Green Science 10 Final PDF (2076)

Magnesium + Chloride Magnesium chloride

Calcium carbonate heat Calcium oxide + Carbon dioxide

Zinc + Sulphuric acid Zinc sulphate + Hydrogen

Potassium Chlorate heat Potassium chloride + Oxygen

Formula Equation

A formula equation or chemical equation is the chemical reaction expressed by writing
symbols and molecular formulae of reactants and products. It is more informative than a
word equation. Some examples of formula equations are given below:

2Na + Cl2 2NaCl
2H2 + O2
Mg + Cl2 2H2O
CaCO3 MgCl2
Zn + H2SO4 CaO + CO2
2KClO3
ZnSO4 + H2
2KCl + 3O2

Reactants and Products

Reactants are chemical substances which take part in a chemical reaction. They are written
on the left hand side of the arrow.

Products are the chemical substances which are produced after a chemical reaction. They
are written on the right hand side of the arrow.

Examples ZnCl2 + H2 Do You Know
Zn + HCl (Products)
(Reactants) Reactants are written on the left hand side
HCl + NaOH NaCl + H2O and products are written on the right hand
(Reactants) (Products) side of the arrow.

Unbalanced Chemical Equation

The chemical equation in which the total number of atoms of each element in reactants
and products are not equal is called an unbalanced chemical equation.

Examples H2O
H2 + O2 ZnCl2 + H2
Zn + HCl Na2SO4 + H2O

NaOH + H2SO4

GREEN Science (Chemistry) Book-10 151

Balanced Chemical Equation

A balanced chemical equation can be defined as the chemical equation written by balancing
the total number of atoms of each element in reactants and products.

Examples 2H2O
2H2 + O2 ZnCl2 + H2
Zn + 2HCl Na2SO4 + 2H2O

2NaOH + H2SO4

Methods of Writing Balanced Chemical Equation

1. First of all, a chemical change is written correctly in the form of word equation.

For example: Sodium + Chlorine Sodium chloride

2. A word equation is written correctly in the form of formula equation.

For example: Na + Cl2 NaCl

3. The number of atoms of each element are balanced by using suitable coefficient

without changing the molecular formulae of reactants and products.

For example: 2Na + Cl2 2 NaCl

Information obtained from a Balanced Chemical Equation

We can get the following information from a balanced chemical equation
1. The names of reactants and products
2. Symbols and molecular formulae of reactants and products
3. Total number of atoms or molecules of reactants and products
4. Type of chemical reaction
5. Atomic weight and molecular weight of reactants and products

Types of Chemical Reaction

There are different types of chemical reaction on the basis of process of formation of
products from the reactants. Some of the common types of chemical reactions are as
follows:
1. Combination reaction 2. Decomposition reaction
3. Displacement reaction 4. Acid- base reaction

1. Combination reaction

When carbon (C) burns in air (oxygen), it forms carbon dioxide (CO2). In the chemical
reaction, carbon (C) combines with oxygen (O2) and forms a product carbon dioxide

152 GREEN Science (Chemistry) Book-10

(CO2). This type of chemical reaction is called combination reaction. The chemical reaction
in which two or more reactants combine to form a single product is called combination
reaction. It is also called addition or synthesis reaction. Various factors like heat, light,
pressure, etc. are responsible for combination reaction.

Some examples of combination reaction are given below:

1. Carbon + Oxygen burn Carbon dioxide

C + O2 CO2 [Combination]

2. Potassium + Chlorine Potassium chloride

2K + Cl2 2KCl [Combination]

3. Sodium + Chlorine Sodium chloride

2Na + Cl2 2NaCl

4. Aluminium + Nitrogen Aluminium nitride

2 Al + N2 2AlN

5. Nitrogen + Hydrogen Ammonia

N2 + 3H2 2NH3
6. Iron + Sulphur Iron (Ferrous) sulphide

Fe + S FeS

7. Iron + Oxygen Ferric (iron) oxide

4Fe + 3O2 2Fe2O3

8. Calcium carbonate + Carbon dioxide + Water Calcium bicarbonate

CaCO3 + CO2 + H2O Ca(HCO3)2
9. Sodium + Oxygen Sodium Oxide

4Na + O2 2Na2O

10. Hydrogen + Oxygen Water

2H2 + O2 2H2O
11. Iron + Sulphur Ferrous Sulphide

Fe + S FeS

GREEN Science (Chemistry) Book-10 153

2. Decomposition reaction

When calcium carbonate (CaCO3) is heated, Do You Know
it breaks down into two products, viz.
calcium oxide (CaO) and carbon dioxide Decomposition reaction is also called
(CO2). Here, a single reactant decomposes dissociation or analysis reaction.
into two products due to action of heat.

When lead nitrate [Pb(NO3)2] is heated, it decomposes into three products, viz. Lead oxide
(PbO), Nitrogen dioxide (NO2) and Oxygen (O2). These are the examples of decomposition
reaction. So, the chemical reaction in which a single reactant decomposed into two or

more products is called decomposition reaction. Various factors like heat, light, electricity,

catalyst, etc. are responsible for decomposition reaction.

1. Water electricity Hydrogen + Oxygen

2H2O 2H2 + O2

2. Calcium carbonate heat C alcium oxide + Carbon dioxide

CaCO3 D CaO + CO2

3. Ammonium hydroxide heat Ammonia + Water

NH4OH NH3 + H2O

4. Copper carbonate heat Copper oxide + Carbon dioxide

CuCO3 D CuO + CO2
5. Potassium chlorate heat Potassium chloride + Oxygen

2KClO3 D 2KCl + 3O2
6. Hydrogen peroxide MnO2 W ater + Oxygen
2H2O2 MnO2 2H2O + O2

7. Silver nitrate heat Silver + Nitrogen dioxide + Oxygen

2AgNO3 D 2Ag + 2NO2 + O2

8. Copper nitrate heat Copper + Nitrogen dioxide + Oxygen

2Cu (NO3)2 D 2Cu + 2NO2 + O2

9. Lead nitrate Lead monoxide + Nitrogen dioxide + Oxygen

2Pb (NO3)2 2PbO + 4NO2 + O2
10. Silver oxide Silver + Oxygen

2Ag2O 4Ag + O2

154 GREEN Science (Chemistry) Book-10

3. Displacement reaction

When zinc (Zn) reacts with dilute sulphuric Do You Know
acid (H2SO4), zinc displaces Hydrogen
of the acid, i.e. Sulphuric acid and forms Displacement reaction is also called

Zinc sulphate (ZnSO4) and Hydrogen replacement reaction.
gas (H2). This type of chemical reaction is
called displacement reaction. The chemical

reaction in which an atom or radical of a compound is displaced by another element to

form new products is called displacement reaction.

Displacement reaction is of two types, viz. single displacement reaction and double
displacement reaction.

i. Single displacement reaction

It is the displacement reaction in which one atom or a radical is displaced by another
element. Some examples of single displacement reaction are as follows:

1. Potassium + Hydrochloric acid Potassium chloride+ Hydrogen

2K + 2HCl 2KCl + H2

2. Magnesium + Hydrochloric acid Magnesium chloride + Hydrogen

Mg + 2HCl MgCl2 + H2

3. Calcium + Sulphuric acid Calcium sulphate + Hydrogen

Ca + H2SO4 CaSO4 + H2

4. Iron + Copper sulphate Iron sulphate + Copper

Fe + CuSO4 FeSO4 + Cu

5. Magnesium + Zinc chloride Magnesium chloride + Zinc

Mg + ZnCl2 MgCl2 + Zn

ii. Double displacement reaction

It is the chemical reaction in which an atom or radical of a compound is mutually displaced
by a radical or an atom of another compound. Some examples of double displacement
reaction are as follows:

1. Sodium chloride + Silver nitrate Sodium nitrate + Silver chloride

NaCl + AgNO3 NaNO3 + AgCl
or, Na+ Cl– + Ag+ NO3–

2. Magnesium chloride + Silver nitrate Magnesium nitrate + Silver chloride

MgCl2 + 2Ag NO3 Mg(NO3)2 + 2AgCl

GREEN Science (Chemistry) Book-10 155

3. Sodium hydroxide + Ferrous chloride Sodium chloride + Ferrous hydroxide

2NaOH + FeCl2 2NaCl + Fe(OH)2

4. Calcium chloride + Silver nitrate Calcium nitrate + Silver chloride

CaCl2 + 2Ag NO3 Ca(NO3)2 + 2AgCl

5. Sodium sulphate + Lead nitrate Lead sulphate + Sodium nitrate

Na2SO4 + Pb (NO3)2 PbSO4 + 2NaNO3

6. Magnesium nitride + Water Magnesium oxide + Ammonia

Mg3N2 + H2O 3MgO + 2NH3

7. Mercuric chloride + Potassium iodide Potassium chloride + Mercuric iodide

HgCl2 + 2KI 2KCl + HgI2

4. Acid - base reaction

When Hydrochloric acid (HCl) reacts Do You Know
with Sodium hydroxide (NaOH), it

forms Sodium chloride (NaCl) and Acid-base reaction is also called neutralization

water (H2O). It is an example of an acid- reaction. However, all acid-base reactions are
base reaction. The chemical reaction not neutralization reaction.

which takes place between an acid and

a base to form salt and water is called acid-base reaction.

Some examples of acid-base reaction are given below:

Acid + Base Salt + Water

1. Hydrochloric acid + Potassium hydroxide Potassium chloride + Water

HCl + KOH KCl + H2O

2. Nitric acid + Potassium hydroxide Potassium nitrate + Water

HNO3 + KOH KNO3 + H2O

3. Sulphuric acid + Sodium hydroxide Sodium sulphate + Water

H2SO4 + 2NaOH Na2SO4 + 2H2O

4. Sulphuric acid + Calcium oxide Calcium sulphate + Water

H2SO4 + CaO CaSO4 + H2O

5. Sulphuric acid + Ferrous oxide Ferrous sulphate + Water

H2SO4 + FeO FeSO4 + H2O

6. Acetic acid + Sodium hydroxide Sodium acetate + Water

CH3COOH + NaOH CH3COONa + H2O

156 GREEN Science (Chemistry) Book-10

7. Hydrochloric acid + Ammonium hydroxide Ammonium chloride + Water

HCl + NH4OH NH4Cl + H2O

In acid-base reaction, both acid and base lose their properties and form two neutral

substances, viz. salt and water. Therefore, acid-base reaction is also called neutralization

reaction.

Factors that bring out chemical reaction

Various conditions are required to bring out chemical reaction. Some of them are discussed
below:

1. Simple contact

Some chemical reactions take place when the reactants are brought in contact. For
example, when sodium is brought in contact with chlorine, chemical reaction takes place.
As a result, sodium chloride is formed.

Na + Cl2 2NaCl

2. Heat

Some chemical reactions take place when reactants are heated. Heat energy increases the
kinetic energy of the molecules of reactants which brings the reacting molecules in close
contact. As a result, chemical reaction takes place. For example,

CaCO3 CaO + CO2
2KClO3 2KCl + 3O2

3. Light

Some chemical reactions take place when the reactants are exposed to light. Light energy
makes reactant molecules active which brings out chemical change. For example,

H2 + Cl2 Sunlight 2HCl

4. Electricity

Some chemical reactions take place when electricity is passed through solution state
or fused state of reactants. Electricity helps the ions move towards oppositely charged
electrodes and chemical reaction takes place. For example, when water is electrolyzed, it

decomposes into hydrogen and oxygen.

