Science and Environment 10

Reasonable Fact-6

Potassium is more reactive than sodium although they both belong to the same group, why?
Ans: Potassium is more reactive than sodium because the atomic size of potassium atom
is larger than that of the sodium. So the valence electron of potassium can more easily be
taken by other reacting atoms as compared to that of sodium. Hence, potassium is more
reactive than sodium.

Reactivity of non-metals in a group

The chemical reactivity of non-metals decreases on moving from top to bottom in a group
of the modern periodic table. When we move from top to bottom in a group of non-metals,
the atomic size increases gradually. Hence, the nuclear attraction for incoming electron/s
decreases. As a result, the tendency of the non-metallic atom to gain electron/s decreases.
Therefore, the chemical reactivity of non-metals decreases on moving from top to bottom in
a group.
Example: In group VIIA/17, the chemical reactivity of halogens decreases from fluorine (F) to
Astatine (At).

Group VIIA/17

F Most reactive non-metal

Cl

Reactivity of non-metals

Br decreases on moving from
top to bottom in a group

I

At Least reactive non-metal

Reasonable Fact-7

Elements of group VA, VIA and VIIA (or group 15, 16, and 17) are less reactive as they
go down in the group of the Modern periodic table, why?
Ans: The atomic size of elements increases as we move downwards in groups. The elements
of group VA, VIA and VIIA (or group 15, 16 and 17) are non-metals. The chemical reactivity
of non-metals decreases when the atomic size of those elements increases. Therefore, the
elements of group VA, VIA and VIIA (or group 15, 16 and 17) are less reactive as they go
down in the group of the Modern periodic table.

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Reactivity of elements in a period

The atomic size of elements decreases gradually on moving from left to right in a period of the
Modern periodic table. So the reactivity of metallic elements decreases and the reactivity of
non-metals increases on moving from left to right in a period.

Elements of 3rd period Na Mg Al Si P S Cl

Chemical reactivity Most Least Most
reactive reactive reactive

→ Reactivity decreases → → Reactivity increases →

In the given table, Na is the most reactive metal among Na, Mg, and Al. Si is the least reactive
element. Similarly, Cl is the most reactive non-metal among P, S and Cl.

Reasonable Fact-8

Why fluorine is more active than chlorine although both lie in the same group?
Ans: Since the atomic size of fluorine is smaller than that of chlorine, nuclear attraction is
more on the valence shell of the fluorine atom. As a result, fluorine can get one electron
more easily during chemical reaction. Therefore, fluorine is more reactive than chlorine
although both of them belong to the same group.

Reasonable Fact-9

Chlorine is more reactive than bromine, why?
Ans: Chlorine is more reactive than bromine because the atomic size of cholorine is smaller
than that of bromine. Due to the smaller atomic radius of the chlorine atom, nuclear attraction
is greater on the valence shell of chlorine than that of bromine. Therefore, chlorine is more
reactive than bromine.

Worked out Example 1

Answer the following questions from the given electronic configuration.

A 1s2, 2s2 2p6, 3s1
B 1s2 , 2s2 2p6, 3s2 3p6, 4s1
C 1s2 , 2s2 2p6, 3s2 3p5
i. Name the elements A, B and C.
ii. Identify the period, group, valency, metal/non-metal and block of the given elements.
iii. Which compounds will be formed due to the chemical reaction between A with C and B with C?
What types of compounds are these and why?
iv. Which is more reactive out of A and B and why?

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Solution:
i. The elements are:
A - Sodium (Na)
B - Potassium (K)
C - Chlorine (Cl)
ii.

Element Period Group Valency Metal/Non-metal Block

A 3 IA/1 1 Metal s - Block

B 4 IA/1 1 Metal s - Block

C 3 VIIA/17 1 Non-metal p - Block

iii. Compounds formed due to chemical reaction between
A and C is NaCl (Sodium chloride)
B and C is KCl (Potassium chloride)
They are ionic (or electrovalent) compounds because they have an ionic

(or electrovalent) bond.

iv. B is more reactive than A because B has larger atomic size, less nuclear power and
hence it loses electron more easily than A during chemical reaction.

Variation in periodic properties in periods and groups

S.N. Properties Along a period Down in a group
(While moving from left
(While moving from top to
to right) bottom)

1. Atomic size (radius) Decreases Increases

2. Valency 1 to 4 and 4 to zero Remains the same

3. Metallic character Decreases Increases

4. Non-metallic character Increases Decreases

5. Ionization potential Increases Decreases

6. Electron affinity Increases Decreases

7. Electronegativity Increases Decreases

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SUMMARY

• Grouping of elements on the basis of their similarities and dissimilarities is
called classification of elements.

