reactions: (i) For the fifth minute (ii) For the overall reaction (c) Based on the plotted graph, calculate the rate of reaction for the second minute. 221
theme 3
Interaction between Matter
Solution
(a)
Volume of gas (cm3)
60
∆V
40
20
∆t
0
2
4
6 Time (minute)
8
10
(b) (i) Average rate of reaction for the 5th minute Total volume of gas collected from the 4th minute to the 5th minute = Time taken (63 – 58) cm3 = The average rate of reaction (5 – 4) minute from the 4th to the 5th minute. 5 cm3 = 1 minute = 5 cm3 minute–1 (ii) The overall average rate of reaction Total volume of gas collected = Time of reaction 70 cm3 = The reaction ends at the 7th 7 minute
= 10 cm3 minute–1
minute and not 10th minute.
(c) The rate of reaction at the second minute = Gradient of the tangent to the curve at the second minute ΔV = Δt (64 – 20) cm3 = (4 – 0.2) minute = 11.58 cm3 minute–1 220
Rate of Reaction
CHAPTER 7
Activity 7.3 CT Solving numerical problems related to rate of reactions In the presence of manganese(IV) oxide, MnO2, hydrogen peroxide, H2O2 decomposes to water and oxygen. The oxygen gas released is collected in a gas syringe and the volume recorded at intervals of 0.5 minute. The data collected is shown in Table 7.3.
Table 7.3 Time (minute)
0.0
0.5
Volume of gas (cm3)
0.0
13.5 22.0 28.0 33.0 37.0 40.5 43.0 45.0 47.0 48.0 49.0 50.0 50.0 50.0
1.0
1.5
2.0
2.5
3.0
3.5
4.0
4.5
5.0
5.5
6.0
6.5
7.0
1. Based on Table 7.3, plot a graph of volume of gas against time. 2. Calculate the following average rate of reaction: (a) For the first minute (b) For the fifth minute (c) For the whole reaction 3. Calculate the rate of reaction at the following time: (a) 1.5 minute (b) 4.0 minute
TestYourself
7.1
1. Explain the meaning of rate of reaction. 2. Classify the following reactions as fast or slow: (a) Photosynthesis (c) Rusting of iron gate (b) Combustion of petrol in car engines (d) Explosion at oil factory 3. State the observable and measurable change(s) to determine the rate of reaction in the following examples of reactions: (a) 2HCl(aq) + CaCO3(s) ˜ CaCl2(aq) + H2O(l) + CO2(g) (b) H2SO4(aq) + Na2S2O3(aq) ˜ Na2SO4(aq) + H2O(l) + SO2(g) + S(s) (c) 2SO2(g) + O2(g) ˜ 2SO3(g) (d) H+(aq) + OH–(aq) ˜ H2O(l) Table 7.4 4. Metals react with acids at different rates. Metal + Acid → Salt + Hydrogen gas
Metal
Time (s)
P
60
Q
95
Three different metals, P, Q and R react separately with 100 cm R of acid. The time taken to collect 50 cm3 of hydrogen gas for each of the metal is recorded in Table 7.4 (a) Calculate the average rates of reaction for each of the metal with acid. 3
20
(b) Based on your answer in 4(a), arrange the three metals in order of descending reactivity. Explain your answer. 221
theme 3
Interaction between Matter
7.2
Factors Affecting Rate of Reactions
Different chemical substances have different chemical properties. As a result, different chemicals have different reactions and occur at different rates. What are the factors that affect the rate of reactions?
Size of Reactants
Solid reactants can undergo change in sizes. A piece of marble chip can be cut into smaller pieces. The total surface area of all the smaller pieces is larger than the total surface area of the original piece of marble as shown in Figure 7.11
cut
2 cm 2 cm
2 cm
2 cm 1 cm
Total surface area = 24 cm2
2 cm
Total surface area = 2 × 16 cm2 = 32 cm2
g Learnin tandard S At the end of the lesson, pupils are able to: 7.2.1 Investigate factors affecting the rate of reactions through experiment, based on: • Size of reactants • Concentration • Temperature • Presence of catalyst
cut 1 cm 2 cm 1 cm Total surface area = 4 × 10 cm2 = 40 cm2
Figure 7.11 Total surface area of different sizes of reactant
For a fixed mass, the powdered form has a larger total surface area than the original pieces of the solid. Experiment 7.1 shows the effect of size of reactants on the rate of reaction.
Experiment
7.1
Aim: To study the effect of size of reactants on the rate of reaction. Problem statement: How can size of reactant affect the rate of reaction? Hypothesis: The smaller the size of the marble chips, CaCO3, the higher the rate of reaction. Variables: (a) Manipulated : Size of marble chips, CaCO3 (b) Responding : Rate of reaction (c) Fixed : Mass of marble chips, CaCO3, temperature, volume and concentration of hydrochloric acid, HCl Materials: 0.1 mol dm–3 hydrochloric acid, HCl, large pieces of marble chips, CaCO3 and small pieces of marble chips, CaCO3 Apparatus: Conical flask 250 cm3, retort stand with clamp, burette, basin, 100 cm3 measuring cylinder, rubber stopper, delivery tube, electronic balance and stopwatch 220
Rate of Reaction
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Delivery tube Procedure: 3 -3 1. Put 80 cm of 0.1 mol dm hydrochloric acid, Carbon HCl into a conical flask. dioxide Hydrochloric 2. Fill the burette with water and invert it gas, CO2 acid, HCl into a basin filled with water. Clamp the Burette burette vertically. Conical flask Basin 3. Adjust the water level in the burette so that Water 3 the water level reading is 50 cm . Marble chips, CaCO3 4. Set up the apparatus as shown in Figure 7.12. Figure 7.12 5. Weigh 5 g of large pieces of marble chips, CaCO3 and add them into the conical flask. 6. Immediately close the conical flask with the rubber stopper which is connected to a delivery tube. At the same time, Acids are corrosive. Wear start the stopwatch. gloves and safety glasses when handling acids. 7. Slowly swirl the conical flask throughout the experiment. 8. Record the burette reading at intervals of 30 seconds until the burette is completely filled with gas. 9. Repeat steps 1 to 8 by using smaller pieces of 5 g of marble chips, CaCO3. 10. Record all the data in table form.
Interpreting data: 1. Based on the data obtained, plot two graphs of volume of gas against time on a same set of axis. 2. Based on the graph, determine: (a) The tangent gradient at t = 0 (initial rate of reaction). (b) The time taken for the complete reaction.
Conclusion: Is the hypothesis acceptable? What is the conclusion of this experiment? Prepare a complete report after carrying this experiment.
Concentration
The concentration of a solute in a solution can be changed. The concentration of the solution can be changed by adding solvent or solute. Adding solvent
Low concentration
Adding solute
High concentration
Figure 7.13 Concentration of solution
What is the effect of concentration of reactant on the rate of reaction?
231 221
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Interaction between Matter
Experiment
7.2
Aim: To investigate the effect of concentration of reactants on the rate of reaction. Problem statement: How does the concentration of the reactants affect the rate of reaction? Hypothesis: The higher the concentration of sodium thiosulphate solution, Na2S2O3, the shorter the time taken for the ‘’ mark to disappear from view. Variables: (a) Manipulated : Concentration of sodium thiosulphate solution, Na2S2O3 (b) Responding : Time taken for the ‘’ mark to disappear from view (c) Fixed : Temperature, total volume of mixture, concentration and volume of sulphuric acid, H2SO4 and the size of the conical flask Materials: 1.0 mol dm–3 sulphuric acid, H2SO4, 0.2 mol dm–3 sodium thiosulphate solution, Na2S2O3, distilled water and white piece of paper with a ‘’ mark at the centre Apparatus: 150 cm3 conical flask, stopwatch, 10 cm3 and 50 cm3 measuring cylinders Procedure: 1. Put 45 cm3 of 0.2 mol dm–3 sodium thiosulphate solution, Na2S2O3 into a conical flask. Sodium thiosulphate, 2. Place the conical flask on the ‘’ mark on the white Na2S2O3 paper as shown in Figure 7.14. ‘ ’ mark + Sulphuric acid, H2SO4 3. Swiftly, pour 5 cm3 of 1.0 mol dm–3 sulphuric acid, H2SO4 into the conical flask carefully and at White paper the same time start the stopwatch. Figure 7.14 4. Swirl the conical flask gently and place it again on the ‘’ mark. 5. Observe the ‘’ mark vertically from the mouth of the conical flask. 6. Stop the stopwatch once the ‘’ mark disappears from view. Record the time taken. 7. Repeat the experiment by using 0.2 mol dm–3 sodium thiosulphate solution, Na2S2O3 that has been diluted with distilled water as given in Table 7.5. The volume of 1.0 mol dm–3 sulphuric acid, H2SO4 is fixed at 5 cm3. 8. Record all data in Table 7.5. Results:
Table 7.5 Experiment
I
II
III
IV
V
Volume of sodium thiosulphate solution, Na2S2O3 (cm3)
45
40
30
20
10
Volume of distilled water (cm3)
0
5
15
25
35
Volume of sulphuric acid, H2SO4 (cm3)
5
5
5
5
5
Total volume of mixture (cm3)
50
50
50
50
50
Time taken for the ‘’ mark to disappear from view (s)
220
Rate of Reaction
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Interpreting data: 1. The concentration of dilute sodium thiosulphate solution, Na2S2O3 is calculated using the formula M1V1 = M2V2. M1 = The original concentration of sodium thiosulphate solution, Na2S2O3 V1 = The volume of sodium thiosulphate solution, Na2S2O3 used M2 = Concentration of the dilute sodium thiosulphate solution, Na2S2O3 V2 = Volume of the dilute sodium thiosulphate solution, Na2S2O3 Use the given formula and the data collected to calculate the concentration of the dilute sodium thiosulphate solution, Na2S2O3. 2. In the experiment, the rate of reaction is inversely proportional to the time taken for the ‘’ mark to disappear from view. Thus, rate of reaction = 1 . Use this formula and the data time collected to calculate the rate of reaction for all the five experiments. 3. Record all answers from (1) and (2) in Table 7.6. Table 7.6
Experiment
I
II
III
IV
V
Concentration of dilute sodium thiosulphate, Na2S2O3 solution (mol dm–3) Rate of reaction,
1 –1 time (s )
4. Use the data in Table 7.6 to plot a graph of rate of reaction, 1 against concentration of time sodium thiosulphate solution, Na2S2O3 , M2. 5. Based on the graph, state the relationship between rate of reaction and concentration of sodium thiosulphate solution, Na2S2O3. Conclusion: Is the hypothesis acceptable? What is the conclusion of this experiment? Discussion: 1. Why does the solution in the conical flask turn cloudy? 2. Name the substance that causes the solution to turn cloudy. 3. The ‘’ mark disappears from view when the solution in the conical flask reaches a certain level of cloudiness. What are the steps required in this experiment so that the same level of cloudiness is achieved in all the five experiments? 4. What are the changes being measured in the experiment to determine the rate of reaction? Prepare a complete report after carrying out this experiment.
