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Introducing chemistry concepts

Chemical reactions and energy changes

PS548 SCIENCE FOR PRIMARY TEACHERS

Theopen University

INTRODUCING CHEMISTRY CONCEPTS

1 CHEMISTRY: GENERAL INTRODUCTION 3

2 UNDERSTANDING THE CHEMISTRY IN A PRIMARY SCIENCE TOPIC 4 2.1 Food and diet 4 2.2 Primary school activities associated with food and diet 4 Activity 1 : food and healthy living 4 2.3 Chemical concepts 5 2.4 Food as an energy source 5 Summary of Section 2 6

3 CHEMICALS AND ATOMIC THEORY 6 3.1 Recognizing chemical changes 3.2 Mixtures and compounds 3.3 Atoms and elements 3.4 The masses of atoms 3.5 Chemical formulae 3.6 Chemical equations 3.7 Valency Summary of Section 3

4 METHODS FOR SEPARATING SUBSTANCES FROM OTHER SUBSTANCES 4.1 Chemical purity 4.2 Soluble and insoluble substances Activity 2: separation of a mixture of a soluble and an insoluble substance Activity 3: Growing crystals Activity 4: Extensions to crystal growing 4.3 Separation by chromatography Activity 5: Experiments with Smarties 4.4 Separation by distillation 4.4.1 Solids, liquids and gases, the three states of matter 4.4.2 Simple distillation 4.4.3 Separation by freezing

5 ATOMIC STRUCTURE 23 5.1 A simple atomic model 23 5.2 Sub-atomic particles 24 5.3 Nuclear symbols 24 5.4 What are isotopes? 25 5.5 How are electrons arranged in atoms? 26 5.6 The Periodic Table 28 Summary of Section 5 28

6 CHEMICAL BONDING 29 6.1 Noble gas electron configurations 29 6.2 Ionic bonding 29 6.3 Covalent bonding 31

6 6.4 Electronegativity 32 7 6.5 Intermolecular forces 34 8 Summary of Section 6 36

11 ITQ ANSWERS AND COMMENTS 37 12 12 NOTES 38

CENTRE FOR SCIENCE EDUCATION

SCIENCE FOR PRIMARY TEACHERS: CONTRIBUTORS

Bany Alcock (human biology, Nene College, Northamptonshire) Mary Aitken (chemistry, consultant author) Fiona Allen (reader, Hillside Infants School, Northwood, Middlesex) Bob Allgrove (chemistry, Chichester College of Technology) Matthew Baird (advisory teacher, London Borough of Enfield) Steven Baker (Earth sciences, Droitwich High School) Chris Brown (Earth sciences, consultant author) Sue Browning (advisory teacher, EPSAT, Essex) Andrew Coleman (editor) Hazel Coleman (editor) Chris Culham (advisory teacher, EPSAT, Essex) , Carolyn Dale (advisory teacher, Buckinghamshire) Myra Ellis (secretary, electronic publishing, The Open University) Graham Farmelo (physics, The Open University) Stuart Freake (physics, The Open University) David Gamble (county adviser, science, Buckinghamshire) Jack Gill (senior science inspector, Essex) Jackie Hardie (adviser, London Borough of Enfield) Linda Hodgkinson (CO-director, Science for Primary Teachers, The Open University) Barbara Hodgson (ET, The Open University) Anne Jones (deputy headteacher, Simpson Combined School, Milton Keynes) Hilary MacQueen (biology, consultant author) Baird McClellan (consultant author) Catherine Millett (chemistry, consultant author) Peter Morrod (chemistry, The Open University) Shelley Nott (illustrator, En-igma Design) Katharine Pindar (information officer, The Open University) Jane Savage (Institute of Education, University of London) David Sayers (Science INSET co-ordinator, North London Science Centre) John Slade (chemistry, consultant author) Freda Solomon (advisory teacher, London Borough of Enfield) Valda Stevens (biology, consultant author) David Surnner (physics, Tarragon Press) Liz Swinbank (physics, consultant author) Margaret Swithenby (biology, The Open University) Peter Taylor (chemistry, The Open University) Jeff Thomas (biology, The Open University) Susan Tresman (CO-director, Science for Primary Teachers, The Open University) Liz Whitelegg (academic liaison adviser, The Open University) Margaret Williams (advisory teacher, Buckinghamshire) Geoff Yarwood (electronic publishing, The Open University)

The Pilot Project for Science for Primary Teachers was made possible by funding from the Department of Education and Science and from National Power plc and Nuclear Electric plc.

The Open University, Walton Hall, Milton Keynes MK7 6AA.

First published 1991.

Copyright O 1991 The Open University. 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 permission in writing from the publisher or a licence from the Copyright Licensing Agency Limited. Details of such licences (for reprographic reproduction) may be obtained from the Copyright Licensing Agency Ltd., 33-34 Alfred Place, London WClE 7DP.

Printed in the United Kingdom by H. Charlesworth & Co. Ltd. Huddersfield. Further information on this and other Open University courses may be obtained from the Learning Materials Sales Office, The Open University, P.O. Box 188, Walton Hall, Milton Keynes, MK7 6DH.

ISBN 0 7492 5032 1

1.1


1 CHEMISTRY: GENERAL INTRODUCTION The way of presenting chemistry in Science for Primary Teachers is different from that used for physics, Earth sciences and biology. Here we have departed from using S102 Units as the main source of information with Study Commentaries and will be using more or less free standing texts that refer occasionally to Units 11-18 of S102.

We begin by looking at the topic of food and diet and where chemistry concepts can be helpful in understanding this more fully. We then go on to look in more detail at a number of these concepts from atomic structure to the structure of giant molecules-polymers.

Section 4 of 'Introducing chemistry concepts' is very practical in its approach. The experiments only require very simple apparatus and materials that can usually be obtained in the home, at school or cheaply from shops. We recommend that you try the experiments yourself and then use them in classwork, with any modifications you think necessary. Most of the topics in this Section are not covered to any extent in S102. If you wish to pursue them further we suggest you look at GCSE (or 0-level) chemistry texts. These topics relate directly to key stages 1 and 2 of AT 1 and AT 6. To carry out the experiments in Section 4 you will find it useful to collect the equipment and materials listed in Table 3 of the Introduction to the Study Commentaries before you start.

AT6 deals with the properties of materials and the way properties determine use and underpin classification. The methods used to separate different chemical substances from each other provide ways of examining similarities and differences between materials (key stage 1 of AT6). AT1, as you know, is concerned with the processes of scientific enquiry.

The practical work included in 'The chemistry of carbon compounds' should be camed out with the help of an experienced chemist. Here, the aim is to provide you with the opportunity and experience of handling laboratory chemicals safely and to help your understanding of the concepts introduced.

The chemistry section of the Course is concluded with work relating to three different topics-Materials, Fuels and Clean science. Many of you use the topic approach to teach areas of the primary curriculum. Here, we show how chemistry can be considered through topics. The emphasis is to show how such class topics could include the different scientific concepts and skills that you have covered. Fuels and Clean science are topics for you to work through to help your own understanding of some of the scientific concepts and encourage you to pose new questions. However, you will have no difficulty in realizing the classroom and INSET potential. Clean science involves some practical work for which you will need to obtain some special materials. Details are included in the Introduction to the Study Commentaries.

As you work through the chemistry section of the Course the concepts introduced will enhance your subject knowledge and skills relating to ATs 2, 3, 5, 6, 7, 8 and 13.

SCIENCE FOR PRIMARY TEACHERS

2 UNDERSTANDING THE CHEMISTRY IN A PRIMARY SCIENCE TOPIC

2,l FOOD AND DlET Food and diet are familiar topics in the primary school. Their importance has been enshrined within the statutory requirements for the national curriculum for science of all children in state schools in England and Wales.

For example, the programmes of study state:

for key stage l-'Children should be finding out about themselves, developing their ideas about how they . . . feed.'

for key stage 2-'Children should develop further an awareness of the role and importance of science in everyday life. ... Domestic contexts should be introduced . . . as starting points for children's work in science.'

AT 3 includes the statements:

Level 2-'Pupils should know that . . . food . . . [is] important.'

Level 3-'Pupils should know that the basic life processes [including] feeding . . . are common to human beings.'

Level 4-'Pupils should know about the factors which contribute to good health and body maintenance, including . . . balanced diet.'

@ Level 5-'Pupils should understand malnutrition.'

2.2 PRIMARY SCHOOL ACTIVITIES ASSOCIATED WITH FOOD AND DlET

ACTIVITY 1 : FOOD AND HEALTHY LIVING Prepare an outline scheme of work for the children on the subject of food and healthy living.

Your scheme should take into account the ages, abilities, ethnic backgrounds and previous experiences of the children. It may include some of the following ideas:

Making displays of food packets, adverts, labels, etc.

This will show the variety of foods, and introduce ways of grouping them. Harvest Festival might provide a good opportunity to display the food itself.

Categorizing foods by specific contents (e.g. sugar-containing and sugar- free foods; foods rich in protein/carbohydrate/fat/particular vitaminslmin- eralslfibre), so that children are aware that different foods share common components.

Considering the effects of malnutrition and good nutrition (e.g. the problem of scurvy; problems of famine and of plenty). This helps children identify the contribution made by different components of food to growth, health and energy.