• According to Mendeleev's periodic law, "The physical and chemical properties
of elements are a periodic function of their atomic weights."

• In Mendeleev's periodic table, there are seven horizontal rows, i.e., periods and
eight vertical columns, i.e., groups.

• Modern periodic law states, "Physical and chemical properties of the elements
are a periodic function of their atomic numbers."

• The table which is obtained after arranging elements on the basis of increasing
atomic numbers is called the Modern periodic table.

• The 14 elements from cerium (58Ce) to lutetium (71Lu) after lanthanum (57La) are
called Lanthanides and other 14 elements from thorium (90Th) to lawrencium
(103Lr) after actinium (89Ac) are called Actinides.

• Aufbau principle states, "The electrons in an atom are so distributed that they
occupy shells in the order of their increasing energy."

• The atomic radius of elements increases gradually on moving from top to
bottom in a group of the Modern periodic table.

• The electrons present in the outermost shell (or valence shell) of an atom are
called valence electrons.

• In a period, the valency increases from 1 to 4 and then decreases to zero (0).
But in a group, valency of all elements remains the same.
• A metal which loses electron/s easily is called an active metal and a non-metal

which gains electron/s easily is called an active non-metal.
• The chemical reactivity of metals increases on moving down in a group of the

Modern periodic table.

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Exercise

Group-A

1. What is periodic table?

2. Name the scientist who propounded “atomic theory”.

3. State Mendeleev’s periodic rule.

4. Write down the number of group and period in Mendeleev’s periodic table.

5. State the Modern periodic law.

6. Which period is the longest period of Modern periodic table? How many elements are
there in that period.

7. What is electronic configuration?

8. What is duplet state? Write with one example.

9. How many periods are there in the Modern periodic table?

10. Write down the position of metals and non-metals in the Modern periodic table.

11. What is the position of alkali metals in the Modern periodic table?

12. What are alkaline earthmetals?

13. Write down the position of alkaline earthmetals in the Modern periodic table. What is
their valency?

14. What are halogens? Give one example.

15. What are inert gases? Give one example.

16. Write down the position of hydrogen in the Modern periodic table.

17. How many electrons are found in the outermost shell of transitional elements?

18. Write down the common name of the elements shown in the given table. F
19. What is a metalloid? Give one example. Cl
20. Name any two elements that give positively charged ions. Br
22. Name any two elements that give negatively charged ions. I

23. Name any two elements that can take or give equal number of electrons.

24. In which groups of the Modern periodic table are there alkali metals and inert gases?

25. What are transition metals?

26. What are Lanthanides?

27. Define Actinides.

18. In which periodic table elements are arranged on the basis of increasing atomic number?

29. Write down two factors that determine the reactivity of elements.

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30. Write the name of the groups, in which very active metals and very active non-metals are
placed in modern periodic table.

31. In which period do potassium and calcium lie in the Modern periodic table?
32. What is s-block element? Give an example.
33. What is the position of lanthanides in the Modern periodic table?
34. What is the position of halogens in the Modern periodic table?
35. Write the name of two elements of group 17 of the Modern periodic table.

Group-B
1. Compare and contrast between Mendeleev’s periodic table and Moseley’s periodic table.
2. What type of elements are kept in p-Block? Write with examples.
3. What is the valency of halogens? Why?
4. Why are halogens kept in the group 17 of the Modern periodic table?
5. Sodium is called an alkali metal, why?
6. Magnesium is called an alkaline earth-metal, why?
7. Lithium is less reactive than sodium, why?
8. Argon atom can exist freely in nature, why?
9. Metals of group 1 become more reactive as we go down in the periodic table, why?
10. Distinguish between Modern periodic table and Mendeleev’s periodic table in two

points.
11. Out of magnesium and calcium, which element is more reactive? Give reason.
12. Why does the reactivity of elements increase on moving from top to bottom in group 1

of Modern periodic table?
13. What difference in chemical reactivity of metals of second period occur while moving

from left to right in Modern periodic table?
14. Elements of group 15, 16 and 17 are less reactive as they go down in the group of the

Modern periodic table.
15. What is the change in chemical reactivity of very active non-metals when their atomic

size increases? Describe.
16. Why is potassium more reactive than sodium although they both belong to group 1?
17. Why is fluorine more active than chlorine although both lie in the same group?
18. Write any two reasons that hydrogen is kept in group 1 in the Modern periodic table.
19. On what factors the modern periodic table is different from that of the Mendeleev’s

periodic table? Write any two reasons.
20. Why is potassium more reactive than calcium although both lie in the same period?