Temperature
Most reactions occur faster at high temperatures, that is, the rate of reaction increases with increasing temperature. For reactions that occur at room temperature, each increase of 10 °C will increase the reaction rate by two times. 221
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Interaction between Matter
Experiment
7.3
Aim: To investigate the effect of temperature on the rate of reaction. Problem statement: How does temperature affect the rate of reaction? Hypothesis: The higher the temperature of the sodium thiosulphate solution, Na2S2O3, the shorter the time taken for the ‘’ mark to disappear from view. Variables: (a) Manipulated : Temperature of sodium thiosulphate solution, Na2S2O3 (b) Responding : Time taken for the ‘’ mark to disappear from view (c) Fixed : Volume and concentration of sulphuric acid, H2SO4 Materials: 1.0 mol dm–3 sulphuric acid, H2SO4, 0.2 mol dm-3 sodium thiosulphate solution, Na2S2O3 and a piece of white paper with an ‘’ mark in the middle Apparatus: 150 cm3 conical flask, 10 cm3 and 50 cm3 measuring cylinders, stopwatch, thermometer, Bunsen burner, wire gauze and tripod stand Procedure: 1. Put 50 cm3 of 0.2 mol dm–3 sodium thiosulphate solution, Na2S2O3 into a conical flask. Leave it for 5 minutes. 2. Record the temperature of the 0.2 mol dm–3 sodium thiosulphate solution, Na2S2O3. 3. Place the conical flask on the ‘’ mark of the white paper. 4. Quickly, add in 5 cm3 of 1.0 mol dm–3 sulphuric acid, H2SO4 into the conical flask carefully. At the same time, start the stopwatch. 5. Swirl the conical flask gently and place it again on the ‘’ mark. 6. Observe the ‘’ mark vertically from the mouth of the conical flask. 7. Stop the stopwatch once the ‘’ mark disappears from view. 8. Record the time taken when the ‘’ mark disappears from view. 9. Repeat steps 1 until 8 by using 50 cm3 of 0.2 mol dm–3 sodium thiosulphate solution, Na2S2O3 that has been heated to 40 oC, 45 oC, 50 oC, and 55 oC. Interpreting data: 1. Use the data obtained to plot a graph of rate of reaction, 1 against temperature of time sodium thiosulphate solution, Na2S2O3. 2. Based on the graph, state the relationship between rate of reaction and the temperature of sodium thiosulphate solution, Na2S2O3. Conclusion: Is the hypothesis acceptable? What is the conclusion of the experiment? Discussion: 1. Write the ionic equation for the reaction between sodium thiosulphate solution, Na2S2O3 and sulphuric acid, H2SO4. 2. Can sulphuric acid, H2SO4 be replaced with hydrochloric acid, HCl? Explain.
Prepare a complete report after carrying out this experiment. 220
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Presence of Catalyst
Catalysts are chemical substances that alter the rate of chemical reactions without undergoing any chemical changes at the end of the reaction. Although the chemical properties of the catalyst does not change, but its physical properties can change. For example, a lump of catalyst can turn into powder. Catalysts do not change the quantity of products. Does the addition of a catalyst increase the rate of reaction?
Experiment
7.4
Aim: To investigate the effect of catalyst on the rate of reaction. Problem statement: How does the presence of catalyst affect the rate of reaction? Hypothesis: Presence of catalyst increases the rate of reaction. Variables: (a) Manipulated : Presence of catalyst (b) Responding : Rate of reaction (c) Fixed : Mass of manganese(IV) oxide, MnO2, temperature and volume of hydrogen peroxide solution, H2O2 Materials: 20-volume hydrogen peroxide solution, H2O2, manganese(IV) oxide powder, MnO2 and distilled water Apparatus: 10 cm3 measuring cylinder, test tubes, test tube rack, glowing wooden splinter, filter funnel, filter paper, 150 cm3 beaker, spatula and electronic balance Procedure: Concentration of 1. Label two test tubes as I and II. hydrogen peroxide 2. Put 5 cm3 hydrogen peroxide solution, H2O2 into test tube I solution and test tube II separately. http://bit.ly/ 3. Place the two test tubes in the test tube rack. 2W0iIKD 4. Add 0.5 g manganese(IV) oxide powder, MnO2 into test tube II. Place a glowing wooden splinter into the mouth of both test tubes quickly. 5. Observe the changes that occur to the wooden splinter and record your observations. Conclusion: Is the hypothesis acceptable? What is the conclusion of this experiment? Discussion: 1. What are the changes observed and measured in this experiment? 2. Explain how the observation in (1) can assure your hypothesis. Prepare a complete report after carrying out this experiment.
Chemistry The mass of catalyst does not change before and after the reaction. You can compare the mass of manganese(IV) oxide, MnO2 before and after Experiment 7.4 for confirmation.
221
Interaction between Matter
Literacy Tips
Diagram 7.15 summarises the effect of concentration, size of reactants, temperature and catalyst on the rate of reaction.
Acronym CSTP can help you to remember the factors that affect the rate of reaction.
Concentration of reactants increases Size of solid reactants decreases
Rate of reaction increases
Temperature increases
Effect of pressure on the rate of reaction
Presence of catalyst Figure 7.15 Factors affecting rates of reactions
http://bit.ly/33zfz6O
Chemistry Pressure is another factor that affects the rates of reactions. Increasing the pressure on reactions involving gases will affect the rate of reaction. When a gas is compressed at constant temperature, the gas particles are pushed into a smaller volume of space. Increasing pressure, increases the concentration of the gas and the rate of reaction. Changes in pressure does not affect the rate of reaction involving solids and liquids reactants as the volume does not change with pressure.
Example
Increasing pressure
2
Reactive metals such as potassium reacts with water to release hydrogen gas. Ca(s) + 2H2O(l) → Ca(OH)2(aq) + H2(g) Two experiments are carried out to determine the rate of reaction between 0.7 g of calcium and 200 cm3 of water at different temperatures. Experiment I is carried out at room temperature. In experiment II, the temperature of water is increased by 10 °C. The diagram below shows the graph of volume of hydrogen gas against time for experiment I. Volume of hydrogen gas (cm3)
theme 3
420 I
0
Time (s)
(a) What is the total volume of hydrogen gas produced in experiment II? Explain your answer. (b) Copy the graph above and sketch the curve for experiment II. (c) What is the effect of temperature on the rate of reaction? 220
Rate of Reaction
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Solution
Volume of hydrogen gas (cm 3)
(a) Total volume of hydrogen gas produced in experiment II = 420 cm3. The quantity of reactants (calcium and water) is the same in both experiments. So, the results of (hydrogen gas) reaction must be the same. (b) 420 II I
Time (s)
0
(c) Increase in water temperature increases the rate of reaction.
Example
3
Gastric is caused by the production of too much acid in the stomach. Doctors use antacid tablets to neutralise the acid in the stomach. Table 7.7 shows the time taken by each antacid tablet to react completely in excess hydrochloric acid, HCl under different conditions. Table 7.7 Experiment
Volume of Concentration of Temperature of hydrochloric acid, hydrochloric acid, hydrochloric acid, HCl (cm3) HCl (mol dm‒3) HCl (oC)
Time of reaction (s)
I
50
1.0
30
120
II
50
2.0
30
60
III
100
2.0
30
60
IV
50
2.0
40
30
(a) Why the time of reaction is different for experiment I and II? (b) Which of the following experiment shows that the change in volume of hydrochloric acid, HCl does not affect the rate of reaction? (c) Why is the rate of reaction for experiment IV higher than experiment II? (d) Other than temperature and concentration of the hydrochloric acid, HCl, what changes can be made to increase the rate of reaction in experiment I? Solution
(a) Experiment I uses 1.0 mol dm–3 hydrochloric acid, HCl while experiment II uses 2.0 mol dm–3 hydrochloric acid, HCl. The concentration of hydrochloric acid is different. (b) Experiment II and III. (c) The temperature of the hydrochloric acid, HCl in experiment IV is higher than in experiment II. (d) The size of the antacid tablet. Crush the tablet into smaller pieces so as to increase the total surface area. 221
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Interaction between Matter
Example
4
Adnan carries out an experiment to investigate the decomposition of hydrogen peroxide, H2O2. He records the volume of oxygen gas released. At the 5th minute, he adds one spatula full of black powder into the hydrogen peroxide solution, H2O2. The diagram shows the graph of volume of oxygen gas released against time. (a) What is the effect of the black powder on the rate of reaction? (b) What is the function of the black powder?
Volume of oxygen gas (cm3)
0
5
Solution
10 15 20 Time (minute)
(a) The addition of the black powder increases the rate of reaction. (b) The black powder acts as a catalyst.
Activity 7.4 CT Discussing the solution involving the rate of reactions and determine the variables involved in the reactions As a group, solve the following problems: Volume of carbon dioxide gas (cm3) 1. Figure 7.16 shows the graph of volume of carbon dioxide against time for the two experiments, I and II. Table 7.7 shows the conditions of the two experiments. II
Table 7.7
I
Experiment
Reactants
I
1.0 g of marble chips, CaCO3 + 50 cm3 of 0.5 mol dm–3 hydrochloric acid, HCl at room temperature
II
Marble, CaCO3 + 50 cm of 0.5 mol dm hydrochloric acid, HCl 3
0
Time (s)
–3
Figure 7.16
Suggest two ways to carry out experiment II so that a similar graph as in Figure 7.16 can be obtained. State all the variables involved. 2. Hydrogen peroxide, H2O2 decomposes slowly at room temperature to produce water and oxygen. The decomposition of hydrogen peroxide, H2O2 can be accelerated by the presence of a catalyst. Three experiments are carried out to determine the effect of three different catalysts on the complete decomposition of 50 cm3 of 10-volume hydrogen peroxide solution, H2O2. Table 7.8 shows the results of the experiments. Table 7.8 Type of catalyst
220
Time(s)
10
20
30
40
50
60
Volume of oxygen gas collected (cm ) 3
Manganese(IV) oxide, MnO2
57
82
93
100
100
100
Copper(II) oxide, CuO
12
19
25
28
30
31
Iron, Fe
33
47
55
58
59
60
Rate of Reaction
(a) Write the chemical equation for the decomposition of hydrogen peroxide, H2O2. (b) Explain briefly how the experiment was conducted. Include the following in your explanation: (i) Problem statement (ii) Hypothesis (iii) All the variables (iv) Diagram for the set-up of the apparatus (c) Which catalyst is more effective in speeding up the rate of decomposition of hydrogen peroxide, H2O2? Explain your answer.
