Matter

Form 1 Chemistry

Materials can be solid, liquid or gaseous, depending upon the arrangement and freedom of movement of

these particles.

5.1. THE CONCEPT OF MATTER

Matter is anything that has mass and occupies space. Therefore, anything around us provided it has mass and can occupy the space, is termed as matter. There are many kinds of matter. Can you mention some? The word matter is used to cover all the substances and materials from which the earth and universe is

composed of. These include all materials around us such as water, soil, plants, animals, air, clothes, etc.

Any particular kind of matter is called a substance. Substances include elements and compounds. An element is a substance which is the limit of chemical analysis. When two or more elements are combined chemically, a compound is formed. Matter is made up of atoms, ions or molecules. You will learn more about this later.

5.2. STATES OF MATTER

Any chemical substance we study exists in any of the three forms (or physical states). The three different states of matter are solid, liquid and gaseous states. So, each of the many millions of substances around us can be classified as a solid, a liquid or gas. Look around you and name substances that are solids, liquids and gases. The state in which any matter exists depends on temperature and sometimes pressure conditions. One substance may exist as a solid in one condition and as a liquid or gas under a different

condition. Water is an example of such substances. This change is called a change in the state of matter.

The three physical states of matter differ in the way they respond to temperature and pressure. All three states can increase in volume (expansion) when the temperature is increased. They decrease in volume (contraction) when the temperature is decreased. Gases are easily compressed. Liquids are only slightly

compressible. Solids are incompressible. They are not affected by change in pressure.

Experiment 5.1. Investigation of the compressibility of solids, liquids and gases

Procedure

1. Take three new syringes and fill them with sand, water and air respectively (figure 5.1).

2. Try to push in the end of each syringe.

3. Observe what happens.

MCHORO

Figure 5.1. Compressibility of solids, liquids and gases

Observation

Which of the substances under investigation can compress into a smaller volume?

Findings

You should have found that a solid (sand) and a liquid (water) cannot be compressed but a gas (air) is

easily compressed.

The three states of matter differ in their physical properties. These differences in properties are

summarized in table 5.1.

Table 5.1. Differences in properties of the three states of matter

Property Physical state

Solid Liquid Gas

Shape has a definite no definite shape, takes no definite shape, occupy

shape shape of the container whole container

Volume

has a fixed volume

has a fixed volume variable (depending on temperature and pressure)

Fluidity does not flow generally flows easily flows easily

Expansion on heating

low medium high

Compressibility incompressible almost incompressible highly compressible

Motion of particles

slow high very high

Density high moderate to high low

Tangibility tangible tangible intangible

Visibility visible visible invisible

5.2.1 Change in states of matter

We have seen that matter exists in three different states - solids, liquids and gases. We can use the kinetic theory of matter to explain how a substance changes from one state to another. Basically, changes from one state to another are caused by alterations in temperature and pressure. Normally molecules, ions or

atoms of a substance move faster when the temperature is increased.

Melting and freezing

Melting is a change from solid to liquid state. When solids are heated, their constituent particles (atoms, molecules or ions) get energy and vibrate more violently. Vibrations of these particles overcome (exceed) their binding forces. The particles become mobile. The crystalline structure of solid is destroyed. A liquid state is reached and the particles are free to move. The temperature at which this happens is calledmelting

point of the solid.

Figure 5.2. Change in state from solid to liquid

The melting point of a solid tells us something about the strength of forces holding its constituent particles 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 a change from liquid to solid state. Freezing is the opposite of melting. The process is reversed at the same temperature if a liquid is cooled. The temperature at which a substance turns to a solid is called freezing point. The melting point and freezing point of any given substance are both the same. For example, the melting and freezing of pure water takes place at 0C. Melting is not affected by any

changes in atmospheric pressure.

Evaporation and boiling

Boiling is a change from liquid to vapour state at a particular temperature. Evaporation is the change

from liquid to vapour state at any given temperature.

If a liquid is exposed to open air, it evaporates. Splashes of water evaporate at room temperature. After rain, small pools of water dry up. When a liquid changes into a gas at any temperature, the process is called evaporation. Evaporation takes places from the surface of the liquid. The larger the surface area, the

faster the liquid evaporates. The warmer the liquid is, the faster it evaporates. Thus, surface area and

temperature affects the rate of evaporation of a liquid.

When a liquid is heated, its molecules get more energy and move faster. They knock into each other violently and bounce further apart. As the heating goes on, its molecules vibrate even faster. Bubbles of gas (due to air dissolved in water) appear inside the liquid. The whole process is called boiling. The

temperature at which a liquid boils is called boiling point.

Figure 5.3. Change in state from liquid to gas

The molecules at the surface of the liquid gain enough energy to overcome the forces holding them together. They break away from the liquid and from a gas (vapour). As more of the liquid molecules

escape to form a gas, a liquid is said to evaporate. This occurs at the boiling point of a liquid.

The temperature at which a liquid boils explains how strong the forces holding its particles (molecules) together are. Liquids with high boiling points have strong forces of attraction between their molecules

than those liquids with low boiling points.

The boiling point of a liquid can change if the surrounding pressure changes. If the surrounding pressure falls, the boiling point also falls. The boiling point of water at standard pressure (760 mmHg) is 100C. On a high mountain, where pressure is low, it is lower than 100C. If the surrounding pressure is increased,

the boiling point rises. The same behaviour is experienced by a gas when the pressure is either increased

or decreased.

Table 5.2. The melting and boiling points of some common chemical substances at standard temperature and pressure (s.t.p)

Substance Physical state at room temperature (20C)

Melting point (C)

Boiling point(C)

Oxygen gas -219 -183

Nitrogen gas -210 -196

Ethanol (alcohol) liquid -117 78

Water liquid 0 100

Sulphur solid 115 444

Common salt (sodium chloride)

solid 801 1465

Copper solid 1083 2600

Carbon dioxide gas

sublimation point (C):

-78

From the above explanation, obvious differences between evaporation and boiling can be detected. See

table 5.3.

Table 5.3. Differences between evaporation and boiling

Evaporation Boiling

1.Occurs at all temperatures Occurs at one particular temperature (boiling point)

2.Occurs on the surface of the liquid Occurs both inside and on the surface of the liquid

3.Takes place slowly Takes place faster

4.Bubbles are not necessarily formed Bubbles are formed

Therefore, the two terms can be defined as follows:

Evaporation is a change in state of a substance from liquid to gas (vapour) state at any temperature.

Boiling is a change in state of a substance from liquid to gas at a particular temperature and pressure.

Condensation and solidification

The reverse of evaporation is condensation. This is brought about by cooling. When a gas is cooled down, its particles lose energy. They move more and more slowly. When they knock into each other, they do not have enough energy to bounce away again. They stay close together and a liquid forms. This process is called condensation. When the liquid is cooled further, the movement of the particles slows down even

more. Eventually, they stop moving and a solid forms. This is called solidification.

Condensation can be defined as a change in state of a substance from gas (vapour) to liquid.Solidification is

a change from liquid to solid state of a substance. Solidification is the same as freezing.

Sublimation

A few solids do not melt when they are heated. Instead, they change directly from the solid to gaseous state without passing through the liquid state. This change in state is called sublimation. When a solid

changes directly into gas, it is said to sublime. Iodine, solid carbon dioxide ("dry ice") and ammonium chloride are examples of solids that sublime. Like melting, sublimation also occurs at one particular temperature for each pure solid.

5.2.2. Importance of change in state

The following points summarize the importance of change in state:

1. Separation of mixtures

Different mixtures can be separated through such processes as distillation, sublimation, evaporation and condensation. Let us have a look at an example of distillation. This process involves boiling, evaporation and condensation. Distillation as a process can be applied in separation of a mixture (solution) of two or more substances. A mixture of two or more substances with different boiling points e.g. water and alcohol can be separated by this means. In such a case, a container with the mixed-up liquids is heated. The liquid with a low boiling point evaporates and condenses first, leaving the one with a high boiling point in the container. The distillate (liquid with low boiling point) is collected, cooled down and

transferred into another container.

2. Industrial manufacture of products

Industrially, the process of distillation is applied in the production of pure substances such as beer and other alcoholic drinks such as wine, vodka, konyagi, etc. The manufacturing process involves boiling,

evaporating and condensation.

3. Refining of petroleum (crude oil)

Crude oil contains organic liquid components, each with a different boiling point. In the refinery, the components with lower boiling points evaporate first and get separated out, leaving those with higher boiling points behind. In this way, we get various types of oil components (fractions) such as petrol,

diesel, kerosene, lubricating oil, etc.

