I l l u s t r a t i v e

C H E M I S T R Y

F O R S E C O N D A R Y S C H O O L S

F O R M S 1 & 2

J O H N M E S H A C K M A N D I G A

D i p ( E d ) , B s c ( E d ) w i t h H o n o u r s

First printed in June, 2013

Printed and published by

Pepawaks Publishing House

Kampala, Uganda

© John Meshack Mandiga

P.O.Box 30484

Vikuge - Kongowe

Kibaha - Pwani

Mobile:

+255 (0)765 883 640

+255 (0)654 130 099

+255 (0)788 573 277

Email: mandiga1976@yahoo.com

Cover design, photos and typsetting and graphics design by

Michael Ivan Senyonga

Melany Stationaries

Dar es salaam - Tanzania.

ISBN: 978-99-7498-14

All rights reserved. No part of this publication may be reproduced, stored in a

retrieval system, or transmitted, in any form or by any means, without the prior

permission in writting from Pepawaks Publishing House, Kampala - Uganda.

TABLE OF CONTENTS

Chapter 1 - Introduction to chemistry

1. 1 The concept of chemistry .................................................................................... 1

1.2. Branches of chemistry ........................................................................................ 1

1.3 Application of chemistry ..................................................................................... 1

1.4 Importance of chemistry ..................................................................................... 2

1.5 Chemistry for professional development ........................................................... 4

Chapter 2 - Laboratory techniques and safety

2.1 Introduction ........................................................................................................... 7

2.2 Rules and safety precautions in a chemistry laboratory .................................. 7

2.3 Chemical warning signs ....................................................................................... 10

2.4 Chemistry apparatus ............................................................................................. 12

2.5 First aid ................................................................................................................... 21

2.6 First aid procedures .............................................................................................. 23

Chapter 3 - Heat sources and flames

3.1 Heat sources ........................................................................................................... 33

3.2 The Bunsen burner ................................................................................................ 35

3.3 Types of flame ........................................................................................................ 36

3.4 Characteristics of luminous and non-luminous flames ................................... 38

Chapter 4 - scientific procedure

4.1. Significance of the scientific procedure ............................................................. 42

4.2 Main steps of the scientific procedure. .............................................................. 42

4.3 Application of the scientific procedure ............................................................. 46

Chapter 5 - Matter

5.1 Concept of matter ................................................................................................. 48

5.2 States of matter ...................................................................................................... 48

5.3 Physical and chemical changes. .......................................................................... 54

Chapter 6 - Elements, compounds and mixtures

6.1 Elements and symbols .......................................................................................... 60

6.2 Compounds and mixtures ................................................................................... 64

6.3 Solutions, suspensions and emulsions ............................................................... 66

6.4 Methods of separating mixtures ......................................................................... 70

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Chapter 7 - Air, Combustion, Rusting and Fire fighting

7.1 Composition of air. ............................................................................................... 82

7.2 Combustion.............................................................................................................86

7.3 Fire fighting ............................................................................................................ 88

7.4 Rusting. ................................................................................................................... 92

Chapter 8 - Oxygen

8.1 Concept of oxygen. ............................................................................................... 99

8.2 Laboratory preparation of oxygen. ..................................................................... 99

8.3 Properties of oxygen. ............................................................................................ 102

8.4 Industrial manufacture of oxygen. ...................................................................... 105

8.5 Uses of oxygen. ...................................................................................................... 105

Chapter 9 - Hydrogen

9.1 The concept of hydrogen ...................................................................................... 109

9.2 Laboratory preparation of hydrogen .................................................................. 109

9.3 Properties of hydrogen ........................................................................................ 112

9.4 Industrial preparation of hydrogen..................................................................... 114

9.5 Uses of hydrogen .................................................................................................. 115

Chapter 10 - Water

10.1 Occurrence and nature of water. ....................................................................... 120

10.2 Properties of water ............................................................................................. 122

10.3 Water treatment and purification. .................................................................... 132

10.4 Importance of water treatment.......................................................................... 135

10.5 Uses of water. ....................................................................................................... 135

Chapter 11 - Fuels and Energy

11.1 Sources of fuel ..................................................................................................... 138

11.2 Categories of fuels ............................................................................................... 142

11.3 Uses of fuels ......................................................................................................... 142

11.4 Conservation of energy. ..................................................................................... 143

11.5 Biogas as renewable energy................................................................................ 144

Chapter 12 - Atomic Structure

12.1 The concept of atom ............................................................................................ 150

12.2 Sub atomic particles ........................................................................................... 151

12.3 The arrangement of electrons in an atom ........................................................ 152

12.4 Atomic number, mass number and isotopes .................................................. 155

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Chapter 13 - Periodic classification

13.1 The periodic table ............................................................................................... 164

13.2 Periodicity ............................................................................................................ 165

13.3 General periodic trends ..................................................................................... 167

Chapter 14 - Formula, bonding and nomenclature

14.1 Chemical formulae .............................................................................................. 174

14.2 Valency ................................................................................................................ 175

14.3 Oxidation states .................................................................................................. 175

14.4 Radicals ............................................................................................................... 178

14.5 Ions ........................................................................................................................ 179

14.6 Writing a chemical formulae ............................................................................. 181

14.7 Empirical and molecular formulae ................................................................... 183

14.8 Bonding ............................................................................................................... 186

14.9 Nomenclature of binary inorganic compounds .............................................. 193

Questions for Revision and Practice ........................................................................ 201

Appendix A..................................................................................................................... 203

Appendix B..................................................................................................................... 204

Appendix C..................................................................................................................... 207

Glossary ....................................................................................................................... 208

Bibliography ................................................................................................................ 216

Index ............................................................................................................................. 217

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PREFACE

Illustrative Chemistry for Secondary Schools: Forms 1 and 2, is a book that has been

written in such a way that is easily readable and understandable to learners. This

book covers Chemistry topics for forms 1 and 2 as prescribed in Tanzania syllabus.

Chemistry knowledge and skills can easily be acquired through reading this book.

Experiments, activities and exercises are well covered in this book to cater for

learner-centered approach. This book covers the following topics:

1. Introduction to Chemistry

2. Laboratory techniques and safety

3. Heat sources and flames

4. The scientific procedure

5. Matter

6. Elements, Compounds and mixtures

7. Air, Combustion, rusting and fire fighting

8. Oxygen

9. Hydrogen

10. Water

11. Fuels and energy

12. Atomic structure

13. Periodic classification

14. Formula, bonding and Nomenclature

I wish you all the best as you interact with this book.

John Meshack Mandiga,

Kibaha- Pwani,

June - 2013.

vii

ACKNOWLEDGEMENT

I sincerely extend my gratitude to the following people for their help in one way or

another in writing this book.

1. Mr. Michael Senyonga (Graphics designer; Pepawaks Publishing House)

2. Mr. Michael Mogendi Nyanchini (Teacher and author of educational

materials)

3. Pr. Leonard O. Metobo (Theologian and advisor)

4. Mr. Nashon Fanuel Rhobi (Headmaster - Heritage Secondary School)

5. All other individuals who contributed towards successful production of this

book.

John Meshack Mandiga,

Kibaha - Pwani,

June - 2013

viii

Chapter 1

Introduction to

Chemistry

1. 1 THE CONCEPT OF CHEMISTRY

The word Chemistry comes from

By definition, Chemistry is the

the word Alchemy. Alchemy is an branch of science which deals with

earlier set of practices that encompassed

composition, decomposition, structure

elements of Chemistry, metallurgy, and properties of matter. The people who

philosophy, astrology, mysticism and study Chemistry are called Chemists.

medicine. The word alchemy in turn Hundreds of years ago people were

is derived from the Arabic word “al-

interested in what we call Chemistry.

kimia” or “kimi” which is Egyptian These people are called alchemists.

name, meaning “cast together” .

1.2. BRANCHES OF CHEMISTRY

There are many branches of Chemistry. In this level of study the five (5) major

branches of Chemistry will be considered,

➢ Organic Chemistry: is the study of carbon and its compounds.

➢ Inorganic Chemistry: is the study of other substances not containing carbon.

➢ Analytical Chemistry: is the study of separation, identification and composition of materials and the development of tools used to measure properties of matter.

➢ Physical Chemistry: is the study of the physical characteristics of materials and the mechanisms of their reactions.

➢ Biochemistry: is the study of chemical processes that occur inside of living organisms.

1.3 APPLICATION OF CHEMISTRY

Most of the products we use fact that the knowledge of Chemistry is

are produced through application of applied in order to get those products.

Chemistry. There is no way in life we Examples of products made through

can interact with industrial and non-

application of Chemistry are fertilizers,

industrial products and do away with the

pesticides, drugs, vaccines, detergents,

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toothpaste, insecticides, soft drinks, water treatment and purification. The

common salt, paints, cement, clothes, places where Chemistry is applied

fuels, lubricants and grease.

among others are hospitals, homes,

Chemistry is an important subject factories, laboratories, research centres,

that is applied in different fields such universities,water treatment plants and

as mining, agriculture, medicine, mining centres. In our homes, Chemistry

manufacturing, education, food and may be used in baking, cooking and

beverage industry, home care and washing clothes with soap.

cosmetics industry, film industry,

1.4 IMPORTANCE OF CHEMISTRY

The application of Chemistry knowledge directly relates with the importance

of Chemistry. Chemistry is applied in different fields to produce very important

products or rendering helpful services.

1. WATER TREATMENT

➢ The knowledge of Chemistry is

is a useful disinfectant that is used

used to produce chemicals like

in swimming pools to kill bacteria.

water guard that kill germs present

Potassium Aluminum Sulphate is

in water.

a chemical which when added into

➢ Liquid chlorine: is another

water fine particles found in water

chemical which is added in water

settles down to allow the process of

to kill harmful bacteria. Chlorine

sedimentation.

2. TRANSPORT AND COMMUNICATION

➢ Liquid fuels like gasoline (petrol),

the knowledge of Chemistry.

diesel and kerosene are used by

➢ Non-petroleum fossils like biodiesel

different means of transport. These

and alcohols are produced through

fuels are produced by chemical

application of Chemistry.

processes

➢ Chemical processes are applied to

➢ The gaseous fuels like natural gas

produce papers, and wires. These

(methane), liquefied petroleum gas

products are very essential in

(LPG) are also processed by using

communication.

3. AGRICULTURE

➢ Agriculture is the science of

fertilizers, hormones, and growth

livestock keeping and production

agents are important in the field of

of crops.

agriculture. All the named products

➢ Products like pesticides,

are produced through chemical

insecticides, herbicides, fungicides,

processes. Farmers use the products

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I n t r o d u c t i o n t o C h e m i s t r y

in order to get better agricultural

➢ Other products used by farmers are

yields.

weed kil ers and animal vaccines.

4. FOOD AND BEVERAGE INDUSTRY

➢ Biochemical products like

can be employed industrially in

carbohydrates, lipids, proteins,

canning and bottling of foods.

vitamins, minerals, enzymes,

➢ Soft drinks like coca cola and Pepsi

food additives, flavors and colours

are produced through chemical

are produced through chemical

processes like carbonation.

processes.

➢ Carbonation is also involved in

➢ The food industry has also benefited

production of beer, wine, tonic

from chemical processes like food

water and many others.

preservation. Food preservation

5. MANUFACTURING INDUSTRY

➢ Manufacturing is the production

much on Chemistry knowledge.

of goods for use or sale using labor

➢ Products like cement, cars, plastic

and machines, tools, chemical and

containers, chemicals, textiles,

biological processing or formulation.

paper, rubber, glass, computers,

Manufacturing relies on Chemistry

mobile phones and many others

and its chemical processes. The

are produced by application of

raw materials and products depend

Chemistry knowledge.

6. HOME CARE AND COSMETICS INDUSTRY

➢ Home care products like soaps,

application of chemical processes.

detergents, disinfectants, air

➢ Cosmetics like lip stains, face

fresheners, paints, polish and

powder, bronzer, eye lash curler,

vacuum cleaner are used to make the

nail polish, eye liner, lotions and

home and its surrounding cleaner

creams are produced chemically.

and more comfortable to live. All

Other beauty products are

these products are made through

deodorants and primers.

7. MEDICINE

➢ Chemically produced substances

ensures our well-being.

like drugs, vaccines and food

➢ In general the field of medicine

supplements are very important

backed up by Chemistry knowledge

in our lives. Prevention of illness

grants us with healthier living.

and diseases, treatment of diseases

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8. FILM INDUSTRY.

➢ A film camera catches the picture

emulsion consists of silver halide

using chemicals on film. The first

grains suspended in a gelatin

popular photographs were captured

colloid, in the case of colour film.

on copper plates in the 1840s.

➢ Development chemicals applied to

➢ Video recorder, digital cameras and

an appropriate film can produce

video tapes are produced by the aid

either a positive (showing the same

of Chemistry knowledge.

densities and colours as the subject)

➢ A piece of film consists of a light

or negative image (with dark

sensitive emulsion applied to

highlights, light shadows).

a tough, transparent base. The

9. MINING INDUSTRY

➢ Mining is the extraction of valuable

processes especially electrolysis are

minerals or other geological

very essential in converting metal

materials from the earth. Extraction

oxides or sulphides into a pure

metallurgy is the practice of

metal.

removing valuable metals from an

➢ The knowledge of Chemistry is very

ore and refine the extracted raw

important in all these processes.

materials into pure form. Chemical

1.5 CHEMISTRY FOR PROFESSIONAL DEVELOPMENT

The study of Chemistry is very important in many professions or careers.

Professionals like geologists, engineers, nurses, medical doctors, farmers,

horticulturalists, floriculturists, pharmacists, laboratory technicians, researchers and science teachers depend on the knowledge of Chemistry in their careers. Through

studying Chemistry, skills are acquired by such professionals

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I n t r o d u c t i o n t o C h e m i s t r y

SUMMARY

(a). In Chemistry we study the substances which make up the earth, the living things

and the universe in general.

(b). People who study Chemistry are called chemists.

(c). People who were interested in Chemistry hundreds of years ago are called

alchemists.

(d). Chemistry deals with the composition, decomposition, structure and properties

of matter.

(e). Major five branches of Chemistry are organic Chemistry, inorganic Chemistry,

analytical Chemistry, physical Chemistry and biochemistry.

(f). The importance of Chemistry is revealed in such fields as medicine, agriculture,

mining Industry, water treatment, film industry, home care and cosmetics

industry, transport and communication, food and beverage industry and

manufacturing industry.

REVIEW QUESTIONS

1. Match the items in list A with their corresponding statements in LIST B.

LIST A

i. Agriculture ............................................................................................

ii. Medicine ................................................................................................

iii. Manufacturing industry .........................................................................

iv. Transport industry ..................................................................................

v. Food and beverage industry ..................................................................

LIST B

A. Clothes, dyes

B. Fertilizers, pesticides, weed killers, animal vaccines.

C. Fuels, lubricants, oil, grease, coolant, drugs, animal vaccines, food

supplements.

D. Paints, chemicals, vanishes, cement, plastics

E. Detergents, beauty products, shoe polish, tooth paste, disinfectants,

insecticides.

F. Drugs, vaccines, food supplements

G. Soft drinks, common salt, yeast, baking powder, canned food.

2. Write TRUE or FALSE for the following statements.

i. Matter is anything that has mass and occupies space .............................

ii. The people who study Chemistry are called chemists.............................

iii. We cannot apply Chemistry in our homes when we are baking, cooking

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or even washing our clothes with soap ....................................................

iv. The study of Chemistry is important in professional development ........

v. Few materials and objects are made by application of chemical

processes...................................................................................................

3. Choose the most correct answer for the following questions.

(i) ................ might be involved in finding minerals and crude oil in rocks and in their exploitation.

A. Biologists

B. Mathematicians

C.

Geologists

D.

Alchemists

(ii) Letters, newspapers and magazines are written on .....................that is

manufactured chemically.

A. Rocks

B. Plastic

C. Black board

D. Paper.

(iii) Chemically produced substances in the form of...................................and

...............................are used in the field of medicine.

A. Vaccine, medicines

B. Medicines, immunity

C. Vaccines, fruits

D. Fertilizers, vaccines

(iv) The method of protecting food from spoilage is called..............................

A. Food protection

B. Food preservation.

C. Food additives

D. Food canning.

(v). Chemical substances that are used to destroy unwanted plants which are

harmful to crops are called

A. Pesticides

B. Insecticides

C. Weed killers

D. Animal vaccines.

4. Mention seven examples of substances made by applying chemical methods.

5. One of the major problems in the world today is the spread of HIV/AIDS.

Chemists are working to develop new drugs that will combat HIV/AIDS.

Make a list of other problems that we face in the world today. Can you

suggest anything that Chemistry might be able to do to solve them?

6. Briefly explain how you can apply the knowledge of Chemistry at home.

6

Chapter 2

Laboratory

Techniques and Safety

2.1 INTRODUCTION

The term laboratory come from the same Latin word as the English word

labour, which means hard work.

By definition, a laboratory is a special room or building that is designed and used for scientific experiments. Chemical experiments are carried out in a chemistry

laboratory.

2.2 RULES AND SAFETY PRECAUTIONS IN A

CHEMISTRY LABORATORY

The laboratory is meant to be a quiet and safe place to work in . In view of

this, safety in the laboratory is of great importance to both students and teacher.

The laboratory can be very dangerous place if safety regulation is not followed.

LABORATORY RULES.

Laboratory rules are specified guidelines required to be followed when

working with the laboratory. Here is a set of rules to be followed when dealing

with the laboratory.

➢ Never enter the laboratory without the permission or presence of the teacher.

➢ Always dress appropriately for the laboratory activities

➢ Always keep the windows open for proper ventilation

➢ Always master the location of all exit

➢ Always read instruction carefully before you start any experiment or activity.

➢ Never run in the laboratory

➢ Never eat or drink anything in the laboratory

➢ Never quarrel or fight in the laboratory

➢ Never use laboratory apparatus for drinking or storing food

➢ Never taste or sniff chemicals unless advised on how it should be done

➢ Never throw any solid into the sink or water ways

➢ Always use the fume chamber when carrying out experiments where harmful

I l l u s t r a t i v e C h e m i s t r y f o r S e c o n d a r y S c h o o l s

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gases are produced.

➢ Always wash your hands with soap and water before you leave a laboratory

➢ Always perform the intended experiments.

➢ Always replace covers and stoppers on the container after using the chemicals.

➢ Never spill liquids on the floor.

➢ Always keep the gangways and exits clear.

➢ Always report any breakages or accident to the teacher immediately.

➢ Always keep your bench top clean, dry and well arranged.

➢ Always direct the mouth of the test tube away from you or others when heating

substances.

➢ Always use a clean spatula to remove chemicals from containers.

➢ Always use a lighter or wooden splints to light burners, remember to strike the

match before turning on the gas tap.

➢ Never touch electrical appliance with wet hands.

➢ Always turn off any gas or water taps that are not in use.

➢ Never use dirty, cracked or broken apparatus.

➢ Never heat flammable liquids with a Bunsen burner flame.

➢ Never remove chemicals or equipment from the laboratory.

➢ Always wash off any chemical spillage on your skin or clothes with plenty of

water.

➢ Always keep inflammable substances away from naked flames.

➢ Never casually dispose of chemicals and wastes inappropriately.

➢ Always clean up the equipment and store properly after use.

LABORATORY SAFETY MEASURES

Laboratory safety measures are precautions undertaken in order to minimize

risks when carrying out laboratory activities. Laboratory safety measures can be

explained as follows:

1.

The laboratory should be equipped with protective clothes like

laboratory coat and safety glasses all the time

2.

All people using the laboratory should not put themselves or anyone

else in danger by either smoking, causing panic incase of an unwanted

chemical reaction, damage or injury .The supervisor should be called

in case of such situation.

3.

The fume cupboard should be used whenever working with hazardous

or toxic substances and for every experiment where easily evaporable

chemicals are employed

4.

All flammable substances should not be exposed to open flame

5.

When working with concentrated acids or alkalis protective gloves

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and safety glasses should be weared.

6.

Immediately rinse with water and neutralize if necessary in case a

chemical comes in contact with your skin. In case a chemical come

in contact with your eyes wash with running water and call the

supervisor.

7.

A pipette bulb should be used instead of mouth suctions when

pipetting harmful or evaporable substance.

8.

Organic solvent of extremely hazardous and toxic substance should

be disposed of into the designated waste container.

9.

All flammable liquids should not be mixed with oxidizing substance

as this can cause fire.

10. All chemicals that easily react with each other should be stored

separately.

11. There should be equipment for monitoring contamination in order to

give alert of any possible dangers in the laboratory.

12. Refrigerator and freezers should be used for working with laboratory

chemicals only.

13. Adequate first aid kits should be available in the laboratory

14. There should be clear instruction on how to use fire extinguishers in

case of fire.

15. All persons working in the laboratory should not accidentally get into

contact with harmful chemicals by ensuring that cupboards, storage

cabinet and drawers have locks.

16. There should be regular inspection and checking of stored chemicals

in order to avoid using expired substances.

17. Chemical containers should be stoppered and regularly checked to

ensure that they do not leak.

18. Accidental use of wrong substances should be avoided by labeling

all chemicals.

19. All spills should be cleaned immediately.

20. There should be emergency exits and easily accessible.

21. Gas cylinders should be in good working condition always, they

should be labelled, stored and well supported.

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2.3 CHEMICAL WARNING SIGNS

Chemical warning signs are safety symbols found on chemical containers

especially those used in the laboratory to warn about hazardous materials, location

or object.

All chemicals are potential poisons. Some chemicals are dangerously lethal.

Others harmful or toxic. The chemicals we meet in the school laboratory are not

all completely safe. The containers of modern chemicals carry special chemical

warning signs.

h

i

Harmful

Irritant

Radioactive

Toxic

Corrosive

Oxidant

Flammable

Explosive

Fig 2.1. Chemical warning signs

EXPLOSIVE

Explosive substances are usually detonated before they explode. An explosion

is forceful rapid reaction which involve throwing of particles at high speed.

Explosives which can react without external detonation are particularly dangerous.

Chemicals carrying the sign of explosive may cause explosion if they are not

handled carefully and according to instruction. It is dangerous to keep explosives

in glass containers. In case the explosion occurs glass particles will fly around and

cause injuries to people.

TOXIC

Toxic substances can poison people. They can cause death immediately or after

few days. They are most dangerous when they get into your body. They should not

be allowed to enter your body through the mouth, nose, eyes, skin or ears. Some

chemicals in this group are so dangerous that they can enter you body through the

skin. If that happens, wash it out with plenty of water.

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RADIOACTIVE

Radioactive substance can emit harmful radiations. Radioactive materials

contain unstable elements such as uranium and plutonium. Radioactive material

can be in the form of open sources or sealed source. An open source of radioactive

material is normally used as a tracer in experiments and has the potent ion for

spillage and release if not properly contained. A sealed source is in a form that

is permanently bonded or fixed in a capsule or matrix designed to prevent release

of radioactive material

FLAMMABLE

Flammable chemicals can catch fire easily. They should never be brought near

open flames. If they have to be heated at all, an electric heater may be used. Any

type of sparks may set these chemicals on fire, all burners must be put off before

working with flammable chemicals. Flammable chemicals always evaporate fast.

Their containers must be stoppered immediately after use.

HARMFUL OR IRRITANT

Harmful substances may make you sick or endanger your health. They will

not kill you instantly but may affect you after a long exposure. There is always

a danger of ignoring harmful chemicals because they are not instantly lethal. It is

advised that they should be handled according to the safety instruction provided .

Irritating substances can cause pains on your skin or eyes. They can endanger your

health if they come in contact with your skin or eyes for too long

OXIDANT

This is a chemical or substance which helps a burning substance to burn

faster. In presence of an oxidizing agent, small fires can be made very big. Heating

a mixture of an organic material with an oxidizing agent may cause an explosion.

An example is the heating of potassium permanganate with saw dust.

CORROSIVE

Corrosive substances can burn your skin. They can also corrode the floor and

the desk top. You can turn blind if they get into your eyes. If by accident a corrosive substance comes into contact with your skin, go to the sink and wash with plenty

of water. Corrosive substance may destroy metals .Concentrated mineral acids;

H SO , HCl, HNO and concentrated alkalis; NaOH, KOH, NH are examples of

2

4

3

3

very corrosive substances

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2.4 CHEMISTRY APPARATUS

Apparatus are special tools and equipment that are used in the laboratory for

various purposes such as heating, testing, measuring, filtering and grinding.

Table 2.1 showing chemistry apparatus

Apparatus

Diagram

Explanation

A pipette is a narrow glass tube

into which small amounts of liquid

Pipette

are suctioned for transfer to other

(Made of glass)

containers. It is used to measure

specific volumes of liquid.

Measuring

A measuring cylinder is a glass or

cylinder

plastic container that is marked to

(Made of glass

measure volumes of liquid

or plastic)

A thermometer is an instrument

Thermometer

that is used to measure and indicate

(Made of glass)

the temperatures of substances

A burette is usual y used to

accurately measure and dispense

liquids. It is commonly used in

Burette

titrations.

(made of glass)

Measuring

A measuring syringe is used for

Syringe

sucking and measuring specific

(Made of

volumes of liquids or gases.

plastic)

A triple beam balance is a weighing

instrument. It has three beams that

Triple beam

carry weights. On one side there

balance

is a pan on which the object is

(Made of Iron)

placed. Each of the three weights is

slid along the beams to weigh the

object.

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Apparatus

Diagram

Explanation

An electronic balance is a scale that

Electronic

is usual y used to measure the mass

balance

of chemicals. It gives more accurate

(Made of

readings than the beam balance.

variety of

elements)

This is a special watch that is used

Stop watch

to accurately time laboratory

(Made of

processes.

plastic or iron)

A fractionating column is used

to separate vapors of different

Fractionating

densities.

column (made

of glass)

A beehive shelf is used to support a

Beehive shelf

gas jar during collection of gas.

(Made of clay)

A bell jar is used in gas or air

Bell Jar

experiments.

(Made of glass)

Test tube brush is a device made

Test tube brush

with nylon bristles attached to a

(Made of Iron

twisted wire shaft, used to clean test

and nylon)

tubes and flasks.

Corks

Corks are used as stoppers for

(Made of

bottles

wood)

Capil ary tubes are used to enable

Capil ary tubes

small volumes of solutions to be

(Made of glass)

drawn into the tube. It can also be

used for stirring.

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Apparatus

Diagram

Explanation

Rubber tubing

Rubber tubing are suitable for

(Made of

connecting apparatus and delivery

rubber)

of gases.

Glass tubing are used for blowing

them in different shapes to suit

Glass tubing

in some laboratory preparations

(Made of glass)

involving linkage between

containers in experiments

A rubber bung is used to fit in

Rubber bung

delivery tubes during experiments.

(Made of

rubber)

A desiccator is used for keeping

Desiccator

substances dry

(made of glass)

A glass rod is used for stirring

Glass rod

substances during an experiment

(made of glass)

A Kipp’s apparatus is used for

generating small amounts of gas

continuously. It is also used to

prepare poisonous gases.

Kipp’s

apparatus

(made of glass)

Tweezers

Tweezers are tools used for picking

(Made of iron

up objects too small to be easily

or wood)

handled with the human hands.

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L a b o r a t o r y Te c h n i q u e s a n d S a f e t y

Apparatus

Diagram

Explanation

Safety goggles are used to protect

Safety goggle

the eyes from chemical spil s,

(Made of glass

strong light and harmful vapors in

or plastic)

the laboratory.

A boss head is a device that holds a

Boss head

clamp onto a retort stand.

(made of iron)

A volumetric flask is used to make

up a solution to a total final volume.

Volumetric or

graduated flask.

(Made of glass)

Pipe clay

A pipe clay triangle fits on the top

triangle (Made

of a tripod and support a crucible

of Iron and

when it is heated.

clay)

Test tube

A test tube holder is an instrument

holder (Made

that is used for holding a test tube

of iron and

while heating.

wood)

Retort stand and clamp are used

to hold apparatus such as burettes

Retort stand

during experiments.

and clamp

(made of Iron)

This is an instrument that is used to

Tongs (Made of

hold hot substances and apparatus.

iron)

A reagent bottle is a bottle that is

used to store different chemicals.

Reagent bottle

(Made of glass)

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I l l u s t r a t i v e C h e m i s t r y f o r S e c o n d a r y S c h o o l s

Apparatus

Diagram

Explanation

This container is used to store

Plastic wash

distilled water.

bottle (made of

plastic)

A spatula is used for scooping small

Spatula (Made

quantities of powder or crystalline

of iron or

chemicals.

wood)

A thistle funnel is a glass funnel

Thistle funnel

with a wide top and a long stem. It

(Made of glass)

is used to add reagents into flasks

during experiments.

A Bunsen burner is a source of heat

in the laboratory.

Bunsen burner

(made of iron)

A boiling tube is a large test tube

that is used to heat substances

Boiling tube

that should be heated strongly. It

(Made of glass)

is also used when the amount of a

substance is too large for a test tube.

A tripod stand is a stand or support

that has three legs. It is usual y

placed above the Bunsen burner

Tripod stand

when heating or boiling substances.

(Made of Iron)

A metal gauze is placed on top of

the stand. The container that is

used for heating is placed on the

wire gauze.

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L a b o r a t o r y Te c h n i q u e s a n d S a f e t y

Apparatus

Diagram

Explanation

A wire gauze is usual y placed on a

tripod stand. The flask or beaker is

Wire gauze

placed on the gauze during heating.

(Made of iron)

It helps to spread out the flame and

heat evenly under the container.

A crucible is a container in which

Crucible and

substances can be heated to very

lid (made of

high temperature. It is made of

Porcelain)

porcelain or a non - reactive metal.

An evaporating dish is a shallow

Evaporating

bowl with a curled lip. It is used

dish (Made of

to heat and evaporate liquids and

porcelain)

solutions. It can be heated to very

high temperatures.

A deflagrating spoon is a long -

Deflagrating

handled spoon used to heat small

spoon (Made of

amount of substances inside a gas

iron)

jar.

A filter funnel is a device that is

Filter funnel

wide at the top and narrow at the

(Made of

bottom. It is used to separate solids

plastic or glass)

from liquids.

A filter paper is usual y folded into

Filter paper

a cone and placed in filter funnel to

(made of tissue

separate solids from liquids.

paper)

A mortar is a small hard bowl. A

Mortar and

pestle is a small heavy tool used

pestle

for crushing things. A mortar and

(Made of

pestle are usual y used for crushing

porcelain)

or grinding substances.

A test tube rack is special y

Test tube rack

designed for placing test tubes so

(Made of wood

that they do not roll or break.

or plastic)

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I l l u s t r a t i v e C h e m i s t r y f o r S e c o n d a r y S c h o o l s

Apparatus

Diagram

Explanation

A beaker is a glass or plastic

container that is used for holding,

Beaker

heating and mixing liquids. It can

(Made of glass

not measure liquids accurately

or plastic)

but can be used to estimate their

volumes.

A test tube is used for holding

Test tube

chemicals or for heating substances

(Made of glass)

for short periods of time.

Dropper (made

A dropper is used to add liquids

of glass or

during an experiment, drop by

plastic)

drop.

A conical flask is used for holding

Conical flask

liquids in experiments that heating

(made of glass)

is not necessary.

A flat bottomed flask is used for

Flat - bottomed

holding liquids in experiments

flask (made of

where heating is not necessary.

glass)

A round bottomed flask is used to

Round

hold liquids in experiments where

bottomed flask

heating is needed.

(made of glass)

A distilling flask is used for

holding mixtures to be separated in

distil ation process.

Distilling flask

(made of glass)

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L a b o r a t o r y Te c h n i q u e s a n d S a f e t y

Apparatus

Diagram

Explanation

This is a very shallow circular glass

container that is used as a surface

Watch glass

to evaporate some liquids, to hold

(Made of glass)

substances that are being weighed

or observed, or as a cover for a

beaker.

A gas jar is a glass container that

is used for collecting gases during

Gas jar

experiments.

(Made of glass)

A condenser is used to cool gases

down during distil ation so they

condense and form a liquid. The

condenser consists of two tubes one

inside the other. The hot gases pass

Condenser

through the middle tube and the

(made of glass)

cold water passes through the other

tube.

This is a heavy walled glass or

plastic basin used to contain water

Trough (Made

and support a beehive shelf during

of glass or

preparation of gases.

plastic)

Dispensing bottle is used for

keeping aqueous chemicals for

Dispensing

dispensing.

bottle (Made of

glass)

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Experiment 2.1

Aim : To measure the temperature of liquids

Materials: Thermometer, beakers, tripod stand, wire gauze, stopwatch, pair of tongs, water.

Procedure

1. Pour some tap water into two beakers. Measure the temperature of the water

by dipping a thermometer in each of the beakers for one minute.

2. Remove the thermometer and record the temperature

3. Place one beaker in a fridge, or in a bucket of cool water or ice cubes. Let it

stand there for about ten minutes

4. Remove the beaker from the fridge or bucket and dip a thermometer in the

water for one minute. Record the temperature.

5. Place wire gauze on a tripod stand.

6. Place a Bunsen burner under the tripod stand and light it.

7. Place the second beaker of water on the wire gauze and heat for five minutes

8. Turn off the Bunsen burner. Use tongs to remove the beaker from the wire

gauze to avoid burning yourself.

9. Place the thermometer in the beaker containing hot water. Let it stand for one

minute then remove it . Note the temperature

Discussion question

When the thermometer is placed in tap water, water from the fridge, and heated

water, what will be the respective readings?

Experiment 2.2

Aim: To measure the mass of different substances

Materials: Triple beam or electronic balance, various substances such as salt, sugar, watch glasses.

Procedure

1. Place an empty watch glass on the weighing balance. Note down its mass

2. Place the various items you have on the different watch glasses

3. Place the watch glass on the balance ,one at a time and note down the mass

of each item

Discussion question

What are masses of different items you have measured?

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2.5 FIRST AID

First aid is the help given to a sick or injured person before medical assistance

from the hospital. Accidents happen always unexpectedly. Therefore the knowledge

of first aid is important.

The following are importance of first aid:

1. First aid reduces the chance of infection, pain, bleeding and scarring.

2. First aid relieve suffering.

3. First aid preserve life.

4. First aid prevent the situation from worsening.

5. First aid promote recovery.

6. First aid prevents permanent disability.

CAUSES OF ACCIDENTS IN THE LABORATORY

The following are some of the causes of accidents in the laboratory:

1. Careless handling and storage of flammable and toxic substances may

result in fire and explosions

2. Dealing with new or unfamiliar chemical substance may lead into

existence of special hazard

3. Toxic gases, fumes or liquids may escape from their container or spill

while being handled and cause poisoning ,allergies and respiratory

problems.

4. Uncontrolled or unplanned chemical reactions can cause fire and

dangerous explosions.

5. Wet, uneven or damaged floors can cause slips

6. Dropped or burst glassware can cause severe cuts.

7. Entanglement of clothes ,hair or fingers in rotating equipment such as

centrifuges can cause bodily injury

8. Noise and vibration produced from equipment such as centrifuges and

stirrers can cause hearing loss and stress

9. Lack of working understanding of hazards could lead to poisoning.

10. Improper or unintended use of equipment may cause damage to

equipment or injure yourself

11. Inexperienced person working in the laboratory could lead to spillage

of chemical and even fires.

12. Loss of attention to task may lead to use of wrong reagents.

13. Chemical spills and exposure could lead to burns and damage to body

parts such as eyes.

14. Improper disposal of wastes may result in explosions, burns and fires

15. Poor ventilation in a laboratory may cause poisoning.

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16. Electrocution could occur if electrical appliances are not plugged

properly or are touched with wet hands

FIRST AID KIT

A first aid kit is a box that contains equipment and chemicals needed for the

first aid. The main items in a first aid kit can be explained in the following table.

Table 2.2 Components of the first aid kit and their uses.

ITEM

USE

First aid book

Contains guidelines on how to use the items in the first aid

kit.

Plaster or adhesive

Covering small cuts or wounds.

bandage

Sterile gauze

Covering wounds to protect them from dirt and germs.

Iodine tincture

Cleaning fresh wounds to kill germs.

Soap

Washing hands, wounds and equipment

Pain killers

Relieving pain

Scissors or razor blade

Cutting dressing materials

Safety pins

Securing bandages

Bandage

Keeping dressings in place and immobilizing injured limbs

Cotton wool

Cleaning and drying wounds

Thermometer

Taking body temperature

Disposable sterile gloves Covering the hands to avoid infecting wounds and to

prevent direct contact with a victim’s body fluids.

Petroleum jel y

Soothing chapped skin

Liniment

Reducing muscle pain

Torch

Source of light

Whistle

Blown to call for help

Mild antibiotics

Treating mild bacterial infections on the skin, ear, nose and

mouth

Gentian violet (G.V)

For fungal infections of the skin and mouth. Also used for

the treatment of serious heat burns.

Tweezers

Plucking hair from the face or eye brows

Instant ice packs

Treating injuries, sprains, bumps, bruises, insect bites and

stings. Can also help ease the symptoms of headaches.

Emergency blanket

Preventing and countering hypothermia by reducing heat

loss from person’s body. Covering injured victims in order to

help reduce shock.

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2.6 FIRST AID PROCEDURES

The following are some situations that may require first aid and the procedure

to follow in giving help.

BURNS

A burn is a type of injury to flesh or skin caused by heat, electricity, chemicals,

light, radiation or friction. Burns usually cause blisters on the skin and if severe the skin becomes charred and peels off. Burns caused by hot liquids or gas are called

scalds

Procedure

1. Lay the victim down and protect the burned area from coming into

contact with the ground if possible

2. Hold the burned area under cool running water for 10 or 15 minutes or

until the pain subsides. If this is impractical, immerse the burn in cool

water or cool it with cold compresses. Cooling the burn reduce swelling

by conducting heat away from the skin. Don’t put ice on the burn.

3. Check breathing and pulse and be prepared to resuscitate the victim ,if

necessary

4. Gently remove any jewelry, shoes or burned clothing from the injured

area. Loosen any tight clothing. Do not remove any clothing that is

sticking to the skin.

5. Cover the burn with sterile gauze bandage. Don’t use fluffy cotton or

other material that may get lint in the wound. Wrap the gauze loosely

to avoid putting pressure on burned skin. Bandaging keeps air off the

burn, reduces pain and protects blistered skin.

6. Take an over-the-counter pain reliever. These include aspirin, naproxen

or acetaminophen. Use caution when giving aspirin to children or

teenagers. Though aspirin is approved for use in children older than

2,children and teenagers recovering from chickenpox or flu-like

symptoms should never take aspirin. Seek medical help immediately.

Caution

➢ Do not use ice. Putting ice directly on a burn can cause a person’s body to

become too cold and cause further damage to the wound.

➢ Do not apply egg whites, butter or ointments to the burn. This could cause

infection.

➢ Do not break blisters. Broken blisters are more vulnerable to infection

➢ Burns to the face and in the mouth or throat are serious, as they cause rapid

inflammation of the air passage and may cause suffocation. Seek medical help

immediately.

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SUFFOCATION

Suffocation is the condition of being deprived of oxygen. Foam can appear at

the mouth and nostril of the victim. The victim may eventually lose consciousness.

Some chemical substances can cause suffocation when inhaled, leading to breathing

difficulties

Procedure

1. Remove anything covering the mouth and nose. A plastic bag should

be torn open and attempt to remove anything obstructing the airway.

2. Ensure the victim’s airway is open for air to reach the lungs. Do this

by placing the victim on his or her back. With one hand on the victim’s

fore head and the other on the chin, tilt the head backwards to open the

air way

3. If the casualty is unconscious,place him or her in the lateral position,

check the airway, breathing and pulse and begin expired air resuscitation

(EAR) or cardiopulmonary resuscitation (CPR)if necessary. Continue

until natural breathing is restored.

4. If the casualty is or becomes conscious ,monitor air way, breathing and

pulse.

5. Keep the victim warm using a light blanket.

6. Seek medical help immediately

CHOKING

Choking is the mechanical obstruction of the flow of air from the environment

into the lungs. Choking prevent breathing, and can be partial or complete. Signs of

choking include difficulty in speaking and breathing.

Procedure

1. Encourage the victim to cough up the object.

2. If the object remain stuck deliver five back blows between the person’s

shoulder blades with the heel of your hand

3. If the object is still stuck ,perform the Heimlich manouvre. This procedure involves the following

a. Stand behind the person. Wrap your arms around the waist. Tip the person

forward slightly

b. Make a fist with one hand. Position it slightly above the person’s navel.

c. Grasp the fist with the other hand. Press hard into the abdomen with a

quick upward thrust-as if trying to lift the person up .

d. Perform a total of 5 abdominal thrusts, if needed. If the blockage still isn’t

dislodged, repeat the thrusts until the object comes out.

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BRUISES

A bruise is a mark on the skin formed when small blood vessels break and

leak their content into the soft tissue beneath the skin. When a blow breaks blood

vessels near the skin surface leakage of small amount of blood lead to discoloration

of the skin.