2H2O 2H2 + O2

5. Pressure

Some chemical reactions take place when the reactants are kept under certain pressure.
Pressure brings the reacting molecules closer and chemical reaction takes place. For
example,

3H2 + N2 200 – 900 atm (pressure) 2NH3
500°C/Fe, Mo

GREEN Science (Chemistry) Book-10 157

6. Solution state

Some chemical reactions take place when reactants are mixed in solution state. For
example, when Sodium Chloride and Silver Nitrate are mixed in solid state, chemical
reaction does not occur. But they react in solution state.

NaCl + AgNO3 AgCl + NaNO3

7. Catalyst

Some chemical reactions take place only in the presence of catalyst. For example, Hydrogen

peroxide (H2O2) decomposes in the presence of catalyst Manganese dioxide (MnO2).

2H2O2 MnO2 2H2 + O2
(Catalyst)

Factors Affecting the Rate of Chemical Reaction

The rate of different chemical reaction is different. Some chemical reaction occur very fast

whereas some chemical reaction occur very slowly. Some reactants undergo chemical

reaction when they come in contact whereas some chemical reactions do not take place

without supplying heat, light, electricity,

catalyst, etc. The rate of a chemical Do You Know
reaction depends on the concentration
physical nature and chemical nature of The rate of chemical reaction is defined as the
reactants. Various factors increase or positive change in concentration of a reactant
decrease the rate of a chemical reaction. or a product per unit time.

Some of the major factors that affect the

rate of chemical reaction are given below:

1. Temperature 2. Light 3. Surface area 4. Pressure 5. Catalyst

1. Temperature

The rate of chemical reaction increases on increasing the temperature of reactants. Similarly,
the rate of chemical reaction decreases on decreasing the temperature of reactants. More
heat is supplied to the reacting molecules while increasing the temperature. It provides
more kinetic energy to reacting molecules and frequency of collision of these molecules
increases to give more products. Increase in temperature increases the rate of dissociation
and recombination of reacting molecules. As a result, the rate of chemical reaction
increases. Many chemical reactions do not occur without heating the reactants to a certain
temperature.

Example
2KClO3 360°C 2KCl + 3O2

158 GREEN Science (Chemistry) Book-10

Experiment 1

Objective : To demonstrate that the rate of chemical reaction increases on increasing the
temperature of reactants.

Materials required : Beakers, dilute sulphuric acid, water, potassium permanganate,
potassium thiosulphate, oxalic acid (aq), spirit lamp, stand
Procedure

• Take two beakers and keep one crystal of oxalic acid in each of them.

• Now, add about 10ml of dilute sulphuric acid in each beaker.

• Add 5 ml of potassium permanganate solution in each beaker.

• Stir the solution in each beaker using a glass rod and observe the solution carefully.

• You can see that the solution in both beakers appear pink.

• Now, heat one of the beakers upto 60° – 80° C with the help of spirit lamp.

• Observe carefully, in which beaker does the pink colour disappear faster?

Observation

The pink colour of the solution in the beaker disappears which is heated after some
time. But the pink colour of the solution does not disappear in the beaker which is not
heated. It shows that chemical reaction occurs faster at high temperature.

Conclusion

From this activity, it can be concluded that the rate of chemical reaction increases on
increasing the temperature of reactants.

2. Light

Many chemical reactions take place in the presence of sunlight. The rate of a chemical
reaction increases in the presence of light. Some examples of chemical reactions that take
place in the presence of light (sunlight) are as follows.

i. Photosynthesis in green plants

6CO2 + 6H2O Sunlight C6H12O6 + 6O2

ii. Formation of hydrochloric acid
H2 + Cl2 Sunlight 2HCl
iii. Dissociation of silver bromide
2AgBr Sunlight 2Ag + Br2
iv. Chlorination of methane

CH4 + Cl2 CH3Cl + HCl

v. Formation of Ozone from oxygen

3O2 UV rays 2O3

GREEN Science (Chemistry) Book-10 159

3. Surface area

The rate of a chemical reaction increases if the contact area of reacting molecules is more
and vice-versa. If the contact area is more, many reacting molecules come in contact. As
a result, the rate of chemical reaction increases. The area of contact can be increased by
either of the following methods.

i. By breaking down reactants into small pieces

ii. By using the reactants in powdered from

iii. By using a common solvent

Experiment 2

Objective : To demonstrate that the rate of chemical reaction increase on increasing the
surface area of reacting molecules.

Materials required : Zinc powder, Zinc granules, dilute Hydrochloric acid, beaker (2),
measuring cylinder, top pan balance, watch glass (2), glass rod

Procedure

• Take a measuring cylinder and measure 25 ml of dilute hydrochloric acid.
• Keep 25/25 ml of dilute hydrochloric acid in two beakers.
• Take a top pan balance and measure 2.5 gram of zinc granules and 2.5 gram of zinc

powder. Keep zinc power and zinc granules in a separate watch glass.
• Now, pour the zinc powder into a beaker and zinc granules into another beaker

simultaneously. Stir the mixture by using a glass rod.
• Observe the chemical reaction that occurs in the beakers carefully. In which beaker

does the reaction complete faster?
Observation

The bubbles of gas (hydrogen) evolve earlier in the beaker having zinc powder than
in the beaker having zinc granules. Similarly, chemical reaction completes faster in
the beaker having zinc powder. It shows that chemical reaction takes place faster in
powdered from than the granules, because the area of contact is more in powdered
form than that of the granules.

From this activity, it can be concluded that the rate of a chemical reaction increases on
increasing the area of contact of the reacting molecules.

4. Pressure

The rate of chemical reaction of gases molecules depends on the pressure of the reactants.
The rate of chemical reaction increases on increasing the pressure of the reacting molecules
and vice-versa. Some gases react only in the high pressure of the reacting molecules
and vice-versa. Some gases react only in the high pressure. Increase in pressure brings
the molecules of reacting gases closer and the rate of chemical reaction increases. Some
examples of the chemical reactions that take place due to application of pressure are as
follows.

i. Formation of ammonia gas by Haber's process

N2 + 3H2 2NH3

160 GREEN Science (Chemistry) Book-10

ii. When pressure is applied on the mixture of sulphur and potassium chloride, it
explodes.

iii. Fire crackers get exploded on applying pressure.

5. Catalyst

The chemical substance which is used to increase or decrease the rate of a chemical
reaction is called a catalyst. A catalyst remains chemically unchanged throughout the
chemical reaction but its presence may increase or decrease the rate of a chemical reaction.

The catalyst which increases the rate of a chemical reaction is called positive catalyst. It is
used to speed up the rate of chemical reaction.

A positive catalyst decreases or lowers the energy required to break down the chemical
bond of the molecules of reactants. As a result, the rate of chemical reaction increases.

Examples of positive catalyst

i. Manganese dioxide (MnO2) acts as a positive catalyst during decomposition of
Hydrogen peroxide. (H2O2)

2H2O2 MnO2 2H2O + O2

ii. Manganese dioxide (MnO2) acts as a positive catalyst during decomposition of
Potassium chlorate (KClO3). It means that MnO2 speeds up the decomposition of
KClO3.
2KClO3 MnO2 2KCl + 2O2

The catalyst which decreases the rate of a chemical reaction is called a negative
catalyst. It is used to slow down the rate of a chemical reaction.

Example negative catalyst
Glycerol [C3H5(OH)3] acts as a negative catalyst during decomposition of Hydrogen

peroxide (H2O2). It means that glycerol slows down th decomposition of hydrogen
peroxide.
2H2O2 Glycerol 2H2O + O2

Characteristics of catalyst

1. The mass of a catalyst does not change till the end of the chemical reaction.
2. A catalyst remains chemically unchanged throughout the chemical reaction.
3. A catalyst does not initiate a chemical reaction but increases or deceases the rate of

chemical reaction.

Endothermic reaction and Exothermic reaction

Most chemical reactions occur due to change in heat. Some chemical reactions absorb heat
whereas some chemical reactions evolve heat during the chemical reaction. On this basis,
there are two types of chemical reaction. They are : Endothermic reaction and exothermic
reaction.

GREEN Science (Chemistry) Book-10 161

The chemical reaction that absorbs heat during the chemical change is called endothermic
reaction.

Examples

N2 + O2 + Heat 2NO

CaCO3 + Heat CaO + CO2

2NaCl + Heat 2Na + Cl2

NH4Cl + NaNO2 + Heat NaCl + 2H2O + N2

2KClO3 + Heat 2KCl + 3O2

The chemical reaction that evolves heat during the chemical change is called exothermic
reaction.

Examples CO2 + Heat
C + O2 CH4 + Heat
C + 2H2 ZnCl2 + H2 + Heat
Zn + 2HCl Ca(OH)2 + Heat
CO2 + 2H2O + Heat
CaO + H2O
CH4 + 2O2

Difference between Endothermic and Exothermic reaction

Endothermic Reaction Exothermic Reaction

1. Heat is absorbed during chemical 1. Heat is evolved during chemical

reaction. reaction.

For example, For example,

2KClO3 + Heat 2KCl + 3O2 C + O2 CO2 + Heat

Key Concepts

1. The combination, decomposition or replacement that occurs in the molecules of matter
during a chemical change is called chemical reaction.

2. The chemical bond present in the molecules of reactants breaks due to heat, light,
electricity, etc. during a chemical change.

3. Losing, gaining or sharing of electrons by an atom to gain stable electronic
configuration is the major cause of chemical reaction.

4. The chemical equation in which the total number of atoms of each element in
reactants and products are not equal is called an unbalanced chemical equation.

5. A balanced chemical equation can be defined as the chemical equation written by
balancing the total number of atoms of each element in reactants and products.

6. The chemical reaction in which two or more reactants combine to form a single
product is called combination reaction. It is also called addition or synthesis reaction.

162 GREEN Science (Chemistry) Book-10

7. The chemical reaction in which a single reactant decomposed into two or more
products is called decomposition reaction.

8. The chemical reaction in which an atom or radical of a compound is displaced by
another element to form new products is called displacement reaction.

9. Double displacement reaction is the chemical reaction in which an atom or radical
of a compound is mutually displaced by a radical or an atom of another compound.

10. The chemical reaction which takes place between an acid and a base to form salt
and water is called acid-base reaction.

11. In acid-base reaction, both acid and base lose their properties and form two
neutral substances, viz. salt and water. Therefore, acid-base reaction is also called
neutralization reaction.

12. The rate of different chemical reaction is different. Some chemical reaction occur
very fast whereas some chemical reaction occur very slowly.

13. Various factors increase or decrease the rate of a chemical reaction. Some of the
major factors that affect the rate of chemical reaction are given below:

a. Temperature b. Light c. Surface area d. Pressure e. Catalyst

14. The rate of chemical reaction is defined as the positive change in concentration of a

reactant or a product per unit time.

15. The chemical substance which is used to increase or decrease the rate of a chemical

reaction is called a catalyst.