4. Drying of crops and clothes

When you suspend your clothing on a cloth line to dry, the moisture in it is lost through evaporation. Likewise, farmers in the village often spread crops on the ground to dry. They do this in order to reduce moisture content and hence prevent decaying. The moisture contained in crops leave by evaporation. Therefore, you can notice how evaporation, as a change in state, is important in everyday lives.

5. Cooling of our bodies in hot weather

During hot weather, our bodies perspire a lot. When water evaporates from the body, it takes up heat. This brings about the cooling effect, as heat is lost from the body surface. The cooling effect is more evident when the wind or air is blowing over the body. This is because wind increases the rate of evaporation. In this way, the body gets cooled down.

6. Ice formation in refrigerators

You all like to drink cold water or beverages especially during hot weather. You can use a refrigerator to cool down drinking water or beverages directly. Alternatively, you can freeze water into ice and then use the resulting ice for cooling the beverage. Ice blocks are also saleable. Moreover, one can earn some money if she freezes water into ice blocks and then sells them to beverage vendors. Perishable products such as fish, meat, milk, etc are often packed in ice blocks to prevent them from going bad. Ice, as we

studied early, is formed when water freezes (a change in state from liquid to solid).

7. Melting metals to make alloys

In metallurgical industries, need may arise to mix two or more metals (alloys) together. This is only possible, where two or more metals are first melted at high temperatures into liquids. Then the resulting liquid metals are mixed in appropriate proportions. This is followed by cooling down the mixture to a

solid alloy. Normally alloys have better qualities than individual metals.

8. Testing the purity of substances

The presence of impurity may lower or raise the boiling point of the substance. A pure substance melts and boils at definite temperatures (see table 5.4). The values for the melting point and boiling point are precise and predictable. This means that we can use them to test the purity of a sample. They can also be

used to check the identity of unknown substance.

A typical example

Sea water is impure. It freezes at a temperature well below the freezing point of pure water (0C) and boils at a temperature above the boiling point of pure water (100C). Other substances behave in a similar

manner. So, boiling as a change in state can be used to test for the purity of a substance.

In addition, the impurity also reduces the exactness of the melting or boiling point. An impure substance

melts or boils over a range of temperature, not at a particular point.

Table 5.4. Melting and boiling points of some pure substances

Substance Melting point (C) Boiling point (C)

Water 0 100

Ethanol -117 78

Oxygen -219 -183

Sodium 98 890

Sulphur 119 445

Iron 1540 2900

Diamond 3550 4832

Cobalt 1492 2900

Nitrogen -210 -196

Propane -188 - 42

Ethanoic acid 16 118

9. Formation of rain

Perhaps the most important of all, as far as change in state is concerned, is the formation of rain. Rained is mainly formed through the process of evaporation and condensation. Water vapour, evaporating mostly from water bodies (oceans, seas, lakes, rivers, ponds, etc), land and plants rises up to the sky. As it rises, it cools down and condenses into tiny droplets. On further cooling as they rise up, these droplets form bigger water drops. Owing to gravitational force, these drops fall down as rainfall. Every one of you knows how important rain is to our life. Therefore, you have noticed how evaporation and condensation,

as changes in state, contribute to rain formation.

Fig. 5.4 Rain formation

5.2.3. Kinetic nature of matter

We already know that matter is composed of atoms, ions or molecules. We have not yet considered the reason why the same substance, say water, can exist in more than one form, for example as solid ice,

liquid water, and gaseous steam. But does matter behave like that?

The kinetic theory of matter has been used to explain the way in which the arrangement of the particles of a substance can determine the properties of that substance, and particularly the state in which it is likely to be found under a given set of conditions. The idea is that all matter is made up of tiny moving

particles. The main points of the theory are as follows:

1. All matter is made up of tiny particles (atoms and molecules) that are invisible to the naked eye and to

most microscopes.

2. The particles are moving all the time. The higher the temperature is, the higher the average energy of

the particles.

3. Heavier particles move more slowly than lighter particles at the same temperature.

4. Each substance has unique particles that are different from the particles of other substances.

5. The particles of matter are held together by strong electrostatic forces.

6. There are empty spaces between the particles of matter that are very large compared to the particles

themselves.

The solid state

In the solid state, the particles are so closely packed (see figure 5.5(a). The particles are held together by strong forces of attraction that act like a chemical glue. Free movement of particles cannot take place. They cannot move around freely in this arrangement. Instead, they vibrate about a fixed position. They are arranged in a fixed pattern which form a cluster of vibrating masses. This makes a solid to have a

fixed shape, which cannot be changed except by applying strong external forces.

The liquid state

The particles of a liquid are also closely packed but the forces of attraction between them are weaker than of a solid. These forces of attraction tend to bind them together. The particles have more kinetic energy and they can move around each other. The binding forces are strong when particles come close to one another. It is thought that the particles of a liquid are fairly randomly arranged but consist of "clusters" closely packed together. This property makes a liquid to have a definite volume. However, since the particles are fairly free to move a liquid does not have any characteristic shape (see figure 5.5(b). Thus, a liquid will always take the shape of its container.

The gaseous state

The gaseous state is one in which the particles are moving independently of each other in all directions and at great speeds. The particles of a gas are relatively far apart (see figure 5.5(c). They exert no force of

attraction on each other. They have more energy than the particles of solids and liquids. They move rapidly and randomly, colliding with each other and with the walls of the container. A typical speed for a molecule of hydrogen in air at ordinary temperature and pressure has been found to be approximately 500 ms-1. It has been estimated that a nitrogen molecule makes 109 collisions each second. Thus, a gas will

rapidly spread out to fill any container in which it is placed. A gas cannot have any shape of its own.

Figure 5.4. Three states of matter

5.3. PHYSICAL AND CHEMICAL CHANGES

Depending on the nature of change, all changes that matter undergoes can be classified as either physical

or chemical.

Physical change

Substances may undergo changes in their physical properties e.g. changes in colour, shape (or form), state, density, structure and texture, etc. If you take a stone and break it down into small particles, you will have only changed its form, but it will remain as a stone. Likewise, melting ice to water or freezing water to ice does not change it, but it is still water. The same case happens when you dissolve salt in water to get a solution of salt in water. You can still get back the original salt by evaporation, except that the crystals of the salt obtained will not look exactly the same as those of the original salt.

These changes of state are examples of physical changes. Physical changes such as melting and boiling do not result in new substances being formed. For example, ice and water still contain the same particles whether in solid (ice) liquid (water) or gaseous (vapour) state.

Figure 5.6. Changes in states

Characteristics of a physical change

In the explanation above, we find that in a physical change it is only the physical form, and not the actual nature, of a substance that changes. The changes are brought about by a mere addition or removal of heat, as in the case with water or ice. Such a change is called a physical change. It can be distinguished by

the following characteristics:

1. There is no formation of a new substance. Consider an example given above. The ice, liquid water and

steam are the solid, liquid and gaseous forms of the same substance (water).

2. There is no change in weight of the substance undergoing the change. If you start with 50g of ice, you

will still get the same mass of water and steam (vapour) upon melting and boiling respectively.

3. The changes are readily reversible. You can easily change water back to ice and vapour to water by a

mere subtraction of heat (cooling).

4. It is not accompanied by a great heat change. Just a little heat is required to change ice to water, and water to steam.

Demonstration of physical changes experimentally

Experiment 5.2.

1. Add some common salt (sodium chloride) to distilled water in a beaker. Stir the mixture until the salt disappears and forms a solution with water. Transfer the water into a porcelain dish. Heat the content

until all the water has evaporated off. The salt reappears in its original white solid form.

2. Grind some roll sulphur in a mortar to powder. Put the resultant powder in a test-tube and heat gently, shaking all the time. The sulphur melts to an amber-coloured liquid. On cooling, this liquid

returns to its original condition as a yellow solid.

3. Put a block of ice in a beaker. Heat gently until the whole block melts to form water. Pour the water

formed in a cup and place it in a deep freezer overnight. The water will freeze back to ice.

You will have seen that all the above changes involve only changes in physical forms of the substances. The chemical nature of substances remained unchanged. Therefore, we can define a physical change as a change that does not involve formation of a new substance but involves a change in state or physical

form of the substance and that such a form can be reversed.

Chemical change

Some changes that materials undergo are permanent. Such changes usually involve changes in chemical properties of a substance. For example, when you burn a piece of wood in fire, you get ash. The properties of wood and ash are very different. There is no way you can change ash back to wood. It is practically impossible. A permanent change in chemical properties of a substance is called a chemical change. In a chemical change, a substance losses all its physical and chemical properties.

Characteristics of a chemical change

1. A chemical change results in the formation of a new substance. The new substance has different

chemical and physical properties as compared to the original substance.

2. It is generally not reversible. For example, you cannot turn the ash back to wood.

3. There is a change in weight or mass of the substance undergoing the change. When you burn wood

weighing 5 kg, you cannot expect to get the same weight of ash.

