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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.
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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
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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.
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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
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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,
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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.
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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
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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
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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.
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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
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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.
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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.
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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.
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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
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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)
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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
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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
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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
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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
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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
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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
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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.
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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)
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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
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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
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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
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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
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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
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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.
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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
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= 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.
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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
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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
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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.
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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:
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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