New Creative Science Book 9

Let’s observe the above chemical reactions. In these reactions, either atoms or
molecules are combined together or compounds are decomposed or one atom or one
group of atoms are displaced by other. Thus,

The chemical change which involves either addition or decomposition or
exchange of the chemical substances is called a chemical reaction. This chemical
reaction can be expressed by writing full names of reactants and products or by using
symbols. Thus, there are two ways of expressing chemical reactions.

Word Equation

The chemical change which is expressed by writing full names of reactants and
products is called word equation.

For example:

Sodium + Chlorine Sodium chloride

Nitrogen + Hydrogen Ammonia

Hydrochloric acid + Sodium hydroxide Sodium chloride + Water

Chemical equations (or formula equation)

The chemical change which is expressed by using symbols of reactants and
products is called the chemical equation or formula equation. It is easier and more
scientific in writing.

For example: 2NaCl
2Na + Cl2 NaCl + H2O
HCl + NaOH K2SO4 + 2H2O
2KOH + H2SO4

Reactants and products

Observe the given chemical equation, where two substances are combining to
give two new substances. Therefore, combining substances are reactant and new
substances obtained are called products.

2KOH + H2SO4 K2SO4 + 2H2O

Reactants Products

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The chemical substances which are participating in the chemical reaction are
called reactants. In the above reaction, potassium hydroxide and sulphuric acid
are the reactants. They are kept in the left side of the arrow. Similarly, the chemical
substances which are obtained after the chemical change are called products. For
example; in the above reaction potassium sulphate and water are the products . We
write the products in the right side of the arrow.

Exothermic and Endothermic reactions

Exothermic reactions

Those chemical reactions which evolve heat during their processing are called
exothermic reactions. These reactions do not need external energy. Hence, they occur
on their own spontaneously.

For example:

C + O2 CO2+ Heat
C + 2H2 CH4 + Heat

Endothermic reactions

Those chemical reactions which absorb heat during their processing are called
endothermic reactions. These reactions need to supply energy from outside. Hence,
they do not occur spontaneously.

For example:

2KClO3 heat 2KCl + 3O2
CaCO3 heat CaO + CO2

Reversible and irreversible reactions

When hydrogen and iodine combine they give hydrogen iodide. The hydrogen
iodide after decomposition gives hydrogen and iodine again. Such types of chemical
reactions are called reversible reactions.

H2 + I2 2HI

N2 + 3H2 2NH3

The types of chemical reactions which can be reversed back from products to

reactants are called reversible reactions. In these reactions, reactants convert into

products and again products to reactants. Between reactants and products of these

148 | Chemical Reactions

reactions, we use double headed arrows. Similarly, those chemical reactions which
cannot be reversed back from products to reactants are called irreversible reactions.
Between reactants and products of these reactions, we use single headed arrows.
For example, when potassium chlorate is decomposed, it gives potassium chloride
and oxygen. From these potassium chloride and oxygen, we cannot get potassium
chlorate again.

2KClO3 2KCl + 3O2

Unbalanced (or skeletal) chemical equations

When a word equation is translated into chemical equation, it just shows the
symbols of reactants and products. It does not show the actual number of reactants
and product atoms. This type of equation is called unbalanced equation. Therefore,
the chemical equations which do not have an equal number of atoms of each element
in reactants and products are called unbalanced or skeletal chemical equations. These
equations do not follow the law of conservation of mass.

For example:

N2 + H2 NH3
KClO3 KCl + O2
H2O2 H2O + O2

Balanced chemical equations

Unbalanced or skeletal chemical equations do not show the actual number
of reactants and product atoms. It just shows the symbols of these reactants and
products. This type of equation is called unbalanced equation. Now, we use coefficient
in reactant and product molecules so that it has an equal number of atoms in both
sides. Therefore, the chemical equations which have an equal number of atoms of
each element in reactants and products are called balanced chemical equations. These
equations follow the law of conservation of mass with an equal number of reactants
and product atoms.

For example:

N2 + 3H2 2NH3
2KOH + H2SO4 K2SO4 + 2H2O
2H2O2
2H2O + O2

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Ways of writing balanced chemical equation

To balance a chemical equation, the following steps should be followed

1. Write the word equation to show the chemical change.

For example:

Potassium hydroxide + Sulphuric acid → Potassium sulphate + Water

2. The given word equation is converted into correct formula equation.

For example:

KOH + H2SO4 K2SO4 + H2O

3. Now, the total number of atoms of each element in reactants and products

are counted correctly. We use the suitable coefficient to balance the equation.

For example:

2KOH + H2SO4 K2SO4 + 3H2O

Examples of balanced chemical equations

N2 + 3H2 2NH3

2KOH + H2SO4 K2SO4 + 3H2O

2H2O2 2H2O + O2

H2 + I2 2HI

4P + 5O2 2P2O5

2H2 + O2 2H2O

4Fe + 3O2 2Fe2O3

2KClO3 2KCl + 3O2

Information of the balanced chemical equation

We need to balance the chemical equation to obtain the following information.
1. Balance a chemical equation shows the total number of atoms of reactants

and products molecules.
2. It shows the total number of molecules of each reactant and product.
3. It shows the ratio of molecules of the reactant and the product.
4. It shows the ratio of molecular weight of the reactant and the product.
5. It shows the type of chemical equation.

150 | Chemical Reactions

6. It shows the name and symbols of reactants and product molecules.

For example: Formation of water molecule from hydrogen and oxygen is
given below.

Word equation: Hydrogen + Oxygen Water

Skeletal equation: H2 + O2 H2O
Balanced equation: 2H2 + O2 2H2O
Molecular weight : 4 + 32
36

Ratio of hydrogen to water is 1 : 9

Ratio of oxygen to water is 8 : 9

Type of chemical equation: Addition equation

Limitation of chemical equations
1. Chemical equation does not show the physical state of the reactants and
products.
2. It does not show the conditions required to proceed and to complete the
reaction.
3. It does not show the rate of reaction.
4. It does not show the concentration of reactants and products.
5. It does not show the reversible and irreversible nature of the reactions.
6. It does not show the exothermic and endothermic nature of the reactions.

Modification in a chemical equation
The above limitation are corrected by using the following modification.
1. The physical state of reactants and products are expressed as: Solid is by ‘s’
Liquid is by ‘l’ and Gas is by ‘g’
2. The conditions required for the chemical reaction are kept by writing them
along with the arrow; like heat, light, catalyst, etc.
3. Reversible reactions are identified by a double arrow and irreversible
reactions are identified by a single arrow which can be used in between the
reactants and products.
4. In exothermic reactions, we write heat along with the products and for
endothermic reactions, we write heat along with the reactants.
5. Concentration of the reactants and products are expressed by dilute and
concentrate.

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Catalyst

The chemical substances which increase or decrease the rate of the chemical

reaction are called catalysts. Their concentration remains unchanged till the end of

the reaction. There are two types of catalysts. They are:

a. Positive Catalyst b. Negative Catalyst

Positive Catalyst

The chemical substances which increase the rate of the chemical reaction are

called positive catalysts. For example, MnO2 increases the decomposition of the
potassium chlorate (KClO3) to give potassium chloride (KCl) and oxygen gas(O2).

2KClO3 MnO2 2KCl + 3O2

Negative Catalyst

The chemical substances which decrease the rate of the chemical reaction are
called negative catalysts. For example, glycerine decreases the decomposition of the
hydrogen peroxide (H2O2) to give water (H2O) and oxygen gas (O2).

2H2O2 Glycerine 2H2O + O2

Characteristics of the catalyst
i) Catalysts remain unchanged at the end of the chemical reaction in their
mass and chemical structure.
ii) They do not start the chemical reaction. But they change the rate of the
chemical reaction.
iii) The catalyst is specific to the chemical reaction. It means that one catalyst
can alter the rate of only one chemical reaction.

ANSWER WRITING SKILLS

1. What is chemical reaction? How many ways are there to write chemical
equations?

Ü The chemical change which involves the addition or the decomposition or
exchange of the chemical substances is called chemical reaction. It can be
expressed in two ways.

a. Word equation: When reactants and products are expressed by writing
their full names then it is called a word equation.

152 | Chemical Reactions

For example: Water
Hydrogen + Oxygen

b. Chemical equation: When reactants and products are expressed by using

symbols then it is called a chemical equation.

For example:

2H2 + O2 2H2O

2. Why is it necessary to balance a chemical equation?
Ü According to the law of conservation of mass, the amount of mass can neither be

created nor be destroyed but it remains constant before and after the chemical
reaction. So, to make an equal mass in reactants and products, it is necessary to
balance the chemical equation.

3. Magnetizing and demagnetizing of iron is a physical change. Give suitable
reason.

Ü Magnetizing and demagnetizing of iron is a reversible process, where no new
substance is formed. Hence, it is a physical change.

4. What are the meanings of given symbols? (s), (l), (g) (→), ( ), (↑), (↓), (aq.) and
(D)

Ü The meaning of the above symbols are given below:
(s) = solid state of substance (l) = liquid state of substance

(g) = gaseous state of substance (→) = irreversible reaction

( ) = reversible reaction (↑) = gaseous product

(aq.) = solution in water (D) = heat

(↓) = heavy salt product or residue or (ppt.)

SUMMARY

” The type of temporary and reversible change where no new substances are
formed are called physical changes.

” The irreversible and permanent change in which new substance gets formed
is called a chemical change.

” The chemical change which involves either the addition or the decomposition
or the exchange of the chemical substances is called a chemical reaction.

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” The chemical change which is expressed by writing full names of reactants
and products is called a word equation.