16. The chemical reaction that absorbs heat during the chemical change is called

endothermic reaction.

17. The chemical reaction that evolves heat during the chemical change is called

exothermic reaction.

Sequential General Exercise 1

1. Choose the best answer from the given alternatives.

a. Magnesium burns in air and forms magnesium oxide. What type of chemical
reaction is this?

Combination reaction Decomposition reaction

Displacement reaction Acid base reaction

b. Which of the given chemical reactions is a displacement reaction.
Electrolysis of water
Reaction between Zinc and Hydrochloric acid
Reaction between Sodium and Chlorine
Reaction between Magnesium and Sulphuric acid

GREEN Science (Chemistry) Book-10 163

c. Which of the following reaction is a neutralization reaction?
Reaction between Sulphuric acid and Sodium hydroxide
Reaction between Magnesium and Oxygen
Reaction between Iron and Copper sulphate
Reaction between Potassium and Chlorine

d. Which of the given factors is essential for the chemical reaction during
photosynthesis?

Heat Sunlight Light Electricity

e. Which of the following is a positive catalyst?

Glycerol Sodium chloride

Manganese dioxide Calcium carbonate

2. Answer the following questions.
a. Define chemical reaction with any three examples.
b. What are reactants and products?
c. What is a word equation? Give any three examples.
d. What is a chemical equation? Write any three examples.
e. Name the four types of chemical reaction.
f. Define combination reaction with any three examples.
g. Define decomposition reaction with any three examples.
h. What is meant by displacement reaction? Name its types.
i. Write any three examples of decomposition reaction.
j. What is single displacement reaction? Give any two examples.
k. What is double displacement reaction? Write any two examples.
l. What is acid-base reaction? Write any two examples.
m. What is meant by rate of chemical reaction? Name any three factors that affect
the rate of chemical reaction.
n. How is rate of chemical reaction affected by increase or decrease in temperature?
Describe in brief.
o. What is the effect of light on the rate of chemical reaction?
p. Write down the effect of increase or decrease in surface area of reactants on the
rate of chemical reaction.
q. What is the effect of pressure on the rate of chemical reaction?

164 GREEN Science (Chemistry) Book-10

r. What is a catalyst? Write its types.
s. Define positive and negative catalyst with any one example of each.
t. Define exothermic and endothermic reaction with any two examples of each.

3. Differentiate between:
a. Reactants and Products
b. Word equation and Chemical equation
c. Combination reaction and Decomposition reaction
d. Positive catalyst and Negative catalyst
e. Endothermic reaction and Exothermic reaction

4. Give reason:

a. The chemical reaction between Hydrogen and Oxygen is called combination
reaction.

b. The chemical reaction between Zinc and dilute Hydrochloric acid is called
displacement reaction.

c. The chemical reaction between sulphuric acid and sodium hydroxide is called
neutralization reaction.

d. The rate of chemical reaction increases on increasing the temperature of reactants.

e. Manganese dioxide is called a positive catalyst.

f. Positive catalyst increases the rate of chemical reaction.

g. Increase in pressure increases the rate of chemical reaction on gaseous reactants.

5. Convert following unbalanced chemical equations into balanced chemical equations.

a. Mg + N2 Mg3N2

b. HCl + K2O KCl + H2O

c. K + O2 K2O

d. Fe + CuSO4 FeSO4 + Cu

e. HNO3 + Ca(OH)2 Ca(NO3)2 + H2O

f. Au + Cl2 AuCl3

g. H2SO4 + NaOH Na2SO4 + H2O

h. CaCO3 CaO + CO2

i. AgNO3 Ag + NO2 + O2

j. Na2SO4 + Pb (NO3)2 PbSO4 + NaNO3

6. Write down the given word equations in the form of balanced chemical equations.

a. Nitrogen + Hydrogen Ammonia
b. Hydrogen + Oxygen Water

c. Aluminum + Nitrogen Aluminium nitride

GREEN Science (Chemistry) Book-10 165

d. Magnesium + Chlorine Magnesium chloride

e. Calcium oxide + Water Calcium hydroxide

f. Calcium carbonate Calcium oxide + Carbon dioxide

g. Zinc + Hydrochloric acid Zinc chloride + Hydrogen

h. Copper + Oxygen Copper oxide

i. Potassium chlorate Potassium chloride + Oxygen

j. Nitric acid + Calcium hydroxide Calcium nitrate + Water

7. Describe an activity to demonstrate that the rate of a chemical reaction increases
on increasing the temperature of reactants.

8. Explain an activity to demonstrate that the rate of a chemical reaction increases on
increasing the area of contact of reactants.

Grid-based Exercise 2

Group ‘A’ (Knowledge Type Questions) (1 Mark Each)

1. What is a chemical reaction? Give one example.

2. Define formula equation with one example.

3. State any four factors that bring out a chemical change.

4. What is a precipitate ? Give one example.

5. Define negative catalyst with one example.

6. Define exothermic reaction with one example.

7. What is combination reaction? Give one example.

8. Define decomposition reaction with one example.

9. What is acid-base reaction? Give one example.

10. What is meant by the rate of chemical reaction?

11. Define endothermic reaction with one example.

12. What is displacement reaction? Write with examples.

13. Write a balanced chemical equation of decomposition reaction which is carried out
by the catalyst.

14. Write two information which can be obtained from the balanced chemical equation.

15. Write an example of single displacement chemical reaction.

Group ‘B’ (Understanding Type Questions) (2 Marks Each)

16. Write any two differences between reactants and products.

17. Acid-base reaction is also called neutralization reaction, why?

18. The rate of a chemical reaction increases on increasing temperature, why?

19. Write any two differences between synthesis reaction and dissociation reaction.

20. The rate of a chemical reaction increases on powdering the reactants, why?

21. Write any two limitations of balanced chemical equation.

166 GREEN Science (Chemistry) Book-10

22. Which type of chemical equation is given below? Define it.

Fe + CuSO4 → FeSO4 + Cu
23. How does heat enhance the rate of chemical reaction? Write in short.

24. The rate of a chemical reaction increases on increasing the concentration of reactants,
why?

25. Write any two differences between exothermic chemical reaction and endothermic
chemical reaction.

Group ‘C’ (Application Type Questions) (3 Marks Each)

26. In what condition do sodium chloride and silver nitrate react? Write the balanced
chemical equation of that reaction.

27. Change the given word equation into balanced chemical equation. What type of
chemical reaction is it ? What is the role of MnO2 in this reaction ?
Potassium Chlorate MnO2 Potassium Chloride + Oxygen

28. Write the balanced chemical equation for the following word equations:

i. Iron + Oxygen → Ferric Oxide

ii. Potassium Chlorate Heat Potassium Chloride + Oxygen

29. Give one example of exothermic chemical reaction. Write any two applications of
catalyst.

30. Describe an experiment to demonstrate that the rate of chemical reaction increases on
increasing the surface area of reactants.

Group ‘D’ (Higher Abilities Type Questions) (4 Marks Each)

31. Change the given word equations into formula equation. Also, write down the type
of the chemical equation.

i. Calcium bicarbonate → Calcium carbonate+Water+Carbon dioxide

ii. Aluminium + Hydrochloric acid → Aluminium chloride + Hydrogen

32. How does the concentration of sodium thiosulphate affect the rate of chemical reaction
in between the hydrochloric acid and sodium thiosulphate? Write the chemical
equation of the reaction between more active metal and more active non-metal. What
type of chemical reaction is it ?

33. Write a balanced chemical equation of decomposition reaction which is carried out by
the catalyst. Which type of chemical equation is given below? Define it.

AgNO3 + CaCI2 → AgCI + Ca(NO3)2
34. Describe an experiment to demonstrate that the rate of chemical reaction increases on

increasing the temperature of reactants.

35. Chemical reaction takes place when iron dust is added into Copper sulphate solution
but no reaction takes place when Copper dust is added into Ferrous sulphate solution.
Why ? Describe in brief the effect of physical state of reactants in the rate of chemical
reaction.

GREEN Science (Chemistry) Book-10 167

UNIT Acid, Base and Salt

9

Weighting Distribution Theory : 7 Practical: 2

Before You Begin

Matter can be defined as anything having mass and volume. All
matter have mass and they occupy space. For example, air, soil,
water, milk, stone, brick, wood, smoke, cloud, petrol, kerosene,
iron, gold, plastic, etc. Sound, light, shadow, heat, etc. do not have
mass and volume. So they are not matter. Matter can be soluble
or insoluble, transparent or opaque and good conductor or bad
conductor of heat and electricity. Matter exist in three different
states, viz. solid, liquid and gas. Same matter can exist in three
different states. For example, water can exist in all three states,
viz. solid (ice), liquid (water) and gas (vapour).

Learning Objectives Syllabus
After completing the study of this unit, students will be able to:
i. introduce acid, base and salt with examples. • Acid-Introduction and types

ii. explain the properties and uses of acid, base and salt. • Physical and chemical
properties of acids
iii. state the use of acid, base and salt in our daily life.
• Uses of acids
iv. write neutralization reaction between acid and base
and balance neutralization reactions. • Base and alkali-Introduction

• Physical and chemical
properties of bases/alkalis

• Uses of bases

• Salt: Introduction

• Properties and uses of salts

• Indicators, pH and pH scale

• Neutralization reaction

Glossary: A dictionary of scientific/technical terms

acid : the substances that produces H+ ions when dissolved in water
base : metal oxide or metal hydroxide
alkali : the base that dissolves in water
salt : the substance formed by the reaction between an acid and a base

168 GREEN Science (Chemistry) Book-10

Acids Do You Know

When we drink lemon juice, we feel sour

due to the presence of acid (i.e. citric acid) Most acids are sour in taste. But it is

in lemon juice. Similarly, pickles containing dangerous to touch and taste acids in

vinegar are also sour in taste. It shows that laboratory as acids burn our skin and

acids possess sour taste. The word 'acid' tongue.

refers to a sour substance. However, it does

not mean that all acids are sour. Most acids are sour. The word 'acid' has been derived

from the Latin word 'acidus' which means sour in taste. Citric acid, lactic acid, carbonic

acid, hydrochloric acid, sulphuric acid and nitric acid are the examples of acids.

The foods having acids are sour in taste. Fruits like lemon, orange, etc. taste sour due to
presence of citric acid. Grape fruit tastes sour due to the presence of tartaric acid. Similarly,
sour milk contains lactic acid, vinegar contains acetic acid and vitamin C contains ascorbic
acid.

The chemical substances which give hydrogen (H+) ions when dissolved in water are
called acids. For example: Hydrochloric acid (HCl), Sulphuric acid (H2SO4), Nitric acid
(HNO3), Carbonic acid (H2CO3), etc.

HCl +H2O H+ + Cl–
H2SO4 +H2O 2H+ + SO4– –
HNO3 +H2O H+ + NO3–
H2CO3 +H2O 2H+ + CO3– –

Fig.