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21. Which one is more reactive out of Fluorine and Chlorine? Why?

22. A small portion of Modern periodic table (Group 1) is given. Give two reasons of placing
hydrogen along with metals in this group.

23. Elements of group 1, 2, 13 are more reactive as we go down in the group of the periodic
table, why?

24. Modern periodic table is less defective than Mendeleev’s periodic table. Give two reasons
to justify it.

Group-C

1. Mention any three characteristics of Mendeleev’s periodic table.

2. Write down the three drawbacks of Mendeleev’s periodic table.

3. Write down three characteristics of Modern periodic table.

4. Which group of the periodic table does the element ‘A’ belong to? Which element is more
reactive in between ‘B’ and ‘C’? Give reason.

Element Electronic configuration
A 2, 8, 5
B 2, 3
C 2, 8, 3

5. Identify the groups of the alkali metals, metalloids and inert gases from the given table.
What is the valency of chlorine and neon? Why?

1 2 13 14 15 16 17 18
Li Be B C N O F Ne
Na Mg Al Si P S Cl Ar

6. Atomic number of elements ‘A’ and ‘B’ are 8 and 13 respectively. Answer the following
questions on the basis of that.

i) Give the electronic configuration of element ‘A’ and ‘B’ on the basis of subshell.

ii) Write down the molecular formula of the compound formed by the combination
of above elements.

Group-D

1. Molecular formula of a certain ionic compound is XY2 and ‘X’ is a metal. State group to
which elements X and Y belong to the periodic table. Which group of elements of the

periodic table are kept in the given table? What happens to the chemical reactivity of

elements from top to bottom in the given table? Why?

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2. Answer the following questions on the basis of the given table.

Name of element Electronic configuration

X 1s2 , 2s2 2p6, 3s1

Y 1s2, 2s2 2p6, 3s2 3p5

Z 1s2, 2s2 2p6, 3s2 3p6

i) Write the block of the elements X and Z.
ii) Write the valency and chemical nature of element Z.

iii) Write the balanced chemical equation between elements X and Y.

3. A part of periodic table is given below. Study it and answer the following questions.

Li Be B C N O F Ne

Na Mg Al Si P S Cl Ar

i) How are these elements arranged?
ii) Which one is more active in between Li and Na, and why?

iii) Write the formula of a compound made from Mg and Cl.

4. Give the name and group of an element whose electronic configuration is 1s2, 2s2, 2p6,
3s2, 3p6, 4s1. On what factors the Modern periodic table is different from that of the
Mendeleev’s periodic table? Write any two causes. Which one is more reactive between
Fluorine and Chlorine? Why?

5. Study the part of a periodic table and answer the following questions.

H He

Li Be B C N O F Ne

Na Mg Al Si P S Cl Ar

K C a

i) State the law on which given periodic table based.

ii) Which element belongs to group 17 and period 3?

iv) Which one is more reactive between S and Cl? Write with reason.



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8UNIT Estimated teaching periods

Theory 5

Practical 2

CHEMICAL

REACTION Avogadro

Objectives (1776–1856 AD)

After completing the study of this unit, students will be able to:
• define different types of chemical reactions.
• write chemical changes in the form of chemical equations.
• explain the rate of chemical reaction.
• explain the factors that bring chemical change.
• introduce catalyst with its types.

8.1 Introduction

When hydrogen gas (H2) burns in air (O2), water (H2O) is formed. Here, hydrogen and oxygen
combine to form water. When calcium carbonate (CaCO3) is heated, it decomposes into
calcium oxide (CaO) and carbon dioxide (CO2). Similarly, when iron (Fe) powder is kept in
copper sulphate solution (CuSO4), copper (Cu) and iron sulphate (FeSO4) are formed. Here,
iron displaces copper from copper sulphate solution. These are some examples of chemical

changes.

H2 + O2 burn H2O [Combination]

CaCO3 heat CaO + CO2 [Decompositon]

Fe + CuSO4 FeSO4 + Cu [Displacement]

The combination, decomposition or displacement that occurs in the molecules of matter
during a chemical change is called a chemical reaction. It can be represented by an equation.
The elements or compounds that take part in a chemical reaction are called reactants whereas
the elements or compounds that are formed as a result of chemical change are called products.
Reactants are written on the left side of an arrow and products are written on the right side of
the arrow. The direction of the arrow indicates the reactants and the products. For example,

Reactants Products Fact File - 1
Zinc + Sulphuric acid Zinc sulphate + Hydrogen
The chemical reaction
Zn + H2SO4 ZnSO4 + H2 is represented by word
equation and formula
equation /ɪˈkweɪʒn/ - a statement showing that two amounts are equal (chemical) equation.