” The chemical change which is expressed by using symbols of reactants and
products is called a chemical equation or formula equation.

” The chemical substances which are participating in the chemical reaction are
called reactants.

” The chemical substances which are obtained after the chemical change are
called products.

” Those chemical reactions which evolve heat during their processing are called
exothermic reactions.

” Those chemical reactions which absorb heat during their processing are called
endothermic reactions.

” Those chemical reactions which can be reversed back from products to
reactants are called reversible reactions.

” The chemical substances which increase or decrease the rate of the chemical
reaction are called catalysts.

” The chemical substances which increase the rate of the chemical reaction are
called positive catalysts.

” The chemical substances which decrease the rate of the chemical reaction are
called negative catalysts.

EXERCISE

1. Define the following terms with example.

(a) Chemical reaction (b) Chemical change (c) Physical change
(d) Reversible and irreversible equation (e) Catalyst

2. Give reason:

(a) Rusting of iron is a chemical change but magnetizing and demagnetizing of iron is
a physical change.

(b) Digestion of food is a chemical change but making the salt solution is a physical
change.

(c) MnO2 is a positive catalyst and glycerine is a negative catalyst.
3. Write two differences between the following.

(a) Chemical and physical change.

154 | Chemical Reactions

(b) Balanced and unbalanced chemical equation.

(c) Exothermic and endothermic reactions.

(d) Word and formula equations.

4. Convert the following word equations into balanced formula equations.

(a) Potassium chlorate → Potassium chloride + Oxygen

(b) Zinc + Sulphuric acid → Zinc Sulphate + Hydrogen

(c) Calcium carbonate → Calcium oxide + Carbon dioxide

(d) Hydrogen peroxide → Water + Oxygen

5. Balance the following equation.

(a) Fe + O2 Fe2O3
(b) P2 + O2 P2O5
(c) H2SO4 + KOH K2SO4 + H2O
(d) H2 + I2 HI
(e) Ca(OH)2 + CO2
6. Answer the following questions: CaCO3 + H2O

(a) What are physical and chemical changes? Write their three characteristics of each.

(b) Write any four examples of each physical and chemical change.

(c) What information do you get from a balanced chemical equation?

(d) What are reactants and products? Explain with an example.

(e) What are catalysts? How many types of catalysts are there? Explain with examples.

A

B GLOSSARY
C

Spontaneous : self

Reversible : that occurs in both directions

Product : resultant

Catalyst : chemicals which alter the rate of the chemical reaction

Positive catalyst : they increase the rate of reaction

Negative catalyst : they decrease the rate of chemical reaction



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UNIT

10 SOLUBILITY

About the Scientist Introduction

Daniel Gabriel Fahrenheit Take a glass of water, add some amount of salt and
(1686-1736) stir it. You will get a salt solution in the water. In this
solution, we cannot see the particles of salt. Again take
a glass of water and add some amount of sand in it. You
will get a mixture of sand and water, where sand particles
are seen clearly in the water. It is also clear that we can
add or mix the mixing components at any ratio. The above
examples are the mixtures of two substances but in both
cases there is a great difference. In the former case, solid
particles are uniformly distributed whereas in the later
case solid particles are not distributed uniformly. On the
basis of the above fact, they are named as homogeneous
and heterogeneous mixtures.

D.G. Fahrenheit was a Homogeneous mixture

German physicist and

instrument maker, though Bring a glass of water, add
some amount of common salt
he lived in Holland for and stir it. After sometimes, you
will get a mixture of common
most of his life. He was a salt in the water. In this mixture,
a new observer cannot identify
merchant by profession, whether we have mixed salt or
sugar. Thus, the mixture of two Homogeneous mixture
but his interest in natural or more substances, in which the mixing components are
distributed uniformly, and if they cannot be identified by
sciences led him to naked eyes they are called homogeneous mixtures. For
example, sugar solution, salt solution, alloy, etc.
conduct plenty of scientific

experiments. In 1714, he

invented the mercury-

in-glass thermometer

and in 1724 designed

the Fahrenheit scale, a

temperature scale in use

even today.

Heterogeneous mixture

Bring a glass of water, add Heterogenous mixture
some amount of sand and stir it.
After sometime, you will get a
mixture of sand in the water. In
this mixture, we can easily identify
water and sand. Thus, the mixture
of two or more substances, in which

156 | Solubility

the mixing components are not distributed uniformly, and if they can be identified
easily they are called heterogeneous mixtures. For example, sand in water, husk in
rice, etc.

Differences between homogeneous and heterogeneous mixture

Homogeneous mixture Heterogeneous mixture

1. It is a mixture of two or more 1. It is a mixture of two or more

substances, in which the mixing substances, in which the mixing

components are distributed components are not distributed

uniformly. uniformly.

2. In the homogenous mixture, the 2. In the heterogeneous mixture, the

components of mixture cannot be components of the mixture can be

separated and recognized by our separated and recognized by our

naked eyes. naked eyes.

Solution

Let’s see the atmosphere. It is a mixture of different types of gases, dust particles,
germs, pollen grains of flowers, etc. Similarly, observe the sugar solution, the salt
solution, the solution of copper sulphate, etc. In all the above examples, two or more
components are mixed uniformly. This uniform mixture is called a homogeneous
mixture. The homogeneous mixture of two or more substances that can be formed
between solid in solid or liquid in solid or liquid in liquid where the components of
the mixture are mixed up to a certain limit it is called a solution. In the given solution,
there must have two components; they are solutes and solvents. In this mixture of
solute and solvent components are distributed uniformly and the size of particles
are so small (10–7cm or less in diameter) that cannot be identified by our naked eyes.
Hence, the solution is a resultant homogeneous mixture of solutes and solvents. For
example; the sugar solution, the salt solution, the solution of copper sulphate, etc. In
the homogeneous mixture of sugar and water, the amount of sugar is less than the
amount of water. Here sugar is called a solute and water is called a solvent. Thus, a
solute is a component of solution which gets dissolved in the solvent and is present
relatively less in amount and the solvent is a component of the solution which is
present relatively large in quantity and dissolves the solute.

Characteristics of a solution
1. A solution is a uniform and homogenous mixture of two or more substances.
2. While making a solution, the mixing components do not lose their own
identity and lie just together.
3. In a solution, the size of solute particles ranges about 10–7cm in diameter or
less than this.
4. It is difficult to separate the components of a solution by applying simple
physical methods.

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Colloid

Observe milk, blood and gum. What do you see there? We can see the simply
homogeneous solution. But, basically they are the examples of heterogeneous
mixtures. In the above mixtures, the size of solute particles is between 10–7 to 10–5 cm
in diameter. Hence, they seem to be homogeneous mixtures.

Colloids are those heterogeneous mixtures which look like homogeneous and
the size of solute particles is in between 10–7 to 10–5 cm in diameter.

In colloids, the solute and solvent particles cannot be separated easily by applying
simple physical methods. Examples of colloids are blood, gum, milk, wax, etc.

Characteristics of colloids
1. Colloids are heterogeneous mixtures but they seem as if they are homogenous
in nature.
2. The range of the size of colloidal particles is between 10–7 to 10–5 cm in
diameter.
3. We cannot separate the components of colloids by applying simple physical
methods. But they can be separated by the ultra-filtration method as well as
the centrifugation method.

MEMORY TIPS
A colloid is a mixture of solid in liquid.

QUESTIONS >>
# Define a solute and a solvent. Write two differences between them.
# Blood is a heterogeneous mixture but salt solution is a homogeneous mixture. Justify the

answer.
# Can we separate a solute and a solvent by filtration? If not give its reason.

Suspension

Observe the atmosphere in the clear sky. What do you see there? Can you see the
fine dust particles which are present there? No, we cannot see the fine dust particles
because it is a heterogeneous mixture but looks like homogeneous. The size of mixing
particles in suspension is about 10–5 cm in diameter or larger than that.

Those heterogeneous mixtures which seem like homogeneous mixtures where
the size of mixing particles ranges between 10–5 cm in diameter or larger than this are
called suspension.

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In suspension, the size of mixing particles is bigger and heavier. So, they can be
settled down easily, and can be separated by the filtration method.

Characteristics of suspension
1. It is a type of heterogeneous mixture but looks like homogeneous.
2. The size of solute particles is 10–5 cm or larger in diameter.
3. The components of a heterogeneous mixture can be separated by the
filtration method. For example: muddy water, dust particles in air, etc.

Types of homogeneous solution

a) Unsaturated solution
Take some amount of common salt and a glass of water. Add little amount of salt

in the water and stir it continuously. Salt gets dissolved easily in water. Add more
salt in the solution. It also gets dissolved. This phenomenon proves that the given
solution can dissolve more solute without changing temperature. Hence, it is called
an unsaturated solution.

The solution which can dissolve more amount of solute without increasing
temperature is called unsaturated solution.

b) Saturated solution
In the above prepared solution, go on adding the salt. What do you get? A stage

will come when no more salt can dissolve in it without increasing the temperature.
This stage of solution is called saturated solution.

The solution which cannot dissolve more amount of solute at particular
temperature is called saturated solution.

Saturated solution has more density as well as more saturation. It can be
precipitated on cooling and becomes unsaturated on heating.

MEMORY TIPS
Saturated solution becomes unsaturated on heating and becomes super saturated
on cooling.

c) Supersaturated solution
In the above saturated solution, we cannot dissolve more solute. But on heating

it can dissolve more amount of solute. This resultant solution no longer remains
saturated but on cooling it tends to emit the excess solute in the form of crystals. This
solution is called super saturated solution.

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The saturated solution at higher temperature is called super saturated solution.