Finding out about the treatment of food. (This could be a major topic in its own right, and might include ways of preparing and preserving different foods.)

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Modelling or presenting posters of balanced meals, to show how foods with different components can be combined to produce a balanced diet.

Designing and marketing a new foodstuff (e.g. a muesli-type fruit and cereal mixture). This requires the application of what has been learned and understood about food and diet.

Your outline scheme will probably be based on the following facts:

A balanced diet keeps us healthy.

Some foods are better for us than others.

Food helps growth, maintains, health and gives energy.

Potatoes, rice and pasta contain carbohydrates, for energy; meat, milk and fish contain proteins, for growth; and fat, which can also provide energy; fruits and vegetables contain mineral salts and vitamins, for health.

Cooking makes food more tasty and more digestible, and kills bacteria.

Digestion begins in the mouth, with the action of saliva, and continues in the stomach and small intestine, where the useful constituents are absorbed into the blood. The waste material is excreted. (You will learn more about this in Unit 23.)

The shelf-life of food can be extended by, for example, cooking, freezing, canning and refrigerating.

2.3 CHEMICAL CONCEPTS In order to have a full knowledge and understanding of food in the human diet it is important to be familiar with certain chemistry concepts.

For example:

1 Proteins, carbohydrates, fats, etc., are chemical compounds; they are composed of atoms of different elements.

2 Chemical compounds may involve covalent or ionic bonding (which can be explained in terms of atomic structure and electronic conJiguration).

1 and 2 are introduced in Sections 5 and 6 of 'Introducing chemistry concepts'.

3 Food can provide energy because of the chemical reactions that occur during digestion, and within the cells of the body.

4 Cooking and refrigeration show the effects of heating and cooling on rates of reaction.

3 and 4 will be considered in 'Chemical reactions and energy changes'

5 Carbohydrates are carbon compounds, with properties determined by their functional groups. They are giant molecules, which are broken down by cooking and digestion into the simple molecules of water and carbon dioxide.

These compounds are considered in 'The chemistry of carbon compounds'.

2,4 FOOD AS AN ENERGY SOURCE The chemical composition of food is important in the growth and repair of body tissues, and for the regulation of body processes. However, the other major role of food is to provide energy for action. An appropriate diet will vary according to the energy demands of our life-style, but on average about 70% of our energy will be provided by carbohydrates. This is why breakfast cereals, which are rich in carbohydrates, include on their packaging a statement of their energy content in kilojoules (H) or kilocalories (kCal). Even fats are useful in moderation, for they provide the body's long-term energy reserves.

It is possible to demonstrate that food is a sort of fuel that can be burnt in air to produce heat and light-indeed a powdered carbohydrate such as custard powder can bum with explosive force in a confined space. However, these are very crude reactions compared with the series of gentle, controlled stages that occur within the body as the food chemicals are made to release their energy gradually. . In common with other energy stores, such as reservoirs high in the mountains, which obtain their energy from the hydrological cycle, or tightly coiled springs, which acquire energy by being physically coiled, we need an energy source for the production of food.

It is common knowledge that most plants grow as a result of nutrients from the soil and carbon dioxide from the air being converted into carbohydrates by the action of sunlight in the presence of a chemical called 'chlorophyll'. Energy from the Sun is thus stored in the carbohydrates and can eventually be released for our benefit.

SUMMARY OF SECTION 2 l Topics, such as food and diet, which are important areas of the primary school curriculum, can be better understood with some knowledge of the underlying chemistry.

2 Almost all foods are mixtures of complex chemical compounds.

3 A chemical change occurs when food is digested and converted into different compounds, which may be required for body maintenance, growth and energy production.

3 CHEMICALS AND ATOMIC THEORY

RECOGNIZING CHEMICAL CHANGES Every substance is a chemical: from the air you breathe and the food you eat to the intricate collection of matter that makes a human body. Chemistry is the study of all the materials in the natural world and of the ways in which they can be transformed into new substances. Chemists spend much of their time observing materials, trying to bring about changes and looking for signs that changes have occurred. Not all changes result in the formation of new substances, so the ability to recognize a chemical change and distinguish it from a physical change is an important skill for a chemist to acquire.

A chemical change occurs when new substances are formed. But what is meant by the term 'new substance'? When you have learned more about the particles that make up matter and the forces that hold them together, you may wish to think again about chemical change.

How can you tell when new substances are being formed? Think, for example, about an ice-cube melting to form liquid water. Everyone knows that ice and water are made of the same substance, which has the chemical formula H20. So changing from solid to liquid is merely a physical change-a change of state. How would you recognize such a change in an unfamiliar material? If you refreeze liquid water, it will turn back into solid ice. A physical change is usually easy to reverse.

Now think about boiling an egg. The yolk and white of a raw egg are both liquids. Heat will change them to solids but no amount of cooling will cause the egg to revert to its original form. New substances have been formed and a chemical change has taken place.

Observations of some other changes are shown in Table 1. -


TABLE 1 Some examples of change

Process Observations

Striking a match

Ripening a green tomato

Adding liver salts to water

The match bums, giving out heat and yellow light. A spent match cannot be re-used

The tomato turns red. Once ripe, its colour will not revert to green

The solution fizzes. Some matter is lost as gas, so the mass of the remaining solution decreases

Adding boiling water to instant coffee A brown solution is formed. The powder coffee powder can be recovered by

evaporating the water from the solution

Passing an electric current through a coil of A magnetic field is produced, but this wire disappears as soon as the electric

current is turned off

0 What characteristics are associated with a chemical change?

A chemical change can involve a change in energy or colour. New substances, such as gases, may be easily recognizable. There may be an apparent loss in mass. The changes are hard to reverse.

0 What characteristics are associated with a physical change? The change may be easy to reverse and no new substances are formed.

The difference between chemical and physical changes is by no means clear cut. The dissolving of instant coffee powder in water as described in Table 1 seems to be a physical change. But in most cases dissolving is accompanied by an energy change, and most chemists would consider dissolving to be a sort of chemical change, even though it is possible to recover the original components by physical means. They reach this decision by making use of modern theories of structure and bonding, some of which are described later.

3,2 MIXTURES AND COMPOUNDS A mixture contains more than one substance. Instant coffee is a mixture of the chemicals in coffee granules and hot water. You might also add milk and sugar. How much of these ingredients you choose to add depends on personal taste-the drink remains, unmistakably, coffee. Adding more coffee powder will make the drink taste more bitter. Adding more sugar will make the drink taste sweeter. A mixture can have a variable composition and its properties are those of its components. It is possible to separate a mixture using physical techniques such as filtration and evaporation. In the case of instant coffee, you would boil away the water and get the coffee grains back again.

A compound also contains more than one substance. (The substances making up compounds are called elements. You will learn about these in Section 3.3.) Water, for example, is a compound-it contains hydrogen and oxygen. But water is not a mixture of those two substances, its properties are unique and quite different from those of a mixture of hydrogen and oxygen. For example, at room temperature water is a liquid whereas a mixture of hydrogen and oxygen is gaseous.


It is possible to separate a mixture of these two gases by cooling. Oxygen liquefies at a temperature of -183 "C while hydrogen is still a gas, so the liquid oxygen could be tapped off from the mixture. If you cool water, it freezes to form a single solid substance, ice, at 0 "C. Water can be broken down into its components by passing an electric current through it. This process, called electrolysis, brings about a chemical change, which is easily recognizable because the products are the two gases hydrogen and oxygen. New substances have been formed.

Making a compound involves a chemical change. To make water from hydrogen and oxygen you would have to supply energy in the form of a spark. The mixture of gases would explode, giving out more energy in the form of light, sound and heat. If you cooled the gas left after the explosion, it would condense to form a colourless liquid-water. Compounds also differ from mixtures in that they containfied proportions of their components.

Water contains two parts hydrogen to one part oxygen. Sparking a mixture of equal volumes of each gas would produce water, but half the oxygen would be left unre8cted.

Using words is a long-winded way of describing the nature of mixtures and compounds. Scientists tend to explain such ideas in terms of models that can be represented pictorially. In this case, they would consider all matter to be made of particles. Look at the diagrams in Figure 1.

FIGURE 1 Mixtures and compounds.

How many particles are there in each diagram? There are eight particles in each one.

Which diagram represents the mixture and which the compound? A is the mixture and B the compound.

How many different types of particle are there in the mixture? There are three types of particle in the mixture.

How many different types of particle are there in the compound? There is one type of particle in the compound.

Each of the particles in the compound is identical-the compound is a pure substance. On the other hand, the three types of particle in the mixture represent different substances, which can be separated by physical methods and which can be present in any proportions.

3.3 ATOMS AND ELEMENTS Sometimes scientists need to look at matter in more detail and regard it as more than a collection of single particles. You will have noticed that some of the particles in Figure 1 are in fact made of two smaller ones. These smaller particles are called atoms. The idea that matter consists of atoms is very old: it was first considered by the ancient Greeks. The philosopher Democritus described a Universe made up of atoms and the void, the atoms being different in size and shape and incapable of change. These ideas about atoms are not very far removed


from our modern ones, though they were based on philosophy rather than science.

John Dalton, the English chemist, revived ideas about atoms in the early nineteenth century. In his atomic model, he regarded atoms as the smallest indivisible particles of elements, rather like tiny, hard spheres. He used atomic theory to explain the results of experiments in which he investigated the composition of the products of some chemical reactions.