TestYourself
7.2
1. The rate of reaction is affected by various factors. (a) State four factors that would affect the rate of reaction. (b) Zinc, Zn reacts with excess sulphuric acid, H2SO4 according to the following equation: Zn(s) + H2SO4(aq) → ZnSO4(aq) + H2(g)
State four ways to speed up the reaction. In your answer, state the manipulated and the fixed variables.
2. Four experiments to study the reaction between 2 g of marble, CaCO3 with 15 cm3 of hydrochloric acid, HCl is shown in Figure 7.17. 1.0 mol dm–3 hydrochloric acid, HCl + water at 30 °C
5.0 mol dm–3 hydrochloric acid, HCl + water at 30 °C
Lumps of marble, CaCO3 Experiment I Experiment II
1.0 mol dm–3 hydrochloric acid, HCl + water at 60 °C
5.0 mol dm–3 hydrochloric acid, HCl + water at 60 °C
Powdered marble, CaCO3 Experiment III
Experiment IV
Figure 7.17
(a) State the variable that can be observed and measured to determine the rate of reaction. (b) What is the manipulated variable in the experiment? (c) Which experiment has the highest initial rate of reaction? Explain your answer.
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theme 3
Interaction between Matter
7.3
Application of Factors that Affect the Rate of Reaction in Daily Life
Rate of reactions are important in daily life whether at home or in industry. Have you ever wondered how to cook food fast?
g Learnin tandard S At the end of the lesson, pupils are able to: 7.3.1 Explain with examples the application of factors that affect the rate of reaction in daily life
Activity 7.5
Century Solving problems in various daily life activities 21st Skills 1. Carry out the Role-Play activity. 2. Collect information from various sources to solve the following problems: (a) How to cook food fast? (b) How to maintain the freshness of milk? (c) How to remove a blood stain on a shirt? 3. Discuss and prepare the script with suitable equipment. 4. Present your group performance in front of the class within the time allocated.
CT
The Size Factor Action of medicines Antacid tablets are used to treat gastric. Doctors advise patients to chew the tablet instead of swallowing. Breaking up the tablet into smaller pieces increases the total surface area exposed and increases the rate of reaction between the medicine and the acid in the stomach. Cooking food Potatoes are cut into thin slices or long strips so that it can be cooked faster. Thin slices or long strips increases the total surface area exposed to the cooking oil compared to uncut potatoes.
Photograph 7.4 Potato strips
Concentration Factor Corrosion due to acid rain Buildings made of iron that are located near the industrial areas will corrode fast due to acid rain. The atmosphere around industrial areas contains a high concentration of sulphur dioxide. When the concentration of the acidic pollutants increases, the level of acid rain increases and the rate of corrosion increases. 220
Rate of Reaction
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Combustion of petrol in car engines Petrol vapour and air are compressed in the car engine combustion chamber before being burned. The compression increases the concentration of the petrol vapour allowing the petrol to burn very quickly until it explodes. The energy released from the combustion of petrol will make the car move.
Photograph 7.5 Combustion of petrol in a car engine
Temperature Factor Cleaning Washing clothes using detergent powder and hot water combines two factors that increase the rate of reaction. The process of washing clothes will be even quicker in this situation. Cooking food
Photograph 7.6 Washing clothes
Other than decreasing the size, food also cooks faster at high temperatures. Water boils at 100 °C while cooking oil would not boil even the temperature reaches 180 °C. Therefore, frying food in oil will cook the food even faster.
Catalyst Factor Catalytic converter Modern cars are fitted with catalytic converters as shown in Photograph 7.7 to cut down atmospheric pollution. Exhaust gas from car engines contains pollutants. Figure 7.18 shows how catalytic converters change pollutants into non-harmful products that are safe to be released into the atmosphere in the presence of platinum catalyst, Pt. Platinum catalyst, Pt Unburned hydrocarbon, CxHy Carbon monoxide, CO Nitrogen oxide, NOx
Photograph 7.7 Catalytic converter Water, H2O Carbon dioxide, CO2 Nitrogen, N2
Catalytic converter
Figure 7.18
Making alcohol Ethanol, C2H5OH, is the main ingredient in alcoholic drinks. Ethanol is produced through the fermentation of glucose with the help of enzyme in yeast as a catalyst at 37 °C. 221
Interaction between Matter
Activity 7.6 CT Discussing the application of the knowledge on the factors that affect the rate of reactions in daily activities 1. Carry out the activity in groups. 2. Find information from various reading sources and the internet related to the application of the knowledge on the factors that affect the rate of reactions in the following daily activities. (a) Burning of coal (b) Food storage in the refrigerator (c) Cooking using pressure cooker (d) Fermentation process in the making of ‘tapai’ 3. Based on the information collected, start a forum titled ‘Rate of Reaction in Daily Life’.
TestYourself
7.3
1. Fill in the blanks. sizes. (a) Meat cooks faster if it is cut into temperature decreases the growth of bacteria that (b) In the refrigerator, the causes food decay. (c) Food cooks faster in a pressure cooker because of high . (d) Coal burns faster in small chips because of the large surface area that is . 2. In industry, ammonia, NH3 is produced by the direct combination between nitrogen, N2 and hydrogen, H2. Ammonia, NH3 is used to make nitrogeneous fertilisers. Figure 7.19 shows the percentage yield of ammonia, NH3 under different conditions. (a) State the effect on the percentage yield of
ammonia, NH3 with increasing:
% yield of ammonia, NH3
theme 3
70 60 50 40 30 20 10 0
350 °C 400 °C 450 °C 500 °C 550 °C 50 100 150 200 250 300 350 400 450 Pressure (atm)
Figure 7.19 (i) Temperature (ii) Pressure (b) Explain the advantages and disadvantages of using 350 °C and 550 °C in the process of producing ammonia, NH3. How can the disadvantages be overcome?
3. Car exhaust gas contains polluting gases formed from the burning of fossil fuels in the car engine. (a) Name three polluting gases in the car exhaust. (b) In the car engine, nitrogen, N2 combines with oxygen, O2 to produce nitrogen monoxide, NO. At room temperature and pressure, this reaction occurs very slow. Why can this reaction occur in the car engine?
220
Rate of Reaction
7.4
CHAPTER 7
Collision Theory
According to the kinetic theory of matter, matter is made up of tiny and discrete particles that are constantly moving; vibrating at fixed positions for solid, and moving freely for liquids and gasses. As a result, particles collide with one another. During collision, transfer of energy occurs. Fast moving particles transfer some of their energy to slow moving particles and increases their kinetic energy. This process is repeated with other particles. As a result, particles do not have the same kinetic energy and are constantly changing. The collision theory explains how reactant particles interact with one another to cause reactions to occur and form products. According to the collision theory, • reactant particles must collide with one another for reaction to occur • the rate of reaction depends on the frequency of effective collisions Not all collisions between reactant particles result in reactions and the formation of products. Only effective collisions would cause reactions to occur. To produce effective collisions, the reactant particles must have energy equal to or more than the activation energy and collide in the correct orientation.
Activation Energy
g Learnin tandard S At the end of the lesson, pupils are able to: 7.4.1 Describe the collision theory 7.4.2 Explain activation energy using examples 7.4.3 Interpret an energy profile diagram for exothermic reaction and endothermic reaction
Chemistry Collision between particles with energy less than the activation energy or in the wrong orientation is called an ineffective collision.
Collision theory http://bit.ly/35RRzhl
Reactant particles need to have enough energy to initiate a reaction. In other words, activation energy is required to start a reaction. Activation energy is like an energy barrier that needs to be overcomed before a reaction can take place. Photograph 7.8 shows the analogy of such a barrier. The energy required for a horse to overcome the barrier is identical to the activation energy that the particles require to initiate a reaction. Reactant particles need to acquire the minimum energy known as activation energy so as to break the bonds in the reactant particles and form new bonds in the products. Different reactions have different activation energy.
Photograph 7.8 Analogy of activation energy
221
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Interaction between Matter
The activation energy is represented by the symbol Ea. In the energy profile diagram, the activation energy is the difference in energy between the energy level of the reactants and the energy at the peak of the curve in the graph. Exothermic reaction
Endothermic reaction
Energy
Energy
Reactants E a Reactants
Ea
Products
Products Reaction pathway
Reaction pathway
Figure 7.20 Energy profile diagram
In an exothermic reaction, the total energy of the reactants is higher than the total energy of the products. In an endothermic reaction, the total energy of the products is higher than the total energy of the reactants.
Collision Orientation
Reactant particles must be in a specific orientation to result in effective collisions. Figure 7.21 shows the collision between reactant particles in the correct orientation to allow old bonds to be broken and new bonds to form.
→
+
Collision theory (Part 2)
+
←
http://bit.ly/ 35Q8DnZ
Figure 7.21
If the reactant particles collide in incorrect orientations, the particles will bounce back and no reaction occurs. Figure 7.22 shows the collision between reactant particles in the wrong orientation.
→
+
+
←
Collision theory (Part 3) http://bit.ly/ 35RRRor
Figure 7.22
Effective Collision and Rate of Reaction
The rate of reaction depends on the rate of successful collisions between the reactant particles. The higher the frequency of collision between reactant particles with enough energy and in the correct orientation, the faster the reaction occurs. In other words, the rate of reaction depends on the frequency of effective collisions. 220
Rate of Reaction
CHAPTER 7
The higher the frequency of effective collisions, the higher the rate of reaction. The lower the frequency of effective collisions, the lower the rate of reaction.
Effect of Concentration on the Rate of Reaction
Based on Figure 7.23, which has a higher frequency of collision?
Reactant particles
(a) Low concentration
(b) High concentration
Figure 7.23 The effect of concentration of reactants on the rate of reaction
When the concentration of the reactant particles increases, • the number of particles per unit volume increases • the frequency of collisions between particles increases • the frequency of effective collisions between particles increases • the rate of reaction increases For reactions involving gases, changes in pressure is the same as changing the concentration of the gas. What is the effect of gas pressure on the rate of reaction? When the pressure of a gas increases, • the number of particles per unit volume increases • the frequency of collisions between particles increases • the frequency of effective collisions between particles increases • the rate of reaction increases
Effect of Size of Reactant on the Rate of Reaction
When a large piece of a solid reactant is broken up into smaller pieces, the total volume of the substance remains the same. However, the total surface area of the reactant increases. Figure 7.24 explains the effect of particle size on the rate of reaction. 221
theme 3
Interaction between Matter
Only the outer layer of the pink particles of the solid reactants can collide with the green particles.
The increasing surface area enables more pink particles to collide with the green particles.
Broken down
Small fragments of solid reactant
Block of solid reactant
Figure 7.24 The effect of size of reactant on the rate of reaction
When the total surface area of the reactant increases, • the total surface area exposed to collision increases • the frequency of collisions between particles increases • the frequency of effective collision increases • the rate of reaction increases
Effect of Temperature on the Rate of Reaction
When the temperature increases, the rate of reaction increases. This phenomenon is explained in Figure 7.25.