Procedure

1. Elevate the injured area

2. Apply ice or a cold pack several times a day for a day or two after the injury

3. Rest the bruised area, if possible

4. Consider pain relief drugs and others which reduce swelling

Caution

For the first 48 hours after an injury, avoid things that might increase swelling

such as hot showers, hot tubs, hot packs or alcoholic beverages

SHOCK

Shock is a life –threatening condition that occurs when the body is not getting

enough blood flow. This can damage multiple organs. Shock requires immediate

medical treatment and can get worse very rapidly. Shock ultimately leads to cellular

death, progressing to organ failure, and finally if untreated whole body failure and

death. In general when a person is in shock, his or her vital organs like heart , the

lungs and the brain are not getting enough blood or oxygen.

A person experiencing shock may have pale or gray skin ,the pulse is weak

and rapid, the person may be nauseated and vomit, the eyes lack luster and may

seem to stare (sometimes the pupils are dilated) the person may be unconscious. If

conscious, the person may feel faint or be very weak or confused. Shock sometimes

causes a person to become overly excited and anxious

Procedure

1. Have the person lie down on his or her back with feet about a foot higher

than the head. If raising the legs will cause pain or further injury, keep him

or her flat. Keep the person still.

2. Check for signs of circulation (breathing, coughing or movement) and if

absent; begin cardiopulmonary resuscitation (CPR).

3. Keep the person warm and comfortable by loosening any belts or tight

clothing and covering the person with a blanket. Even if the person

complains of thirst, give nothing by mouth.

4. Seek medical help immediately.

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ELECTRIC SHOCK

An electric shock is a condition that occurs when a person comes into contact

with an electric energy source. An electric shock is usually caused by contact with

poorly insulated wires or ungrounded electrical equipment, by using electrical

equipment while in contact with water, or by being struck by lightning.

Procedure.

1. Turn off the source of electricity.

2. Move the source of electric shock away from the person using a dry, non

conducting object made of cardboard, plastic or wood.

3. Check for signs of circulation (breathing, coughing or movement). If

absent, begin cardiopulmonary resuscitation (CPR) immediately.

4. If the person is breathing but unconscious, put him or her in the recovery

position.

5. Administer first aid for burns, shock or other injuries the victim may have

sustained.

6. Seek medical help immediately.

Caution.

➢ Do not touch the person with bare hands if he or she is still in contact with

the electrical current.

➢ Do not get near high – voltage wires until the power is turned off. Stay at

least 20 feet away – farther if wires are jumping and sparking.

➢ Do not move a person with an electrical injury unless the person is in

immediate danger.

FAINTING.

Fainting is a sudden and temporary loss of consciousness resulting from a

reduction in blood flow to the brain. The episode lasts less than a couple of minutes

and a victim recovers from it quickly and completely. A victim may feel light headed

or dizzy before fainting. A longer, deeper state of unconsciousness is often called

a coma.

Procedure.

1. Check the person’s airway and breathing. If necessary begin rescue

breathing through cardiopulmonary resuscitation (CPR)

2. Loosen tight clothing around the neck.

3. Raise the person’s feet above the level of the heart (about 12 inches).

4. If the person has vomited, turn onto his or her side to prevent choking.

5. Keep the person lying down for at least 10 – 15 minutes, preferably in a

cool and quiet space. If this is not possible, sit the person forward with the

head between the knees.

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L a b o r a t o r y Te c h n i q u e s a n d S a f e t y

6. Seek immediate medical help if the victim does not recover in a few

minutes.

BLEEDING.

Bleeding is the loss of blood from the circulatory system. Bleeding can occur

internally, where blood leaks from blood vessels inside the body or externally,

either through a natural opening such as the mouth, nose, ear, urethra, vagina, anus

or through a break in the skin. Bleeding may be light or severe. Excessive loss of

blood can cause death.

Light bleeding.

1. Place the victim in a comfortable position.

2. Elevate the injured part

3. Gently clean the wound using clean water and antiseptic or common salt. Cover

the wound using sterile gauze. Gently clean the surrounding skin and dry it

using sterile dressing.

4. Dress the wound and bandage it.

5. Take the person to hospital in case the bleeding continues.

Severe bleeding.

Procedure.

1. Have the injured person lie down and cover the person to prevent loss of body

heat. If possible, position the person’s head slightly lower than the trunk or

elevate the legs and elevate the site of bleeding.

2. While wearing gloves, remove any obvious dirt or debris from the wound. Do

not remove any large or more deeply embedded objects.

3. Apply pressure directly on the wound until the bleeding stops. Use a sterile

bandage or clean cloth and hold continuous pressure for atleast 20 minutes.

Maintain pressure by binding the wound tightly with a bandage or clean cloth

and adhesive tape.

4. Don’t remove the gauze or bandage. If the bleeding continues and seeps through

the gauze or other material you are holding on the wound, don’t remove it,

instead add more absorbent material on top of it.

5. Squeeze a main artery if necessary. If the bleeding does not stop with direct

pressure, apply pressure to the artery delivering blood to the area.

6. Immobilize the injured body part once the bleeding has stopped. Leave the

bandages in place and get the injured person to the emergency room as soon

as possible.

7. Seek medical help immediately.

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Nose bleeding.

Procedure.

1. Let the victim sit upright and lean forward. Remaining upright reduce blood

pressure in the veins of the nose. Sitting forward helps to avoid swallowing

blood. Swallowing blood can irritate the stomach.

2. Let the victim pinch his/her nose by using the thumb and index finger. Let him/

her breath through the mouth and continue to pinch for five to ten minutes.

Pinching sends pressure to the bleeding point on the nasal septum and often

stops the flow of blood.

3. Place a wet piece of cloth at the back of the victim’s neck.

4. Let the victim prevent re-bleeding by avoiding picking or blowing the nose.

Advice the victim not to bend down for several hours after the bleeding episode.

5. If re-bleeding occurs, take the victim to hospital.

POISONING

A poison is a substance which, if taken into the body in sufficient quantity,

may cause temporary or permanent damage. Poisons can be swallowed, absorbed

through the skin, inhaled, splashed into the eyes or injected. Once in the body, they

may enter blood stream and can be carried swiftly to all organs and tissues.

Signs and symptoms of poisoning include burns or redness around the mouth

and lips from drinking certain poisons, breath that smell like chemicals, vomiting,

difficulty breathing, sleepiness and confusion.

Procedures

1. Call for medical assistance immediately

2. If the person has been exposed to poisonous fumes such as carbon monoxide,

get him or her into fresh air immediately.

3. If the person swallowed the poison, remove anything remaining in the mouth.

4. If the suspected poison is a household cleaner or other chemicals, read the label

and follow instructions for accidental poisoning.

5. If the poison spilled on the person’s clothing, skin or eyes, remove the clothing.

Flush the skin or eyes with cool or lukewarm water.

6. Make sure the person is breathing. If not, start cardio pulmonary resuscitation

(CPR) and rescue breathing.

7. Take the poison container (or any pill bottle) with you to the hospital.

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VOMITING.

Vomiting is the forceful expulsion of the contents of the stomach through the

mouth sometimes the nose. Vomiting may result from many causes, ranging from

gastritis or poisoning to brain tumors, or elevated intracranial pressure. The feeling

that one is about to vomit is called nausea, which usually precedes, but does not

always lead to vomiting.

Procedure.

1. Reassure the victim and advise him or her to take slow and deep breaths.

2. Use a damp cloth to wipe the face of the victim

3. Let the victim drink gradually large amounts of clear liquids, including

oral rehydration drink.

4. Do not give the victim solid food until vomiting has stopped.

5. Provide the victim with bananas, rice, apple sauce without sugar, toast,

potatoes and pasta after 24 hours.

6. Get medical assistance if

a. Vomiting goes on for longer than one day.

b. There is blood in the vomit.

c. The vomiting is occurring because of a known injury, like head

trauma or infection.

d. The victim acts confused, lazy or lethargic and is less alert than

usual.

e. Vomiting and diarrhea are present.

f. Severe abdominal pain is present.

g. Severe headache or stiff neck is present.

h. Feelings of nausea last for longer than one week.

FALLING

Falling is an act of moving from a higher level to a lower level rapidly and

without control. Falling may result after loss of balance and collapsing.

Procedure.

1. Cover the victim’s body with a blanket to reduce the chances of going into

shock.

2. Look for signs of a fractured skull such as unequal pupils in their eyes, bleeding

from the ear or clear fluid running from the nose.

3. Check for swelling or limbs sticking out at an unusual angle. If a bone is broken,

keep it still and wrap it with towels to support it.

4. If you cannot see any broken bones or you do not think the victim have got a

head or neck injury, put the victim in recovery position.

5. If the victim is conscious and there are no obvious signs of serious injury apply

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a cold compress made from a cloth dipped in water to the areas that were hit

for 10 minutes to help with swelling.

6. If the victim is unconscious seek medical help immediately.

SUMMARY

(a). A laboratory is a special room or building that is designed and used for

scientific experiments.

(b). Apparatus are special tools and equipment that are used in the laboratory.

(c). Laboratory rules are specified guidelines required to be followed when

working with the laboratory.

(d). Chemical warning signs are safety symbols found on chemical containers,

especially those used in the laboratory.

(e). First aid is the help given to a sick or injured person before medical assistance from the hospital.

(f). First aid kit is a box that contains equipments and chemicals needed for the

first aid.

REVIEW QUESTIONS.

1. Choose the best answer from the choices given.

i. Temperature can be measured with

(A) Vernier caliper

(B) Measuring cylinder.

(C) Measuring syringe.

(D) Thermometer.

ii. ………… is good for measuring volumes quickly but not very accurately.

(A) Measuring cylinder.

(B) Beaker

(C) Volumetrick flask.

(D) Pipette.

iii. ………………….. Is used to make up a solution to a total final volume.

(A) Burette.

(B) Gas syringe.

(C) Volumetric flask.

(D) Conical flask.

iv. …………….. Can be used to measure the volume of a gas produced during

an experiment.

(A) Gas syringe.

(B) Round bottomed flask

(C) Flat bottomed flask

(D) Volumetric flask.

v. A condenser is used to cool gases down during …… so they will condense

and form a liquid.

(A) Evaporation

(B) Decantation

(C) Filtration

(D) Distillation.

30

L a b o r a t o r y Te c h n i q u e s a n d S a f e t y

2. Indicate TRUE or FALSE for the following statements.

i. A mortar and pestle is used to grind lumps into a powder………………..

ii. The funnel is used when transferring a liquid from one container to

another………….

iii. A pipe clay triangle fits on the top of a tripod and supports a crucible when

it is heated……………

iv. A combustion spoon is used to burn a large amount of a solid, perhaps

sulphur in a gas jar of oxygen………….

v. Students are advised to pick up reagent bottles by holding the

stopper………………..

3. Match the items in List A with their corresponding statements in LIST B.

List A

List B

i. Harmful

A. This substance is dangerous and can cause death.

ii. Irritant

B. This substance explodes easily

iii. Radioactive

C. This substance catches fire easily

iv. Corrosive

D. This substance reacts easily with oxygen.

v. Toxic

E. Corrodes surfaces as well as the human body.

vi. Oxidant

F. Emits harmful radiations.

vii. Flammable

G. Irritate parts of the body.

viii. Explosive

H. This substance is harmful or poisonous

I. Emits harmful radiations and reacts easily with oxygen.

J. Catches fire and explodes easily.

4. Define first aid and first aid kit respectively.

5. Some beakers, syringes, measuring cylinders and troughs are made of plastic.

What are the advantages of plastic containers.

6. Chemical apparatus can be divided into that which holds chemicals (containers)

and that used to support other pieces of apparatus (supporters) copy and

complete the following table putting in the names of the following pieces of

apparatus. Flask, tripod stand, test tube rack, test tube holder, beaker, trough.

Containers

Supporters

7. The following list contains the names of pieces of apparatus with the letters

in each name jumbled up. Rearrange the letters to produce the names of ten

pieces of apparatus.

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I l l u s t r a t i v e C h e m i s t r y f o r S e c o n d a r y S c h o o l s

A.

Lupstaa

B.

Lefnum

C.

Tptipee

D. Romtmreehet

E. Bulceric

F. Faskl

G.

Tertbeu

H.

Rebake

I.

Gotsn

J. Gothur

From the list you have obtained.

i. Name seven pieces of apparatus made of glass.

ii. Name four pieces of apparatus which can be used as containers.

iii. Name two pieces of apparatus which can be used to measure volumes

of liquids.

iv. Name one instrument that measures temperature.

8. Copy and complete the following table.

Apparatus

Made of

Uses

Gas jar

Measuring cylinder

Funnel

Tongs

Forceps

Watch glass

Evaporating basin

Burette

Test tube rack

Spatula

32

Chapter 3

Heat Sources

and Flames

3.1 HEAT SOURCES

Heat is energy transferred from one body to another by thermal interactions.

It is very useful in our day to day life. It is important to have good sources of

heat in the laboratory. Chemical reactions which would proceed slowly at normal

temperatures will take place faster when they are heated.

Heat may be obtained from electricity or from natural fuels like coal, charcoal,

gas, kerosene, spirit etc. For each type of fuel used, a different kind of burner is

needed. The common types of burners used in laboratories are: Bunsen burner, spirit

burner, kerosene burner, gas stove, charcoal burner and candle. The Bunsen

burner is the best of them all because it is convenient to handle and it produces a

range of flames including a very hot flame (approximately 1000˚C).

A spirit burner uses methylated spirit as a fuel. Candle may be used where the

experiment does not need heating to a high temperature. Burning candle produces

a lot of soot and smoke. A kerosene burner can reach high temperatures (1000˚C

or more). A gas stove uses natural gas as a source of fuel. If the Bunsen burner

and kerosene burner are not available, a charcoal burner is the next choice simply

because a charcoal burner can burn for a long time. The burner can produce high

temperatures. A tin lamp is a small kerosene burner which can be used if soot is not

a major handicap. This burner can supply a high temperature with use of blow pipe.

A sand bath is common piece of laboratory equipment made from a container

filled with heated sand. It is used to provide even heating for another container,

most often during a chemical reaction.

A water bath can be used to keep a reaction on vessel at the temperature of

boiling water until all water is evaporated. A thermometer is usually kept in the

water to monitor the temperature.

I l l u s t r a t i v e C h e m i s t r y f o r S e c o n d a r y S c h o o l s

33

I l l u s t r a t i v e C h e m i s t r y f o r S e c o n d a r y S c h o o l s

Spirit lamp

Bunsen burner

Kerosene burner

Charcoal burner

Gas stove

Tin lamp

Sand

Water

Metal dish

Heat

Heat

Water bath

Sand bath

Fig. 3.1. Heat sources in the laboratory

34

H e a t s o u r c e s a n d f l a m e s

3.2 THE BUNSEN BURNER

A Bunsen burner is a small laboratory heat source consisting of a vertical metal

tube connected to a gas source and producing a very hot flame from a mixture of

gas and air let in through adjustable holes. These holes allow air to enter the tube

and mix with the gas in order to make a very hot flame.

The gas from the source can be natural gas (which is mainly methane) or a

liquefied petroleum gas such as propane, butane or a mixture of both. The Bunsen

burner is named after a German chemist and physicist, Robert Wilhelm Bunsen,

who invented it in 1855. Consider the diagram of a Bunsen burner below;

Chimney/barell

Collar/metal ring

Jet

Gas

Base

Fig: 3.2 Bunsen burner.

HOW THE BUNSEN BURNER WORKS

The gas enters the burner through a tube connected to a jet inside the base.

Air enters the burner through the air holes at the base. The amount of air coming

in can be varied by turning the collar. At the top of the barrel, the mixture of air

and gas burns.

Activity 3.1; to light and adjust a Bunsen burner.

1. Dismantle the Bunsen burner. Start by unscrewing the barell or chimney. Next

remove the collar or metal ring. Use diagram to identify the parts.

2. Reassemble the Bunsen burner.

3. Carry out the following instructions for lighting a Bunsen burner.

a. Connect the Bunsen burner by the rubber tube to the gas mains.

b. Close the air hole.

c. Allow plenty of gas to enter the burner by turning the gas tap on.

35

I l l u s t r a t i v e C h e m i s t r y f o r S e c o n d a r y S c h o o l s

d. Immediately bring a flame to the top of the barrel. You may use a lighted

taper or a match stick. You will get a tall shapeless yellow flame. This is

called the luminous or yellow flame.

e. Adjust the gas supply until the flame is about 10cm high.

f. Turn the collar until the air hole is fully open. You now have the non-

luminous or blue flame. For safety reasons a lit Bunsen burner should

always be left on a luminous flame.

3.3 TYPES OF FLAME

A flame is a zone of burning gases that produces heat and light. It is visible gaseous part of a fire. The flame is formed as a result of burning a fuel. The colour

and temperature of flame depend on the type of fuel burning and the source of the

flame.

Flames can be classified as luminous or non-luminous. A luminous flame is yellow. Although it uses oxygen when burning, the oxygen supply is usually not

enough to completely burn up the fuel. Therefore it produces a black substance

known as soot.

The non-luminous flame is blue. This is because there is an adequate supply

of oxygen and so the fuel burns efficiently. Such a flame doesn’t produce any soot.

It also produces more heat than the luminous flame.

Different heat sources produce different types of flames. Example, the candle

and the tin lamp produce luminous flames, where as the gas stove produces a non-

luminous flame.

On lighting the Bunsen burner, the type of flame obtained depends on how

much oxygen is available for burning. When the air holes are closed, there is less

oxygen and the flame is luminous. When the air holes are partly open, there is more

oxygen and the flame is medium. When the air holes are fully open, there is enough

oxygen and the flame is non-luminous.

When the air holes of the Bunsen burner are closed, the flame is wavy and

yellow. This luminous flame is called a safety flame because it is easily seen and so

is less likely to cause accidents. When the air holes of a Bunsen burner are half open, the flame becomes difficult to see. This type of flame can easily cause accidents in

the laboratory since a person may not be aware that the burner is on. When the air

holes of a Bunsen burner are fully open, the flame is cone shaped and burns with

a roaring sound.

36

H e a t S o u r c e s a n d f l a m e s

Consider below diagrams for luminous and non-luminous flames.

Thin outer zone

Luminous and yellow zone

Dark zone of unburnt gas

Blue Zone

Fig. 3.3 Luminous flame

Pale blue or purple outer zone

Blue green middle zone

Dark zone of unburnt gas

Fig. 3.4 Non-luminous flame

37

I l l u s t r a t i v e C h e m i s t r y f o r S e c o n d a r y S c h o o l s

3.4 CHARACTERISTICS OF LUMINOUS AND NON-

LUMINOUS FLAMES

Table 3.1; characteristics of luminous and non-luminous flames.

Luminous flames

Non - luminous flames

i. Yellow in colour.

i. Blue in colour.

ii. Produces soot.

ii. Does not produce soot.

iii. Has a wavy flame.

iii. Have a triangular flame.

iv. Produces less heat.

iv. Produces more heat.

v. Burns quietly.

v. Burns with a roaring noise.

vi. Produced when air holes of a Bunsen

vi. Produced when air holes of a Bunsen

burner are closed.

burner are open.

vii. Has four zones.

vii. Has three zones.

USES OF LUMINOUS AND NON - LUMINOUS FLAME

Luminous Flame.

The luminous flame is mainly used for lighting. It is yellow flame that easily

brightens a room. Since it is not very hot, it is safer for lighting than the non-

luminous flame. Some heat sources that can produce a luminous flame are candle,

tin lamp and hurricane lamp.

Non-Luminous Flame.

The following are uses of non-luminous flame.

1. Non-luminous flame is used for heating purposes because it gives a lot

of heat. Most heating in the laboratory is done using a Bunsen burner.

Non-luminous flame of a Bunsen burner has three parts; the pale blue or

purple outer zone, the blue green middle zone and the inner dark zone of

unburnt gas.

2. The hottest part of the non-luminous flame is at the tip of the middle zone.

3. A non-luminous flame is used in the flame test of certain chemical

substances. A flame test involves introducing a sample of the desired

substance to the non-luminous flame and observing the colour change.

The colour can indicate what components the substance is made of.

4. Welding is the joining together of metal pieces or parts by heating their

surfaces and pressing them together. A non-luminous flame is suitable for

this task because it is very hot.

5. A non-luminous flame is suitable for cooking since it gives off enough

heat and it does not produce soot.

38

H e a t s o u r c e s a n d f l a m e s

Activity 3.2

To compare the heating ability of luminous and non-luminous Bunsen flames.

1. Measure 100 cm3 of water in a measuring cylinder and pour this into a 250

cm3 beaker.

2. Put a thermometer into the water and take the temperature.

3. Stand the beaker on a tripod and gauze.

4. Light a Bunsen burner and adjust until you have a blue flame 10cm high.

Do not touch the gas tap now throughout the experiment.

5. Put the lit Bunsen burner under the tripod and heat the water.

6. Stir the water constantly with a glass rod and take the temperature every

half minute for 5 minutes.

7. Empty the water away and allow the apparatus to cool.

8. Repeat the experiment using a yellow flame.

9. Draw two graphs on the same grid showing the temperature of the water

during the experiment.

Thermometer

Glass rod

Water

Figure 3.5; Apparatus used to demonstrate the heating ability of luminous and

non-luminous flames.

SUMMARY

(a). Heat is energy transferred from one body to another by thermal interactions.

(b). Common types of burners used in laboratories are Bunsen burner, spirit

burners, kerosene burners, gas stove, charcoal burner and candle.

(c). The Bunsen burner is the best source of heat in the laboratory.

(d). A Bunsen burner is a small laboratory heat source consisting of a vertical

metal tube connected to a gas source and producing a very hot flame from

the mixture of gas and air let in through adjustable holes.

(e). The Bunsen burner is named after a German chemist and physicist, Robert

Wilhelm Bunsen.

(f). Flames can be classified as luminous and non-luminous flames.

(g). Luminous flame has four zones whereas non-luminous flame has three zones.

39

I l l u s t r a t i v e C h e m i s t r y f o r S e c o n d a r y S c h o o l s

REVIEW QUESTIONS

1. Match the items in LIST A with their corresponding statements in LIST B.

List A

List B

i. Spirit lamp

A. Keep a reaction vessel at the temperature of boiling water.

ii. Kerosene

B. Container filled with heated sand.

burner

C. Small kerosene burner which produces lot of soot.

iii. Charcoal

D. Can burn for a long time when lit and produce high

burner

temperature.

iv. Gas stove

E. Uses natural gas as a source of fuel.

v. Tin lamp

F. Does not heat to a high temperature of 1000˚C or more.

vi. Water bath

G. Does not heat to a high temperature.

vii. Sand bath

H. Uses methylated spirit as fuel.

viii. Candle

I. The best source heat in the laboratory.

2. Write TRUE or FALSE for the following statements

i. Heat is very useful to our day to day life,.................

ii. Heat increases the speed of chemical reactions.,.................

iii. The Bunsen burner is named after a German chemist and physicist

Robert Wilhelm Bunsen.,.................

iv. Methane is the good example of natural gas.,.................

v. Air enters the Bunsen burner through air holes.,.................

vi. For safety reasons a lit Bunsen burner should always be left on a

luminous flame.,.................

3. Choose the most correct answer for the following questions.

i. ........................... is a zone of burning gases that produces heat and light.

A. Fuel

B. Heat

C. Flame

D. Combustion.

ii. The colour and temperature of flame depend on:

A. Type of fuel burning and the source of flame

B. Type of fuel burning and heat involved

C. Bunsen burner capacity.

D. Quantity of heat.

iii. A black substance resulted after incomplete combustion.

A. Soot

B. Rust

C. Smoke

D. Flame

iv. When there is an adequate supply of oxygen:

A. The fuel burns slowly

B. The fuel burns smoothly

C. The fuel burns efficiently.

D. The fuel produces soot.

40

H e a t s o u r c e s a n d f l a m e s

v. ........................and .......................... produces luminous flames

A. Kerosene burner and blue flame of burner

B. Candle and gas stove

C. Candle and tin lamp

D. Candle and Bunsen burner

4. Outline the differences between luminous and non-luminous flames

5. Draw well labeled diagrams of luminous and non-luminous flames respectively.

6. Briefly answer the following questions

a. Name six sources of heat that can be used in the laboratory.

b. Name two sources of heat that can produce a luminous flame

c. Describe the process of lighting of Bunsen burner.

d. The non-luminous flame is used in different areas. Give four such areas

and explain why the flame is preferred for those uses

e. Give two reasons why a luminous flame not a non-luminous flame, is

preferred for lighting.

41

Chapter 4

Scientific Procedure

The scientific procedure is the organized set of guidelines used for investigating

problems, acquiring new knowledge or correcting and integrating previous

knowledge. The main characteristic which distinguishes the scientific method from

other methods of acquiring knowledge is that scientists seek to let reality speak for

itself, support a theory when a theory’s predictions are confirmed and challenging

a theory when its predictions prove false.

4.1. SIGNIFICANCE OF THE SCIENTIFIC PROCEDURE

We carry out experiments to find answers to scientific questions. The findings

from the experiments help us to solve scientific problems. The following are some

of significance resulted from application of scientific procedure;

➢ The scientific procedure allows for

the side of the experimenter.

one to prove, or disprove a theory

➢ The scientific procedure provides

or hypothesis. It allows for results

an objective, standardized approach

from an experiment to be repeated

to conducting experiments and

or validated.

improve the results.

➢ The scientific procedure is a rational

➢ The scientific procedure is used by

method for examining the universe

scientists to answer questions and

around us, making observations

solve complex problems.

about it, and hypothesizing rules or

➢ The scientific procedure help

models for how it operates.

scientists to detect changes in our

➢ The scientific procedure attempts to

environment and suggest ways in

minimize the influence of bias on

which the effects may be avoided.

4.2 MAIN STEPS OF THE SCIENTIFIC PROCEDURE.

The scientific method is a logical way to solve a problem. It is a process

that is used to find answers to questions about the world around us. The scientific

method is used by researchers to support or disprove a theory. Below are steps of

the scientific procedure.

I l l u s t r a t i v e C h e m i s t r y f o r S e c o n d a r y S c h o o l s

42

S c i e n t i f i c P r o c e d u r e

1. Identification of the problem.

4. Data collection and analysis.

2. Formulation of Hypothesis.

5. Data interpretation.

3. Experimentation and

6. Drawing conclusion.

Observation.

1. IDENTIFICATION OF THE PROBLEM

A problem statement is a question that compares variables. Careful observation

leads to this question. This is when scientists raise questions based on what they

have observed. When writing the problem statement think what you want to know

or explain. Use observation you have made to write a question about the problem

or topic you want to investigate. An example of a problem could be “does eating

of greasy food cause pimples?”

2. FORMULATION OF HYPOTHESIS

A hypothesis statement is a statement that expresses what the expected answer

to the experiment will be. This is what you think the results of the experiment will

show. A hypothesis is an intelligent guess which predict the answer to the problem

at hand. An example of a hypothesis could be “eating greasy food cause pimples”

3. EXPERIMENTATION AND OBSERVATION

An experiment is a planned way to test a hypothesis and find out the answer to

the problem statement. To conduct an experiment a scientist develops and follows

a procedure (steps to the experiment). The procedure also includes a detailed materials list.

In the experiment there are things that might change: These are called variables.

There are three types of variables. These are:

➢ Dependent variable - This is

as to obtain different values for

the factor in the experiment that

comparison.

changes its value when the values

➢ Controlled variable - This is the

of the other variables change. It is

factor in the experiment that does

the factor being measured.

not change, or is kept constant. It

➢ Independent variable - This is

the factor that is manipulated so

does not affect the outcome of the

experiment.

In the experiment on investigating whether eating greasy food causes pimples

or not:

➢ Pimples are the factor under investigation. It is the dependent variable.

➢ Eating greasy food is the factor that is varied. It is the independent variable.

➢ The number of people involved as experimental sample is kept constant. It is

the controlled variable.

43

I l l u s t r a t i v e C h e m i s t r y f o r S e c o n d a r y S c h o o l s

It is often important to carry out an experiment where all the independent

variables are kept constant. This is called a control experiment.

A control experiment is an observation designed to minimize the effects of

variables other than the single independent variable.

Example: In the experiment on whether eating greasy food causes pimples or not, there can be two equal groups of people where by one group will be provided

with greasy food while the other group won’t be provided with greasy food but

rather food which is greasy free. The group provided with greasy free food is the

control experiment.

4. DATA COLLECTION AND ANALYSIS

Data is the information you get when you test the variable. Confirm the results

by retesting and modify the procedure if needed. Did you get the same results each

time you retested? If not, why? Is there a change you can make that would make

results more consistent?. Include tables, graphs and photographs that show your

data.

Consider the table below;

Table. 4.1; number of people with pimples at different food amounts

Amount (grams) of greasy food

Number of people with pimples

50

150

200

250

300

350

400

450

500

5. DATA INTERPRETATION

After recording and analyzing the data, you look for trends or patterns and

explain why they occur that way. This is what is known as interpretation of data.

These trends will help you to make your conclusion.

Example: in the above experiment, it may be noticed that as you increase the amount of greasy food the number of people with pimples increases. This is because

increasing the amount of greasy food increases the effect of having pimples.

44

S c i e n t i f i c P r o c e d u r e

6. DRAWING A CONCLUSION

A conclusion statement is a statement that presents the findings of an experiment,

what the data shows, and states if the hypothesis was correct (supported) or incorrect

(negated). Conclusion statements also make recommendations for further study and

possible improvements to the procedure.

Example, in the experiment on whether the greasy food causes pimples or not, you may notice that increasing the amount of greasy food increases the number of

people with pimples.

From this experiment you may conclude that “eating greasy food cause

pimples” This supports and approves the formulated hypothesis.

REPORTING RESULTS

Scientists usually communicate by publishing their findings in scientific

journals and other publications. They can also present their results to the scientific

community at seminars, symposia and meetings.

Consider the below figure that shows a summary of scientific procedure.

Ask

questions

Do

background research

Construct

Think try

hypothesis

again

Test with an

experiment

Analyze

results draw

conclusion

Hypothesis is true

Hypothesis is false or

partially true

Report

results

Figure 4.1; steps in scientific procedure.

45

I l l u s t r a t i v e C h e m i s t r y f o r S e c o n d a r y S c h o o l s

4.3 APPLICATION OF THE SCIENTIFIC PROCEDURE

The scientific procedure have proven to be so fruitful and it is applied in many

fields. The following are the common areas where it is applied.

1. In hospitals and health centres.

Diagnosis and treatment of diseases mainly depend on scientific procedure.

Systematic steps undertaken by a medical practitioner leads to identification and

subsequent treatment of illness in general.

2. In project work.

Careful and persistent studying of a problem over a period of time leads to

planning of a project work thereby solutions to a problem can be found.

3. In field work

Field work is an investigation or search for material or data, made in the field

as opposed to the classroom, laboratory or official headquarters. In field work,

finding answers to problems and testing of hypothesis are important elements.

4. In research centres

A research centre is an establishment endowed for doing research. Research

centres may specialize in basic research or may be oriented to applied research.

Scientific research is the term mainly used to activities undertaken in research

centres. Scientific research involves scientific methods.

5. In school laboratory

A laboratory is a special room or building designed and applied for scientific

experiments. Scientific experiments relies heavily on scientific procedures.

SUMMARY

(a). The scientific procedure is the organized set of guidelines used for

investigating problems, acquiring new knowledge or correcting and

integrating previous knowledge.

(b). We carry out experiments to find answers to scientific questions.

(c). Main steps of the scientific procedure are: identification of the problem,

formulation of hypothesis, experimentation and observation, data collection

and analysis, data interpretation and drawing conclusion.

(d). The three types of variables are dependent variable, independent variable

and controlled variable.

(e). The scientific procedure has proven to be so fruitful and it is applied in many

fields.

46

S c i e n t i f i c P r o c e d u r e

REVIEW QUESTIONS

1. Choose the most correct answer for the following questions;

vii. ................... is the organized set of guidelines used to answer questions or solve problems.

A. Scientific experiment

B. Scientific research

C. Scientific procedure

D. Hypothesis

viii. The second step of the scientific procedure is ...................

A. Data interpretation.

B. Drawing conclusion

C. Indentificaton of the problem.

D. Formulation of hypothesis.

ix. ...................is a suggested answer to a question or problem.

A. Hypothesis B. Diagnosis C. Prognosis

D. Inquiry

x. A test done to find out if the formulated hypothesis is true or false is called

A. Research B. Experiment C. Observation D. Investigation.

xi. ...................are factors that affect the problem being investigated.

A. Variance

B. Variables

C. Objects

D. Data.

2. In an experiment to find out what causes rusting of iron you might carry out a

control experiment. How would you do this and why would you do it?

3. What are the advantages of following scientific procedure over trial and error

methods?

4. Write TRUE or FALSE for the following statements.

i. If the results you obtain from an experiment do not support your hypothesis,

give ideas for further testing to find a solution.

ii. The sixth step in the scientific procedure is drawing conclusion.

iii. Factors in an experiment that can be manipulated to get desired results are

called independent variables.

iv. A theory can be derived without any experimental results.

v. A law can only be claimed when all scientists agree with the experiment

results and procedures.

47

Chapter 5

Matter

5.1 CONCEPT OF MATTER

Matter is anything that occupies space and has mass. Mass is a measure of a quantity of matter in an object. Mass is usually measured in kilograms or grams.

Matter is also a general term for the substance of which all observable physical

objects consists.

5.2 STATES OF MATTER

Properties describe matter. A block of wood, milk and air all have properties.

All the material on earth exist in three states; solid, liquid and gas. The state of

matter refers to the group of matter with the same properties.

SOLID

A solid has a definite size and shape. A solid is usually hard and not easily

deformed. Examples of solids are wood, computer, desk and floor.

LIQUID

Liquids do not have a definite shape. Liquids have definite volume and take

the shape of its container. Liquids can flow, be poured and spilled. Examples of

liquids are milk, water and juice.

GAS

A gas is matter that has no specific shape or size of its own. Some one can feel

gas when the wind blows. The wind is moving air. Air is many gases mixed together.

Gases occupy all the space in the container holding them. Examples of gases are

oxygen, hydrogen and chlorine.

CHANGES OF STATES OF MATTER

Matter can change its state from one form to another .Matter exist in three

states: Solid, Liquid and Gas. Changes of states of matter are caused by alterations

in temperature and pressure.

I l l u s t r a t i v e C h e m i s t r y f o r S e c o n d a r y S c h o o l s

48

M a t t e r

MELTING AND FREEZING

Melting is the change from solid to liquid state .When a solid is heated the

particles gains energy and vibrate faster. Eventually, they break free from their

fixed positions and begin to move in clusters. A solid will change to a liquid at the

melting point. The temperature at melting point must remain constant until all the

solid has changed to liquid.

The melting point of a solid tells us how strongly its particles are held together.

Substances with high melting points have strong forces between their particles.

Those with low melting points have weak forces between their particles.

Freezing is the opposite of melting. For any particular substance, the melting

point is the same as its freezing point. The freezing point of pure water, 0˚C, is the

melting point of ice. It is not affected by atmospheric pressure.

BOILING AND EVAPORATION.

Boiling is the change of a liquid to vapour. When a liquid boils, it evaporates.

When a liquid is heated, the molecules gain energy and move faster. Some of the

molecules at the surface of the liquid attain enough energy to escape into the air. As

more of the liquid molecules escape to form a gas, the liquid is said to evaporate.

This happen at the boiling point of the liquid.

The temperature at which a liquid boils tells us how strongly the particles are

held together in the liquid. Liquids with high boiling points have stronger forces

between their particles, than liquids with low boiling points.

Consider the below figure which shows the relationship between the three

states of matter.

GAS

Reverse sublimation

Sublimation

ation

Condensation apor

Ev

LIQUID

Melting

SOLID

Freezing

Fig. 5.1; relationship between the three states of matter.

49

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Condensation, reverse sublimation and freezing are exothermic processes

where by evaporation, sublimation and melting are endothermic processes.

Exothermic processes gives off heat energy while endothermic processes

requires heat energy.

Experiment 5.1

Aim: To observe the change of matter from one state to another

Apparatus: condenser, source of heat, beaker. Other materials: ice cubes.

Procedure

1. Put some ice in a beaker and heat. Carefully observe the change that

take place.

2. Heat the substance further until it changes into vapour.

3. Pass the vapour through a condenser. Record your observations.

Discussion questions

1. How do ice look like before heating?

2. Explain the shape of ice when it melts and when it vapourizes.

3. Explain what happens when vapour is passed through a condenser.

PARTICULATE NATURE OF MATTER.

For many years scientists were not sure whether matter was continuous or

particulate. Later, it was discovered that matter is made up of particles. This was

proved by a phenomenon known as Brownian motion. It can be observed in liquid and gaseous substances. Therefore it is now known that matter is not continuous.

Matter is made up of many tiny particles put together. The tiny particles

which make up matter are either ions, atoms or molecules. We can not see atoms or

molecules by our eyes. They are very small. We can prove that matter is particulate

by doing experiments. If you open a bottle containing a strong perfume, the smell

of perfume will soon be detected in other parts of the room. If you weigh a bottle

of perfume before and after opening a bottle you will not be able to detect any mass

change. The particles that were lost must be very small as there is no detectable

mass change.

You may dissolve some sugar in water. The sugar seems to disappear

completely. When you taste the water, it will be sweet, meaning that the sugar is

still there. But in the water, the sugar is divided into such small particles that we

can’t see them with our eyes.

BROWNIAN MOTION

Brownian motion is an irregular motion of tiny particles suspended in a fluid.

In 1827, a Botanist called Robert Brown observed through microscope that pollen

grains suspended in water moved short distances in an irregular manner.

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M a t t e r

Experiment 5.2

Aim: to observe Brownian motion

Materials: Pollen grains, distilled water, microscope, microscope slide, dropper.

Procedure:

1. Put a drop of distilled water on a microscope slide.

2. Put some pollen grains on top of the drop of water.

3. Place the slide under a microscope and observe any movement of the pollen

grains.

Discussion questions

1. What kind of movement have you observed?

2. Why do the grains move this way?

Pollen grains move in a random zigzag motion. This is because they are

constantly bombarded by water particles which cannot be seen even through a

laboratory microscope. This shows that matter is particulate in nature and not

continuous in form.

Activity 5.1: To show that a coloured compound is made up of small coloured

molecules.

Procedure:

1. Dissolve about 1.0g of purple potassium manganate (VII) crystals in

1dm3 of water. Stir until it all dissolves. Label it solution A (figure 5.2).

2. Measure 1000 cm3 of solution. Label this solution B.

3. Repeat the process four more times, taking 100 cm3 of the subsequent

dilute solution at each step. You will finally get solution samples of

different colour intensity A,B,C,D,E and F (Figure 5.2).

1.0g of potassium manganate (VII)

100 cm3

100 cm3

100 cm3

100 cm3

100 cm3

A

B

C

D

E

F

figure 5.2; observing the colour change when a solution of potassium

manganate (VII) is diluted several times over.

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When potassium manganate (VII) dissolved, its particles spread throughout

the water. The resulting solution was dark purple. When the purple solution was

diluted, the particles spread further apart, and there were fewer in each subsequent

diluted solution. As the solution is diluted the colour gets less intense paler and

paler pink. But you will still see a pink colour. Hold piece of white paper behind

the solution if you are not sure.

The experiment shows that a small amount of a coloured substance can

give colour to a large volume of water. It also suggests that molecules are extremely

small.

Activity 5.2: to show that liquids are made of particles

Procedure

1. Place a few drops of ether on a watch glass and observe what happens

after a short time.

2. Spray perfume at one corner of the classroom.

Discussion questions

1. Can you detect the smell of ether from a distance?

2. If you stay in any corner of classroom can you detect the smell of perfume?

Spread of the ether and perfume so that they are detected further away from

their original positions is called diffusion. Diffusion is the movement of particles from an area of high concentration to the area of low concentration. Therefore

liquids are made of particles.

KINETIC NATURE OF MATTER.

Enough evidence shows that matter exists as particles. These particles are

atoms, ions or molecules, and are in constant motion. All matter is always in motion.

Solids appear to be stationary, but their particles keep on vibrating all the time. At

high temperatures, the movement of particles is faster. Temperature changes may

alter the state of matter. Matter exists in three states; solids, liquids and gaseous. At very low temperatures, matter tends to exist as solid. As temperature rises, solids

change to liquids, and liquids to gases. The reverse is possible when temperature

is lowered.

The gaseous state is one in which the particles are moving independently of

each other. They move in all directions and at great speeds. Gas particles maintain

large distances from each other. They collide as they move. A gas will speed out to

fill all the space in which it is contained. A gas cannot have any shape of its own.

It is possible to compress a gas because the particles are far apart.

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M a t t e r

Particles in a liquid are much closer than those in a gas. Forces of attraction

exist between the molecules. Molecules in a liquid can move in all directions but

not as fast as gas molecules. A liquid may have a definite volume, but does not have

a definite shape. A liquid will always take the shape of its container.