9.1 Sulphuric acid (H2SO4) Nitric acid (HNO3)
Hydrochloric acid (HCl)

There are two types of acids on the basis of source or chemical nature. They are
(i) Inorganic acids and (ii) Organic acids.

i. Inorganic acids

The acids which are obtained from minerals are called inorganic acids. Acids like
Hydrochloric acid (HCl), Sulphuric acid (H2SO4), Nitric acid (HNO3), Carbonic acid
(H2CO3), etc are inorganic acids. Inorganic acids are also obtained from minerals. So
they are called mineral acids. They are commonly used in laboratories and industries.
Inorganic acids are strong in nature.

GREEN Science (Chemistry) Book-10 169

ii. Organic acids

The acids which are obtained from living organisms (plants and animals) are called
organic acids. Citric acid, Acetic acid, Tartaric acid, Formic acid, Lactic acid, etc. are
examples of organic acids. These acids are weak in nature. Citric acid is found in lemon.
Ascorbic acid and tartaric acid are found in fruits and vegetables. Similarly, formic acid
is found in ant bite.

There are two types of acids on the basis of strength. They are:

(i) Strong acids and (ii) Weak acids.

i. Strong acids

The acids which produce a high concentration of hydrogen (H+) ions when dissolved in
water are called strong acids. They are more corrosive in nature. Examples: Hydrochloric
acid (HCl), Sulphuric acid (H2SO4), Nitric acid (HNO3), etc. These acids produce high
concentration of hydrogen ions (H+) when dissolved in water. They have low pH value
and are good conductors of electricity.

ii. Weak acids

The acids which produce a low concentration of hydrogen (H+) ions when dissolved in
water are called weak acids. They are less corrosive in nature. Examples: Carbonic acid
(H2CO3), Acetic acid (CH3COOH), Formic acid (HCOOH), etc. These acids produce a low
concentration of H+ ions when dissolved in water. They have high pH value. They are
poor conductors of electricity.

Physical properties of acids Do You Know
1. Most acids are sour in taste.
Most acids are sour in taste but acids like
2. They change blue litmus paper into red Boric acid, Stearic acid and Salicylic acid
and methyl orange into red. are not sour.

3. They are corrosive in nature.

4. They do not change the colour of phenolphthalein.

5. They burn our skin.

Chemical properties of Acids

1. Acids reacts with active metals like Zn, Mg, Na, etc. and form hydrogen gas.

Dilute acid + Metal Salt + Hydrogen gas

2HCl + Zn ZnCl2 + H2

H2SO4 + Zn ZnSO4 + H2

2HCl + Mg MgCl2 + H2

H2SO4 + Mg MgSO4 + H2

2HNO3 + Zn Zn(NO3)2 + H2

170 GREEN Science (Chemistry) Book-10

2HCl + Na 2NaCl + H2

6HNO3 + 2Al 2Al(NO3)3 + 3H2

2. Acids react with bases and form salt and water.

Acid + Base Salt + Water

H2SO4 + KOH K2SO4 + H2O

HCl + NaOH NaCl + H2O

2HCl + CaO CaCl2 + H2O

H2SO4 + 2NaOH Na2SO4 + 2H2O

H2SO4 +CaO CaSO4 + H2O

HNO3 + KOH KNO3 + H2O

2HNO3 + Ca(OH)2 Ca(NO3)2 + 2H2O

HCl + KOH KCl + H2O

3. Acids dissolve in water and produce hydrogen ions.

HCl +H2O H+ + Cl–

H2SO4 +H2O 2H+ + SO4– –
H+ + NO3–
HNO3 +H2O H+ + CH3COO –

CH3COOH +H2O

4. Acids react with carbonates and form salt, water and carbon dioxide gas.

Acid + Carbonates Salt + Water + Carbon dioxide

2HCl + Na2CO3 2NaCl + H2O + CO2
2HCl + MgCO3 MgCl2 + H2O + CO2
H2SO4 + MgCO3 MgSO4 + H2O + CO2

5. Acids react with bicarbonates and form salt, water and carbon dioxide gas.

Acid + Bicarbonates Salt + Water + Carbon dioxide

HCl + KHCO3 KCl + H2O + CO2

2HCl + CaCO3 CaCl2 + H2O + CO2

H2SO4 + 2NaHCO3 Na2SO4 + 2H2O + 2CO2

H2SO4 +Mg(HCO3)2 MgSO4 + 2H2O + 2CO2

H2SO4 + Ca(HCO3)2 CaSO4 + 2H2O + 2CO2

2HCl + Ca(HCO3)2 CaCl2 + 2H2O + 2CO2

2HNO3 + Na2CO3 2NaNO3 + H2O + CO2

GREEN Science (Chemistry) Book-10 171

Uses of Acids

1. Hydrochloric acid, sulphuric acid and nitric acid are used in science laboratories and
industries.

2. Sulphuric acid is used for making chemical fertilizers, drugs and detergents.
3. Hydrochloric acid is used in tanning and printing industries.
4. Nitric acid is used for making plastics, dyes and explosives.
5. Carbonic acid is used in soft drinks like coca-cola, soda water, beer, etc.
6. Acetic acid (vinegar) is used in pickles.
7. Carbolic acid (phenol) is used to kill germs.
8. Citric acid is used as a source of vitamin C.

Some acids of daily use and their sources are given below.

S.N. Acids Sources
1. Lemon, tomato
2. Citric acid Grape fruit
3. Tartaric acid Red ant
4. Formic acid Milk, curd
5. Lactic acid Sour fruits
6. Ascorbic acid Chari amilo
Oxalic acid

Activity 1

Take a test tube and keep 10 ml of dilute hydrochloric acid. Now, keep, a piece of
magnesium.

Take another test tube and keep 10 ml of acetic acid in it. Keep a piece of magnesium.
What do you observe? Write down the conclusion of this activity.

Bases Do You Know

Bases are metallic oxides and metallic The bases that dissolve in water are called
hydroxides which react with acids and alkalis. But bases like CuO, HgO, BaO,
produce salt and water. For example, PbO, etc. do not dissolve in water. So,
Sodium oxide (Na2O), Calcium oxide all alkalis are bases but all bases are not
(CaO), Magnesium oxide (MgO), Sodium alkalis.
hydroxide (NaOH), Calcium hydroxide
[Ca(OH)2], etc. The bases that dissolve in
water and produce hydroxyl (OH–) ions are
called alkalis.

Na2O + H2O 2NaOH (Alkali)

NaOH +H2O Na+ +OH–

172 GREEN Science (Chemistry) Book-10

K2O + H2O 2KOH

KOH +H2O K+ + OH–

Sodium hydroxide (NaOH), Magnesium hydroxide [Mg(OH)2], Potassium hydroxide
(KOH), Calcium hydroxide [Ca(OH)2], etc. are examples of alkalis. Bases are bitter in taste.

Fig.

9.2 Magnesium hydroxide Calcium hydroxide
Sodium hydroxide

Differences between Acids and Bases

Acids Bases

1. Acids produce hydrogen (H+) ions 1. Bases produce hydroxyl (OH–) ions

when dissolved in water. when dissolved in water.

2. They turn blue litmus paper into red. 2. They turn red litmus paper into blue.

Types of bases

On the basis of strength, there are two types of bases, viz. (i) Strong bases and (ii) Weak
bases.

a. Strong bases

The bases that produce a high concentration of hydroxyl (OH–) ions in water are called
strong bases. Examples: Sodium hydroxide (NaOH), Potassium hydroxide (KOH),
Magnesium hydroxide [Mg(OH)2], etc. Strong bases have a high pH value. Their rate of
decomposition is more than that of weak bases.

b. Weak bases Do You Know

The bases that produce a low concentration Strong bases/alkalis burn our skin, So,
of hydroxyl (OH–) ions in water are called we should not touch bases/alkalis in a
weak bases. Examples: Ammonium laboratory.
hydroxide (NH4OH), Copper hydroxide
([Cu)OH)2], Ferric hydroxide [Fe(OH)3], etc. Ammonium hydroxide (NH4OH) is
Weak bases have a low pH value. Their rate called a weak base because it produces
of decomposition is less than that of strong a low concentration of hydroxyl ions
bases or alkalis. (OH–) when dissolved in water.

GREEN Science (Chemistry) Book-10 173

Physical Properties of Bases/Alkalis

1. Most bases are bitter in taste.

2. Their solutions have a soapy touch.

3. They turn red litmus into blue, methyl orange into yellow and phenolphthalein into
pink.

4. Strong alkalis dissolve oil and grease.

5. They burn our skin. So, we should not touch strong alkalis in science laboratories.

Chemical Properties of Bases/Alkalis

1. Bases react with acids and form salt and water.

Base/Alkali + Acid Salt + Water

2NaOH + H2SO4 Na2SO4 + H2O
NaOH + HCl
KOH + HCl NaCl + H2O
2NaOH + H2SO4
KOH + HNO3 KCl + H2O
MgO + 2HCl
Ca(OH)2 + 2HNO3 Na2SO4 + H2O
Ca(OH)2 + 2HCl
KNO3 + H2O

MgCl2 + H2O

Ca(NO3)2 + 2H2O

CaCl2 + 2H2O

2. Bases/Alkalis react with carbon dioxide and form carbonate and water.

Base/Alkali + Carbon dioxide Carbonate + Water

LiOH + CO2 Li2CO3 + H2O
2NaOH + CO2 Na2CO3 + H2O
Mg(OH)2 + CO2 MgCO3 + H2O
2KOH + CO2 K2CO3 + H2O
Ca(OH)2 + CO2 CaCO3 + H2O

3. Alkalis dissolve in water and produce hydroxyl (OH–) ions.
+H2O
NaOH Na+ + OH–

Mg(OH)2 +H2O Mg++ + 2OH–
KOH +H2O K+ + OH–

Ca(OH)2 +H2O Ca++ + 2OH–
NH4OH +H2O NH4+ + OH–

174 GREEN Science (Chemistry) Book-10

4. Alkalis react with Ammonium salts and produce salt, water and ammonia gas.

Alkali + Ammonium salt Salt + Water + Ammonia gas

Ca(OH)2 + 2NH4Cl CaCl2 + 2H2O + 2NH3
NaOH + NH4Cl NaCl + 2H2O + NH3
Ca(OH)2 + (NH4)2CO3 CaCO3 + 2H2O + 2NH3
Mg(OH)2 + (NH4)2CO3 MgCO3 + 2H2O + 2NH3

5. Alkalis separate insoluble metal hydroxides when they are kept in heavy metal salts.

Alkali + Heavy metal salts Salt + Insoluble hydroxide

2NaOH + CuSO4 Na2SO4 + Cu(OH)2
3NH4OH + FeCl3 3NH4Cl + Fe(OH)3
2KOH + ZnCl2 2KCl + Zn(OH)2

Uses of Bases/Alkalis

1. Sodium hydroxide (NaOH) or Caustic Soda is used for making soft soap, paper, etc.
It is also used to purify petroleum products.

2. Potassium hydroxide (KOH) or Caustic potash is used for making soft soap, chemical
fertilizers, etc.

3. Aluminium hydroxide [Al(OH)3] is used to reduce hyperacidity.
4. Calcium hydroxide [Ca(OH)2] or lime water is used in laboratory and to reduce

hardness of water. It is also used to reduce acidity of soil and to prepare bleaching
powder.