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8.2 Word Equation

A chemical reaction expressed by writing the full names of reactants and products is called
word equation.

Examples:

Magnesium + Oxygen Magnesium oxide

Hydrochloric acid + Sodium hydroxide Sodium chloride + Water

8.3 Chemical Equation

A chemical reaction expressed by writing the symbols and molecular formulae of reactants
and products is called a chemical equation.

Examples:

2Mg + O2 2MgO
HCl + NaOH NaCl + H2O

8.4 Balanced Chemical Equation

A chemical equation written by balancing the total number of atoms of each element on the
reactant side and product side is called a balanced chemical equation.

Examples:

2KClO3 ∆ 2KCl + 3O2
∆ K2SO4 + 2H2O
H2SO4 + 2KOH

Methods of writing a balanced chemical equation

1. First of all, the chemical reaction should be written correctly in the form of a word
equation.

Example: Nitrogen + Hydrogen Ammonia

2. The word equation is written in the form of a formula or chemical equation.

Example: N2 + H2 NH3

3. The number of atoms of each element should be made equal by using a suitable coefficient
without changing the molecular formula.

Example: N2 + 3H2 2NH3

Some more examples of balanced chemical equations

1. Word equation : Potassium chlorate Heat Potassium chloride + Oxygen

Chemical equation : KClO3 ∆ KCl + O2
Balanced equation : 2KClO3 ∆ 2KCl + 3O2

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2. Word equation : Hydrogen peroxide Catalyst Water + Oxygen

Chemical equation : H2O2 MnO2 H2O + O2
Balanced equation : 2H2O2 MnO2 2H2O + O2

Information obtained from a balanced chemical equation
A balanced chemical equation provides the following information:

1. The names and symbols of reactants and products

2. The total number of atoms or molecules of reactants and products

3. The ratio of molecular weight of reactant and product molecules

4. The type of chemical reaction

The above information is illustrated by the given example:

Information Reactants Product
Chemical equation
N2 + 3H2 2NH3
Name Ammonia
Total molecules Nitrogen Hydrogen
Molecular weight
Type of reaction 13 2

28 6 34

Addition reaction or combination reaction

The above example also states that 28g of nitrogen reacts with 6g of hydrogen to give 34g
of ammonia. The ratio of molecular weight of nitrogen to ammonia is 14:17 and ratio of
molecular weight of hydrogen to ammonia is 3:17.

Limitations of a balanced chemical equation
A balanced chemical equation cannot provide the following information:

1. The physical state of reactants and products
2. The concentration of reactants
3. Conditions required for the reaction like heat, light, pressure, catalyst, etc.
4. The rate of chemical reaction
5. Time taken to complete the reaction
6. Whether the reaction is reversible or irreversible
7. Whether the reaction is exothermic or endothermic

8.5 Modification of Chemical Equation

To make a chemical equation more informative the following modifications are done.
1. The physical state of the reactants and products is denoted by 's' for solid, 'l' for liquid,

'g' for gas and 'aq' for aqueous solution.

reversible / r ɪ v 3 ː s ɪ b l / - a process that can be changed so that sth returns to its original state

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2. Concentration of reactants is denoted by 'dil.' for dilute and 'conc'. for concentrated.

3. The conditions like temperature, pressure, light, catalyst, etc. are written above or below
the arrow.

4. For reversible reaction, a double way arrow ( ) and for irreversible reaction, a single
way arrow (→) is used.

5. The symbol ∆ indicates heat, ↑ indicates gas and ↓ indicates the precipitate.

For example,

2Na(s) + 2H2O (l) 2NaOH (aq) + H2↑

N2(g) + 3H2(g) heat / pressure 2NH3 ↑

catalyst / promoter

2HgO (s) 2Hg (l) + O2 ↑

NaCl (aq) + AgNO3 (aq) NaNO3 (aq) + AgCl ↓

8.6 Reversible Reaction

A chemical reaction in which the products can recombine to give back the reactants is called a
reversible reaction. For example,

H2 + I2 heat 2HI

When hydrogen (H2) and iodine (I2) are heated, it forms hydrogen iodide (HI). When hydrogen
iodide is heated in a closed vessel, it also forms hydrogen and iodine. Therefore, the given

reaction is a reversible reaction. Reversible reactions are written by giving a double way arrow

between the reactants and products as follows:

H2 + I2 2HI
2NH3
N2 + 3H2 2H2O

2H2 + O2

8.7 Irreversible Reaction

A chemical reaction in which the products cannot recombine to give back the reactants is
called an irreversible reaction.