4. The change is accompanied by a considerable heat change. For wood to burn to ash a lot of heat must

be supplied.

We can therefore define a chemical change as the one in which a new substance is formed and that such a

form cannot be reversed.

Demonstration of chemical changes experimentally

Experiment 5.3.

1. Strongly heat some roll sulphur on a deflagrating spoon until it melts and begins to burn with a blue flame. If you continue heating, it gradually decreases in amount and finally the spoon will be left empty. The disappearance of sulphur is due to the formation of a new gaseous substance that is invisible. The presence and existence of a gas in air can be defected by its irritating smell. The gas can also be detected by burning the sulphur in a gas jar to which some blue litmus solution has been added. The gas formed,

sulphur dioxide, will turn the blue litmus paper into a red one.

2. With the aid of tongs, subject a piece of magnesium ribbon to a Bunsen burner flame. The ribbon

burns to produce a new substance, white ash of magnesium oxide.

3. Wrap a wet cotton wool around an iron nail. Keep it in a test tube for 3 days. By the 3rd day, some brown marks of rust will appear on the surface of the nail. Rust is hydrated iron (III) oxide. This is quite a new substance compared to iron nails.

Table 5.5. Differences between physical and chemical changes

Physical change Chemical change

1. Produces no new kind of matter Always produces a new kind of matter

2. There is no change is mass or weight of the substance

2. There is a substantial change in the weight of the substance

3. The change can be reversed 3. The change cannot be reversed

4. Little heat is absorbed or evolved 4. Heat changes may be large

5. The change involves only a change in physical properties of a substance

5. Both physical and chemical properties are changed.

5.4. ELEMENTS AND SYMBOLS

5.4.1. Elements

Our world and universe are made up of millions of different substances. We have already seen how these substances can be classified into solids, liquids and gases. However, on close examination, we find that

these substances are made up of a number of small elements.

An element can be defined as a substance that cannot, by any known chemical process, be split into two or more simpler substances. This means that elements cannot, by any chemical process, be made to yield substances simpler than themselves. An element is a substance because it has the same composition

throughout.

In 1803, a scientist called John Dalton suggested that each element was made up of its own kind of particles. He called these particles the atoms. Therefore, an element is a substance that is made up of

only one kind of atoms.

There are over 105 different elements known. Of these, 90 have been obtained from the Earth's crust and the atmosphere, and 15 have been artificially made by scientists. From this small band of elements, all other substances on earth are made. Table 5.6 shows the approximate percentage composition by mass of the elements in the earth's crust, the oceans, and the atmosphere. Can you notice the abundance of oxygen? Analysis of the earth's crust, the oceans, and the atmosphere, reveals that oxygen is the most

abundant element on earth, accounting for half the total mass.

Table 5.6. Percentage by mass of elements in the earth's crust, oceans and atmosphere

Percentage by mass of elements in the earths crust

Percentage by mass of elements in the oceans

Percentage by mass of elements in the atmosphere

Oxygen 47 Oxygen 86 Nitrogen 75.5

Silicon 28 Nitrogen 10.9 Oxygen 23

Aluminium 7.8 Chlorine 1.8 Argon 1.4

Iron 4.5 Sodium 1.0 Hydrogen 0.02

Calcium 3.5 Magnesium 0.1 Carbon 0.01

Sodium 2.5 Calcium 0.05

Others

( total)

0.07

Potassium 2.5 Sulphur 0.05

Magnesium 2.0 Potassium 0.04

Titanium 0.5 Nitrogen 0.02

Hydrogen 0.2 Bromine 0.01

Carbon 0.2 Carbon 0.01

Others

( total)

1.3 Others (total) 0.02

5.4.2. Names and Symbols of Elements

A chemical symbol is the way of representing an element using initial letter(s). There are many different elements as you have seen above. Every element has a name and a symbol to represent it. Some symbols are just a single capital letter, such as H. Others have two letters, the first of which is always a capital, such as Mg.

Rules for assigning chemical symbols to elements

1. Each element is given a different symbol to represent it. 2. Some elements are represented by two letters e.g. Ca (for calcium), Cl (for chlorine), etc. 3. If two letters represent the element, the first letter is always a capital and the second letter is always a

small letter e.g. argon (Ar) and helium (He). 4. In order to avoid confusion, some elements have their chemical symbols derived from Latin names (See

table 5.8)

All symbols are recognized and are used by all scientists all over the world. Some examples of elements

and their symbols are given in the table below:

Table 5.7. Names and Symbols of some Elements

Element Symbol Element Symbol

Aluminium Al Bromine Br

Copper Cu Carbon C

Iron Fe Chlorine Cl

Lead Pb Hydrogen H

Magnesium Mg Nitrogen N

Mercury Hg Oxygen O

Potassium K Phosphorus P

Silver Ag Sulphur S

Sodium Na Silicon Si

Calcium Ca Iodine I

Manganese Mn Fluorine F

Tin Sn Gold Au

Chromium Cr

Zinc Zn

Nickel Ni

It is easy to remember that the symbol for aluminium is Al, and for carbon is C. But some symbols are harder to remember because they are taken from Latin names. For example, potassium has the symbol, K from its Latin name Kalium. Sodium has the symbol, Na from its Latin name Natrium. See the complete

list in the following table.

Table 5.8. Elements with Latin names

English name Latin name Chemical symbol

Sodium Natrium Na

Gold Aurum Au

Potassium Kalium K

Mercury Hydragyrum Hg

Silver Argentum Ag

Antimony Stibium Sb

Lead Plumbum Pb

Tin Stannum Sn

Iron Ferrum Fe

Copper Cuprum Cu

Tungsten Wolfram W

The reason for assigning some elements with Latin names was to avoid confusion among scientists when representing different elements. For example, the symbol for silicon is Si. It could be impossible to represent silicon by the symbol S and at the same time represent the element sulphur by the very symbol, S. Similarly, potassium could not be represented by the symbol P that was assigned to phosphorus. So in order to avoid such confusion, scientists decided to use Latin names to represent some elements. In so doing, the anticipated and unnecessary contradiction among scientists from different parts of the world

was avoided.

Symbols are particularly useful when more than one atom is present in a substance. For example, hydrogen gas consists of pairs of hydrogen atoms joined together. So hydrogen gas is shown as H2. When more than one atom is joined together like this, we call the substance formed a molecule. Atoms making up gases such as hydrogen, oxygen, nitrogen, etc, always exist as molecules. Sulphur exists as a hexagonal ring of eight atoms. Phosphorus exists as a tetrahedron of four atoms. Table 5.9 shows some

elements that exist as molecules.

Table 5.9. Elements that exist as molecules

Element Atomic symbol Molecular symbol

Oxygen O O2

Nitrogen N N2

Hydrogen H H2

Sulphur S S2

Phosphorus P P2

Chlorine Cl Cl2

Fluorine F F2

Bromine Br Br2

Iodine I I2

5.4.3. Classification of Elements

Elements can be classified as either metals or non-metals. Metals and non-metals have different physical and chemical properties. This is the criterion used for classification of these elements into metals on one hand and non-metals on the other hand. Table 5.10 summarizes the physical and chemical properties of

some common elements.

Table 5.10. Properties of some common elements

Element

Date of discovery

Metal or non-metal?

Solid,

liquid or

gas?

Melting point(C)

Boiling point(C)

Density

( g cm-3 )

Oxygen 1774 Non-metal Gas -219 -183 0.00132

Nitrogen 1772 Non-metal Gas -210 -196 0.00117

Carbon Ancient Non-metal Solid 3500 4827 22

Iron 1735 Metal Solid 1540 3000 7.9

Copper Ancient Metal Solid 1080 2500 9.0

Lead Ancient Metal Solid 327 1744 113

Gold Ancient Metal Solid 1060 2700 193

Silver Ancient Metal Solid 961 2200 10.5

Hydrogen 1766 Non-metal Gas -259 -253 0.00008

Aluminium 1825 Metal Solid 660 2450 27

Zinc 1746 Metal Solid 419 910 7.1

Mercury Ancient Metal Liquid -39 357 13.6

Iodine 1811 Non-metal Gas 114 183 4.9

Chlorine 1774 Non-metal Gas -101 -35 0.003

Sodium 1807 Metal Solid 98 890 0.97

Potassium 1807 Metal Solid 64 760 0.86

Sulphur Ancient Non-metal Solid 119 444 2.1

Phosphorus 1669 Non-metal Solid 44 280 1.8

5.4.4. Metals and non-metals

There are 94 naturally occurring elements. Some of them are very rare. Francium, for instance, has never been seen. The radioactive metals neptunium and plutonium, which we make artificially in quite large amounts, only occur in very small (trace) quantities naturally. Most of the elements can be classified as metals. The rest are non-metals. To understand these elements better, refer to the Periodic Table of

Elements at the back of this book.