ACTIVITY

Objectives
To identify the given sample of copper sulphate solution with the help of copper sulphate crystals.
Theory
Unsaturated solution can dissolve more amount of solute at given temperature. Saturated solution
cannot dissolve more amount of solute at the same temperature. But, super saturated solution
starts to exit excess solute in form of crystals.
Materials required
Beakers of solution, glass rod, copper sulphate.
Procedure

1. Take three beakers each containing the copper sulphate solution.
2. Add an equal amount of copper sulphate crystals in each of the beakers and observe them.
Observation
Let suppose
• Solution of beaker ‘A’ dissolves the added solute easily on stirring.
• Solution of beaker ‘B’ does not dissolve the added solute.
• Solution of beaker ‘C’ starts to give large size of copper sulphate crystals on adding the

solute.
Result

(A) (B) (C)
The above activity proves that solution ‘A’ is unsaturated solution, solution ‘B’ is saturated solution
and solution ‘C’ is supersaturated solution.

Dilute and Concentrated solution

The amount of solute which is present in a given volume of solvent is called its
concentration. If the amount of solute is less in comparison to the amount of solvent,
it is called dilute solution and it is denoted by (dil.) in the reaction. If the amount of
solute is relatively large in amount, it is called concentrated solution and it is denoted
by (conc.) in the reaction.

160 | Solubility

Differences between dilute and concentrated solution

Dilute solution Concentrated solution

1. Dilute solution contains 1. Concentrated solution contains

comparatively less amount of comparatively more amount of

solute. solute.

2. The density and saturation of dilute 2. The density and saturation of

solution is comparatively less than concentrated solution is comparatively

concentrated solution. more than dilute solution.

Importance of solution

1. Fine root hairs of plants absorb food and minerals from the soil in the
solution form.

2. The nutrients, vitamins and minerals of the digested food are absorbed by
the intestine of our body in the form of solution.

3. The animals and plants which are present in water ( aquatic organisms) take
their food and oxygen from the solution.

4. There are so many chemical reactions which start, continue and are
completed in the solution state.

5. Medicines (specially syrup) are made in the solution state.

Solubility

Take some amount of copper sulphate crystals and a glass with 100 ml of water.
Add this copper sulphate crystals in water and stir it constantly. This solute gets
dissolved in the water easily. Go on adding the solute till no more solute can dissolve.
Now, we will get a saturated solution. The amount of solute which gets dissolved
here is called solubility.

The amount of solute which gets dissolved in 100 gram of solvent at a fixed
temperature to make a saturated solution is called solubility.

Since, we know that different chemical substances have different physical and
chemical properties. That is why they have different solubility. The solubility of a
substance at a particular temperature can be calculated by using the given formula.

Solubility = Weight of solute × 100%

Weight of solvent

A list of solubility of some substances at 20°C is given below

S.N. Substances Solubility at
20°C
1. Common salt (NaCl) 36
2. Copper sulphate (CuSO4 ) 21
3. Sodium nitrate (NaNO3 ) 88

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4. Lead nitrate [Pb(NO 3)2 ] 56.5
5. Sodium sulphate (Na 2SO 4) 44
6. Calcium sulphate (CaSO4 ) 0.20

MEMORY TIPS
Solubility has no unit because it is the simple ratio of two same units.

| SOLVED NUMERICALS |

1. Calculate the solubility of copper sulphate in which 10 gram of it is dissolved
in 25 g of water at 28°C.

Solution: Given, Amount of solute = 10 gram
Amount of solvent = 25 gram
Temperature = 28°C
Solubility = ?

Now, we know that,

Amount of solute in gm
Solubility = Amount of solvent in gm × 100%

10
= 25 × 100%

= 40

Hence, the solubility of copper sulphate at 28°C is 40.

2. Calculate the amount of sugar which is required to make saturated solution in
20 gram of water if the solubility of sugar at 20°C is 204.

Solution: Given, Amount of solvent = 20 gram

Solubility = 204

Temperature = 20° C

Amount of solute = ?

Now, we know that,

Solubility = Amount of solute × 100%

Amount of solvent

x
204 = × 100%

20

204
x = 5 = 40.8 gram
Hence, the amount of solute is 40.8 gram.

162 | Solubility

3. Calculate the amount of water required to make saturated solution of 20 gram
of copper sulphate at 25°C if the solubility of copper sulphate at 25°C in 20.

Solution: Given, amount of solute = 20 gram

Temperature = 25°C

Amount of solvent = ?

Now, we have,
Solubility = Amount of solute in gm × 100%
Amount of solvent in gm

20 = 20 × 100%
x

x = 100 gram

Hence, required amount of water is 100 gram.

4. If 10 gram of saturated solution of copper sulphate at 75°C is cooked down to
25°C. How much amount of copper sulphate will be precipitated if solubility
of copper sulphate at 75°C is 90 and 25°C is 20 respectively?

Solution:

Calculation at higher temperature (i.e. 75°C)

Let,amount of solute = x gram
Amount of solvent = 10 – x gram
Solubility at 75°C = 90
Now, we have

Amount of solute in gm
Solubility = Amount of solvent in gm × 100%

x
90 = 10 — x × 100%
x = 4.7 gram
So, the amount of solvent = 10 – x = 10 – 4.7 = 5.3 gram

Calculation at lower temperature (i.e. 25°C)

Let, amount of solute = y gram
Amount of solvent = 5.3 gram
Solubility at 25°C is = 20
Now, we have,

Amount of solute in gm
Solubility = Amount of solvent in gm × 100%

y
20 = × 100%

3.5

New Creative Science, Class 9 | 163

y = 1.06 gram
Therefore, the amount of copper sulphate precipitated is x – y gram

= 4.7 — 1.06 = 4.64 gram
The amount of copper sulphate precipitated is 3.64 gram.

MEMORY TIPS

Solubility of sodium nitrate at 25°C is 88. It means 88 gram of sodium nitrate gets
dissolved in 100 gram of water at 25°C temperature.

Factors affecting the solubility

The following are the factors that affect the solubility of a substance.

a) Temperature

Temperature plays an important role in the solubility of a substance. As we
increase temperature, the kinetic energy of the solvent molecules increases. As a
result of this kinetic energy, molecules of the solvent move apart creating large inter
molecular space. Hence, a large amount of the solute can be dissolved in the same
amount of the solvent. Likewise, as we decrease the temperature, the kinetic energy
of the solvent molecules decreases. In that condition solvent molecules come near to
each other decreasing the size of intermolecular space. Hence, less amount of solute
can be dissolved.

In a liquid, the solubility of the solid materials increases as we increase
temperature but the solubility of gas decreases. For example, as the temperature
increases, oxygen comes out from the water. There is deficiency of oxygen gas in the
water. Thus, fish and other aquatic animals come to the surface for their breathing.
Similarly, cold drinks release gas as we increase temperature.

S.N. Substances Solubility at 0°C Solubility at Solubility at
20°C 60°C
1. Potassium nitrate 14
2. Sodium chloride 36 21 40

36.5 38

Solubility ∝ Temperature

MEMORY TIPS

Solubility increases with increasing temperature and decreases with decreasing
temperature but it has no role during the calculation of numerical.

164 | Solubility

b) Dimension of solute particles
A small size of solute particles can be fitted easily within the small intermolecular

space. Hence, the rate of solubility is more for finely powdered solute than the bulge.

MEMORY TIPS

Finely powdered solute has greater surface area than the bulge. So, it has a faster
rate of solubility.

c) Thrilling of solution
Thrilling increases the rate of solubility. Fresh portion of the solute particles comes

in contact with the vacant intermolecular space of the solvent. Hence, it increases the
rate of formation of solution.

Solubility curve

Different kinds of substances have different physical and chemical properties.
Solubility of a substance depends upon its physical and chemical properties. That is
why the solubility of different substances is different. Not only this, the solubility of
the same substance differs in different temperatures. Generally, solubility increases
with the rise in temperature and decreases with the fall in temperature. This type
of variation can be studied easily with the help of a graph which is also known as
solubility curve.

The curve which is plotted between solubility and temperature is called a
solubility curve.

Since, the solubility of most of the substances increases with the rise in temperature
but the effect is not the same for all. For example, the solubility of potassium nitrate
and sodium chloride at different temperatures is given below.

S.N. Substances Solubility Solubility Solubility
at 0°C at 20°C at 60°C
21 40
1. P o t a s s i u m 14
nitrate 36.5 38
130

2. S o d i u m 36 120
chloride
110

100
(y-axis) Solubility
Potassium nitrate90

80

The above data shows that the solubility of different 70
substances is different and the effect of temperature is not
the same for both. This type of comparison can be studied 60
with the help of a solubility curve.
50 Magnesium chloride
40
30
20
10 Calcium sulphate

0 10 11 12 13 14 15 16 17 18 19
(x-axis) Temperature in °C

New Creative Science, Class 9 | 165

MEMORY TIPS
If solubility increases with increasing temperature it shows the continuous solubility
curve. If solubility starts to decrease after certain temperature then it gives a
discontinuous solubility curve.

Information obtained from the solubility curve
1. A solubility curve is the best way to calculate the solubility of a substance at
a particular temperature.
2. The variation of solubility along with different temperatures can be shown
by using a solubility curve.
3. With the help of a solubility curve, we can compare the solubility of different
substances at a particular temperature.
4. The amount of precipitate or residue or crystals obtained after cooling the
saturated solution from a higher temperature to the lower temperature can
be calculated with the help of a solubility curve.

Crystals

We can see diamond, quartz, calcite, rock salt, etc. They are the solids which are
made by repeating the fixed shaped particles called crystals.