For example,

CaCO3 ∆ CaO + CO2↑

2Na+2H2O 2NaOH + H2↑

2KClO3 ∆ 2KCl + 3O2↑

precipitate /prɪˈsɪpɪteɪt/ - a solid substance that has been separated from a liquid in a chemical process

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8.8 Exothermic Reaction

A chemical reaction which evolves heat during the chemical change is called an exothermic
reaction. For example,

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

CH4 + 2O2

Zn + 2HCl

8.9 Endothermic Reaction

A chemical reaction which absorbs heat during the chemical change is called an endothermic
reaction.

For example,

N2 +O2 Heat 2NO ↑
Heat
2KClO3 2KCl + 3O2↑
CaCO3 Heat CaO + CO2↑
Heat NaCl + 2H2O + N2↑
NaNO2 + NH4Cl

8.10 Factors Which Affect a Chemical Reaction

There are several factors which play an important role during a chemical reaction. Some
important factors are discussed below:

1. Heat

Heat provides kinetic energy to the reactant molecules. Due to the absorption of heat
energy, the frequency of collision of these molecules increases to give more products.
There are certain reactions which do not initiate without heat. So, heat plays an important
role as a reaction initiator. Some chemical reactions that occur in the presence of heat are as
follows:

CaCO3 ∆ CaO + CO2 ↑
2 KClO3 ∆ 2KCl + 3O2 ↑

2. Light

Certain chemical reactions occur in the presence of sunlight. For example, photosynthesis,
chlorination of methane, decomposition of silver bromide, formation of hydrogen
chloride, etc.

6H2O + 6CO2 Sunlight C6H12O6 + 6O2↑
CH4 + Cl2 UV-Rays CH3Cl + HCl
2AgBr 2Ag + Br2
Light 2HCl
H2 + Cl2 Light

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3. Pressure

The reactant molecules having more volume than product cannot collide easily with each
other to give a product. Hence, to accelerate such a reaction, high pressure is needed. For
example, when pressure is applied on a mixture of potassium chloride and sulphur, it
explodes. A fire cracker explode on applying pressure. Similarly, the synthesis of ammonia
by Haber's process needs high pressure, i.e., about 200-500 atmospheric pressure.

N2 + 3H2 200-500 atm, 500 0C 2NH3↑
4. Catalyst
Fe/Mo

A catalyst is a chemical substance which increases or decreases the rate of chemical
reaction out itself remaining chemically unchanged. The mass and chemical nature of a
catalyst does not change during the chemical reaction. There are two types of catalysts.

a. Positive catalyst

A catalyst which increases the rate of chemical reaction is called a positive catalyst.

For example, manganese dioxide (MnO2) acts as a positive catalyst during the
decomposition of hydrogen peroxide.

2H2O2 MnO2 2H2O + O2↑
(catalyst)

b. Negative catalyst

A catalyst which decreases the rate of chemical reaction is called a negative
catalyst. For example, glycerine acts as a negative catalyst and decreases the rate
of the given chemical reaction.

2H2O2 Glycerine 2H2O + O2↑
5. Solution

Acid doesn't show any effect without water. Similarly, salt doesn't give ions without

water. This shows that a solution also plays an important role in many chemical

reactions. When silver nitrate and calcium chloride are brought in contact, they do not

show any reaction. But if they are dissolved in water, they give corresponding ions; and

after exchanging ions, products are formed.

AgNO3 (aq) + NaCl (aq) NaNO3 (aq) + AgCl↓

2AgNO3 (aq) + CaCl2 (aq) Ca(NO3)2 (aq) + 2AgCl↓

6. Direct contact

No contact, no collision, no reaction. Some active substances react on coming in direct
contact with others. For example, when sodium is kept in chlorine solution, it gives
sodium chloride (NaCl).

2Na + Cl2 2NaCl
7. Electricity

Electricity is also one of the important factors that brings about chemical reaction. When

electricity is passed in acidified water, it decomposes and gives hydrogen (H2) and
oxygen (O2) gas.

H2O Electricity 2H2 + O2

collision /kəˈlɪʒ(ə)n/ - the situation in which two things crash into each other

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8.11 Types of Chemical Reaction

There are basically four types of chemical reactions on the basis of their nature, which are
described below:

1. Combination or addition or synthesis reaction

A chemical reaction in which two or more reactants combine together to give a single
product is called addition reaction. This type of chemical reaction occurs either in the
presence or absence of heat, light, pressure, electricity, catalyst, etc.