Classification of elements into metals and non-metals is based on differences between their physical and

chemical properties. Differences between metals and non-metals are shown in table 5.11.

Table 5.11. The differences between metals and non-metals

A: Physical properties

Metals Non-metals

1. Have high densities except for sodium and potassium

Have low densities

2. Shine and can be polished

Are dull and cannot be polished

3. Are malleable and ductile, that is, they can be

Are brittle

hammered into sheets and drawn into wires

4. Have high tensile strength

Have low tensile strength

5. Have high melting points except for sodium, potassium and mercury (which is a liquid at room temperature)

Have low melting points and many of them are gases

6. Are good conductors of heat and electricity

Are poor conductors of heat and electricity except carbon, in the form of graphite, which conducts electricity

7. They make a ringing sound when struck (they are sonorous)

They are not sonorous

B: Chemical properties

1. Give basic oxides, that is, oxides which react with acids

Give acidic oxides, that is, oxides which react with bases

2. Replace hydrogen in acids to form salts

Do not react with acids in this manner

3. Form positive (+) ions

Form negative (-) ions

4. Form electrovalent chlorides which are stable in water

Form covalent chlorides which react with water.

5. Do not react with hydrogen

Form stable compound with hydrogen

The properties discussed above are of a general nature and exceptions do occur. Hence, some elements may appear to be intermediate between metals and non-metals. These are called metalloids or semi-conductors. Others may differ from the two groups in just one or two cases. Such elements have some of

the properties of metals and others that are more characteristic of non-metals. See the Periodic Table at

the back of this book for illustration.

Compound and Mixtures

Form 1 Chemistry

Matter contains a mixture of many elements or compounds. In gaseous phase the atmosphere (air) is a

mixture of many gases like Nitrogen, oxygen, sulfur di oxide, nitrogen oxides, oxides of carbon, ozone

and traces of inert gases along with water vapor. In liquid state the water contains many dissolved salts

of metals like Na,K,Mg,Zn,Al etc, Even organic liquid like petroleum is a mixture of many compounds

5.5. COMPOUNDS AND MIXTURES

A compound is a substance that contains two or more elements chemically combined together. A mixture is something that contains two or more elements not combined chemically.

It is always difficult to identify a mixture from a compound. Before going any further into this topic, let us start by looking at the differences between compounds and mixtures. These differences are

summarized in the table below.

Table 5.12. Differences between mixtures and compounds

Mixtures Compounds

1. The components of a mixture can be separated by physical means, e.g. filtering, magnetic separation, decantation, etc

The components of a compound can be separated by chemical means only

2. The composition of a mixture can vary widely, e.g. a mixture of 20g of sand with 1g of salt or vice versa.

Compounds are fixed in their compositions by mass of elements present, e.g. there are always 2 atoms of hydrogen to 1 atom of oxygen in a molecule of water (H2O)

3. Mixing is not usually accompanied by external effects such as explosion, evolution of heat, or volume change (for gases)

Chemical combination is usually accompanied by one or more of these effects

4. Properties of a mixture are the sum of the properties of the individual constituents of the mixture.

The properties of a compound are quite different from those of its constituent elements. For example, water is a liquid whereas its constituent elements, hydrogen and oxygen, are both gases.

5. No new substance is produced as the mixture forms

A new substance is always produced when a compound forms.

5.5.1. Compounds

A compound is a substance that contains two or more elements chemically combined together. This is a very important difference from mixtures. Mixtures can contain more than one element but the elements are not chemically combined. The number of chemical substances known is approximately four millions. All compounds on earth are made from about one hundred simple materials. Such compounds range from simplest substances, like water, which contains only two elements, to those complex materials of which our own bodily tissues are composed. The following is a short list of common compounds and the

elements they are made of.

Table 5.13. Elemental composition of some compounds

Compound Constituent elements

Water hydrogen and oxygen

Carbon dioxide carbon and oxygen

Ethanol carbon, hydrogen and oxygen

Sugar (sucrose) oxygen, hydrogen and carbon

Sodium chloride

(common salt)

sodium and chlorine

Marble (calcium carbonate) calcium, carbon and oxygen

Sulphuric acid hydrogen, sulphur and oxygen

Sand silicon and oxygen

Clay aluminium, oxygen and hydrogen

Compounds have different properties from the elements that make them up. For example:

1. Water (H2O) is a colourless liquid at room temperature but the elements that make it, hydrogen and oxygen are both gases.

2. Sodium chloride is a white solid made of sodium and chlorine. Sodium is a solid, highly reactive metal, and chlorine is a greenish yellow gas with a chocking smell. Preparation of a binary compound

A binary compound refers to a compound made up of two different elements. In a binary ionic compound, the total numbers of positive charges and negative charges must be equal. The following

experiment demonstrates a typical preparation of a binary compound.

Experiment 5.4. Preparation of iron sulphide

Procedure:

1. Weigh 56g of iron filings and 32g of sulphur.

2. Put the two elements in a mortar, grind them thoroughly and mix uniformly.

3. Put the mixture into a dry test tube. Heat the test tube, at the bottom, with a small flame. The mixture

will glow.

5. When it glows remove the flame. The glow will then spread slowly through the mixture without further heating.

6. Allow the test tube to cool, and then break it away from the mass of material left.

Result:

A dark grey, almost black, solid will be formed. This is iron sulphide. The reaction that took place can be

presented as follows

Fe(s)+S(s) FeS(s)

Here 56g of iron react with 32g of sulphur to produce 88g of iron sulphide.

5.5.2. Mixtures

A mixture is something that contains two or more substances not combined chemically. The substances

may mix up completely or they may remain separate.

Our environment is a mixture of all forms of matter. For example, the earth's crust is a mixture of soils, rocks, minerals, and water. Sea, river, and lake waters contain dissolved gases, living organisms and, sometimes, salt. Air consists of gases, water vapour, and dust particles. The components of each of these mixtures could be elements such as oxygen, nitrogen, sulphur or gold. Alternatively, the mixture might

consist of elements and compounds such as hydrocarbons (e.g. petroleum), water, metallic oxides or salts.

Other substances that can form mixtures when placed or mixed together include sand and sugar, maize

and bean seeds, soil and table salt, water and mud, etc.

Classification of mixtures

Mixtures can be classified as solutions, suspensions or emulsions. This classification is based on whether the mixed substances dissolve completely or not. It also depends on the nature of the mixtures that result

upon mixing. Let us look at each category in detail.

Solutions

A solution is a uniform mixture of two or more substances. Such mixtures may be a solid in a liquid, a liquid in a liquid, a liquid in a gas and, very rarely, a gas in a gas. (See table 5.14). We most often think of a solution as being made of a solid dissolved in a liquid. For example, solutions of sugar or salt in water are quite common. A solid that dissolves in a liquid is called a solute while the liquid in which that solid

dissolves is called a solvent. For example, sugar and salt are solutes and water is a solvent.

However, other substances that are not normally solids can be found dissolved in a liquid. For example, the gases, carbon dioxide and oxygen, dissolved in water are important for life to continue in oceans,

seas, lakes, rivers, etc.

Less obvious perhaps, but quite common, are solutions of one liquid in another. Alcohol mixes (dissolves) completely with water. Beer, wine and whisky do not separate into layers of alcohol and water (even when the alcohol content is quite high). Alcohol and water are completely miscible, that, is they make a

solution.

Solutions of gases in gases are very uncommon. Technically, air could be described as a solution of several gases in nitrogen, though this could be unusual everyday use of the term. However, it is

interesting to note that different gases always mix completely with each other.

Table 5.14 Examples of types of solutions

Solutes

Solid Liquid Gas

Solvents

Gas Naphthalene slowly sublimes in air to form a solution

Water vapour in air Oxygen and other gases in the air

Liquid Sucrose (sugar) in water and salt in water

Ethanol (alcohol) in water and various hydrocarbons in each other (petroleum)

Carbon dioxide in water (carbonated water)

Solid Steel and other metal alloys

Mercury in gold and hexane in paraffin wax

Hydrogen in metals

Suspensions

A suspension is a cloudy mixture of solid particles suspended in a liquid. A solid is said to be in

suspension in a liquid when small particles of it are contained in a liquid, but are not dissolved in it.

If the mixture is left undisturbed, the solid particles will slowly settle to the bottom of the containing

vessel, leaving the pure liquid above them.

Muddy water is a typical suspension. The mud would settle after a time if left undisturbed leaving brown residue on the bottom of the containing vessel and clear water above. The particles of mud would be

retained by filtering whilst the water (and any solids in solution) would pass through.