Reactant particles
(a) Low temperature
(b) High temperature
Figure 7.25 The effect of temperature on the rate of reaction
When temperature increases, • the kinetic energy of the reactant particles increases • more particles have energy to overcome the activation energy • the frequency of effective collisions between particles increases • the rate of reaction increases
Effect of Catalyst on the Rate of Reaction
Temperature is a measure of the average kinetic energy of the particles.
A catalyst is involved in a reaction but remains chemically unchanged at the end of the reaction. The catalyst allows the reaction to occur by providing an alternative pathway with lower activation energy, Ea as compared to the original activation energy, Ea. 220
Rate of Reaction
Exothermic reaction
CHAPTER 7
Endothermic reaction Energy
Energy Without catalyst
Ea Reactants
Ea'
Without catalyst
With catalyst
With catalyst
Ea Reactants
Ea'
Products
Products Reaction pathway
Reaction pathway Key: Ea = Activation energy without catalyst
Ea´ = Activation energy with catalyst
Figure 7.26 The effect of catalyst on the magnitude of the activation energy
In the presence of a catalyst, • the catalyst provides an alternative pathway by lowering the activation energy • more reactant particles can achieve the activation energy • the frequency of effective collisions between the particles increases • the rate of reaction increases
Activity 7.7
Century CT Conceptualising the collision theory in reactions that are affected 21st Skills by temperature, reactant size, pressure, concentration and catalyst 1. Carry out the Gallery Walk activity. 2. Get the following information from various reading and search the Internet. (a) Collision theory (b) Use of the collision theory to explain the effect of the following factors on the rate of reaction
3. Discuss with your group members and prepare an interesting presentation. 4. Present your group’s work in the class. Move in groups to see the works of other groups. 5. Write two stars and a wish about the works of the other groups on sticky notes and paste on the work.
Exothermic Reaction and Endothermic Reaction
Changes in energy can occur during chemical reactions. Reactions that release heat energy to the surroundings are called exothermic reactions. Otherwise, reactions that absorb heat energy from the surroundings are called endothermic reactions. All exothermic reactions or endothermic reactions have activation energy, Ea that must be overcomed by the reactant particles. 221
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Interaction between Matter
Exothermic reaction
Endothermic reaction
Energy
Literacy Tips
Energy
Reactants
Ea
∆H (Negative)
Reactants
Ea
You will learn more about exothermic and endothermic reactions in the topic of thermochemistry in Form 5.
Products ∆H (Positive)
Products
Reaction pathway
Reaction pathway
Figure 7.27 Energy profile diagram for exothermic reaction and endothermic reaction
The change in energy content that occurs when reactants are changed into products is known as the heat of reaction and is represented by the symbol ΔH.
TestYourself
7.4
1. The Kinetic Theory of Matter states that the particles in matter are constantly moving. Mark (3) for true statements and (7) for false statements. (a) At constant temperature, all particles move with the same velocity. (b) Particles in solids are moving freely. (c) Collisions between particles are random. (d) The kinetic energy of particles increases with increasing temperature. 2. Scientists use the collision theory to explain how chemical reactions occur. State two important conditions for effective collisions. 3. Catalysts can help to speed up chemical reactions. How does a catalyst speed up a chemical reaction? 4. All reactions including exothermic reactions and endothermic reactions have activation energy that must be overcome by reactant particles. (a) What do you understand by the term activation energy? (b) Mark and label the activation energy in Figure 7.28. (c) Complete the following statements: heat (i) Exothermic reactions . to heat (ii) Endothermic reactions . from 220
Energy
Products
Reactants
Reaction pathway
Figure 7.28
Rate of Reaction
Concept Mass of reactants (g)
Rate of Reaction
m
Tangent gradient = Instantaneous rate of reaction, t1
definition
Change in the quantity of reactants or products Time taken
Quick
0 explained by
http://bit.ly/ 33X1ksG
V
Reactant particles must collide with one another Frequency of effective collisions between particles increases, the rate of reaction increases
• Concentration of reactant , rate of reaction • Number of particles per unit volume
• Size of reactant , rate of reaction • Total surface area exposed to collision
t2
Time (s)
ΔV Average rate = –– Δt
Tangent gradient = Instantaneous rate of reaction, t1 0
affected by
Size of reactant
t1
Volume of gas released (cm3)
Collision Theory
Concentration
Δm Average rate = –– Δt
Temperature • Temperature , rate of reaction • Temperature , kinetic energy of particles
t1
t2
Time (s)
Presence of catalyst • Presence of catalyst, rate of reaction • Alternative pathway with lower activation energy, Ea.
Self Reflection 1. What new knowledge have you learned from Rate of Reaction? 2. Which is the most interesting subtopic in Rate of Reaction? Why? 3. Give a few examples of the application of Rate of Reaction in daily life. 4. Rate your performance in Rate of Reaction on a scale of 1 to 10; 1 being the lowest while 10 the highest. https://bit.ly/ Why would you rate yourself at that level? 32lEzyv 5. What can you do to improve your mastery in Rate of Reaction? 221
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theme 3
Interaction between Matter
7
Achievement
1. Rate of reaction measures the change in the quantity of reactants or products per unit time. (a) State three units for measuring the quantity of substances. (b) For each unit of measurement for the quantity of substances in 1(a), state the corresponding unit for the rate of reaction. 2. There are chemical reactions that are fast or slow in daily life. Some examples are given below. Arrange the reactions in the descending order of the speed of reaction. (a) Banana decaying (d) Combustion of domestic gas (b) Baking a cake (e) Rusting of iron nails (c) Boiling eggs 3. A protein-digesting enzyme is used to study the effect of temperature on the rate of protein digestion in milk. The set-up of the apparatus is shown in Figure 1. The time measured is the time taken for the enzyme to digest all the protein and the milk becomes clear, that is until the ‘’ mark is visible. Table 1 shows the results of the experiment.
Milk + enzyme X
‘ ’ mark is drawn on a white paper and put behind the boiling tube
Figure 1 Table 1 Temperature ( C) o
Time taken for the ‘’ mark to be visible (minute)
15.0 25.0 35.0 45.0 55.0 65.0 12.0
7.0
2.5
4.0
7.0
19.0
1 (minute–1) time
(a) State the changes that are observed and measured. (b) State the manipulated variable and the fixed variables. (c) The time of reaction is determined by measuring the time taken for the milk solution to turn clear. (i) What is the relationship between the rate of reaction and time? (ii) Complete Table 1 by calculating the value of (iii) Plot a graph of
1 . time
1 against temperature. time
(iv) What conclusion can be formulated based on the graph in (c)(iii)? Explain your answer. 220
Rate of Reaction
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4. The collision theory explains how reactant particles interact for reaction to occur and products formed. (a) What are the two main principles of the collision theory? (b) Dinitrogen monoxide, N2O, known as laughing gas, is produced from the following reaction: NH4NO3(aq) → N2O(g) + H2O(l) + Heat (i) Balance the above equation. (ii) Draw the fully labelled energy profile diagram, including the activation energy, for the reaction. (c) Dinitrogen monoxide, N2O reacts with nitrogen monoxide, NO to produce nitrogen, N2 and nitrogen dioxide, NO2. N2O(g) + NO(g) → N2(g) + NO2(g)
The reactants, dinitrogen monoxide, N2O and nitrogen monoxide, NO have to collide in the correct orientation to produce effective collisions and for the reaction to occur. Figure 2 shows the atomic arrangement of the reactants and products. Dinitrogen monoxide
Nitrogen
Nitrogen monoxide
Nitrogen dioxide
Key: Atom N Atom O
Figure 2
Draw the orientation of the reactant particles, dinitrogen monoxide, N2O and nitrogen monoxide, NO that results in effective collisions.
Enrichmen Corner 1. 1.0 g marble powder, CaCO3 is added simultaneously to 50 cm3 of hydrochloric acid, HCl of _ _ concentrations 0.5 mol dm 3 and 1.0 mol dm 3. Excess acid is used in each beaker. Which reaction proceeds fastest? Explain your answer using the collision theory. 2. Acids reacts with metals to produce salt and hydrogen gas. One example of the reaction is shown below: Tin + Hydrochloric acid → Tin(II) chloride + Hydrogen
The time of reaction can be obtained by recording the volume of gas released at fixed time intervals. Plan an experiment to determine the effect of particle size on the rate of reaction. In your answer, include the followings: (a) Sketch the set-up of the apparatus (b) Volume and concentration of the acid used Check Answers (c) Mass and the physical state of tin used (d) Temperature of the reaction (e) Procedure of the experiment https://bit.ly/ 35RsfIt (f) Appropriate graphs (g) Explanation on the conclusion obtained 221
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8
Keywords
Manufactured Substances in Industry
Alloy Pure metal Superconductor Glass Ceramic Composite material
What will you learn? 8.1 8.2 8.3 8.4 252
Alloy and Its Importance Composition of Glass and Its Uses Composition of Ceramics and Its Uses Composite Materials and Its Importance
Bulletin Developments in the automobile industry from time to time has brought about many astounding changes. Try to visualise yourself in the cockpit of the most advanced car in the future. With the help of Artificial Intelligence (AI), the car can automatically move according to the driver’s instructions. Besides traditional synthetic materials, the usage of advanced materials such as nano materials can lead to the invention of high capability systems to execute AI programs in the vehicle. The University of Kuala Lumpur Malaysian Spanish Institute, UniKL MSI has successfully found a way to produce nanoparticles at a low cost. This ideal innovation indirectly leads to the invention of the most advanced car.
Why is pure aluminium not used to make the body of aeroplanes? What are the two main elements found in all types of glass? What is the property of ceramic that makes it suitable to be used in the construction of houses?
253
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Industrial Chemistry
8.1
g Learnin tandard S
Alloy and Its Importance The National Monument was built in 1966 using alloys. What is an alloy? An alloy is a mixture of two or more elements where the main element is a metal. Did you know that many objects around you are made from alloys? Look at several examples of alloys as shown in Figure 8.1.
At the end of the lesson, pupils are able to: 8.1.1 Describe briefly alloy with examples 8.1.2 Compare the properties of an alloy with its pure metal through experiment 8.1.3 Justify the usage of alloys based on their composition and properties
The National Monument of Malaysia was built using bronze.
Photograph 8.1 The National Monument
Duralumin is used to make the body of an aeroplane.
Pewter is used to make souvenirs.
Alloy
Bronze is used to make medals.
Stainless steel is used to make cutlery.
Brass is used to make keys. Steel is used to make the body of a car. Figure 8.1 Examples of alloys and their importance
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Comparison of Properties of Alloys and Pure Metals
The strength and hardness of alloys are based on the arrangement of particles. Are alloys stronger and harder compared to pure metals or vice versa?