Look again at 'the diagrams in Figure 1.

0 How many different types of atom are shown? B Just two types of atom are shown, 0 and 0 .

The compound is made up of two types of atom, joined together as a result of a chemical reaction. The mixture, however, happens to include some particles that contain only one type of atom, and 0. Substances that contain only one type of atom are known as elements.

There are 92 naturally occumng elements, and scientists have managed to make about 20 others. Some elements occur on Earth in great abundance, others are very rare. Each element is made from atoms that are chemically similar, but which are different from the atoms of every other element. Elements are assigned symbols, as a form of shorthand, some of which are shown in Table 2.

You will find a full list of elements and symbols in Appendix 1 (p. 72) of Units 11-12 of S102.

TABLE 2 Some common elements and their symbols

Element Symbol

argon aluminium barium beryllium boron carbon calcium chlorine chromium cobalt fluorine gold helium hydrogen iodine iron lead lithium magnesium manganese mercury neon nickel nitrogen oxygen phosphorus silicon silver sodium sulphur tin


The reasons for choosing most of the symbols is obvious, but some refer to Latin names for the elements. For instance, the Latin for sodium is natrium, for potassium, kalium, for iron, ferrum and for copper, cuprum. In the list of elements shown in Table 2 you will probably recognize some everyday substances. Others will be less familiar to you. Each element has its own characteristic chemical and physical properties. Some elements are solids at room temperature and pressure, others are liquids or gases. Some elements are necessary for life, others are poisonous. You can probably pick out some that you know to be metals, whilst others you will recognize as non-metals.

Elements, then, are made of atoms of the same kind, but compounds and mixtures contain atoms of more than one kind.

It is surprising that of the 92 naturally occurring elements, only four-carbon (C), oxygen (0), hydrogen (H) and nitrogen (N)-are present in those components of food that account for most of its mass. Of these, carbohydrates and fats are composed of carbon, oxygen and hydrogen; proteins are compounds of all four elements.

However, other elements are also essential for life and must therefore be present in the diet, even though the amounts needed are small. Calcium (Ca) and phosphorus (P) account for 75% of the mass of the eight 'mineral elements' present in human beings (little more than 1.5 kg in total); the other 25% comprises potassium (K), sulphur (S), sodium (Na), chlorine (Cl), magnesium (Mg) and iron (Fe)-(about 5 g in an average male).

Finally, there are the trace elements: copper (Cu), zinc (Zn), manganese (Mn) and molybdenum (MO).

A B

FIGURE 2 For use with ITQ 1.

B FIGURE 3 For use with ITQ 2.


I T Q 1 Which of the diagrams in Figure 2 represents:

(a) A mixture of two elements? (b) A mixture of two compounds? (c) A mixture of an element and a compound?

ITQ 2 Which of the diagrams in Figure 3 represents:

(a) The decomposition of a compound into its constituent elements? (b) The combination of two elements to form a compound?

ITQ 3 Classify the substances below as mixtures, compounds or elements, giving a reason for your choice in each case.

(a) soil (b) carbon dioxide

( 4 gold (d) tap water

3.4 THE MASSES OF ATOMS Individual atoms are too small to be seen with an optical microscope, even the most powerful. They are also too small to be weighed, even with the most sensitive balance. However, since the early part of the twentieth century, it has been possible to make indirect measurements of the masses o f atoms using a technique called mass spectrometry. By this method, the mass of an atom of neon, for example, is found to be 3.3 X 10-26 kg, which is very small indeed. For convenience, scientists often express the masses of different atoms relative to one another. They call the value arrived at in this way the relative atomic mass (A,). It would seem sensible to choose as a reference the lightest atom, hydrogen, and set its relative mass equal to 1. 9

0 Explain why relative atomic masses have no units. They have no units because they represent only the number of times heavier an atom is than an atom of hydrogen.

If you want to know more about mass spectrometry, which enables us to find values for relative atomic masses, read Section 2.2.1 of Units 11-12 of S102.

Values for relative atomic masses of some common elements are shown in Table 3. A full list can be found in Appendix 1 of Units 11-12.

TABLE 3 Some common elements and their relative atomic mass

Atom Relative atomic mass


ITQ 4 How many times heavier is the atom of the first-named element than an atom of the second-named element? (a) Mg and C (b) Br andCa (C) Cu and S (4 NandH (e) Cu and 0

3.5 CHEMICAL FORMULAE Symbols are used by chemists to summarize the composition of compounds by means of chemical formulae.

Consider the familiar formula H20. What do you think the formula suggests?

Chemical analysis has shown that the masses of hydrogen and oxygen that combine to produce water (which are the same as the masses of each element produced by the decomposition of water) are in the ratio of 1 : 8. The mass of a single atom of hydrogen is approximately l unit, and that of a single atom of oxygen (measured by a mass spectrometer) is 16 units. If we assume that atoms are not split, the simplest combination of atoms of hydrogen and oxygen in water must be two hydrogen atoms per oxygen atom, or H20. The formula H20 represents a molecule of water.

In a complete reaction between sodium and chlorine to produce salt (sodium chloride) the sodium and chlorine combine in the proportions by mass of 23 : 35.5.

0 Using Table 3, what do you think will be the ratio of atoms of sodium (Na) and chlorine (Cl)? The ratio is clearly 1 : 1, and the simplest formula is NaC1.

Section 2, about food and diet, began by referring to some of the basic chemical constituents of foods such as carbohydrates and proteins. These have far more complicated formulae than water and sodium chloride.

Consider one of the carbohydrates : the sugar glucose.

Its formula can be written as:

C64206

You may be wondering why the formula is not CH,O. Certainly the ratio 6 : 12 : 6 simplifies to 1 : 2 : 1, but the instrument called a 'mass spectrometer' measures the mass of one molecule of glucose as approximately 180 units, whereas the mass would be only 30 units if each molecule contained only 1 atom of carbon and oxygen combined with 2 atoms of hydrogen.

So a molecule of glucose consists of 6 atoms of carbon, 12 of hydrogen and 6 of oxygen. The number of atoms is shown by a subscript.

3.6 CHEMICAL EQUATIONS It is a short step from using symbols to summarize the composition of molecules of chemical compounds to using them to summarize an entire chemical reaction.

Consider once again sodium chloride. One possible way to make sodium chloride is by burning sodium in chlorine. The reaction could be summarized as follows :

Na + C1 = NaCl


However, this representation would not correspond to the chemist's analysis of chlorine gas. In chlorine, and in many other gaseous elements, the atoms are paired up into molecules, as in H2, O2 and Cl2. The correct formula for the reaction is therefore

Cl What do you think 2Na represents, and why was it needed? 2Na represents the two atoms of sodium that are needed to make the equation balance-there are two atoms on the right hand side of the equation in the compound 2NaC1 means two molecules of sodium chloride.

When glucose is converted in the body using oxygen that has been breathed in, the eventual products are carbon dioxide (breathed out) and water (lost in urine and sweat). The overall reaction can be summarized as follows:

The sugar was itself made, in plants, by complicated reactions, requiring energy from sunlight and the presence of the green pigment chlorophyll. The actual process may be complicated but the overall reaction can be summarized as the reverse of the one above:

When writing equations, the number in front of a molecule represents the number of molecules; for example, the 6 in ' 602 ' means six molecules of oxygen.

3.7 VALENCY The empirical formulae of very many chemical compounds contain small whole numbers representing the relative numbers of atoms in the compound. There are important theoretical reasons for these ratios to be small whole numbers, which we will discuss in Section 5. Here we focus on, the relationship between the numbers of atoms that combine in different compounds. An examination of this relationship leads to a property called valency, which is helpful in predicting the formulae of compounds.

Consider the following three formulae:

Notice that in HCI the elements hydrogen and chlorine combine with each other in the atomic ratio of 1 : 1. Each of these two elements combines with nitrogen in the same ratio of 3 : 1 in NH3 and NC13. There appears to be a relationship between the three formulae. This kind of observation for sets of compounds led chemists to suggest that the formula of a compound is determined by some inherent property of the constituent elements. This property, called valency, determines the relative numbers of atoms of elements in compounds. If each element had only a single value of valency, the formulae of compounds could be readily predicted from a knowledge of those valencies. We will assume, for the moment, that this is theqcase. The rules for determining the valency of an element are then quite simple, particularly in binary compounds (that is, compounds that contain two elements). Because valency is based on formulae, it is concerned with the relative numbers of atoms in compounds. So, some arbitrary reference element is chosen as the basis of values of valency. This element is hydrogen, and by definition its valency is taken to be one.


If an element forms a binary compound with hydrogen, its valency is equal to the number of atoms of hydrogen that combine with one atom of that element. For example, in HC1, one atom of chlorine combines with one atom of hydrogen. So the valency of chlorine is one.

0 What valency do you deduce for nitrogen on the basis of its compound with hydrogen, NH,? In NH, one atom of nitrogen combines with three atoms of hydrogen, so the valency of nitrogen is three.