5. Ammonium hydroxide (NH4OH) is used for making nitric acid, chemical fertilizers,
dyes and plastics.

6. Calcium oxide (CaO) or quick lime is used for softening hard water, for making
cement and in purification of sugar.

7. Magnesium hydroxide [Mg(OH)2] is used as an antacid to reduce hyperacidity.

Activity 2

Take a magnesium ribbon and burn it. The magnesium burns in air with a very bright
flame and forms a white powder, i.e. Magnesium oxide.

– Collect the Magnesium oxide in a beaker and add a few drops water.

The magnesium oxide combines with water and forms an alkali, i.e., Magnesium
hydroxide [Mg (OH)2].

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

MgO(s) + H2O (l) Mg (OH)2 (aq)
(Alkali)

In this way, we can produce an alkali from magnesium ribbon.

GREEN Science (Chemistry) Book-10 175

Salt

A salt is a chemical substance formed by the chemical reaction between an acid and a base.
Examples: (i) Sodium chloride (NaCl), (ii) Potassium chloride (KCl), (iii) Sodium sulphate
(Na2SO4), (iv) Copper sulphate (CuSO4), (v) Ammonium chloride (NH4Cl), etc.

Fig.

9.3 Copper sulphate Ammonium chloride
Sodium chloride

Salt can also be defined as a substance formed by partial or complete replacement of hydrogen
atom by a metal or ammonium radical.

KOH + H2SO4 KHSO4 + H2O (Partial displacement)

2KOH + H2SO4 K2SO4 + H2O (Complete displacement)

Most salts are neutral but some are acidic and some are basic in nature. The process in which

an acid reacts with a base and forms two neutral substances, i.e. salt and water is called

neutralization reaction.

A salt consists of two types of radicals. They are acid radical or non-metallic radical and basic
radical or metallic radical. The radical obtained from an acid is called an acid radical or non-
metallic radical. Similarly, the radical obtained from a base is called basic or metallic radical.
In salt NaCl, Na+ is a basic radical as it comes from the base i.e. Na2Oand Cl– is an acid radical
as it comes from acid, i.e. HCl.

There are different types of salt on the basis of chemical nature and method of formation. The
different types of salts are given below:

1. Neutral salt 2. Acidic salt

3. Basic salt 4. Acid salt

5. Base salt 6. Hydrated salt

1. Neutral salt or Normal salt

Neutral salt is the salt which is formed by the chemical reaction between a strong acid
and a strong alkali or a weak acid and a weak base. Neutral salts are formed by complete
displacement of hydrogen from acids.

Examples of neutral salts or normal salts:

i. Sodium chloride (NaCl) ii. Magnesium chloride (MgCl2)
iii. Potassium chloride (KCl) iv. Sodium nitrate (NaNO3)

176 GREEN Science (Chemistry) Book-10

v. Potassium nitrate (KNO3), etc.
Base/Alkali + Acid Neutral salt + Water

NaOH + H2SO4 Na2SO4 + H2O
KOH + HCl KCl + H2O
KOH + H2O
+ HNO3 KNO3

2. Acidic salt

Acidic salt is the salt which is formed by chemical reaction between a strong acid and a weak
base (alkali).

Examples of acidic salts: ii. Copper chloride (CuCl2)
i. Copper sulphate (CuSO4) iv. Lead nitrate [Pb(NO3)2]
iii. Lead chloride (PbCl2) Acidic salt + Water
v. Lead sulphate (PbSO4), etc.
Strong acid + Weak acid

H2SO4 + CuO CuSO4 + H2O
H2SO4 + 2NH4OH (NH4)2SO4 + H2O
2HNO3 + CuO Cu(NO3)2 + H2O
HCl + NH4OH NH4Cl + H2O

3. Basic salt

Basic salt is the salt which is formed by the chemical reaction between a weak acid and a
strong alkali (base).

Examples of basic salts:

i. Sodium carbonate (Na2CO3) ii. Potassium carbonate (K2CO3)
iii. Magnesium carbonate (MgCO3) iv. Sodium acetate (CH3COONa), etc.
Weak acid + Strong alkali Basic salt + Water

H2CO3 + 2NaOH Na2CO3 + 2H2O
CH3COOH + KOH
CH3COOH + NaOH CH3COOK + H2O

CH3COONa + H2O

4. Acid salt

Acid salt is the salt which is formed by partial replacement of hydrogen (H+) ion by a metal.
Examples of acid salts:

i. Sodium bicarbonate (NaHCO3) ii. Potassium bicarbonate (KHCO3)

iii. Magnesium bicarbonate [Mg (HCO3)2] iv. Calcium bicarbonate [Ca (HCO3)2], etc.

NaOH + H2CO3 NaHCO3 + H2O

KOH + H2CO3 KHCO3 + H2O

Mg(OH)2 + 2H2CO3 Mg(HCO3)2 + H2O

NaOH + H2SO4 NaHSO4 + H2O

GREEN Science (Chemistry) Book-10 177

5. Base salt

Base salt is the salt formed by partial replacement of hydroxyl (OH–) radical of a base by an
acidic radical.

Examples of base salts:

i. Zinc hydroxychloride [Zn(OH) Cl]

ii. Lead hydroxychloride [Pb(OH)Cl], etc.

HCl + Zn(OH)2 Zn(OH)Cl + H2O
HCl + Pb(OH)2 Pb(OH)Cl + H2O

Activity 3

Take five test tubes and prepare a solution of NaCl, CuSO4, Na2CO3, NaHCO3 and
Zn(OH)Cl separately.

Identify normal salt, acid salt, base salt, acidic salt and basic salt using indicators
(litmus paper, phenolphthalein and methyl orange).

6. Hydrated salt

Hydrated salt is the salt which contains certain molecules of water associated with it. These
water molecules are called water of crystallization or water of hydration.

Examples of hydrated salts:

i. Sodium carbonate decahydrate (Na2CO3. 10 H2O)
ii. Calcium sulphate heptahydrate(CaSO4 . 7H2O)
iii. Copper II sulphate pentahydrate (CuSO4. 5H2O) or Blue vitriol
iv. Ferrous sulphate heptahydrate (FeSO4. 7H2O)
v. Zinc sulphate heptahydrate (Zn SO4. 7H2O) or White vitriol
vi. Magnesium sulphate heptahydrate (MgSO4.7H2O) or Epsom salt
vii. Sodium sulphate decahydrate (Na2SO4.10H2O)

Fig.

9.4 Sodium carbonate
Zinc sulphate

178 GREEN Science (Chemistry) Book-10

Properties of salts

1. Most salts are bitter in taste. Some are salty (e.g. NaCl) and others are tasteless.

2. Most salts are neutral but some may be acidic or basic.

3. Most salts dissolve in water. All salts of Na, K and NH4 are soluble in water. All nitrate
and bicarbonate salts also dissolve in water. All chloride salts are water soluble except
chloride salts of Ag and Pb. All sulphate salts are water soluble except sulphates of Pb
and Ba.

4. They may be white, colourless or colourful. The salts of Cu, Fe, Mn, Cr, etc. are
coloured.

5. They conduct electricity in solution or molten state.

6. They have high melting point and boiling point.

Uses of salts Do You Know
1. Common salt (NaCl) is used in our foods.
Salts of metals like Na, K, Mg, Ca, Al
It is also used as preservative and in and Ba are white or colourless.
the manufacture of sodium hydroxide,
hydrochloric acid and washing soda. Salts of metals like Cu, Co, Mn, Fe, Ni
and Cr are colourful.

2. Copper sulphate (CuSO4) is used to The salts that dissolve in water can be
make fungicides and for electroplating.
electrolysed.

3. Calcium sulphate (CaSO4.7H2O) is used
for making cement, chalk and plastering of fractured bones.

4. Ammonium sulphate [(NH4)2SO4] and Potassium nitrate (KNO3) are used for making
chemical fertilizers.

5. Sodium bicarbonate (NaHCO3) is used for making baking powder and to make fire
extinguisher.

6. Magnesium sulphate (MgSO4) is used for treating constipation.
7. Sodium carbonate (Na2CO3) is used for making glass, soap and detergent.
8. Ammonium chloride (NH4Cl) is used in dry cells as electrolyte.
9. Zinc sulphate (ZnSO4) is used to make white pigment.
10. Anhydrous Ferrous sulphate (FeSO4) is used to make medicines for anaemia.
11. Copper sulphate (CuSO4) is used for electroplating and preserving woods.
12. Sodium carbonate (Na2CO3) is used to make soap, detergent and glass.
13. Silver nitrate is used as a laboratory reagent.

Indicators

We cannot identify whether a given chemical substance is an acid, base or salt just
by observing it. We use some chemicals to identify them. These chemicals are called
indicators. These chemical substances are used to indicate whether a substance is acidic,
basic or neutral in nature. For example, litmus paper (red and blue), methyl orange,
phenolphthalein, etc. These indicators are called ordinary indicators.

GREEN Science (Chemistry) Book-10 179

Fig.9.5 Blue litmus paper Methyl orange
Fig. Red litmus paper

Indicators are obtained from different parts of plants like roots, flowers, leaves, etc. These
parts are collected, crushed and mixed with organic solvents to get indicators.

Following table shows the various indicators and their effects in acid, base and salt.

S.No. Indicators Colour in acid Colour in basic Colour in neutral
solution solution salt solution
1. Litmus paper (red)
No change in Changes into No change in
2. Litmus paper (blue) colour blue colour

3. Methyl orange Changes into No change in No change in
red colour colour
4. Phenolphthalein
Changes into Changes into No change in
red yellow colour

No change in Changes into No change in
colour pink colour

Universal Indicator

Ordinary indicators can indicate whether a substance is acidic, basic or neutral in nature
but cannot measure the strength. Therefore, a special kind of indicator is used to measure
the strength of the given substance which is called the universal indicator. So, a universal
indicator is a special kind of indicator which is used to measure the strength of acidity or
alkalinity of a solution. It changes colour when kept in an acidic, basic or neutral solution
which is strength of the solution. A universal indicator is prepared by mixing several
ordinary indicators of different colour.

0 1 2 3 4 5 6 7 8 9 10 11 12 13 14

9.6 acidic neutral alkaline

pH colour chart

180 GREEN Science (Chemistry) Book-10

pH and pH Scale

A pH is the measure of hydrogen ion concentration present in a solution. It is measured
by using pH paper and pH meter.

A pH scale is the standard scale which is used to Fig.
measure the strength of acidic or alkaline solution. It
consists of numbers 1 to 14 with their corresponding 9.7
colours in the scale. The pH value 1 to 6 represents
acidity, pH value 7 represents neutrality and pH value
8 to 14 represents the basicity or alkalinity. The solution
having pH value 1 is the strongest acid and that having
14 is the strongest alkali.

pH meter

Acidity increases Neutral Alkalinity increases
7
pH 1 23 456 8 9 10 11 12 13 14
Red Green
Rose Yellow Light green Greenish blue Blue Deep blue

The pH value of some common chemicals

Acidic Chemicals pH value
Neutral Hydrochloric acid (HCl) 1
Basic/Alkaline Sulphuric acid (H2SO4) 1.2
Lemon juice (citric acid) 2.5
Carbonic acid, vinegar 3
Common salt solution
Water 7
Sugar solution
Human blood 7.3
Baking soda 8.5
Washing soda 11.5
Sodium hydroxide (NaOH) 13

Activity 4

Take red and blue litmus paper and three test tubes. Mark the test tubes 1, 2 and 3.
Keep the solution of acid in test tube 1, solution of alkali in test tube 2 and solution

of common salt in test tube 3.
Now, take red litmus papers and immerse one litmus paper separately in each test

tube. Observe the change in colour.
Take blue litmus papers and repeat the above activity.
Prepare a chart after your observation.