Experiment
8.1
Aim: To investigate the comparison between the properties of alloys and pure metals. Problem statement: Is an alloy more resistant to corrosion and harder than a pure metal? Materials: Stainless steel plate, iron plate, distilled water, bronze block and copper block Apparatus: 100 cm3 beaker, 100 cm3 measuring cylinder, sandpaper, steel ball bearing, 1 kg weight, retort stand with clamp, meter ruler and cellophane tape Resistance to corrosion Hypothesis: Stainless steel is more resistant to corrosion than iron. Variables: (a) Manipulated : Type of plate (b) Responding : Corrosion of plate (c) Fixed : Size of plate and volume of distilled water Procedure: 1. Clean the surfaces of stainless steel and iron plates by using a sandpaper. Observe the surfaces of both plates. Record your observations. 2. Immerse both plates in a beaker containing 80 cm3 of distilled water as shown in Figure 8.2. Stainless steel Iron plate plate Beakers 80 cm3 of distilled water
Figure 8.2
3. Leave both beakers aside for one week. 4. After one week, remove both plates and observe the conditions of their surfaces. Record your observations in Table 8.1. Results:
Table 8.1 Condition of plate’s surface Type of plate
After cleaning with sandpaper
After immersing in distilled water for one week
Stainless steel Iron
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Industrial Chemistry
Hardness of substances
Hypothesis: Construct a suitable hypothesis for this experiment. Variables: State the variables involved in this experiment. Procedure: Retort stand 0 cm
1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30
1 kg weight Cellophane tape Steel ball bearing Bronze block
Figure 8.3
1. Fix a steel ball bearing on the surface of the bronze block using a cellophane tape. 2. Hang a 1 kg weight on the retort stand at 50 cm above the surface of the bronze block, as shown in Figure 8.3. 3. Release the weight onto the steel ball bearing. 4. Measure the diameter of the dent formed on the surface of the bronze block. 5. Repeat steps 1 to 4 three times but on different surfaces of the bronze block to obtain an average diameter of the dent formed. Record the reading in Table 8.2. 6. Repeat steps 1 to 5 by replacing the bronze block with a copper block. Results:
Type of block
1
Table 8.2 Diameter of dent (cm) 2 3 4
Average
Bronze Copper
Conclusion: Is the hypothesis acceptable? What is the conclusion of this experiment? Discussion: 1. Why must the stainless steel plate and iron plate be cleaned with a sandpaper? 2. Compare the rate of corrosion of the stainless steel and iron plates. 3. Which block formed a dent with a larger diameter? 4. State the relationship between the diameter of the dent and the hardness of the substance. Prepare a complete report after carrying out this experiment.
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Figure 8.4 shows the comparison between the properties of an alloy and a pure metal. The alloying process helps to prevent the corrosion of metal and alter the properties of a pure metal to make it harder and stronger. How did these changes happen? Alloy
Pure metal
Figure 8.4 Comparison between the properties of an alloy and a pure metal
A pure metal is made up of one type of atom that is of the same size and arranged in an orderly arrangement. When force is applied, the layers of atoms in the metal easily slide over each other. This causes pure metals to be ductile and easily pulled into fine wires.
Atom of pure metal Force
Figure 8.5 Pure metals are ductile
Force
Figure 8.6 Pure metals are malleable
An alloy is formed when foreign atoms are mixed with the pure metal. These foreign atoms are different in size compared to the atoms in the pure metal. Hence, the orderly arrangement of atoms in a pure metal is disrupted. This makes it difficult for the layers of atoms in an alloy to slide over each other when force is applied.
Pure metals are also malleable and its shape can be easily changed. There are empty spaces between the atoms in a pure metal. When force is applied, the layer of atoms in a metal will slide to fill the empty spaces and form a new structure.
Atom of pure metal
Foreign atom
Force
Figure 8.7 Arrangement of atoms in an alloy
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Activity 8.1
Century CT 21st Skills Competition to build a model of particle arrangement in an alloy 1. Carry out this activity in pairs. 2. Discuss with your partner and build a model to illustrate the arrangement of particles in an alloy using balls or spheres. 3. Present your group work by explaining how the model can form an alloy with sturdy arrangement of particles.
Justify the Uses of Alloys based on Their Composition and Properties
The uses of alloys is based on the composition and properties of the alloy produced. Alloys are made for specific purposes. Scientists change the elemental composition to produce alloys with different properties. For example, steel and stainless steel originate from the same pure metal, which is iron. However, both alloys have different properties and are used for different purposes. Table 8.3 Composition, properties and the uses of alloys Alloy
Composition
Properties
Uses
Duralumin
• 93% Aluminium • 3% Copper • 3% Magnesium • 1% Manganese
• Stronger than pure aluminium • Low density • Does not rust
• Body of aeroplanes • Electric cables • Racing bicycles
Bronze
• 90% Copper • 10% Tin
• Stronger than pure copper • Does not rust • Shiny
• Medals • Monuments • Trophies
Brass
• 70% Copper • 30% Zinc
• Stronger than pure copper • Does not rust • Shiny
• Musical instruments • Doorknobs • Keys
Steel
• 98% Iron • 0.2 – 2% Carbon
• Also known as carbon steel • Stronger and harder • Malleable • There are three types of steel, which are low-carbon steel, average-carbon steel and high-carbon steel
• Structure of buildings • Railway tracks • Body of cars
Stainless steel
• 73% Iron • 18% Chromium • 8% Nickel • 1% Carbon
• Stronger than pure iron • Resistant to corrosion
• Cutlery • Sinks • Surgical instruments
Pewter
• 95% Tin • 3.5% Antimony • 1.5% Copper
• Stronger than pure tin • Does not rust • Shiny
• Decorative ornaments • Trophies • Souvenirs
* The composition percentage of metals in each alloy may vary 258
Manufactured Substances in Industry
Photograph 8.2 Maglev train
A superconductor is an example of an alloy used in electrical transportation such as the Maglev train. Superconductors do not have electrical resistance at very low temperatures. This alloy is used to make a magnet that can levitate the train and move it at a very high speed.
CHAPTER 8
Superconductors were discovered in 1911 when mercury that was cooled at 4 K by Heike Kamerlingh Onnes did not have any electrical resistance.
Activity 8.2
Century CT Making a poster to relate the properties and suitability of alloys in 21st Skills daily life 1. Carry out the Gallery Walk activity. 2. Obtain information from various reading sources and the Internet regarding the characteristics and properties of the following alloys and their uses in daily life.
3. Discuss with your group members and design an interesting poster. 4. Display your group work in the class. Move with your group to see the work of other groups. 5. Write comments on other groups̕ work on sticky notes and paste them.
TestYourself
8.1
1. Atoms in pure iron are arranged in an orderly manner and in layers. (a) What is the effect of atom arrangement in pure iron on the ductility and malleability of the metal? (b) Iron alloy can be made by mixing some carbon into molten iron. How does the arrangement of atoms in the iron alloy affect the hardness of the alloy? 2. The purity of gold is measured in carats (K). 24 carat gold is pure gold without the addition of any other metal whereas 18 K gold is a mixture comprising of 18 units by mass of gold with 6 units by mass of other metal such as copper. (a) What is the role of copper in 18 K gold? (b) Calculate the composition percentage by mass in 24 g of 18 K gold. 3. Justify the following statements: (a) Stainless steel is used to make washing machines (b) High voltage electric cables are made from aluminium alloy (c) 18 K gold is used to make rings 259
THEME 4
Industrial Chemistry
8.2
Composition of Glass and Its Uses
Rainbow Skywalk in Georgetown, Penang is a bridge made of glass. Do you know of any other structures that are made from glass? What about the method to make glass?
g Learnin tandard S At the end of the lesson, pupils are able to: 8.2.1 Describe briefly with examples the type of glass, their composition, properties and uses
Photograph 8.3 Rainbow Skywalk
When silica is heated together with other chemicals, various types of glass with different properties are obtained. However, all types of glass have the same basic properties.
Hard but brittle
Chemically inert Basic properties of glass
Electrical insulator Heat insulator
Transparent
Waterproof
Figure 8.8 Basic properties of glass
Types of Glass
In this chapter, you will learn about four types of glass, which are fused silica glass, soda-lime glass, borosilicate glass and lead crystal glass.
Fused silica glass is made from silica (silicon dioxide, SiO2) without adding any other chemical. Silica, SiO2 requires high temperature around 1800 °C to melt. Hence, fused silica glass has a high melting point. This glass does not expand nor contract much when there is a large change in temperature. Fused silica glass is suitable to be used in making telescope lens. 260
Photograph 8.4 Telescope
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Soda-lime glass is made from silica, SiO2, soda (sodium carbonate, Na2CO3) and limestone (calcium carbonate, CaCO3).
Photograph 8.5 Glass containers
Soda, Na2CO3 lowers the melting point of silica, SiO2. Hence, soda glass has a low melting point, around 1000 °C. This glass is easily moulded and used to make glass containers such as bottles and jugs. However, this glass cannot withstand high temperatures and can easily crack when subjected to sudden temperature change.
Borosilicate glass is made from silica, SiO2, soda, Na2CO3, limestone, CaCO3, boron oxide, B2O3 and aluminium oxide, Al2O3. Many laboratory glassware such as beakers and flasks are made from borosilicate glass because of its resistance to heat. These glassware do not crack easily when subjected to thermal stress because of its low expansion coefficient. Borosilicate glass can be removed from the refrigerator and heated immediately without cracking.
Photograph 8.6 Laboratory glassware
Lead crystal glass is made from silica, SiO2, soda, Na2CO3 and lead(II) oxide, PbO. Lead, Pb replaces calcium, Ca to produce glass that is softer and denser. Lead glass is heavier and has a high refractive index. This glass is suitable to be used in making prisms. Photograph 8.7 Prism
Activity 8.3 Making a multimedia presentation about types of glass, composition, CT properties and uses 1. Carry out this activity in groups. 2. Gather information from various reading sources or the Internet about types of glass. 3. Interpret the data regarding the composition, properties and uses of glass. 4. Present the information obtained to the class in the multimedia presentation. 261
THEME 4
Industrial Chemistry
TestYourself
8.2
1. Silica is used to make all types of glass. (a) State the type of glass that is made up of only silica. (b) Soda-lime glass contains alkali metal ions. Name the ion. 2. A sample of borosilicate glass has composition of 80% silica, SiO2,15% boron oxide, B2O3 and 5% alumina, Al2O3. Calculate the mass of each component in the borosilicate glass sample with mass of 1 kg. 3.
Aini’s mother : Aini, do not store food in lead crystal glass containers. Aini : Why not, mother? Based on the conversation above, explain the advantages and disadvantages of using containers made from lead crystal glass.