Those homogenous solid particles which have a definite geometrical shape are
called crystals.

Different substances have different shapes of crystals and they are arranged in a
specific three dimensional pattern.

Characteristics of crystals
1. Crystals are pure substances arranged in a fixed pattern.
2. They have definite geometrical shapes with close a sharp edges.
3. Different crystalline substances have different types of crystals. They are
arranged in a specific three dimensional pattern.
4. They have an accurate melting point.

Crystallization

Prepare a saturated solution of copper sulphate at a definite temperature and
filter it. Allow the filtrate to cool down for a day. Now, observe the solution. What do
you see at the bottom of this container? Here, we can see a number of small and tiny
particles lying at the bottom. These small and tiny particles are called crystals and the
process by which we get crystals is called crystallization.

The process by which we get fine crystals from the saturated solution of a given
substance is called crystallization.

166 | Solubility

To get big crystals the saturated solution at a higher temperature is allowed to
cool down for a few days.

ACTIVITY

Objective
To prepare the crystals of copper sulphate.
Theory
The process by which we get fine crystals from the saturated solution of a given substance is
called crystallization.
Materials required
Copper sulphate, beaker, water, China dish, wire gauze, source of heat, tripod stand, etc.
Procedure
Prepare a saturated solution of copper sulphate at 50°C and let it to cool down for a day after its
filtration. Pick up the crystals and observe them under the microscope.
Observation
Crystals of definite shapes are seen under the microscope. If we want to get big crystals, the
saturated solution is cooled down slowly for a few days.
Conclusion
Crystals are obtained after the process of crystallization.

ANSWER WRITING SKILLS
1. What is supersaturated solution? How does it form?
Ü The type of saturated solution which contains more amount of solute than it

requires is called supper saturated solution. Saturated solution at a higher
temperature is cooled down slowly to get the super saturated solution. Super
saturated solution throws out the excess solute in the form of crystals.
2. What is meant by solubility? Write down its unit. What is the use of temperature
in the numerical problems of solubility?
Ü The amount of solute which can be dissolved in 100 g of solvent at a fixed
temperature is called its solubility.

New Creative Science, Class 9 | 167

Solubility = Amount of solute × 100
Amount of solvent

Since, both the solute and solvent are taken in gram, it has no unit.

During solving the numerical problems, there is no use of temperature but in
general solubility increases with increasing temperature.

3. What is a solubility curve? What information do you get from the solubility
curve?

Ü The graph which is plotted between solubility and temperature is called a
solubility curve. It is different for different substances. The pieces of information
obtained from the solubility curve are given below.
1. A solubility curve is the best way to calculate the solubility of a substance at
a particular temperature.
2. The variation of solubility along with different temperature can be shown
by using a solubility curve.
3. With the help of a solubility curve, we can compare the solubility of different
substances at a particular temperature.
4. The amount of precipitate or residue or crystals obtained after cooling the
saturated solution from a higher temperature to the lower temperature can
be calculated with the help of a solubility curve.

4. You are given three solutions and some amount of copper sulphate. Identify
which one is saturated, unsaturated and supersaturated solution with the help
of this copper sulphate.

Ü Add some amount of copper sulphate crystals in each of the given solutions and
observe it one by one.
i) If the copper sulphate crystals get dissolved easily, it is an unsaturated
solution.
ii) If the copper sulphate crystals do not get dissolved without increasing
temperature, it is a saturated solution.
iii) After adding a little amount of copper sulphate crystals, if we get large sized
crystals in the solution, it is a supersaturated solution.

5. How do you get big crystals during crystallization?

Ü When saturated solution at a higher temperature is cooled down slowly in
the room temperature then we get big sized crystals in the container.

6. How does sodium chloride dissolve in water?

Ü Sodium chloride is an electrovalent compound. It has positively charged sodium
ion and negatively charged chloride ion. Sodium ions get attracted to oxygen
atom of water as it has partial negative charge and chloride ions get attracted
towards positively charged hydrogen atom of water. As a result of this, sodium
and chloride ions are separated from their bond and fit inside the intermolecular
space of the water.

168 | Solubility

7. Show the solution, colloid and suspension with a figure.

SUMMARY

” A mixture is an impure mass obtained by mixing two or more substances in
any quantity by their weight.

” There are two types of mixtures; they are homogeneous mixture and
heterogeneous mixture.

” Solution is a homogenous mixture of solute and solvent.

” Colloid is a mixture in which the size of solute particles is between 10–5 cm to
10–7cm in diameter.

” The size of solute particles in suspension is 10–5 cm or large in diameter.

” Supersaturated solution is prepared at a high temperature and it throws out

the excess solute in the form of crystals.

” The amount of solute which can be dissolved in 100 gram of solvent at a fixed
temperature is called solubility.

Solubility = Weight of solute × 100

Weight of solvent

” Solubility is directly proportional to the temperature. The more the temperature
the more will be the solubility.

” A solubility curve is plotted between temperature and solubility.

” Solubility has no unit because it is the simple ration of two same units.

” Crystals are the homogeneous solid particles which have a definite geometrical
shape.

” Crystallization is the process by which we get fine crystals after cooling the
saturated solution slowly.

New Creative Science, Class 9 | 169

EXERCISE

1. Define the give terms.

(a) Homogeneous mixture (b) Heterogeneous mixture

(c) Solution (d) Colloid

(e) Suspension (f) Supersaturated solution

(g) Concentration of solution (h) Solubility

(i) Solubility curve (j) Crystals

2. Define supersaturated solution. How is it prepared? Describe it in brief.

3. Solubility has no unit. Why?

4. Write two differences between homogeneous and heterogeneous mixtures.

5. What is solution? Write its importance.

6. Write two differences between saturated and unsaturated solutions.

7. Write a short note on the process of crystallization.
8. Milk is a heterogeneous mixture. How?

9. Write two differences between solute and solvent.

10. Compare the size of solute particles in solution, colloid and suspension.

11. Define a solubility curve. What information do you get from it?

12. Write two differences between dilute and concentrated solution.

13. The solubility of sugar is 204 at 25°C. What does it mean.
14. Temperature has no significance during the calculation of numerical. Why?
15. Define the process of crystallization.

16. Write down any three characteristics of solution.

17. Solve the given numerical problems:

(a) Calculate the solubility of sucrose when 50 gram of it is dissolved in 40 gram of
water at 25°C temperature.

(b) Solubility of common salt at 40°C is 35. Calculate the amount of water which is
required to make the saturated solution of 500 gram of common salt.

(c) Compute the solubility of copper sulphate if 25 gram of saturated solution contains
10 gram of copper sulphate crystals at 48°C

(d) 25 gram saturated solution of potassium nitrate at 95°C is cooled down to 55°C, then
how much gram of crystals of potassium nitrate will be separated if the solubility of
potassium nitrate at 95°C is 100 and at 55°C is 25 correspondingly.

A

B GLOSSARY
C

Concentrated : dissolved in less amount of water
Dilute : dissolved in more amount of water
Supersaturated : saturated solution at a high temperature
Crystals : the solid substances which have a de nite shape and a size
Bond : the force of attraction between two or more atoms or ions
Kinetic energy : the energy possess by the moving molecules



170 | Solubility

UNIT

11 SOME GASES

About the Scientist Introduction

Antoine-Laurent de Out of nine planets, the earth is a single one, in which
Lavoisier there is suitable climate which is fit for living beings. Due
to the presence of strong gravity, the mixture of different
Antoine-Laurent de Lavoisier types of gases is present around it. Thus, the mixture of
(26 August 1743 – 8 May 1794) different types of gases which is present around the earth
was born to a wealthy family surface is called atmosphere. In this mixture, the maximum
in Paris, Antoine Laurent percentage is of nitrogen. It includes about seventy eight
Lavoisier inherited a large percentage of the total volume of the atmosphere. The
fortune at the age of five with second largest volume is occupied by oxygen. This is about
the passing of his mother. He twenty one percentage of the total volume. Other gases
attended the College Mazarin like carbon dioxide, argon, neon, etc. are also present in
from 1754 to 1761, studying the atmosphere.
chemistry, botany, astronomy,
and mathematics. His education The above mentioned gases are very important to
was filled with the ideals of the all living beings. Oxygen is used for respiration and
French. combustion, carbon dioxide is used by green plants for
Lavoisier also demonstrated the the process of photosynthesis. The dispersal of seeds, the
role of oxygen in the rusting of transfer of pollen grains and microbes, etc. take place with
metal, as well as oxygen’s role in the help of the atmosphere.
animal and plant respiration.
Working with Pierre-Simon The different gases with their percentage composition
Laplace, Lavoisier conducted is shown in the given table.
experiments that showed that
respiration was essentially a Name of gas Percentage Composition of different gases
slow combustion of organic Nitrogen 78
material using inhaled oxygen. Oxygen 20.9
Lavoisier’s explanation of Carbon 0.03
combustion disproved the dioxide
phlogiston theory, which Argon 0.9
postulated that materials Other gases 0.17
released a substance called
phlogiston when they burned.

New Creative Science, Class 9 | 171

Hydrogen

Symbol =H

Atomic mass = 1 0n 0n
1p 1p

Atomic no =1 Molecular structure of hydrogen
Valency =1

Molecular formula = H2

Hydrogen was observed in the 16th century. When iron responded with
sulphuric acid, a combustible gas was produced; it was hydrogen. In 1766 AD, Henry
Cavendish, an English scientist prepared the hydrogen gas by reacting various
metals with acids. He found that,. it was burning in nature. So, he named it as an
“inflammable air”. Later in 1783 AD, Lavoisier gave the name hydrogen to this gas.