Examples:

N2 + 3H2 2NH3
C + O2 CO2
C + 2H2 CH4
2 Na + Cl2 2NaCl
Fe + S
FeS
CaO + H2O
Ca (OH)2
CaCO3 + H2O + CO2 Ca(HCO3)2

2. Decomposition or dissociation or analysis reaction

A chemical reaction in which a single reactant is broken down into two or more products
is called decomposition reaction. Such type of chemical reaction occurs in the presence
of heat, light, catalyst, electricity, etc.

Examples:

CaCO3 ∆ CaO + CO2

2KClO3 ∆ 2KCl + 3O2

CuCO3 ∆ CuO + CO2

2Ag2O ∆ 4Ag + O2

2HgCO3 ∆ 2Hg + 2CO2 + O2

2Pb(NO3)2 ∆ 2PbO + 4NO2 + O2

2Cu(NO3)2 ∆ 2CuO + 4NO2 + O2

2 AgNO3 ∆ 2Ag + 2 NO2 + O2

3. Displacement reaction or replacement reaction
A chemical reaction in which an atom or a radical of a compound is displaced by another

element is called displacement or replacement reaction. It is of two types.

i. Single displacement reaction

A chemical reaction in which one atom or one radical is displaced by another
element is called single displacement reaction.

Examples:

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Zn + 2HCl ZnCl2 + H2
Zn + ZnSO4 + H2
2KBr + H2SO4 2KCl + Br2
2KI + Cl2 2 KCl + I2
Zn + Cl2 ZnSO4 + Cu
CuSO4

Mg + ZnCl2 MgCl2 + Zn

Fe + CuSO4 FeSO4 + Cu
Cu + 2AgCl CuCl2 + 2Ag

ii. Double displacement reaction

A chemical reaction in which an element or a radical of a compound is mutually
displaced by an element or a radical is called double displacement reaction.

Examples:

AgNO3 + NaCl NaNO3 + AgCl

CaCl2 + 2AgNO3 Ca(NO3)2 + 2AgCl

HgCl2 + 2KI 2KCl + HgI2

Pb (NO3)2 + Na2SO4 PbSO4 + 2NaNO3

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

4. Acid-base reaction

A chemical reaction in which an acid and a Reasonable Fact-1

base react together to give salt and water is Why is acid-base reaction called
called acid-base reaction. Here, acidic and neutralization reaction?
basic nature of the compounds is neutralized
during the chemical reaction, so it is also called Ans: In acid-base reaction, both acid
neutralization reaction. However, all acid-base and base lose their properties during
reactions are not neutralization reactions. the chemical reaction. As a result, a
neutral substance (i.e., salt and water)
are formed. So acid-base reaction is
also called neutralization reaction.

Examples:

Acid + Base Salt + Water

HCl + NaOH NaCl + H2O
CaO CaSO4 + H2O
H2SO4 + 2NaOH Na2SO4 + 2H2O
H2SO4 + CuO CuCl2 + H2O
2HCl + Ca(OH)2 Ca(NO3)2 + 2H2O
2KOH K2SO4 + 2H2O
2HNO3 +
H2SO4 +

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8.12 Gram Atomic Weight

The atomic weight of an element expressed in gram is called gram atomic weight of that
element. For example, the gram atomic weight of oxygen is 16 gram because the atomic weight
of oxygen is 16. Similarly, the gram atomic weight of sodium is 23 gram because the atomic
weight of sodium is 23.

8.13 Gram Molecular Weight

The molecular weight of a substance expressed in gram is called gram molecular weight of
the substance.

1. Gram molecular weight of CaCO3 = Ca × 1 + C × 1 + O × 3
40 × 1 + 12 × 1 + 16 × 3
=

= 40 + 12 + 48 = 100 gram

∴ The gram molecular weight of CaCO3 is 100 gram.

2. Gram molecular weight of H2SO4 = H×2+S×1+O×4

= 1 × 2 + 32 × 1 + 16 × 4

= 2 + 32 + 64

= 98 gram

∴ The gram molecular weight of H2SO4 is 98 gram.

8. 14 One Mole

The amount of a substance which contains 6.023 × 1023 atoms, ions, molecules or particles is
called one mole. Thus, a mole is a quantity of any material which contains one Avogadro's
number, i.e., 6.023 × 1023 particles.

Fact File -2

One mole of a substance = 6.023 × 1023 atoms, ions, molecules or particles. One mole of a gas
has 22.4 liter volume.

Fact File -3

Avogadro's number is the number of particles (atoms, molecules or ions) in one mole of a
substance. It is 6.023 × 1023. It is denoted by NA.

Worked out Example 1

1. Calculate the amount of carbon dioxide and calcium oxide produced when 200 g of
limestone is heated so that it decomposes completely.