8.3
Composition of Ceramics and Its Uses
g Learnin tandard S
The Great Pyramid of Giza in Egypt is believed to be made from ceramic. What is ceramic? Is ceramic a type of element or compound? Photograph 8.8 The Great Pyramid of Giza
At the end of the lesson, pupils are able to: 8.3.1 Describe briefly with examples of ceramics, their composition, properties and uses 8.3.2 Expain the uses of ceramics in daily life
A ceramic is a solid made up of inorganic and non-metallic substances. Ceramic is produced through the process of shaping and hardening by using heating technique at a high temperature. Most ceramics are made up of metal compounds, non-metal compounds or metalloid compounds. Photograph 8.9 shows examples of ceramic products.
Aluminium oxide, Al2O3
Titanium carbide, TiC
Photograph 8.9 Examples of ceramic products
262
Silicon carbide, SiC
Manufactured Substances in Industry
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All ceramics have the same basic properties. Do ceramics have the same properties as alloys and glass or vice versa? High thermal resistant Heat insulator
Basic properties of ceramics
Hard and strong
Break easily
Chemically inert Electrical insulator
The atoms in ceramics are bonded by strong covalent bonds and ionic bonds. Hence, ceramics only melt at very high temperatures, are hard and resistant to compression. When force is applied, the atoms in ceramics cannot slide over each other because these atoms are strongly bonded in indefinite arrangement. The energy from the force will be used to break the bonds between the atoms. Hence, ceramics are brittle and weak towards stretching. The electrons in ceramics cannot move freely to conduct electricity or heat.
Figure 8.9 Basic properties of ceramics
Types of Ceramics
Did you know that ceramics can be classified into two groups, that is traditional ceramics and advanced ceramics?
Activity 8.4 Century CT Classifying ceramics into traditional ceramics and advanced ceramics 21st Skills 1. Carry out this activity in groups. 2. Obtain information from various reading sources or the Internet about several examples of ceramics and the classification of ceramics into traditional ceramics or advanced ceramics. 3. Based on the information obtained, discuss with your group members and produce a mind map. 4. Present your group work to the class.
Traditional ceramics are made from clay such as kaolin, Al2O3.2SiO2.2H2O. Clay is mixed with water to produce a soft, mouldable mixture. The mixture is then heated at a very high temperature. Generally, traditional ceramics are used to make bricks, pottery and crockery.
Brick Pottery Bowl
Photograph 8.10 Examples of traditional ceramic use
263
THEME 4
Industrial Chemistry
Advanced ceramics are made from inorganic compounds such as oxides, carbides and nitrides. Advanced ceramics have higher resistance to heat and abrasion, more chemically inert and have superconductivity properties. Advanced ceramics such as silicon carbide are used to make cutting discs due to its hard and strong properties. Silicon carbide is also used to make brake discs because it can withstand thermal shocks and has high resistance to heat. Advanced ceramics are also used to make tungsten carbide rings because it is hard and resistant to abrasion. What other properties and use of traditional ceramics and advanced ceramics that you know?
Cutting disc
Tungsten carbide ring
Brake disc
Photograph 8.11 Examples of advanced ceramic use
Activity 8.5 CT Making a multimedia presentation about the classification, properties and uses of ceramics 1. Carry out this activity in groups. 2. Gather information from various reading sources or the Internet about the classification, properties and uses of ceramics. 3. Use Table 8.4 to organise the information that you have obtained.
Table 8.4
Ceramic …
Classification of ceramic Traditional
Advanced
3
Properties
Uses
…
…
4. Present the information obtained in an interesting multimedia presentation to the class.
Ceramic Uses Application
Examples of ceramic uses
Medicine
• Zirconia ceramic is used in dental implants. • Alumina ceramic is used to make knee bone. • Ceramic is used in Magnetic Resonance Imaging (MRI) machines because it has superconductivity properties.
Transportation • Engine components in jet planes are made from ceramics. Energy production
264
Table 8.5 Ceramic uses
• Ceramic is used to make electrical insulators in high voltage areas such as power stations.
Manufactured Substances in Industry
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Activity 8.6
Century Explaining the use of ceramics in house construction based on the CT 21st Skills properties of ceramics 1. Carry out the Three Stray One Stay activity. 2. Watch a video clip about house construction from the Internet. 3. Based on what you have watched: (a) List ceramic materials used in the construction of houses (b) Discuss the properties of ceramics involved 4. Present your discussion using a suitable mind map. 5. Select a representative to explain the use of ceramic in the construction of houses based on their properties. The other members move to observe and obtain information from the work of other groups.
TestYourself
8.3
1. Fill in the blanks: solid element and non compound. (a) Ceramic is a non (b) Give three examples of substances that make up ceramics. 2. Silicon carbide is an example of advanced ceramic that has a hard structure and diamond-like properties. Can silicon carbide be used to make drinking glass? Explain. 3. Kaolin is a white clay used to make white pottery. What substance must be added to produce green coloured pottery? Explain.
8.4
Composite Materials and Its Importance
g Learnin Standard
Most traditional houses in Malaysia are built using wood. Wood is a natural composite material that is strong and sturdy. What is the meaning of composite material? Why is wood classified as a composite material? Photograph 8.12 Traditional house
At the end of the lesson, pupils are able to: 8.4.1 State the meaning and properties of composite materials 8.4.2 Describe with examples the uses of composite materials 8.4.3 Compare and contrast the properties of a composite material with its constituent materials
265
THEME 4
Industrial Chemistry
A composite material is a material made from combining two or more non-homogeneous substances, that is matrix substance and strengthening substance. The matrix substance surrounds and binds the strengthening substance together.
Matrix substance
Strengthening substance
Chemistry & Us Teeth is a composite material made up of hydroxyapatite and collagen.
Composite material
Cellulose fibres
Lignin
Cellulose fibres
Lignin
Wood
Figure 8.10 Example of matrix substance, strengthening substance and composite material
Both components of a composite material have different physical and chemical properties. When combined, the composite material formed has better properties than the original components.
Composite Materials and Their Uses
Composite materials are widely used in the development and advancement of technology nowadays.
Career Ki o s k A materials scientist studies the structure and chemical properties of natural and synthetic substances to develop new, enhanced materials.
Reinforced concrete is produced when steel bars or wire mesh (strengthening substance) is immersed in concrete (matrix substance). Reinforced concrete is widely used in the construction of bridges, dams and buildings.
Bridge
Building
266
Dam
Manufactured Substances in Industry
CHAPTER 8
Fibre glass is produced when plastic (matrix substance) is strengthened with glass fibres (strengthening substance). This composite material is used to make helmets, car bumpers and printed circuit boards.
Helmet
Car bumper
Printed circuit board
Optical fibre consists of three layers. The innermost layer is the core that is made up of silica glass fibres (strengthening substance). The core is encased in a second layer or cladding that is made up of glass or plastic (matrix substance). The outermost layer is made of plastic that acts as a protective jacket (matrix substance). Optical fibres are used to transmit information and data in the form of light. Light moves through the optical fibre (core section) in a series of total internal reflection. The core and cladding have different refractive indexes to enable them to carry data in large capacity and to not be influenced by electromagnetic disturbances. This composite material has replaced copper wires in video cameras and connects computers in Local Area Network (LAN).
Protective jacket
Optical fibre
Cladding layer
Video camera Core
Cables in computer network
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THEME 4
Industrial Chemistry
Photochromic glass is formed when glass (matrix substance) is combined with silver chloride, AgCl and copper(I) chloride, CuCl (strengthening substance). When exposed to sunlight, photochromic glass darkens. This is due to the formation of silver atoms, Ag that prevents the passage of light. In dim light, copper(I) chloride, CuCl in photochromic glass catalyses the reverse process so that glass becomes transparent again. Photochromic Photochromic glass protects the user glass from UV rays and is suitable for use in car windows, building windows http://bit.ly/ 33wWTEO and camera lenses.
Car window
Camera lens
Building window
Superconductors such as yttrium barium copper oxide, YBCO ceramic is a composite material that has superconductivity properties other than alloys. This superconductor is used to make electromagnets, that are superconductor magnets or supermagnets. Superconductor magnets are light and have strong magnetic force. Superconductor magnets are used in particle accelerators and involved in Nuclear Magnetic Resonance (NMR) machines and Magnetic Resonance Imaging (MRI) machines.
Magnetic Resonance Imaging (MRI)
268
Nuclear Magnetic Resonance (NMR)
Particle accelerator
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Comparison and Difference in Properties of Composite Materials and Their Original Components
Composite materials have different properties compared to their original components. What is the difference in the properties of a composite material and its original components? Concrete can withstand high compression forces but will break if subjected to high stretching forces. Combination of concrete and steel bars or wire mesh can increase the ability of the concrete to withstand compression forces. Properties High compression strength Low stretching strength Resistant to corrosion Properties High stretching strength Corrodes easily
Concrete +
Reinforced concrete
Steel bars or wire mesh
Properties High compression strength High stretching strength Resistant to corrosion
Figure 8.11 Comparison in properties of reinforced concrete with its original components
Plastic matrix consists of plastic that is weak, soft and easily burned. The properties of plastic are reinforced by adding glass fibres. Properties Low stretching strength Low heat and electrical conductivity Resistant to corrosion Durable
Plastic
Properties High stretching strength Low heat and electrical conductivity
Glass fibre
+
Fibre glass
Properties High stretching strength Heat and electrical insulator Resistant to corrosion Durable
Figure 8.12 Comparison in properties of fibre glass with its original components
Optical fibre has high compression strength although the original component that is glass fibre, is brittle. Properties High compression strength Flexible
Plastic
Properties Low compression strength Hard
Glass fibre
+
Optical fibre
Properties High compression strength Flexible
Figure 8.13 Comparison in properties of optical fibre with its original components
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Industrial Chemistry
Glass is transparent and does not absorb UV rays. Silver halide salt crystals such as silver chloride, AgCl is transparent to visible light and absorbs UV rays at the same time. Properties Transparent Does not absorb UV rays Not sensitive to light Properties Transparent to visible light Absorbs UV rays Sensitive to light intensity
Glass
+
Photochromic glass
Silver chloride
Properties Transparent Absorbs UV rays The absorption of UV rays depends on light intensity
Figure 8.14 Comparison in properties of photochromic glass with its original components
Superconductors can conduct electrical current without any resistance at very low temperature whereas its original components cannot. Yttrium(III) carbonate + Property High electrical resistance at room temperature
Copper(II) carbonate + Barium carbonate +
Superconductor (YBCO)
Property No electrical resistance at very low temperature
Oxygen Figure 8.15 Comparison in properties of superconductor with its original components
Activity 8.7 Making a multimedia presentation about the properties, examples and CT comparison between composite materials and their original components 1. Carry out this activity in groups. 2. Gather information from various reading sources or the Internet about composite materials in terms of: (a) Properties (b) Examples (c) Comparison of properties with the original components 3. Present your group work to the class in the form of multimedia presentation. 270
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CHAPTER 8
Activity 8.8
CT STEM Building a composite material 1. Carry out this activity in groups. 2. Read and understand the following passage: The recycling of old newspapers has saved the environment. When 1 tonne of paper is recycled each year, 7000 gallons of water, 4200 kilowatt-hour of energy and 17 trees are saved. Old newspapers can also be reused in the production of products such as baskets, bag handles and paper maché. Starting from this simple effort, you have played a role in building a safe, healthy and beautiful world.