Hydrogen is the simplest element. It is the lightest among the gases present
in the atmosphere. Hydrogen is usually combined with other elements in a variety
of compounds. It is richly present on the earth in combined state. Free hydrogen
is less common because of its inflammability. It is present in the upper part of the
atmosphere in very little percentage. A little amount of hydrogen gas occurs during
the volcanic eruptions and in petroleum productions.

General methods of hydrogen preparation

1. From water

Hydrogen gas can be prepared by the process of electrolysis of water. For
speeding this process, different kinds of acids may use as catalysts.

2H2O 2H2 ↑ + O2 ↑

2. By the reaction of acids with metals

Dilute acids like sulphuric acid, hydrochloric acid, etc. when react with metals
like iron, copper, zinc, etc. give hydrogen gas.

Mg + dil. HCl MgCl2 + H2 ↑

Zn + dil. HSO4 ZnSO4 + H2 ↑

3. By the reaction of metals with alkalis

Aluminium, zinc, etc. when react with alkalis give hydrogen gas.

Zn + 2NaOH Na2ZnO2 + H2 ↑

(Sodium zincate)

172 | Some Gases

4. By the reaction of active metals with water

Active metals like sodium, potassium, magnesium, etc. when react with water at
an ordinary temperature, they give hydrogen gas.

2Na + 2H2O 2NaOH + H2 ↑
2K + 2H2O 2KOH + H2 ↑
Mg + 2H2O Mg(OH)2 + H2 ↑

Laboratory preparation of hydrogen gas

Principle

In the laboratory, when impure zinc granules react with dilute hydrochloric acid
or dilute sulphuric acid we get hydrogen gas.

Zinc + dil. Hydrochloric acid Zinc chloride + Hydrogen

Zn + dil. 2HCl ZnCl2 + H2 ↑
Zinc + dil. Sulphuric acid Zinc sulphate + Hydrogen

Zn + dil. H2SO4 ZnSO4 + H2 ↑

Apparatus required

i) Woulfe’s bottle ii) Thistle funnel iii) Delivery tube

iv) Beehive shelf v) Water trough vi) Gas Jar

Chemicals required
i) Zinc granules (Zn)
ii) Dilute hydrochloric acid or dilute sulphuric acid

Procedure

First the apparatus according to the diagram. Take some pieces of impure zinc
granules and keep them in the woulfe’s bottle and assemble the apparatus in the
correct way. With the help of a thistle funnel, transfer the dilute hydrochloric acid or
dilute sulphuric acid in the woulfe’s bottle. Make sure that the lower end of the thistle
funnel is dipped by an acid. As acid comes in contact with zinc granules, it reacts and
gives hydrogen gas. This gas is passed through the delivery tube and collected in the
gas jar by the downward displacement of water.

New Creative Science, Class 9 | 173

Laboratory preparation of hydrogen gas

Precautions during the laboratory preparation of hydrogen gas
i) The apparatus should be assembled airtight.
ii) Zinc should be in granulated and be in impure form.
iii) Hydrogen is inflammable gas so the experiment should be carried out far
away from the fire.
iv) We should discharge dilute acid till the lower end of the thistle funnel is dipped
in the acid.

Test of hydrogen gas
Hydrogen gas burns with a pale blue flame at the mouth of the gas jar. But the

splinter does not burn inside the gas jar in an atmosphere of hydrogen. So, the flame
is extinguished with pop sound.

Manufacture of hydrogen gas

For commercial use, hydrogen gas should be prepared in large scales in the
industry. That can be done by the following methods.

i) By the reaction of methane gas with steam
The mixture of methane gas and the steam of water is heated strongly in the

presence of about 1200°C temperature, about 30 atmospheric pressure and in the
presence of a nickel catalyst. As a result of this process, we get hydrogen gas along
with carbon monoxide. The mixture of hydrogen gas and carbon monoxide is called
water gas. It is used in industries as a fuel.

174 | Some Gases

Methane + Water vapour 1200°C, 30 atm Carbon monoxide + Hydrogen
Ni
CH4 + H2O 1200°C, 30 atm
Ni CO + H2 ↑

ii) From the electrolysis of water
When we pass electricity in the acidified water,
these water molecules disassociate into opposite
ions. These opposite ions form corresponding
hydrogen and oxygen gas.

H2O Electricity H+ + OH–

Mechanism of getting hydrogen gas

H+ + e— H (Nascent hydrogen) H+ H+

H+H H↑

Physical properties of the hydrogen gas

i. Hydrogen is a colourless, odourless and tasteless gas.
ii. It is neutral in nature. Therefore, it does not change the colour of indicators.
iii. It is the lightest gas of the atmosphere.
iv. It is insoluble or very less soluble in the water.
v. Hydrogen is combustible but not the supporter of combustion. It burns with a

blue flame.
v. Hydrogen can be liquefied below –252°C and can be solidified below – 260°C.

Chemical properties of the hydrogen gas

1. Reaction with non-metals

Different types of non-metals like oxygen, nitrogen, sulphur, etc. react with
hydrogen to give different products. For example;

Hydrogen + Oxygen Heat water

2H2 + O2 2H2O
2NH3
Hydrogen + Nitrogen Ammonia
Hydrogen fluoride
3H2 + N2 500°C, 200 atm
Fe, Mo

Hydrogen + Fluorine

H2+ F2 In dark 2HF

Hydrogen + Chlorine Hydrogen chloride

H2 + Cl2 In light 2HCl

New Creative Science, Class 9 | 175

2. Reaction with metals

Certain metals like sodium, calcium, barium, strontium, etc. combine with
hydrogen to give their corresponding metallic hydrides. For example;

2K + H2 2KH (Potassium hydride)
Ca + H2 CaH2 (Calcium hydride)
Ba + H2 BaH2 (Barium hydride)

iii. Reaction with metallic oxides (Reducing property of hydrogen)

As we pass hydrogen gas over heated metal oxides like zinc oxide, lead oxide,
copper oxide, etc., they get reduced to free metals.

ZnO + H2 Heat Zn + H2O
PbO + H2 Heat Pb + H2O
CuO + H2 Heat Cu + H2O

iv. Hydrogenation

Unsaturated hydrocarbons under suitable conditions, directly combine

with hydrogen to form saturated hydrocarbons. This process is also known as

hydrogenation.

For example:

C2H4 + H2 Ni C2H6

(Ethylene) (Ethane)

Uses of hydrogen gas

i. Since hydrogen is the lightest gas, it was used to fill within the balloons and
ships in previous years.

ii. It is used in the process of hydrogenation to convert vegetable oils into ghee.
iii. It is used as a reducing agent to reduce the metal oxides into free metals.
iv. A hydrogen torch is made using hydrogen gas.
v. In the rockets, hydrogen gas is used as a fuel.
vi. Oxygen and hydrogen flames are used for cutting and welding metals.

ANSWER WRITING SKILLS

1. What is nascent hydrogen?
Ü The atomic form of hydrogen is called nascent hydrogen. It is more reactive and

less stable.

2. What is water gas? Write down its use.
Ü The mixture of carbon monoxide and hydrogen gas is called water gas. It is used

176 | Some Gases

as a source of heat in industries.

3. Why is hydrogen collected by the downward displacement of water?
Ü Hydrogen is the lightest gas and least soluble in water. So, it is collected in a gas

jar by the downward displacement of water.

4. Why should we take impure zinc?
Ü Impurity of the zinc plays an important role to increase the rate of the reaction as

impurities work like a positive catalyst.

EXERCISE

1. What is atmosphere?
2. Write down the principle of the laboratory preparation of hydrogen gas with a balanced

chemical equation.
3. Draw a well labeled diagram of the laboratory preparation of hydrogen gas.
4. Enlist the different types of gases present in the atmosphere.
5. We cannot use concentrated sulphuric acid during the laboratory preparation of

hydrogen gas. Why?
6. How do you test the hydrogen gas in the laboratory? Describe it in brief.
7. Describe any two general methods of manufacturing hydrogen gas with a chemical

reaction.
8. Write down any three physical properties of hydrogen gas.
9. Write down any three uses of hydrogen gas.
10. Hydrogen is collected in a gas jar by the downward displacement of water. Why?
11. Write three chemical properties of hydrogen gas with a balanced chemical equation.
12. We should use impure zinc during the laboratory preparation of hydrogen gas. Why?
13. Write the precautions of the laboratory for preparation of hydrogen gas.
14. Hydrogen is an elemental gas. Why?
15. What happens when(Give chemical equation)?

(a) Hydrogen reacts with ethylene in the presence of a nickel catalyst
(b) Hydrogen gas is passed over the heated copper oxide.
(c) The mixture of hydrogen and nitrogen under pressure is passed over heated iron.
(d) Hydrogen reacts with calcium at temperature of about 300°C.
(e) The mixture of methane gas and water vapour is heated at a temperature of about

1200°C, 30 atmospheric pressure and in pressure of a nickel catalyst.
(f) Hydrogen burns with chlorine gas.

New Creative Science, Class 9 | 177

Oxygen

Symbol =O

Atomic mass = 16 8n 8n
Atomic number =8 8p 8p

Valency =2

Molecular formula = O2 Molecular structure of oxygen

Scheele was the first scientist, who prepared oxygen in 1772 AD. Joseph Priestly
had given its name as a perfect gas and finally Lavoisier coined it oxygen. Oxygen
is one of the lively and plentiful elements found on the earth. About one fifth of the
atmosphere is occupied by oxygen gas. It also occurs in a combined form in many
compounds. The rocks of the earth contain about fifty percentage of oxygen and
water is made up of almost ninety percentage of oxygen. It is also present in fat,
protein, sugar, etc.