GREEN Science (Chemistry) Book-10 181

Activity 5

Take solution of acids, bases and salts in different test tubes.
Measure the pH value of each by using a pH meter.

Neutralization Reaction

The chemical reaction that takes place between an acid and a base to from neutral
substances, i.e. salt and water is called neutralization reaction. During chemical reaction,
both acid and base lose their properties and form two neutral substance, i.e. salt and
water. So, acid-base reaction is called neutralization reaction

Examples:

Strong acid + Strong base Neutral salt + Water

HCl + NaOH NaCl + H2O
H2SO4 + 2KOH K2SO4 + 2H2O
HNO3 + NaOH NaNO3 + H2O
H2SO4 + 2NaOH Na2SO4 + 2H2O

Weak acid + Weak base Neutral salt + Water

H2CO3 + Cu(OH)2 CuCO3 + 2H2O
H2CO3 + Pb(OH)2 PbCO3 + 2H2O
H2CO3 + NH4OH (NH4)2CO3 + H2O

Application of Neutralization Reaction

1. Neutralization reaction is utilized to treat hyperacidity or gastritis. Magnesium
hydroxide [Mg (OH)2] is used as an antacid to neutralize the acidity caused by
Hydrochloric acid (HCl) in our stomach.

2. Calcium oxide or lime (CaO) is used by farmers to neutralize the acidity of soil. If
the soil is basic, it is treated with compost made of rotting vegetables or leaves. The
acidic gas formed from the decomposition of compost neutralizes the alkalis in the
basic soil.

3. Soap and baking powder are used to neutralize the effect of acidic poison of sting of
red ant and bees.

4. Acetic acid or vinegar is used to neutralize the acidity caused by the sting of yellow
bumble bee and wasp.

5. Tooth decay occurs due to acid produced during decomposition of food particles in
the mouth. It is neutralized by brushing teeth with alkaline toothpaste.

182 GREEN Science (Chemistry) Book-10

Key Concepts

1. The word 'acid' has been derived from the Latin word 'acidus' which means sour in
taste.

2. The chemical substances which give hydrogen (H+) ions when dissolved in water
are called acids. For example: Hydrochloric acid (HCl), Sulphuric acid (H2SO4),
Nitric acid (HNO3), Carbonic acid (H2CO3), etc.

3. The acids which are obtained from minerals are called inorganic acids. Acids like
Hydrochloric acid (HCl), Sulphuric acid (H2SO4), Nitric acid (HNO3), Carbonic acid
(H2CO3), etc are inorganic acids.

4. The acids which are obtained from living organisms are called organic acids. Citric
acid, Acetic acid, Tartaric acid, Formic acid, Lactic acid, etc. are examples of organic
acids.

5. The acids which produce a high concentration of hydrogen (H+) ions when dissolved
in water are called strong acids.

6. The acids which produce a low concentration of hydrogen (H+) ions when dissolved
in water are called weak acids.

7. Bases are metallic oxides and metallic hydroxides which react with acids and
produce salt and water. For example, Sodium oxide (Na2O), Calcium oxide (CaO),
Magnesium oxide (MgO), Sodium hydroxide (NaOH).

8. The bases that produce a high concentration of hydroxyl (OH–) ions in water are
called strong bases. Examples: Sodium hydroxide (NaOH), Potassium hydroxide
(KOH), Magnesium hydroxide [Mg(OH)2], Calcium hydroxide [Ca(OH)2], etc.

9. The bases that produce a low concentration of hydroxyl (OH–) ions in water are
called weak bases. Examples: Copper hydroxide ([Cu)OH)2], Ferric hydroxide
[Fe(OH)3], etc.

10. Salt can also be defined as a substance formed by partial or complete replacement of
hydrogen atom by a metal or ammonium radical.

11. Neutral salt is the salt which is formed by the chemical reaction between a strong acid
and a strong alkali or a weak acid and a weak base.

12. Acidic salt is the salt which is formed by chemical reaction between a strong acid
and a weak base (alkali).

13. Basic salt is the salt which is formed by the chemical reaction between a weak acid
and a strong alkali (base).

14. Acid salt is the salt which is formed by partial replacement of hydrogen (H+) ion by
a metal.

15. Base salt is the salt formed by partial replacement of hydroxyl (OH–) radical of a
base by an acidic radical.

16. Hydrated salt is the salt which contains certain molecules of water.

17. A special kind of indicator is used to measure the strength of the given substance
which is called the universal indicator.

18. The chemical reaction that takes place between an acid and a base to from neutral
substances, i.e. salt and water is called neutralization reaction.

GREEN Science (Chemistry) Book-10 183

Sequential General Exercise 1

1. Choose the best answer from the given alternatives.
a. Which of the given substance is a weak acid?

HCl H2SO4 HNO3 H2CO3

b. ....................... is used as a source of vitamin C.

Carbonic acid Citric acid

Sulphuric acid Acetic acid

c. Bases react with carbon dioxide and form ...................... and water.

nitrates carbonates sulphates oxides

d. Which of the following is a neutral salt?

NaHSO4 CuSO4 NaCl NaHCO3

e. Which of the given alkalis is used as an antacid?

Mg(OH)2 KOH NaOH Ca(OH)2

2. Answer the following questions.

a. Define acids with any four examples.

b. What are strong acids? Give any two examples of weak acids.

c. Define organic acids. How do they differ from inorganic acids?

d. Write any three physical properties of acids.

e. Write any two chemical properties of acids.

f. Write down the uses of given acids.

i. Acetic acid ii. Nitric acid

iii. Carbonic acid iv. Sulphuric acid

v. Tartaric acid vi. Formic acid

g. Define bases with any five examples.

h. What are alkalis? Give any three examples of the bases that dissolve in water.

i. Write any three physical properties and two chemical properties of bases (alkalis).

j. Write any four uses of alkalis.

k. Define salt and write any five examples.

l. Write any four properties and four uses of salt.

m. Define indicator and universal indicator.

184 GREEN Science (Chemistry) Book-10

n. What is pH? Write down the pH value of the strongest acid, neutral salt and the
strongest alkali.

o. What is a pH-scale?
3. Give reason.

a. We should not touch and taste acids in a science laboratory.
b. Sodium chloride is called a neutral salt.
c. All bases are not alkalis but all alkalis are bases.
d. Methyl orange is called an indicator.
e. Universal indicator is better than an ordinary indicator.
f. We eat aluminium hydroxide to reduce hyperacidity.
g. Acids are sour in taste.
4. Differentiate between:
a. Inorganic acids and Organic acids
b. Base and Alkali
c. Acids and Alkalis
d. Ordinary indicator and Universal indicator
5. All alkalis are bases but all bases are not alkalis. Justify this statement.
6. Write down the effects of litmus paper, methyl orange and phenolphthalein on
acid, base and salt.
7. What are hydrated salts? Give any three examples.
8. Define neutralization reaction. Explain with examples.
9. Neutralization reactions are highly applicable in our daily life. justify this
statement giving any three examples.

GREEN Science (Chemistry) Book-10 185

Grid-based Exercise 2

Group ‘A’ (Knowledge Type Questions) (1 Mark Each)

1. Define acid with any two examples.

2. Name the acid found in each of the given substances:

i. Juice of lemon ii. Vinegar

3. Define organic acids with any two examples.

4. Mention any two physical properties of acids.

5. Define weak acid with one example.

6. Define base and give any two examples.

7. What is an alkali? Give any two examples.

8. Name any two alkalis that react with skin.

(Ans: NaOH, KOH)

9. What is a salt? Give any two examples.

10. Define acidic salt with one example.

11. Define hydrated salt with one example.

12. Write any two physical properties of bases.

13. Define strong alkali with one example.

14. Name the bases which are used for given activities:

i. To soften hard water ii. To make soft soap [Ans: (a) CaO (b) KOH]

15. Write down a name of alkali which is used to balance the pH of human stomach.

16. Name the bases which are used for given activities:

i. To purify sugar

ii. To purify petroleum products [Ans: (a) CaO (b) NaOH]

Group ‘B’ (Understanding Type Questions) (2 Marks Each)

17. Write any two differences between acid and base.

18. All alkalis are bases but all bases are not alkalis, why?

19. Sodium hydroxide is called a base but sodium chloride is called a salt. Why?

20. It is dangerous to touch or taste acids, why?

21. Write any two differences between alkali and base.

22. Write any two differences between acidic radical and basic radical.

23. Explain why water can be considered as an acid as well as a base?

186 GREEN Science (Chemistry) Book-10

24. Write any two differences between neutral salt and basic salt.
25. Acetic acid is called weak acid and Sulphuric acid is called strong acid, why?
26. Sodium hydroxide is called a strong alkali, why?
27. Write any two differences between strong acid and weak acid.
28. What is meant by neutralization reaction? Give one example.

Group ‘C’ (Application Type Questions) (3 Marks Each)

29. Name any three acids which are used in our daily life. Also, write an application of
each.

30. Write down any three examples of neutralization reaction applied in our daily life.

31. Write any three uses of acids in our daily life.

32. Give an application of each of the given compounds.

i. Calcium sulphate ii. Sodium hydroxide iii. Sulphuric acid

33. Write any three uses of salts in our daily life.

34. How are neutral salt and acid salt formed? Write with one example of each.

Group ‘D’ (Higher Abilities Type Questions) (4 Marks Each)

35. Mention any two chemical properties of alkalis with chemical equations. ‘

36. How is salt prepared? Mention any two methods.

37. A compound gives hydrogen ion and chlorine ion in the solution state:

i. Write down the name and molecular formula of the compound.

ii. In what colour is methyl orange changed when it is treated with above compound ?

iii. Write down the name of salt formed by the chemical reaction of above compound
with zinc. Write chemical equation.

38. Write the balanced chemical equation of the reaction between strong base and weak
acid, and also mention the type of salt obtained in the reaction.

39. Name the compound which gives hydrogen and sulphate ion in solution. Write
down the balanced chemical equation of the chemical reaction occurred when above
compound is treated with Sodium hydroxide.

GREEN Science (Chemistry) Book-10 187

UNIT Some Gases

10

Weighting Distribution Theory : 5 Practical: 2

Before You Begin

Air consists of different types of gases like nitrogen, oxygen, carbon
dioxide, argon, neon, etc. Nitrogen occupies about 78.1% of the air
by volume, oxygen gas occupies about 20.9% of the air by volume.
Similarly, carbon dioxide gas occupies about 0.03% of air by volume.
Green plants use carbon dioxide during photosynthesis and all
animals and plants release carbon dioxide gas while breathing. When
dead bodies of plants and animals decay and decompose, two gases,
viz. carbon dioxide and ammonia are released in air. In this unit,
we will study about laboratory preparation, properties and uses of
carbon dioxide and ammonia gas.