If you mix flour or chalk dust in water, it forms a suspension. Their particles are simply dispersed (spread) throughout the water and would eventually settle down to the bottom of the vessel if left

undisturbed for sometime.

Table 5.15 shows the differences between solutions and suspensions

Table 5.15 Differences between solutions and suspensions

Solutions Suspensions

Homogeneous Heterogeneous

Transparent/clear Opaque/not clear

Particles completely dissolved Particles separate on standing

Components separated by evaporation Components separated by filtration

Emulsions

An emulsion is a cloudy mixture of tiny droplets of one liquid suspended in another liquid. Sometimes two immiscible liquids will not separate out into two layers when mixed together. One of the liquid may form droplets and spread throughout the other to form an emulsion. Cooking oil and water do not mix but

they will form an emulsion when they are mixed and shaken. Droplets of oil will spread throughout the water. Unlike pure liquids, emulsions are cloudy (opaque). So you cannot see through them. The

emulsion will not settle like a suspension. Which other liquids you know can form suspensions?

Formation of mixtures

Mixtures can be formed from different substances in two major ways.

The first type constitutes homogenous mixtures, where the substances are totally mixed together uniformly.

Examples include solutions of salts and sugars in water.

The second type constitutes heterogeneous mixtures, where the substances remain separate and one

substance is spread throughout the other as small particles, droplets, or bubbles. All emulsions and suspensions fall under this category. Examples include suspensions of insoluble solids or oil droplets in water.

5.6. SEPARATION OF MIXTURES

To make use of the materials around us, we need methods for physically separating the many and varied mixtures that we come across. One of the distinctive characteristics of a mixture of substances is that it is usually possible to separate the constituents by physical means. There are many different physical methods used to separate a wide variety of mixtures. The particular method employed to separate any given mixture depends upon the nature of its constituents. The following are some of the methods in wide use.

5.6.1. METHODS OF SEPARATING MIXTURES

1. Filtration

This method is best applicable in separation of components of mixtures called suspensions. A mixture of chalk dust or flour with water can be separated by filtering the suspension. The suspended particles get trapped in the filter paper. The trapped particles are called the residue. The water is called the filtrate.

Figure 5.8. Decantation of muddy water

3. Evaporation

This method is used to separate substances that form a solution. In such a mixture, the solute is completely dissolved in a solvent to make a uniform solution. To separate these substances, the solution is heated so that the solvent evaporates, leaving the solid residue behind. A mixture of salt or sugar in water can be separated by applying this method.

5. Fractional distillation

Separating the liquids from a mixture of two (or more) miscible liquids is again based on the fact that liquids will have different boiling points. However, the boiling points are closer together than for solid-in-liquid solutions. It is difficult to separate mixtures of liquids whose boiling points differ by only a few

degrees. In this case, fractional distillation is used.

For example, ethanol boils at 78C whereas water boils at 100C. When a solution of ethanol and water is heated, ethanol and water vapours enters the fractionating column. Evaporation and condensation take place as the vapours rise up the column. Ethanol passes through the condenser first as the temperature of the column is raised above the boiling point. Water condenses in the column and flows back into the flask

because the temperature of the column is below its boiling point of 100C.

The temperature on the thermometer stays at 78C until the ethanol has distilled over. Eventually, the thermometer reading rises above 78C. This is a sign that all the ethanol has been separated, so heating can be stopped. By watching the temperature carefully, the two liquids (fractions) can be collected

separately.

Various forms of fractionating column can be used. Their general purpose is to provide surfaces, e.g. flat discs, on which ascending vapour can condense. Glass beads in the column provide a large surface area

for condensation.

Figure 5.12. Sublimation of ammonium chloride

7. Chromatography

This method is commonly used to separate a mixture of coloured substances (solids or dyes). An example of this is the separation of dyes that make up black ink. Chromatography works better when a solvent is used. The commonest solvent is water, though other solvents such as ethanol or ether may be used for those substances that do not dissolve in water. There are two types of chromatography, namelycolumn chromatography and paper chromatography. The two types of chromatography follow the same principle,

but paper chromatography is the simplest form to set up, and hence is more commonly used.

On which principle does chromatography work? Let us consider an example of separating dyes that

make up black ink. In this case, water is used as a solvent.

Experiment 5.5. Separating dyes in ink

Procedure

1. Put a small spot of the water-soluble ink onto a strip of filter paper as shown in figure 5.13.

2. Place the filter paper in a beaker of water. Make sure the level of the water is below the level of the ink

spot.

3. Leave the filter paper until the water has risen to the top of the paper.

4. Remove the paper and allow it to dry.

5. Note the colours the ink contains.

Uses of Chromatography

Chromatography is used in many different ways. The following are some of the application of chromatography:

1. It can be used to find out the components of a liquid or solid, or even to identify different substances. 2. It can be used by security agents and medical personnel to analyse blood and urine samples. 3. Causes of pollution in water and in animals that live in water can also be detected using chromatography. 4. In chemistry, chromatography is used to test the purity of substances and in separation of mixtures.

8. Layer separation

Mixtures of two immiscible liquids can be separated with a separating funnel. The mixture is placed in a separating funnel and allowed to stand. The liquids separate into two different layers. The lower denser

layer is then "tapped" off at the bottom.

For example, when a mixture of kerosene and water is poured into the funnel, the kerosene floats to the top as shown in figure 5.15. When the tap is opened, the water runs out. The tap is closed again when all water has gone, leaving the kerosene in the funnel.

The solvent extraction works on two principles:

1. One solid in the solution must be more soluble in the extracting solvent than the other.

2. The extracting solvent must not be miscible with the solvent in which the mixture of solids is

dissolved. Neither should it react with it.

10. Centrifugation

A centrifuge is used to separate small amounts of suspension. Centrifugation is used with insoluble solids where the particles are very small and spread throughout the liquid. In centrifugation, test tubes containing suspensions are spun round very fast. The solid gets thrown to the bottom. Here, it is no longer the force of gravity on the solid that causes settling. Instead, there is a huge centrifugal force acting on the particles due to the high speed spinning of the samples. This causes the solid to be deposited at the bottom of the centrifuge tube.

11. Magnetic separation

If the solid mixture contains iron, the iron can be removed using a magnet. This method is used to separate scrap iron from other metals. Magnetic iron ore can be separated from other material in the crushed ore by using an electromagnet. In the process of recycling metals, iron objects can be picked out

from other scrap metals using electromagnets.

12. Crystallization

This process involves evaporation but the speed of evaporation is much slower. In principle the salt solution can be left in the evaporating basin for a long period until all the water has evaporated but in

practice this takes longer time.

The process begins by evaporating away the liquid. However, because the crystals are needed, evaporation is stopped after the solution has been concentrated enough. The concentrated solution is

allowed to cool slowly and crystallize. The crystals so formed can be filtered off and dried.

A similar process is used to extract salt from the sea. Salty sea water is placed in wide basins and put in the sun. Water evaporates off, leaving the salt crystals in basins.

13. Winnowing or threshing

This is a method used to separate grains from husks or bran. The process makes use of the differences in density of the constituents in the mixture. When the winnower is shaken around, grains, being denser than husks or bran, sink to the bottom of the winnower. The less dense husks or bran moves to the top. They are then blown off the winnower by wind or breath, or sometimes picked by hand and separated

from the grains.

5. (i) Paper chromatography is very useful in analysis of substances present in a solution. For example, it can

tell whether a substance has become contaminated or otherwise. This can be very important, because

contamination of food or drinking water, for instance, may be dangerous to our health.

(ii) Chromatography has proved very useful in the analysis of biologically important molecules such as sugars, amino acids, and nucleotide bases. Molecules such as amino acids can be seen if the paper is

viewed under ultra- violet light.

(iii) Paper chromatography is the test that can be used to check for the purity of a substance. If the sample is

pure, it should only give one spot when run in several different solvents (see figure 5.13).

6. Other separation methods are also used to check whether purification has been successful. Samples obtained by distillation can be re-distilled. The purity of crystals can be improved by re-crystallisation. A water sample can be tested for amount of dissolved material by evaporating a certain amount of water to

dryness. The solid waste can be weighed. This would give the amount of dissolved solid in the water.

The process of purification is of crucial importance in many areas of chemical industry. Medical drugs (pharmaceuticals) must be of highest possible degrees of purity. Any contaminating substances even in

very small amounts may have harmful side effects.

7. (i) Separation of cream from whole milk is done by the process of centrifugation. As the milk is spun,

the heavier contents are forced down and the lighter cream rises up. After centrifugation, the cream is poured off the top by decantation. This is the initial stage of milk constituent separation, after which other

components such as milk proteins (cheese) are separated.