If the valency of each of the elements does not vary, it should be possible to predict the formula of a binary compound. For example, we have concluded that the valency of nitrogen is three (in NH,), and the valency of chlorine is one (in HC1). With this information, we would expect nitrogen and chlorine to form a compound with the formula NC13 which indeed is the case. Thus, valency has important predictive properties. Moreover, the observation that nitrogen and chlorine combine as NC13 shows that it is not necessary to rely on a knowledge of the formula of a compound of hydrogen in order to determine the valency of an element. For example, if the valency of chlorine in HC1 is the same as that of hydrogen (one), the valency of nitrogen is equal to the number of atoms of chlorine that combine with one atom of nitrogen. In NC13 therefore, the valency of nitrogen is three. To take another example, hydrogen and oxygen combine as water, which has the formula H20.

0 What is the valency of oxygen in H20?

One atom of oxygen combines with two atoms of hydrogen, so its valency is two.

As the examples above illustrate, the idea that elements have a fixed valency is useful in predicting formulae. In general, if element A has valency X, and element B has valency y, then if A and B combine to form a compound, y atoms of A will combine with X atoms of B.

Unfortunately, many elements have two or three common values of valency, which reduces the predictive power of the concept. Nevertheless, it is still possible to predict the formulae of compounds that might be formed by choosing formulae that are consistent with the known values of valency.

However, there are limitations to the generality of these predictions and to the information that valency provides. For example, valency alone gives no indication of the nature of the interaction between atoms.

ITQ 5 In forming compounds, silicon (Si) generally exhibits a valency of four. What is the empirical formula of the simplest substance that silicon can form with hydrogen?

ITQ 6 Sulphur (S) forms an oxide in which sulphur has a valency of six. What is the empirical formula of that oxide?

I T Q 7 Carbon and hydrogen form a compound, methane, with the formula CH,. On the basis of its valency in methane, what are the formulae of the simplest compounds that carbon should form with (i) oxygen and (ii) chlorine?

SUMMARY OF SECTION 3 1 All substances are made out of atoms.

2 Atoms can combine together to form larger particles.

3 All substances can be classified into one of three categories: elements, compounds or mixtures.

4 These categories can be distinguished by using a model of atoms as tiny spheres.

(a) Elements contain only one type of atom; compounds and mixtures contain more than one type of atom.


(b) Compounds contain atoms that.have combined together in fixed proportions as a result of a chemical reaction.

(C) The particles in a mixture are not identical and can be present in any proportions.

5 Although atoms are very small, chemists can compare their masses by using the idea of relative atomic mass.

6 The composition of compounds can be represented using symbols to show the numbers of each type of atom present in the chemical formula.

7 Chemical equations can be used to represent chemical reactions.

8 The valency of an element is equal to the number of hydrogen atoms that combine with one atom of that element. Elements have either a single value or a small number of values of valency. So valencies of an element are useful in the prediction of the formulae of the possible compounds that the element might form.

4 METHODS FOR SEPARATING SUBSTANCES FROM OTHER SUBSTANCES

4,l CHEMICAL PURITY To a chemist a 'pure' substance is a single substance with no trace of anything else mixed with it. Although in practice this cannot be completely achieved, much of the work of a manufacturing chemist involves purification of impure substances. Therefore, methods for separating out unwanted contaminants are important.

4.2 SOLUBLE AND INSOLUBLE SUBSTANCES A substance in which another substance dissolves is called a solvent. The substance which dissolves is called the solute. Not all substances dissolve to the same extent in a given solvent. For example, sand does not dissolve in water, whereas sodium chloride readily dissolves in water to form a solution. In this solution, water is the solvent and sodium chloride is the solute.

What precisely do we mean by the term 'dissolve'? (This can form the basis of class investigations into which common substances dissolve and which do not.)

Try shaking about a teaspoonful of sand with about half a cupful of distilled or deionized water in a clear plastic bottle. Then put the bottle down and watch what happens. The largest particles will settle quickly leaving a cloudy liquid above. This is a suspension and is cloudy because the undissolved particles are big enough to reflect light. Examination of this liquid under a microscope reveals the individual particles. They are unchanged by the water-they are insoluble. With prolonged standing even these very small particles fall to the bottom of the container, leaving clear water above. None of the sand has disappeared.

Now repeat the above experiment using sodium chloride (cooking salt) and compare the results. Make sure the comparison is 'fair'-use the same volume of water and the same volume (or mass) of salt as you used of sand. You should see the salt dissolve completely to give a clear colourless solution. The salt has disappeared. (If you use table salt your solution may be slightly cloudy. Table salt often contains. the anti-caking agent magnesium carbonate, which does not


dissolve. Another common anti-caking agent, sodium hexacyanoferrate(II), does not have this effect.)

What do we mean on the molecular level by saying salt dissolves? Why is the liquid not cloudy? single atoms or small groups of atoms become attached to water molecules-in other words, the soluble substance separates into particles much too small to see. dissolved particles are so small that they do not reflect light. In terms of Unit 10 of S102, the particles are very much smaller (by a factor of about 10-~) than the wavelength of light. This results in the light passing by them unaffected. In contrast, the insoluble particles of sand are very much bigger than the wavelength of light. They are able to reflect the light and hence they are visible.

ACTIVITY 2: SEPARATION OF A MIXTURE OF A SOLUBLE AND AN INSOLUBLE SUBSTANCE You will need:

cooking salt

sand

distilled water

funnel

filter paper

2 containers e.g. jam-jar

flat s'hallow dish

solution and sand

residue (sand)

filtrate (solution of salt)

(a) (b) RGURE 4 (a) A folded filter paper in a funnel; and (b) a filtration.

Shake about a teaspoonful of a 1 : 1 mixture of cooking salt and sand with about half a cupful of distilled water as described in Section 4.2. After shaking you will have a solution of salt and a residue of sand. Filtration through porous paper can be used to separate these. Although special filter papers which filter rapidly and efficiently can be bought, you can improvise with household items to find a substitute. Kitchen paper, folded into a cone often works well, as do ready-shaped coffee filter papers. Place your folded filter paper (or substitute) in a funnel, as shown in Figure 4a.


Support the funnel in the top of a jam jar and pour the sand/salt/water mixture into the filter paper, taking care that it does not overflow and run down between the paper and the funnel (see Figure 4b).

This will separate the undissolved solids, called the residue, from the solution, which is known as the filtrate. If the filtrate is still cloudy it is because the pores in the paper are too big to prevent some solid particles passing through. You would see the same effect if you tried to separate fine food particles from a sauce using a sieve with a coarse mesh.

To recover the solid sodium chloride place the filtrate in a flat shallow dish (e.g. a glass casserole lid) and leave it for several days until the water evaporates. You may be lucky and get some big crystals that clearly show the cubic shape of sodium chloride crystals or you may just get a lot of very small crystals.

This activity shows the basis of sea-salt production. If you live near the coast you could collect some seawater, filter it to remove suspended matter and leave the filtrate in a flat dish to evaporate off the water leaving crystals of salt. Alternatively (and probably much safer if the children are going to taste it) you could make some 'seawater' by dissolving about 4 g of salt in about l litre of water.

If you cheat a little and make it more concentrated you will get your crystals by

. evaporation more quickly.

ITQ 8 The technique you used in Activity 2 removes insoluble impurities from the seawater and hence from the salt. What do you think happened to the soluble impurities?

ITQ 9 Will the product be contaminated with impurities that are not soluble in water?

Some substances (sodium chloride is not one of them) are very much more soluble in hot water than in cold. This affords a method of crystal production without having to wait for the water to evaporate.

Substances that exhibit this property include potassium chromium sulphate (chrome alum, a purple salt), potassium aluminium sulphate, (potash alum, a colourless salt) and copper sulphate (a blue salt). (CAUTION-These salts are highly toxic and must not be tasted; they must be handled carefully and hands washed after use or gloves worn.) They can be easily obtained and produce spectacular crystals. (You can dispose of small quantities. down the sink or lavatory.) You may decide that you would rather demonstrate the technique to classes, though, as heating is involved.

ACTIVITY 3: GROWING CRYSTALS You will need:

saucepan of hot water

distilled water

2 Pyrex beakers or jugs

potassium chromium sulphate, potassium aluminium sulphate or copper sulphate

stirrer e.g. glass rod

You will need some way of keeping water hot in a transparent container. A Pyrex beaker or jug stood in a saucepan of hot water is suitable. Put about 50 cm3 (a normal size yoghurt pot is 150 cm3) of distilled water in the beaker and allow to reach approximately 60 "C. Then add about 30 g of one of the

above salts and stir with a glass rod to dissolve the salt whilst keeping the temperature at about 60 "C.

You will eventually find that no more of the solid will dissolve. You have produced a saturated solution, the most concentrated stable solution that is possible at that temperature. Remove the beaker from the heating bath and, while it is still hot, decant the solution into another beaker, leaving behind any undissolved solid. Set it aside to cool slowly and stand overnight. The slower it cools the bigger the crystals that will grow. If crystals do not form overnight try dropping a tiny crystal of the solid into the solution. This is called 'seeding' and provides a centre on which crystals can grow. If this does not work it is likely that your hot solution was not saturated, so reheat it and attempt to dissolve more solid into it. Decant and cool again. Retain your crystals for further experiments.

ITQ 10 Suppose the original substance that you have crystallized contained (i) soluble impurities, (ii) insoluble impurities. Where are the impurities after you have crystallized the substance?

If you, or your classes, have enjoyed crystal growing you may like to try two further experiments.