Learning Objectives Syllabus

After completing the study of this unit, students will be able to: • Introduction to carbon
dioxide and ammonia gases
i. prepare carbon dioxide gas in the laboratory and
explain its properties and uses. • Occurrence of carbon dioxide
and ammonia
ii. prepare ammonia gas in the laboratory and explain its
properties and uses. • Laboratory preparation of
carbon dioxide and ammonia

• Manufacture of carbon
dioxide and ammonia

• Properties of carbon dioxide
and ammonia

• Uses of carbon dioxide and
ammonia

Glossary: A dictionary of scientific/technical terms

photosynthesis : the process of making food by green plants in the presence of sunlight
carbogen
urea : the mixture of 10 – 15 % oxygen and carbon dioxide gas

dry ice : a chemical fertilizer produced by heating carbon dioxide and ammonia
under high pressure

: the white solid form of carbon dioxide obtained after cooling carbon
dioxide to – 78°C

188 GREEN Science (Chemistry) Book-10

A. Carbon dioxide

Carbon dioxide is a compound gas having molecular formula CO2. It means that one
molecule of carbon dioxide (CO2) is made of one atom of carbon and two atoms of oxygen.
The molecular weight of carbon dioxide is 44 amu. This gas is very essential for living
beings as green plants use CO2 gas to prepare food during photosynthesis.

Discovery of carbon dioxide

Carbon dioxide gas was discovered by Van Helmont in 1630 AD by burning wood. In
1755 AD, Joseph Black prepared this gas by burning magnesium carbonate (MgCO3).
Similarly, in 1783 AD, Lavoisier proved that carbon dioxide is the compound made of
carbon and oxygen.

Occurrence of Carbon dioxide

In nature, carbon dioxide gas is found in free as well as in combined state. In atmosphere,
carbon dioxide occupies 0.03% by volume. All animals and plants release carbon dioxide
in air while breathing. Some amount of carbon dioxide is found dissolved in water. So
some carbon dioxide is found in water of sea, river, lake, pond, etc.

When carbon compounds like wood, coal, petrol, diesel, oil, fat, wax, etc. burn, they release
carbon dioxide in air. In compound state, carbon dioxide is found in mineral carbonates
such as calcium carbonate (CaCO3), magnesite (MgCO3), etc.

Laboratory Preparation of Carbon dioxide

Principle

In laboratory, carbon dioxide gas is prepared by the chemical reaction between pieces of

marble or limestone (CaCO3) with dilute hydrochloric acid (HCl).

CaCO3 + 2HCl CaCl2 + H2O + CO2

(dil.)

Materials required Do You Know
i. Apparatus
Woulfe's bottle The pure form of calcium carbonate is the
Thistle funnel limestone found in nature.

Corks

Delivery tube

Gas jar

ii. Chemicals

Pieces of limestone or marble

Dilute hydrochloric acid

GREEN Science (Chemistry) Book-10 189

Procedure

• Some pieces of limestone are kept in a Woulfe's bottle.

• The apparatus is set as shown in the figure and dilute hydrochloric acid is poured in
the Woulfe's bottle till the acid covers the pieces of limestone.

• Chemical reaction takes place between dilute hydrochloric acid and limestone. Brisk
effervescence can be seen during the chemical reaction. As a result, carbon dioxide
gas is produced.

• Carbon dioxide thus produced is collected in the gas jar by upward displacement of
air as carbon dioxide gas is heavier than air.

Dil. hydrochloric acid Delivery tube
Thistle funnel

Gas jar

Fig. Pieces of CaCO3 CO2 gas

10.1

Laboratory Preparation of Carbon dioxide gas

Precautions

1. The apparatus should be made air tight. Do You Know

2. Carbon dioxide gas should be collected Carbon dioxide is heavier than air and
in the gas jar by upward displacement soluble in water. So, it can be collected in
of air. the gas jar by upward displacement of air.

3. The lower end of the thistle funnel
should be immersed in the acid.

4. The lower end of delivery tube inside the round bottom flask should not touch the
acid.

5. Carbon dioxide gas dissolves in water. So it should not be collected in the gas jar by
passing it through water.

Test of carbon dioxide

1. When a burning piece of wood is inserted in the gas jar, it extinguishes. It shows that
the gas is carbon dioxide because carbon dioxide is non-supporter of combustion.

2. When a moist blue litmus paper is kept in the gas jar containing carbon dioxide, the
litmus paper turns red because carbon dioxide gas is acidic in nature.

190 GREEN Science (Chemistry) Book-10

3. When carbon dioxide is passed through clear solution of lime water, the lime water
turns milky due to formation of water insoluble calcium carbonate.

CO2 + Ca(OH)2 CaCO3 + H2O

Water insoluble

Some more methods of preparation of carbon dioxide gas

1. Carbon dioxide is prepared by burning carbon in sufficient oxygen.

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

2. When hydrocarbons like methane, ethane, propane, butane, etc. burn in oxygen,

carbon dioxide is formed.

CH4 (g) + 2O2 (g) CO2(g) + 2H2O (l)

2C2H6 (g) + 7O2 (g) 4CO2 (g) + 6H2O(l)

3. Carbon dioxide gas is prepared by heating calcium carbonate in a kiln.

CaCO3 (s) CaO (s) + CO2 (g)

4. Carbon dioxide gas can be prepared by the reaction of carbonates or bicarbonates

with acids.

Na2 CO3 (s) + 2HCl (aq) 2NaCl (aq) + H2O (l) + CO2 (g)
Ca (HCO3)2 (s) + 2HCl (aq) CaCl2 (s) + 2H2O (l) + CO2 (g)
Mg (HCO3)2 + 2HCl (aq)
MgCl2(s) + 2H2O (l) + CO2 (g)

Manufacture of carbon dioxide gas

In industrial scale, carbon dioxide gas is prepared by heating limestone or calcium
carbonate in a kiln. When limestone is heated in a kiln at high temperature, it decomposes
into calcium oxide and carbon dioxide gas.

CaCO3   CaO + CO2

Carbon dioxide gas thus produced is used for industrial purpose and calcium oxide

is used for white washing. When calcium oxide is mixed with water, it forms calcium

hydroxide.

CaO + H2O Ca(OH)2

Physical properties of carbon dioxide Do You Know

1. Carbon dioxide is colourless and The clear solution of calcium
odourless gas. hydroxide is called lime water.

2. It is slightly acidic in taste when Calcium oxide is called quicklime
dissolves in water. whereas calcium hydroxide is called
slaked lime.
3. It is about 1.5 times heavier than air.

4. It turns moist blue litmus paper into red.

GREEN Science (Chemistry) Book-10 191

5. It changes into white solid when cooled down to – 78° C, which is commonly known
as dry ice.

6. It does not support combustion and extinguishes burning substances.

7. It is a non-poisonous gas. However, animals die in the atmosphere of carbon dioxide
due to suffocation.

Chemical properties of carbon dioxide gas

1. Carbon dioxide reacts with alkali solution (e.g. potassium hydroxide) and forms
corresponding carbonate and water.

CO2 + 2KOH K2 CO3 + H2O

CO2 + 2 NaOH Na2 CO3 + H2O

2. When carbon dioxide gas dissolves in water, it forms carbonic acid. It is used in cold

drinks to make them sour.

CO2 + H2O H2CO3

3. Carbon dioxide neither burns itself nor supports combustion. But burning magnesium

ribbon keeps on burning when inserted in the gas jar containing carbon dioxide and

forms white solid powder (MgO) and carbon (C).

2Mg + CO2 2MgO + C

4. When carbon dioxide is passed in the clear solution of lime water for a while, the lime

water turns milky due to formation of water insoluble calcium carbonate.

CO2 + Ca (OH)2 CaCO3 + H2O

When carbon dioxide is passed in lime water for a long time, milky colour disappears

due to formation of water soluble calcium bicarbonate.

CO2 + CaCO3 + H2O Ca (HCO3)2

5. Green plants use carbon dioxide gas to prepare food in leaves in the presence of the
sunlight. This process is called photosynthesis.

Sunlight

6CO2 + 6H2O chlorophyll C6H12O6 + 6O2

6. Carbon dioxide reacts with red hot coke at about 900°C and forms carbon monoxide.

CO2 + C 900°C 2CO

Activity 1

Take a beaker and prepare lime water using quick lime (Calcium oxide). Take a straw
and blow air into the lime water. What do you observe? What is its reason? Discuss
among friends and write down the chemical reaction involved in this process.

192 GREEN Science (Chemistry) Book-10

Uses of carbon dioxide gas Do You Know

1. Carbon dioxide is used in cold drinks Carbon dioxide gas does not extinguishes
like coca-cola, beer, soda water, etc. fire itself but helps to reduce flame
by displacing oxygen supply as it is a
2. Green plants used carbon dioxide gas heavier gas.
for photosynthesis.

3. Liquid carbon dioxide is used in sugar
mills for carbonation process.

4. It is used for manufacturing urea Do You Know
(NH2CONH2) and sodium carbonate or
washing soda (Na2CO3.10H2O). Carbon dioxide along with calcium
hydroxide helps to remove gum, colour
5. It is used for making dry ice to preserve and amino acid impurities.
meat, fish, fruits and vegetables.

6. The mixture of 95% oxygen and 5% Do You Know
carbon dioxide (carbogen) is used to
stimulate breathing to treat pneumonic Carbogen is also called Meduna's mixture
patients. after its inventor Ladislas Meduna.

7. Carbon dioxide is used in fire
extinguishers.

In a fire extinguisher, sodium bicarbonate or sodium carbonate
and conc. sulphuric acid are kept separately. During fire, the
fire extinguisher is inverted to prepare a large amount of carbon
dioxide by mixing those chemicals. The CO2 gas thus produced is
used to extinguish fire.

2NaHCO3 + H2SO4 (conc.) Na2SO4 + 2H2O + 2CO2 Fig.
Na2CO3 + H2SO4(conc.) Na2SO4 + H2O + CO2

B. Ammonia 10.2
Fire extinguisher

Ammonia is a compound gas having molecular formula NH3. It means that one molecule
of ammonia is formed by combination of one atom of nitrogen and three atoms of

hydrogen. The molecular weight of ammonia is 17 amu. This gas has strong pungent

odour and highly soluble in water.

Discovery of ammonia

Ammonia gas was disovered by Lavoisier by heating the mixture of ammonium chloride
(NH4Cl) and Calcium hydroxide [Ca(OH)2]. The composition of ammonia gas was
discovered by Davy and Berthecol.

GREEN Science (Chemistry) Book-10 193

Occurrence of ammonia

Ammonia gas is found in free as well as in the form of mixture in nature. Some amount of
ammonia is found in air and soil in free state. It is formed when nitrogenous compounds
decay in the absence of air (oxygen). In the form of compound, ammonia is found in
Ammonium nitrate (NH4NO3), Ammonium sulphate [(NH4)2SO4], etc.

Laboratory preparation of ammonia gas

Principle

In laboratory, ammonia gas is prepared by heating two parts of ammonium chloride

(NH4Cl) and one part of calcium hydroxide [Ca(OH)2].