3. Discuss with your group members and draw a composite material. This drawing must combine at least two of the following substances: 4. Write the procedure and build the creation. 5. Present the group work in class. The other group members can obtain information about other groups’ work to improve on the creation.
TestYourself
8.4
1. Concrete is an example of composite material that has been used since ancient times. (a) What is the meaning of composite material? (b) Is concrete suitable to be used to build the pillars of buildings? Explain. (c) Explain how concrete may be reinforced. (d) State two uses of reinforced concrete. 2. Fibre glass is made by immersing glass fibre in molten plastic. (a) Name the matrix substance and strengthening substance used to make fibre glass. (b) Explain why fibre glass is suitable to be used to make water storage tanks. 3. Optical fibre has replaced copper wire in the transmission of information and data. (a) Name three structures that make up an optical fibre. (b) How does optical fibres transmit information and data? (c) Compare the usage of optical fibres and copper wires in high definition cable TV network. 4. Photochromic glass is a composite material which is always used to make car windows. (a) State the main component in photochromic glass. (b) Which component in 4(a) is sensitive to UV light? (c) State two other uses of photochromic glass.
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272
Lead crystal glass
Borosilicate glass
Soda-lime glass
Fused silica glass
• Harder • Stronger • Resistant to corrosion
Pewter
Stainless steel
Steel
Brass
Bronze
Duralumin
examples
Silica
made from
types
properties
Concept
Ceramic
types
• Hard but brittle • Chemically inert • Transparent • Waterproof • Heat and electrical insulator
properties
Glass
examples
Combination of two or more non-homogeneous substances
meaning
Composite material
properties
defined as
Advanced ceramic
Traditional ceramic
Manufactured Substances in Industry
Alloy
meaning
Mixture of two or more elements where the main element is a metal
http://bit.ly/ 2oWu4mS
Quick
Superconductor
Photochromic glass
Optical fibre
Fibre glass
Reinforced concrete
• Heat and electrical insulator • Hard and strong • Chemically inert • High heat resistance • Break easily
Non-metal solid element or inorganic compound made up of metal, metalloid or non-metal
THEME 4 Industrial Chemistry
Manufactured Substances in Industry
CHAPTER 8
Self Reflection 1. What new knowledge have you learned in Manufactured Substances in Industry? 2. Which is the most interesting subtopic in Manufactured Substances in Industry? Why? 3. Give several examples of application of Manufactured Substances in Industry in daily life. 4. Rate your performance in Manufactured Substances in Industry on a scale of 1 to 10; 1 being the lowest and 10 the highest. Why would you rate yourself at that level? https://bit.ly/ 5. What can you do to improve your mastery in Manufactured 2BiDzPy Substances in Industry?
Achievement
8
1. The addition of coke (carbon) in the extraction process of iron is to remove oxygen from iron ore. The iron and carbon mixture will form steel. Table 1 shows two types of steel with different percentage of carbon. Table 1 Steel
Carbon %
Cast iron
4.0
High-carbon steel
0.8
(a) Cast iron is brittle whereas high-carbon steel is hard and strong. Based on Table 1, calculate the percentage of carbon that must be removed from cast iron to produce high-carbon steel. (b) Stainless steel is produced from a mixture of chromium, nickel and carbon. (i) State the percentage of chromium, nickel and carbon in stainless steel. (ii) Stainless steel is suitable to be used to make high quality knife blades. Explain. 2. Lead crystal glass can be used to make spectacle lenses. (a) What is the composition of lead crystal glass? (b) Explain the advantages and disadvantages of using lead crystal glass to make spectacle lenses. (c) Nowadays, spectacle lenses are made from polycarbonate polymer. The properties of polycarbonate are as follows: • Low density and easily moulded • Absorbs UV rays and is very transparent • High impact resistance
You need a pair of new spectacles. Will you choose lenses made from lead crystal glass or polycarbonate? Explain your answer. 273
Industrial Chemistry
3. Traditional ceramics are made from clay such as kaolin. (a) Name two oxide compounds found in kaolin. (b) Give the formula of the ion that produces brown colour in clay. (c) State two uses of traditional ceramics. 4. The various unique properties of ceramics are modified in its use in various fields. State the property of ceramic involved in the manufacture of the following objects: (a) Car engine (b) Spark plug 5. Metals can conduct electricity. Ceramic materials can also be processed to conduct electricity and be made superconductors. Figure 1 shows the change in the electrical resistance value of two conductors against the temperature. Non-superconductor metal
Resistance
THEME 4
Superconductor X
0K
4K
Temperature
Figure 1
(a) Give an example of a ceramic that shows superconductivity properties. (b) Explain the difference in the electrical conductivity properties of the two conductors in Figure 1. (c) How do scientists create a very cold condition to investigate the superconductivity phenomena?
Enrichmen Corner 1. Silicon carbide, SiC is a hard and strong substance that melts at 2700 °C. Silicon carbide, SiC is suitable to be used as an abrasive. Explain why this substance is hard and has high melting point. Carbon, C
Check Answers
Silicon, Si
274
https://bit.ly/ 367eOnU
Figure 1
275
Rb
Sr
Y
57
Actinides
89 – 103
Lanthanides
Yttrium 89 57 – 71
Ti
V
Db
Tantalum 181 105
Ta
Niobium 93 73
Nb
Vanadium 51 41
23
59
Pr
Cr
Mn
Tc
Manganese 55 43
25
Fe
26
8
Pa
Protactinium 231
Ru
Iron 56 44
60
Nd
Seaborgium
Sg
Tungsten 184 106
W
61
Pm
Bohrium
Bh
Rhenium 186 107
Re
U
Np
Neptunium
Metal
Uranium 238
Plutonium
Pu
Samarium 150 94
62
Sm
Hassium
Hs
Osmium 190 108
Os
Molybdenum Technetium Ruthenium 96 101 74 76 75
Mo
Chromium 52 42
24
Praseodymium Neodymium Promethium 141 144 93 91 92
Key:
Thorium 232
Th
Cerium 140 90
58
Ce
Rutherfordium Dubnium
Rf
Hafnium 178.5 104
Hf
Zirconium 91 72
Zr
Titanium 48 40
22
7
Co
27
9
Ni
28
10 29
11
12
Ds
Platinum 195 110
Pt
Palladium 106 78
Pd
Nickel 59 46
Cu
Rg
Gold 197 111
Au
Silver 108 79
Ag
Copper 64 47
Cn
Mercury 201 112
Hg
Cadmium 112 80
Cd
Zinc 65 48
Zn
30
Americium
Am
Europium 152 95
63
Eu
Cf
Dysprosium 162.5 98
66
Dy
Es
Holmium 165 99
67
Ho
Nihonium
Nh
Thallium 204 113
Tl
Indium 115 81
In
Gallium 70 49
Ga
Aluminium 27 31
Al
Fermium
Fm
Erbium 167 100
68
Er
Flerovium
Fl
Lead 207 114
Pb
Tin 119 82
Sn
Germanium 73 50
Ge
Silicon 28 32
Si
N
7
15
117
Mendelevium
Md
Thulium 169 101
69
Tm
Nobelium
No
Ytterbium 173 102
70
Yb
Lawrencium
Lr
Lutetium 175 103
71
Lu
Tennessine
Ts
116
Lv
Astatine
At
Iodine 127 85
I
Bromine 80 53
Br
Chlorine 35.5 35
17
Cl
Fluorine 19
F
9
17
Polonium
Po
Tellurium 128 84
Te
Selenium 79 52
Se
Sulphur 32 34
S
16
Oxygen 16
O
8
16
Moscovium Livermorium
Mc
Bismuth 209 115
Bi
Antimony 122 83
Sb
Arsenic 75 51
As
Phosphorus 31 33
P
15
Nitrogen 14
Non-metal
Berkelium Californium Einsteinium
Bk
Terbium 159 97
65
Tb
Semi-metal
Curium
Cm
Gadolinium 157 96
64
Gd
Meitnerium Darmstadtium Roentgenium Copernicium
Mt
Iridium 192 109
Ir
Rhodium 103 77
Rh
Cobalt 59 45
(Source: International Union of Pure and Applied Chemistry, IUPAC)
Actinium
Lanthanum 139 89
Ac
Radium
Ra
Barium 137 88
Ba
Strontium 88 56
Actinides series
Francium
Fr
Caesium 133 87
Cs
Rubidium 85.5 55
Sc
Scandium 45 39
La
7
6
5
Ca
Calcium 40 38
Potassium 39 37
K
Magnesium 24 20
4
21
6
14
5
C
6
13
Mg 4
B
5
14
12
3
Name of the element
13
Carbon 12
Relative atomic mass
H
Hydrogen 1
Proton number
Boron 11
Be
Symbol of the element
1
Beryllium 9
4
2
The Periodic Table of elements
Sodium 23 19
Na
11
Lithium 7
Li
Hydrogen 1 3
1
H
3
2
1
1
Group
Lanthanides series
Period
Oganesson
118
Og
Radon
Rn
Xenon 131 86
Xe
Krypton 84 54
Kr
Argon 40 36
18
Ar
Neon 20
Ne
Helium 4 10
2
He
18
The data table of elements Element Aluminium Silver Argon Barium Beryllium Boron Bromine Iron Fluorine Phosphorus Helium Hydrogen Iodine Potassium Calcium Carbon Chlorine Cobalt Crypton Chromium Copper Lithium Magnesium Manganese Sodium Neon Nickel Nitrogen Oxygen Lead Rubidium Caesium Silicon Scandium Tin Sulphur Titanium Vanadium Xenon Zinc
276
Symbol Al Ag Ar Ba Be B Br Fe F P He H I K Ca C Cl Co Kr Cr Cu Li Mg Mn Na Ne Ni N O Pb Rb Cs Si Sc Sn S Ti V Xe Zn
Proton number 13 47 18 56 4 5 35 26 9 15 2 1 53 19 20 6 17 27 36 24 29 3 12 25 11 10 28 7 8 82 37 55 14 21 50 16 22 23 54 30
Relative atomic mass 27 108 40 137 9 11 80 56 19 31 4 1 127 39 40 12 35.5 59 84 52 64 7 24 55 23 20 59 14 16 207 85.5 133 28 45 119 32 48 51 131 65
Melting point (°C) 660 962 -189 710 1285 2030 -7 1540 -220 44 -270 -259 114 63 839 3500 -101 1495 -156 1860 1084 180 650 1250 98 -248 1455 -210 -219 327 39 29 1410 1540 232 115 1670 1920 -122 420
Boiling point (°C) 2350 2160 -186 1640 2470 3700 59 2760 -188 280 -269 -252 184 777 1490 4827 -34 2870 -152 2600 2580 1360 1100 2120 900 -246 2150 -196 -183 1760 705 670 2620 2800 2270 445 3300 3400 -108 913
Density (g cm-3) 2.70200 10.50000 0.00170 3.59000 1.84800 2.47000 3.11900 7.89400 0.00160 1.82000 0.00017 0.00008 4.95000 0.86200 1.55000 2.26000 0.00300 8.90000 0.00350 7.19000 8.93000 0.53400 1.73800 7.47000 0.97100 0.00080 8.90000 0.00120 0.00130 11.35000 1.53000 1.87300 2.33000 2.99000 7.28000 1.96000 4.51000 6.09000 0.00550 7.14000
Abrasion
Resistance towards surface friction that occurs in a structure.