Solution:

CaCO3 CaO + CO2
12 + 32
40 + 12 + 48 40 + 16 44
56
100

CHEMISTRY Oasis School Science - 10 159

According to the above reaction,

100g of CaCO3 gives 44g of CO2

1g of CaCO3 gives 44 g of CO2
100

200 g of CaCO3 gives 44 × 200
100

= 88 g of CO2

Again, 100g of CaCO3 gives 56 g of CaO

1 g of CaCO3 gives 56 g of CaO
100

200 g of CaCO3 gives 56 ×200 of CaO
100

= 112 g of CaO

∴ When 200 gram of CaCO3 is decomposed completely, it produces 88 g of CO2 and 112 g
of CaO.

8.15 Rate of Chemical Reaction

The positive change in the concentration of a reactant or a product per unit time is called the
rate of chemical reaction. It is calculated by the given formula:

Rate of chemical reaction = Change in concentration of a reactant or a product

Time taken for the change

The SI unit of the rate of chemical reaction is moles per liter per second (mol/l/s).

Various factors affect the rate of chemical reaction. Some major factors are:
i) Concentration of reactants
ii) Temperature
iii) Physical nature of reactants

iv) Chemical nature of reactants

i) Concentration of reactants

The rate of chemical reaction Reasonable Fact-2
increases on increasing the
concentration of reactants. Write down the effect of concentration of

It is because increasing the reactants on the rate of chemical reaction.
Ans: The rate of a chemical reaction increases
concentration of reactants increases on increasing the concentration of reactants.
the number of reacting molecules. It is because increasing the concentration of

There will be more collision reactants increases the number of reacting

between the reacting molecules, molecules. There will be more collision between
and hence the rate of chemical the reacting molecules and hence the rate of
reaction increases. chemical reaction increases.

160 Oasis School Science - 10 CHEMISTRY

ii) Temperature

Increase in temperature increases the rate of a chemical reaction whereas decrease in
temperature decreases the rate of chemical reaction. The energy of the reacting molecules
increases on increasing the temperature of the reactants. Due to increased energy, the
frequency of collision increases, and finally the rate of chemical reaction increases.

iii) Catalyst

A positive catalyst increases the rate of chemical reaction. In the presence of a positive
catalyst, the reaction takes place faster and at low temperature.

For example,

2KClO3 360 0C 2KCl + 3O2↑
240 0C 2KCl + 3O2↑
2KClO3
MnO2

Reasonable Fact-3

Manganese dioxide (MnO2) is called a positive catalyst, why?
Ans: Manganese dioxide (MnO2) increases the rate of chemical reaction. So, it is called a
positive catalyst.

From the above example, it becomes clear that potassium chlorate (KClO3) can be
decomposed into potassium chloride (KCl) and oxygen (O2) by heating at about 360 0C,
but the same reaction takes place at 240°C by using MnO2 as a positive catalyst.

iv) Physical nature of reactants

The reaction takes place in the contact area of the reacting molecules. So the rate of
chemical reaction can be increased by increasing the area of contact between the
reactants. Similarly, the rate of chemical reaction increases by using a common solvent if
the reactants are not soluble in one another. The common solvent helps bring the reacting
molelcules closer, and hence the rate of the reaction increases.

v) Chemical nature of reactants

Chemical nature of reactants also determines the rate of chemical reaction. Some
reactants are more active and some are less reactive. The elements of group IA and VIIA
are more reactive than the elements of other groups. Similarly, the reactions between
ionic compounds are faster than the reactions between covalent compounds.

CHEMISTRY Oasis School Science - 10 161

SUMMARY

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

• A chemical reaction expressed by writing the full names of reactants and products is
called a word equation.

• A chemical reaction expressed by writing the symbols and molecular formulae of the
reactants and products is called a chemical equation.

• A chemical equation written by balancing the total number of atoms of each element
on the reactant side and product side is called a balanced chemical equation.

• A chemical reaction in which the products can recombine to give back the reactants is
called a reversible reaction.

• A chemical reaction in which the products cannot recombine to give back the reactants
is called a irreversible reaction.

• A chemical reaction which evolves heat during the chemical change is called an
exothermic reaction.

• A catalyst is a chemical substance which increases or decreases the rate of chemical
reaction, but itself remaining chemically unchanged.

• A chemical reaction in which two or more reactants combine together to give a single
product is called addition reaction.

• A chemical reaction in which a single reactant is broken down into two or more
products is called decomposition reaction.

• A chemical reaction in which an atom or a radical of a compound is displaced by
another element or radical is called displacement or replacement reaction.

• A chemical reaction in which an acid and a base react together to give salt and water
is called acid-base reaction.