General methods of preparation of oxygen gas

(i) By heating metal oxides

A chemical compound made up of oxygen and metal is called metallic oxide.
These compounds after heating give oxygen gas.

2ZnO Heat 2Zn + O2 ↑
2Ag2O Heat 4Ag + O2 ↑
2HgO Heat 2Hg + O2 ↑

(ii) From the electrolysis of water

As we have discussed in the general preparation of hydrogen gas, similarly when
we pass electricity in the acidified water, the water molecules dissociate into opposite
ions. From these opposite ions we get hydrogen and oxygen.

H2O H+ + OH—

Mechanism of oxygen formation

OH— — e— OH

4OH 2H2O + O2 ↑

(iii) By the reaction of metal peroxide with water

Metal peroxides like sodium peroxide, potassium peroxide, etc. give oxygen gas
when they are treated with water.

2Na2O2 + 2H2O 4NaOH + O2 ↑

178 | Some Gases

Laboratory preparation of oxygen gas

Oxygen gas can be prepared in the laboratory by the following two ways.
(A) By using heat
(B) Without using heat

(A) Laboratory preparation of oxygen gas by using heat

Principle

In the laboratory, oxygen gas can be prepared by heating the mixture of potassium
chlorate and manganese dioxide at about 250°C.

Potassium chlorate Manganese dioxide Potassium chloride + Oxygen

2KClO3 MnO2 2KCl + 3O2 ↑

Apparatus required ii) Delivery tube
i) Hard glass test tube iv) Water trough
iii) Beehive shelf

v) Gas jar vi) Bunsen burnervii) Stand viii) Cork

Chemicals required
i) Potassium chlorate (KClO3)
ii) Manganese dioxide (MnO2)

Procedure Stand O2 gas O2 gas Gas jar
KClO3 + MnO2 bubbles
Prepare a mixture of Cork
potassium chlorate and
manganese dioxide at

about the of ratio of 4 : 1

and keep it inside a hard Delivery Water
tube trough
glass test tube or directly

prepare the mixture in Burner
the hard glass test tube.

Arrange the apparatus

according to the above

diagram. The source of heat is adjusted to supply heat. Mostly, nowadays Bunsen

burner is used. As we supply heat, potassium chlorate is decomposed into potassium

chloride and oxygen gas. The produced oxygen gas is delivered through the delivery

tube and collected in the gas jar by the downward displacement of water.

Precautions
i) An apparatus like a hard glass test tube should be airtight.
ii) A pure form of manganese dioxide (MnO2) should be used.
iii) Continuous and low strength flame should be supplied from the Bunsen
burner.

New Creative Science, Class 9 | 179

(B) Laboratory preparation of oxygen gas without using heat
Principle

In the laboratory, oxygen gas can be prepared by the decomposition of hydrogen

peroxide in presence of manganese dioxide.

Hydrogen peroxide Manganese dioxide Water + Oxygen

2H2O2 MnO2 2H2O + O2 ↑

Apparatus required ii) Thistle funnel
i) Conical flask iv) Water trough
iii) Delivery tube vi) Beehive shelf
v) Gas Jar viii) Cork
vii) Stopper

Chemicals required
i) Hydrogen peroxide (H2O2)
ii) Manganese dioxide (MnO2)

Procedure

In a conical flask, keep some amount of pure manganese dioxide. Arrange
the required apparatus as given in the diagram. With the help of a thistle funnel
discharge hydrogen peroxide in the conical flask drop by drop by a using stopper. As
hydrogen peroxide comes in contact with manganese dioxide, it is decomposed into
water and oxygen gas. The resultant oxygen is delivered through the delivery tube
and collected in the gas jar by the downward displacement of water.

180 | Some Gases

Test of the oxygen gas

Place a burning candle or splinter or match stick in the gas jar full of oxygen.
What do you see? The candle or splinter or match stick begins to burn brightly. This
experiment proves that the gas jar is containing oxygen and also shows that oxygen
is the supporter of combustion.

Manufacture of oxygen

In a large scale for commercial purposes oxygen gas can be manufactured in the
following ways.

By the electrolysis of acidified water

When we pass electricity in the acidified water, the water molecules dissociate
into hydrogen and hydroxyl ions. Hydroxyl ions are collected at anode and lose their
extra electron. Now, from the neutral hydroxyl species, we get oxygen gas from the
anode. With the help of a receiver we collect oxygen gas from there.

H2O Electricity H+ + OH—

Mechanism of oxygen formation at anode

OH— — e— OH

4OH 2H2O + O2 ↑

Physical properties of oxygen gas

i. Oxygen is a colourless, odourless and tasteless gas.

ii. It does not have acidic and basic nature. So that, it does not change the
colour of the indicators.

iii. It is a good supporter of combustion. So that all types of firing is possible.
But it itself does not burn.

iv. It is faintly soluble in water. So aquatic life is possible there.

v. It can be converted into liquid below the temperature of about –182°C and
into solid below the temperature of about –218°C.

Chemical properties

i) Combustion

Different types of reactive metals like sodium, potassium, calcium, magnesium,
etc. burn in oxygen to give their respective metallic oxides.

4Na + O2 Heat 2Na2O (Sodium oxide)
4K + O2 Heat 2K2O (Potassium oxide)

New Creative Science, Class 9 | 181

2Mg + O2 Heat 2MgO (Magnesium oxide)
2Ca + O2 Heat 2CaO (Calcium oxide)

ii) Reaction with non-metals

Oxygen reacts with non-metals to give non-metallic oxides.

C + O2 Heat CO2 (Carbon dioxide)
S + O2 Heat SO2 (Sulphur dioxide)

iii) Reaction with oxides of non-metals

Higher oxides are obtained when oxygen reacts with the oxides of non-metals.

2CO + O2 2CO2 (Carbon dioxide)
2SO2 + O2 2SO3 (Sulphur trioxide)

iv) Reaction with hydrogen
Oxygen reacts with hydrogen to give water.

2H2 + O2 2H2O

v) Reaction with organic compounds

Different types of organic compounds such as carbohydrate, protein, fat, etc.
when react with oxygen give water, carbon dioxide and energy.

C6H12O6 + 6O2 6CO2 + 6H2O + Energy

(Glucose)

Uses of Oxygen gas

i. Oxygen is a vital gas. So, all living organisms use oxygen for the process of
respiration. It is said that “no oxygen no life”.

ii. Persons suffering from pneumonia or other diseases are given oxygen.
iii. It is used for combustion of wood, paper, fuel, etc. No oxygen no firing.
iv. In the region of low concentration of oxygen, mountaineers, sea divers,

astronauts, etc. use an oxygen cylinder for breathing.
v. In the metal industry, the cutting and welding of metals are done by using

oxy-hydrogen flames.

ANSWER WRITING SKILLS

1. What is nascent oxygen?

Ü The atomic form of oxygen is called nascent oxygen. It is more reactive and less
stable.

2. Why is oxygen called a vital gas?
Ü There is no life without oxygen. So, it is called a vital gas.

182 | Some Gases

3. Why is oxygen collected by the downward displacement of water?
Ü Oxygen is a lighter gas and is less soluble in water. So, it is collected in a gas jar

by the downward displacement of water.

4. Why should we take MnO2 ?
Ü MnO2 plays an important role to increase the rate of the reaction as it works as a

positive catalyst.

EXERCISE

1. Write down the principle and balanced chemical equation for the laboratory preparation
of oxygen gas without using heat

2. Draw a well labelled diagram for the laboratory preparation of oxygen gas without using
heat.

3. Write down the principle and balanced chemical equation for the laboratory preparation
of oxygen gas by using heat.

4. Draw a well labelled diagram for the laboratory preparation of oxygen gas by using heat.
5. Write down any four physical properties of oxygen gas.
6. Write down any three chemical properties of oxygen gas with a balanced chemical

equation.
7. Describe the manufacturing of oxygen gas.
8. Write any four uses of oxygen gas.
9. How do you test the oxygen gas in the laboratory?
10. Draw the molecular structure of oxygen gas.
11. What happens when (Give balanced chemical equation)?

(a) Sulphur dioxide reacts with oxygen.
(b) Glucose reacts with oxygen
(c) Calcium reacts with oxygen
(d) Phosphorus reacts with oxygen.
(e) Carbon monoxide reacts with oxygen.
12. Give suitable reason:
(a) Oxygen is collected by passing through water.
(b) Manganese dioxide (MnO2 ) is added during the laboratory preparation of oxygen.
(c) It is difficult to fire inside a cave.
(d) Oxygen is collected by the downward displacement of water.
(e) Oxygen is also called a vital gas.

New Creative Science, Class 9 | 183

Nitrogen

Symbol =N

Atomic mass = 14 7n 7n
7p 7p
Atomic number = 7
Molecular Structure of Nitrogen
Valency = Generally 3

Molecular formula = N2
Molecular weight = 28

Many scientists played a vital role in the discovery of nitrogen. Daniel Rutherford,
a Scottish scientist discovered nitrogen in 1772 AD and called it “Phologisticated
air”. It was studied by Scheele and Lavoisier. Later they called it “Azote”. The name
nitrogen was suggested by Chapter in 1823 A.D. In the structure of saltpeter, there
is a nitrogen element. In Greek language saltpeter is called nitro. Therefore from
saltpeter the name of nitrogen has been derived. About 78% of the atmosphere
consists of nitrogen. Thus, the atmosphere constitutes an infinite source of nitrogen.
In addition to free nitrogen in the air, the element is also found in the combined state
as potassium nitrate, which is present in most of the soils and as sodium nitrate,
which is found in huge deposits in Chile and Peru. It also occurs in proteins, which is
a complex organic compound.