Acidic oxide
Oxide compound formed from the reaction between non-metals and oxygen.
Antiseptic
Chemical substance used to kill or prevent the growth of bacteria.
Aqueous solution
A solution which solvent is water.
Atom
The smallest particle in an element that takes part in a reaction.
Basic oxide
Oxide compound that is formed from the reaction between metals and oxygen.
Boiling point
Constant temperature at which a substance changes from liquid to gas at a particular pressure.
Chemical equation
A way of writing that describes a chemical reaction in the form of words or chemical formula.
Chemotherapy
Treatment that involves the use of certain chemicals to treat and control certain diseases, especially cancer.
Coefficient
The number in front of a chemical formula in a chemical equation.
Colouring agent
Natural of synthetic colouring substance that is added to enhance the final colour of processed food.
Crystallisation
The process or forming crystals from a saturated solution.
Degree of dissociation
Mole ratio of reactants that has dissociated to products in 1 litre of solution.
Delocalised electron
Electron that moves freely and is not owned by any atoms or ions.
Disinfectant
Germ-killing agent used to sterilise objects such as laboratory equipment.
Duplet electron arrangement
Stable arrangement of two electrons in the valence shell.
Electron
Negatively charged subatomic particle that orbits the nucleus of an atom.
Electronegativity
The tendency of an atom to pull shared electrons towards itself in a covalent bond.
Expansion coefficient
The change in particle size to temperature changes ratio. 277
Fungicide
Chemical substance to retard or destroy fungi.
Herbicide
Chemical substance to retard or destroy weeds.
Lewis structure
Diagram that shows the bonds between atoms in a molecule and the electron lone pairs in a molecule.
Maglev
Train transportation system that uses magnets to levitate and propel trains.
Metalloid
Element with metal and non-metal properties.
Neutron
Neutral subatomic particle in the nucleus of an atom.
Octet electron arrangement
Stable arrangement of eight electrons in the valence shell.
Pesticide
Poison that is used to kill pests.
Pharmaceutical
Field related to medicine.
Precipitate
Solid that does not dissolve in a solvent and deposited at the bottom.
Preservative
Food additive that delays or prevents the growth and reproduction processes of microorganisms that spoil food.
Proton
Positively charged subatomic particle in the nucleus of an atom.
Reactant
Starting material in a chemical reaction.
Reaction pathway
Refers to a reaction coordinate diagram that shows the change of energy when a reaction occurs.
Refractive index
The ratio of the speed of light in vacuum to the speed of light in a particular medium.
Saturated solution
Solution that contains maximum amount of solute and cannot dissolve any more solute at a particular temperature.
Semiconductor
Substance that has electrical insulator and conductor properties.
Solute
Substance that dissolves in a solvent.
Stoichiometry
Study that is related to the quantity of substance involved in a chemical reaction.
STP
Abbreviation for standard temperature and pressure, which is 0 °C and 1 atm respectively.
Volatility
The property of a liquid that rapidly changes to vapour.
278
Ameyibor, K. & Wiredu, M.B. (2006). Chemistry for Senior Secondary Schools. Oxford: Macmillan Education. Cheng, E., Chow, J., Chow, Y.F., Kai, A., Lai, K.K. & Wong, W.H. (2010). HKDSE Chemistry – A Modern View 1. Hong Kong: Aristo Educational Press Ltd. Cheng, E., Chow, J., Chow, Y.F., Kai, A., Lai, K.K. & Wong, W.H. (2010). HKDSE Chemistry – A Modern View 2. Hong Kong: Aristo Educational Press Ltd. Department of Information. (2018). Rukun Negara. Accessed on 24th July 2019, from http://www. penerangan.gov.my/dmdocuments/rukun_negara_2018/mobile/index.html#p=4 Department of Occupational Safety and Health. (2013). Occupational Safety and Health (Classification, Labelling and Safety Data Sheet of Hazardous Chemicals) Regulations 2013. Accessed on 24th July 2019, from http://www.dosh.gov.my/index.php/legislation/eregulations/regulations-underoccupational-safety-and-health-act-1994-act-514/1125-01-occupational-safety-and-healthclassification-labelling-and-safety-data-sheet-of-hazardous-chemicals-regulations-2013/file Gallagher, R., & Ingram, P. (2015). Complete Chemistry for Cambridge IGCSE. Oxford: Oxford University Press. Harwood, R., & Lodge, I. (2014). Cambridge IGCSE Chemistry Coursebook with CD-ROM. Cambridge: Cambridge University Press. Hayworth, R.M. & Briggs, J.G.R. (2006). All About Chemistry ‘O’ Level. Singapore: Pearson Education. Hayworth, R.M. & Briggs, J.G.R. (2010). Chemistry Insights ‘O’ Level (2nd e.d.). Singapore: Pearson Education. Honeysett, I., Lees, D., Macdonald, A. & Bibby, S. (2006). OCR Additional Science for GCSE. Oxford: Heinemann Educational Publishers. Jones, M., Harwood, R., Lodge, I., & Sang, D. (2017). Cambridge IGCSE Combined and Co-ordinated Sciences Coursebook with CD-ROM. Cambridge: Cambridge University Press. Lim, K.C., Yeo, P.C., Tan, S.H., Cheong, S.L., Chin, S.M., Yabi, S., Guoh, S.L. & Mohamad, K. (2012). Buku Teks KBSM Kimia Tingkatan 5. Selangor: Pan Asia Publications Sdn. Bhd. Low, S.N., Lim, Y.C., Eng, N.H., Lim, E.W. & Ahmad, U.K. (2011). Chemistry Practical Book KBSM Form 4. Kuala Lumpur: Abadi Ilmu Sdn. Bhd. Low, S.N., Lim, Y.C., Eng, N.H., Lim, E.W. & Ahmad, U.K. (2011). Chemistry Textbook KBSM Form 4. Kuala Lumpur: Abadi Ilmu Sdn. Bhd. National Institute of Occupational Safety and Health. (t. t.). General Rules in the Laboratory. Accessed on 24th July 2019, from http://www.niosh.com.my/publication/poster/item/204-peraturan-am-didalam-makmal Oon, H.L. & Chia, L.S. (2010). Chemistry Expression – An Inquiry Approach. Singapore: EPB Pan Pacific. Ryan, L., & Norris, R. (2014). Cambridge International AS and A Level Chemistry Coursebook with CDROM. Cambridge: Cambridge University Press. Tan, Y.T., Chen, L.K. & Sadler, J. (2010). Discover Chemistry Normal (A) 5N Textbook. Singapore: Marshall Cavendish Education. Tan, Y.T., Chen, L.K., Sadler, J. & Sadler, E. (2015). Chemistry Matters for GCE ‘O’ Level (2nd e.d.). Singapore: Marshall Cavendish Education.
279
Acid basicity 137 Activation energy 243, 244, 246, 247 Alkali metal 81, 87, 88, 90, 91 Alloy 254-259, 263, 268 Anion 67, 68, 112 Atomic structure 35 Avogadro constant 50, 73 Catalyst 235, 236, 241, 246, 247 Cation 67, 68, 112 Ceramic 262-264, 268 Chemical bond 110, 126, 128 Chemical equation 69, 70, 72 Chemical formula 59, 60, 65, 67, 70 Chemical technology 6 Composite material 265-269 Covalent bond 110, 114-117, 119, 120, 126, 128, 263 Covalent compound 119, 122, 125-129 Dative bond 120 Double bond 114, 115 Double decomposition reaction 180, 186, 187 Effective collision 243-247 Electronic arrangement 35, 82, 83, 86, 110, 112-116, 120 Electrostatic attraction force 113, 121, 125-127 Empirical formula 60, 61, 64, 65,190 Energy profile diagram 244, 248 Freezing point 26 280
Giant molecule 128 Glass 260, 261, 267-270 Group 80, 82, 83 Halogen 92-95 Hydrogen bond 117-119 Indicator 170, 172 Insoluble salt 179, 180, 186, 188, 190, 200, 221 Ionic bond 110, 111, 113, 116, 126, 127, 263 Ionic compound 67, 113, 122, 125-127, 129, 174, 175, 178 Isotope 37, 38, 45, 91, 243 Lewis structure 115 Matter 4, 24-26, 30, 70 Melting point 26, 85, 88, 91, 93, 101, 126-128 Metallic bond 88, 121 Metalloid 97, 99 Molar mass 52, 73 Molar volume 54, 56 Molarity 158, 159, 160, 164, 172 Mole 49, 50, 52-54 Molecular formula 60, 65 Natural abundance 37 Neutralisation 167-170, 172, 174 Noble gas 84, 85, 110 Nucleon number 32, 34, 37 Period 81-83, 96-99 Periodic Table of Elements 80-83, 101 pH 15, 16, 143-148, 152
pOH 144 Product of reaction 70, 72, 221, 226, 235, 243, 244,247, 248 Protective equipment 12 Proton number 32, 34, 35, 37, 81-83 Pure metal 255, 257, 258 Qualitative analysis 197, 199, 200, 212 Rate of reaction 221, 224-226 Reactant 221, 222, 226, 227, 230, 231, 236, 243-248 Recrystallisation 184 Relative atomic mass 37, 4448, 52 Relative formula mass 48 Relative molecular mass 45, 47 Scientific method 8-10 Simple molecule 127, 128 Single bond 114 Standard solution 162-166 Strength of acid and alkali 149 Superconductor 259, 268, 270 Theory of collision 243 Titration 170 Transition elements 101-104 Triple bond 114, 115 Valence electron 35, 82, 83, 91, 110, 112, 115, 116, 121 Van der Waals attraction force 119, 127, 128
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