Particles in a solid are very close together. The forces of attraction between

solid particles are large. Movement of particles from one place to another is not

allowed. Rotations and vibrations of solid particles are allowed. In a solid the

particles are arranged in a fixed pattern. A solid can therefore have fixed shape.

Strong force must be applied to alter the shape. You can’t compress a solid but you

can hammer it into a different shape.

Consider the below figure which shows particles in solid, liquid and gaseous

state

Solid state

Liquid state

Gaseous state

Figure 5.3 particles in solid, liquid and gaseous state

The way in which particles behave in solids, liquids and gases is called kinetic

molecular behavior.

Activity 5.3 to demonstrate the kinetic nature of matter.

Procedure:

1. In groups of about twenty, sit closely together so that your shoulders

touch.

2. Can you individually move about freely?

3. Move into any empty classroom and walk about freely. How far are

you able to move around?

4. Go into the field and move about. How far are you able to move?

Discussion question

Which of your movements in the above steps compares to movement in solids,

liquids and gases respectively?

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Students’ movement is limited when they sit close together. This can be

compared to particles of a solid that can only vibrate but remain in their respective

positions. When the students move about in a classroom, they have a measure of

freedom to move, but cannot go beyond the walls of the room. Their movement is

similar to that of particles in a liquid, which can move about freely, but only within

the container holding the liquid.

The students in the field can move far apart from each other within the field.

If the field is not fenced, they can move beyond its boundaries. This is similar to

the movement of gaseous particles which are far apart from each other and fill the

container holding the gas. If the container is opened, the gas particles move out of it.

5.3 PHYSICAL AND CHEMICAL CHANGES.

Physical changes are temporary changes in the physical properties of a

substance. Chemical changes are permanent changes in the chemical properties of

a substance.

CHARACTERISTICS OF PHYSICAL AND CHEMICAL

CHANGES.

Physical changes are concerned with energy and states of matter. A physical

change does not produce a new substance. Changes in state or phase (melting,

freezing, evaporation, condensation, sublimation) are physical changes. Examples

of physical changes include crushing a can, melting an ice cube and breaking a

bottle.

Chemical changes take place on the molecular level. A chemical change

produces a new substance. Examples of chemical changes include combustion

(burning ), cooking an egg, rusting of iron, mixing hydrochloric acid and sodium

hydroxide to make salt and water.

Consider the below table which shows examples of physical and chemical

changes.

Table 5.1; examples of physical and chemical changes

Physical changes

Chemical changes

Melting an ice cube.

Burning wood

Crumpling a sheet of paper

Ripening of green tomato

Casting silver in a mold

Milk goes sour

Aluminum foil is cut in half

Jewelry tarnishes

Butter melts on warm toast

Bread becomes toast

Water evaporates from the surface of the

Rust forms on a nail left outside.

ocean

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Physical changes

Chemical changes

A juice box in the freezer freezes.

Gasoline is ignited.

Dissolving salt in water

Food scraps are turned into compost in a

compost pile.

Water boiling

You take an antacid to settle your

stomach

Cutting carrots

Digestion of food

Grinding meat to make hamburger

Frying an egg

Frost forming on the window

Creation of a metal alloy

Fog forming at night

Baking a cake

Making tea

A battery turning on flash light

Dissolving sugar in water

Mixing vinegar with baking soda.

Milo dissolves into hot milk

Fireworks explode to form colorful light

and loud sounds.

A plate is dropped and shatters

An egg is cooked to become a white and

yellow solid.

Grass is mowed

Spoiling food

Metal knife is sharpened

Burning leaves

Breakfast cereal goes soggy

Burning of candle

Heating an iron bar until it glows red hot.

Burning of camphor

Melting of wax

Rancidification of butter

Heating of zinc oxide

Photosynthesis

Heating of camphor

Combustion of fuel

Heating of ammonium chloride

Electrolysis of water.

Dissolution of sulphur in carbon disulphide Hydrogen burning in oxygen

Cutting of wood

Decomposition of potassium chlorate.

Experiment 5.3

Aim: To demonstrate types of changes in matter.

Apparatus: Beaker, evaporating dish, source of heat, crucible.

Other Materials: Sugar, water, candle, wax, paper.

Procedure;

1. Pour water into a beaker. Dissolve some sugar in it. Pour the solution in an

evaporating dish and evaporate the solvent. Let the resultant product cool.

What is the resultant product? Does it look like the sugar you dissolved?

2. Heat some candle wax in a crucible until it melt. Let it cool. Is the final

product different from the initial one?

3. Burn a piece of paper until it is completely burnt. What is the final product?

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Is it the same as the initial substance?

4. Put some sugar in a crucible, heat and let it melt completely until it changes

color. Is the final substance the same as the initial product?

Discussion questions

1. Explain which of the changes observed were physical and chemical

respectively?

2. What are characteristics of physical and chemical changes?

Consider the below table which shows differences between physical and

chemical changes.

Table 5.2 Differences between physical and chemical changes;

Physical change

Chemical change

Alters only the physical properties but

Alters the chemical properties and molecular

the molecular composition remains

compositions.

totally unaltered.

No new substances are formed as a

A new substance is always formed as a result

result of physical change

of chemical change.

The change is temporary

The change is permanent

The change is reversible

The change is irreversible

No energy change take place as a result Energy changes take place as a result of

of physical change

chemical change.

Mass of a substance remains total y

Mass of substances gets total y altered during

unaltered

chemical change.

SUMMARY

(a). Matter is anything that occupies space and has mass.

(b). Mass is a measure of a quantity of matter in an object.

(c). Matter exists in three physical states of gas, liquid and solid.

(d). Matter can change its state from one form to another.

(e). Matter is particulate and not continuous in form.

(f). Brownian motion is an irregular motion of tiny particles suspended in a fluid.

(g). Diffusion is the movement of particles from an area of high concentration to

the area of low concentration.

(h). All matter is always in motion.

(i). Physical changes are temporary changes while chemical changes are

permanent changes.

(j). At high temperature, the movement of particles is faster while at low

temperatures the movement of particles is slow.

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REVIEW QUESTIONS

1. Choose the most correct answer for the following questions;

i. ..................... Is anything that has mass and occupies space.

A. Weight

B. Element

C. Matter

D. Atom

ii. A measure of the quantity of matter in an object is called..........................

A. Mass

B. Weight

C. Space

D. Gram

iii. Matter exists in....................... different physical states

A. One

B. Three

C. Two

D. Four.

iv. A gas can be ..................... to form a liquid

A. Condensed

B. Evaporated

C. Solidified D. Neutralized.

v. Matter is ___________ in form

A. Continuous

B. Particulate

C. Solidified D. Neutralized.

2. Match the correct item from List B with the corresponding one in list A.

List A

List B

i. Solid

A. Has no definite shape

ii. Liquid

B. Take the shape of the container

iii. Gas

C. Retains a fixed shape

iv. Melting point

D. The temperature at which a liquid boils to form a gas

v. Boiling point

E. The temperature at which a solid melts to form a liquid

vi. Sublimation

F. A change from solid directly to a gas

vii. Brownian motion

G. Movement of particles from high concentration to low

viii. Diffusion

concentration

H. Irregular motion of tiny particles suspended in a fluid

I. The point of melting substances.

J. The point of liquefaction.

3. Write TRUE or FALSE in the space provided against each question.

i. The kinetic molecular behavior is the way particles behave in solids, liquids

and gas

ii. Physical changes are irreversible, they only affect chemical properties of

substances

iii. Chemical changes are reversible, they affect physical properties of a

substance......................

iv. A physical change result in the production of new substances while a

chemical change does not.............................

v. Matter is made of particles that are able to display some form of

movement...................

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4. Copy and fill the following table. One of the item has been done for you.

PROPERTY

GAS LIQUID

SOLID

Shape

Movement of particles

Compressibility

Not easily compressible

Spaces between particles

Flow

Volume

5. Outline the differences between physical and chemical changes.

6. Copy the following diagram

SOLID

e

b

a

f

c

GAS

LIQUID

d

Name the process a-f above.

7. Categorize the following changes as either chemical or physical.

a. Ammonium carbonate is dissolved in water gradually with stirring. The

solution becomes very cold.

b. Exactly 6g of copper filings are reacted, the mass is found to be 5.55g.

c. A small piece of sodium was dropped into water. It dated about on the water

surface and finally it burst into flames.

d. Copper(ii) sulphate is heated strongly; it changes from blue to white. On

addition of water to the white substance, it changes back to blue.

e. When ethanol is mixed with water, a homogenous solution is formed. When

the solution is distilled, the two liquids are obtained.

f. If water is kept in a deep freezer it solidifies to ice. If the ice is kept in the

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sun, it liquefies to water.

g. Water is added drop wise to dry calcium oxide on a watch glass. Heat

is developed, steam is formed, the oxide cracks and puffs up and finally

crumble to a powder about three times as bulky.

8. Find out about how a refrigerator works. It contains a liquid called a refrigerant

that has a low boiling point. Explain why a refrigerator is cold inside but hot at

the back outside.

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Chapter 6

Elements, Compounds

and Mixtures

6.1 ELEMENTS AND SYMBOLS

An element is a pure substance which cannot be split into simpler substances by simple chemical means. An element is a substance because it has the same

composition throughout. When two or more elements combine they form a

compound.

There are about 105 known chemical elements and some of them are extremely

rare. Some elements do not exist in the earth at all and they are man made in very

small quantities. Examples of elements are; sulphur, iron, copper, aluminium, tin,

lead and magnesium.

NAMES AND SYMBOLS OF ELEMENTS.

Each element has a short hand symbol to represent it. These symbols are used

internationally. In 1808, Dalton published a list of elements in which the symbols

were very difficult to write. In 1811 Swedish chemist Johns Jacob Berzelius

published a system of chemical symbols based on letters and this system is still

used today.

Hydrogen

Strontian

Sulphur

Lead

Azore

Barytes

Magnesia

Silver

Carbon

Iron

Soda

Gold

Oxygen

Zinc

Lime

Platina

Phosphorus

Copper

Potash

Mercury

Figure 6.1; Dalton’s elements symbols.

I l l u s t r a t i v e C h e m i s t r y f o r S e c o n d a r y S c h o o l s

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The rules by which the symbols were assigned can be explained as follows:

1. An element can be represented by a symbol that is derived from the first

letter of its English name. Examples; Iodine(I), Fluorine (F), Hydrogen

(H), Nitrogen (N), Phosphorus (F), and Sulphur (S).

2. Names of different elements may have the same first letter, example copper

and calcium. It is thus necessary to differentiate the elements. Therefore,

another letter, usually the second from the name is used with the first one.

The first letter will be capital while the second will be a small letter.

Examples of elements whose symbols have two letters are; chlorine (Cl),

cobalt (Co), magnesium (Mg), manganese (Mn), and aluminum (Al).

3. In some cases, the symbols can be derived from Latin names instead of the

common English names. Examples of elements whose symbols are derived

from their Latin names are: sodium is natrium in Latin and it’s symbol

is Na, potassium is kalium in Latin its symbol is K, copper is cuprum in

Latin and its symbol is Cu, iron is ferrum in Latin and its symbol is Fe,

mercury is hydrargyrum in Latin and its symbol is Hg, silver is argentum

in Latin and its symbol is Ag, gold is aurum in Latin and its symbol is Au,

tin is stannum in Latin and its symbol is Sn and lead is plumbum in Latin

its symbol is Pb.

Consider the table below which shows symbols of first twenty elements.

Table 6.1 the symbols of first twenty elements.

Element

Symbol

Hydrogen

H

Helium

He

Lithium

Li

Beryllium

Be

Boron

B

Carbon

C

Nitrogen

N

Oxygen

O

Fluorine

F

Neon

Ne

Sodium

Na

Magnesium

Mg

Aluminium

Al

Silicon

Si

Phosphorus

P

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Element

Symbol

Sulphur

S

Chlorine

Cl

Argon

Ar

Potassium

K

Calcium

Ca

IMPORTANCE OF CHEMICAL SYMBOLS

The following are importance of chemical symbols in the field of chemistry.

1. It is easy to understand elements quickly instead of having to memorize full

names.

2. It is now possible to write equations in short form instead of having to write

each element with its full name.

3. The quantity of the element being referred to can be known readily. Example,

the symbol Ag, could refer to one atom or more mole of silver.

Activity 6.1; to test the electrical conductivity of elements.

Procedures;

1. Connect four dry cells and a 2.5v bulb in a series circuit. The dry cells

will produce about 6.0 volts.

2. Connect the terminals as shown in figure 6.2

3. Put the element to be tested in a beaker. The following substances will

be needed: Sulphur powder, iron filings, copper, aluminium, tin, lead,

magnesium.

4. Hold the two terminals of the circuit so that they touch the element to be

tested.

5. Observe whether the bulb lights or not. Compare the intensity of the light

as different elements are tested.

6. Which of the following elements conduct electricity?

Complete the fol owing table with yes or no.

Element

Electrical conductivity

Sulphur

no

Iron

Copper

Aluminium

Lead

Tin

Magnesium

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Battery

Bulb

Carbon rod terminals

Element being tested

Figure 6.2; electrical circuit for testing elements

Discussion questions

1. Between metals and non-metals which are poor electric conductors?

2. Can you explain the reason for the answer you have given above?

MOLECULES

A molecule is an electrically neutral group of two or given atoms held together by chemical bonds. Atoms of oxygen always exist in pairs. Sulphur exists as rings

of eight atoms and phosphorus as tetrahedral of four atoms,

Consider the below diagrams for molecules;

Oxygen

Sulphur

Phosphorus

Figure 6.3; oxygen, sulphur and phosphorus molecules.

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Table 6.2 examples of molecules

Elements

Atomic symbol

Molecular symbol

Oxygen

O

O2

Sulphur

S

S8

Phosphorus

P

P4

Nitrogen

N

N2

Fluorine

F

F2

Chlorine

Cl

Cl2

Bromine

Br

Br2

Iodine

I

I2

Hydrogen

H

H2

6.2 COMPOUNDS AND MIXTURES

A compound is a pure substance that is made up of two or more elements

chemically combined together. The combination is always in a fixed ratio. Examples

of compounds are carbon dioxide, water, table salt and limestone.

A mixture is a combination of two or more substances physically combined

together. Combination in a mixture is in any ratio. Since mixtures are not chemically

combined, they can be separated by physical means. Examples of mixtures are: oil

and water, muddy water, sand and salt.

Mixtures can be homogeneous or heterogeneous. A homogenous mixture has uniform composition, appearance and properties. Example a small amount of sugar

dissolved in a cup of water. After stirring, every section of the solution is identical in composition, appearance and physical properties.

A heterogeneous mixture has different compositions, appearance and properties

at various points in the mixture. Example in a mixture of ice and water, at one point

we find a solid while at another we find a liquid.

Activity 6.2; to prepare a mixture of iron and sulphur.

Procedure;

1. Measure approximately 5g of iron filings and 5g of sulphur powder. Mix

them in a mortar and grind them with a pestle to mix well.

2. Observe the mixture with a hand lens. Can you distinguish the specks of

iron from the specks of sulphur? Does the mixture retain the properties of

iron and those of sulphur?

3. Put some of the mixture on a piece of paper. Move a strong magnet below

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E l e m e n t s , C o m p o u n d s a n d M i x t u r e s

the paper. What do you observe?

4. Put ¼ of spatulaful of the mixture in a test tube. Add dilute hydrochloric

acid. If the gas is evolved ask your teacher how to identify it. Enter your

observations in the table.

Mixture of iron and

Sulphur

Iron filings

sulphur

1 Appearance

2 Effects of a magnet

3 Action of dilute hydrochloric acid

You will notice that the properties of the mixture of iron and sulphur are

similar to the properties of the separate elements.

Activity 6.3; to prepare a compound of sulphur and iron.

Procedures;

1. Prepare about 6g of a mixture of sulphur powder and iron filings by

measuring out 3g of each. Grind them separately using a pestle and

mortar then mix them together thoroughly.

2. Put this mixture in a hard glass test tube.

3. Heat the mixture strongly until you are convinced that the reaction is

over.

4. Cool the product. Empty the product into a mortar and break it into

small pieces.

5. Put one of the pieces of the product on a watch glass and observe it with

a hand lens. Is it different from the original mixture?

6. Put another piece on a piece of paper and bring a magnet near it. Is it

affected by the magnet?

7. Put a third piece of the product into a test tube and add dilute hydrochloric

acid to it. If a gas is evolved, ask your teacher to show you how to

identify it.

8. Enter your observations in a copy of the table below:

Compound of iron

Mixture of iron

and sulphur

and sulphur

1

Appearance

2

Effects of a magnet

3

Action of dilute hydrochloric acid

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Note;

When hydrochloric acid is added to iron sulphide, hydrogen sulphide gas is

produced. This is a very toxic gas and teachers must ensure that this is done on a

very, very small scale preferably in a fume cupboard, a very well ventilated room

or even outside. All remaining iron sulphide should be collected and disposed off.

If it is left in a sink it can react later if acid is put down the sink.

Discussion question

Have you noted any difference in properties between iron sulphide and a mixture

of iron and sulphur? Explain the differences you have noted.

Table 6.3; Differences between compounds and mixtures.

Compounds

Mixtures

Components cannot be seen separately

Components may be seen separately.

The constituent elements can not be

Components can be separated by physical

separated by physical methods

methods.

Compounds have definite compositions Mixtures may vary widely in composition.

by mass of the elements. The proportions The elements are mixed in any proportion.

can not be changed.

A chemical change is involved when new No chemical change takes place when a

substances are formed.

mixture is formed.

The properties of compounds are very

The properties of a mixture are those of the

different from those of the individual

individual elements.

elements

6.3 SOLUTIONS, SUSPENSIONS AND EMULSIONS

SOLUTIONS

A solution is a homogenous mixture of two or more substances in which one of the substances is a solute and the other a solvent. A solvent; Is the component of a solution that dissolves a solute. A solute; is the component in a solution that is dissolved in a solvent. Example; table salt dissolves in water to form a solution.

The salt is a solute and water is a solvent.

When common salt (sodium chloride) is added to water and stirred, it can be

seen to disappear from view. The salt is still there but it is dissolved in water. What can only be seen is the solution. The mixture is said to be homogenous because it is the same throughout.

When sand is mixed with distilled water, neither of the two disappears. The sand

can be seen to spread throughout the water. The mixture is said to be heterogeneous, because it is not uniformly mixed. Both components of the mixture can be seen.

Sand is not dissolved in the water. It is said to be suspended in the water.

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TYPES OF SOLUTIONS

A solution can be unsaturated, saturated or supersaturated.

➢ Unsaturated solution is the solution that can still dissolve more solute, at a given temperature.

➢ Saturated solution; is the solution that can dissolve no more solute at a given temperature.

➢ Supersaturated solution; is the solution that temporarily holds more solute than the saturated solution at a given temperature.

Experiment 6.1

Aim: To distinguish between saturated, unsaturated and super saturated solutions,

Materials;

Common salt, beaker, Bunsen burner, tripod stand, spatula, glass rod, trough.

Procedure;

1. Measure about 100 cm3 of water and pour it in a beaker.

2. Add about a spatulaful of the common salt and stir. Does all the salt

dissolve?

3. Continue to add more salt to the solution, stirring until no more salt can

dissolve.

4. Place the solution on the tripod stand and heat gently, as you continue

stirring.

5. Stop heating when the salt dissolves.

6. Place the beaker in a trough that is half filled with cold water and allow

cooling for 5 minutes. Record your observation.

7. Note that saturation depends on temperature.

Discussion questions.

1. What type of solution do you get when a spoonful of table salt dissolves

in 100 ml of water?

2. What solution do you have, at room temperature when no more salt can

dissolve?

3. What name do you give to the final solution?

On adding a spoonful of salt to a beaker of water, unsaturated solution is

formed. As you dissolve more salt in the unsaturated solution, a point is reached

when the water loses its ability to dissolve any more salt at room temperature. It

becomes a saturated solution. If you heat the water, it increases its solubility and

more salt can be added. If this solution is cooled, the salt particles reappear in the solution. This solution is a supersaturated solution.

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CLASSIFICATION OF SOLUTIONS INTO THREE STATES

OF MATTER

Solutions may be solids, liquids or gases, consider the below table which

expresses examples of types of solutions.

Table 6.4; examples of types of solutions.

SOLUTES

SOLID

LIQUID

GAS

Gas

Ammonium chloride

Water vapour in air

Nitrogen and

sublimes in air to form a

other gases in air

solution

Liquid Glucose in water and table Ethanol in water and

Carbon dioxide

salt in water

various hydrocarbons in water

ts

in each other.

en Solid Steel and other metal alloys Mercury in gold and Hydrogen in

Solv

hexane in paraffin wax metals.

SUSPENSIONS

A suspension is a heterogeneous mixture of liquid and fine particles of a solid in which the solid does not dissolve but is suspended in the liquid. The suspended

particles are slightly visible and will settle over time if left undisturbed. Examples

are muddy water and chalk powder in water.

When suspended particles stay spread out in the water and do not settle, a jelly

like substance is formed and it is called a colloid solution. Liquid droplets or fine solid particles floating in a gas are called Aerosols.

Experiment 6.2;

To prepare suspensions. Materials; mud, chalk powder, water, copper

carbonate, beakers.

Procedure;

1. Put each of the solid substances in a separate beaker.

2. Add enough water to cover and stir until the solid mixes with the water.

3. Leave the mixture to settle. What do you notice?

Discussion questions

1. What place in real life situations can we find mixtures?

2. Can you differentiate mixtures from solutions?

Examples of suspensions are:

Syrup, some paints, body sprays, insecticides and blood in our bodies.

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DIFFERENCES BETWEEN SUSPENSIONS AND

SOLUTIONS

Consider the table below:

Table 6.5; differences between suspensions and solutions.

Solutions

Suspensions

Solutions are homogenous

Suspensions are heterogeneous

Solutions are transparent

Suspensions are opaque or not clear

Particles completely dissolve in solutions

Particles separate on standing

Components in solutions can be separated Components in suspension can be

by evaporation

separated by filtration.

The particles of a solution are at the ion

The particles of a suspension can be seen

or molecular level and can not be seen by by the naked eye

naked eyes

Light can pass through a solution

Light cannot pass through a suspension

EMULSIONS

An emulsion is a mixture of liquids that does not dissolve each other well.

When two liquids are mixed together, they may dissolve in one another, or remain

as separate layers. Liquids which mix together completely are said to be miscible.

Miscible liquids form homogenous mixtures. Ethanol and water mix together

completely to form a single homogenous solution.

Other liquids will never mix even if they are shaken together. They will

separate out into two layers, where the less dense is on top. Water and kerosene are

such liquids. They are said to be immiscible.

Examples of emulsions are ethanol and cooking oil mixture, milk, paint.

Activity 6.4; to prepare a simple emulsion

Procedures;

1. Add warm water to the jar until it is a quarter full.

2. Add a few drops of olive oil or cooking oil.

3. Replace the lid and shake the jar thoroughly. What happens when you leave

the jar to stand?

4. Now add a couple of drops of soap solution and shake again. What happens

now when you leave the jar to stand?

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6.4 METHODS OF SEPARATING MIXTURES

Mixtures contain useful substances mixed with unwanted materials. In order to

obtain these useful substances, chemists often has to separate them from impurities.

The formation of a mixture tends to occur easily and is said to be spontaneous.

Separating mixtures is more difficult to do. One of the jobs done in chemistry is to

separate mixtures into pure substances. Some of the methods of separating mixtures

are; decantation, filtration, evaporation, distillation, sublimation, chromatography

and centrifugation.

DECANTATION

Decantation is the process of separating a heterogeneous mixture of liquid

and solid by pouring out the liquid only and leaving the solid at the bottom of the

container.

The method of decantation works well when the solid component is made up

of large particles.

Consider the diagram below;

Beaker

Sediment left at bottom

Decantation is usually used in blood tests which require the clearer part of the

blood separated from its solid components.

Experiment 6.3

Aim; to separate substances by decantation

Materials; muddy water and beakers

Procedures;

1. Put some muddy water in a beaker and let it stand.

2. Carefully pour out the clear water from the top into another beaker.

Discussion questions

1. What is observed when the muddy water is left to stand for some time?

2. Explain how you can apply decantation in real life situation.

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FILTRATION

Filtration is the method used to separate a heterogeneous mixture of a solid and

a liquid. The solid is separated from the liquid using a porous filter, such as filter

paper. The solid obtained is the residue and the liquid is the filtrate

Activity 6.5;

To separate chalk and salt from their mixture by filtration

1. Grind some chalk in mortar using a pestle.

2. Mix one spatulaful of the chalk powder with two spatulafuls of salt in a 250cm3

beaker.

3. Add water and stir until all the salt has dissolved.

4. Follow the steps as outlined in figure 6.5 to fold a filter paper.

5. Put the filter paper in a funnel, and add some drops of water to fix the filter

paper on the walls of the funnel.

6. Set up the apparatus as in figure 6.6.

7. Pour the contents of the beaker onto the filter paper carefully. Don’t touch the

filter paper. Don’t stir.

Consider the below diagrams;

Step I

Step II

Step III

Step IV

Figure 6.5; folding a filter paper.

Filter funnel

Filter paper

Mixture of sand

Residue

and water

Filtrate

Figure 6.6 filtration.

Discussion questions

1. What is the nature of the filtrate obtained?

2. What is the residue?

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During the filtration, the solid chalk remains on the filter paper and the salt

solution passes through the small holes in the filter paper.

EVAPORATION

Evaporation is the process of separating a solute from a liquid solution. During evaporation the solvent is converted from liquid to gas through heating and the

solute remains as residue. Evaporation is the opposite process of condensation.

Experiment 6.4

Aim; to separate substances by evaporation.

Apparatus; evaporating dish, beaker, tripod stand, Bunsen burner, wire gauze.

Chemicals; salt solution.

Procedures;

3. Pour some salt solution into the evaporating dish and place on a wire gauze

on the tripod stand.

4. Heat the solution until almost all the water evaporates.

5. Let the mixture cool completely.

Discussion questions

1. What is the purpose of heating the solution?

2. What do you observe when the solution cools?

When a solution of common salt in water is heated until the water evaporates,

white solid crystals of the salt are obtained.

Figure 6.7 evaporation.

DISTILLATION

Distillation is the process of separating out a pure liquid from a mixture by

boiling in which a liquid vapourizes and cooled into a distillate. There are two types

of distillation. These are simple distillation and fractional distillation.

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simple Distillation

Simple distillation involves heating the solution to its boiling point and cooling

the vapour to form a liquid thereby the solid solute remains behind the distilling

flask. An example is distilling water from copper (II) sulphate Solution.

Activity 6.6; to obtain pure water from copper (ii) sulphate solution.

1. Put a spatulaful of copper (ii) sulphate crystals in a beaker.

2. Add water and stir until all the crystals have dissolved.

3. Transfer the blue solution into a side arm 250 cm3 distilling flask.

4. Add a few china fragments to the flask; bubbles of vapour form on their

rough surface and ensure smooth boiling.

5. Insert a rubber bung carrying a thermometer as shown in figure 6.8.

6. Connect the side arm of the flask to a Liebig condenser.

7. Set up the apparatus as seen in figure 6.8

8. Allow cold water into the condenser through B and out of it through C.

9. Heat the distilling flask and note the temperature at which the solution

starts to boil.

10. Collect the distillate in a conical flask or a beaker.

Water Out

C

Liebig condenser

Copper sulphate solution

B

China fragments

Water in

Distillate

HEAT

Figure 6.8 simple distillation apparatus.

During the distillation the copper (ii) sulphate remained in the flask. The liquid

collected, called the distillate, is distilled water. The temperature recorded by the

thermometer during the experiment should be 100˚C. The Liebig condenser has

cold water running through an outer tube to cool the steam down and cause it to

condense.

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Fractional Distillation

Fractional distillation is the method of separating a mixture of two or more

liquids that form a homogenous solution by means of fractionating column. The

method works well if the boiling points of the liquids are not too close.

A mixture of ethanol and water is a good example for demonstrating fractional

distillation. These liquids mix completely to form a homogenous solution. The

boiling point of pure ethanol is 78˚C and that of pure water is 100˚C. In simple

distillation only one liquid is involved while in fractional distillation, several liquids may be involved. The liquid with the lower boiling point is the first to be collected

as a distillate.

Experiment 6.5

Aim: To separate two miscible liquids

Apparatus: Liebig condenser, fractionating column, beakers, distillation flasks, retort stands and clamps, thermometer, Bunsen burner, tripod stand, wire

gauze, conical flask.

Chemicals: ethanol, water.

Procedure;

1. Set up the experiment as shown in figure 6.9

2. Heat the ethanol and water mixture. Note the temperature at which the first

liquid is distilled.

3. Record all observation.

Discussion questions

1. Which liquid distills first and at what temperature?

2. What role does the fractionating column play?

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Thermometer

Fractionating

column

Water Out

Clamp

Clamp

Liebig c

Broken porcelain or

ondenser

glass beads

B

Ethanol and water

mixture

Water in

Wire gauze

Distillate

Tripod stand

HEAT

Figure 6.9; fractional distillation.

SUBLIMATION

Sublimation is a direct change from the solid state to the gaseous state usually on heating. The reverse process of direct change from the gaseous state to the solid

state is called reverse sublimation or deposition. The solid that forms after the vapor cools is called a sublimate.

The process of sublimation can be used to separate mixtures where one of the

substances sublimes. Iodine and ammonium chloride are the commonest examples

of such substances.

Activity 6.7; To separate a mixture of sodium chloride and ammonium chloride

by sublimation.

Procedure;

1. Put approximately two spatulafuls each of ammonium chloride and

powdered table salt in a 250 cm3 beaker (hard glass pyrex type). Use a

glass rod to mix the powder.

2. Place the beaker over a gauze mat on a tripod stand.

3. Cover the beaker with a round –bottomed flask containing cold water.

4. Heat the beaker gently. You should avoid forming lots of smoky fumes by

too strong heating.

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You will notice that a white solid forms on the cold flask. This is ammonium

chloride. The sodium chloride does not change when it is heated.

Flask containing cold water

Mixture of ammonium

chloride and sodium chloride

Bunsen burner

Figure 6.10; sublimation of ammonium chloride

FUNNEL SEPARATION

This is a method used for separating immiscible liquids using a separating

funnel. Immiscible liquids are those that do not mix at all when they are kept in

the same container. If they have different densities the most dense liquid settles at

the bottom while the least dense remains at the top of the separating funnel. When

shaken together they make an emulsion but on standing the liquids will separate out

into the original layers. By opening the tap, the liquid at the bottom is let out first.

Activity 6.8 to separate a mixture of kerosene and water by the funnel

separation method.

Materials: separating funnel, kerosene, water, and beaker.

Procedure

1. Pour water into a beaker

2. Add kerosene and stir

3. Pour the mixture into a separating funnel.

4. What do you observe?

5. Open the tap slowly and drain away the liquid at the bottom first, then the

second.

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Stopper

Separating funnel

Retort stand

Clamp

Two distinct layers

Tap

Beaker

Collection

Figure 6.11; separating immiscible liquids.

When kerosene is mixed with water and poured into a separating funnel, it will

move to the top of the funnel. The water is drained out first.

SOLVENT EXTRACTION

Solvent extraction is a method used to separate compounds based on their

relative solubility. This process can be used for extracting essential oil from plant

materials, especially seed, using a liquid that dissolves the oil. The oil is then

distilled and the solvent let to evaporate. This process is also known as liquid-liquid extraction.

Experiment 6.6

Aim: to extract oil from groundnuts

Materials: ethanol, groundnuts, pestle and mortar, fractional distillation equipment, beaker.

Procedure

1. Crush a handful of groundnuts in the mortar.

2. Pour in about 200ml of the ethanol and mix well.

3. Allow the mixture to settle for a while then decant the liquid into a beaker.

4. Use fractional distillation to separate the ethanol from the oil.

Discussion question

Explain the importance of using fractional distillation method in the separation of

ethanol from oil.

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CENTRIFUGATION

Balancing liquid

Centrifugation is a method of

separating a suspension by spinning a

sample around very quickly in a machine

called a centrifuge, which causes

Filtrate

sediment to settle and separate quickly Residue

from the liquid it is mixed. In this process

the sediments go to the bottom with the

liquid above it. The liquid can then be

separated by decantation. A centrifuge is

also used to separate serum from a blood

sample.

Speed control

Figure 6.12 Centrifuge

CHROMATOGRAPHY

Chromatography is a method by Chromatography can also be used to

which different solutes in a solution separate colorless substances in which

can be separated. Frequently they chromatograms can be seen.

are coloured substances in a mixture.

A chromatogram is a pattern

This method enables the different of separated colours formed during

solutes, to be separated and identified. the process of chromatography. Each

This technique was discovered when coloured ring on the filter paper

scientists were extracting coloured dyes corresponds to a different solute in the

from plants. The coloured substances are mixture.

separated by letting them spread across

a piece of filter paper. The substances

Coloured ring

separate out because they move across

the paper at different speeds.

Chromatography works better

when a solvent is used. Solvents like

water, ethanol or ether may be used for

substances which do not dissolve in

water. There are several techniques of

chromatography in which the simplest of

them is called paper chromatography. Figure 6.13 chromatogram

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Activity 6.9, to separate the coloured substances in variegated leaves by using

paper chromatography.

1. Chop small pieces of variegated

the juice at the centre of the filter

leaves into a mortar.

paper. Let it dry completely. When it

2. Add a spatulaful of clean sand

is dry, add another drop of the juice

(which has been washed by water).

and again, let it dry completely.

3. Grind with a pestle and add some

Continue adding a drop at a time at

drops of ethanol as you grind. This

the same spot, and each time letting

process should break down plant

it dry.

cells and release the colour.

8. Use another dropper to add ethanol

4. Squeeze out the green juice as it

onto the coloured spot a drop at a

forms into a small beaker.

time. Allow the first drop of ethanol

5. Continue with the process until you

to spread out before you add the

have at least 20cm3 of the green

next one. Continue adding ethanol

juice from the leaves.

drop wise at the same spot until the

6. Place a filter paper over the top of a

colours have spread out far enough.

petri dish, as shown in figure 6.14.

Let the chromatogram dry.

7. Using a teat pipette put one drop of

Dropper

Filter paper

Petri dish

Figure 6.14; chromatography.

In chromatography the moving analyte. In paper chromatography,

solvent is called a mobile phase and may the stationary phase is a very uniform be a liquid or a gas. The material that absorbent paper. The mobile phase is

absorbs the solvent is called a stationary a suitable liquid solvent or mixture of

phase and is usually a solid, or a liquid solvents.

supported on a solid. The mobile phase

flows through the stationary phase and

carries the components of the mixture

with it. The substance that is separated

during chromatography is called the

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SUMMARY

(a). An element is a pure substance which can not be split into simpler substances

by a simple chemical process.

(b). Chemical symbols are often derived from the Latin or Greek name of the

element and may not have much similarity to the common English name.

(c). A compound is a pure substance that is made up of more than one element in

a chemical combination.

(d). A mixture is a physical combination of two or more substances in any ratio.

(e). A homogenous mixture has uniform composition, appearance and properties.

(f). A heterogeneous mixture has different composition, appearance and

properties at various points in the mixture.

(g). A solution is a homogenous mixture of solute and solvent.

(h). A solution can be unsaturated, saturated or supersaturated.

(i). An emulsion is a mixture of liquids that do not mix well

(j). A suspension is a heterogeneous mixture of liquid and fine particles of a solid

in which the solid does not dissolve but is suspended in the liquid.

REVIEW QUESTIONS

1. Write TRUE or FALSE for the following statements.

i. All known elements have names .............................

ii. An element may be represented by a symbol that is derived from the

first letter of its English name..............................

iii. Hydrogen, oxygen and nitrogen are examples of compounds.......................

iv. A heterogeneous mixture has uniform composition............................

v. A saturated solution is one that can still dissolve more solute at a given

temperature...............................

2. Match items in LIST A with their corresponding statements in LIST B.

List A

LIst B

i. Decantation

A. Used to separate a heterogenous mixture of solid and

ii. Evaporation

a liquid

iii. Distillation

B. separate mixtures using a moving solvent on material

iv. Funnel separation

that absorbs the solvents

v. Sublimation

C. used to separate mixtures by heating a liquid to very

high temperature until it vaporizes

D. separates a solute from a liquid solution

E. Separates immiscible liquids using a separating funnel.

F. Separates a heterogenous mixture of a liquid and a

solid or solids by pouring out of the liquid only and

leaving the solid at the bottom of the container.

G. A solid change state directly to gas, usually on heating.

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3. Choose the most correct answer for the following questions.

i. The substance that is separated during chromatography is called;

A. Analysis

B. Precipitate C. Analyte

D. Stationary phase.

ii. Chromatography can be used in hospitals to:

A. Detect types of drugs in blood B. Analyze blood in crime scenes

C. Testing purity of machines

D. Inject sick people.

iii. One of the following is NOT involved in extracting oil from groundnuts

A. Petroleum

B. Ethanol

C. Groundnuts

D. Fractional distillation apparatus.

iv. The solid obtained during filtration is called_____ and the liquid is called

_____ respectively.

A. Filtrate, residue

B. Residue, analyte

C. Analyte, residue

D. Residue, filtrate

4. Butter is water in oil emulsion. What is meant by water in oil emulsion?

5. If you were stranded in a desert with no water, how could you collect a supply

of drinkable water? You only have a spade, a plastic sheet, a few heavy stones

and a plastic cup.

6. Mention five elements that derive their symbols from their Latin names.

7. Name the process that can be used to separate the following substances:

a. Kerosene and water

b. Iodine and sand

c. A mixture of petrol and diesel

d. A mixture of sulphur and iron filings.

e. Oil from groundnuts.

8. Here are some different mixtures:

Shaving cream, hair lotion

Instant coffee, margarine.

Copy and complete the table below:

Type of mixture

Contains

Example

Solution

One substance dissolved in

another

Foam

Gas mixed with a liquid

Emulsion

Two liquids which do not mix

Gel

Liquid mixed with a solid.

9. Briefly explain how you can tell the difference between a solution and a

suspension just by their appearance.

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Chapter 7

Air, Combustion,

Rusting and Fire fighting

7.1 COMPOSITION OF AIR.

Air is a colourless mixture of gases. It is supposed to be odourless and tasteless.

The atmosphere which covers the earth is made of air. More air is near the surface

of the earth than away from it. Near the surface of the earth, the air is denser than it is higher up. Air is composed of several gases like nitrogen, oxygen, carbondioxide,

water vapor, noble gases and methane.

Air may sometimes contain traces of impurities such as carbon monoxide

(CO), sulphur dioxide (SO ), hydrogen sulphide (H S) and other gases. Consider

2

2

the below table which shows compostion of air by volume.

Table 7.1 Composition of air by volume.

Gas

Symbol

Percentage by volume

Boiling point (˚C)

Nitrogen

N

78.0

-196

2

Oxygen

O

21.0

-183

2

Argon

Ar

0.9

-186

Carbon dioxide

CO

0.03

-57

2

Neon

Ne

0.002

-246

Methane

CH

0.0002

-164

4

Helium

He

0.0005

-269

Krypton

Kr

0.0001

-153

Hydrogen

H

0.00005

-252.9

2

Xenon

Xe

0.00001

-108

TESTS FOR GASES IN AIR.

The test for the presence of different gases in air can be carried out using

several procedures.

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Activity 7.1: To determine the presence and percentage of oxygen in air.

Procedure.

1. Clean two syringes of more than 150 cm3

2. Make very small cuttings of copper foil using a pair of scissors. The copper foil

should be cleaned first with sand paper until it shines.

3. Place the copper cuttings at the centre of a hard glass silica tube.

4. Join the silica tube ends to two syringes as shown in figure 7.1

5. At the beginning, one syringe A should be closed while the other B is filled

with 100 cm3 of air.

6. Heat the silica tube strongly while you push air from one syringe to another.

The air passes over the copper, which uses up the oxygen in air to make copper

oxide. Copper + Oxygen → copper oxide.

7. Allow the tube to cool and read the volume on one syringe while the other is

closed.

8. Repeat the heating and cooling until the volume of air does not change any

more.

9. Summarize the results as illustrated in table 7.2.

Hard glass tube

Copper

B

A

100cm3 air

Gas syringe

Heat

Gas syringe

Fig 7.1 finding percentage of oxygen in air.

Black coating of Copper

B

A

80cm3 air

Heat

Fig 7.2, Reaction of copper with oxygen in air sample.

Table 7.2 Summary of results of activity 7.1

Volume of air in the syringe

Actual reading.

Before heating

100cm3

After first heating and cooling

After second heating and cooling

After third heating and cooling

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Discussion questions.

1. What volume of air was used up by copper?

2. Why was the heating and cooling repeated three times?

3. What is the percentage of oxygen in air?

4. Why was it necessary to cool the apparatus before reading the volume?

5. Why was it necessary to push the air from one syringe to the other as heating

of the copper continued?