General methods of preparation of nitrogen

i. Preparation of nitrogen from nitric oxides

The compounds of nitrogen like nitric oxides are reduced into free nitrogen when
they are passed over heated copper.

2Cu + 2NO Heat 2CuO + N2 ↑

ii. Preparation of nitrogen from bleaching powder

Liquid ammonia when reacts with bleaching powder it gives free nitrogen gas.

3CaOCl2 + 2NH3 Heat 3CaCl2 + 3H2O + N2 ↑

iii. From ammonia gas

When ammonia gas reacts with chlorine gas, it gives ammonium chloride and
nitrogen gas.

8NH3 + 3Cl2 6NH4Cl + N2 ↑

Laboratory preparation of Nitrogen gas

Principle
Nitrogen gas can be prepared by heating the mixture of ammonium chloride and

184 | Some Gases

sodium nitrite in their solution state. The chemical reaction is;

Ammonium chloride + Sodium nitrite heat Sodium chloride + Water +
Nitrogen gas

NH4Cl + NaNO2 Heat NaCl + 2H2O + N2 ↑

MEMORY TIPS

Preparation of nitrogen gas takes place in two steps:

 NH4Cl + NaNO2 Heat NH4NO2 + NaCl
2H2O + N2 ↑
 NH4NO2 Heat

Apparatus required ii. Thistle funnel iii. Delivery tube
i. Round bottom flask v. Beehive shelf vi. Gas jar
iv. Water trough viii. Tripod stand ix. Wire gauze
vii. Bunsen burner
x. Stand

Chemicals required
i. Ammonium chloride (NH4Cl)
ii. Sodium nitrite (NaNO2)

Procedure

Bring ammonium
chloride and sodium
nitrite in an equal
amount and make
a mixture. Keep the
prepared mixture in
a round bottom flask
and adjust all the
required apparatus
as shown in the given
figure. After sometime,
the mixture is directly
prepared in the round
bottom flask. Discharge water with the help of a thistle funnel to make the solution.
After sometime, the mixture and solution is directly made inside the round bottom
flask. In this condition the thistle funnel is not required. Now, supply heat from the
Bunsen burner. As a result of the chemical reaction, nitrogen gas is produced. Since,
nitrogen gas is lighter than air it is delivered through the delivery tube and collected
in the gas jar by the downward displacement of water.

New Creative Science, Class 9 | 185

Precaution

The following points should be remembered during the laboratory preparation
of nitrogen.

i) During the preparation of nitrogen, the apparatus should be airtight.
ii) The required two chemicals are taken in equal ratio.
iii) Continuous and moderate heat should be supplied.
iv) Supply of heat should be stopped when nitrogen gas starts to evolve.

Test of nitrogen gas in laboratory

During the laboratory preparation of nitrogen gas, the following tests are
performed to check whether the collected gas is nitrogen or not.

a) When burning magnesium is introduced in a gas jar of nitrogen, it continues
to burn producing a white ash. If water is added to this ash and heated,
smell of ammonia is felt. This test demonstrates that nitrogen gas is present
in the gas jar.

b) The burning match stick is extinguished in the presence of nitrogen gas.

Manufacturing of nitrogen gas

a) From the air

Nitrogen can be obtained in a large quantity by the fractional distillation of
liquid air. The mixture of air is liquefied by condensing it frequently. The liquid air
is allowed to expand through a spout until it becomes a liquid. When this liquid
is allowed to boil, the more volatile nitrogen (boiling point - 196°C) distils off first
leaving behind the less volatile oxygen (boiling point- 183°C).

b) By passing air over heated copper

In this method, carbon dioxide and
moisture are removed from the air by
passing it through potassium hydroxide
solution (to remove carbon dioxide) and
concentrated sulphuric acid (to remove
water vapour) respectively. Then it
is passed through a strongly heated
combustion tube packed with copper
wire gauze. The heated copper combines
with oxygen to form copper oxide. In this way, oxygen gas can also be separated from
nitrogen.

2Cu + O2 2CuO

Nitrogen thus obtained is not pure. It contains, about 1% by volume of rare gases.

The removal of these gases is not possible by chemical methods.

186 | Some Gases

c) By using pyrogallol and caustic soda

Another method of obtaining atmospheric nitrogen is to remove both carbon
dioxide and oxygen together. This is done by shaking air with a solution of pyrogallol
and caustic soda. The caustic soda absorbs the carbon dioxide and the pyrogallol
absorbs the oxygen. Nitrogen, which is left behind is collected.

Physical properties of nitrogen

i) It is a colourless, odourless and tasteless gas.
ii) It is neutral towards the indicator.
iii) It is a noncombustible gas. It also does not support in combustion.
iv) It is noted for its inertness.
v) It is insoluble or very less soluble in water.
vi) It liquefies at –196°C and solidifies at –210°C.

Chemical properties of nitrogen

Under normal condition, it does not enter into any chemical reaction. This is
why, a large quantity of nitrogen is found in nature.

i) Reaction with metals

Some metals like lithium, calcium, magnesium, aluminium, etc. react with
nitrogen at a high temperature to form metallic nitrides.

3Mg + N2 Mg3N2 (Magnesium nitride)
3Ca + N2 Ca3N2 (Calcium nitride)
6Li + N2 2Li3N (Lithium nitride)
2Al + N2 2AlN (Aluminium nitride)

ii) Reaction with oxygen

Nitrogen does not react with oxygen under normal condition. When a mixture
of nitrogen and oxygen is passed through an electric arc, a reaction takes place. So the
reaction gives nitric oxide.

N2 + O2 2NO (Nitric oxide)

iii) Reaction with hydrogen

Nitrogen and hydrogen do not react under normal conditions. But, when a
mixture of nitrogen and hydrogen at the ratio of 1:3 under the pressure of 200-600

New Creative Science, Class 9 | 187

atmosphere and 500°C temperature is passed over finely divided iron, they combine
to give ammonia gas.

N2 + 3H2 2NH3 (Ammonia)

Uses of nitrogen gas

i) Nitrogen is used for the manufacturing of ammonia gas. This gas is used to
make nitrogenous chemical fertilizers.

ii) Nitrogen is used to fill electric lamps.
iii) It is used in gas thermometers, which are used to measure high temperature.
iv) It is used in aeroplanes to prevent from explosion.
v) It is used to create inert atmosphere as it is about inert in nature.

ANSWER WRITING SKILLS

1. Why is nitrogen gas used to fill the electric bulbs?
Ü Nitrogen behaves like an inert gas. It protects the filament of the bulb from being

oxidized. As a result of this, filament runs for a long time.

2. Why does nitrogen behave like an inert gas?
Ü Nitrogen behaves like an inert gas because it is a triple bonded molecule. Between

two nitrogen atoms, there occurs three covalent bonds (N N). It is very difficult
to break this strong triple covalent bond. Thus, nitrogen molecule behaves like
an inert gas.

3. In the atmosphere, there is about 78 percentage of nitrogen, but plants have
the shortage of nitrogen in their body. Why?

Ü Plants cannot absorb atmospheric nitrogen directly. They absorb nitrogen from
the soil in the form of soluble nitrogenous compounds. Thus, it is necessary to
convert atmospheric nitrogen into nitrogenous compounds like nitrate, nitrite,
etc. This is the reason why plants cannot get nitrogen easily and suffer from
nitrogen deficiency.

4. Why is nitrogen require for plants and animals?
Ü Nitrogen is an important element to make protein, enzyme, nucleic acid, etc. in

plants and animals. So, they require nitrogen.

188 | Some Gases

EXERCISE

1. Write down the principle and balanced chemical equation for the laboratory preparation
of nitrogen.

2. Draw a well-labeled diagram for the laboratory preparation of nitrogen gas.
3. Nitrogen is collected by the downward displacement of water. Why?
4. Write down any three physical properties of nitrogen gas.
5. Write down any three chemical properties of nitrogen with a chemical reaction.
6. Describe one of the important methods of the manufacturing of nitrogen gas.
7. Write any three uses of nitrogen gas.
8. Nitrogen along with argon is used to fill the electric bulbs. Why?
9. Describe the test of nitrogen gas in the laboratory.
10. Draw the molecular structure of nitrogen.
11. Reactivity of nitrogen is very little. Why?
12. Write a balanced chemical equation for the following.

(a) When nitrogen reacts with hydrogen under suitable conditions.
(b) When nitrogen reacts with aluminium.
(c) When nitrogen reacts with oxygen.

A

B GLOSSARY
C

Volatile : that has a less boiling point

Compressible : that can be compressed

In ammable : that can burn



New Creative Science, Class 9 | 189

UNIT

12 METALS

About the Scientist Introduction

Joseph John Thomson More than 114 different elements have been discovered
so far. Many of them are found in nature whereas some of
Joseph John Thomson was them have been made by artificial methods and are called
born in Cheetham Hill, a synthetic elements. Based on their physical and chemical
suburb of Manchester on properties, they have been mainly classified into two
December 18, 1856. He was categories, called metals and non-metals. Among these 114
enrolled at Owens College, elements, more than 80% are metals and the rest of them
Manchester, in 1870, and are non-metals, and a very few are metalloids. Most of the
in 1876 entered Trinity metals are occurring in the form of scums whereas very few
College, Cambridge as a like gold, silver and platinum are present in free form. After
minor scholar. Thomson’s the complete study of these metals and non-metals, they are
early interest in atomic using for the benefit of human beings.
structure was reflected in
his treatise on the Motion Metals
of Vortex Rings, which
won him the Adams Prize Metals are the elements, which are electropositive in
in 1884. His Application nature, malleable, ductile and good conductors of heat and
of Dynamics to Physics electricity.
and Chemistry appeared in
1886, and in 1892 he had Lithium (Li), Sodium (Na), Potassium (K), Magnesium
his Notes. (Mg), Calcium (Ca), Gold (Au), Silver (Ag), Mercury (Hg),
Nickel (Ni), Iron (Fe), Copper (Cu) etc. are some examples
of metals.