2NH4Cl (s) + Ca (OH)2 (s) CaCl2 (s) + 2H2O (l) + 2NH3 (g)

Materials required
i. Apparatus

Hard glass test tube
Bunsen burner
Stand with a clamp (2)
Cork
Delivery tube
Lime tower
Gas jar
Match box or gas lighter

ii. Chemicals required

Ammonium chloride (NH4Cl)
Calcium hydroxide [Ca (OH)2]

Procedure
• First of all, a mixture of Ammonium chloride (NH4Cl) and Calcium hydroxide
[Ca (OH)2] is made in ratio 2:1 and kept in a hard glass test tube.
• Apparatus is set as shown in the figure and the mixture is heated with the help
of a Bunsen burner.

• When the mixture of Ammonium chloride and Calcium hydroxide is heated,
ammonia gas is produced which is passed through lime tower.

• Dry ammonia gas is collected in the gas jar by downward displacement of air.

194 GREEN Science (Chemistry) Book-10

Mixture of NH4Cl Gas jar Stand
and Ca(OH)2 Ammonia gas Moist red
litmus
Hard glass test tube Delivery tube paper

Bunsen burner Lime tower

CaO

Fig.

10.3
Laboratory Preparation of Ammonia (NH3) gas

Precautions
i. The apparatus is made airtight.

ii. The gas is collected in the gas jar by downward displacement of air because ammonia
is lighter than air and highly soluble in water.

iii. The hard glass test tube should be kept in inclined position facing the mouth of
the test tube downwards to prevent it from cracking due to evaporation of water
produced during the chemical reaction.

iv. Ammonia gas should be passed through lime tower to get dry ammonia because
calcium oxide absorbs moisture present in the gas.

Test of ammonia gas

i. When a moist red litmus paper is inserted in the gas jar containing ammonia, the
litmus turns into blue because ammonia is basic in nature.

NH3 + H2O NH4OH

ii. Ammonia gas can be identified from its strong pungent odour.

iii. White fumes of Ammonium chloride (NH4Cl) are formed when a glass rod
dipped in conc. Hydrochloric acid is kept in the gas jar containing ammonia gas.

NH3 + conc. HCl NH4Cl

Some other methods of preparation of ammonia gas

1. Ammonia gas can be prepared by heating ammonium salts like Ammonium
carbonate, Ammonium sulphate, Ammonium chloride, etc.

(NH4)2 CO3 (s) H2O (l) + CO2 (g) + 2 NH3(g)
(NH4)2 SO4 (s) H2O(l) + CO2 (g) + 2NH3 (g)

GREEN Science (Chemistry) Book-10 195

NH4Cl (s) HCl(g) + NH3 (g)

2. Ammonia gas can also be prepared by reacting ammonium salts with strong bases/

alkalis.

(NH4)2 SO4 (s) + 2NaOH (aq) Na2 SO4 (aq) + 2H2O (l) + 2 NH3 (g)
NH4 Cl (s) + KOH (aq) KCl (s) + H2O (l) + NH3 (g)
NH4NO3 (s) + NaOH (aq)
NaNO3 (s) + H2O (l) + NH3 (g)

Manufacture of ammonia gas

In industrial scale, ammonia gas is prepared by heating one part nitrogen gas and three
parts hydrogen gas under high temperature and pressure. This process is called Haber's
process.

500°C, Fe/Mo

N2 + 3H2 200 – 900 atm. 2NH3 + Heat

This process is reversible. So this reaction is very slow under normal conditions. Therefore,
following conditions are required to increase the rate of chemical reaction.

Temperature about 500° C

Pressure 200 – 900 atm. (atmospheric pressure)

Catalyst Iron (Fe)

Promoter Molybdenum (Mo) Do You Know

The method of manufacture of ammonia was More ammonia can be prepared under
discovered by German Chemist Haber. So this high pressure. However, it is dangerous to
process is called Haber's Process. apply high pressure as it may explode.

Physical Properties of Ammonia gas Do You Know

1. Ammonia is a colourless gas.

2. It has a strong pungent odour which A promoter is a substance that enhances
may produce tears in eyes. the function of a catalyst. For example,
Molybdenum.
3. It is highly soluble in water.

4. It is lighter than air.

5. It turns moist red litmus paper into blue as it is basic in nature.

6. It neither burns itself nor supports burning.

7. It solidfies at –78°C and liquifies at –33.4°C.

Chemical Properties of Ammonia gas

1. Ammonia gas is highly soluble in water. It forms Ammonium hydroxide when
dissolved in water.

NH3 (g) + H2O (l) NH4OH (aq)

196 GREEN Science (Chemistry) Book-10

2. Ammonia reacts with acids and produces salt.

2NH3 (g) + H2SO4 (aq) (NH4)2SO4 (aq)
NH3 (g) + HNO3 (aq) NH4NO3 (aq)

3. Ammonia reacts with conc. Hydrochloric acid and forms solid Ammonium chloride.

NH3 (g) + conc. HCl (aq) NH4Cl (s)

4. Ammonia reacts with oxygen and forms greenish yellow flame which contains
nitrogen and water.

4NH3 (g) + 3O2 (g) 6H2O (l) + 2N2 (g)

5. Ammonia reacts with carbon dioxide at about 1500° C and under certain pressure,
(30 atm.), it forms urea and water.

NH3 (g) + CO2 (g) 1500°C NH2 CONH2 (s) + H2O (l)
Pressure

6. Ammonia solution, i.e. Ammonium hydroxide reacts with acid and forms salt and
water.

NH4OH + HCl (aq) NH4Cl (aq) + H2O (l)
2NH4OH (aq) + H2SO4 (aq) (NH4)2 SO4 (aq) + 2H2O (l)
NH4 OH (aq) + HNO3 (aq) NH4NO3 (aq) + H2O (l)

Uses of Ammonia gas

1. Liquid ammonia is used in refrigerator for cooling purpose.
2. It is used for manufacturing nitric acid, plastic, washing soda, alkalis, etc.
3. It is used to develop blue print of maps.
4. It is used as a cleansing agent to remove oil, grease, etc.
5. It is used for making chemical fertilizers like urea, ammonium sulphate, ammonium

chloride, ammonium nitrate, etc.
6. It is used for manufacturing ammonium salts like NH4Cl, (NH4)2 SO4, etc. that are

used in medicines.
7. It is used for making dyes, rayon, nylon, explosives, etc.
8. It is used in cold stores for cooling purpose.

9. It is used in water and waste water treatment such as pH control.

10. It is used in rubber, leather and paper industries.

11. It is used as a source of nitrogen for yeast and microorganisms in food and beverage
industries.

GREEN Science (Chemistry) Book-10 197

Key Concepts

1. Carbon dioxide is a compound gas having molecular formula CO2.
2. Carbon dioxide was discovered by Van Helmont in 1630 AD by burning wood.
3. In atmosphere, carbon dioxide occupies 0.03% by volume. All animals and plants

release carbon dioxide in air while breathing.
4. In laboratory, carbon dioxide is prepared by the chemical reaction between pieces

of limestone (CaCO3) with dilute hydrochloric acid (HCl).
5. Carbon dioxide gas should be collected in the gas jar by upward displacement of

air.
6. Carbon dioxide is heavier than air. So, it can be collected in the gas jar by upward

displacement of air.
7. When a burning piece of wood is inserted in the gas jar, it extinguishes. It shows that

the gas is carbon dioxide because carbon dioxide is non-supporter of combustion.
8. In industrial scale, carbon dioxide gas is prepared by heating limestone or calcium

carbonate in a kiln.
9. When carbon dioxide is passed in the clear solution of lime water for a while, the

lime water turns milky due to formation of water insoluble calcium carbonate.
10. When carbon dioxide is passed in lime water for a long time, milky colour disappears

due to formation of water soluble calcium bicarbonate.
11. Green plants use carbon dioxide gas to prepare food in leaves in the presence of the

sunlight. This process is called photosynthesis.
12. Ammonia is a compound gas having molecular formula NH3.
13. Ammonia gas was discovered by Lavoisier by heating the mixture of ammonium

chloride (NH4Cl) and Calcium hydroxide [Ca(OH)2].
14. Ammonia gas is found in free as well as in the form of mixture in nature. Some

amount of ammonia is found in air and soil in free sate.
15. Ammonia gas should be passed through lime tower to get dry ammonia because

calcium oxide absorbs moisture present in the gas.
16. Ammonia gas can be prepared by heating ammonium salts like Ammonium

carbonate, Ammonium sulphate, Ammonium chloride, etc.
17. In industrial scale, ammonia gas is prepared by heating one part nitrogen gas and

three parts hydrogen gas under high temperature and pressure. This process is
called Haber's process.
18. Ammonia reacts with carbon dioxide at about 1500° C and under certain pressure,
(30 atm.), it forms urea and water.
19. Ammonia is used for manufacturing nitric acid, plastic, washing soda, etc. It is also
used to develop blue print of maps.

198 GREEN Science (Chemistry) Book-10

Sequential General Exercise 1

1. Choose the best answer from the given alternatives.
a. The molecular weight of carbon dioxide is ...............

42 amu 44 amu 17amu 6 amu

b. When limestone pieces are heated in a kiln, it forms .................

calcium hydroxide and water

calcium oxide and water

calcium oxide carbon dioxide

carbon oxide and water

c. Which of the following gases is used to extinguish fire?

CO2 NH3 N2 O2

d. Which of the given gases convert moist red litmus into blue?

NH3 CO2 O2 H2

e. What is the ratio of ammonium chloride and calcium hydroxide to prepare
ammonia gas in laboratory?

1:2 2:1 2:3 3:1

2. Answer the following questions.

a. Write down the molecular formula and molecular weight of carbon dioxide.

b. Where is carbon dioxide gas found in nature?

c. How is carbon dioxide gas prepared in laboratory? Write with the balanced
chemical equation.

d. Draw a neat and labelled figure showing the laboratory preparation of carbon
dioxide.

e. Write any three methods of preparation of carbon dioxide gas.

f. How is carbon dioxide gas prepared in industries? Describe in brief.

g. Write any four physical properties of carbon dioxide.

h. Write any four chemical properties of carbon dioxide with balanced formula
equations.

GREEN Science (Chemistry) Book-10 199

i. Write down the major uses of carbon dioxide.

j. Where is ammonia gas found in nature? Write its molecular formula and
molecular weight.

k. How is ammonia gas prepared in laboratory? Write with the balanced chemical
equation.

l. How is ammonia gas prepared in industries?

m. Write any three physical properties of ammonia gas.

n. Write any four chemical properties of ammonia gas.

o. Write down the major uses of ammonia gas.

3. Study the given figure and answer the following questions:

i. Which gas is collected in the gas jar?
ii. Write down the balanced chemical equation involved in this process.

iii. Why is the gas jar kept erect?
iv. How can you test this gas?

Dil. hydrochloric acid Delivery tube

Thistle funnel

Gas jar

Pieces of CaCO3

4. Study the given figure and answer the following questions.
i. Which gas is collected in the gas jar?
ii. Write down the balanced chemical equation involved in this process.
iii. How can we test this gas?
iv. Write down one sure test of this gas.
v. What is the method of collection of this gas?

200 GREEN Science (Chemistry) Book-10