Activity 7.2 To show that air contains carbon dioxide and water vapour.

Procedure.

1. Grind some blue crystals of copper (II) Sulphate with a pestle in a mortar

to a very fine powder.

2. Heat the powder strongly while stirring until the entire blue colour is lost.

You should remain with a white powder. Put this powder in test tube A.

3. Put lime water in test tube B.

4. Set the apparatus as seen in figure 7.3

5. Put ice cubes in a beaker around test tube A. Add common salt on the ice

cubes. Add some water in the beaker.

6. Pass air through the system by using a filter pump or an aspirator.

7. Observe any color changes in test tube A and test tube B.

8. Stop the reaction when you think there is no more color change.

To filter pump

air

Test tube B

Test tube A

Lime water

Ice

Water

White CuSO4

Fig. 7.3 showing that air contains carbon dioxide and water vapor.

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Discussion questions.

1. What colour change did you observe on white powder in test tube A?

2. What caused the color change in test tube A?

3. Where did the dew in test tube A come from?

4. Why was ice water mixed with salt?

5. What colour change did you observe in test tube B?

6. What caused the colour change in test tube B?

SEPARATION OF AIR INTO ITS CONSTITUENT GASES.

The task of separating air into its constituent elements is not simple. It cannot

be done in laboratory. It is done in the industry. The chemical industry needs the

gases from the air in their pure form. The separation is carried out by fractional

distillation of liquid air. Air is liquefied by subjecting it to high pressure and low

temperature.

There are three main stages in the process of liquefaction of air.

a. Carbon dioxide and water vapour must be removed first. By cooling the air to

very low temperatures, ice and solid carbon dioxide are removed.

b. The remaining air is then compressed to about 150 times the atmospheric

pressure. As the compressed air gets very hot, it has to be cooled.

c. The compressed cooled air is allowed to expand rapidly. The rapid expansion

cools the air to very low temperature, and liquid-air drops out. At -200˚C,

only helium and Neon remain as gases. These cold gases are used to cool the

compressed air.

The liquid – air is then subjected to fractional distillation. The liquid air is

allowed to warm up. Nitrogen boils off first because it has a lower boiling point,

-196˚C. Oxygen follows by boiling at -183˚C.

AIR AS A MIXTURE.

Air is a mixture and not a compound due to the following reasons.

1. The composition of air does not correspond to any simple chemical formula.

Any compound must have a specific formula.

2. The composition of the air varies slightly in different localities. In the same

locality the composition is different at different times. If it were a compound,

it would always have a definite composition by mass everywhere and all

the time.

3. When the constituents of the air are mixed in their right proportions, there is

no energy (heat or light) change. A compound is always made by a chemical

change involving energy changes.

4. The gases in air can be separated by purely physical means. (i.e. fractional

distillation.) Elements of a compound can never be separated by physical

means.

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USES OF THE GASES IN AIR.

1. Nitrogen.

This gas is present in air in the largest proportion. It is not directly available

to plants, or animals. The gas is not essential for breathing, although its presence

in the air helps to dilute the oxygen. Nitrogen in plants and animals is obtained

by indirect methods e.g. nitrogen fixation. Many of the fertilizers are products of

reaction between nitrogen and other elements.

2. Oxygen.

All plants and animals need oxygen for breathing and respiration. Aquatic

animals use the oxygen dissolved in water. Ocean divers and mountain climbers

must carry their own oxygen in tanks. Oxygen is necessary for burning of fuels and

rusting. Green plants supply oxygen to the air during photosynthesis.

3. Carbon dioxide.

Respiration in plants and animals produce carbon dioxide. All fires from

burning fuel produce the gas. Automobiles and all other fuel engines produce carbon

dioxide in their exhaust. Green plants take carbon dioxide from the atmosphere

during photosynthesis. They use it to make food (sugars and starch). The gas is

used extensively in fire extinguishers.

4. Noble gases.

Argon, helium, krypton, neon and xenon comprise a “family” or group of

elements present in air known as noble gases.

Noble gases are never prepared in the laboratory. They are obtained by

fractional distillation of liquid – air. Helium is obtained from natural gases. It is

used to fill certain photographic flash bulbs.

Argon is used for filling electric light bulbs. It is the cheapest and most abundant

of the noble gases. If air is used instead, the filament may burn out.

Neon is extensively used for filling electric discharge tubes for illumination

and for advertising signs.

Helium is second to hydrogen as the lightest gas known. As the gas for filling

balloons, it has replaced hydrogen.

7.2 COMBUSTION.

Combustion is a chemical reaction that involves the burning of a substance in the presence of oxygen. Combustion is accompanied by liberation of heat and

light. In this sense, combustion has the same meaning as burning. In a complete

combustion reaction, a compound reacts with an oxidizing element such as oxygen

or fluorine, and the products are compounds of each element in the fuel with the

oxidizing element.

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Example.

CH + 2O

CO

+ 2H O

+

energy

4(g)

2(g)

2(g)

2 (g)

Methane

Oxygen

Carbondioxide Water vapour

A simple example can be seen in the combustion of hydrogen and oxyen which

is a commonly used reaction in rocket engines whereby the result is water vapour.

2H + O

2H O

+ heat

2(g)

2(g)

2 (g)

Combustion does not necessarily involve oxygen. Sodium burns in chlorine to

form sodium chloride. In combustion, one of the reactants must be a gas. Ignition

is the process of starting the combustion of fuel. Ignition is very applicable in the

cylinders of an internal – combustion engine.

Experiment 7.1

Aim: To find the products of combustion of a candle in air.

Materials: Mortar and pestle, candle, two test tubes, funnel, delivery tube, rubber, beaker, corks.

Other materials: blue copper (II) Sulphate, ice water, lime water.

Procedure.

(i). Grind blue copper (II) Sulphate with a pestle in a mortar to get a fine powder.

Heat this powder very strongly until it turns completely white.

(ii). Put this powder in the test tube A.

(iii). Connect a delivery tube to a funnel using a rubber connection.

(iv). Keep test tube A in a beaker of ice water

(v). Put lime water in test tube B.

(vi). Set the apparatus as seen in figure 7.4

To filter pump

or aspirator

Rubber connection

Test tube A

Test tube B

Ice

Lime water

air

air

White CuSO4

Figure 7.4. Investigating the products of the combustion of a candle

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When the air has passed through the apparatus;

➢ The white anhydrous copper (II) Sulphate turns blue. This shows that water

was formed when the candle burned.

➢ The lime water turns milky. This shows that carbon dioxide was formed when

candle burned. When substances burn they actually gain in mass.

APPLICATION OF COMBUSTION IN REAL LIFE.

Combustion is applied in many areas such as ;

1. Industries. In engines, in large boilers, incinerators for burning wastes, during welding and smelting, Ovens.

2. Domestic activities. Cooking, heating homes, burning wastes, lighting lamps.

3. Laboratory. Sterilization, during experiments, lighting burners.

7.3 FIRE FIGHTING

Fire fighting is the extinguishing of harmful fires. Fire is the state of combustion

in which ignited material combines with oxygen and gives off light, heat and flame.

Fires can be dangerous. The most dangerous fires are those which are caused by

explosions. Explosions may shatter the whole building and kill many people.

HOW TO PREVENT FIRES.

1. Extinguish all glowing substances before you dispose them.

2. Use the fuel which is appropriate for a particular burner. Petrol should not

be used in a kerosene burner.

3. Cool a burner completely before refilling it with a liquid fuel.

4. Flammable chemicals should be locked out of the working laboratory.

5. After taking a flammable liquid from its container for an experiment,

remove the container from the laboratory before you start the experiment.

6. Containers of flammable liquids should be well stoppered always.

7. Never heat flammable liquids with direct flames.

8. Never pour flammable liquids into the sink.

HOW TO CONTROL FIRE.

A fire will start or continue to burn if three factors are satisfied.

1. There should be a combustible material. Examples of materials are solid,

liquid and gaseous fuels. These substances can catch fire and burn. You can

extinguish the fire by removing the combustible materials.

2. There should be a supply of oxygen. Fuels will only burn if there is enough

oxygen. You can extinguish the fire by blocking air or oxygen supply to

the fire.

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3. The temperature should be at the kindling point of the fuel or above it.

Every fuel has its own kindling point. Below the kindling point, the fuel

will not catch fire. You can extinguish the fire by lowering the temperature

below the kindling point.

Heat

Oxygen

Fuel

Fig.7.5 fire triangle

CLASSIFICATION OF FIRES.

Classification of fires depends on what is actually burning.

(a) Class A fires. The burning materials are solids such as wood, coal, paper, rubber, plastics, clothes, furniture. Water is the best extinguisher of these

fires. If water is not available use any other type of fire extinguisher except

carbon dioxide.

(b) Class B fire. The burning material is a flammable liquid e.g. petrol, kerosene, alcohol, ether, varnishes, fats, etc. If the fire is small a fire blanket or sand may

be used. If the fire is large, use foam or carbon dioxide extinguishers. You

should never use water on class B fires.

(c) Class C fires. The burning material is a liquefied gas. Example acetylene, hydrogen or coal gas. Foam extinguishers and carbon dioxide extinguishers

are the best. It is very important to turn off the source of gas. Spray water on

the gas tank to cool it.

(d) Class D fires. The burning material is a metal. Alkali metals e.g. sodium or potassium may catch fire when they come in contact with water and oxygen.

At high temperatures many metals react with oxygen vigorously. Dry powder

extinguisher and foam extinguisher are the best in class D fires.

(e) Class E fires. These are fires on electrical equipment. You must first find the main switch and put it off. You may use carbon dioxide extinguisher to put off

the fire. You should never use water on electric fires. Water increases electrical

conductivity.

(f) Class F fires. The burning materials are cooking appliances with oils and fats at high temperatures. Use wet chemical extinguishers to handle such fires.

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TYPES OF PORTABLE FIRE EXTINGUISHERS.

A portable fire extinguisher is one that can be easily moved from one place

to another. It is usually hung in an upright position on a wall. It consists of a metal container that contains the extinguishing substance stored at high pressure. Fire

extinguishers are classified according to the type of chemicals they contain.

(a) Liquid carbon dioxide extinguisher.

This is an extinguisher which uses liquid carbon dioxide. The liquid is contained

in a metal container. When the safety pin is removed, carbon dioxide evaporates

as solid “Snow”, (carbon dioxide sublimes). The snow settles on the fire and

smothers it.

(b) Soda – Acid fire extinguisher

This is an extinguisher which uses concentrated sulphuric acid and sodium

hydrogen carbonate solution. The solution is contained in a metal container with

a metal knob at the top of the acid bottle. When it is hot, the knob breaks the

bottle, the acid and/soda solution react to form carbondioxide. The metal case

is turned upside down. The gas mixes with the liquid and the froth (bubbles of

carbon dioxide) comes out of the jet. The froth covers the burning substances

to keep air from reaching it.

(c) Foam type fire extinguisher.

This extinguisher also contains a solution of sodium hydrogen carbonate mixed

with aluminium sulphate and saponin. Saponin is a substance which forms

bubbles very readily. When the three substances mix, they form a foam of

carbodioxide bubbles. The foam also keeps air away from the burning material.

(d) Powder – type fire extinguisher.

This extinguisher uses powdered sodium hydrogen carbonate, and a gas kept at

high pressure. When the gas cartridge is broken using the top cap, the carbon

dioxide under pressure propels the powders. The powder forms a layer over the

burning materials to keep air away.

Fig 7.6 Liquid carbondioxide extinguisher

Sulphuric acid

Jet

Fig 7.7, soda -acid fire extinguisher

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CHEMICAL COMPOSITION OF DIFFERENT AGENTS IN

PORTABLE FIRE EXTINGUISHERS

1. Air pressurized water (APW)

This consists of ordinary tap water pressurized with air. It is suitable for class

A fires and unsuitable for Class B, C and D fires.

2. Dry chemical (DC).

Consist of fine sodium bicarbonate powder pressurized by nitrogen. It is suitable

for class A, B, C and E fires. It is unsuitable for class D fires. When used indoors

it can obscure vision.

3. Carbon dioxide (CO ).

2

It consist of non – flammable carbon dioxide gas under extreme pressure.

Suitable for class B, C and E fires and unsuitable for class A fires.

4. Halon.

Consist of Bromochloro – difluoro – Methane. It is suitable for class A and E

fires and unsuitable for class B and C fires.

5. Foam.

Consist of proteins and fluoro – proteins. Suitable for class A and B fires and

unsuitable for class E fires.

6. Wet chemical.

It consist of potassium acetate. Suitable for class F fires and unsuitable for

class E fires.

7. ABC.

Consists of mono- ammonium phosphate with a nitrogen carrier. Suitable for

class A, B and C fires and unsuitable for electronic equipment.

USING A PORTABLE FIRE EXTINGUISHER.

Use the PASS method when using a portable fire extinguisher.

➢ P = Pull the pin

➢ A = Aim hose at the base of the fire.

➢ S= Squeeze the handle

➢ S= Sweep back and forth with the extinguisher.

Remember these fire fighting tips.

1. Most of the fire extinguishers are emptied in less than a minute

2. Do not attempt to fight a large fire

3. Always call the fire department when a fire breaks out.

4. Always leave yourself a way out – keep your back to an exit.

5. Keep a reasonable distance from the fire as it may suddenly change

direction.

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6. Never use a portable extinguisher on people instead use a fire blanket.

7. Do not test a portable extinguisher to see if it works. It may leak and

afterwards fail to work during an emergency.

8. Do not return a used portable extinguisher to the wall.

Activity. 7.3:

Visit the nearest fire station. (Fire brigade). Ask the officials to tell you:

a. How to establish fire protection at your home

b. What to do in case of a large fire in your school laboratory.

c. How to rescue a friend who is caught in a fire.

d. The best way to fight fires using sand buckets, a water hose and an

asbestos blanket.

7.4 RUSTING.

Rusting is a chemical process that occurs in iron or steel after reaction with water and oxygen from the air. Rust is the reddish – brown coating on the surface

of iron or steel.

Rusting causes a lot of damage economically. The rusting of bridges, corrugated

iron sheets on roofs, containers, rails, etc causes economic problems. If a ship is not protected from rusting, it will not last long in the ocean. All efforts must be made

to stop iron or steel items from rusting. This can be done if we know the conditions

which causes or favour rusting.

Rusting and burning have something in common; they both require oxygen.

During burning, magnesium combines with oxygen of the air, forming magnesium

oxide.

2Mg + O

2MgO

(s)

2(g)

(s)

During rusting, iron combines with oxygen of the air in the presence of water

to form brown hydrated iron (III) oxide, ‘rust’

4Fe + 3O + 2xH O

2Fe O ● xH O

(s)

2(g)

2 (l)

2 3

2 (s)

The only difference between rusting and combustion is in the time required for

the two processes. Heat is generated during rusting just as it is during burning, but it is dissipated without attracting notice because of its much slower rate of production.

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CONDITIONS FOR RUSTING.

There are several conditions that are necessary for rusting to take place.

Experiment. 7.2

Aim: To find the conditions needed for iron to rust.

Procedure.

1. Clean four test tubes and label them W,X,Y and Z. Put them in a test

tube rack.

2. Use sand paper to clean 8 new iron nails. Put two of the nails in each

test tube.

3. Add dry anhydrous calcium chloride in test tube W until the nails are

completely covered. Stopper the test tube using a rubber bung.

4. Cover the nails in test tube X with a wet cotton wool. Cover the mouth

of the test tube with a dry cotton wool.

5. Boil water in a beaker until all air in it is driven off. Pour this water into

test tube Y, until the nails are submerged in the water. Pour oil on top

of the boiled water.

6. Boil some oil until you are sure it contains no water or air.

7. Pour the boiled oil in the test tube Z until the nails are submerged in

the oil.

8. Keep the set up aside for two days, after which you observe what will

have happened to the nails in the four test tubes.

Boiled

oil layer

water

Boiled Oil

Air

Air

Anhydrous

calcium

chloride

W

Wet cotton

X

Y

Z

wool

Fig 7.8 Conditions for Iron rusting.

Discussion questions.

1. Why was it necessary to clean the iron nails with sand paper?

2. What was the use of the anhydrous calcium chloride in test tube W? What

other substance can be used instead of calcium chloride for the same purpose?

3. What was the use of the oil layer in test tube Y?

4. Why was the water in test tube Y boiled first?

5. Why was the oil in test tube Z boiled?

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6. In which test tube was air present but water absent?

7. In which test tube was water present but air absent?

8. In which test tube was both air and water absent ?

9. In which test tube was both air and moisture present?

10. In which test tube did rusting occur most?

11. In which test tubes did rusting not occur?

Within two days, the nails in test tube W will not have formed rust. This is due

to the fact that anhydrous calcium chloride absorbs moisture and also the test tube

was stoppered to prevent entrance of air.

The nails in test tube X will undergo rusting because of the presence of

moisture and air.

Nails in test tube Y will not undergo rusting because there is no dissolved

oxygen in boiled water. Also the presence of oil prevents oxygen from dissolving

in the test tube.

Nails in test tube Z will not undergo rusting because of the absence of both

water and oxygen.

METHODS OF PREVENTING RUSTING.

To stop iron and steel from rusting we must protect them from water (moisture)

and oxygen (air). Some of the methods used to prevent rusting are mentioned below.

1. Painting.

An oil base paint is applied on the surface of ships, frills, cars, machines, roofs

etc. The layer of paint keeps oxygen and water away.

2. Oiling.

Moving parts should not be painted. The layer of paint will be easily scratched

off. Instead they are oiled or greased. The greasing also helps lubrications.

3. Alloying.

It is possible to mix iron and steel with other metals to form alloys. Stainless

steel is an alloy which does not form rust. The stainless steel contains chromium,

nickel and manganese mixed with iron.

4. Galvanization.

This is the process of coating iron or steel with zinc, a metal that does not rust.

Zinc is inexpensive and therefore economical to use.

5. Anodizing.

This is a method used to protect large structures such bridges, ships and

pipelines from rusting. The method is based on a process known as sacrificial

protection. The iron or steel structure to be protected from rusting is joined to a

reactive metal (e.g. magnesium) by a wire. The magnesium is corroded instead

of the iron or steel, hence the term sacrificial protection. The magnesium metal

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becomes the anode and the structure becomes the cathode; hence anodizing. The

anode is sacrificially corroded while the cathode is protected.

6. Use of silica gel.

Silica gel is a compound that absorbs water and thus renders its surroundings

dry. It is a granular form of silicon encased in plastic. Some fragile instruments such as cameras have iron and steel parts. These cannot be protected by painting, oiling,

galvanizing etc. Most often, a small bag of silica gel is put inside the bag carrying

the instrument to absorb moisture in order to avoid rusting.

7. Tin plating.

Tin is a useful metal for the food processing industry since it in non – toxic,

ductile and corrosion resistant. Tin plating is the coating of iron with tin, Tin cans

are made of steel, but the inside of the can is coated with a thin layer of tin. This is why they are suitable for canning of foods.

8. Use of plastic.

Plastic can be used to make parts of some machines or instruments. These

instruments need not to be made by using iron or steel. By using plastic, rusting is

prevented and the resultant product is less expensive.

9. Electroplating.

Electroplating is a process that uses electrical current to reduce dissolved metal

cations so they form a coherent metal coating on an electrode. Electroplating can

also prevent rusting. Articles like car bumpers, taps and kettles are chromium plated.

Often these articles are nickel plated first to give protection against rusting and

finally given a very thin coating of chromium. This give a very shiny surface.

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SUMMARY

(a). Air is a colourless mixture of gases.

(b). Air is composed of several gases like nitrogen, oxygen, carbon dioxide, water

vapour, noble gases and methane.

(c). Separation of air into its constituent gases is carried out by fractional

distillation of liquid air.

(d). Air is a mixture and not a compound

(e). Combustion is a chemical reaction that involves the burning of a substance in

the presence of oxygen.

(f). Combustion can be applied in industries, domestic activities and in

laboratories.

(g). Fire fighting is the extinguishing of harmful fires.

(h). Fire is the state of combustion in which ignited material combines with

oxygen and gives off light, heat and flame.

(i). Classification of fire depends on what is actually burning.

(j). Fire triangle consists of oxygen, heat and fuel.

(k). Rusting is a chemical reaction that occurs in iron or steel after reaction with

water and oxygen from the air.

REVIEW QUESTIONS.

1. Copy and complete the table below. One of the items has been done for you.

Class Burning material

Appropriate extinguisher

A

Ordinary solid combustible material Water, any type of portable extinguisher

such as paper, wood and clothing.

except carbon dioxide.

B

C

D

E

F

2. Choose the best answer to complete each statement or to answer each statement

or to answer each question.

(i) …………………… is a colourless mixture of gases.

A Hydrogen

B Oxygen

C. Air

D. Helium

(ii) Air may sometimes contain traces of …………… such as carbon dioxide and

sulphur dioxide.

A.

Impurities

B.

Purities

C. Pure substances

D. Noble gases.

(iii) The chemical reaction that involves the burning of a substance in the presence

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of oxygen is called……......................

A. Combination

B. Synthesis

C. Combustion

D. Reduction.

(iv) The state of combustion in which ignited material combines with oxygen and

gives off light, heat and flame is called…......……....…..

A. Fire

B. Fire fighting

C. Ignition

D. Reaction.

(v) ……….. is the extinguishing of harmful fires.

A. Fire prevention

B. Fire protection

C. Fire rescuing

D. Fire fighting.

3. Write TRUE for statement that is correct and FALSE for statements that are

not correct.

i. A fuel is a substance that releases energy when burnt.....…………….

ii. Hydrogen, heat and fuel are components of fire triangle……..……….

iii. A portable fire extinguisher is one that can be easily moved from one

place to another…………….

iv. Portable fire extinguishers should not have ratings or coding on

them……………..

v. The procedure for how to use a fire extinguisher is easy to remember

using the acronym “PASS”…………………….

vi. Rusting causes a lot of damage economically………………..

vii. When using a fire extinguisher, do not return a used portable extinguisher

to the wall…………………

4. Match the items in list A with the responses in List B by writing the letter of

the correct response beside the item number.

List A

List B

i. Oiling

A. The coating of items using a special pigment paint

ii. Painting

B. The coating of iron with tin

iii. Galvanization

C. Coating iron with oil

iv. Tin plating

D. Used in case some parts of machines do not have to

v. Silica gel

be made of steel or iron.

vi. Use of plastic

E. A chemical containing silicon found in sand and

vii. Electroplating

rocks.

viii. Anodizing

F. Coating iron or steel with zinc.

ix. Alloying

G. Sacrificial protection.

H. Nickel plating and finally coating with a thin layer

of chromium

I. Mixing iron with rust

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5. Give the difference between rust and rusting.

6. What do you understand by ignition?

7. Copy and complete the table below. One of the items have been done for you.

Gas

Symbol

Percentage by Volume

Nitrogen

Oxygen

Neon

Argon

Carbondioxide

Methane

Helium

Krypton

Hydrogen

Xenon

Xe

0.00001%

8. With the aid of a chemical equation explain the process of rusting.

9. Explain the similarities and differences between burning and rusting.

10. Mention three components of a fire triangle.

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Chapter 8

Oxygen

8.1 CONCEPT OF OXYGEN.

Oxygen is a gas that forms about 21% by volume of the air. Oxygen derives its

name from two Greek words; “oxys” which means acid and “gonos” which means

producer. At the time of naming, it was mistakenly thought that all acids required oxygen in their composition.

Oxygen combines with non – metals to form non – metallic oxides and with

metals to form metallic oxides.

PREPARATION OF OXYGEN.

Oxygen can be prepared in a laboratory or at industrial scale. Some of the

methods used to prepare oxygen are;

1. Fractional distillation of air.

2. Heating potassium chlorate

3. The action of manganese (IV) oxide on hydrogen peroxide.

4. Heating mercuric oxide.

5. Heating potassium permanganate.

8.2 LABORATORY PREPARATION OF OXYGEN.

The common methods of preparation of oxygen in the laboratory include the

following.

1. Decomposition of hydrogen peroxide.

2. Heating of potassium chlorate in the presence of a catalyst

3. Heating compounds rich in oxygen.

The most commonly used methods for laboratory preparation of oxygen is

by using hydrogen peroxide and potassium chlorate. In each case manganese (IV)

oxide is used as a catalyst. With potassium chlorate, heat is needed to activate the

decomposition. No heat is required when hydrogen peroxide is used to prepare

oxygen.

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Experiment 8.1

Aim: laboratory preparation of oxygen by the action of manganese (IV) oxide on hydrogen peroxide.

Apparatus: beehive shelf, flat – bottomed flask, delivery tube, trough, thistle funnel, gas jar, two holed rubber bung.

Chemicals: Hydrogen peroxide solution, manganese (IV) oxide and water.

Procedure

1. Arrange the apparatus as shown in figure 8.1. Make sure that the beehive shelf

is completely immersed in water.

2. Put a small quantity of manganese (IV) oxide into a flat bottomed flask.

3. Fill the gas jar with water and invert it over a beehive shelf in a trough of

water. Make sure the water in the gas jar has no bubbles.

4. Slowly run in hydrogen peroxide solution from the thistle funnel. The reaction

is normally vigorous.

5. Allow the first bubbles of the gas to escape then collect several gas jars of

oxygen and keep them well covered.

6. Open one of the jars and smell some of the gas.

Discussion questions

1. Why is oxygen collected over water?

2. Why are the first few bubbles of gas not collected?

Hydrogen peroxide

Oxygen

Tap

Delivery tube

Thistle Funnel

Gas jar

Cork

Flat bottomed flask

Trough

Hydrogen peroxide

water

Beehive shelf

Manganese (IV) oxide

Fig 8.1 Laboratory preparation of oxygen

100

O x y g e n

Hydrogen peroxide decomposes slowly at room temperature but the reaction

is speeded up by a catalyst known as manganese (IV) oxide.

Word equation

MnO

Hydrogen peroxide

2

water + Oxygen

Chemical Equation

MnO

2H O

2

2H O + O

2 2(l)

2 (l)

2(g)

A catalyst alters the rate of a chemical reaction without itself undergoing any

chemical change. It is possible to collect oxygen by downward displacement of

water because it is only slightly soluble in water.

Experiment 8.2

Aim: Laboratory preparation of oxygen from potassium chlorate.

Apparatus;

Trough, beehive shelf, gas jar, delivery tube, retort stand, pyrex test tube and

Bunsen burner.

Chemicals;

Water, potassium chlorate and manganese (IV) oxide.

Procedure

1. Crush some potassium chlorate crystals in a mortar and grind them by pestle

to get fine powders.

2. Place the mixture of potassium chlorate and manganese (IV) oxide in a pyrex

test tube.

3. Fit the apparatus as shown in figure 8.2 and make sure that the beehive shelf is

completely immersed in water.

4. Heat the mixture strongly, moving the flame to and fro.

5. A gas is readily given off and can be collected over water. Potassium chloride

is left in the pyrex test tube.

Manganese (IV) oxide (MnO ) is used as a catalyst as in previous preparation.

2

The catalyst makes potassium chlorate break down quickly at lower temperature

and thus speed up the rate of a chemical reaction.

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Clamp

Pyrex test tube

Oxygen

Delivery tube

Mixture of potassium

Gas jar

Chlorate and

manganese (IV) oxide

Cork

Trough

Retort stand

Heat

Water

Beehive shelf

Fig.8.2 Laboratory preparation of oxygen from potassium chlorate

Word equation

MnO

Potassium chlorate

2

Potassium chloride + Oxygen

Chemical Equation

MnO

2KClO

2

2KCl + 3O

3(s)

Heat

(s)

2(g)

TEST FOR OXYGEN.

Oxygen gas rekindles a wooden glowing splint. The glowing splint bursts into

flame and burns vigorously when lowered in a gas jar with oxygen. Oxygen itself

does not burn.

8.3 PROPERTIES OF OXYGEN.

Physical properties.

1. Oxygen gas is colourless, tasteless and odourless.

2. Oxygen is slightly soluble in water.

3. Oxygen is about 1.5 times denser than air.

4. The boiling point of oxygen is -183˚C.

5. The freezing point of oxygen is -218˚C.

Chemical properties.

1. Oxygen supports combustion. This is why oxygen relights a glowing splint.

2. Oxygen is a very strong oxidizing agent.

CH

+ 2O

CO + 2H O

4(g)

2(g)

2(g)

2 (l)

Methane Oxygen Carbondioxide water

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O x y g e n

2H S

+ O

2H O + 2S

2 (g) 2(g)

2 (l)

(s)

Hydrogen Oxygen water Sulphur

Sulphide

3. Oxygen reacts with metals to form basic oxides.

2Mg

+ O

2MgO

(s)

2(g)

(s)

Magnesium Oxygen Magnesium oxide

2Ca

+

O

2CaO

(s)

2(g)

(s)

calcium

Oxygen

Calcium oxide

4. Oxygen reacts with non - metals to form non metal oxides (acidic oxides)

C

+ O

CO

(s) 2(g)

2(g)

Carbon Oxygen Carbon dioxide

S

+ O

SO

(s) 2(g)

2(g)

Sulphur Oxygen Sulphur dioxide

Experiment 8.3 To investigate properties of oxygen.

1. Label the gas jars of oxygen you have collected as H,I, J,K and L.

2. Examine the gas in jar for colour and smell. Waft the gas towards your nose

and inhale.

3. Put a small candle on a combustion spoon. Light the candle and lower it into

jar I. Observe the candle flame. See figure 8.3

4. Light a wooden splint. Let it burn for some time and then blow out the flame.

It should keep glowing.

5. Lower the glowing splint into jar J (figure 8.3 (b) ) does the glowing splint

relight?

6. Clean a piece of magnesium ribbon using sand paper. Fix the ribbon on a

combustion spoon as shown in figure 8.3 (c)

7. Pour a little distilled water into Jar K.

8. Light the free end of the magnesium ribbon and put it in jar K immediately.

When the reaction is over, shake the gas jar well. Put red litmus paper in the

jar and leave it there for five minutes. Record any colour change.

9. Put a little Sulphur into a combustion spoon.

10. Pour a little distilled water into jar L.

11. Set the Sulphur alight in the combustion spoon and put in jar L immediately.

When the reaction is over, shake the gas jar well. Put blue litmus paper in the

jar and leave it there for five minutes. Record any colour change.

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Glowing splint

gas jar

Oxygen

Gas jar

Candle

Oxygen

Cover

(a)

(b)

Oxygen

Combustion spoon

Magnesium ribbon

water

Red litmus

(c)

Fig. 8.3 Investigating some of properties of oxygen

Discussion questions

1. What is the colour of oxygen gas?

2. Does oxygen have any smell?

3. Was the candle flame brighter in the oxygen jar?

4. What happened to the glowing splint in the oxygen Jar J.

5. What was the color change to the litmus paper in Jar K? What does this

show about the product formed when magnesium burns in oxygen?

6. What was the colour change to the litmus paper in Jar L? What does this

show about the product formed when sulphur burns in oxygen?

7. What is a word equation for the reaction of oxygen with magnesium and

the reaction of oxygen with sulphur respectively?

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O x y g e n

8.4 INDUSTRIAL MANUFACTURE OF OXYGEN.

The most commonly used method is the fractional distillation of liquefied air.

FRACTIONAL DISTILLATION.

Liequefaction of air; Air is liquefied by, first of all, being filtered to remove dust, and then cooled to -200˚C. At this temperature it is a liquid. As the air liquefies;

➢ The water vapour condenses and is removed using special filters.

➢ The carbon dioxide freezes at – 79˚C and is removed.

➢ The oxygen liquefies at - 183˚C

➢ The nitrogen liquefies at - 196˚C

Distillation:

Liquid nitrogen and oxygen mixture is then separated into the two components

by fractional distillation. The liquid mixture is passed into a fractionating column

from the bottom. Since the column is warmer at the bottom than at the top, the

liquid nitrogen boils at the bottom of the column. The Gaseous nitrogen rises to

the top where it is piped off and stored. The liquid oxygen collects at the bottom

of the column.

8.5 USES OF OXYGEN.

Life on earth depends very much on the presence of oxygen. Animals and

plants, both on land and in water need oxygen gas. Some important industrial

processes need oxygen.

1. Oxygen in breathing and respiration. Oxygen is the only respiratory gas used by both plants and animals. In respiration, oxygen is used to burn food to

produce energy, carbon dioxide and water vapour.

Oxygen is used in hospitals to people who are sick. Some diseases damage the

lungs and stop them from taking enough oxygen from the air. Oxygen is given to

people under anaesthetics and having operations. People who have almost drowned

or have breathed poisonous gases are often given oxygen. Premature babies are

kept in oxygen tents. Mountain climbers and airmen who fly very high use oxygen

in tanks. Ocean divers use oxygen in tanks. Ocean divers use stored pure oxygen.

2. Oxygen in welding. Much oxygen is used in oxy-acetylene and oxy –

hydrogen blow pipes. The oxy – acetylene flame is intensely hot (over 2000˚C)

3. Oxygen in purification of iron. Oxygen is blown into molten iron to remove impurities e.g. carbon or phosphorus. Such impurities are removed as gases.

4. Oxygen in space rockets. In outer space there is no oxygen. Rockets must carry their own supplies of liquid oxygen. Liquid oxygen is used to burn fuels in

rockets. An oxygen hydrogen liquid mixture is the most known powerful propellant.

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5. Oxygen in burning. Oxygen is essential for burning and enables vital combustion processes to take place. Without oxygen, we could not have fires or

we could not be able to burn fuel in order to drive machines and cars.

Burning is a similar process to respiration. Their main difference is that burning

take place faster while respiration is a slow process.

Respiration does not produce light energy. Also, burning can occur in any

material while respiration occurs only in living cells organisms.

SUMMARY

(a). Oxygen is a chemical element with symbol O and atomic number 8 and it

makes up about 21% by volume of air.

(b). The melting point of oxygen is -218˚C and boiling point is – 183˚C

(c). At standard temperature and pressure, oxygen is a colorless, odorless gas with

the molecular formula O .2

(d). The common allotrope of elemental oxygen on earth is called dioxygen, O .2

(e). Trioxygen, O is usually known as Ozone and is a very reactive allotrope of

3

oxygen that is damaging to lung tissue.

(f). Naturally occurring Oxygen is composed of three stable Isotopes, 16O 17O and

18O with 16O being the most abundant (99.762% natural abundance).

(g). Industrial manufacture of Oxygen is mainly by fractional distillation of liquid

air.

(h). Oxygen is prepared in the laboratory by mainly two methods; decomposition

of hydrogen peroxide with Manganese (IV) oxide as a catalyst and heating

potassium chlorate with manganese (IV) oxide as a catalyst.

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O x y g e n

REVIEW QUESTIONS.

1. Choose the most correct answer for the following questions.

i. …………………….. is the most abundant element on the earth’s outer crust.

A. Hydrogen

B. Oxygen

C. Nitrogen

D. Carbon dioxide.

ii. ……………… alters the rate of a chemical reaction without itself undergoing

any chemical change.

A. Matter

B. Element

C. A catalyst D. Compound.

iii. The chemical test for oxygen is…………………

A. Oxygen does not support combustion

B. Oxygen burns in air with a pop sound

C. Oxygen reacts with hydrogen to form water.

D. Oxygen gas rekindles a glowing splint.

iv. …………… is the breaking down of a chemical compound into elements or

smaller compounds.

A. Composition B. Decomposition

C. Dissolution D. Degradation.

v. ………….. burn in oxygen to produce basic oxides while …………. burns

in oxygen to produce acidic oxides.

A. Metals, non metals.

B. Non metals, metals.

C. Metals, metals

D. Non metals, non metals.

2. Match the items in list A with their correct statements in List B.

List A

List B

i. Physical propertis of

A. 16O,17O and 18O

oxygen.

B. 16O and 15O

ii. chemical properties of C. -183˚C

oxygen.

D.180˚C

iii. industrial manufacture E. Oxygen support combustion and it is a very

of oxygen.

strong oxiding agent

iv. the boiling point of

F. Oxygen gas is colourless, tasteless and

oxygen.

odourless

v. Isotopes of oxygen.

G. Fractional distillation of liquid air

H. 100˚C

3. Write TRUE for a correct statement and FALSE for an incorrect statement.

vi. Oxygen is an essential element for chemical processes like combustion,

rusting and respiration………….

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vii. It is not possible to collect oxygen by downward displacement of water

because it is only slightly soluble in water…………..

viii.

Basic oxides react with water to form basic solutions………………..

ix. Acidic oxides react with water to form acidic solutions……………

x. Industrial manufacture of oxygen is mainly by heating compounds rich

in oxygen……………..

4. Hydrogen peroxide can be a dangerous chemical. A bottle of hydrogen peroxide

was left standing on a sunny window sill. The bottle had a screw top. Suddenly

one day the bottle exploded.

(a) Why did the bottle explode?

(b) How would you suggest the hydrogen peroxide is stored safely?

5. Describe briefly four uses of oxygen.

6. The non – metal phosphorus burns in oxygen to form phosphorus (V) oxide.

(a) Write a word equation for the reaction

(b) What colour could you expect litmus paper to turn when phosphorus (V)

oxide is tested? Explain your answer.

7. With the aid of well labeled diagram and a balanced chemical equation, outline

the laboratory preparation of oxygen from hydrogen peroxide.

8. How does oxygen react with sulphur, charcoal, magnesium, calcium and sodium?

Name specifically the products obtained in each reaction.

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Chapter 9

Hydrogen

9.1 THE CONCEPT OF HYDROGEN

Hydrogen is the lightest and most abundant element in the universe and make up about 90% of the universe by weight.

Hydrogen is derived from two Greek words “hydro” which means water and

“genes” which means generator. Hydrogen was discovered by Henry Cavendish in 1766 in London. Cavendish described it as “inflammable air from metals” because

he collected it over mercury.

PREPARATION OF HYDROGEN

Hydrogen gas can be prepared industrially or in a laboratory. Among others,

methods used to prepare hydrogen gas are:

1. Electrolysis of water.

2. Thermal decomposition of ammonia

3. The reaction of water with certain metals

4. The reaction of water with hot carbon.

5. The reaction of dilute acids with moderately reactive metals

9.2 LABORATORY PREPARATION OF HYDROGEN

The common method for preparing hydrogen in the laboratory is by the action

of moderately reactive metal on a dilute acid. The common moderately reactive

metals are zinc, iron and magnesium. The common choice is zinc metal and dilute

hydrochloric acid or dilute sulphuric acid. Zinc is preferred as it reacts smoothly

with either hydrochloric acid or sulphuric acid.

Experiment 9.1

Aim; laboratory preparation of hydrogen.

Apparatus; Thistle funnel, gas jar, trough, beehive shelf, two-holed rubber bang, flat-bottomed flask.

Chemicals; Dilute hydrochloric acid, Zinc granules, water.

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Procedure;

1. Arrange the apparatus as shown in figure 9.1

2. Put some pieces of zinc granules into a flat-bottomed flask.

3. Fill a gas jar with water and invert it over the beehive stand in the trough.

4. Add dilute hydrochloric acid through the thistle funnel until the zinc is covered

with the acid and the end of the thistle funnel is immersed in the acid.

5. Collect the gas over water, ensuring you only remove the gas jar when it is full.

6. Collect several jars of the hydrogen gas.

Dilute hydrochloric acid

Hydrogen gas

Tap

Delivery tube

Thistle Funnel

Gas jar

Cork

Flat bottomed flask

Trough

Water

dilute hydrochloric acid

Beehive shelf

Zinc granules

Fig 9.1 laboratory preparation of hydrogen gas

Discussion questions

1. Why were the few bubbles of gas not collected?

2. What is the colour of the gas?

3. Why is it possible to collect the gas using this method?

4. Suggest one problem with collecting hydrogen this way.

Word equation

Zinc + Hydrochloric acid

Zinc Chloride + Hydrogen gas

Balanced chemical equation

Zn

+ 2HCl

ZnCl + H

(s)

(aq)

2(aq)

2(g)

Hydrogen can be collected over water because it does not dissolve in water. It

can also be collected by upward delivery (downward displacement of air) because

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H y d r o g e n

hydrogen is much less dense than air.

If a sample of dry hydrogen is required the gas has to be passed through a

drying agent before collection. Then the gas could either be collected in a gas

syringe or by upward delivery.

Experiment 9.2; To demonstrate that hydrogen is less dense than air.

1. Prepare hydrogen as shown in figure 9.2

2. Collect some amounts of hydrogen in a test tube. Light a glowing splint and

hold it near the mouth of the test tube. You will hear a pop sound as hydrogen

burns explosively with the oxygen in the air to form water vapour. This is the

test to identify hydrogen.

3. Collect the hydrogen in an inverted test tube labeled A.

4. Allow the gas to escape from test tube A to another test tube B as shown in figure

9.2. The test tube B should be inverted directly over the mouth of test tube A.

B is filled with air.

5. After few minutes test the contents of the test tube A and B using a burning splint.

By the end of experiment you will find that; when testing A with lighted splint

there was no pop sound while on testing B with lighted splint there was a pop

sound.