ACTIVITY 4: EXTENSIONS TO CRYSTAL GROWING

Growing a big crystal

You will need:

items as for Activity 2

Prepare a saturated solution of one of the three salts, as before, and allow it to cool to room temperature. Before it crystallizes, select the largest crystal that you grew previously and tie it to a length of cotton. Suspend it in the middle of your new solution and leave overnight.

Growing a silica garden

You will need:

water-glass (sodium silicate)

distilled W ater

glass container

stirrer e.g. glass rod

Make a 10% solution of water-glass (sodium silicate) in distilled water by stirring about 25 g water-glass with 225 cm3 water in a glass container until the thick syrupy liquid dissolves. (CAUTION-This solution is rather caustic so wash it off immediately if you get it on your skin.)

Wait until the solution is still and then add some of your largest crystals of the three salts. Do not stir. Plant-like growths will be seen growing from the crystals within a few minutes. The effect can be impressive, especially if left overnight or longer.

The principle on which this works is that the salts start to dissolve in the water, but then react with the sodium silicate to produce insoluble substances (copper silicate, chromium silicate or aluminium silicate). The crystal thus becomes encapsulated within a thin membrane. Water passes through the membrane by diffusion (osmosis) and the pressure inside the membrane eventually bursts it. You will see the 'plants' grow jerkily as this happens. Each time the membrane


bursts a jet of solution is fired into the sodium silicate solution. This reacts to produce more of the insoluble silicate in the new position. The randomness of the process results in the plant-like growths. Note that the whole process relies on the fact that sodium silicate and the three other salts are soluble, while the other silicates, e.g. copper silicate, are insoluble but permeable to water. If you have access to other salts you could try adding these to your garden.

4.3 SEPARATION BY CHROMATOGRAPHY There are many variations on the technique of chromatography, the best one for school use being paper chromatography using coloured substances. Another technique, gas chromatography, is used in the S102 TV programme 'Organic molecules in action'.

ACTIVITY 5: EXPERIMENTS WITH SMARTIES You will need:

filter paper

Smarties

distilled water

container

The colours of Smarties are created by adding food dyes to their outer layer. These colours can easily be removed with water. This forms the basis of a popular experiment.

Obtain some porous paper. Filter paper is definitely the best material, but if you are unable to get this try 10 cm diameter circles of white blotting paper. Cut the paper as shown in Figure 3a. Dip the edge of a Smartie in water. Remove it and wait a few seconds for the water to soften the Smartie's surface. Then gently press the wet edge of the Smartie on to the paper to leave a spot of colour, as shown in Figure 5a. Bend the strip down and place the circle of paper on a beaker of water with the strip dipping into the water as shown in Figure 5b.

FIGURE 5 Setting up a chromatograph experiment.

The water will soak up the strip and then spread radially when it reaches the horizontal paper.

If the Smartie was coloured by a mixture of dyes the water will carry the different dyes with it at different rates as it spreads, thus separating the chromatograph, see Figure 6.


blue band

yellow band

red band

extent of solvent spread

FIGURE 6 A chromotograph.

The principle behind this technique for separating a mixture relies on the fact that the individual dyes are attracted to both the water and the cellulose fibres in the paper but to differing degrees. As the water moves through the porous paper there is competition for the dye-the water is moving it, the paper is retaining it. As these attractive forces are different for different dyes they move at different rates. In the example shown in Figure 6b the blue dye is more attracted to water than is the red dye and hence moves further than the red dye in a given time. Do not let the wetted area reach the edge of the paper. If you do the water stops flowing outwards and colour separation ceases.

An interesting experiment can be done to see if a given colour is always produced by the same dye in the Smarties. For example, is the blue component of the dye mixture in brown Smarties the same dye as in blue Smarties?

Compare chromatographs of the dyes from two Smarties of different colour, that both contain blue dye. Figure 7 shows a typical result. Measure the distance ( X )

travelled by blue dye whilst the water travelled distance y. The value xly is called the retardation factor (Rf) for that particular dye. Repeat this measurement for the blue dye from the other Smartie (Figure 7b) and compare the two values of Rf. If they are the same it is likely that the dyes are the same. If the Rf values are different the dyes are not the same.

Experiments similar to those with Smartie dyes can be done with almost any water-soluble dyes. Felt-tipped pens give very intense colours. If you use small bottles of concentrated food dyes you may be able to identify what the dyes in Smarties actually are.

FIGURE 7 (a) Chromograph from a blue Smartie; (b) Chromograph from a brown Smartie.


You could link this topic with work on diet. Do the Smarties taste any different once the dye has been washed off? Can you tell the difference between the taste of Smarties of different colours if you do not know what the colour is? Could the dyes be harmful? Is any benefit gained by using food dyes? A useful source of information about the various dyes used in food is E for Additives by M . Hanssen (published by Thorsons, 1984).

4.4 SEPARATION BY DISTILLATION

4.4.1 SOLIDS, LIQUIDS AND GASES, THE THREE STATES OF MAllER Many substances, for example iron and air, are only familiar to us in one of their three possible physical states. On the other hand, we are familiar with water in all its three states-ice, water and steam. What happens on the molecular level as a lump of ice is heated so that it melts to water and then boils to steam?

In ice, the water molecules are arranged regularly in a crystal lattice and are held together by electrostatic forces called hydrogen bonds (see Section 6) . Because the molecules are close together, these hydrogen bonds are strong enough to give a rigid structure, in other words it is solid.

Figure 8 illustrates how this changes when the ice melts. When the ice is heated, increased molecular vibrations occur.

Its volume does not change much, but if the temperature of the ice reaches 0 "C the vibrations are sufficient for enough hydrogen bonds to be broken for all the molecules to be free of the rigid lattice. The substance therefore loses its own shape and takes on the shape of the containing vessel. We say it has melted.

FIGURE 8 (a) An ice cube has shape and volume; (b) water has volume but no shape; (c) gas has no shape and no separate volume.

If the liquid water is now heated further, more hydrogen bonds are broken. Eventually, at 100°C, so many have broken that the molecules not only adopt the shape of the container but also its volume. The substance is now gaseous. In an open vessel the gas will escape. (CAUTION-Never attempt to boil a liquid in a closed vessel-it is likely to explode.)

The temperatures at which these processes take place are called the melting point and the boiling point of the substance. Each substance has its own value for each of these.

4.4.2 SIMPLE DISTILLATION Going back to our solution of sodium chloride in water (Section 4.2). you may now be able to see how we can separate and collect the water from the mixture. We simply need to boil the solution and condense the steam back to water; the


sodium chloride will be left behind because its boiling point is very much higher than that of water.

You can carry out a similar procedure quite simply by holding a cold dinner plate at an angle in the steam coming from an electric kettle, as shown in Figure 9. Tap water is a solution of one or more minerals (which ones depends on where you live); when the water is boiled the minerals remain in the kettle so the distilled water collected is much purer. (CAUTION-It is advisable to hold the dinner plate in a clamp to avoid scalding your hands in the steam. You should not allow children to do this experiment themselves.)

It is interesting to compare the taste of tapwater and distilled water. The absence of minerals in distilled water results in a completely uninteresting taste often described as 'flat'. What do you think the work 'pure' means on a bottle labelled 'Pure Mineral Water'? It certainly does not mean chemically 'pure'.

the distillate - (distilled water)

tap water FIGURE 9 Distillation of tap water.

This principle of boiling and condensing is widely used for separating and purifying liquids. Where the mixture contains more than one liquid (e.g. crude oil) each one boils at a different temperature and the procedure is known as fractional distillation. This is an important industrial process, which is used in the refining of crude oil into its many useful components.

An interesting teaching extension to this topic, which links to work on weather (AT9), is to relate the process of boiling and condensation to the formation of rain. Given that water can slowly evaporate to steam at less than 100 'C, it is easy to make the analogy between the kettle and the sea. The clouds of steam represent the clouds in the sky and the cold dinner plate the effect of cooling the clouds at high altitude. (This is covered in detail in the Study Commentary for Unit 27.)

The build up of dissolved salts in the sea is analogous to the deposits left in the kettle. The sea-clouds-rain-river-sea cycle (fuelled by the Sun) is really no more than a big distillation apparatus.

Misting of windows and 'steamy breath' in cold weather are also caused by condensation.

4.4.3 SEPARATION BY FREEZING The compounds that give a drink such as apple juice its characteristic flavour have much lower melting points than water. If apple juice is placed in a plastic beaker (glass or china will crack) in a deep freeze or the ice compartment of a refrigerator, crystals of ice will form. If most of the liquid is frozen the remaining syrup contains the flavours in a concentrated form and the liquid tastes very strongly of apples. This is an experiment that should appeal to children.


SUMMARY OF SECTION 4 1 Some substances dissolve in a given solvent, such as water, others do not. Soluble substances can be separated from insoluble substances by filtering the solution. There is an upper limit to the amount of a soluble substance that can be dissolved in a given amount of a given solvent. A solution containing this amount of solute is called a saturated solution. Attempts to exceed this limit result in crystallization. This can be used to purify substances.

2 Substances in solution can be separated from each other and sometimes identified using paper chromatography. This process depends on the different attractive forces exerted on the substances by water and paper.

3 Each liquid has a temperature at which it boils. This can be used to separate liquids from dissolved solids (simple distillation) or from other liquids (fractional distillation).