(ii) Centrifugation is applicable in blood analysis, where the solid part of blood is separated from the liquid part by centrifugation. Blood is a suspension containing microscopic blood cells (corpuscles) in a liquid called plasma. If blood is centrifuged in a test tube, the blood cells are flung to the bottom, leaving

the liquid plasma on top.

8. Knowledge of separation of two immiscible liquids can be applied in the extraction of metals such as iron from their ores. For example, at the base of the blast furnace, the molten slug forms a separate layer on top of the liquid iron. The two can then be "tapped" off separately. The method is very useful in

organic chemistry as part of the process called solvent extraction.

9. Evaporation process is used in the extraction of common salt from seawater whereby the sun

evaporates water molecules from salty water, leaving crystals of the salt behind.

10. Layer separation technique is applied in the recovery of liquids from contaminants.

11. Solvent extraction process is applied in the extraction of certain edible oils from seeds, and in the

extraction of some metals from sludge mixture.

Basic Chemistry Laboratory Apparatus

Form 1 Chemistry

Find ring stands and accessories, burettes, glass tubing, digital balances, tongs and clamps, corks and

rubber stoppers, distillation equipment, and other chemistry lab equipment.

Instruments used for carrying out different experiments in the laboratory are called laboratory apparatus.

Laboratory apparatus can be classified according to their uses as:

apparatus for holding things e.g. test-tube holder, retort stand and clamp, test-tube rack, tongs and tweezers;

apparatus for taking measurements e.g. thermometer, burette, pipette, measuring cylinder, measuring flask, beam balance, electronic balance, common balance, measuring syringe, beaker and stop watch;

apparatus for heating substances e.g. boiling tube, pipeclay triangle, crucible and lid, wire gauze, deflagrating (combustion) spoon, Bunsen burner, spirit lamp, tripod stand, evaporating dish, wire gauze and stove;

apparatus for doing chemical reactions (or testing) e.g. beaker, test tube, dropper, flask, watch glass, gas jar and thistle funnel;

apparatus for filtering e.g. filter funnel, filter paper and cotton wool;

apparatus for grinding e.g. mortar and pestle;

apparatus for storage e.g. reagent bottles and wash bottle;

apparatus for scooping e.g. spatula; and

apparatus for safety e.g. goggles and hand gloves.

The apparatus can also be classified based on materials they are made of. Most of the apparatus are made

of glass. Others are made of metal, plastic or wood. Just a few are made of clay and asbestos.

Table 2.3 summarizes some common laboratory apparatus and their uses.

Table 2.3. Composition and uses of some chemistry laboratory apparatus

Apparatus Material Uses

1. Test tube Glass Holding chemicals or, heating substances

2. Funnel Glass or plastic Leading liquids into containers, and for filtration purposes

3. Beaker Glass or plastic Holding, heating, and mixing liquids

4. Flask Glass Holding, heating, and titrations

5. Retort stand Metal (iron) Holding apparatus during heating

6. Tripod stand Metal (iron) Holding apparatus during experiments

7. Gas jar Glass Gas collection

8. Wash bottle Plastic Washing

9 Crucible Ceramic or non-reactive metal

Heating

10 Test tube holder

Metal and plastic or wood

Holding test tubes while heating

11. Weighing balance

Metal and plastic Measuring weight (or mass)

12. Spatula Metal Scooping small quantities of powder or crystalline chemicals

13. Condenser Glass Cooling hot liquids

14. Pipette Glass Accurate measurement of specific volumes of liquids for titrations

15. Burette Glass Titrations

16. Trough Glass Assists in gas collection

17. Tongs Metal Picking and holding hot substances and apparatus

18. Measuring jar Glass Measuring volumes of liquids

19. Thistle funnel Glass Leading liquids into containers and apparatus

20. Dropper Glass and rubber Dropping indicators into reagents

21. Mortar and pestle

Clay Crushing or grinding substances

22. Wire gauze Metal Even distribution of heat during heating

23. Spring balance Metal Measuring weight

24. Distillation flask

Glass Distillation

25. Combustion spoon

Metal Burning powder in jars

26. Thermometer Glass and liquid metal

Measuring temperature

27. Delivery tube Glass Allowing gases pass through

28. Bunsen burner Metal Heating substances

29. Separating funnel

Glass Separation of immiscible liquid mixtures

30. Measuring cylinder

Glass or plastic Measuring volumes of liquids

31. Measuring syringe

Plastic Sucking in and measuring specific volumes of liquids

32. Stopwatch Plastic or glass and metal

Accurate measurement of time

33. Watch glass Glass Used as a surface to evaporate some liquids, to hold substances being weighed or observed, or as a cover for a beaker

34. Boiling tube Glass Is a large test tube used to heat substances requiring strong heating, or when the sample is too large for a test tube

35. Evaporating dish

Ceramic Heating and evaporating liquids and solutions

36. Filter paper Paper Filtration

37. Test tube rack Wood or plastic Placing test tubes

38. Reagent bottle Glass Storing different chemicals

39. Wash bottle Plastic Storing distilled water

40. Safety goggles Glass Protecting eyes from chemical spills, strong light and harmful vapours

41. Bell jar Glass Keeping gases, moisture, air, etc. or creating vacuums

Figures of some chemistry laboratory apparatus

Using some chemistry laboratory apparatus

With the help of your chemistry teacher or technician, practise the use of the apparatus for measuring the

following physical quantities:

1. Volumes of liquids and gases 2. Masses of solids 3. Temperature 4. Time 5. Length

Activity 2.1

Aim: To measure volume of liquids using different apparatus

Materials: pipettes, burettes, measuring cylinders, water, beakers.

Procedure

1. Pour some water into a graduated measuring cylinder with a capacity of 100 cm3. Add the water, one

drop at time, up to a 25-cm3 mark.

2. While adding water, position yourself at eye-level with the mark on the cylinder. This will enable you to obtain the most accurate measurement. To simplify the work of reading the level of the water, you may

use coloured water.

3. Select a volumetric flask measuring 50 cm3. Pour the water into the flask until it reaches the mark on the flasks neck.

4. Position yourself at eye-level with the mark. You will obtain the most accurate reading when the mark appears straight rather than elliptical. To obtain this, put a flask on a flat table.

5. Add water one drop at a time. Do so until the bottom of the curved surface of the water exactly

matches the mark on the flask.

Your teacher will guide you how to measure the volume of liquids using the other apparatuses.

Questions for discussion

1. Did you manage to measure the required volumes correctly?

2. Which of the two apparatus is easy to use to measure volumes of liquids? Why?

Activity 2.2

Aim: To measure the masses of solid substances

Materials: chemical, electronic or spring balance, watch glasses, various substances such as sand, sugar,

salt, flour, stones, fruits.

Procedure

1. Put an empty watch glass on the weighing balance. Note down its mass. Record this as mass M1.

2. Place the various items you have on the watch glass, one item at a time. Note down the mass. Record

this as M2.

Note: to obtain the mass of an object, we subtract the mass of an empty watch glass from the mass of the

watch glass and the substance. That is, M2 - M1.

For example

Weight of watch glass = 2.6 grams

Watch glass + st grams

Therefore, mass of st 2.6) = 42.4 grams

Questions for discussion

1. What is the mass of each item that you have measured?

2. Record your measurements in a table like the one shown below.

Name of substance Mass in grams

Activity 2.3

Aim: To measure the temperature of liquids

Materials: thermometer, beakers, tripod stand, wire gauze, stopwatch, a pair of tongs, water

Procedure

1. Pour some tap water into two beakers. Dip a thermometer in each of the beakers. Let it stand there for

one minute.

2. Remove the thermometer from the water and record the temperature.

3. Place one beaker in a fridge and leave it there for about half of an hour.

4. Remove the beaker from the fridge. Dip a thermometer in the water for one minute. Record the

temperature.

5. Place a 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 ten minutes.

8. Turn off the 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 there for one minute. Remove it

from the beaker and note down the temperature.

Questions for discussion

1. What is the reading on the thermometer when it is placed in:

1. tap water? 2. water from the fridge? 3. heated water?

2. For what ranges of temperature can the thermometer give readings

2.5. CHEMICAL WARNING SIGNS

Chemical warning signs are safety symbols found on containers, especially those used in the laboratory. The symbols are also found on tanks or containers that are used to carry, store or transport certain chemicals. Containers holding flammable fuels such diesel, petrol and natural gas, as well as those containing toxic chemicals normally bear warning symbols. These symbols indicate the danger (hazard)

likely to be caused by the chemicals they contain if carelessly handled.

When performing experiments in the laboratory it is important to read the safety signs on chemical containers. This will minimize the chances of causing accidents in the laboratory.

All chemists now have to follow strict rules when handling chemicals. These rules must be obeyed whether you are working in an industry, a research laboratory or a school laboratory.