Discussion question

What do these result show? And why does this happen?

Initially filled with air

B

A

Test tube with hydrogen

Figure 9.2; “Pouring” hydrogen upwards.

Experiment 9.3; demonstration on filling the balloon with hydrogen.

1. Add dilute sulphuric acid to a reagent bottle containing about 20g of zinc

granules.

2. Tie a balloon to the mouth of the bottle using a cotton thread at point A; see

figure 9.3

3. Allow the balloon to be fully inflated with hydrogen.

4. Tie another long cotton thread at point B around the mouth of the balloon.

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5. Unite the thread A, and allow the balloon to rise in the air. Hold it in the air

with the long thread.

You will see the balloon rise upward toward the ceiling. Leave the balloon

inflated for few days. You will see the balloon eventually shrink as the hydrogen

diffuses into the air.

Balloon

B

A

Hydrogen gas

Hydrochloric acid

Zinc granules

Figure 9.3; filling a balloon with hydrogen

9.3 PROPERTIES OF HYDROGEN

Physical properties

1. Hydrogen is a colourless gas

2. Pure hydrogen is odorless

3. Hydrogen is neutral to litmus paper or solution.

4. Hydrogen is the lightest gas

5. Hydrogen is insoluble in water

6. Hydrogen does not support combustion

Chemical properties of hydrogen

1. Hydrogen combines easily with other chemical substances at high temperatures

2. Hydrogen does not usually react with other elements at room temperature.

3. Hydrogen is highly flammable and burns with a blue flame.

4. A mixture of hydrogen and oxygen explodes when lit.

5. Hydrogen reacts with the oxides and chlorides of many metals to produce

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H y d r o g e n

free metals.

6. Hydrogen reacts slowly with oxygen to produce water. The reaction can be

speeded up by catalyst

7. Hydrogen is neither acidic nor basic.

8. Hydrogen reduces metal oxides to metals.

Experiment 9.4; investigation of the effect of hydrogen on copper (ii) oxide.

1. Set up the apparatus as in figure 9.4 in which the test tube has a hole in the end.

2. Turn on the supply of hydrogen and let the gas pass all through the apparatus.

3. Light the jet to burn excess hydrogen.

4. Heat the tube close to the black copper (ii) oxide and observe the change from

black to brown as the reaction occurs.

5. Leave the apparatus to cool before dismantling.

The reaction which takes place can be described as;

Word equation

Copper (II) oxide + Hydrogen

Copper + hydrogen oxide (water)

Balanced chemical equation

CuO

+ H

Cu

+ H O

(s)

2(g)

(s)

2 (l)

The copper (II) oxide is reduced as it loses oxygen. Reduction of copper (ii) oxide has taken place.

copper

Excess hydrogen burning

(ii) oxide

Clamp

Hydrogen

HEAT

Fig9.4 Reduction of copper (ii) oxide with hydrogen

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9.4 INDUSTRIAL PREPARATION OF HYDROGEN

Hydrogen is a very important industrial substance and has to be manufactured

on a large scale. Some of the methods used to produce hydrogen industrially are

explained as follows;

1. Electrolysis of water

Electrolysis of water is a process which decomposes water into oxygen and

hydrogen gas by means of an electric current. The current is passed through the

water. The electrical power sources are connected to two electrodes placed in the

water. Hydrogen will collect at the cathode and oxygen will collect at the anode.

Reduction at cathode;

2H+

+ 2e

H

(aq)

2 (g)

Oxidation at anode;

2H O (l)

O + 4H+ + 4e

2

2 (g)

(aq)

Hydrogen

Oxgyen

Electrodes

Cathode

Anode

Fig 9.5 electrolysis of water using Hoffman voltameter.

2. The reaction of water with hot carbon (Bosch process).

Water gas is produced when steam is passed over white hot coke.

C + H O (g)

CO (g) + H (g)

(s)

2

2

Coke

steam

water gas

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H y d r o g e n

Hydrogen can be obtained from water gas by mixing the water gas with steam

and passing the mixture over catalyst of Iron (III) Oxide at 450˚C.

CO + H O (g)

CO + H

(g)

2

2(g)

2(g)

carbon monoxide

steam

Carbon dioxide Hydrogen

3. Steam reforming

The most common method of producing commercial large amount of hydrogen

is steam reforming of natural gas, also known as steam methane reforming

(SMR). It is a method of producing hydrogen from organic compounds like methane. At high temperatures (700 to 1100˚C), steam reacts with methane to

yield carbon monoxide and hydrogen. The reaction take place in the presence

of a metal - based catalyst.

Word equation;

Steam + methane

carbon monoxide + hydrogen

Chemical equation;

H O + CH

CO + 3H

2 (g)

4(g)

(g)

2(g)

9.5 USES OF HYDROGEN

The following are main uses of hydrogen;

1. Manufacture of ammonia (Haber process)

Ammonia is a compound of nitrogen and hydrogen. Ammonia can be converted

into a large number of nitrogen fertilizers e.g. Ammonium sulphate. Ammonia

can also be used in production of nitric acid and synthetic (man made) fabrics,

such as polyesters and polyamide fabrics.

3H + N

2NH

2(g)

2(g)

3(g)

Hydrogen Nitrogen ammonia

2. Making margarine

Margarine can be made from natural fats and oils such as sunflower oil or

palm oil. These oils are very runny and have to be hardened before they can be

used. The process is known as Hydrogenation or hardening of oils. Hardening involves reacting the heated oil with hydrogen using a nickel catalyst to speed

up the reaction. Margarine is used as an alternative to butter.

3. Oxy-hydrogen flame

The oxy-hydrogen flame uses oxygen and hydrogen to form a hot flame and can

get to temperatures of up to 3000˚C. This flame can be used for welding and

cutting metals.

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4. Manufacture of hydrochloric acid

Hydrogen is used in the manufacture of hydrochloric acid. It reacts with chlorine

to form hydrogen chloride gas which dissolves in water to form hydrochloric acid.

H + Cl

2HCl

2 (g)

2(g)

(g)

Hydrogen Chlorine

Hydrogen chloride

5. Fuel

Hydrogen is used as a rocket fuel. Hydrogen is used to prepare water gas that

can be burnt to propel rockets.

6. Filling the balloons

Hydrogen is the lightest known gas and thus it is used for filling weather

balloons and airships. However, since hydrogen is inflammable, it has been replaced

by helium gas which is inert.

7. Making petrol

Petrol contains about 14 percent of hydrogen and 5 percent of coal. Coal and

hydrogen reacts to form liquid which contains petrol.

8. Manufacture of methanol

Hydrogen can combine directly with carbon monoxide to form methanol.

2H + CO

CH OH

2 (g)

(g)

3

(aq)

Hydrogen Carbonmonoxide Methanol

9. As a reducing agent

Hydrogen is used to reduce metallic oxides to their metals. The oxide of

metals such as iron, copper and lead are reduced to the metals when heated in a

steam of gas.

CuO + H

Cu + H O

(s)

2(g)

(s)

2 (l)

Copper (II) oxide Hydrogen Copper water

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H y d r o g e n

SUMMARY

(a). Hydrogen is a chemical element with symbol H atomic number 1, and

Atomic weight 1.00794.

(b). At standard temperature and pressure, hydrogen is colourless, odourless

and tasteless.

(c). The most common isotope of hydrogen is protium with symbol 1H.

(d). Industrial production of hydrogen is mainly from the steam reforming

of natural gas, and less often from more energy - intensive hydrogen

production methods like the electrolysis of water.

(e). Other isotopes of hydrogen are deuterium (2H) and tritium 3H.

(f). Hydrogen is used to manufacture ammonia because it readily combines

with elements e.g. nitrogen.

(g). Hydrogen is used in production of oxy-hydrogen flame because it is highly

flammable.

(h). Hydrogen is used to manufacture hydrochloric acid due to the fact that it

readily reacts with other chemical substances.

(i). Hydrogen is used to prepare water because it is highly flammable.

(j). Hydrogen was once used to inflate weather balloons because it is lighter

than air.

(k). Hydrogen is used to manufacture margarine because it is a reducing agent.

(l). The boiling point of hydrogen is -252.9˚C and the melting point is

-259.34˚C

(m). The density of hydrogen is 8.988x10-5g/cm3.

(n). Elemental classification of hydrogen is non-metal.

REVIEW QUESTIONS

1. Choose the most correct answer for the following questions;

i. While oxygen is the most abundant element on earth’s outer crust ................

is the most abundant element in the universe.

A. Nitrogen

B. Hydrogen C. Chlorine

D. Carbon dioxide.

ii. The common method for preparing hydrogen in the laboratory is;..............

A. Action of moderately reactive metal on dilute acid.

B. The action of moderately reactive metal on concentrated acid.

C. Water by electrolysis.

D. Reaction of water with hot carbon.

iii. The test for hydrogen;_________________

A. Ignites with a “pop” sound when a lighted splint is inserted into a test tube

with a gas.

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B. Relights a wooden glowing splint.

C. Changes litmus paper red

D. Changes litmus paper blue.

iv. Physical properties of hydrogen gas include;

A. A mixture of hydrogen and oxygen explodes when lit.

B. It is neither acidic nor basic.

C. It is tasteless, colourless and odourless.

D. Combines easily with other chemical substances.

v. Among others, uses of hydrogen include;

A. In the oxy-carbon dioxide flame

B. Manufacture of carbon dioxide

C. As an oxidizing agent

D. Fuel, manufacture of methanol and synthesis of hydrochloric acid

2. Match the items in LIST A with their corresponding statements in LIST B.

List A

List B

x. The addition of oxygen to a substance

A. Protium (1H) and tritium

or the removal of hydrogen from the

(3H)

substance..............

B. Electrolysis of water

xi. The removal of oxygen from a substance

C. Liquid hydrogen and

or the addition of hydrogen to a

gaseous hydrogen

substance..............

D. Hydrogenation

xii.

Is the process of passing hydrogen through E. Hydro oxidation

liquid oil to harden it...............

F. Reduction

xiii.

............. is a process which decomposes G. Oxidation

water into oxygen and hydrogen gas, with

H. Combined reduction and

the aid of an electric current.

oxidation.

xiv.

Isotopes of hydrogen.......................

3. Write TRUE for a correct statement and FALSE for an incorrect statement; i. Hydrogen can be collected over water because it does no dissolve in water........

ii. Hydrogen can also be collected by upward delivery because hydrogen is

much less dense than air..................................

iii. Hydrogen is nowadays used for weather balloons and airships.....................

iv. Hydrogen is widely used as an oxiding agent.................................

v. Ammonia is a compound of nitrogen and hydrogen............................

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H y d r o g e n

4. (a) List any three physical properties and any three chemical properties of

hydrogen

(b) Mention three properties of hydrogen that make it different from oxygen

5. Describe three methods for the industrial manufacture of hydrogen.

6. Hydrogen is a very attractive source of energy but its use as a major source

of energy is not practical. Explain this in terms of its storage, safety and

production.

7. What is a chemical test for hydrogen gas?

8. Why is zinc a better choice than sodium and potassium as a metal for reacting

with acid to give hydrogen?

9. Explain the laboratory preparation of hydrogen gas from zinc and dilute

sulphuric acid.

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Chapter 10

Water

10.1 OCCURRENCE AND NATURE OF WATER.

Water is a compound of Hydrogen and Oxygen. It is a very important

commodity because it is essential for the sustenance of all living things. Water is

also a home for some animals and plants.

Development depends on availability of water. In areas where there is plenty of

water, plants and livestock flourish. In dry desert areas, plants, animals and people

are scarce.

Three quarters of the earth’s surface is covered by oceans and seas. On the

average, oceans and seas water contains 3. 6 percentages by mass of dissolved

solids as follows; common salts 2.8 %, calcium salts 0.1 %, magnesium salts 0.6%,

potassium salts 0.1%. Bromine, Iodine, Copper, Silver and even Gold and Radium

are often found in the sea.

THE WATER CYCLE

Water is continually moving above and below the earth, as water vapour, liquid

water and ice. The water is never lost but it continually being recycled all around

the globe in a system called the water cycle (hydrological cycle).

The water cycle is made up of main five parts.

1. Evaporation.

2. Transpiration

3. Condensation

4. Precipitation

5. Collection.

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Wa t e r

Fig 10.1 the water cycle.

1. Evaporation

Evaporation is the process in which water vapour leaves the streams, rivers,

lakes or oceans and goes into the atmosphere. The sun provides the driving

force by heating up the water bodies and turns the water into vapour.

2. Transpiration

This is the process in which plants lose water through their leaves. This is

another way in which water gets back in the atmosphere. This vapour undergoes

condensation to form clouds.

3. Condensation

This is the process in which water vapour in the air cools and changes back

into liquid, thereby forming clouds.

4. Precipitation

This is the process in which clouds get heavy and water falls in form of rainfall,

hail or snow.

5. Collection

Collection can take place when water falls back to earth as precipitation. The

water may go into oceans, lakes or rivers, or on land. When water falls on land,

It normally soak into the soil and becomes part of the ground water which

plants and animals use.

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RELATIONSHIP BETWEEN WATER CYCLE AND

ENVIRONMENTAL CONSERVATION.

People must observe principles and rules so that the environment can be

conserved. Conservation of Natural resources like plants, water and land refers to

environmental conservation. Land clearing especially near sources of water should

be avoided. Drying of ponds and rivers would in turn affect the normal circulation

of water. Tree planting and land reclamation programs should be encouraged in

order to conserve the natural environment. The use of natural gas as the source of

fuel should be encouraged.

10.2 PROPERTIES OF WATER

The physical properties of water involve aspects such as colour, taste and

smell. They also involve melting, freezing and boiling points. Chemical properties

of water involve the results that occur when water reacts with other substances.

Experiment 10.1

Aim: to measure the melting and boiling points of water.

Apparatus; Bunsen burner, tripod stand, thermometer, retort stand and clamp, two test tubes.

Procedure;

1. Put some ice cubes in one test tube and water in another test tube.

2. Clamp the thermometer in a vertical position

3. Set the apparatus as illustrated in fig 10.2. Adjust the thermometer so that it

dips into the ice cubes, but does not touch the test tube. Record the temperature

of the ice cubes.

4. Heat the water in another test tube. Record the temperature after water forms

steam.

Thermometer

Steam

Water

Ice

HEAT

Fig 10.2 finding the melting point and boiling point of water

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Discussion questions.

1. Why was the thermometer not allowed to touch the test tube?

2. What was the temperature of the ice cubes?

3. What was the temperature of the boiling water?

The thermometer should not touch the test tube in order to avoid altering the

reading on the thermometer.

At atmospheric pressure, pure water boils at 100˚C and freezes at O˚C. Changes

in altitude or pressure alter the boiling point of water.

Experiment 10.2

Aim: To show the presence of water.

Materials; Distilled water, blue and red litmus papers, cobalt chloride paper, watch glasses, anhydrous Copper (II) Sulphate.

Procedure;

1. Pour some amount of distilled water onto a watch glass.

2. Dip a strip of blue litmus paper into the water on the watch glass.

3. Repeat the steps 1 and 2 using red litmus paper and cobalt chloride paper

respectively.

4. Put some amount of anhydrous Copper (II) Sulphate on a watch glass and add

some little distilled water.

Discussion questions.

1. What colour changes were obtained on the blue litmus, red litmus and cobalt

chloride papers?

2. What was the colour change observed on anhydrous Copper (II) Sulphate

when some distilled water was added?

Water is neutral i.e. neither acidic nor basic and therefore it has no effect on

either red or blue litmus papers.

Water reacts with cobalt chloride paper and the paper changes from blue to

pink. The cobalt chloride paper is used to test for the presence of water in substances.

White anhydrous Copper (II) Sulphate turns blue when water is added. The

blue colour indicates the hydrated Copper (II) Sulphate is formed.

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Experiment 10.3

Aim: the synthesis of water.

Materials: U-tube apparatus, filter pump, thistle funnel, test tube.

Procedure;

1. Pass a stream of Hydrogen through a U – tube containing anhydrous calcium

chloride to ensure that no water is present.

2. Allow all the air to be displaced from the apparatus.

3. Ignite the gas coming out of the apparatus. When the sample burns quietly with

blue flame, it indicates that all the air has been displaced and the gas coming

off is pure Hydrogen.

4. Assemble the rest of the apparatus as shown in figure 10.3 below.

5. Ignite the Hydrogen so that the flame is about 3cm3 high.

6. Draw the products of the combustion of Hydrogen through the apparatus by

means of filter pump for about 10 minutes.

7. Test any liquid which collect in the test tube.

Hydrogen

burning at jet

Hydrogen

Anhydrous

To filter pump

calcium chloride

Cold water

Fig. 10.3 Synthesis of water

By the end of experiment, a colourless liquid collects in the test tube which

can be shown to be water by testing it with anhydrous Copper (II) Sulphate. Water

turns anhydrous white Copper (II) Sulphate blue.

When few drops of the liquid formed are added to the blue cobalt chloride

paper the paper will turn pink and this shows that the liquid formed is water.

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Experiment 10.4: to study the reactions of sodium and potassium with cold

water.

1. Pour cold water into a trough until it is half filled. Take sodium out of a bottle

with dry tongs. Use a dry knife and a dry white tile to cut four small pieces

each about the size of a rice grain. Return the rest of the sodium to the bottle

and replace the lid. What colour is the metal when freshly cut, and how does it

change if left in air? Scrape off any white or brown crust that is on the metal.

Never touch sodium or potassium with your fingers.

2. Drop a single piece of sodium into the water in the basin, keeping your face

away in case an explosion occurs. Does the sodium sink or float? Does it move

around in the water? What is its shape and appearance? Does it burn at any

time and if so, what is the colour of the flame? What is the colour of and what

happens to the last bit of sodium?

3. Add two more pieces of sodium to the water, one at a time.

4. Feel the solution between your fingers. How does it feel?

5. Add some universal indicator to the solution in the trough. What colour change

occurs?

6. Repeat the whole test with a fresh sample of cold water and tiny pieces of

potassium, which reacts more vigorously and dangerously than sodium. What

other differences are there between the reactions of sodium and potassium on

water?

Potassium is a bright silvery metal. It turns dull quickly in air as it reacts with

Oxygen. It is therefore kept under paraffin oil. It is so soft that it can be cut easily with a knife. Its density is low so it floats on water.

The reaction between potassium and water is like that between sodium and

water.

Potassium is more reactive and so reacts more quickly and the Hydrogen

formed at once burns with a purple flame.

Sodium is a silvery, shiny metal. It is malleable and soft and can be cut with a

knife. The cut surface at once turns dull because sodium reacts with Oxygen in the

air. The density of sodium is low so it floats on water.

When sodium is added to water, it first melts to a silvery ball and moves

about on the surface of the water; and eventually gives off Hydrogen and gradually

becomes smaller as it reacts.

The solution left feels soapy and greasy when rubbed between the fingers, and

it turns universal indicator purple.

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Experiment 10.5 Collecting the gas formed when sodium reacts with water.

1. Cut three pieces of sodium as described in the previous experiment. Wrap them

in wire gauze so that it forms a cage around them. Leave room for the sodium

to move in the cages, if it cannot move there is more danger of an explosion

during the experiment.

2. Fill a boiling tube with water and invert it in a trough of water.

3. Drop the wire cage into the beaker and move the boiling tube over it to collect

any gas formed.

4. Remove the tube and test the gas with a burning splint. What happens?

Caution;

Gas

Do not attempt this experiment

with potassium

Bubbles

Water

Sodium in wire gauze

Fig.10.4; Collecting the gas formed

when sodium reacts with water

Experiment 10.6 To study the actions of calcium and magnesium on water and

collect any gas formed.

(a) Calcium

1. Fill a beaker with water. Use tongs to drop three small pieces of calcium into

the water. What happens?

2. Pour some of the solution into a test tube and test it with universal indicator.

What change occurs?

3. Filter some of the solution into a test tube. Blow through glass tubing into the

filtrate. What change occurs? What do you think the solution is?

4. Collecting the gas.

a. Fill a boiling tube with water and invert it in the beaker.

b. Add pieces of calcium to the beaker and cover them with filter funnel.

Move the boiling tube over the funnel to collect any gas formed. The stem

of the funnel must be short.

c. Remove the tube and test the gas with a burning splint. What happens?

(b) Magnesium

1. Rub the magnesium ribbon with sand paper until it is bright and shiny, roll the

ribbon into a ball and drop it into water. Leave for about 10 minutes. What do

you observe?

2. Place a boiling tube filled with water over the magnesium and leave it for 2 – 3

days.

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Does any gas collect in the tube? if so, test it with a burning splint.

Hydrogen burning at the jet

Steam

Magnesium ribbon

Safety tube

HEAT

Steam

Boiling water

HEAT

Fig. 10.5 the reaction of magnesium and steam.

Experiment10.7 to study the reaction of magnesium with steam.

1. Rub magnesium ribbon with sandpaper until it is bright and shiny.

2. Set up the apparatus in figure 10.5 and heat the flask containing water until there is a stream of steam going through the apparatus.

3. Then heat the test tube until the magnesium starts to burn.

4. Light the jet of Hydrogen gas leaving the test tube.

Calcium is a shiny, silvery metal. Usually it looks white because calcium

oxide forms on its surface. A piece of calcium sinks in water and reacts quietly to

form bubbles of Hydrogen. The calcium gradually disappears and forms calcium

hydroxide solution. The water becomes cloudy because the calcium hydroxide is

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not very soluble and it forms a suspension.

Magnesium reacts very slowly with cold water but faster with hot water.

Heated magnesium catches fire in steam and burns brightly. The reaction releases

a lot of energy and the Hydrogen formed catches fire. White magnesium oxide is

also formed. Magnesium hydroxide is not formed because the temperature is too

high, but it is formed when the metal reacts with hot or cold water.

Experiment 10.8 to examine the action of steam on heated iron.

1. Place iron powder in the middle part of a combustion tube.

2. Arrange the apparatus as shown in figure 10.6 (but without the gas jar in

position)

3. Boil the water gently. Warm the end of the combustion tube where the steam

enters so that it does not condense. When all the air is out of the apparatus,

place the gas jar in position. What causes the noise in the trough?

4. Heat the iron and keep the steam passing slowly. What happens to the iron?

What collects in the gas jar? The safety tube prevents “sucking back” at the

end of the experiment when the apparatus is cooling. How does it do this? It

also allows water and steam to escape if the apparatus becomes blocked.

5. Test the gas with a burning splint. What happens? What is the gas? Describe

the solid residue.

Steam reacts slowly with heated iron to form Hydrogen and a blue – black

Iron oxide.

Steam

Safety tube

Reduced iron or iron powder

Gas

HEAT

Boiling water

Combustion tube

Cold Water

Tripod stand

Beehive shelf

HEAT

Fig. 10.6 The action of steam on heated iron

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According to the experiments performed previously. The properties of water

can be described as follows;

Physical properties.

1. Water is colourless, odourless, and tasteless.

2. Water is the only substance that occurs naturally in all three states of matter.

3. The freezing point of water is 0˚C and the boiling point is 100˚C

4. Water is the universal solvent because it dissolve more substances than any

other liquid.

5. Water has a high surface tension. It is sticky and elastic and tends to clump

together in drops rather than spread out in a thin film.

6. Water has a high specific heat index. It can absorb a lot of heat before it begins

to get hot.

7. Water expands when it freeze. Ice is therefore less dense than liquid water.

8. Water is miscible with many liquids example ethanol.

Chemical properties.

1. Pure water is neutral that is it is neither acidic nor basic

2. Cold water reacts with some metals to form metal hydroxides and liberate

Hydrogen.

3. Some metals react with steam to give the respective metal oxide and Hydrogen

gas.

TEST FOR WATER

There are two main methods which can be used to test water in laboratory.

i. Water changes cobalt chloride paper from blue to pink.

CoCl + 6H O

CoCl ●6H O

2(s)

2 (l)

2

2 (s)

Cobalt (ii) chloride (pink)

(Blue)

ii. Water changes anhydrous Copper (ii) sulphate from white to blue.

CuSO + 5H O

CuSO ●5H O

4(s)

2 (l)

4

2 (s)

Copper (ii) sulphate

(blue)

(White)

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REACTION OF WATER WITH METALS

(a) Potassium.

Potassium is bright silvery metal. The reaction between potassium and water is

like that between sodium and water. Potassium is more reactive and so reacts more

quickly and the Hydrogen formed at once burns with a purple flame.

Word equation:

Potassium + water

potassium hydroxide + Hydrogen

Chemical equation:

2K + 2H O

2KOH + H

(s)

2 (1)

(aq)

2(g)

Potassium is kept under paraffin oil because it quickly react with Oxygen in

air and turns dull.

(b) Sodium.

Sodium is a silvery, shiny metal. It is malleable and soft and can be cut with a knife.

When added to water sodium;

i. First melt to a silvery ball.

ii. Move about on the surface of the water.

iii. Gives off Hydrogen and gradually become smaller as it reacts.

The solution left feels soapy or greasy when rubbed between the fingers, and

it turns universal indicator purple.

Word equation:

Sodium + water

Sodium hydroxide + Hydrogen

Chemical equation:

2Na + 2H O

2NaOH + H

(s)

2 (1)

(aq)

2(g)

(c) Calcium.

Calcium is a shiny, silvery metal. Usually it looks white because calcium oxide

form on its surface. A piece of calcium sinks in water and reacts quietly to form

bubbles of Hydrogen

Word equation:

Calcium + water

Calcium hydroxide + Hydrogen

Chemical equation:

Ca + 2H O

Ca(OH) + H

(s)

2 (1)

(2)(s)

2(g)

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(d) Magnesium.

Magnesium reacts very slowly with cold water but faster with hot water. Heated

magnesium catches fire in steam and burns brightly. The reaction release a lot

of energy and the Hydrogen formed catches fire. White magnesium oxide is also

formed.

Word equation:

Magnesium + Steam

Magnesium oxide + Hydrogen

Chemical equation:

Mg + H O

MgO + H

(s)

2 (1)

(s)

2(g)

Magnesium hydroxide is not formed as the temperature is too high, but it is

formed when the metal reacts with hot or cold water.

Word equation:

Magnesium + Water

Magnesium hydroxide + Hydrogen

Chemical equation:

Mg + 2H O

Mg(OH) +H

(s)

2 (1)

2(aq)

2(g)

(e) Iron.

Steam reacts slowly with heated iron to form Hydrogen and a blue-black iron oxide.

Word equation:

Iron + Steam

Iron oxide + Hydrogen

Chemical equation:

Fe + H O

FeO +H

(s)

2 (g)

(s)

2(g)

(f) Zinc.

Zinc does not react with hot or cold water. If zinc is heated to redness in a currect

of steam, Hydrogen is formed.

Word equation:

Zinc + Steam

Zinc oxide + Hydrogen

Chemical equation:

Zn + H O

ZnO + H

(s)

2 (g)

(s)

2(g)

Zinc Oxide

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10.3 WATER TREATMENT AND PURIFICATION.

Water treatment; is the process of making water usable for domestic, industrial, medical and other purposes. The aim of the treatment process is to remove existing

contaminants in the water, thus improve it for safe use. Treatment processes may

be physical such as settling, chemical such as disinfection, or biological such as

slow sand filtration.

Water purification; is the removal of contaminants from treated water to

produce drinking water that is pure enough for human consumption. Substances that

are removed include bacteria, algae, and fungi, minerals such as iron and sulphur

and human made chemical pollutants.

DOMESTIC WATER PURIFICATION.

Simple and diverse methods used to treat water to make it safe can be explained

as follows.

1. Filtering.

Filtering will remove solid particles. The size of the holes of a filter is important.

There are two basic types of filters.

Membrane filters use thin sheets with precisely sized pores that prevent objects

larger than the pore size from passing through.

Depth filters use thick porous materials such as carbon or ceramic to trap

particles as water flows through the materials. Activated carbon filters also remove

a range of organic chemicals and heavy metals.

There is a difference between a water filter and a water purifier. Filters do not

filter out viruses, but there are water purifiers that pass the water through both a

filter and an iodine compound that kills any smaller organisms that have passed

through the filter.

2. Boiling.

Boiling is the most efficient way of killing all microorganisms. Water

temperatures above 70˚C kill all pathogens within 30 minutes and above 85˚C

within a few minutes. So in the time it takes for the water to reach the boiling point

(100˚C) from 70˚C, all pathogens will be killed, even at high altitude.The boiled

water is then allowed to cool before being filtered using a clean cloth.

3. Adding water treatment chemicals.

There are two types of chemicals; those using iodine and those using chlorine.

There is a variety of products on the market, so follow the directions on the bottle.

Be advised that many of the tablets have an expiry date and become ineffective

after that point. Remember that chemical purification methods may only be partially

effective, depending on the water temperature.

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Iodine treatment.

Iodine is light sensitive and must always be stored in a dark bottle. It works

best if the water temperature is over 21˚C. Iodine has been shown to be more

effective than chlorine based treatments in inactivating Giardia Cysts.

Chlorine treatment.

Chlorine can be used to treat water by killing microorganisms in the water.

Halazone is an example of a chlorine tablet product. To use, follow the manufacturer’s

instructions.

Experiment 10.9

Aim; to filtrate river water in the laboratory.

1. Using a filter paper collect some water with a bucket from a nearby river, pour

some water into flask and filter this water by means of a filter paper.

2. Using sand filter, instead of a filter paper, carefully carry some large pieces of

gravel to the bottom of the funnel. Arrange a layer of smaller pieces of gravel

on top of large pieces of gravel, followed by a layer of fine sand. Pour a sample

of river water in a beaker.

It can be observed that in both cases the filtrate may appear relatively clear.

The sand filter bed method is also used to purify some public water supplies. These

filter beds are filled with fine layers of gravel and then sand. The water is allowed

to percolate slowly through these layers and both organic and inorganic debris are

retained by them.

Sand

Fine gravel

Coarse gravel

Fig.10.7; sand filter

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URBAN WATER TREATMENT.

River water passes through a series of treatment stage before it is fit for urban

consumption. The stages can be explained as follows;

1. Chlorination and Aeration. After filtering, the water may be aerated.

Water is pumped through fountains and spouts up into the air. Aeration kills many

dangerous aquatic bacteria. Another chemical added in water to kill harmful bacteria

is liquid sodium hypochlorite, which releases chlorine in water. Chlorine is such a

useful disinfectant that is used in swimming pools to kill bacteria.

2. Sedimentation. Dirty water from a river is pumped and stored in large reservoirs where solid particles are allowed to settle to the bottom. The process

takes long time for most of the particles to settle.

3. Chemical treatment. Even after a long time colloidal materials will remain suspended. The deposition of colloidal materials such as clay is affected by the

addition of a chemical precipitant.

Aluminum sulphate, in the form of potash alum, and slaked lime (calcium

hydroxide) are added to precipitate the suspended clay matter. The two chemicals

react to form aluminium hydroxide and calcium sulphate.

The aluminium hydroxide is bulky and sticky, the calcium sulphate is far

denser than water. Both the solid products settle to the bottom.

Aluminium sulphate + calcium hydroxide → aluminium hydroxide + calcium

sulphate.

4. Filtration. Water from the reservoirs is passed through a sand filter. The filter is made up of layers of sand, charcoal and gravel. The charcoal is to remove

colored matter and noxious smells from the water. Filtration beds are expensive

to install and require considerable labour to maintain. The above processes aim at

making water suitable for drinking and hence the water is said to be potable.

Storage

Disinfection/

Ready for use

soft water

chlorination

stored ready

for use

Raw water from

source

Screening

Reservoirs

Primary

Secondary

to remove

water is

filtration

filtration

floating

stored

filtered

substances

high up

through

so it flows

coarse sand.

through

Addition of

gravitation

aluminium

sulphate

Fig.10.8 Process of urban water treatment

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10.4 IMPORTANCE OF WATER TREATMENT

The following are some of the reasons for why water needs to be treated.

1. Untreated water may contain harmful bacteria and other parasites that can

cause diarrhoea, typhoid, cholera and other diseases.

2. Treated water is the best for use in laboratories and medical facilities to ensure

accurate results from experiments and effective treatment.

3. Treated water is suitable to use in factories in order for manufactured products

to be safe for consumption.

4. Treated water is more efficient to use for cleaning in industries and in domestic

settings. Untreated water will usually lead to usage of greater amounts of

certain substances such as soaps for cleaning.

10.5 USES OF WATER.

There are many ways in which water is useful and that is why it is very

important to conserve water. Among them the following are uses of water.

1. Drinking, washing, cleaning and cooking.

2. Irrigation

3. Transportation of people and goods.

4. Generation of hydroelectric power (electricity)

5. Filling swimming pools for recreation activities.

6. Fire fighting.

7. In industries as steam for running machines

8. Cooling engines and parts of machines.

SUMMARY

(a). Water is important on day to day basis for various purposes such as drinking,

cooking, cleaning, washing our bodies and clothes.

(b). Water is also used for transportation and recreation activities.

(c). Water is the most abundant compound on Earth’s surface, covering about

75% (percent) of the planet.

(d). In nature water exists in liquid, solid and gaseous states.

(e). The density of water is 1,000.00 kg/m3

(f). Water is a chemical compound with the chemical formula H O.

2

(g). A water molecule contains one oxygen atom and two hydrogen atoms

connected by covalent bonds.

(h). Pure water has a low electrical conductivity, but this increases with the

dissolution of a small amount of ionic material such as sodium chloride.

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REVIEW QUESTIONS

1. Choose the best answer to complete each statement.

i. A compound of Hydrogen and Oxygen is……………………….

A. Carbon dioxide

B. Water

C. Sulphuric acid

D. Liquid nitrogen

ii. The water cycle is also referred to as

A. Hydrological cycle

B. Hydrogen cycle.

C. Carbon cycle

D. Nitrogen cycle.

iii. The water on earth occurs in ………. main states.

A.

Two

B.

One

C.

Four D.

Three

iv. The boiling point of water is…………….

A. 0˚C

B. 90˚C

C. 100˚C

D. 95˚C

v. Most of the Earth’s water is…………………

A. Pure

B. Not pure

C. More pure D. Clean.

2. Match the following items by writing a letter of their correct statement in the

spaces provided.

List A

List B

i. Evaporation

A. The process in which water vapour cools and changes back

ii. Transpiration

into liquid.

iii. Condensation

B. The process in which plants lose water through their leaves.

iv. Precipitation

C. Take place when water fal s back to earth and go into oceans,

v. Collection

lakes, rivers and on land.

D. The process in which clouds get heavy and water fal s in form

of rainfal , hail or snow.

E. The process in which water vapour leaves the streams, rivers,

lakes and oceans and goes into the atmosphere.

F. The process in which clouds get heavy.

3. Write TRUE for a correct statement and FALSE for an incorrect statement.

i. Water is used for irrigation, in animal dips and for watering flowers……….

ii. In the extraction of certain minerals, water is used as a solvent, or as a means

of carrying away impurities.

iii. Large water bodies, especially rivers and artificial lakes are used to generate

electrical energy…………..

iv. Water bodies such as oceans, lakes and rivers are not used for fishing…………

v. Water treatment is the removal of contaminants from treated water to produce

drinking water that is pure enough for human consumption.

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4. (a) Define water.

(b) Name four different kinds of natural water.

5. Unknown liquid boils at 100˚C and freezes at 0˚C. It has a density of 1g/cm3 at

4˚C

(a) What is the name of that liquid

(b) Give one test of that liquid.

6. Define potable water.

7. (a) Draw the water cycle and describe the stages in the process

(b) Explain the importance of water cycle.

8. You are given a sample of colourless liquid and you carry out some tests. Here

are the results;

Test

Result

Add anhydrous Copper(ii)Sulphate

Turns from white to blue

pH

7

Evaporate to dryness

White residue

Boiling temperature

105˚C

(a) Which is the best test to tell you water is in the liquid?

(b) What does the test suggest about the liquid?

9. (a) Explain the meaning of these terms as far as water treatment is concerned;

filtration, sedimentation, chlorination.

(b)List down four methods of purifying water for cities and towns.

137

Chapter 11

Fuels and

Energy

Fuel is any substance that stores and releases usable energy when its chemical or physical structure is converted.

Fuel releases its energy either through chemical means such as burning or

nuclear means such as nuclear fission and nuclear fusion. An important property

of fuel is that its energy can be stored to be released only when needed and the

release of energy is controlled in such a way that the energy can be harnessed to

produce work.

Energy is defined as the capacity to perform work. Energy exists in several

forms such as heat, kinetic or mechanical energy, light, potential energy, electrical

energy or other forms.

11.1 SOURCES OF FUEL

The energy we use every day comes in many different forms from many

different sources, but they can be categorized as two basic types: non-renewable

(those that cannot be replenished in a short length of time) and renewable (those

that can be replenished in a short length of time).

NON - RENEWABLE FUELS

These are fuels that cannot be replenished in a short length of time. They

are considered non-renewable due to the many years it took to create them and

mankind’s inability to synthesize similar fuels readily. Examples of non-renewable

energy sources are: crude oil, liquefied petroleum gas (LPG), natural gas, coal and

uranium. All except uranium are called fossil fuels because they originate from

decaying plants and animal matter.

RENEWABLE FUELS

These are fuels that can be replenished in a short length of time. The renewable

fuels are abundant and constantly replenished.

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F u e l s a n d E n e r g y

FOSSIL FUELS

Fossil fuels are fuels formed by natural processes such as anaerobic

decomposition of buried dead organisms.

➢ The age of the organisms and their resulting fossil fuels is typically millions

of years.

➢ Fossil fuels contain high percentages of carbon.

➢ Fossil fuels include coal, petroleum and natural gas.

➢ The energy that is released when fossil fuels are burned originates from the

sun. It was captured by the plants through photosynthesis. The same principle

also applies to charcoal and wood.

EFFICIENCY OF A FUEL

During releasing energy a fuel usually does not release all of the stored energy.

Much of the energy is wasted. In a normal fire used for cooking by employing fire

wood or charcoal, only about one fifth of the energy stored is actually released in

a useful form to the surroundings. The efficiency of the energy source is the ratio

of the useful energy released to the total amount of energy stored. It is usually

expressed as a percentage, so the efficiency of the fire is about 20%. The efficiency

of a car is 40%, a lamp is 10% and that of a microwave cooker is 80%.

CHARACTERISTICS OF GOOD FUELS

There are characteristics which determine good fuel because fuels differ in

quality. Good fuels have the following characteristics:

1. They ignite easily

2. They burn well without explosion

3. They give out a lot of heat

4. They are not expensive

5. They have low smoke and ash content

6. They can be stored and transported easily.

7. They do not give off dangerous by-products like poisonous fumes.

8. They are easily controlled.

Activity 11.1; to compare the energy values of foods

Collect small samples of food e.g. bread and biscuit.

1. Clamp the boiling tube as shown in figure 11.1

2. Use a measuring cylinder to transfer 20cm3 of water to the boiling tube.

3. Put a thermometer into the water and take the temperature of the water.

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4. Fasten the first food sample to a long needle and set the food alight with a

Bunsen burner flame.

5. Immediately hold the flame under the boiling tube. Stir the water with a

thermometer.

6. Record the highest temperature reached.

7. Repeat the procedure with each of the foods.

8. The food that produces the largest temperature rise will have released the

most energy.

Thermometer

HEAT

Figure 11.1; comparing the energy values of food.

Consider the table below which shows energy content of some foods.

Table 11.1: energy content of some food stuffs

Food

Energy content (Kj/100g)

Fish

300

Butter

3000

Fresh fruit

200

Bread

1000

It can be seen from the above table that foods that contain carbohydrates (e.g.

bread), and fats (e.g. butter) are the ones which contain greatest amounts of stored

energy.

Activity 11.2; To carry out destructive distillation of wood in the laboratory.

Procedures;

1. Set up the apparatus as shown in figure 11.2

2. Fill the beaker with cold water and insert in it a test tube containing a mixture

of acid liquor and wood tar as shown in figure 11.2. Do not insert the trough in

the position at this stage.

3. Heat the saw dust gently and light the wood. The gas escapes from the jet.

4. If no gas is escaping try to heat more strongly.

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5. Fill in the trough and collect the wood gas over water as shown in figure 11.2.

Collect several gas jars if possible.

6. Go on heating strongly; when there is no further change, remove the trough

of water.

7. Remove the source of heat and allow the apparatus to cool.

8. Examine the products, and test them in second test tube with blue litmus paper.

Saw dust

Wood gas

HEAT

Acid liquor

Wood tar

Water

Figure 11.2; destructive distillation of wood.

Destructive distillation of wood can be carried out locally. Consider the

following diagram

Wood kiln

Figure 11.3; local method of destructive distillation of wood.

Discussion questions

1. What fuels do you use at home? Can you find out the different costs of fuels?

2. What fuel do you think is most efficient for use at home?

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11.2 CATEGORIES OF FUELS

Fuels can be categorized into three states of matter as gaseous fuel, liquid fuel

and solid fuel. The efficiency of a fuel can be explained in terms of the amount of

heat energy that can be produced when that fuel is completely burned.