5 ATOMIC STRUCTURE

A SIMPLE ATOMIC MODEL So far we have used the model of the atom as a solid sphere, like a tiny billiard ball. This has been sufficient to explain simple phenomena and to describe elements, compounds and mixtures. There has been no need to worry about the internal structure of atoms. However, such a simple model is not detailed enough to describe how atoms join together.

Scientists have not yet been able to look inside atoms to see what they are like, but the results of lots of experiments have allowed us to develop a good working model of atomic structure. Scientific models are not necessarily the 'truth' or the 'right answer'. They owe their credibility to their success at explaining our observations of the physical world and guiding our thinking in productive directions.

Many chemical and nuclear processes can be explained by a model in which atoms are thought to be made of three types of sub-atomic particle: protons, neutrons and electrons. Protons and neutrons form the nucleus (or centre) of atoms. Electrons are arranged around the nucleus in a way that will be partly explained soon.

Figure 10 illustrates 'this simple model.

FIGURE 10 A simple model of the structure of an atom. Not to scale.

The nucleus is tiny compared with the volume occupied by the electrons. If you imagined the atom to be the size of Wembley Stadium, the nucleus would be the size of a golf ball on the centre spot! The radius of most atoms is between 0.1 and 0.2nm ( l X 10-l0 to 2 X 10-''m), which means the nucleus must be incredibly small-about 10-l5 m radius, in fact.

5.2 SUB-ATOMIC PARTICLES Some of the properties of protons, neutrons and electrons are summarized in Table 4.

0 What relationship can you see between the charges of protons, neutrons and electrons?

W Protons and electrons have equal but opposite electrical charges. Neutrons are electrically neutral.

0 What relationship can you see between the masses of protons, neutrons and electrons?

W Protons and neutrons have equal masses, and are much more massive than electrons. (Actually the masses of protons and neutrons are very slightly different.)

TABLE 4 Some properties of sub-atomic particles

Particle Mass on relative Relative charge Location atomic mass scale

proton 1 +l in nucleus

neutron 1 0 in nucleus

electron 0.00055 -1 around nucleus

D How do the mass and size of the nucleus compare with the mass and size of the whole atom? The nucleus accounts for almost all the mass of the atom and hardly any of its volume-it is as if there is almost no mass in most of the atom.

0 For an atom to be neutral, that is, to have no overall charge, what is the relationship between the number of electrons and protons?

W For atoms to be electrically neutral there must be an equal number of electrons around the nucleus.

If you want to know more about the background to our present model of atomic structure, you can read Section 3 (pp. 12-18) of Units 11-12 of S102.

5.3 NUCLEAR SYMBOLS In the simple model of the atom illustrated in Figure 10, different atoms can be distinguished by just two numbers-the atomic number (symbol Z) and the mass number (symbol A). The atomic number is the number of protons in the nucleus. For atoms to be electrically neutral there must be an equal number of electrons around the nucleus. Atomic number is the same for every atom of an element: for example, Z = 6 for all carbon atoms because carbon has six protons.

The mass number is the number of protons plus neutrons in the nucleus. If the number of neutrons is given the symbol N, then

A = Z + N


5.4 WHAT ARE ISOTOPES? Whereas the atomic number is the same for every atom of an element, the mass number is not. Atoms of the same element that have different masses are called isotopes. Since the number of protons is the same for all atoms of an element, the differences in mass must arise from different numbers of neutrons.

We denote different isotopes in the following way: %e, ;iNe, ;;Ne, where the upper number is the mass number and the lower number is the atomic number. Sometimes the atomic number is omitted (20~e ) , since the chemical symbol tells us which element it is anyway. The symbol 2 0 ~ e is pronounced 'neon-20'.

ITQ 1 1 Fill in the missing details on Table 5.

TABLE 5 For use with ITQ 11

Isotope Symbol Atomic number Mass number Number of neutrons

nitrogen- 14 14 7

Most elements exist as a mixture of isotopes. The relative atomic mass is an average of the masses of the isotopes which also takes account of the proportions of isotopes present. This is why relative atomic masses are seldom integral (whole) numbers. The isotopes of chlorine and bromine are given in Table 6, together with their percentage abundances in a naturally occurring sample.

0 How many protons, neutrons and electrons are present in atoms of: (a) 7 9 ~ r (b) * ' ~ r (a) 35 protons, 44 neutrons, 35 electrons.

(b) 35 protons, 46 neutrons, 35 electrons.

TABLE 6 Isotopes of chlorine and bromine

Element Isotope Abundance Element Isotope Abundance

35 ,,c1 75% 79 35Br 50% chlorine bromine

37 ,,C1 25% B 50%

If you looked at the early part of Section 2 of Units 11-12 of S102 when you were reading about relative atomic masses, you might now like to look at Section 2.2.2, which gives more background information about isotopes.

Although most of the mass of an atom is concentrated in its nucleus, the electrons are far from insignificant. When atoms combine chemically, it is the electrons on the outside that interact with one another. So chemists need to know much more about electrons than about the nuclei of atoms.


5.5 HOW ARE ELECTRONS ARRANGED IN ATOMS? The basis of our present understanding of the electronic structure of atoms is a theory proposed by the Danish scientist Niels Bohr. Bohr was trying to explain some observations of the light given out when hydrogen gas is made to glow, rather like the sodium vapour in a yellow street lamp.

To understand how electrons interact with each other, Figure 10 (page 23), is too simple a representation of the atom. A better model, using a sodium atom as an example, is shown in Figure 11. Remember that this is only a two-dimensional representation. In this model, electron shells surround the nucleus, so in three dimensions the shells in sodium would be more like a squash ball inside a tennis ball inside a football.

The number of electrons arranged in shells is called the electron shell con$guration of an atom. Table 7 shows the electron shell conjgurations for the first 20 elements.

Remember that a model can help us understand the behaviour of electrons but may not represent the 'real' situation that exists in an atom. This latest model is still a greatly simplified view of the atom, but is sufficient to help our understanding of the behaviour of the electrons that take part in chemical reactions.

TABLE 7 Electron shell configurations for the first 20 elements

Element Atomic 1st shell 2nd shell 3rd shell 4th shell number

hydrogen

.helium

lithium

beryllium

boron

carbon

nitrogen

oxygen

fluorine

neon

sodium

magnesium

aluminium

silicon

phosphorus

sulphur

chlorine

argon

potassium

calcium


shell 3 containing l electron

shell 2 containing 2 electrons

. shell l containing 8 electrons

FIGURE 1 1 Electron shells in a sodium atom.

Much chemistry is determined only by the outer-shell electrons, and chemists can explain many of the properties of atoms without needing to use a detailed theory of atomic structure. One very useful model regards the atom as composed of a core of nucleus plus inner shells surrounded by an outer shell.

Figure 12 shows how a sodium atom would be thought of using this model. Notice that it retains some of the simplest features of Figure 10. Only the outer shell is looked at in the detail of Figure 11.

outer shell ron

core of shell 1

/

' nucleus, and shell 2

FIGURE 12 Core and outer-shell model for a sodium atom.

0 Look at Table 8. What appears to be the maximum number of electrons that can occupy: (a) shell 1 (b) shell 2 (c) shell 3?

However, the filling of shell 3 is not straightforward. In fact, it can hold a maximum of 18 electrons but the other 10 are not put in until the Jirst two electrons have been added to shell 4. This is a consequence of the more detailed inner structure of electron shells. This detail need not concern us here.


0 There seems to be a pattern in the way electrons are put into shells. What is it? When an atom has several electrons, they are arranged so that shell 1 must always have 2 electrons before any electrons occupy shell 2. Then shell 2 must have 8 electrons before any electrons occupy shell 3. Then shell 3 must have 8 electrons before any electrons occupy shell 4.

5.6 THE PERIODIC TABLE There is an intricate relationship between the number of electrons in the outermost shell of an atom and its chemical behaviour. Elements with the same number of outermost electrons show similarities in their chemical properties and can be grouped together in the arrangement known as the Periodic Table of the elements.

A copy of the Periodic Table is to be found on the back cover of S102 Units 11 to 18. You will see that the elements are arranged from left to right in order of increasing atomic number.

Find the elements lithium (Li), sodium (Na) and potassium (K) in the Periodic Table. Now look at Table 8 and note how many electrons each of these elements has in its outermost shell. You should have found that Li, Na and K each have 1 electron in the outermost shell.

Now find the elements fluorine (F) and chlorine (Cl). How many outermost electrons do they have? 7.

In the Periodic Table, the roman numeral at the top of each column of elements indicates the number of electrons in the outermost shell of the atom (with the exception of Group 0, which has 8 electrons). Vertical columns of elements are known as Groups, and are numbered 0 to VII. The members of a Group show similarities in their chemical properties.

Horizontal rows of elements are known as Periods, for example, the elements from sodium (Na) across to argon (Ar).

The way in which atoms join together is determined by the outer-shell electrons, and this will be developed in Section 6.

Elements are classified as metals, semi-metals and non-metals (see Figure 13 of Units 13-14 of S 102).

SUMMARY OF SECTION 5 1 Atoms are composed of three types of sub-atomic particles: protons, neutrons and electrons. The protons and neutrons form the nucleus-the tiny dense centre of the atom. The electrons are arranged around the nucleus.