Before your teacher does any experiment with you, he will have to check for possible hazards and will warn you of these. It is important that you follow all instructions that you are given. There are many

hazard signs but the most common ones are shown below. These signs are called the Hazchem Code.

Figure 2.11. Some chemical w

warning signs

Before conducting any experiment in the laboratory you must be aware whether the chemical you want to use is toxic, corrosive, flammable, oxidant, explosive or harmful. This information will help you know

how to handle the chemicals safely. Proper handling of chemicals enables you avoid unnecessary

accidents. Below is an explanation pertaining to some hazard labels represented by the symbols above.

Toxic

Toxic substances include those that can poison you or the other person working close to you in the laboratory. These substances can kill within a short time or after some few days. They should not be allowed to get into your body through body orifices (month, nose, eyes, ears, etc). Neither should they be allowed to contact your skin. They become even more dangerous when they get into the body. If it

happens that these substances touch your skin accidentally, wash it immediately with ample water.

Corrosive

Corrosive substances refer to those chemicals that can burn or corrode (eat away) your skin. They can also corrode wood or metals. One can become blind if such substances accidentally get into his/her eyes. If they contact your skin, wash it immediately with a lot of water. Examples of corrosive substances commonly found in a school laboratory are concentrated mineral acids such as sulphuric acid, hydrochloric acid and nitric acid, and concentrated alkalis such as sodium hydroxide, potassium

hydroxide and ammonia.

Flammable

These chemicals catch fire easily. For this case, they should be kept away from flames or fires. They can be set into fire by any kind of sparks, be it from welding or fire. When working with flammable chemicals in the laboratory all burners must be put off. These chemicals are usually very volatile. The containers used to carry them must be stoppered immediately after every use. Examples of flammable chemicals are

methylated spirit, ether, acetone and methanol.

Explosive

Explosive chemicals are those that explode rapidly upon detonation (set into fire or ignited). Because the reaction is rapid, it results into throwing off particles at a high speed. For this reason, they should not be kept in glass containers. This is because during explosion the particles will disperse around and cause serious injuries to people. Those explosive chemicals that can react without external detonation are even

more dangerous

Oxidizing agents

These chemicals can stimulate a burning substance to burn efficiently and faster. Therefore, they must be

kept away from fires no matter how small that fire may be. An example of oxidizing agent is oxygen gas.

Harmful or irritant

Harmful substances are those that can impair your health or make you fall sick. They do not normally kill instantly but have detrimental effects following a long exposure to them. These chemicals do not kill immediately. However, care must be taken when handling or dealing with them. Irritating substances cause pains when in contact with the body. They are dangerous to health when in contact with the body

surface for a long period of time.

Air, Combustion, Rusting and Firefighting

Total volume of air present in atmosphere consists of 78% nitrogen 21% oxygen and remaining 1% is

made up of other gases such as argon, neon, helium, krypton, xenon and radon.

6.1. AIR

Air is a mixture of different gases. The gases that make up the air include nitrogen, oxygen, carbon dioxide, noble gases (argon, helium, neon, krypton and xenon) and a little water vapour. Air may also contain traces of impurities such as carbon monoxide (CO), sulphur dioxide (SO2), hydrogen sulphide (H2S) and other gases. The presence of these gases in air results in air pollution. Table 6.1 shows the

composition of air by volume. The proportion of water vapour and impurities in air is very variable.

Table 6.1. The percentage composition of air by volume

Gas Approximate percentage

Nitrogen 78.00%

Oxygen 21.00%

Noble (rare) gases mainly argon 0.94%

Carbon dioxide 0.03%

Water vapour 0 4%

The composition of air is not exactly the same everywhere. It changes slightly from day to day and from place to place. There is more water vapour in the air on a damp day and in air above water bodies such as oceans, seas, lakes, rivers, etc. Over busy cities and industrial areas there is more carbon dioxide. But the uneven heating of the earth's surface by the sun causes the air to move continually, resulting in winds. The resultant winds spread the pollutants around.

6.1.1. The composition of air by mass/weight The determination of air by mass was carried out by Dumas in 1841. The apparatus used consists of three units as shown in figure 6.2

Figure 6.2. Determination of the composition of air by weight

The three parts of the apparatus include the following:

1. Several U-tubes containing potassium hydroxide pellets to remove carbon dioxide (only one tube

shown in the figure for simplicity).

2. Another set of U-tubes containing concentrated sulphuric acid to remove water vapour (only one tube

shown in the figure).

3. A heated, weighed glass tube containing finely divided copper to absorb oxygen.

The three parts of the apparatus would, therefore, remove all carbon dioxide, water vapour and oxygen contained in air. The remaining gas which enters the weighed evacuated flask (globe) will be atmospheric

nitrogen and, of course, plus the rare gases.

The copper will have reacted with all oxygen to form copper (II) oxide. The increase in mass of the copper will give the mass of oxygen. The increase in weight of the globe will be due to the weight of nitrogen and the rare gases. If we neglect the weight of carbon dioxide, the percentage of oxygen by mass (weight)

in dry, pure air is 23.2% and the remaining 76.8% is the percentage of nitrogen and rare gases.

The presence of nitrogen in air

In order to demonstrate the presence of nitrogen in air, we need to carry out an experiment that will convert the nitrogen of the air into a chemically recognizable substance. This is easily done by strongly

heating magnesium in the residual gas from the above experiment. Magnesium and nitrogen will react

thus:

3Mg(s) + N2 (g) Mg3N2(s)

Upon treatment with water, magnesium nitrite gives ammonia gas. The gas can be recognized by its

characteristic smell and its action of turning red litmus paper to blue.

The presence of oxygen in air

Oxygen is known as the active portion of the air because it supports combustion and combines with many other substances. Its presence and composition in air can be determined by using these properties. Any of the following two (2) experiments can be used to determine the composition, by volume of

oxygen contained in air.

1. Experiment 6.1. Determination of the presence and proportion of oxygen in air by combustion of a candle

Method

1. Place a small candle on a plastic lid or any object that can float. Then set up the apparatus as shown in figure 6.3. Sodium hydroxide is used in order to absorb the carbon dioxide gas produced by a burning candle.

2. Light the candle and place the measuring cylinder over the top. Note the level of sodium hydroxide solution in the measuring cylinder at the start. A candle will stop burning (go off) once all the oxygen in

the cylinder is used up.

3. When the candle goes off, leave the apparatus to cool to room temperature. The purpose of cooling is to let the heated and expanded air to return to its normal condition. Then note the level of sodium

hydroxide solution in the measuring cylinder.

Figure 6.3. Determining the presence and percentage composition of oxygen in air by burning a candle

Observation and findings

The oxygen in air enclosed in the measuring cylinder is used to burn the candle to produce carbon dioxide gas. The carbon dioxide so produced dissolves in sodium hydroxide solution. The dissolved carbon dioxide causes the level of sodium hydroxide solution to rise up. The oxygen gas used to burn the candle is practically equal to the amount of carbon dioxide produced. This fact is, therefore, used to

calculate the percentage of oxygen in air.

Model results

In the experiment, the initial volume of air was found to be 70.5 cm3 and the final volume was 55 cm3. The

percentage of oxygen in the air is calculated in two steps:

1. To find the volume of oxygen used up to burn the candle (which is practically equal to the volume of carbon dioxide produced and then absorbed by sodium hydroxide), we subtract the final volume of air from the initial volume, i.e.

Volume or oxygen used = Initial volume of air final volume of air

= 70.5 cm3 55.0 cm3

= 14.7 cm3

Therefore, the volume of oxygen used for combustion of the candle = 14.7 cm3.

2. The percentage composition of oxygen in the air = Volume of oxygen used

Initial volume

=14.7 100 70.5

= 20.8%

Alternatively, the volume of oxygen used up can be calculated by subtracting the initial volume of

sodium hydroxide solution from the final volume. That is:

Volume of oxygen used = final volume of sodium hydroxide initial volume of sodium hydroxide =

Volume of carbon dioxide dissolved in sodium hydroxide.

Therefore, the percentage of oxygen =

Volume of carbon dioxide dissolved in sodium hydroxide 100

Initial volume of calcium hydroxide

In practice, it is difficult to get an accurate result with the above experiment. This is due to a number of reasons such as:

1. Not all the carbon dioxide is absorbed by the sodium hydroxide.

2. The candle may go out (stop burning) before all the oxygen is used up due to accumulation of carbon dioxide in the cylinder.

3. The heating of the air inside the measuring cylinder causes the gases to expand. This is why it is essential that the gases be allowed to cool to room temperature before reading the level.

Experiment 6.2 gives the more accurate results than the combustion of the candle. The copper reacts with

oxygen in the air to give copper (II) oxide.