Characteristics of metals

(A) Physical characteristics
i) Metals are electropositive in nature, i.e., they lose
electrons.
ii) Metals are generally malleable, i.e., they can be
beaten into thin sheets.
iii) Metals are generally ductile, i.e., they can be drawn
into wires.
iv) Metals in pure state possess luster, i.e. they have a
shining surface.

190 | Metals

v) Metals are generally hard. The hardness varies in different metals.
vi) Metals are good conductors of heat and electricity.
vii) Mostly, metals possess high melting points.
viii) Metals are sonorous, i.e., they produce sounds on striking their surface.
ix) Mostly they have high tensile strength.
x) Generally, metals have high density.
xi) Metals are solid at the room temperature except mercury.

Metallic bond

In all metals, they have incomplete outer shells
and they have loosely bonded electrons. These electrons
jump from one shell to another shell if they get energy.
In the same way, metallic atoms lose their outermost
electrons. After losing electrons, metallic atoms become
positively charged. The lost electrons accumulate and
make the electrons sea. Electrons sea has a negative
charge and metal atoms have a positive charge. Between
these opposite charges, there occurs a force of attraction which is called a metallic
bond. This metallic bond is responsible to hold the atoms in the metal.

Why do metals shine?
In all metals, they have incomplete outer shells and they have loosely bonded
electrons. The loosely bonded electrons absorb energy and jump in the higher shells.
But these electrons do not remain in the higher shells for a long time. They come back
in the same shell after releasing the absorbed energy. Due to this released energy,
metals shine.

Why are metals malleable in nature?
Metallic atoms lose their outermost electrons. After losing electrons, metallic atoms
become positively charged. As we hammer them with the help of a hammer, these
positively charged metallic atoms come close to each other. But, due to the opposite
charge, these atoms repel to each other and try to move apart. In this process, in every
strike, atoms try to get a new place. As a result, metals turn into thin sheets.

Why are metals ductile in nature?
As we stretch metals, positively charged metal atoms along with a required amount
of electrons from the electron sea also move. Finally, the long and thin wire also has
the metallic character.

Metals have a high melting and boiling point. Why?

Metals have strong a metallic bond. As a result of this strong metallic bond, all the
atoms of the metals are strongly held to each other. To break this metallic bond, a large
amount of energy is required. Thus, metals have a high melting and boiling point.

New Creative Science, Class 9 | 191

ACTIVITY

Take two objects, an iron sheet and soap. Take an iron nail and scratch at over the iron sheet. It is
hard to scratch the iron sheet. Flow the same process to scratch the soap. It is easy to scratch the
soap. The soap also has a solid shape but not as hard as the sheet of iron.

ACTIVITY

Take a knife. Observe its shinning area. It has a plastic or wooden handle and a metal surface.
Which one shines more, non-metals or metals?

ACTIVITY

Take two rods; one made up of metal and another made of wood. Heat one end of these rods for
some time with a burner. Which one becomes red and hot? Which rod heat more while touching
it? Share your answer with your friends.

ACTIVITY

Connect two dry cells with each other by using a copper wire. Fix a bulb with two ends of the wire.
Break a wire near the dry cell as shown in the figure. Place different materials like iron nails, coins,
plastic, carom coins, wooden blocks in the breaking points of the wire, so that it joins two wires.
The bulb in the circuit glows on keeping the coins and iron nails in the breaking point. Why does n’t
the bulb glow when carom coins, wooden blocks and a plastic wire are used?

192 | Metals

ACTIVITY

Take an iron nail and a glass bottle. Hit them with a hammer. What do you find? Does the iron nail
become flat on hitting it for a long time with the hammer? What happens when a glass bottle is
hammered?

ACTIVITY

Take a metallic rod and a wooden rod of the same size. Hold these rods one by one from their
endpoints and try to break them? Which one breaks easily, the wooden or metallic rod?

ACTIVITY

Take two plates, one metallic plate and another china clay plate. Strike the plate with the same
rod. The metallic plate produces a sharp and twinkling sound but a wooden or china clay plate
produces a blunt and hoarse sound.

(B) Chemical characteristics
i) Most of the metals react with oxygen to give metallic oxides.

2Mg + O2 2MgO (Magnesium oxide)

4Na + O2 2Na2O (Sodium oxide)

ii) Most of the metal oxides are insoluble in water.

iii) Most of the metal oxides are basic in nature. They dissolve in water to give

an alkaline solution. However, some are amphoteric in nature.

iv) Metals react with halogen to give salt.

2Na + Cl2 2NaCl

New Creative Science, Class 9 | 193

v) Metals react with hydrogen to give metal hydrides.

2Na + H2 2NaH (Sodium hydride)

vi) More reactive metals displace less reactive metals from their salt solution.
For example,

Fe + CuSO4 FeSO4 + Cu

Non-metals

The elements which are electronegative in nature, non-conductors of heat and
electricity, non-malleable and non-ductile in nature and do not possess luster are
called non-metals.

Some examples of non-metals are Hydrogen (H), Oxygen (O), Nitrogen (N),
Chlorine(Cl), Phosphorus (P), Carbon (C), Bromine (Br), Sulphur (S), etc.

Characteristics of non-metals

(A) Physical characteristics
i) Non-metals are electronegative in nature, i.e.; they gain electrons.
ii) They are non- malleable.
iii) They are non-ductile.
iv) They are generally bad conductors of heat and electricity except graphite.
v) Non-metals do not possess luster except iodine.
vi) They are soft and brittle, i.e., they break into pieces when hammered except
diamond.
vii) They are non-sonorous.
viii) They generally have low melting and boiling points except boron, graphite
and diamond.
ix) Generally, they have low density.
x) They have low tensile strength.
xi) They may be solid, liquid and gas at the room temperature.

(B) Chemical characteristics of non-metals
i) Non-metals form oxides when heated with oxygen.

C + O2 CO2
S + O2 SO2

194 | Metals

ii) They form acidic oxides. When they are dissolved in water they give acid.

CO2 + H2O H2CO3 (Carbonic acid)

SO2 + H2O H2SO3 (Sulphurous acid)

iii) Generally, they do not react with water.

iv) They do not react with acid to produce hydrogen gas.

v) They react with chlorine to form chlorides

H2 + Cl2 2HCl (Hydrogen chloride)

P4 + 6Cl2 4PCl3 (Phosphorus trichloride)

vi) They combine with hydrogen to form covalent hydrides.

8H2 + S8 8H2S (Hydrogen sulphide)

N2 + 3H2 2NH3 (Ammonia)

vii) A more reactive non-metal displaces the less reactive non-metal from its

salt solution.

2NaBr + Cl2 2NaCl + Br2

Metalloids

Besides metals and non-metals, there is a third category of elements which show
the properties of both metals and non-metals. These elements are called metalloids.
Silicon (Si), Germanium (Ge), Arsenic (As), Antimony (Sb) , tellurium (Te), etc. are the
examples of metalloids.

MEMORY TIPS
(i) Only 22 are non-metals.
(ii) Silver is the best conductor of heat and lead is the poorest conductor of heat.

The sequence of conductivity is given below.

Ag > Cu > Au > Al > W > Hg
(iii) Tungsten (W) has highest melting point

Alloys

We know that iron is the most widely used metal. But it is not used in the
pure state. This is due to this reason that iron is a very soft metal. It stretches easily.
However, if it is mixed with a small amount of carbon (0.05%), it becomes hard and
strong. This mixture is called steel. If instead of carbon, iron is mixed with nickel and

New Creative Science, Class 9 | 195

chromium, we get stainless steel. It is very hard and does not get rust. Thus, we can
conclude that the properties of any metal can be changed if it is mixed with some
other metals or non-metals. This mixture is called an alloy.

The homogeneous mixture of two or more metals or metals and non-metals is
called an alloy.

Why are alloys prepared?

Alloys are prepared to meet the given properties
i) To increase the hardness.
ii) To increase the tensile strength.
iii) To lower the melting point.
iv) To modify the chemical reactivity.
v) To reduce the electrical conductivity.
vi) To modify colour.
vii) To increase the resistance to corrosion.

MEMORY TIPS

An alloy which contains the mercury as one of the constituents is called amalgam.
For example, zinc amalgam, sodium amalgam, etc.

Some common alloys and their composition

Alloy Composition Properties Uses

1. Steel Iron (99.95%), Hard, tough and Construction of
Carbon (0.05%) strong ships, bridges,
vehicles, etc.

Iron (74%), Hard and does not For making cutlery,
2. Stainless steel Chromium (18%), get rust utensils and surgical
instruments.
Nickel (8%)

Copper (80%), Zinc Malleable, strong, For making utensils,
(20%)
3. Brass resists corrosion, can screws, nuts and

be easily cast bolts

4. Bronze Copper (90%), Tin Very strong and For making statues,
(10%) highly resistant to coins, medals, ship’s
corrosion propeller etc.

5. Solder Lead (50%), Tin Has lower melting For joining electrical
(50%) point than either wires together.
lead or tin

196 | Metals