2 Atoms of the same element contain the same number of protons. If they differ in the number of neutrons, the atoms of different mass are called isotopes.

3 The electrons are arranged around the nucleus in electron shells.

4 There are definite rules about the maximum number of electrons that can be held in any shell.

5 Most chemistry, including how atoms bond together, is determined by the outer shell electrons only.


6 CHEMICAL BONDING

6.1 NOBLE GAS ELECTRON CONFIGURATIONS In Section 5.5 we looked at the ways electrons are arranged in atoms. There is s?mething especially stable about electron configurations that place two electrons in shell 1 and eight electrons in shells 2 or 3. The elements with these outer-shell arrangements are very unreactive: they are the noble gases helium, neon and argon.

Their existence as single atoms results from their reluctance to combine chemically with other elements or themselves. Elements with one electron more than a noble gas have one electron in their outermost shell-that is, the next shell starts to fill.

This observation is the starting-point for our theory of chemical bonding. We assume that when atoms react their atoms try to achieve a noble gas electron configuration by losing or gaining an electron or electrons.

IONIC BONDING' The metal atoms in Groups I and I1 of the Periodic Table have few outer-shell electrons. An noble gas configuration can be reached if these electrons are lost to form positively charged ions. For example:

where Na+ represents a positively charged sodium ion and e- an electron. A positively charged ion is known as a cation.

Most non-metal atoms have more than three outer shell electrons. One way they can reach an inert gas configuration is by gaining electrons to form anions- negatively charged ions. For example:

C1 + e- + Cl-

There are limits to how many electrons an atom can gain or lose. Anions with a charge of 3- are unusual. Cations with a charge greater than 3+ are almost unknown.

0 Refer to Table 8 to find the electron configurations of the following ions? To which noble gas electron configurations do they correspond? (a) Na+

' (b) ca2+ (C) ~ 1 ~ +

(6) s2- (e) N3

(a) 2.8, neon

(b) 2.8.8,argon (c) 2.8, neon (d) 2.8.8, argon (e) 2.8, neon

When metals react with non-metals, noble gas configurations can be achieved if electrons are transferred from the metal atoms to the non-metal atoms.

We can summarize the way atoms bond together by drawing ~ e w i s structures, sometimes known as dot-cross diagrams. In these, the outer-shell electrons of one atom are represented by dots and those of the other atom by crosses, and the electrons' are grouped in pairs. Figure 13 shows Lewis structures for the formation of potassium fluoride and magnesium chloride. The dots and crosses


represent only the outer-shell electrons. The numbers underneath the chemical symbols show the shell configurations for all the electrons. The square brackets indicate that the structure is an ion.

T X X

K * + . F : - [K] X X X X

(2,8,8,1) ( ~ 7 ) (2,8,8) (229 potassium fluorine potassium fluoride

atom atom ion ion

(2,821 (2,8,7) (2,8,7) (2,8) (2,8,7) (2,8,7) magnesium two chlorine atoms magnesium two chloride ions

atom ion FIGURE 13 Lewis structures for two ionic compounds.

What electronic changes occur when potassium fluoride is formed?

Each potassium atom loses one electron and each fluorine atom gains one electron, so the compound formed has a formula KF.

What electronic changes occur when magnesium chloride is formed?

Each magnesium atom loses two electrons but each chlorine atom gains only one electron, so the formula for magnesium chloride is MgC12.

usually think of ions as being spheres with electrical charge spread evenly around them. A cation can therefore attract anions from any direciion and vick versa.

The ions in ionic compounds are arranged in lattices-very large numbers of cations and anions in fixed positions in a regular pattern. The pattern found in the sodium chloride lattice is shown in Figure 14.

FIGURE 14 The NaCl lattice.

Any sodium ion in this structure will be attracted to the Cl- ions that are around it. It will be repelled by other Na+ ions which are a bit further away, but attracted by some more Cl- ions, which are even further away, and so on.

The overall result of all these attractions and repulsions is what holds the ionic compound together. So the ionic bond is a complicated bond. It is not, for example, simply the attraction between single Na' and Cl- ions. Because ionic forces are strong and hold so many particles together, ionic compounds are usually solids with high melting points.


Consider what happens when calcium reacts with chlorine-a Ca atom loses two electrons to form a ca2+ ion. Two atoms of chlorine each gain an electron.

0 Draw a Lewis structure to represent CaC12. B

6.3 COVALENT BONDING Ionic compounds are usually formed when metals react with non-metals. But there are many compounds containing only non-metallic elements. These cannot be bonded ionically because some of the atoms would have to lose too many electrons. Instead, inert gas configurations are achieved by sharing electrons. Shared electrons count as part of the outer shell of both atoms in the bond. The Lewis structure for the H2 molecule is shown in Figure 15.

FIGURE 15 Electron sharing in the hydrogen molecule.

Bonds formed by shared electrons are called covalent bonds. If a pair of electrons is involved, the bond is called a single covalent bond, or more simply a single bond.

Examples of Lewis structures for two more covalent compounds are shown in Figure 16.

FIGURE 16 Lewis structures for NH, and H20

Electron pairs that form bonds are called bonding pairs. Pairs of electrons not involved in bonding are called non-bonded pairs (sometimes known as lone pairs). Both water (H20) and ammonia (NH3) have lone-pair electrons.

When two pairs of electrons form a covalent bond, it is called a double bond: The bonds in molecular oxygen and carbon dioxide are double covalent bonds, (see Figure 17a).

When three pairs of electrons form a bond, it is called a triple bond. The bond in molecular nitrogen (N2) is an example (see Figure 17b).

molecular carbon molecular oxygen, O2 dioxide, CO2 nitrogen, N,

(a) (b)

FIGURE 17 (a) Double covalent bonds. (b) A triple bond.

Lewis structures are useful for representing individual electrons in chemical bonds. Chemists often draw structures in a simpler way by using lines to represent a pair of electrons shared between two atoms. A single line represents a single covalent bond. Double and triple lines represent double and triple covalent bonds.


The examples used earlier are shown in this way in Figure 18.

H I

N H-H O=O H'- \ H

molecular molecular ammonia hydrogen oxygen

water molecular carbon dioxide nitrogen

FIGURE 18 Chemical structures showing bonds.

0 Explain what you understand by the term molecule. A molecule is a group of atoms held together by covalent bonds. All the molecules of a pure substance must be identical: they must contain the same number of the same atoms bonded together in the same arrangement. A molecular substance will contain a very large number of these identical molecules.

ITQ 12 Draw Lewis structures for the following covalent molecules.

(a) F2, molecular fluorine

(b) pH3, phosphine

(c) H2S, hydrogen sulphide

(d) CS2, carbon disulphide

(e) HCN, hydrogen cyanide

6.4 ELECTRONEGATIVITY When atoms of chlorine combine with each other, gaseous covalent Cl2 molecules are formed. But when chlorine atoms combine with sodium atoms, the ionic solid NaCl is produced. Figure 19 shows Lewis structures for these two substances.

gaseous chlorine ionic solid, molecules, Cl, NaCl

FIGURE 19 Lewis structures for Cl, and NaCl.

For both these structures the outer electrons are grouped in pairs and, in both structures, the formation of a chemical bond involves the formation of a new electron pair in the outer shell of chlorine. In Cl,, because the two atoms are identical, the electron pair must be equally shared between the two atoms. In NaC1, by contrast, it resides completely on the chloride ion. According to this picture, ionic and covalent bonding are just different aspects of a common process, because both involve the formation of electron pairs: the difference between them lies only in the extent to which those electron pairs are shared between atoms.

Understanding ionic and covalent bonding can be improved using the concept of electronegativity. The electronegativity of an element is a measure of the power of an atom of an element to attract electrons to itself when it is bonded to


another atom. In the Cl2 molecule, the two atoms are identical. They have equal electronegativities, so the electron pair is shared equally between them. Now consider sodium chloride.

0 Which atom is the more electronegative, sodium or chlorine? Chlorine: in sodium chloride, the electron pair has been completely taken over by the chlorine which forms a chloride ion.

/

Table 8 lists some electronegativity values that were calculated by Linus Pauling, a US chemist and winner of the Nobel Prize for chemistry in 1954.

TABLE 8 Pauling electronegativity values for some common elements

Atom Electronegativity

Refer to Table 8 and the Periodic Table on the back cover of S102 Units l l to

What change in electronegativity occurs as one moves down Group V11 from fluorine to iodine? The electronegativity values decrease down the Group.

Do other Groups show the same pattern in electronegativity values? The sake pattern is observed for 0 and S, and for N and P.

What change in electronegativity occurs as one moves across a Period, e.g. from Na to Cl? Electronegativity increases from left to right across a Period.

Where in the Periodic Table are the most electronegative elements? At the top right-hand corner-ignoring the noble gases helium, neon and argon, which form no compounds. The most electronegative elements are first fluorine, followed by oxygen, then nitrogen and chlorine.

The trends in electronegativity can help us predict whether one substance is more likely to be ionic than another. For example, sodium and hydrogen both form chlorides: NaCl and HCl.

0 In which of the two compounds is there a large electronegativity difference? In NaC1.


0 Which of the two compounds is more likely to be ionic? NaCl: the larger difference in electronegativity means that the chlorine is more likely to win control of electrons and form chloride ions.