2. Experiment 6.2. Determination of the presence and proportion of oxygen in air by the combustion of copper in air

Method

1. Set up the apparatus as shown in figure 6.4. Syringe A should contain 100 cm3 of air, syringe B should

be empty.

2. Heat the copper strongly and pass the air from syringe A back and forth (by pushing the piston of the syringe inward and outward) over the copper turnings a few times. Allow the air to cool and measure the

volume of air in syringe A.

3. Repeat the heating and cooling until the volume of air that remains in syringe A is constant. The copper

is heated and cooled several times to ensure that it reacts with all oxygen in the sample of air.

Figure 6.4. Determining the presence and percentage composition of oxygen in air by heating copper

Observations and findings

1. The oxygen in the air reacts with copper to form copper (II) oxide, a black solid

Copper + oxygen copper (II) oxide

2Cu(s) + O2(g) CuO(s)

brown metal black solid

2. The final volume of air in the syringe, at the end of the experiment, is less than that of the original volume. This is because oxygen in the original air has combined with copper

Model results

The volume of air in the syringe at different heating and cooling is as shown below:

Initial volume before heating = 100

Volume after first heating and cooling = 82

Volume after second heating and cooling = 79

Volume after third heating and cooling = 79

The volume of oxygen used up = Initial volume of air before cooling - volume of air after the last heating

and cooling

= 100 - 79

= 21

The percentage of oxygen in air = 21 100

100

= 21%

The presence of carbon dioxide in air

Carbon dioxide is present in air to the extent of 0.03% by volume. The gas is formed during the combustion of all common fuels wood, coal, coke, natural gas, petrol, diesel, paraffin oil, etc, all of

which contain carbon.

C (s) + O2 (g) CO2 (g)

It is breathed out as a waste product of respiration by all animals. All sorts of combustion and burning produce carbon dioxide. The gas produced by all these processes accumulates in air. However, the amount of carbon dioxide in air remains constant instead of the tremendous quantities released into the atmosphere. This is because plants take up carbon dioxide. They then convert it into complex starchy

compounds during photosynthesis. The gas also dissolves in ocean water and other water bodies.

The presence of carbon dioxide in air can be shown by passing air through a test tube containing some limewater (figure 6.5). After a time, the limewater turns milky. This shows the presence of carbon

dioxide.

The reaction involved is as follows:

Ca (OH) 2 (aq) + CO2 (g) CaCO3 (s) + H2O (l)

Slaked lime White solid suspended in water

Figure 6.5. Testing for the presence of carbon dioxide in air

The presence of water vapour in air

Water vapour is present in air in varying quantities. It is given off by evaporation from the oceans, lakes

and rivers.

The presence of water vapour in air can be demonstrated by exposing deliquescent substances to the air on a watch glass. These are substances which when exposed to air tend to absorb much moisture from the air, dissolve in that moisture, and finally form a solution. Examples of deliquescent substances include calcium chloride, sodium hydroxide and phosphorous pentoxide.

The resulting solution is distilled. The colourless liquid obtained from distillation may be proved to be water by various water tests such as use of cobalt chloride paper or anhydrous copper (II) sulphate. The cobalt chloride paper turns from blue to pink in the presence of water. The white anhydrous copper (II)

sulphate turns blue. Any of the two tests confirms the presence of water.

Alternatively, one may expose the anhydrous copper (II) sulphate salt to open air straight away for quite some time and then observe any change in its colour and/or form. Upon absorption of water vapour

from the air, the white, powdery and anhydrous copper sulphate salt turns into hydrated blue crystals.

The noble (rare) gases

About 1% of the air by volume is made up of the noble gases. The most abundant of the noble gases is argon. Others are neon, xenon, krypton and helium. The proportion of these four is very minute. Argon and neon are used in gas-filled electric light bulbs and coloured neon electrical signs. They are

obtained from liquefied air (see figure 6.6).

Air pollutants

The air always contains small quantities of many gases. Such gases include hydrogen sulphide, sulphur dioxide, as well as dust and other solid particles, especially in industrial areas. These gases are given off

during the combustion of coal, and the fuels resulting from coal.

SEPARATION OF AIR INTO ITS CONSTITUENT GASES

The air we breathe is necessary to keep us alive. It is also a chemical resource. Oxygen is used in steel making, and nitrogen is used in making fertilizers. To use these gases in this way, they must be separated from the atmospheric air. Air, as we studied in chapter 5, is a mixture of different gases. The method used to separate its constituent gases is fractional distillation. The gases have to be liquefied so that the mixture

can be fractionally distilled.

The process of separating the air into its constituent gases is difficult. It cannot be done in the laboratory.

It is only done in industry. The chemical industry needs the gases from the air in their pure form.

The fractional distillation of air involves essentially two stages:

1. First, the air must be cooled until it turns into a liquid.

2. Then, the liquid air is allowed to warm up again. The various gases boil off at different temperatures.

More details are given in figure 6.7

Stage 1: Liquefaction of air

Air is filtered to remove any dust particles (purification).

The air is cooled to -180C to remove the water vapour and carbon dioxide.

The air is then compressed to 100-150 atmospheres. As the compressed air gets very hot, it has to be cooled.

The compressed cooled air is allowed to expand rapidly. The rapid expansion cools the air to very low temperatures, and the liquid drops out. At -200C, only helium and neon remain as gases. The cold gases are used to cool the compressed air.

Stage 2: Fractional distillation of liquid air

The air is cooled and compressed to form liquid air. The liquid air is allowed to warm up. Nitrogen boils off first because it has a low boiling point, -196C. Argon follows by boiling at -186C and finally oxygen

at -183C (figure 6.6). Figure 6.7 illustrates all the steps that take place during the process.

Figure 6.7. Fractional distillation of liquid air

6.2. COMBUSTION

6.2.1. The concept of combustion

Combustion of a substance in oxygen or air is so common that it becomes almost a habit to use the word "combustion" as if it referred to this kind of reaction alone. In real sense, it may be applied to any chemical reaction accompanied by light and heat in which one or more of the reactants are gaseous.

Many common substances burn in air. Substances such as coal, wood, kerosene, petrol, etc, burn in air. Any substance that burns is called a combustible material. The air or oxygen that supports the combustion is called a supporter of combustion. This is because we live in an atmosphere of air that contains oxygen, which is a very reactive gas. The gas surrounds any burning material. Oxygen is regarded as a supporter of combustion. However, it can sometimes combine chemically with the burning substance to produce a new substance, as we shall see later.

Combustion of a substance involves its reaction with oxygen and the release of energy. These reactions are exothermic and often produce a flame. An exothermic reaction is the one that is accompanied by release of heat to the surrounding environment. Combustion in which a flame is produced is described as

burning. During burning energy is given out in the form of heat, light and sound.

6.2.2. Combustion of different substances in air

Many different substances burn in air to produce different products. Here are examples of combustion of

some common substances:

Sulphur: This is a yellow powder. When burnt in air, it gives misty fumes of sulphur dioxide gas.

Sulphur powder + air (oxygen) sulphur dioxide gas

S(s) + O2(g) SO2(g)

Copper: When a piece of copper foil in a pair of tongs is held in a Bunsen flame, it becomes red-hot. On cooling, a black layer of some substance is observed. This black substance is copper oxide. The reaction

occurs thus:

Cooper + air (oxygen) Copper oxide

2Cu(s) + O2(g) 2CuO(s)

Magnesium: When one end of a piece of magnesium ribbon in tongs is placed in a Bunsen flame, it burns

with a dazzling flame leaving a white ash. This white ash is magnesium oxide.

Mg(s) + O2(g) 2MgO(s)

Hydrocarbons: These are substances containing carbon and hydrogen only. The burning of these organic substances produces carbon dioxide and water vapour as the main products. If oxygen supply is low,

combustion is incomplete and carbon monoxide may be formed.

Candle wax is a hydrocarbon. When it burns in air, the carbon and hydrogen of the wax react with the

oxygen of the air to give carbon dioxide and water vapour respectively.

C(s) + O2(g) CO2(g)

2H2(g) + O2(g) 2H2O(g)

Coal: Coal is a solid fuel that will burn in air to give the following products:

Coal ash + soot smoke + gases (carbon dioxide and steam)

6.2.3. Application of combustion in real life

1. The combustion of a natural gas is an important source of energy for homes and industry. Natural gas

is mainly methane. Its complete combustion produces carbon dioxide and water vapour.

CH4(g) + O2(g) CO2(g) + H2O(g)

Substances like methane, which undergo combustion readily and give out large amount of energy, are known as fuels.

2. There are some reactions where fuels and other substances burn to produce a flame. These are combustion reactions. There are also other combustion reactions (exothermic) where no flame is evident. T