An efficient fuel produces a lot of heat energy when a small amount is used.

It has also a smooth combustion which does not produce a lot of gaseous wastes

into the environment. Example kerosene produces very hot flame and less soot

when used at home. Fire woods burn slowly and produce smoky flame. The heat

produced is not high. Charcoal is more efficient than firewood and produces reliable

heat energy.

In industries, coke and carbon monoxide are important and efficient fuels

as they produce much heat energy. Gas fuels like biogas, coal gas and ethylene

are important and efficient fuels for large production of goods in industries. For

classification, gaseous fuels are more efficient, followed by liquid fuels and lastly

the least efficient are solid fuels.

Fuels can also be categorized on the basis of their occurrence, fuels can be

classified into natural (primary) and artificial (secondary) fuels. Natural fuels occur in nature. Examples are; wood, coal, peat, petroleum and natural gas.

Artificial fuels are either manufactured in industries or derived from the primary

fuels by refining. Examples are coke, kerosene, petrol, coal gas and producer gas.

Consider the following table.

Table 11.2 categories of fuels according to the physical states.

Solid fuels

Liquid fuels

Gas fuels

Coal, charcoal and fire

Diesel, kerosene and petrol Acetylene for welding, water

wood

gas and producer gas for

industries, natural gas and

town gas for domestic uses.

11.3 USES OF FUELS

The following is the list of uses of fuels by human being in several activities:

1. Fuels are used to run machines in industries and motor vehicles. Example

petrol, diesel and coke.

2. Fuels are used for cooking, boiling, and warming homes. Example charcoal,

kerosene, fire wood, coke and coal gas.

3. Fuels are used for lighting at home especially in rural areas where electricity

is unavailable. Example kerosene.

4. Fuels are used for drying. Example, tobacco leaves are dried in kilns by

burning woods.

5. Fuels are used for rocket propulsion and space ships travels. Example hydrogen

and uranium or plutonium.

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EFFECT OF FUELS ON ENVIRONMENTAL POLLUTION.

The fuels used mostly nowadays release harmful gases into the air that cause

pollution. Living things are being affected by polluted air. Acidic rains, green house

effect and health effects are all that can result due to animals come in contact with

polluted air. When burning coal and petrol, gaseous pollutants like sulphur dioxide,

and carbon monoxide can be released into the atmosphere. These gases lead to

acidic rain when dissolved in rain water. Green house effect can result after releasing of several gases into the atmosphere. Example carbon dioxide and chlorofluoro

carbons. In view of that, we must consider choosing a fuel to use in order to avoid

environmental pollution. In addition to that, extensive harvest of plants for firewood

and charcoal production would lead to effects like; desertification, disturbance on

water cycle, destruction of homes for wild life and eventually the ecosystem.

In order to avoid environmental hazards due to deforestation and pollution an

alternative source of fuel must be encouraged and this means that renewable energy

sources emphasized.

11.4 CONSERVATION OF ENERGY.

There are different forms of energy. Energy can be potential or kinetic. Potential

energy is energy at rest while kinetic energy is energy in motion.

Energy in motion is illustrated by winds, running water, ocean waves, moving

machines or falling bodies. Potential energy is stored in different forms. E.g.

Petroleum, coal, natural gas, muscles and batteries. Such energy does not do work

when it is stored.

The potential energy is capable of doing work when it is changed to other

forms of energy. Example heat and light. All energy changes which occurs during

chemical and physical changes must conform to the law of conservation of energy.

The law of conservation of energy states that.

“ Energy can neither be created nor destroyed but it can be converted

from one form to another”.

Example; the energy in a fuel is not destroyed when it burns but it simply changes

to other forms like heat and light.

CONVERSION OF ENERGY.

Energy can be converted from one form to another as it can be listed below.

1. Conversion of solar energy into light energy by using solar panel

2. Conversion of sound energy into electrical energy. Example using a

microphone.

3. Conversion of electric energy into sound energy. Example radio speaker.

4. Conversion of electric energy into light energy. Example electric bulb.

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5. Conversion of mechanical energy into sound energy. Example drum beating.

6. Conversion of chemical energy into mechanical energy. Example muscles.

7. Conversion of chemical energy into electrical energy. Example battery.

Activity 11.3 to convert chemical energy into electrical energy.

Procedure.

1. Clean pieces of copper and magnesium ribbon with sand paper.

2. Connect the battery and bulb with the magnesium and copper using connecting

wires with crocodile clips.

3. Dip the two pieces of metal into a beaker containing dilute sulphuric acid.

You should see the bulb light up for a short time, showing that chemical

energy is transferred into electrical energy.

11.5 BIOGAS AS RENEWABLE ENERGY

Biogas refers to a gas produced by breakdown of organic matter in the absence

of oxygen. Organic wastes such as dead plant and animal material, animal faeces

and kitchen waste can be converted into a gaseous fuel called biogas. Biogas is

produced by the anaerobic digestion of biodegradable materials such as biomass,

manure, sewage, municipal waste, green waste, plant material and crops.

Biogas comprise primarily methane (CH ) and carbon dioxide (CO ) and may

4

2

have small amount of hydrogen sulphide (H S), moisture and siloxanes.

2

Biogas is a renewable energy which is environmentally friendly. In creating

biogas the air should be absent and water present. The organic matter being cattle

dung, poultry wastes, night soil, crop residues, food processing and paper wastes.

Residue from the agricultural sector, such as spent straw, hay, and cane trash need

to be cut into small pieces in order to facilitate their flow into the digester reactor as well as to increase the efficiency of bacterial action. Succulent plant material

yields more gas than dry matter does.

THE WORKING MECHANISM OF A BIOGAS PLANT.

The biogas plant consists of two components; A digester or fermentation tank

and a gas holder.

The digester is a cube-shaped or cylindrical water proof container with an

inlet into which the fermentable mixture is introduced in the form of a semi-liquid

mixture of fine particles of manure.

The gas holder is normally an air proof steel container that, by floating like a

ball on the fermentation mix, cuts off air to the digester and collect the gas generated.

The organic waste matter is generally animal or cattle dung. These waste

products contain carbohydrates, proteins and fat material that are broken down by

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bacteria. These materials are soaked in water to give the bacteria a proper medium

to grow. Absence of air or oxygen is important because bacteria then take oxygen

from the waste material itself and in the process break them down.

The fermentation materials should not be too diluted or too concentrated,

because this may lead to low biogas production and insufficient fermentation activity

respectively. The ratio of raw material to water should be 1:1 i.e. 1500 kg of excrete

to 1500 kg of water. The loading of the raw material in the plant is determined by the

influent solid contents, retention time, and the digester temperature. High loading

is usually used in high ambient temperature. The plant is seeded with adequate

population of both the acid forming and methanogenic bacteria. These accelerate

the process of digesting sludge. Sometimes addition of lime or other alkali is done

in case the overloading of the seeding material is done. For the successive action of

the bacteria on the sludge, the pH in the plant must be maintained between 6.0 – 8.0.

This is because efficient digestion occurs at a pH near neutrality. The temperature

should be maintained at a range that usually favours the digestion process. At high

temperature biodigestion occurs faster, reducing the time requirement. A normal

period for the digestion of dung would be two to four weeks.

Consider the below diagram which shows the biogas plant.

Outlet of

Inlet

biogas

Outlet

Slurry

mixing

Effluent pit

Inlet pipe

Floating

gas cap

Outlet pipe

Dung and water mixture

Partition wall

Underground digester tank

Support for pipe

Fig. 11.4 the biogas plant

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Activity 11.4 A class to construct a small scale model of biogas plant.

Materials needed:

Biogas plant model, water, pipes, concrete, sewage, cow dung, animal and plant

waste.

Procedures

1. Select a place outside the classroom to construct the biogas plant

2. Using the materials you have collected construct a model of biogas plant.

Consider figure 11.4

3. The teacher to assist students in each step to construct the biogas plant.

4. Feed the biogas plant with organic residues and let them decompose for 2 – 3

weeks.

5. After three weeks collect the gas and demonstrate its uses.

Discussion questions.

1. What problems did you face during construction of the model?

2. How efficient was the gas generated?

THE USE OF BIOGAS AND ENVIRONMENTAL

CONSERVATION.

When biogas is used, many advantages arise. Utilization of biogas would

generate enough electricity and also could potentially help reduce global climatic

change. Biogas is a combustible gas, that when used produce no air pollutants on

burning. Wastes and biodegradable residue may be accompanied by a variety of

pollutants, but this can be corrected by remedying the carbon/nitrogen ratio of

manure through the addition of shredded bagasse or straw or by dilution.

Consider the below table which explains what 1 cubic meter of biogas can do.

Table 11.3 application of 1 cubic metre of biogas.

Application

1 cubic meter of biogas

Application

One 100w light bulbs for 6 hours

Cooking

Cooks 3 meals for family of 5-6 people

Fuel replacement

0.7 kg of petrol

Power

Can run a 1 horse power motor for 2 hours

Electricity generation

Can generate 1.25 kwh of electricity

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SUMMARY

(a). We use energy everyday in homes, industries, schools and in our bodies

(b). A fuel is a store of energy and when a fuel is used energy is released.

(c). Energy is measured in the unit of Joule (J)

(d). Fossil fuels include coal, natural gas and petroleum

(e). When a fuel releases energy it usually does not release all of the stored

energy. Much of the energy is wasted.

(f). The efficiency of the energy source is the ratio of the useful energy released to the total amount of energy stored.

(g). Energy can never be created nor destroyed but it can be converted from one

form to another.

(h). Biogas is the gas produced by the fermentation of organic matter, including

manure and other solid wastes under anaerobic conditions i.e. absence of air.

(i). A fuel can be in any of the three states of matter.

(j). Sources of fuels can be renewable or non – renewable.

REVIEW QUESTIONS.

1. Choose the most correct answer for the following questions.

i. The process of heating wood in the absence of air is called……………

A. Destructive distillation

B. Destructive condensation

C. Distillation

D. Energy conversion

ii. A fuel can be in …………….. states of matter

A.

Two

B.

Four C.

Three D.

Five

iii. …………… affects living things in different ways.

A. Atmospheric air

B. Polluted air

C. Non-poisonous gases

D. Fuel

iv. ………… comes from the heat energy present inside the earth.

A. Solar energy

B. Wind energy

C. Electric energy

D. Geothermal energy.

2. Match the items in list A with their corresponding statements in LIST B.

List A

List B

i. Most efficient category of fuel……

A. Coal, charcoal and fire wood.

ii. An example of renewable B. Wood

resources……

C. Natural gas and uranium

iii. Non renewable energy source…....

D. Remains of plants and animals that

iv. Fossil fuel………

lived millions of years ago

v. Solid fuels………………

E. Gas fuel

F. Liquid fuel.

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3. Write TRUE for a correct statement and FALSE for an incorrect statement.

i. Good fuel should have high heat content…………….

ii. Good fuel should be very expensive. …………….

iii. Good fuel should not be easily controlled…………….

iv. Burning of coal and petrol produces gaseous pollutants like sulphur dioxide

gas into the air. …………….

v. Deforestation leaves the land bare to agents of erosion like moving water,

animals and wind…………….

vi. Solar energy is available in plenty and cheap…………….

vii. Renewable energy is an environmental friendly because it does not pollute

the environment…..

viii. Potential energy is energy in motion while kinetic energy is energy at

rest..........

ix. Biogas is a combustible gas that when used produce no air pollutants on

burning……………

x. Use of biogas is cost effective and environmental friendly…………….

4. (a) What do you understand by renewable energy?

(b) What is the importance of renewable energy source?

(c) What do you understand by biogas plant?

(d) Mention organic materials that can be used in biogas plant.

5. (a) State the law of conservation of energy.

(b) Mention at least four ways by which energy can be converted from one form

to another.

6 (a) State the classes of fuels according to their physical states. Give three examples in each case.

(b) Classify fuels according to their efficiency.

7. How do you obtain fuel locally from your environment?

8. Describe an experiment that could be used to compare the amount of energy

released when wood, charcoal and coal are burned.

9. The table below gives the energy stored in different fuels, A-E and the energy

released when these are burned.

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(a) Copy and complete the table.

(b) Which fuel has the greatest efficiency?

(c)The same mass of fuel is burned in each case. Which of these statements is

true and which is false?

➢ The efficiency depends upon the amount of fuel burned.

➢ The more fuel burned the greater the energy released.

Energy

Energy stored in

Fuel Mass of sample of fuel

released by

Efficiency%

fuel

fuel

A

100g

1250kJ

450 kJ

36

B

100g

1500 kJ

600 kJ

C

100g

1000 kJ

450 kJ

D

100g

1800 kJ

450 kJ

E

100g

2000 kJ

400 kJ

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Chapter 12

Atomic structure

12.1 THE CONCEPT OF ATOM

The name atom comes from the Greek word “Atomos” which means

“Indivisible” . The two terms of “Atomos” in Greek are “a” which means “not”

and “Temno” which means “cut”. As early as 400 BC some people had thought about the structure of matter. Democritus, an ancient Greek thinker, suggested that

all matter was made of particles. However, he had no scientific evidence for this.

More powerful people like Aristotle disagreed with Democritus and so the ideas

were largely forgotten.

By definition, an atom is the smallest particle of an element that can take part in a chemical reaction. The concept of an atom was first studied by John Dalton in

his theory named Dalton atomic theory.

DALTON ATOMIC THEORY.

In 1803, Dalton developed his theory about the atom. The five main points of

Dalton’s atomic theory are;

1. All matter is made up of tiny indivisible particles called atoms.

2. Atoms can neither be created nor destroyed.

3. Atoms of any one element are identical and have the same mass.

4. Atoms of different elements have different masses.

5. Atoms may combine to form compounds, but the combination must involve

whole numbers of atoms.

Dalton’s atomic theory was the first step towards the formulation of modern

atomic theory. Dalton’s theory has been subjected to a lot of experimentation. This

has led to slight modifications to Dalton’s theory. The modifications are;

1. Atoms can be created or destroyed or split by means of nuclear reactions

2. Atoms of the same element may have different masses. These are called

‘Isotopes’, the existence of which is a challenge to Dalton’s ideas. Some atoms of the same element can be quite different in their mass.

3. An atom is made up of even smaller sub atomic particles called protons,

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neutrons and electrons.

4. Atoms of different elements may combine in many different ways. Example

the formation of carbon dioxide from carbon and oxygen.

12.2 SUB ATOMIC PARTICLES

The early years of the twentieth century saw many scientists working to

establish the structure of the atom. Prominent names in this area are; Thomson,

Rutherford, Chadwick, Bohr and many others. An atom is now thought to consist of

a very small and dense region, called the nucleus, surrounded by orbiting electrons.

The main sub atomic particles are protons, neutrons and electrons.

THE ELECTRON.

This is a particle with a unit of negative charge. It has a very small mass, about

1/1800 that of the proton. Its symbol is e

0

. At this basic level, we shall assume

-1

that the electrons orbit around the nucleus in electron shells at high speeds.

THE PROTON

This is a particle which carries a unit of positive charge. It has a mass

approximately the same as that of hydrogen atom, i.e. one atomic mass unit (amu).

Protons and neutrons form the nucleus of an atom. The symbol of proton is 1+1 P

THE NEUTRON

1

Neutron carries no charge at all. It is represented by letter 0 n. This mass of neutron is the same as that of proton. The table below gives a summary of the

properties of sub – atomic particles of an atom.

Table 12.1 Properties of sub – atomic particles.

Sub atomic

Relative

Symbol

Location Charge

Real mass (g)

particle

mass

Proton

+

1

In the

1

1.6726 x 10-24

1

+1 P

nucleus

Neutron

In the

0

1.6750 x 10-24

1

1

0 n

nucleus

Electron

e

-

0

Outside

1

9.109 x 10-28

1/1840

-1

the

nucleus

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Nucleus (Protons and neutrons)

Electron cloud

Fig.12.1; structure of an atom

12.3 THE ARRANGEMENT OF ELECTRONS IN AN ATOM

Electrons in an atom move in the space around the nucleus and are at different

distance from it. Electrons are arranged in shells called energy levels. An energy level is considered as a region around the nucleus most frequently occupied by

electrons of approximately the same energy. Each shell can hold only a certain

number of electrons. Each shell can hold a maximum number of electrons given

by the formula 2n2, where n is the number of the shell starting from that nearest to

the nucleus.

According to the formula, the first, second and third shells can contain a

maximum of 2,8 and 18 electrons respectively. The arrangement of electrons around

the nucleus is also known as the electronic configuration in which electrons are arranged in different energy levels.

Electron

Nucleus

Fig.12.2 Electron diagram of hydrogen

32

18

8

2

K L

M

N

1st

2nd

3rd

4th

Fig.12.3 Electron shells

Fig 12.4 Electronic arrangement of an element

With

atomic

number

19

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Table 12.2 electron arrangement of the first 20 elements.

Number of Number of Electronic

Element

Symbol

Period Group

electrons

shells

configuration

Hydrogen

H

1

1

1

1

Helium

He

2

2

2

1

0

Lithium

Li

3

2

2:1

2

1

Beryllium

Be

4

2

2:2

2

2

Boron

B

5

2

2:3

2

3

Carbon

C

6

2

2:4

2

4

Nitrogen

N

7

2

2:5

2

5

Oxygen

O

8

2

2:6

2

6

Fluorine

F

9

2

2:7

2

7

Neon

Ne

10

2

2:8

2

8

Sodium

Na

11

3

2:8:1

3

1

Magnesium Mg

12

3

2:8:2

3

2

Aluminium Al

13

3

2:8:3

3

3

Silicon

Si

14

3

2:8:4

3

4

Phosphorus P

15

3

2:8:5

3

5

Sulphur

S

16

3

2:8:6

3

6

Chlorine

Cl

17

3

2:8:7

3

7

Argon

Ar

18

3

2:8:8

3

0

Potassium

K

19

4

2:8:8:1

4

1

Calcium

Ca

20

4

2:8:8:2

4

2

Activity 12.1 To make the model of electronic diagram.

Materials: Flexible material (wire or thin stick), threads, clay soil or any 3 soft round small objects or fruits with different colours, water and gloves.

Procedures:

1. Make clay balls with a hole, or if you have soft round small objects or fruit

make a hole at the centre.

2. Put one on a flexible material (example wire or thin stick) like beads on its

thread.

3. Bend that material slowly to make a circle and tie it with a thread or plastic

string.

4. Make another clay ball and on it put other small balls one at the top another

at the bottom

5. Put a circle you have made on the table or ground and at the centre of it put

the clay ball you have prepared in (4) above. Draw that model. Repeat the

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exercise but this time making separate models as follows.

a. One circle with two balls on it and clay ball in (4) at the centre.

b. A model with two circles, inner circle (2 balls) and the outer (6 balls)

and one clay ball prepared in stage (4) at the centre.

c. A model with three circles, inner circle (2 balls). As normal put a clay

ball prepared in stage (4) at the centre. Lay them on the ground and

draw them.

Discussion questions.

1. What is the name of the diagram(s) you have drawn?

2. Basing on that diagram, what do

a. Balls on the string or soft wire represent?

b. String or wire represent?

c. Clay ball at the centre represent?

Activity 12.2

Draw the electronic configuration and electronic diagrams of the following atoms

a. Neon with atomic number 10

b. Sodium with atomic number 11

c. Hydrogen with atomic number 1

d. Magnesium with atomic number 12

e. Chlorine with atomic number 17

Activity 12.3

Students to role play on electronic arrangement of an atom. In this activity students will act as electrons and the teacher as the nucleus of an atom.

Procedures:

1. Students are assigned K,L,M and N shells.

2. The teacher should be at the centre and the students should go around in their

respective orbits (or shells)

3. Students should touch one another by their hands, so that they form cycles

around the teacher.

4. Students now in their shells, extend outwards as possible as they can.

5. Now the teacher should not move. Let students start moving around the teacher

in their respective shells. Relate this with electrons in their respective shells.

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Discussion questions.

1. What was the aim of the activity?

2. How many electrons are found in the K,L,M and N shells?

3. What was the aim of a teacher to be at the centre?

12.4 ATOMIC NUMBER, MASS NUMBER AND ISOTOPES

ATOMIC NUMBER (Z)

Atomic number is the number of protons in the nucleus of an atom. It is also

the number of electrons in a neutral atom. It is also part of the atomic structure

which identifies the element. Example the atomic number of hydrogen is 1 since it

has only one proton. A magnesium atom has 12 protons in the nucleus and therefore

its atomic number is 12.

Since the number of protons is equal to the number of electrons in the atoms,

the atomic number also indicates the number of electrons in the atom.

The official symbol for atomic number is Z. Atomic number is usually written

as a subscript in front of the symbol for the atom. Example Chlorine, where Z =

17, is written as Cl.

17

MASS NUMBER (A)

Mass number is the sum of the numbers of protons and neutrons in the nucleus

of an atom. The mass number is usually written as a superscript infront of the

symbol for the atom. Example Nitrogen atom, where A=14 is written 14N. The

official symbol for mass number is A.

It is possible to get the number of neutrons and number of electrons of an atom

if its mass number and atomic number are given.

Example 12.1

Atom M has a mass number of 16 and an atomic number 8. What is the neutron

number? What is the number of electrons in atom M?

Solution.

Mass number = 16, atomic number = 8.

(a) Neutron number = Mass number – Atomic number

= 16 – 8 = 8

(b) Number of electrons = number of protons

= Atomic number = 8

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Example 12.2

Complete the table below by indicating the number of protons, electrons and

neutrons of the atoms shown. The atomic number and mass numbers are given;

Atomic

Mass

Atom

Protons

Electrons

Neutrons

number

number

Sulphur

16

32

Chlorine

17

35.5

Argon

18

40

Potassium

19

39

Calcium

20

40

NUCLIDE NOTATION.

Atoms of different elements can be represented by symbols that indicate their

respective atomic numbers and mass numbers. Using an arbitrary element X, the

mass number (A) is placed on its upper left end while its atomic number (Z) is

placed on the lower left end. Then element X is represented as A

. This is known

Z X

as nuclide notation..

Example 12.3

39

19 X

You are given an atom

a. Write its electronic configuration.

b. Identify the number of neutrons and protons in the atom.

c. Draw its electronic diagram.

d. Identify the element X by its symbol and give the nuclide notation.

e. Show the representation of the nucleus of the atom.

Solution:

(a) Electronic configuration 2:8:8:1

(b) The number of neutrons = Mass number – Atomic number.

= 39 – 19 = 20.

Number of protons = atomic number = 19.

(c) Electronic diagram

156

A t o m i c S t r u c t u r e

(d) Element X is potassium; K and its nuclide notation is 39

.

19 K

(e) 19P

20n

ISOTOPES

Isotopes are atoms of the same element with the same number of protons but

different number of neutrons. Isotopes have the same number of protons, the same

number of electrons and the same atomic number. They have the same chemical

properties but may have slightly different physical properties. Isotopes have different mass numbers because they have different numbers of neutrons.

Chlorine for example has two common isotopes which are explained in the

table below.

Table 12.3 isotopes of chlorine.

Number of

Element Atomic number

Number of neutrons

Atomic mass

protons

Chlorine

17

17

18

35

17

17

20

37

37

The isotopes of chlorine can thus be written as 35

and

which are

17 Cl

17 Cl

known as chlorine -35 and chlorine -37 respectively.

17P

17P

18n

20n

Chlorine -35

Chlorine -37

Fig. 12.5; Isotopes of chlorine

Having learned about Isotopes, it is also important to understand the concept of

Isotopy. Isotopy is the tendency of atoms of the same element to contain the same

number of protons but different number of neutrons or mass number. Examples of

common elements that display isotopy are hydrogen, oxygen, chlorine and neon.

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I l l u s t r a t i v e C h e m i s t r y f o r S e c o n d a r y S c h o o l s

Consider the below table showing common isotopes.

Table 12.4 Examples of common isotopes.

Element

Symbol Atomic number

Isotopes

Abundance

1

(Protium)

99.99%

1 H

Hydrogen

H

1

2

(Deuterium)

0.01%

1 H

3

(Tritium)

very rare

1 H

12

98.9%

6 C

Carbon

C

6

13

1.1%

6 C

14

trace

6 C

Cl

35

75%

17

Chlorine

Cl

17

Cl

37

25%

17

16

99.8%

8 O

Oxygen

O

8

17

0.37%

8 O

18

0.20%

8 O

Ne

20

90.5%

10

Neon

Ne

10

21

0.3%

10 Ne

22

9.2%

10 Ne

RELATIVE ATOMIC MASS (R.A.M)

The relative atomic mass of an element is the average mass of one atom of the

element relative to 1 th the mass of one atom of carbon -12.

12

R.A.M = Average mass of atom of an element

1 th the mass of carbon -12 atom.

12

The relative atomic mass has no units.

The carbon atom was chosen as the standard atom (reference atom) and its mass

was arbitrarily chosen as 12 units. Then using a machine called mass spectrometer,

all the other atoms were compared to this standard atom. This reference is called

the carbon -12 scale. For example it was found that:

i. The magnesium atom was twice as heavy as the reference atom, so its

mass was put at 24.

ii. The hydrogen atom was 1/12 as heavy as the reference atom, so its mass

158

A t o m i c S t r u c t u r e

was put at 4.

iii. The helium atom was 1/3 as heavy as the reference atom, so its mass was

put at 4.

Example 12.4

A sample of chlorine gas contains 75% of the isotope 35

and 25% of the

17

37

Cl

other isotope, which is

. What is the relative atomic mass of chlorine?

17 Cl

Solution

Relative atomic mas (R.A.M) is given by the formula

R.A.M = Mass number Relative abundance

Mass number Relative abundance

of an Isotope

for an isotope

of an Isotope

for an isotope

R.A.M = 35 x 75

37 x 25

100

100

= (75 x 35) + (25 x 37) = 35.5

100

Therefore, the R.A.M of chlorine is 35.5

Example 12.5

12

Carbon has two main isotopes 13

and

with abundance of 1.11% and

6 C

6 C

98.89% respectively. Calculate its relative atomic mass (R.A.M).

Solution

R.A.M = 13 x 1.11

12 x 98.89

100

100

=

(13 x 1.11) + (12 x 98.89) = 12.1

100

Therefore, the R.A.M of Carbon is 12.01

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I l l u s t r a t i v e C h e m i s t r y f o r S e c o n d a r y S c h o o l s

Example 12.6

20

21

22

Neon has three isotopes;

(90.5%),

(0.3%) and 10Ne (9.2%).

10 Ne

10 Ne

Calculate the relative atomic mass of neon.

Solution.

20 x 90.5

21 x 0.3

22 x 9.2

RAM = 100

100

100

= (20 x 90.5) + (21 x 0.3) + (22 x 9.2) = 20.187

100

Therefore, atomic mass of Neon is 20.187.

SUMMARY

(a). An atom is the smallest particle of an element that can take part in a chemical

reaction.

(b). The concept of atom was first studied by John Dalton in his theory named

Dalton atomic theory.

(c). The three sub atomic particles are protons, neutrons and electrons.

(d). The arrangement of electrons around the nucleus is also known as electronic

configuration.

(e). Each shell in an atom can hold a maximum number of electrons given by the

formula 2n2

(f). Atomic number is the number of protons in the nucleus of an atom.

(g). The mass number is the sum of the numbers of protons and neutrons in the

nucleus of an atom.

(h). Isotopes are atoms of the same element with the same number of protons but

different number of neutrons.

(i). Relative atomic mass of an element is the average mass of one atom of the

element relative to 1/12th the mass of one atom of carbon -12.

160

A t o m i c S t r u c t u r e

REVIEW QUESTIONS.

1. Choose the best answer for each of the following questions.

(i) The smallest particle of an element that can take part in a chemical reaction is

called.

A. Matter

B. Substance

C. Atom

D. Element.

(ii) The first modern scientist to come up with ideas about building blocks of matter

which he gave the name “atoms” was……………..

A. John Dalton

B. Michael Faraday

C. Mendeleev

D. Thomas Edison.

(iii) The three sub – atomic particles are

A. Electron, nucleus and protons.

B. Protons, nucleus and electrons.

C. Electrons, nucleons and protons.

D. Protons, neutrons and electrons.

(iv) Electronic arrangement in atoms is also known as

A. Electronic configuration

B. Electronic arrangement

C. Electronic shells

D. Electron cloud.

(v) Is the tendency of atoms of the same element to contain the same number of

protons but different number of neutrons or mass number.

A. Isotopes

B. Isotopy

C. Allotropy

D. Atomic number.

2. Match the items in List A by writing a letter of the correct statement of List B

in the space provided.

LIST A

LIST B

i. Proton

37

35

A. Cl and Cl

ii. Neutron

17

17

B. Negatively charged particle of an atom

iii. Electron

C. Positively charged particle of an atom.

iv. Carbon-12

D. Neutral charged particle of an atom.

v. Isotopes

E. The reference for relative atomic mass.

F. Partial y charged particle of an atom.

3. Write TRUE for a correct statement and FALSE for an incorrect statement.

i. Atoms can be created or destroyed or split by means of nuclear

reactions……………….

ii. An atom is made up of even smaller sub atomic particles called protons,

neutrons, and electrons.

iii. The atoms always combine in simple ratios……………….

iv. The atoms of any one element are identical and have the same chemical

properties and the same mass………………..

v. Dalton never imagined that any one would ever be able to see an atom.

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I l l u s t r a t i v e C h e m i s t r y f o r S e c o n d a r y S c h o o l s

4. (a) What is electronic configuration?

(b) An atom has mass number of 24 and neutron number 12;

(i) What is its atomic number.

(ii) Write down its electronic configuration.

5. Describe the changes in properties across the period from sodium to argon.

6. A certain element P, was found by mass spectrometry – a sophisticated technique

for measuring the relative masses of element or compounds – to have Isotopes.

20 P

4

, 20 P

6

, 20 P

7

and 20 P

8

.

Their relative abundance were, respectively 2,24,22 and 52, as determined by

a mass of the element P, can you suggest a name for the element? (P is not the

real symbol of the element).

7. Draw the structure of an atom according to

(a) Dalton

(b) Rutherford.

8. State the number of protons, neutrons and electrons in the following Isotopes.

(i) 12

(ii) 14

(iii)

3

(iv)

87

(v)

238

92 U

38 Sr

1 H

6 C

6 C

9. An Isotope of carbon has a mass number of 13 and an atomic number of 6.

(i) Write its nuclide notation.

(ii) How many neutrons does it have?

(iii) How many electrons does it have?

10. Copy and complete the table below;

Atomic

Relative

Number of

Number of

Element

number

atomic mass

protons

neutrons

Hydrogen

1

1

Helium

2

4

Lithium

6.9

Beryllium

4

9

Boron

5

10

Carbon

6

12

Nitrogen

14

Oxygen

16

Fluorine

9

19

Neon

10

162

A t o m i c S t r u c t u r e

Atomic

Relative

Number of

Number of

Element

number

atomic mass

protons

neutrons

Sodium

11

23

Magnesium

12

Aluminium

13

27

Silicon

Phosphorus

15

31

Sulphur

16

32

Chlorine

17

Argon

18

Potassium

19

Calcium

20

163

Chapter 13

Periodic classification

Periodic classification is the systematic grouping of elements into periods and groups in the periodic table according to their similarities and differences in electro negativity, ionization energy, melting point, boiling point, density, atomic radius

and reactivity,metallic character and non - metallic character.

Many elements were known since the ancient times but up to the middle

of seventeenth century, the ideas of elements were disordered and the elements

classification did not exist.

13.1 THE PERIODIC TABLE

Periodic table is the arrangement in which all known elements are arranged in

their respective periods and groups in order of increasing atomic number. The first

periodic table was devised by the Russian chemist Mendeleev in 1869. He devised

it to help his students understand chemistry

Mendel eev did the following:

➢ He arranged the known elements in groups according to their properties in

order of increasing atomic mass, but left gaps for undiscovered elements.

➢ He predicted the properties of the missing elements from the properties of the

elements above and below them in the table.

➢ He proposed a periodic law, which states

“The properties of elements are a periodic function of their atomic masses”

When elements were arranged in order of increasing atomic weights, the

properties of elements fell into repeating patterns. The vertical columns of

similar elements were called groups. The horizontal rows were called periods.

➢ He listed separately (in group 8) some elements which did not appear to fit into

any group i.e. iron, cobalt, nickel etc.

I l l u s t r a t i v e C h e m i s t r y f o r S e c o n d a r y S c h o o l s

164

P e r i o d i c C l a s s i f i c a t i o n

Table 13.1; Mendel eev ‘s periodic table

I

II

III

IV

V

VI

VII

VIII

Group

A

B

A

B

A B A

B A

B

A

B A

B

Period 1

H

Period 2 Li

Be

B

C

N

O

F

Period 3 Na

Mg

Al

Si

P

S

Cl

Period 4 K

Ca

T

V

Cr

Mn

Fe,Co,Ni

Period 5

Cu

Zn

As

Se

Br

Period 6 Rb

Sr

Nb

Mo

Ru,Rh,Pd

Period 7 Ag

Cd

In

Sn

Sb

Te

I

Limitations of Mendel eev ’s classification;

➢ The appearance of two families in one group, e.g. K and Cu, Ca and Zn.

➢ Some elements had to be placed in reverse order of their atomic weight so as to

fit into groups having similar properties. (e.g. Ar and K, I and Te, Co and Ni)

MODERN PERIODIC TABLE

The modern periodic table of elements is a table of elements arranged

systematically according to their increasing atomic numbers. It is a result of several modifications to Mendeleev’s periodic table. The modifications were made as new

elements were discovered and new theories developed to explain the chemical

behavior of elements.

The modern periodic law states that;

The properties of elements are a periodic function of their atomic numbers.

13.2 PERIODICITY

Periodicity is the regular periodic changes of elements due to their atomic

number.

FEATURES OF THE MODERN PERIODIC TABLE.

The main features of the modern periodic table are;

➢ The elements are placed in order of their atomic numbers instead of their atomic

masses.

➢ A new group of ‘noble gases’ has been discovered and included on the right

hand side of the table as group O.

➢ Gaps in the table left by Mendeleev, have been filled as new elements e.g.

germanium have been discovered. Man-made elements have also been

accommodated in the table.

➢ Some groups, have names as well as numbers. Group I is called the alkali

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I l l u s t r a t i v e C h e m i s t r y f o r S e c o n d a r y S c h o o l s

metals. Group II is called the alkaline earth metals. Group (VII) is called the

halogens. Group O is called the noble gases.

➢ Metals have been clearly separated from non-metals.

➢ There are five blocks of similar elements in the modern periodic table.

s

Non-metals

ase g

Reactive

Transition

metals

‘elements’

Noble

Metalloids

Figure 13.1 Blocks in the modern periodic table.

Table 13.2, position and electronic arrangement of the first twenty elements

in the Periodic Table.

Groups

I

II

III

IV

V

VI

VII O

1

2

Period 1

2He

1 H

Period 2 3 Li 4 Be 5 B

6 C 7N 8O 9F

Ne

10

2.1

2.2

2.3

2.4

2.5

2.6

2.7

2.8

Period 3

Na

11

12 Mg Al

13

Si

14

P

15

S

16

Cl

17

Ar

18

2.8.1

2.8.2

2.8.3

2.8.4

2.8.5

2.8.6

2.8.7

2.8.8

19

20

Period 4

K Ca

2.8.8.1

2.8.8.2

The modern periodic table can be divided into blocks as it can be explained

as follows:

(a) Reactive metals

These are metals of Group I and Group II. They include reactive metals such

as potassium, sodium, calcium, and magnesium. They are the alkali metals and

alkaline earth metals.

(b) Transition metals

Metals in this block resemble each other slightly across the period and down

the group. They are less reactive than the alkali metals, but they have greater

tensile strength.

(c) Metalloids

These are semi-metals which are fairly unreactive. They have some metallic

and some non-metallic properties. Examples of these metals are: gallium,

germanium and indium

166

P e r i o d i c C l a s s i f i c a t i o n

(d) Non-metals

Non-metals form a group of elements which are also reactive. They have

properties opposite to metals. They react by receiving electrons. With this

behavior they have negative ions. Members of this group include nitrogen,

chlorine, sulphur and bromine.

(e) Noble gases

Noble gases are also known as inert gases. They are gaseous non-metals which

are very unreactive. They form the last group in the periodic table. These

elements are stable and unreactive because they have complete outermost shell.

Examples are Helium, Neon and Krypton.

13.3 GENERAL PERIODIC TRENDS

The general periodic trends can be explained as follows;

1. Melting point

Melting point is the temperature at which a solid melts to form a liquid. Down

a group the melting points increases and decreases along a period from left to

right of a periodic table.

2. Boiling Point

Boiling point is the temperature at which a liquid boils to form vapour. Generally

metals have high boiling points than non-metals. The boiling point increases

down a group and decreases along a period from left to right of the periodic table.

3. Density

Is the degree of compactness of a substance which means it is the mass per unit volume of a substance. Density increases down the group and decreases across

the period from left to right.

4. Electronegativity

Electronegativity is the ability of an atom to attract an electron towards itself.

Across the period from left to right electronegativity increases. Down the group

electronegativity decreases.

5. Ionization Energy

Ionization energy is the energy required to remove an electron from an atom or

ion. Across the period from left to right ionization energy increases and down

the group ionization energy decreases.

6. Atomic Radius

Atomic radius is the distance between the nucleus of an atom and the outermost

stable energy level. The atomic radius of elements across the period decreases

from left to right. Atomic radius increases down the group as successive energy

levels are filled.

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I l l u s t r a t i v e C h e m i s t r y f o r S e c o n d a r y S c h o o l s

7. Metallic Character

Element in group I, II, III and IV are metals. The metallic character of an element

is its ability to lose electrons more easily from the valency shell. When an atom

lose electrons from the valency shell it becomes positively charged. The metallic

character decreases along a period from left to right. Down the group metallic

character increase.

8. Non-Metallic Character

Non-metallic character is the tendency of an atom to accept an electron more

easily into its valency shell. Across the period from left to right non-metallic

character increases and down the group metallic character decreases.

SPECIFIC GROUP TRENDS.

Arranging the elements in the periodic table results in the elements to fall

in groups of similar properties. Example, all elements in group two have the

same number of electrons in their outermost shell. This explains why elements

have similar properties. If elements of the same group are studied closely, some

differences in properties are noted, some properties increase or decrease gradually

toward one direction.

1. Group One Elements (Alkali Metals)

This group consists of five metals which are lithium (Li), sodium (Na),

potassium (K), rubidium (Rb) and caesium (Cs). These elements have one electron

each in their outer energy level.

Francium (Fr) is also an alkali metal but rarely included in the group. It is

among the rarest naturally occurring elements.

Lithium, sodium and potassium all react very readily with water or air and are

stored under paraffin oil. The three metals have the following properties:

Physical Properties;

i. They are good conductors of heat and electricity.

ii. They are soft metals

iii. They have low density.

iv. They have shiny surface when freshly cut.

Chemical Properties;

i. They burn in oxygen or air with a characteristic flame colour to form white

solid oxides.

Metal + oxygen→ metal oxides.

The oxides dissolve in water to form alkaline solutions of the metal hydroxide.

Metal oxide + water → metal hydroxide.

ii. They react vigorously with water to give the alkaline solution and hydrogen gas.

Metal + water → metal hydroxide + hydrogen.

168

P e r i o d i c C l a s s i f i c a t i o n

Table 13.3 trends in group I

)

y

)

er(Z)

˚C

)3

˚C

ol)

t (

tivit

t (

ation

adius

es)

tion

oin

oin

onic

y (g/cm

o nega

tr

ometr

gy (kj/mol)

tr

tomic numb

tomic r

Ioniza

elting p

Name(symb

A

Elec

configur

A

(pic

st 1 ener

M

Densit

Elec

Boiling p

Lithium (Li)

3

2:1

152

526

180 0.54 1.0

1347

Sodium (Na)

11 2:8:1

186

504

98

0.97 0.9

883

Potassium (K)

19 2:8:8:1

231

425

64

0.86 0.8

774

Rubidium (Rb)

37 2:8:18:8:1

244

410

39

1.5

0.8

688

Caesium (Cs)

55 2:8:18:18:8:1

262

380

29

1.9

0.7

678

2. Group Two Elements (Alkali Earth Metals).

Group II consist of six metals which are beryllium (Be), magnesium (Mg),

calcium (Ca), strontium (Sr), barium (Ba) and radium (Ra).The elements have

two electrons each in their outer energy level. These metals have the following

properties.

Physical Properties

i. They are harder metals than those in group I.

ii. They are silvery grey in colour when pure and clean but they tarnish quickly

when left in air due to the formation of the respective metal oxide.

iii. They are good conductors of heat and electricity.

Chemical Properties

i. They burn in oxygen or air with a characteristic flame colour to form a solid

white oxide

Metal + oxygen → metal oxide.

ii. They react with water but much less vigorously than the elements in group 1.

Metal + water → metal hydroxide + hydrogen.

iii. The metals become more reactive as we move down the group.