But what happens in a compound like HC1 which is covalent but where there is a small difference between the electronegativities of the elements?

The situation is part way between the extreme forms of bonding in Cl2 (which is purely covalent) and NaCl (which is fully ionic). The H electron is not fully acquired by the C1 atom: the sharing is unequal, and the electron pair is attracted more towards the C1 than the H.

The situation can be represented by saying that in HC1 molecules the C1 carries a partial negative charge and the H carries an equal partial positive charge. These partial charges are symbolized as follows: 6+ 6-

H-C1

In general, different atoms will attract bonding electrons unequally. One atom will acquire a partial negative charge because it has a greater share of the bonding electrons. The other atom will become partially positively charged because it has lost some of its share in the bonding electrons. Bonds like this are called polar bonds.

Polar bonds are like covalent bonds with a bit of ionic character in them. The ionic and covalent models are extreme forms of bonding: polar bonds are somewhere between the two (see Table 9, overleaf).

d

TABLE 9 Characteristics of bonding

Ionic bonding Polar covalent bonding Covalent bonding

electrons lost by one atom electrons shared electrons shared equally and gained by the other unequally-one atom between two atoms

makes a partial gain of the electrons

Oxygen, like chlorine, is one of the most electronegative elements. So oxygen in H20, like chlorine in HC1, carries a partial negative charge. The resulting charge distribution is shown in Figure 20. Note that because there are two hydrogen atoms, and because the total charge must be zero; the partial charge of oxygen is doubled.

FIGURE 20 Charge distribution of an H 2 0 molecule.

If you wish to take your study of chemical bonding a little further, you will find Section 7 of Units 13-14 of S 102 useful.

INTERMOLECULAR FORCES Whether a substance is a solid, liquid or gas is determined by the average energy of the molecules (see also 'Chemical reactions and energy changes') and the cohesive forces between the molecules. These cohesive forces can arise in a number of ways, and are known collectively as intermolecular forces.

Consider the series of molecules made from elements in Group V11 of the Periodic Table. Fluorine (F2) and chlorine (Cl2) are gases at room temperature,


bromine (Br,) is a liquid and iodine (I2) is a solid. If cooled to sufficiently low temperatures, chlorine and fluorine will liquify and eventually solidify. -There are attractive forces between the molecules of all four elements, but what are these forces? They are electrical in origin, although all form molecules that are electrically neutral and both atoms making up each molecule have an equal share of the bonding electrons.

These forces, called London forces, arise because although on average the electrons in each molecule are evenly distributed, at any instant their distribution is not uniform (see Figure 21).

FIGURE 21 The origin of London forces: (a) shows a molecule made up of two atoms with the electrons (denoted by shading) distributed evenly around the two positive nuclei; (b) show.how the electrons can transiently be distributed unevenly. Though on average the distribution of charge is even, the molecule can none the less have a temporary dipole, that is, it can be slightly positive at one end and slightly negative at the other. The transient dipole in one molecule can then be attracted by the transient dipole in an adjacent molecule, (c). The result is a net attractive force between the molecules (dashed lines).

This results in a momentary separation of positive and negative charge giving rise to an electric dipole. These transient dipoles give rise to short-range attractive forces between molecules.

The size of the London forces depends on the relative molecular mass of the molecule: the higher the relative molecular mass, the greater the attraction between molecules.

Therefore, iodine molecules are more strongly attracted to one another than are chlorine molecules.

The sequence of boiling temperatures is:

London forces act between all molecules, but for some molecules, there are additional attractive forces. Hydrogen chloride (HC1) has a relative molecular mass of 36.5 and that of fluorine (F2) is 38. If London forces alone were operating these two molecules might be expected to have rather similar boiling temperatures. However, hydrogen chloride boils at -85 "C. Compare this with the boiling temperature of fluorine. Additional electrical forces operate between hydrogen chloride molecules.


As you saw in Section 6.4, molecules of HC1 are polar. Because the molecule has two charges, negative at one end and positive at the other, we say it possesses a dipole. It is a permanent dipole: the partial charges are always there in the same atoms because of the electronegativity difference.

When a large number of HC1 molecules are put together, as in a sample of gas, the molecules will be attracted to one another because the 6+ charges on some molecules interact with the 6 charges on others.

We call these interactions dipole-dipole forces. They are weak forces-much weaker than the forces between fully charged ions-so, at room temperature, hydrogen chloride molecules are not locked together into a crystalline solid. Dipole-dipole forces will exist whenever a substance contains dipolar molecules.

A particularly strong and special form of a dipole-dipole force arises when the partial charges are quite large, and the interacting atoms are small and able to approach each other closely. This situation arises whenever hydrogen atoms are bonded to the small and highly electronegative N, 0 and F atoms. The example of hydrogen fluoride is illustrated in Figure 22.

The partial positive charge on a hydrogen atom bonds to the partial negative charge on a fluorine atom to form a hydrogen bond. These are represented by the dashed lines in Figure 22.

FIGURE 22 Hydrogen bonding between HF molecules.

Hydrogen bonding is present in water, and it is responsible for the structural order of H20 molecules in ice.

Hydrogen atoms are bonded to oxygen and nitrogen atoms in many naturally occumng molecules. Hydrogen bonds are found extensively in nature, forming important links in carbohydrates, proteins, DNA and other molecules that are vital for life. You will learn more about hydrogen bonding in Units 22 and 24 of S102.

ITQ 13 Which of the bonds in the following list will be polar? In cases where there is a polar bond, show which atom is positive and which is negative using the 6+, 6 convention.

C-F C-H C-S H-C1 H-N S-Br C - 0

SUMMARY OF SECTION 6 1 Elementary bonding theory describes how atoms bond together in such a way as to achieve a noble gas electron configuration.

2 In ionic bonding, electrons are transferred from atoms of one type to atoms of another type. Cations carrying a positive charge and anions carrying a negative charge result.

3 In covalent bonding, electrons are shared between atoms. Neutral molecules result. A shared pair of electrons is called a single bond, two shared pairs constitute a double bond.

4 Electronegativity is a measure of the power of an atom in a molecule to attract bonding electrons to itself.

5 Ionic and covalent bonding can be thought of as extreme models of bonding. If there is a large electronegativity difference, the bond is ionic; if the difference is small, the bond is covalent.

6 Intermediate electronegativity differences result in polar covalent bonds.


7 In samples of molecular substances, atoms are held together in the molecules by interactions between molecules as well as by covalent bonds. These interactions are called intermolecular forces. There are three types: London forces, dipole-dipole forces and hydrogen bonds.

8 London forces act between all molecules-whether they possess a permanent dipole or not. They are weak but they are also vital for the proper functioning of many of the chemicals that are essential for life.

9 Dipole-dipole forces act between the permanent dipoles on molecules.

10 The hydrogen bond is a particularly strong dipole-dipole force. It occurs whenever a hydrogen atom is bonded to an atom of nitrogen, oxygen or fluorine. Hydrogen bonds are very important in naturally occurring molecular substances.

ITQ ANSWERS AND COMMENTS ITQ 1 (a) C; (b) A; (c) B.

Elements are the basic building blocks of matter and contain only one type of atom. They cannot be further broken down chemically into more basic parts. A compound is a substance made up of elements in fixed proportions. It is not a mixture of the elements but a new substance with new, often very different, properties.

lTQ 2 (a) B; (b) A.

I T Q 3 (a) Soil is a mixture. It contains many components in variable amounts.

(b) Carbon dioxide is a compound. It contains two elements and in the ratio of two atoms of oxygen to each atom of carbon,

(c) Gold is an element. It contains only one kind of atom.

(d) Tap water is a mixture. The water may contain other substances, such as those minerals that make it 'hard' (causing fur or scale in kettles) and chlorine, which is added to kill harmful bacteria.

ITQ 4 (a) 24/12 = 2; (b) 80140 = 2; (c) 64/32 = 2; (d) 1411 = 14; (e) 64/16 = 4.

ITQ 5 SiH4. If the valency of silicon is four, one atom of silicon combines with four atoms of hydrogen.

ITQ 6 SO,. The valency of oxygen is two, so as the valency of sulphur is six, the ratio of sulphur atoms to oxygen atoms is 1 : 3.

ITQ 7 (a) CO2 (b) CC14.

The valency of carbon in methane is equal to the number of hydrogen atoms that combine with one atom of carbon, which is four. Taking the valency of oxygen as two, we get the formula CO2 by setting the valency of carbon equal to twice the number of oxygen atoms that combine with one atom of carbon.

ITQ 8 They are present in the solid sodium chloride crystals that are left when the water evaporates. (If you use genuine seawater the crystals will become sticky on exposure to the air. They contain some magnesium chloride, which absorbs moisture from the air and becomes sticky.)

ITQ 9 No. They will have been removed when you filtered the seawater.

ITQ 10 (a) In the solution; (b) In the solid product.

(Insoluble impurities can be removed by filtering the solution while it is still hot before cooling to crystallize. It may crystallize in the filter paper unless you can keep it hot whilst it filters.)


ITQ 1 1

Isotope Symbol Atomic number Mass number Number of neutrons

ITQ 12

ITQ 13 S+ S- S+ S- C -F H-N

S+ S- C-S

The C-H bond is not polar.

NOTES

introducing chemistry concepts chemical reactions … · 3.1 recognizing chemical changes ......

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