169

I l l u s t r a t i v e C h e m i s t r y f o r S e c o n d a r y S c h o o l s

Table 13.4 trends in group II

gy

)

y

)

˚C

˚C

ol)

er (Z)

t (

tivit

t (

ation

adius

er)

tion ener

oin

oin

onic

y (g/cm3)

o nega

tr

omet

tr

tomic numb

tomic r

ioniza

J/mol)

elting p

Name (symb

A

Elec

configur

A

(pic

st 1 (K

M

Densit

Elec

Boiling p

Beryllium (Be)

4

2:2

112

899

1287 1.85

1.5 2469

Magnesium (Mg) 12 2:8:2

160

738

649

1.74

1.2 1090

Calcium (Ca)

20 2:8:8:2

197

590

839

1.55

1.0 1484

Strontium (Sr)

38 2:8:18:8:2

215

549

769

2.63

1.0 1382

Barium (Ba)

56 2:8:18:18:8:2

217

503

1850 3.51

0.9 1897

3. Group Seven Elements (Halogens)

The halogens in the periodic table consist of five chemically related elements,

fluorine (F), chlorine (Cl), bromine (Br), iodine (I) and astatine (At)

Table 13.5 trends in group VII

)

y.

)

˚C

ol)

er (Z)

t(

tivit

˚Ct(

ation

adius

er)

tion

oin

oin

onic

y (g/cm3)

o nega

tr

omet

gy (kj/mol)

tr

tomic numb

tomic r

ioniza

elting p

Name (symb

A

Elec

configur

A

(pic

st 1 ener

M

Densit

Elec

Boiling p

Fluorine (F) 9

2:7

71

1681.0 -219.62 0.017

3.98

-188.12

Chlorine (Cl) 17

2:8:7

99

1251.2 -101.5

0.0032

3.16

-34.04

Bromine (Br) 35

2:8:18:7

114

1139.9 -7.3

3.1028

2.96

58.8

Iodine (I)

53

2:8:18:18:7 133

1008.4 114

4.933

2.66

184.3

Physical Properties;

i. Halogens range from solid (I ) to liquid (Br ) to gaseous (F and Cl ) at room

2

2

2

2

temperature.

ii. Fluorine is a poisonous pale yellow gas, chlorine is a poisonous pale green gas,

bromine is toxic and caustic brown volatile liquid, and iodine is shiny black

solid which easily sublimes to form a violet vapour on heating.

iii. At room temperature all the halogens exist as diatomic molecules.

iv. The melting points, boiling points, atomic radii and ionic radii all increase on

descending the group.

170

P e r i o d i c C l a s s i f i c a t i o n

Chemical Properties;

i. They are highly reactive especially with alkali metals and alkaline earth metals

forming stable ionic crystals.

ii. They have very high electronegativities.

iii. The halogens oxides are acidic and the hydrides are covalent.

iv. The halogens form organic compounds which are best known for their industrial

and environmental impact. Example PVC and DDT.

v. They oxidize metals to form halides.

SUMMARY

(a). In 1817, a German scientist, Dobereiner, noted that there was connection

between relative atomic mass and chemical properties of an element.

(b). Dobereiner arranged elements into groups of three and called ‘Dobereiner’s

triads’

(c). In Dobereiner’s triads the relative atomic mass of the middle element was an

average of the mass of the other two.

(d). It was noted that very few elements formed a group of three and therefore

the Dobereiner classification was discouraged.

(e). In 1866, a British chemist, John Newlands, thought of the idea of arranging

atomic masses.

(f). Newlands noticed that an element tended to display characteristics similar to

the 8th element in front of it.

(g). Newlands arranged the elements in columns according to a law he called the

Law of octaves

(h). The Newland classification was unfortunate since he grouped together

certain elements which had very different characteristics. Example oxygen

was placed in the same group as iron and sulphur and therefore his ideas

were rejected.

(i). A Russian chemist, Dimitri Mendeleev, later developed Newland’s ideas and

convinced other chemists to use them.

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REVIEW QUESTIONS

1. The figure below is a skeleton of the periodic table. Only a few elements have

been represented. The letters used are not the official chemical symbols of the

elements. You may use these letters in attempting this question.

(a) write the letters of three elements in the same period.

(b) write the letters of the elements commonly known as;

(i) Alkali metals

(ii) Alkaline earth metals

(iii) Halogens

(iv) Noble gases

(v) Transition metals.

(c) write the letters of;

(i) Seven metals and

(ii) Six non-metals.

Z

P

U

T

Y G

Q

R

V

J

M

H

W

2. Choose the most correct answer for the following questions;

i. The first periodic table was devised by the Russian chemist known as;............

A Dobereiner

B Newland

C Medvedev D Mendeleev.

ii. Newland introduced the...................... which states that “elements with similar chemical properties have atomic numbers which differ by 7”.

A. Law of octaves

B. law of elements

C. law of octate

D. Periodicity.

iii. ______is the regular periodic changes of elements due to their atomic number.

A. Periodic law

B. Periodic table

C. Periodic classification D. Periodicity

iv. The vertical columns of elements in the periodic table are called;__________

A. Groups

B. Periods

C. Metals

D. Non-metals

v. The metallic elements in Group I of the periodic table are called;_______

A. Alkali elements

B. Alkali earth metals

C. Alkali compounds

D. Alkali metals.

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3. Match the items in LIST A with their correct statements in LIST B;

List A

List B

i. The modern periodic law

A. The temperature at which a liquid boils to

ii. Group II metals

form a gas.

iii. Elements which display

B. The temperature at which a solid melts to

metallic and non-metallic

form a liquid.

characteristics

C. The properties of elements are a periodic

iv. Melting point

function of their atomic numbers.

v. Boiling point

D. Alkaline metals

E. Alkaline earth metals

F. Metalloids

4. Write TRUE for a correct statement and FALSE for an incorrect statement.

i. The atomic radii of elements in a periodic table decrease from left to right. . .... .

ii. Elements to the left of the periodic table show non-metallic properties...........

iii. Electronegativity decreases from left to right across the period......................

iv. Densities of elements increases down the group...............................................

v. Electronegativity and ionization energy decrease down the group.................

5. Explain how electronegativity, atomic radius, metallic properties and non-

metallic properties vary;

(a) Across a period from left to right of the periodic table.

(b) down a group

6. What do you understand by;

(a) modern periodic table

(b) periodicity

(c) transition elements

(d) modern periodic law

(e) law of octaves

7. A metallic element Y reacts slowly with water to give a strongly alkaline solution.

In which group of the periodic table would you place Y?

8. With two examples briefly explain the meaning of metalloids.

9. (a) Define noble gases.

(b) Name the group which they belong in a periodic table.

(c) Explain briefly how noble gases;

(i) Are arranged in electronic configuration

(ii) Undergo chemical reaction.

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Chapter 14

Formula, Bonding

and Nomenclature

14.1 CHEMICAL FORMULAE

A formula is a method of representing a molecule of a compound. Example

the chemical formula of limestone is CaCO . The formula of a compound expresses

3

the composition of a substance in terms of the relative numbers of atoms that have

chemically combined together. In a formula each atom present is represented by its

symbol. The symbol of an element represents one atom of the element. A symbol is

shown by one or two letters, the first of which is always a capital.

One atom of chlorine is represented by the symbol Cl. When two atoms of

chlorine combine, the molecule formed is represented by the formula Cl . Water is

2

denoted by the formula H O which contains two atoms of hydrogen and one atom

2

of oxygen combined. Two molecules of water are written 2H O; the two infront of

2

the formula multiplies everything that follows it.

Qualitatively H O means a water molecule is made of hydrogen and oxygen

2

atoms. The quantitative meaning of water gives us the amounts:

➢ One molecule of water

➢ Two hydrogen atoms for every one

➢ One mole of water molecule

oxygen atom.

➢ 18g of water

Consider the table below showing chemical formulae of some compounds.

Table 14.1 Chemical formulae of some compounds.

Compound

Chemical formula

Calcium hydroxide

Ca(OH)2

Zinc chloride

ZnCl2

Sodium carbonate

Na CO

2

3

Copper(II)sulphate

CuSO4

Sulphuric acid

H SO

2

4

Hydrogen peroxide

H O

2 2

Potassium chlorate

KClO3

Nitric acid

HNO3

I l l u s t r a t i v e C h e m i s t r y f o r S e c o n d a r y S c h o o l s

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14.2 VALENCY

Valency is the combining capacity of an element. Atoms lose, gain or share

electrons which are equal to their valencies in forming compounds.

Consider the following examples.

(1). All elements which have one electron in their outermost shell have a valency

of 1. E.g. sodium (2:8:1) and potassium (2:8:8:1). Elements with two, three

and four electrons in their outer most shells have valencies of 2, 3 and 4

respectively e.g. calcium (2:8:8:2) has two electrons in its outermost shell,

therefore its Valency is 2, Aluminium (2:8:3) has three electrons in its outer

most shell and its valency is 3 and carbon (2:4) has four electrons in its outer

most shell and its valency is 4.

(2). Elements with electrons greater than 4 in the outer most shells, their valency

can be calculated by using the formula; V= 8 -n.e

Where by;

V = Valency

8 = represents 8 electrons in the fulfilled outermost shell.

n.e = number of electrons in the outermost shell.

(a). The electronic configuration of oxygen is 2:6 the number of electrons in

the outermost shell is 6.

V = 8-n.e

V = 8-6

V = 2

Therefore, Valency of oxygen is 2.

(b). Electronic configuration of fluorine is 2:7. The number of electrons in the

outermost shell is 7.

V = 8-n.e

V = 8-7 = 1

Therefore, valency of fluorine is 1.

14.3 OXIDATION STATES

Oxidation states (also oxidation number) tell us the number of electrons a

particular element has lost, gained or shared in forming a compound. When an

element is alone, its oxidation state is zero.

There is a close relationship between valency and oxidation state, but they

are not the same thing. The oxidation state of an atom in a compound is assigned

relative to the other elements in that compound. While oxidation states are an

arbitrary assignment, valency is a fixed value. Oxidation state may be assigned a

positive or a negative number.

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VARIABLE VALENCY

Some elements have more than one valency. Such elements can change their

oxidation states. Examples are iron and copper.

Compounds in which iron has oxidation number +3 e.g. FeCl and those which

3

the oxidation number is +2 e.g. FeCl are well known. Copper has oxidation number

2

+2 in many of its compounds e.g. CuSO and only a few compounds with the

4

oxidation number +1 e.g. Cu O. If an element has a variable valency, its valency

2

must be shown by a roman numeral in the names of all its compounds.

Consider the table below which show a summary on valencies;

Table 14.2 Summary of valencies

Valency 1

Valency 2

Valency 3

Metals

Potassium,K, Sodium,Na

Calcium, Ca

Aluminium ,Al

silver,Ag

Magnesium, Mg Iron (III)

Ferric, Fe.

Non-metals Hydrogen,H, Chlorine ,Cl

Oxygen, O

Nitrogen,N phosphorus

Sulphur, S

,P

Radicals

Ammonium NH+

2-

Phosphate(V)oxide,

4

Carbonate CO

3

3-

Hydroxide OH-

Sulphate SO2-

PO4

4

Nitrate NO-

Sulphite, 2-

3

SO3

Nitrite NO-2

Hydrogen carbonate HCO-3

Example 14.1;

Give the oxidation number of Cr in Cr O2-

2

7

Solution

Total charge on the dichromate ion = -2

For oxygen -2x7 = -14

Then, 2Cr-14 = -2

2Cr = +12

Cr = +6

Therefore, the oxidation number of Cr is +6.

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Example 14.2

Find the oxidation state of chlorine in the compound KClO .3

Solution

The oxidation number of potassium is +1.

The oxidation number of oxygen is -2. With three oxygen atoms the oxidation

number is given by; (-2x3) = -6

Therefore, +1 + Cl-6 = 0

Cl = 6-1 = +5

Therefore, Oxidation number of chlorine in KClO is +5

3

Example 14.3

Calculate the oxidation number of sulphur in H SO .

2

4

Solution

Total oxidation state of H SO = 0

2

4

Oxidation state for H= +1; O = -2; S =?

H SO is a neutral compound and therefore the oxidation number of the molecule

2

4

is zero.

Therefore;

2x (+1) + S + 4x (-2) = 0

+2 + S - 8 = 0

S - 6 = 0

S = +6

Therefore, the oxidation number of sulphur in H SO is +6.

2

4

Example 14.4

Calculate the oxidation state of nitrogen in NH +4

Solution

Total oxidation state of NH+4 = +1

Oxidation state for H = +1, N=?

Therefore;

N + (4X (+1)) = +1

N + 4 = +1

N + 4 - 4 = +1 - 4

N = - 3

Therefore, the oxidation state of nitrogen in NH+ is -3

4

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Example 14.5

What is the oxidation number of sulphur in the sulphate ion, SO2- .

4

Solution;

The total charge on the sulphate ion is -2.

The oxidation number of oxygen is -2.

Then;

S + (-2x4) = -2

S-8 = -2

S = 8-2

S = +6.

Oxidation state of sulphur in SO2- is +6.

4

Example 14.6

Calculate the oxidation state of sulphur in HSO-3

Solution

Total oxidation state of HSO- = -1

3

Oxidation state for H=+1, O=-2, S=?

HSO-3 = -1

+1 +S + 3x(-2) = -1

+1 + S - 6 = -1

S = +4

Therefore, the oxidation state of sulphur in HSO- is +4

3

14.4 RADICALS

A radical is a group of elements which acts like a single atom in forming

compounds. Radicals do not have independent existence.

Most radicals are negatively charged. However, a few like ammonium are

positively charged.

Consider the table below which shows the radicals and their names, symbol,

oxidation state, valency and their compounds.

Table 14.3; the common radicals.

Oxidation

Example of compound in

Name of radical

Symbol

Valency

state

which it is found.

Ammonium

NH +

+1

1

Ammonium chloride NH Cl

4

4

Sulphate

SO2-

-2

2

Calcium sulphate CaSO

4

4

Carbonate

CO 2-

-2

2

Zinc carbonate ZnCO

3

3

Hydrogen

HCO -

-1

1

Sodium hydrogen carbonate

carbonate

3

NaHCO3

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Oxidation

Example of compound in

Name of radical

Symbol

Valency

state

which it is found.

Hydrogesulphate

HSO -

-1

1

Potassium hydrogen sulphate

4

KHSO4

Nitrate

NO-

-1

1

Silver nitrate AgNO

3

3

Hydroxide

OH-

-1

1

Potassium hydroxide KOH

Phosphate(V)

PO3-

-3

3

Phosphoric (V) acid H PO

4

3

4

Chlorate(V)

ClO-

-1

1

Sodium chlorate(V) NaClO3

3

Chlorate(I)

ClO-

-1

1

Potassium chlorate(I) KClO

Sulphite(IV)

SO2-

-2

2

Sulphuric (IV) acid H SO

3

2

3

14.5 IONS

Metal atoms lose electrons to become ions. The atom will lose as many

electrons as its valency. Metal ions are always positively charged.

The general equation for the loss of electrons by metal atoms is:M→Mn++ ne

Where;

M = Metal atom

Mn+ = Metal ion which carries a charge of n+

n = number of electrons lost

e = the negatively charged electron.

Non – metal atoms gain electons to become ions. The atom will gain as many

electrons as its valency. Ions of non – metals are always negatively charged.

The general equation for the gain of electrons by a non – metal atom is;

X + n e

Xn-

Where by;

X = non - metal atom

Xn- = non - metal ion which has negative charge.

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Consider the below table showing ions of common metals and non – metals.

Table 14.4 Ions of common metals and non metals.

Ionic name Symbol Charge

Valency Example of a compound

Hydrogen

H+

+1

1

Hydrogen chloride, HCl

Potassium

K+

+1

1

Potassium bromide, KBr

Sodium

Na+

+1

1

Sodium chloride, NaCl

Copper (I)

Cu+

+1

1

Copper (I) oxide, Cu O

2

Nickel

Ni2+

+2

2

Nickel sulphate, NiSO4

Silver

Ag+

+1

1

Silver chloride, AgCl

Copper (II)

Cu2+

+2

2

Copper (II) sulphate, CuSO4

Barium

Ba2+

+2

2

Barium sulphate, BaSO4

Calcium

Ca2+

+2

2

Calcium nitrate, Ca(NO )32

Magnesium Mg2+

+2

2

Magnesium oxide, MgO

Zinc

Zn2+

+2

2

Zinc sulphate, ZnSO4

Iron

Fe2+

+2

2

Iron (II) chloride, FeCl2

Lead

Pb2+

+2

2

Lead (II) sulphate PbSO4

Mercury (II) Hg2+

+2

2

Mercury (II) chloride, HgCl2

Aluminium Al3+

+3

3

Aluminium oxide, Al O

2 3

Ammonium NH+

+1

1

Ammonium nitrate, NH NO

4

4

3

Chloride

Cl-

-1

1

Aluminium chloride AlCl3

Bromide

Br-

-1

1

Potassium bromide KBr

Iodide

I-

-1

1

Silver iodide AgI

Oxide

O2-

-2

2

Magnesium oxide MgO

Sulphide

S2-

-2

2

Lead (II) sulphide, PbS

Hydroxide

OH-

-1

1

Ammonium hydroxide, NH OH

4

Nitrate

NO-

-1

1

Silver nitrate, AgNO

3

3

Carbonate

CO2-

-2

2

Calcium carbonate, CaCO3

3

Sulphate

SO2-

-2

2

Iron (II) sulphate, FeSO

4

4

Hydrogen

HCO-

-1

1

Sodium hydrogen carbonate

carbonate

4

NaHCO3

Hydrogen

HSO-

-1

1

Potassium hydrogen sulphate

sulphate

4

KHSO4

Sulphite

SO2-

-2

2

Sodium sulphite Na SO

3

2

3

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14.6 WRITING A CHEMICAL FORMULAE

The points to remember when writing chemical formulae are explained as

follows.

1. Positively charged ions are written before the negatively charged ions.

2. A radical must be treated as a unit.

3. The name of the cation is the same as the neutral element from which it is

derived (e.g. Na for sodium)

4. Single elements are not bracketed.

5. The valency of 1 is simply assumed and not written in the formula.

Consider the below examples.

Examples 14.7

Write the formula of Copper (II) chloride.

Step 1. Write down the name: Copper (II) chloride.

Step 2. Write down the ions in the compound, complete with their charges.

Cu2+Cl-

Step 3. Draw boxes round the ions and draw arrows to exchange the valencies.

Cu2+

Cl-

Cu

2Cl

Step 4. Write the exchanged valencies as subscripts. A valency of 1 should not

be written. In this step, radicals must be bracketed if necessary. The

complete formula is CuCl . The compound contains two chloride ions

2

for every one Copper (II) ion.

Example 14.8

Write the formula of aluminium Sulphate.

Step 1. Aluminium sulphate

Step 2. Al3+

SO 2-

4

2Al3+

3SO 2-

Step 3.

4

Step 4. Al (SO )

2

4 3

The complete formula is Al (SO )

2

4 3

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Example 14.9

Write formula for calcium oxide.

Step 1: calcium oxide

Step 2: Ca2+

O2-

2Ca2+

2O2-

Step 3:

Step 4: CaO

The complete formula is CaO.

Example 14.10

Write the formula for zinc chloride

Step 1: Zinc chloride

Step 2: Zn2+

Cl-

1Zn

2Cl-

Step 4: ZnCl2

The complete formula is ZnCl .2

Activity 14.1

Write the chemical formula of the compounds formed when the following combine.

(a). Sodium ion and carbonate ion.

(b). Sodium ion and cyanide ion.

(c). Ammonium ion and sulphate ion.

(d). Hydrogen ion and chloride ion.

(e). Copper ion and nitrate ion.

(f). Magnesium ion and chloride ion.

(g). Write the names of each of the following compounds.

(i). CuCl 2

(ii). Na SO

2

4

(iii). NH NO .

4

3

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14.7 EMPIRICAL AND MOLECULAR FORMULAE

Chemical formula can basically be divided into three types, namely empirical

formula, molecular formula and structural formula. For this level we shall

concentrate on the empirical and molecular formulae.

Empirical formula

The empirical formula of a compound is the simplest formula which expresses

its composition by mass.

Molecular formula.

The molecular formula is the formula which shows the actual number of each

kind of atom present in a molecule. The molecular formula is equal to n(empirical

formula), where n is a whole number.

CALCULATING THE EMPIRICAL AND MOLECULAR

FORMULAE.

Example 14.11

Sodium sulphate has the following composition by mass; Sodium 32.4%,

Sulphur 22.5% and Oxygen 45.1%. What is its empirical formula?

Solution

Elements

Na

S

O

% Composition

32.4

22.5

45.1

Divide by relative Atomic mass

32.4

23 = 1.14 22.5

32 = 0.70 45.1

16 = 2.82

Divide by smallest number

1.14

0.70 = 2

0.70

0.70 = 1

2.82

0.70 = 4

2

1

4

Therefore the empirical formula is Na SO

2

4

Example 14.12

What is the empirical formula for a compound of mass 8.1g if it consists of 4.9 g

of magnesium and 3.2g of oxygen?

Solution

Elements

Mg

O

Relative Masses

4.9

3.2

Divide by relative Atomic mass

4.9

24 = 0.20 3.2

16 = 0.20

Divide by smallest number

0.2

0.2 = 1

0.2 = 1

0.2

1

1

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Therefore the empirical formula is MgO

Example 14.13

An experiment showed that 13.88g of calcium chloride were obtained from

the combination of 5g of calcium with unknown relative mass of chlorine. What is

the formula of calcium chloride?

Solution.

Mass of calcium chloride = 13.88g

Mass of calcium = 5.00g

Mass of chlorine =?

Mass of chlorine = mass of calcium chloride – mass of calcium.

Mass of chlorine = (13.88 – 5.00)g

Hence the mass of chlorine = 8.88g.

Elements

Ca

Cl

Relative masses

5.00

8.88

Divide by relative Atomic mass

5.00

40 = 0.125 8.88

35.5 = 0.25

Divide by smallest number

0.125

0.125 = 1

0.25

0.125 = 2

1

2

Therefore the empirical formula is CaCl2

Example 14.14.

A compound contains 15.8% carbon and 84.2% sulphur. Calculate its empirical

formula. If its relative molecular mass is 76, what is its molecular formula?

Solution.

Elements

Carbon (C)

Sulphur (S)

% Composition

15.8

84.2

Divide by relative Atomic mass

15.8

12 = 1.32 84.2

32 = 2.63

Divide by smallest number

1.32

1.32 = 1

2.63

1.32 = 2

1

2

Therefore the empirical formular is CS .2

The molecular formular is given by;

Molecular formulae = n (empirical formula)

Where n is a whole number.

Then; Molecular formula = n (CS )2

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Now Relative molecular Mass = n (sum of relative atomic mass)

76 = n (12+ (2x32))

76 = n (12 + 64)

76 = 76n

n = 1

Hence the molecular formula is CS .2

Example 14.15

A certain gaseous compound contains 82.8% of carbon and 17.2% of hydrogen by

mass. The vapour density of the compound is 29. Calculate its molecular formula.

Elements

C

H

% Composition

82.8

17.2

Divide by relative Atomic mass

82.8

12 = 6.90 17.2

1

= 17.2

Divide by smallest number

6.90

6.90 = 1

17.2

6.90 = 2.5

Multiply by 2 to get whole number only

1 x 2 = 2

2.5 x 2 = 5

Empirical formula is C H

2 5

Relative molecular mass of gas = 2 x Vapour density.

= 2 x 29 = 58

Molecular formula is given by:-

(Empirical

formula)

n

=

relative

molecular

mass

(C H ) n = 58

2 5

[(12x2)+ (1x5)] n = 58

(24 + 5) n = 58

29n = 58

29 29

n = 2

(C H ) = C H

2 5 2

4 10

The molecular formula of the gas is C H

4 10

Example 14.16

A certain compound contains 22.7% of zinc, 11.0% of sulphur, 22.3% of

oxygen and the rest being water of crystallization. Calculate

(i) Empirical formula.

(ii) Molecular formula of a compound if its molar mass is 287g.

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Solution;

Elements

Zn

S

O

H O

2

% Composition

22.7

11.0

22.3

44

Divide by relative Atomic

22.7

mass

63 = 0.36 11.0

32 = 0.34 22.3

16 = 1.39 44

18

= 2.44

Divide by smallest number

0.36

0.34 = 1

0.34

0.34 = 1

1.39

0.34 =4

2.44

0.34 = 7

1

1

4

7

The empirical formula is ZnSO ●7H O

4

2

(ii) Given molar mass 287g

(Empirical formula) n = molar mass

[ ( 65 + 32 + ( 16x4) + ( 7x18) ]n = 287

287n = 287

n = 1

The molecular formula is ZnSO ●7H O

4

2

14.8 BONDING

Bonding is the process of joining two or more things or substances together.

A link that holds two or more things together is called a bond.

Chemical substances are made of atoms that are held together by chemical

bonds.A chemical bond is a force of attraction that holds atoms together to form

molecules.

CHEMICAL BONDING

Chemical bonding is a process by which atoms become more stable by

donating, gaining or sharing electrons. Chemical bonding involves forces of

attraction between electrons in the outermost energy levels of atoms.

When the outermost levels are filled, the atoms are said to be stable. The stable

atoms are generally unreactive and can exist freely as single atoms. Unstable atoms

cannot exist freely as single toms. For unstable atoms to become stable, they should

acquire electronic arrangements similar to those of noble gas. This means that they

either lose, gain or share electrons.

Types of chemical bonding;

➢ There are two main types of chemical bonds, these are ionic and covalent

bonds. Ionic bond involves the transfer of electron from a metal atom to the

non-metal atom. Covalent bond involves the sharing of a pair of electrons in

which, each bonding atom donates sharing electrons.

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ELECTROVALENT BONDING (IONIC BONDING)

An electrovalent bond is the bond due to the transfer ( gain or loss) of one or more electrons from one atom or radical to another. The electrovalence of an

atom or a group of atoms is the number of electrons which it transfers or receives.

Electrovalent bonding usually occurs between a metal and a non-metal.The metal

loses electrons and the non-metal gains electrons.

Example; when sodium and chlorine react to form sodium chloride, the sodium atoms must each lose an electron to acquire a stable neon structure, while chlorine

atoms must each gain an electron to acquire a stable argon structure, this results

in positively charged sodium ions and negatively charged chloride ions to form

sodium chloride molecules.

Consider the following “cross” and “dot” diagrams in which a cross (x) represents the electrons in a non- metal atom, the dot (•) represents the electron

transferred from a metal atom.

11p

17p

12n

18n

Na

Cl

17p

11p

18n

12n

Na+

Cl ¯

Figure 14.1 electrovalent bond in sodium chloride.

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20p

8p

20n

8n

O

Ca

8p

20p

8n

20n

O2-

Ca2+

Figure 14.2; electrovalent bonding in calcium oxide.

Properties of electrovalent compounds

(i). They are generally soluble in water

(ii). Conduct electricity when in solution or molten form, but not when in solid

form

(iii). They are usually crystalline solids at room temperature.

(iv). They have high melting and boiling points.

(v). They are generally insoluble in non-polar solvents.

(vi). They consist of ions that are positively and negatively charged.

Task 14.1

1. Cations are slightly smaller than their neutral atoms, where as anions are slightly larger than their neutral atoms. Explain why.

2. Briefly explain how magnesium and oxygen atoms combine to form magnesium

oxide molecule. In your answer highlight

(a). the charge on a magnesium ion (b) The charge on an oxide ion

3. Illustrate using a dot and cross diagram the electron transfer involved in forming

potassium chloride bond.

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COVALENT BONDING

A covalent bonding is a type of bonding which involves the sharing of electrons between atoms. In covalent bonding, the atoms share electrons and the atoms within

each molecule are held together by the attraction between their positive nuclei and

these shared electrons.

A molecule is thus formed in each atom which has attained an energetically

stable arrangement of electrons. Covalent bond occurs when non-metal atoms

combine with other non-metal atoms. Molecules of covalent compounds are discrete

or distinct. It means that the atoms forming the molecules cannot exist freely in the

compound. They remain bound together in molecules and their electrons are not

free to form bonds with other atoms.

The following examples of covalent bond formation are concerned.

1. Hydrogen atom needs one electron to acquire the helium structure. It can

combine with another hydrogen atom and share the electrons so that each

acquires a stable helium structure. This results in the formation of a hydrogen

molecule. The kind of bond formed between the two atoms is called a covalent

bond.

Shared electrons

1p

1p

1p

1p

or H● + H● → H : H or H - H

Figure 14.3 covalent bonding in hydrogen molecule.

2. Fluorine molecule is formed by means of establishing a covalent bond between

two fluorine atoms. Each of the fluorine atoms donates single electrons from

its outer most shell.

F=2:7

F=2:7

F=2:8

F=2:8

Shared electrons

Figure 14.4 covalent bonding in fluorine molecule.

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In fluorine molecule, each atom has 6 electrons of its own in the outer most orbits plus the shared pair, which revolves around both atoms. i.e. They have a

complete outer octet and the stable neon structure (2:8).

3. Oxygen requires two electrons to acquire the stable atom structure. Therefore

two atoms of oxygen combine and share four electrons between them.

two atoms of oxygen

Oxygen molecule

Or

or O = O

Figure 14.5 covalent bonding in oxygen. (double covalent bond).

4. In the water molecule (H O) each of the unpaired electrons in the outer most

2

shell of oxygen atom is used for covalent bonding with the electrons of the

two hydrogen atoms. Therefore, there are two single covalent bonds in the

water molecule.

O=2:6

H=1

Then, two hydrogen atoms combine with one atom of oxygen to form water

molecule.

O

H

H

Shared electrons

Fig 14.6 covalent bonding in water.

Both oxygen and hydrogen attain a noble gas configuration. O2- = 2:8 and

H- = 2

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5. Consider a triple covalency such as nitrogen gas (N ). It is a diatomic. A triple

2

bond is formed between the two atoms. Electronic configuration for nitrogen

is 2:5

N

N

Na = 2:5

N = 2:5

N

N

N = 2:8

N = 2:8

Shared electrons

N molecule or N N

2

Fig. 14.7 Covalent bonding in nitrogen

Properties of covalent compounds

(i). The melting and boiling points are usually low

(ii). They are usually liquids or gases at room temperature.

(iii). They are generally insoluble in water but soluble in organic solvents such as

ether, ethanol or benzene.

(iv). They do not conduct electricity when solid, liquid or in solution because they

do not have ions.

(v). They consist of molecules with two or more different atoms linked together

by covalent bonds.

Experiment 14.1

Aim; to find out the electrical conductivity of covalent and electrovalent

compounds.

Procedures;

1. Use crucible or beaker to hold the solution which is to be tested.

2. Arrange the electric circuit as in figure 14.8 . A 6-volt battery or any other 6-volt direct current (d.c) supply is connected by a copper wire to a switch and either

a 6-volt bulb in a bulb holder or an ammeter. You can include a variable resistor

if necessary. Two carbon rods take the current into and out of the solution being

tested.

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3. Fill the container with one solution under test. Put on the switch to complete

the circuit. If the bulb glows or the ammeter shows a reading, the solution is a

conductor of electric current. Switch off at once.

4. Use the following solutions; diesel, ethanol, dilute hydrochloric acid, sodium

hydroxide, sodium chloride, sugar solution and dilute sulphuric acid and

clean the two carbon rods after each experiment. If we are to investigate the

conduction of electricity by molten substances we should use the same circuit

as that used for investigation of the conduction of electricity by a variety of

substances in the solid and liquid states. The substances to be used for testing

are such as wax, sugar, lead (II) bromide, potassium iodide, lead (II) iodide.

Switch

6-volt battery

A Ammeter (or bulb)

Crocodile clip

Wooden holder

Crucible

Carbon rods

Solution on test

Figure 14.8; to find out which solids and solutions conduct electricity.

From your results of the experiment, divide the solutions into two groups:

(a). The solutions which conduct electric current and cause light in the bulb

or deflection of the ammeter.

(b). Solutions which neither conduct electric current nor cause the bulb

to glow or no any ammeter deflection. Which of these solutions are

covalent compounds and which are electrovalent compounds?

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Table 14.5. Differences between covalent and electrovalent compounds.

Covalent compounds

Electrovalent compounds.

They are formed by mutual sharing

They are formed by complete transfer of

of electrons

electrons

They are made up of molecules

They are made up of ions

They are usual y liquids or gases

They are usual y crystalline solids

They are soluble in non-polar solvent They are usual y soluble in water but insoluble

like benzene or carbon tetrachloride in non-polar solvents like carbon tetrachloride

and insoluble in polar solvents like

water

They general y have low melting and They have general y high melting and boiling

boiling points

points

They do not conduct electricity.

They are good conductors of electricity in

the molten state and in aqueous solutions but

insulators in the solid state.

14.9 NOMENCLATURE OF BINARY INORGANIC

COMPOUNDS

Nomenclature is a systematic way of assigning names to items or substances that belong to a particular group or classification.

A binary compound is a chemical compound that contains exactly two

different elements. Inorganic compounds fall mainly into two main categories, namely ionic and covalent.

BINARY IONIC COMPOUNDS

A binary ionic compound is a salt consisting of only two elements in which both elements are ions, a cation (which has a positive charge ) and anion (which

has a negative charge).

The following are steps to be considered when naming binary ionic compounds:

(1). Name the metallic ion that appears first in the formula using the name of the

element itself.

(2). The second part of the formula, which is usually an anion in the compound,

will end in “ide”. Example, chlorine become chloride, fluorine become fluoride

and hydroxyl become hydroxide.

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Example 14.17

What is the name of the compound of formula CuS?

Solution;

Step 1.

Let x be the valence of a Cu atom.

Step 2. Sulphur has a -2 charge

Step 3. 1(x) +1(-2) = 0

x = +2 for Cu

Step 4. Write the name copper and place (II) in brackets beside it.

Step 5. Use the name sulphur but change the last two letters to “ide”.

Therefore the compound is copper (II) sulphide.

Example 14.18

What is the name of the compound with formula FeCl ?

3

Solution;

The total charge of the molecule is zero and Cl- has a negative charge.

Then;

Step 1. let x be the valence of the Fe atom

Step 2. 1(x) + 3(-1) = 0

Step 3. x = +3

Step 4. So the Fe is in the +3 oxidation state.

Write the name “iron” and place (III) in brackets beside it.

Step 5. Use the name chlorine but change the last three digits to “ide”.

Hence the name is iron (III) chloride.

More examples of binary ionic compounds are;

CaO is named Calcium oxide

AlN is named aluminium nitride

K S is named potassium sulphide

2

MnO is named manganese (IV) oxide.

2

Note; Manganese can have more than one charge, but each oxide ion has a charge of -2. In order for the compound to be neutral, Mn must have a +4 charge.

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BINARY COVALENT COMPOUNDS.

A binary covalent compound is composed of two different non-metal elements.

Example, a molecule chlorine trifluoride (ClF ) contains 1 atom of chlorine and

3

three atoms of fluorine. A set of prefixes is used to identify the subscript in the

binary covalent compound formula.

Consider the table below: Table 14.6. List of prefixes

Number

prefix

1

mono

2

di

3

tri

4

tetra

5

penta

6

hexa

7

hepta

8

Octa-

9

Nona-

10

Deca-

The steps to be considered when naming binary covalent compounds are;

1. Give the name of the first element

2. Give the name of the second element with the ending changed to –ide.

3. If more than one compound is possible between the two elements, give

prefixes to indicate the number of atoms of each element.

Example 14.19

Name the compound N O .

2 4

(i). Use the prefix “ide” in front of nitrogen

(ii). Use the prefix “tetra” in front of the oxygen.

(iii). We drop “-ygen” and replace with “ide”.

Therefore the name is dinitrogen tetraoxide

Example 14.20

What is the name for PCl ?

3

(i). Since there is one phosphorus atom, we use it as the first part of the name.

(ii). There are three chlorine atoms, so we use “tri” in front of chlorine. We then

drop the “-ine” in chlorine and replace with “ide”.

Therefore the name is phosphorus trichloride.

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Consider the table below which gives the formulae and names of some binary

covalent compounds.

Table 14.7 some binary covalent compounds

Formula

Name

CO

Carbon dioxide

2

CO

Carbon monoxide

P O

Diphosphorus pentoxide

2 5

N O

Dinitrogen monoxide

2

SiO

Silicon dioxide

2

CBr

Carbon tetrabromide

4

SO

Sulphur trioxide

3

PBr

Phosphorus pentabromide

5

ICl

Iodine trichloride

3

NI

Nitrogen triiodide

3

N O

Dinitrogen trioxide

2 3

N O

Dinitrogen tetraoxide

2 4

SO

Sulphur dioxide

2

NO

Nitrogen monoxide

NO

Nitrogen dioxide

2

As O

Diarsenic pentoxide

2 5

PCl

Phosphorus trichloride

3

CCl

Carbon tetrachloride

4

H O

Dihydrogen monoxide (most call it water)

2

SF

Sulphur hexafluoride

6

XeO

Xenon trioxide

3

P S

Tetraphosphorus trisulphide

4 3

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CHEMICAL NAMES OF COMMON SUBSTANCES

Chemical names are usually used to give an accurate description of the

composition of a substance.

The table below gives a list of common names for some chemicals.

Table 14.8 chemical names of common substances

Common name

Chemical name

Common salt/table salt

Sodium chloride (NaCl)

Washing soda/soda ash

Sodium carbonate (Na CO )

2

3

Limestone/marble/chalk Calcium carbonate (CaCO )3

Quick lime /lime

Calcium oxide (CaO)

Lime water

Calcium hydroxide Ca(OH) (aq) solution

2

Slaked lime

Calcium hydroxide (Ca(OH ))

2

Salt petre

Potassium nitrate (KNO )3

Plaster of Paris (POP)

Calcium sulphate -1-water (CaSO •H O)

4

2

Gypsum

Calcium sulphate-2-water (CaSO •2H O)

4

2

Caustic soda

Sodium hydroxide (NaOH)

Caustic potash

Potassium hydroxide (KOH)

Brine

Concentrated sodium chloride solution ((NaCl(aq))

Iodine tincture

A mixture of iodine and sodium iodide dissolved in a 50-50

solution of ethanol and water.

Silica

Silicon (IV) oxide (SiO )2

Blue copper

Copper sulphate-5-water ( CuSO •5H O)

4

2

Rock salt

Natural y occurring sodium chloride (NaCl impure)

Red oxide

Lead tetraoxide (Pb O )

3 4

Chile salt petre

Sodium nitrate (NaNO )3

Sugar

Sucrose (C H O )

12 22 11

Aspirin

Acetyl/salicyclic acid (C H O )

9 8 4

Baking soda

Sodium bicarbonate (NaHCO )3

Vitamin C

Ascorbic acid (C H O )

6 8 6

Asbestos

Magnesium silicate (MgSiO )3

Fluorspar (fluorite)

Calcium fluoride (CaF )2

Soda lime

Mixture of sodium and calcium hydroxide (Ca(OH) + NaOH)

2

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SUMMARY

(a). Valency is the combining capacity of an element.

(b). Oxidation state of an element indicates the number of electrons lost, gained

or shared by an atom of the element with respect to its neutral atom.

(c). The neutral atom has no charge.

(d). A radical is a group of elements which act like a single atom in forming

compounds.

(e). Chemical bonding is a method by which atoms become more stable by

donating, gaining or sharing electrons.

(f). The two main types of chemical bonds are covalent bond and electrovalent

bond.

(g). A covalent bond involves sharing of electrons.

(h). An electrovalent bond involves transfer of electrons.

(i). The empirical formula of a compound is the simplest formula which expresses

its composition by mass.

(j). Molecular formula is the formula which shows the actual number of each kind

of atom present in a molecule.

(k). A binary compound is a chemical compound that contains exactly two

different elements.

REVIEW QUESTIONS

1. Choose the most correct answer for the following questions;

(i). .............. is a method of representing a molecule of a compound.

(A) Valency

(B) Element

(C) A formula

(D) Oxidation

(ii). The combining capacity of an element is called; ...............................

(A) Oxidation state

(B) Valency

(C) Mole

(D) Ion

(iii). A group of elements which acts like a single atom in forming compounds is

called;.................................

(A) Molecule

(B) Compound

(C) Radical

(D) Formula.

(iv). The formula of copper (II) chloride is..............

(A) CuCl

(B) Cu Cl

(C) Cu Cl

(D) CuCl

2

2

2

2

(v). The system of naming compounds is also known as; ...........................

(A) Nomenclature

(B) Naming methodology

(C) Classification

(D) Binary compound

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2. Match the items in LIST A with their correct statement in LIST B

List A

List B

(i). Oxidation state ________

A. Sharing of electrons

(ii). Covalent bond _______

B. Transfer of electrons

(iii). Electrovalent bond_____

C. A chemical compound that contains

(iv). Binary compound _____

exactly two different elements.

(v). Empirical formula_______ D. The simplest formula which expresses its

composition by mass

E. The formula which shows the actual

number of each kind of atom present in

one molecule of a compound.

F. The number of electrons lost, gained or

shared by an atom

3. Write TRUE for a correct statement and FALSE for an incorrect statement.

(i). The melting and boiling points of covalent compounds are usually low.............

(ii). Electrovalent compounds are usually liquids or gases at room temperature

(iii). Electrovalent compounds are generally crystalline solids at room

temperature....................

(iv). Electrovalent compounds have high melting and boiling points ..........

(v). Electrovalent compounds are generally soluble in water ...................

(vi). Covalent compounds consist of ions that are positively and negatively

charged...................

(vii). Covalent compounds are usually crystalline solids at room temperature.......

4. Name the elements which have these symbols.

K, Na, Ag, Pb, Zn, Ba, Sn, Ar, P and Hg

5. Calculate the oxidation number of the underlined elements in the following

compounds or radicals;

a. Na PO

3

4

b. SO 2-

3

c. NO -2

6. Write the formulae of the following compounds;

(a) Iron (III) chloride

(b) Lead (II) oxide

(c) Potassium nitrate

(d) Barium hydroxide

(e) Potassium hydrogen carbonate

7. Here is a list of formulae for some iron compounds;

Fe O , FeSO , FeCl , Fe (SO )

2 3

4

3

2

4 3

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(a) Which compounds are binary compounds

(b)Which compound contains iron in oxidation state +2?

8. From the following list, O , Al, Na+, CH , I-, N select;

2

4

(a) Two atoms

(b) two molecules

(c) two ions.

9. Name the following compounds;

(a) MgI

(b) N O

(c) CCl

2

2 5

4

(d) FeBr

(e) Cu P

(f) H S

2

3 2

2

(g) Hg Cl

(h) K O

(i)

PCl

2

2

2

5

(j) SF6

10. Explain the differences between the properties of electrovalent and covalent

compounds using sodium chloride and water as examples.

11. Copy and complete this table;

Formula

Systematic (IUPAC) name

Cu O

2

FeS

Pb (NO )32

Fe (SO )

2

4 3

KClO3

Cu(NO )32

MnO2

200

Questions for

Revision and Practice

1. (i) Mention at least ten products that can be produced through application of

chemistry knowledge

(ii) Explain how the knowledge of chemistry is used to produce the named

products in (i) above.

2. Explain how the following professionals can apply the knowledge of chemistry in

their respective fields: Geologist, nurse, medical doctor, engineer, science teacher,

laboratory technician, pharmacist, driver, farmer, horticulturists, floriculturists,

and housewife.

3. (i) List at least ten examples of laboratory apparatus.

(ii) Name the nature and use of each apparatus you have listed in (i) above

(iii) Draw all the apparatus you have listed in (i) above

4. Explain how you could go about when rendering first aid to accidents of bruise,

burns, fainting, shock, suffocation, choking, electric shock, severe bleeding, light

bleeding and vomiting respectively

5. (i) Define first aid and first aid kit respectively..

(ii) Mention at least eleven items found in a first aid kit.

(iii) Explain the uses of items you have mentioned in (ii) above.

6. Draw chemical warning signs expected to be seen on the following items:

(i) A can of petroleum

(ii) A bottle of concentrated sulphuric acid

(iii) A bottle of zinc granules

(iv) A bottle of sodium lump

(v) A bottle of uranium metal

(vi) A container of mercury metal

(vii) A container with hydrogen gas

7. (i) Define heat and flame respectively

(ii) Draw well labeled diagrams of luminous and non-luminous flames respectively

(iii) Outline five(5) differences between luminous and non-luminous flames

(iv) Juma is working in a welding industry in Dar es Salaam. What type of flame

is he supposed to use in his work?

(v) Mwajuma is in a sitting room at home when the electric supply is cut off

suddenly. Advise her on what type of flame she might use to il uminate the

sitting room.

8. Explain how you can produce the hottest flame using a Bunsen burner

9.(i) What do you understand by control experiment?

(ii) Design an experiment and explain how you would use scientific procedure to

go about the experiment.

201

10. (i) Outline five (5) differences between Physical and Chemical changes.

(ii) In tabular form, name at least ten (10) examples of physical and chemical

changes respectively.

11. Define:

(i) Brownian motion

(ii) Particulate nature of matter

(iii)

Diffusion

(iv)

Osmosis

(v) Kinetic molecular behavior

12. Define:

(i) Element

(ii) Compound

(iii) Mixture

(iv) Molecule

(v) Homogenous mixture (vi) Heterogeneous mixture

13. Name and define three types of solutions as far as saturation is concerned.

14. (i) Define solutions, suspensions and emulsions respectively.

(ii) Mention at least two examples on the terms you have defined in (i) above

15. Name the method you would use to separate the following mixtures:

(i) Muddy water

(ii) Mixture of salt and water

(iii) Mixture of sand and water

(iv) Mixture of rice and water

(v) Mixture of ethanol and water

(vi) Kerosene and water

(vii) Oil in ground nuts

(viii)Tea leaves and water

(ix) Mixture of dyes

(x) Protein and blood

16. (i) Mention and explain three components of a fire triangle.

(ii) With the aid of a chemical equation, define Rust and Rusting respectively.

(iii) Explain how you would go about in using a portable fire extinguisher in your

school laboratory in case of fire outbreak.

(iv) Mention and explain at least five (5) methods of preventing rusting

17. Explain how Oxygen is important to the following people: housewife, mountain

climber, deep sea diver, rocket scientist, premature baby, factory worker and

laboratory technician.

18. Maduhu is working in a multi-product industry in Tanga. Explain how he will

make use of hydrogen gas to produce different products.

19. (a) Define: (i) Water

(ii) Water cycle

(iii) Water treatment (iv) Water purification

(b) (i) Explain how you can use water in day-to-day activities

(ii) Explain how you can purify domestic as well as you can treat urban water.

Use one page narration.

20. Define Isotope and Isotopy respectively.

21. What do you understand by periodic table?

22. Define:

(i) Bonding

(ii) Chemical formula

(iii) Nomenclature

(iv) Binary compound

202

Appendix A; Periodic table

203

Appendix B. A table of elements

Atomic

M.P

B.P

Discovery

Ionization

Name

Symbol

Density

number

(°C)

(°C)

(Year)

energy (eV)

Hydrogen

H

1

-259

-253

0.09 (g/l)

1766

13.5984

Helium

He

2

-272

-269

0.18 (g/l)

1895

24.5874

Lithium

Li

3

180

1347

0.53 (g/cm3)

1817

5.3917

Beryllium

Be

4

1287

2469

1.85 (g/cm3)

1797

9.3227

Boron

B

5

2300

2550

2.46 (g/cm3)

1808

8.298

Carbon

C

6

3500

4827

2.26 (g/cm3)

Ancient

11.2603

Nitrogen

N

7

-210

-196

1.25 (g/l)

1772

14.5341

Oxygen

O

8

-218

-183

1.43 (g/l)

1774

13.6181

Fluorine

F

9

-220

-188

1.7 (g/l)

1886

17.4228

Neon

Ne

10

-249

-246

0.9 (g/l)

1898

21.5645

Sodium

Na

11

98

883

0.97 (g/cm3)

1807

5.1391

Magnesium

Mg

12

649

1090

1.74 (g/cm3)

1755

7.6462

Aluminium

Al

13

660

2467

2.7 (g/cm3)

1825

5.9858

Silicon

Si

14

1410

2355

2.33 (g/cm3)

1824

8.1517

Phosphorus

P

15

44

280

1.82 (g/cm3)

1669

10.4867

Sulphur

S

16

113

445

1.96 (g/cm3)

Ancient

10.36

Chlorine

Cl

17

-101

-35

3.21 (g/l)

1774

12.9676

Argon

Ar

18

-189

-186

1.78 (g/l)

1894

15.7596

Potassium

K

19

64

774

0.86 (g/cm3)

1807

4.3407

Calcium

Ca

20

839

1484

1.55 (g/cm3)

1808

6.1132

Scandium

Sc

21

1539

2832

2.99 (g/cm3)

1879

6.5615

Titanium

Ti

22

1660

3287

4.54 (g/cm3)

1791

6.8281

Vanadium

V

23

1890

3380

6.11 (g/cm3)

1830

6.7462

Chromium

Cr

24

1857

2672

7.19 (g/cm3)

1797

6.7665

Manganese

Mn

25

1245

1962

7.43 (g/cm3)

1774

7.434

Iron

Fe

26

1535

2750

7.87 (g/cm3)

Ancient

7.9024

Cobalt

Co

27

1495

2870

8.9 (g/cm3)

1735

7.881

Nickel

Ni

28

1453

2732

8.9 (g/cm3)

1751

7.6398

Copper

Cu

29

1083

2567

8.96 (g/cm3)

Ancient

7.7264

Zinc

Zn

30

420

907

7.13 (g/cm3)

Ancient

9.3942

Gallium

Ga

31

30

2403

5.91 (g/cm3)

1875

5.9993

Germanium

Ge

32

937

2830

5.32 (g/cm3)

1886

7.8994

Arsenic

As

33

81

613

5.72 (g/cm3)

Ancient

9.7886

Selenium

Se

34

217

685

4.79 (g/cm3)

1817

9.7524

Bromine

Br

35

-7

59

3.12 (g/cm3)

1826

11.8138

Krypton

Kr

36

-157

-153

3.75 (g/l)

1898

13.9996

Rubidium

Rb

37

39

688

1.53 (g/cm3)

1861

4.1771

Strontium

Sr

38

769

1382

2.63 (g/cm3)

1790

5.6949

Yitrium

Y

39

1523

3337

4.47 (g/cm3)

1794

6.2173

204

Atomic

M.P

B.P

Discovery

Ionization

Name

Symbol

Density

number

(°C)

(°C)

(Year)

energy (eV)

Zirconium

Zr

40

1852

4377

6.51 (g/cm3)

1789

6.6339

Niobium

Nb

41

2468

4927

8.57 (g/cm3)

1801

6.7589

Molybdenum

Mo

42

2617

4612

10.22 (g/cm3) 1781

7.0924

Technetium

Tc

43

2200

4877

11.5 (g/cm3)

1937

7.28

Ruthenium

Ru

44

2250

3900

12.37 (g/cm3) 1844

7.3605

Rhodium

Rh

45

1966

3727

12.45 (g/cm3) 1803

7.4589

Palladium

Pd

46

1552

2927

12.02 (g/cm3) 1803

8.3369

Silver

Ag

47

962

2212

10.5 (g/cm3)

Ancient

7.5762

Cadmium

Cd

48

321

765

8.65 (g/cm3)

1817

8.9938

Indium

In

49

157

2000

7.31 (g/cm3)

1863

5.7864

Tin

Sn

50

232

2270

7.31 (g/cm3)

Ancient

7.3439

Antimony

Sb

51

630

1750

6.68 (g/cm3)

Ancient

8.6084

Tellurium

Te

52

449

990

6.24 (g/cm3)

1983

9.0096

Iodine

I

53

114

184

4.94 (g/cm3)

1811

10.4513

Xenon

Xe

54

-112

-108

5.9 (g/l)

1898

12.1298

Caesium

Cs

55

29

678

1.87 (g/cm3)

1860

3.8939

Barium

Ba

56

725

1897

3.51 (g/cm3)

1808

5.2117

Lanthanum

La

57

920

3469

6.15 (g/cm3)

1939

5.4569

Cerium

Ce

58

795

3257

6.69 (g/cm3)

1803

5.5387

Praseodymium

Pr

59

935

3127

6.64 (g/cm3)

1885

5.473

Neodymium

Nd

60

1010

3127

7.01 (g/cm3)

1885

5.525

Promethium

Pm

61

1100

3000

7.3 (g/cm3)

1945

5.582

Samarium

Sm

62

1072

1900

7.35 (g/cm3)

1879

5.6437

Europium

Eu

63

822

1597

5.24 (g/cm3)

1901

5.6764

Gadolinium

Gd

64

1311

3233

7.9 (g/cm3)

1880

6.1501

Terbium

Tb

65

1360

3041

8.23 (g/cm3)

1843

586.38

Dysprosium

Dy

66

1412

2562

8.55 (g/cm3)

1886

5.9389

Holmium

Ho

67

1470

2720

8.79 (g/cm3)

1867

6.0215

Erbium

Er

68

1522

2510

9.07 (g/cm3)

1842

6.1077

Thulium

Tm

69

1545

1727

9.32 (g/cm3)

1879

6.1843

Ytterbium

Yb

70

824

1466

6.57 (g/cm3)

1878

6.2542

Lutetium

Lu

71

1656

3315

9.84 (g/cm3)

1907

5.4259

Hafnium

Hf

72

2150

5400

13.31 (g/cm3) 1923

6.8251

Tantalum

Ta

73

2996

5425

16.65 (g/cm3) 1802

7.5496

Tungsten

W

74

3410

5660

19.25 (g/cm3) 1783

7.864

Rhenium

Re

75

3180

5627

21.04 (g/cm3) 1925

7.8335

Osmium

Os

76

3045

5027

22.6 (g/cm3)

1803

8.4382

Iridium

Ir

77

2410

4527

22.65 (g/cm3) 1803

8.967

Platinum

Pt

78

1772

3827

21.09 (g/cm3) 1735

8.9587

Gold

Au

79

1064

2807

19.32 (g/cm3) Ancient

9.2255

205

Atomic

M.P

B.P

Discovery

Ionization

Name

Symbol

Density

number

(°C)

(°C)

(Year)

energy (eV)

Mercury

Hg

80

-39

357

13.55 (g/cm3) Ancient

10.4375

Thallium

Ti

81

303

1457

11.85 (g/cm3) 1861

6.1082

Lead

Pb

82

327

1740

11.34 (g/cm3) Ancient

7.4967

Bismuth

Bi

83

271

1560

9.78 (g/cm3)

Ancient

7.2856

Polonium

Po

84

254

962

9.2 (g/cm3)

1898

8.417

Astatine

At

85

302

337

7.0 (g/cm3)

1940

9.3

Radon

Tn

86

-71

-62

9.73 (g/l)

1900

10.7185

Francium

Fr

87

27

677

1.87 (g/cm3)

1939

4.0727

Radium

Ra

88

700

1737

5 (g/cm3)

1898

5.2754

Actinium

Ac

89

1050

3200

10.07 (g/cm3) 1899

5.17

Thorium

Th

90

1750

4790

11.72 (g/cm3) 1829

6.3067

Protactinium

Pa

91

1568

4027

15.4 (g/cm3)

1913

5.89

Uranium

U

92

1132

3818

19.05 (g/cm3) 1789

6.1947

Neptunium

Np

93

640

3902

20.45 (g/cm3) 1940

6.2657

Plutonium

Pu

94

640

3235

19.84 (g/cm3) 1940

6.0262

Americium

Am

95

994

2607

13.67 (g/cm3) 1944

5.9738

Curium

Cm

96

1340

3100

13.5 (g/cm3)

1944

5.9915

Berkelium

Bk

97

986

2900

14.78 (g/cm3) 1949

6.1979

Californium

Cf

98

900

1472

15.1 (g/cm3)

1950

6.2817

Einsteinium

Es

99

860

____

8.84 (g/cm3)

1952

6.42

Fermium

Fm

100

1527

____

____

1952

6.5

Mendelevium

Md

101

827

____

____

1955

6.58

Nobelium

No

102

827

____

____

1958

6.65

Lawrencium

Lw

103

1627

____

____

1961

4.9

Unnilquadium

Unq

104

2100

5500

23 (g/cm3)

1964

5.79

206

APPENDIX C: HAZARD WARNING SYMBOLS

OXIDISING

HARMFUL/IRRITANT

DANGEROUS FOR

THE ENVIRONMENT

GENERAL DANGER

RISK OF FIRE

RISK OF EXPLOSION

TOXIC HAZARD

CORROSIVE SUBSTANCE

RISK OF IONISING

RADIATIONS

RISK OF ELECTRIC

NON-IONISING

SHOCK

LASER BEAM

RADIATION

CHEMICAL WEAPON

BIOHAZARD

NO OPEN FLAME

207

GLOSSARY

1. Acid – a substance that produce

outermost stable energy level.

hydrogen ions (H+) or hydroxonium 20. Balance-Apparatus for weighing

ions (H O+) in solution.

substances

3

2. Acidic oxides – Oxides of non-metals

21. Barrel - the cylindrical tube that is the

which react with water to form acidic

upper part of a Bunsen burner

solution.

22. Base - A metal oxide or hydroxide that

3. Alkali- a soluble base

neutralize an acid to form salt and

4. Alkali earth metals- group II elements

water only.

5. Alkali metals- group I elements

23. Basic oxide-An oxide of metal which

6. Analysis- the process of examining

react with water to form alkaline

data in detail in order to understand it.

solutions.

7. Anhydrous- describe a substance that

24. Beaker-lipped cylindrical glass or

contains no water.

plastic container used for holding,

8. Anion- an ion with a negative charge.

treating and mixing liquids

9. Anode- positively charged electrode.

25. Biogas - Gas produced by the decay of

10. Anthocyanin- the pigment in the juice

organic matter in the absence of air.

of red cabbage that can be used as an

26. Binary compound-A compound which

indicator.

contain two types of ions

11. Aqueous solution- a solution obtained

27. Biotechnology-The use of living

by dissolving a solute in water (solvent).

organism such as bacteria, fungi,

12. Apparatus- special tools and equipment

and cel s from plants and animals for

used in a laboratory.

industrial and scientific purpose to

13. Atmospheric pressure- force per unit

make product

area exerted on object on the earth as

28. Boiling point-The temperature at

a result of the attraction of the earth.

which a liquid boils to form vapour

14. Atmosphere- the mass of air 29. Brownian motion-An irregular motion

surrounding the earth.

of tiny particles suspended in fluids

15. Atom- the smallest particle that retain

30. Bunsen burner-A small adjustable gas

the property of the element.

burner commonly used in laboratories.

16. Atomic number- the number of 31. Cardiopulmonary resuscitation

protons in an atom.

(CPR)-Ultimately applying chest

17. Atomic mass - the weighted average

compression and lung ventilations to

of the atomic masses for the mixture

a person having difficulty in breathing.

of all natural y occurring isotope of an

32. Catalyst - A substance that speeds up

element.

the rate of a chemical reaction but

18. Atomic mass unit (amu)- the unit for

itself remains unchanged at the end of

expressing relative masses of atoms

reaction.

with reference to one- twelfth of the 33. Cathode-The negatively charged

mass of a carbon-12 atom.

electrode.

19. Atomic radius - The distance between

34. Cation-positively charged ion

the nucleus of an atom and the 35. Chemical bonds-force of attraction

208

that hold atoms or ions together in

substance is made of different parts.

chemical compound

49. Compound - a pure substance made

36. Chemical equation-symbolic

up of two or more kinds of elements

representation of a chemical reaction

that are chemical y combined together

showing chemical formulae of

in fixed proportions.

substance present before and after a 50. Compressible- capable of being

chemical reaction.

condensed or reduced in volume.

37. Chemical change –An irreversible 51. Concentrated solution- a solution

change that affects the chemical

containing a relatively large quantity

properties of substances and there is

of the solute.

formation of new products.

52. Concentration- a measure of the

38. Centrifugation-spring a sample around

quantity of solute dissolved in a

very quickly in a machine called a

particular volume of solution.

centrifuge, which causes sediment to 53. Conclusion- summary of the results of

settle and separate quickly from the

an experiment.

liquid

54. Condensation- the cooling of a vapour

39. Chromatography-A method of

to form a liquid.

separating two or more solute from the

55. Contaminant- any potential y harmful

liquid they are dissolved in: Example,

substance in the environment.

separating dyes in colored ink

56. Control experiment- in a control

40. Chemical warning signs-safety symbol

experiment all the independent

found on chemical containers.

variables are kept constant.

41. Chemist-A person who studies 57. Coordinate covalent bond- a chemical

chemistry

bond formed when one atom donates

42. Chemistry-The branch of science

both of the electrons shared in covalent

which deals with composition,

bond.

decomposition, structure and 58. Covalent bond- a pair of electrons

properties of matter.

shared by two atoms in a molecule.

43. Collar –A band near the base of a 59. Crystal- a solid having plane faces

Bunsen burner that can be rotated to

at definite angles and whose atoms,

control air supply to the burner.

ions, or molecules have a regular three

44. Colloid-contains very small particles

dimensional arrangement.

of one substance dispersed through out

60. Control ed variable- the factor in an

another substance usual y in different

experiment that is kept constant and

state. Example: Tiny particles of a solid

does not affect the outcome of the

dispersed in a liquid.

experiment.

45. Combustion- burning of a substance 61. Covalency- the process of sharing

in the presence of oxygen.

electrons by which many non-metals

46. Combustible- a substance which

and sometimes certain metals, join to

ignites and burns easily

form compounds.

47. Combination reaction- two or more 62. Cosmetics- products used for cleaning,

substances combine to form a single

beautifying, promoting attractiveness

compound.

or altering ones appearance.

48. Composition- the way in which a 63. Crystallization- the formation of

209

crystals from a homogenous solution.

liquid from a solute dissolved in it by

64. Crystallization point- the temperature

boiling.

at which crystals form from 79. Drugs- substances used to treat or

homogenous solution.

prevent diseases, diagnose illnesses or

65. Data- information that is analyzed

relieve pain.

so as to make decisions after 80. Effervescence- the escape of gas from

experimentation.

a liquid.

66. Decantation- pouring off a liquid into

81. Electrode- a conductor through which

a separate container, leaving sediment

an electric current enters or leaves an

behind.

electrolyte.

67. Dependent variables -Variables that 82. Electrolysis- the decomposition of an

cannot be changed by an investigator

electrolyte by an electric current.

or researcher.

83. Electronegativity- ability of an atom to

68. Decomposition- breaking down of

attract an electron.

chemical compound into elements or

84. Electron- the negatively charged

smaller compounds

particle of an atom.

69. Decomposition reaction- a compound

85. Electronic configuration- symbolic

breaking down into two or more

representation of the electron

simpler substance

arrangement in sublevels of an atom

70. Density- the degree of compactness of

using the electrons.

a substance.

86. Electronegative elements - elements

71. Detergent- liquid or solid compound

like fluorine, oxygen and nitrogen

or mixture of compounds used to assist

that have a very strong attraction for

in cleaning.

electrons involved in a chemical bond.

72. Diatomic elements - Elements which

87. Electrovalent bonding- force of

form molecules composed of two

attraction between ions.

atoms.

88. Electrolyte- a solution or molten

73. Diffusion- process by which a gas

substance that will conduct electricity

spreads out to fill all available space.

and can be electrolyzed.

A solute dissolved in a solvent also 89. Element- a pure substance that cannot

diffuses.

be split into simpler substances by a

74. Disinfect- to make clean and free

simple chemical process.

from infection, especial y by use of a 90. Electron shel - a set of energy levels

chemical.

outside the nucleus that contains the

75. Dissolving- process by which a solute

atom’s electrons.

breaks down into tiny particles and 91. Electrovalency - the process of

spreads throughout the solvent.

exchange of electrons which metals

76. Displacement reaction- a reaction in

and some non-metals join to form

which one substance takes the place of

compounds.

another.

92. Experiment- controlled investigation

77. Dissociation- the process by which

used to test or obtain facts, or establish

ionic compounds separate or split into

a hypothesis, or illustrate a known

smaller molecules, ions or radicals.

scientific law.

78. Distillation- separating out a pure 93. Empirical formula- the formula that

210

represents the simplest ratio of the

organisms in the ground millions of

atoms or ions in compound.

years ago.

94. Emulsifying agent- a substance 110. Food supplement- a nutrient added to

that can be added to an emulsion to

a food stuff to enhance it.

prevent it from separating into the two

111. Fractional distillation- separation

component liquids.

of two or more liquids of different

95. Emulsion- a mixture of liquids that do

boiling points by boiling and using a

not mix wel .

fractionating column.

96. Energy levels- special regions in which

112. Fractionating column- apparatus used

electrons orbit the nucleus of an atom.

in the distil ation of liquid mixtures to

97. Evaporation- process by which

separate them into their components

molecules, usual y from a liquid, leave

or fractions.

the liquid surface and become gaseous.

113. Fuels- a substance that burns to

98. Environment - the surroundings in

produce energy.

which something exists or lives.

114. Fume chamber- a ventilated enclosure

99. Filtrate- clear solution that passes

in a chemistry laboratory in which

through the filter paper or other

harmful experiments involving

filtration medium during filtration.

poisonous gases and vapours are

100. Filtration- process of separating an

carried out.

insoluble solid from a solution or 115. Galvanize- to coat iron or steel with

liquid.

zinc in order to prevent rusting.

101. Fertilizers- chemicals or natural 116. Gas- a fluid with no definite shape

substances added to soil to increase its

or volume but it is able to expand

fertility.

indefinitely.

102. Fire extinguisher- apparatus used for 117. Group- a vertical column of elements

putting out hazardous fires.

in the periodic table

103. Fire triangle- three components 118. Haber process- the industrial

required to start a fire.

production of ammonia on a large scale

104. First aid- the help given to a sick

by reacting hydrogen with nitrogen.

or injured person before medical 119. Halogens- group VII elements.

assistance from hospital.

120. Hard water - Water that does not easily

105. First aid kit- a small box in which

form lather with soap.

equipment and chemicals needed for

121. Hazard- something that is dangerous

the first aid are kept.

and can cause damage.

106. Flame- a burning gas which produces

122. Heat- energy transferred from

heat and light.

one body to another by thermal

107. Formula unit- the specific group

interactions.

of atoms or ions symbolized and 123. Heimlich manouvre - first aid

expressed in the chemical formula.

procedure for dislodging an

108. Formula- a collection of symbols

obstruction from a person’s throat.

and numbers, representing the 124. Heterogeneous- a substance that is not

composition of a compound.

the same all the way through.

109. Fossil fuel- fuels such as coal and oil 125. Homogenous- a substance that is the

originated from the remains of living

same all the way through.

211

126. Hydrated- a substance which contains

can only proceed in one direction.

water molecules.

141. Isotopes- atoms of a particular kind of

127. Hypothesis- a proposal or idea put

element that have different numbers of

forward to explain something, but not

neutrons and therefore different mass

yet proved.

numbers.

128. Immiscible- liquids that when mixed 142. IUPAC- The abbreviation for the

do not dissolve in each other to form

international union of pure and

a solution.

applied chemistry, which deals with

129. Incinerator- A facility in which the

chemical nomenclature.

combustion of solid wastes take place.

143. Kinetic molecular behavior- the

130. Insoluble - a substance that does not

manner in which particles in matter

dissolve to any easily detectable extent.

behave.

131. Independent variable- the factor in 144. Laboratory- a special room or building

an experiment that is manipulated

used for carrying out scientific

or changed to obtain values being

experiments.

measured.

145. Law- a rule that must be followed in

132. Indicator- a substance that exhibits

scientific studies.

different colours in solutions of 146. Limestone- rocks that are mainly

different acidity and basicity.

composed of insoluble calcium

133. Interpretation- an explanation of the

carbonate.

meaning or importance of something.

147. Liquefy- to make or to become liquid.

134. Ion- the electrically charged particle 148. Liquid- a fluid with a definite volume formed when an atom or group of

and takes the shape of the container

atoms either gains or loses electrons.

holding it.

135. Ionic compound- A chemical 149. Luminous- bright and full of light

substance that is made up of positive

150. Mass- a measure of the quantity of

and negative ions.

matter in an object.

136. Ionic equation- a balanced chemical 151. Mass number- the total number of

equation that shows all water soluble

proton plus neutrons in the nucleus of

ionic substances written in ionic forms,

an atom.

while insoluble solids and covalently

152. Matter- anything that has mass and

bonded substances are written in

occupies space.

molecular form.

153. Metallic bonding- the bonding of

137. Ion exchange resin- small beads made

metal atoms in solids where positive

of special plastic that remove hardness

metal ions are arranged in a regular

of water in an ion exchange.

three-dimensional way and the loosely

138. Ion exchanger- the device that removes

held valence electrons are able to move

both types of hardness of water by

freely throughout the crystal.

removing all calcium and magnesium

154. Metallic solid- a solid that is composed

ions.

of metal atoms.

139. Ionization energy- the energy required

155. Metals- shiny, dense substances that

to remove elections from an atom or

conduct electricity well and usually

ion.

have high melting points.

140. Irreversible reaction- a reaction that 156. Metalloids- elements that are neither 212

completely metallic nor completely

and neutrons.

non-metallic in their properties.

171. Nuclide notation- representation of

157. Miscible-liquids that mix very wel .

an atom of an element by its chemical

158. Mixture- physical combination of two

symbol, having its mass number on the

or more substances in any ratio.

upper left and it atomic number on the

159. Molar mass- the mass in grams of a

lower left of the symbol.

mole of any substance that is the sum

172. Octet of electrons- used to describe the

of the atomic masses of all the atoms

complete set of eight valence electrons

represented in the formula.

that are present in the outer most

160. Molecular formula-the chemical

energy level of noble gases that follow

formula that gives the actual number

helium in the periodic table.

of atoms of each kind present in a 173. Octet rule- the tendency for a non-

molecule of the substance.

metal atom to gain or share electrons

161. Molecular weight- the sum of the

until it has eight valence electrons.

atomic masses of all atoms in one 174. Oxidation- the process that occurs

molecule of a particular compound.

when substance combines with oxygen

162. Molecule- the smallest part of a

or any chemical process where there is

substance that can take part in

a loss of electrons.

chemical reaction and still retain the 175. Oxidation number- (also called

properties of that substance

oxidation state) the charge of a simple

163. Neutralization- the reaction between

ion or the “apparent charge” assigned

an acid and a base to give salt and

to an atom within a compound or

water.

polyatomic ion.

164. Neutron- an electrically neutral 176. Oxidizing agent- a substance that adds

particle with a mass of 1 amount that

oxygen to, removes hydrogen from or

is found in the nucleus of atoms.

loses electrons to, another substance.

165. Noble gases- the chemically stable 177. Oxy-hydrogen flame- a very hot

non-metal elements of group VIII of

flame produced by the combustion of

the periodic table.

mixture of oxygen and hydrogen.

166. Nomenclature - the system of names

178. Particulate nature- relating or in the

and formula used to identify all

form of very tiny particles.

chemicals.

179. Percent by mass- the mass of solute

167. Non-luminous flame- the blue flame

divided by the total mass of solution

of burner with an adequate supply of

times 100%.

oxygen and which produces more heat

180. Percent by volume- the volume of

than the luminous flame.

solute divided by the total volume of

168. Non-metals- group of elements that

solution times 100%.

do not conduct electricity (except 181. Percent by weight- the number of

granite).

grams of a particular substance that

169. Nucleons- sub-atomic particles found

can be found in 100 grams of the

in the nucleus of an atom, namely

sample.

proton, and neutrons.

182. Percent composition- a list of the

170. Nucleus- the tiny dense centre of an

percents by mass of each element in a

atom that contains an atom’s protons

compound.

213

183. Period of elements- a horizontal row of

law in science.

elements in the periodic table.

197. Problem- a scientific question to be

184. Periodic law- there is a periodic

answered.

variation in the physical and chemical

198. Products- the chemical substances

properties of elements when they are

formed during a chemical reaction.

arranged in order of increasing atomic

199. Properties- attributes or settings that

number.

are used to describe an object.

185. Periodic table- an arrangement of al

200. Proton- positively charged particle

known elements in order of increasing

with a mass of 1 a.m.u found in the

atomic number , set up according to

nucleus of all atoms.

the order in which electron shel s are

201. Pure substance- an element or

filled and so that elements with similar

compound composed of the same

properties fall into the same vertical

kind of matter with the same kind of

column.

particles throughout.

186. Periodicity- the regular periodic 202. Radical- a group of combined atoms

changes of elements according to their

that has a positive or negative charge

atomic numbers.

and tends to remain together as unit

187. pH value- a value that indicates the

during chemical reactions.

concentration of hydrogen ions in a 203. Reactants- the starting materials that

solution.

enter into a chemical reaction and

188. Pharmacy- a facility or location where

they appear on the left-hand side of a

drugs and other medical services are

chemical equation.

sold or dispensed.

204. Reactivity- the vigour with which an

189. Physical change- a temporary change

element reacts with other substances.

in the outward or physical properties

205. Reagent- a substance that takes part in

of a substance, usual y fairly easy to

chemical reaction.

reverse.

206. Recovering position- a position used

190. Pol utant- a contaminant introduced

to prevent an unconscious person

into the environment.

from choking in which the body is

191. Potable- describes water that is safe for

placed face downwards and slightly to

people to drink.

the side, supported by bent limbs.

192. Precipitate- a solid that forms and 207. Redox reaction- a reaction in which

separates out of solution as the result

both reduction and oxidation take

of a chemical reaction.

place.

193. Precipitation- water that fal s to the 208. Reducing agent- a substance that

ground in the form of rain, snow or

removes oxygen from, gives hydrogen

hail.

to, or gains electrons from, another

194. Precipitation reaction- a reaction

substance.

in which two soluble compounds 209. Reducing sugar- a sugar that must have

combine to give a soluble compound

available an aldehyde group that can be

and an insoluble compound.

oxidized to an acid by a mild oxidizing

195. Prediction- a kind of forecast which

agent.

suggests what the result is likely to be.

210. Reduction- the process that occurs

196. Principle- has got the same meaning as

when a substance loses oxygen or any

214

chemical process where there is gain

or backward.

of electrons.

217. Rust- The reddish-brown substance

211. Relative atomic mass (R.A.M)- The

that forms on the surface of iron or

mass of an atom compared with the

steel due to a chemical reaction with

mass of an atom of carbon -12, which

water and air.

is taken exactly 12.00000 units.

218. Rusting- the corrosion of iron and steel

212. Reservoir - a container designed to

by water and oxygen from the air.

hold fluids such as water.

219. Sacrificial protection- the process of

213. Residue- the insoluble material that

using some blocks of a more reactive

remains on the filter paper after

metal to protect large objects made of

filtration.

iron or steel from rusting.

214. Respiration- Breakdown of food 220. Saturated solution- a solution that

substances in the cell to release energy.

contains as much solute as it can

215. Resuscitate- Reviving an unconscious

dissolve at a particular temperature.

person.

216. Reversible reaction- a reaction that

can go in either direction of forward

215

BIBLIOGRAPHY

Holderness, A& Lambert, J (1987). A New Certificate of Chemistry, Heineman

Educational Publishers, London

SCSU & MoEVT - Zanzibar (2012); Chemistry for Secondary Schools, Forms 1

and 2, Oxford University Press (T) Limited, Dar es Salaam

Sedrick, C (2008) Comprehensive Chemistry for Secondary Schools, Book 2,

Nyambari Nyangwine Publishers, Dar es Salaam

Tanzania Institute of Education (2009), Chemistry Forms 1 and 2 Students’ Book,

Pearson Longman, Essex

Tanzania Institute of Education (1995), Secondary School Chemistry, Book One,

T.I.E, Dar es Salaam.

216

INDEX

AAbundance 159

Alkaline earth metals 166

Accident 8, 21

Alloys 94

Acid 8

Al uminium sulphate 134

Acid rain 143

Ammonia 115

Aeration 134

Ammonium chloride 75

Air 82,85,86,92,99

Apparatus 8

Alchemist 1

Atomic number 155

Alkali metals 166

Atoms 150,156

BBalloons 116

Breathing 105

Binary compound 193,195

Bonding 186,187,189

Biogas 144,145,146

Bunsen burner 8,16,33,34,35,36

Biogas plant 144,145

Burners 33,34

Boiling 49

Burning 36,86,88,89,105,106

Boiling point 49

CCalcium 62,120,126,130

Chemistry 1,2,3

Calcium chloride 93,124

Chlorine 2

Carbohydrate 3,144

Chromatography 78

Carbon atom 158,159

Cobalt (II) chloride 129

Carbon dioxide 105,143,144

Colloid 4,68

Carbon monoxide 115,143

Coloured substance 78

Catalyst 99,100,101

Combustion 86,102

Celsius (centigrade) 49,85,102

Compound 64,66

Centrifugation 78

Condenser 19,73,75

Chalk 197

Conservation of energy 143

Charcoal 139,142,143

Control experiment 44

Charcoal burner 33,34

Copper 61

Chemical bonding 186

Copper (II) oxide 113

Chemical change 54,56

Copper (II) sulphate 129

Chemical formulae 174,181

Corrosion 11

Chemical properties 54,102,112,129

Covalent bond 189

Chemical treatment 134

DDalton atomic theory 150

Decomposition 1

Decantation 70

Distil ation 72,73,74

EElectron diagrams 152

Electronegativity 164

Electrical conductivity 188,191,192

Electronic arrangement 152

217

Electrons 151,155

Energy sources 138

Electroplating 95

Ethanol 74

Electrovalent bonds 187,188

Evaporation 49,72

Element 60,61

Experiment 8,43,44

Empirical formula 183,184,185,186

Explosive 10

Energy 138, 140,143,144

F

Fertilizers 1,115

Flames 33,36,37

Filter paper 17

Flammable chemicals 11

Filtrate 71

Foods 3

Filtration 71

Fossil fuels 139

Fire extinguisher 90,91

Fraction distil ation 74

Fire triangle 89

Fractionating column 75

Fire fighting 88,91

Freezing 49

Fires 88,89

Fuels 2

First aid 21,22,23

Funnels 17

GGalvanizing 94

Gas 142,144,145

Gases 82

Gas jar 19,100,102,110

Good fuel 139

Goggles 15

Groups 164

Gypsum 197

HHeating 12,38

Hydrogen 109,110,112,113,114,115

Heterogenous 66

Hydrogen peroxide 100

Homogenous 66

Hydrogen sulphide 103

Hydrocarbons 68

Hydroxide 176

Hydrochloric acid 110

Hypothesis 43

IIce 23,49

Ions 179

Immiscible liquids 76,77

Iron 12,13,14,15,16,17

Independent variable 43

Irritant 10

Iodine 22,75

Isotopes 157

Ionic bonds 187

Isotopy 157

Ionic compounds 188,193

KKerosene 2

Kinetic theory of matter 52

Kerosene burners 33,34

LLaboratory 2,7,8,12

Law 143

Layer separation 77

Liebig condenser 73,75

218

Lime water 84

Luminous flame 36,37,38

Liquids 8

MMagnesium 109,131

Metals 166,168

Magnesium oxide 131

Methane 2,144

Manganese (IV) oxide 99,100,101

Milk 69

Margarine 115

Miscible 69

Mass number 155

Mixture 70

Matter 1,48

Modern periodic law 165

Melting 49

Molecular formulae 183

Mendeleev 164,165

Molecules 63

Metalloids 166

Mortar and pestle 17

NNaming compound 193

Nomenclature 174,193

Nature and occurrence of water 120

Non - luminous flame 36,37,38

Neutrons 151

Nucleus 152

Noble gases 166,167

Nuclide notation 156

OObservation 43

Oxidizing agents 11

Oxidation numbers 175

Oxygen 99,102,105

Oxidation state 175

PPaper chromatography 78

Pol utants 143,146

Particulate nature of matters 50

Precipitation 120,121

Periodicity 165

Prediction 43

Pestle 17

Preparation of oxygen 99

Physical change 54,55,56

Proton 151

Pipette 12

Purifier 132

Plating 95

RRadical 178

Renewable 138

Radioactive 10

Reservoir 134

Reagent bottle 15

Respiration 105

Reduction 114

rust 92,94

Relative molecular mass 158

Rusting 92,94

S

Sacrificial protection 94

Scald 23

Safety 7,8

Science 1

Sand 34,133,134

Scientific procedure 42,46

Sand filter 133,134

Sedimentation 2

219

Separation of mixture 70

Stainless steel 94

Shel s 152

States of matter 48

Sodium 130

Steel 94

Solids 42

Subatomic particles 151

Solute 66,68

Sublimation 49,75

Solution 66,68

Suspensions 68

Solvent 66,68

Syringe 12

TTemperature 12,145

Toxic 8,10

Test tube 18

Transition metals 166

Theory 150

Transpiration 120,121

Thermometer12,73,75

Trends 167,168

UUniversal solvent 129

Urban water treatment 134

Uranium 138,142

VVaccine 1,3

Variables 43

Valency 175,176

Volume 48

W

Warning signs 10

Water treatment 2,132

Water 120,122,129,135

Welding 105,115

Water cycle 121

Wet chemical 91

Water gas 114

Z Zinc 94,109,110

220

221

Table of Contents

Chapter 1 - Introduction to chemistry

Chapter 4 - scientific procedure

Chapter 5 - Matter

Chapter 8 - Oxygen

Chapter 9 - Hydrogen

Chapter 10 - Water

Chapter 11 - Fuels and Energy

Chapter 12 - Atomic Structure

Chapter 13 - Periodic classification

Chapter 2

Chapter 3

Chapter 4

Chapter 5

Chapter 6

Chapter 7

Chapter 8

Chapter 9

Chapter 10

Chapter 11

Chapter 12

Chapter 13

Chapter 14