• The amount of a substance which contains 6.023 × 1023 atoms, ions, molecules or
particles is called one mole.

• Avogadro's number is the number of particles (atoms, molecules or ions) in one mole
of a substance. It is 6.023×1023.

• Positive change in the concentration of a reactant or a product per unit time is called
the rate of chemical reaction.

• The rate of chemical reaction increases on increasing the concentration of reactants.

• Increase in temperature increases the rate of a chemical reaction whereas decrease in
temperature decreases the rate of chemical reaction.

• In the presence of a positive catalyst, a reaction takes place faster and at low
temperature.

• The rate of chemical reaction can be increased by increasing the area of contact between
the reactants.

162 Oasis School Science - 10 CHEMISTRY

Exercise

Group-A
1. What is a chemical reaction? Give one example.
2. What is a word equation? Write with an example.
3. Define formula equation with one example.
4. What is a balanced chemical equation? Give one example.
5. What is a precipitate? Give one example.
6. What is a catalyst? How many types of catalyst are there?
7. Define negative catalyst with one example.
8. In which condition do iodine and phosphorus react?
9. Define exothermic reaction with one example.
10. Define endothermic reaction with one example.
11. How many types of chemical reactions are there? Name them.
12. Define decomposition reaction with one example.
13. What is displacement reaction?
14. What is acid-base reaction? Give one example.
15. Write any two pieces of information that can be obtained from a balanced chemical

equation.
16. What is meant by the rate of chemical reaction?
17. List the factors which affect the rate of a chemical reaction.

18. What is gram molecular weight?

Group-B
1. Write any two differences between reactants and products.
2. Write any two differences between synthesis reaction and dissociation reaction.
3. Write any two differences between positive catalyst and negative catalyst.
4. Write down the limitations of a balanced chemical equation.
5. Acid-base reaction is also called neutralization reaction, why?
6. The rate of a chemical reaction increases on increasing the concentration of reactants,

why?
7. The rate of a chemical reaction increases on increasing temperature, why?
8. The rate of a chemical reaction increases on powdering the reactants, why?

Group-C
1. In what condition do sodium chloride and silver nitrate react? Write the balanced

chemical equation of that reaction.

2. 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?

CHEMISTRY Oasis School Science - 10 163

Potassium Chlorate MnO2 Potassium Chloride + Oxygen

3. Write the balanced chemical equation for the following word equation:
i) Potassium chlorate Heat Potassium Chloride + Oxygen

4. Give one example of exothermic chemical reaction. Write any two characteristics of
catalyst.

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

Group-D

1. 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) Aluminimum + Hydrochloric acid Aluminimum chloride + Hydrogen

2. 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?

3. 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.

Fe + CuSO4 FeSO4 + Cu

4. Describe an experiment to demonstrate that the rate of chemical reaction increases on
increasing the temperature of reactants.

5. 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.

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9UNIT Estimated teaching periods

Theory 7

Practical 2

ACID, BASE

AND SALT Svante Arrhenius
(1859–1927 AD)
Objectives

After completing the study of this unit, students will be able to:
• define acid, base and salt.
• explain characteristics of acid, base and salt.
• classify acids on the basis of source and strength.
• differentiate between acids and bases.
• differentiate between bases and alkalis.
• write neutralization reactions between acids and bases and balance them.

9.1 Introduction

We eat different types of vegetables and fruits. Among them, some have a sour taste like an
orange, lemon, grape, apple, etc., some have a bitter taste like the bitter gourd, edible soda,
etc. and some have a salty taste like table salt. In our surroundings, there are several kinds of
compounds which cannot be tasted due to their corrosive and poisonous nature. These edible and
non-edible compounds are divided into three groups, i.e., acids, bases and salts.

9.2 Acid

The word acid is derived from the Latin word "acidus", which means sour in taste. Most acids,
either edible or non-edible, are sour in taste except boric acid, stearic acid and salicylic acid.
But it is dangerous to touch or taste acids in the laboratory. Due to their taste, some acids like
acetic acid, citric acid, ascorbic acid, etc. are used as flavour. According to Swedish chemist
Svante Arrhenius, "acids give hydrogen ions when dissolved in water and conduct electricity".
So, "acids are those chemical substances which give hydrogen ions when dissolved in water."

HCl H+ + Cl–

HNO3 H+ + NO3–

H2SO4 2H+ + SO4– –

CH3COOH H+ + CH3COO–

The hydrogen ions (H+ ) never exist by themselves in a solution. They attach to the water

molecule and form hydronium ions(H3O+). So, hydronium ion is nothing but a hydrated

acid /ˈæsɪd/ - a chemical substance that gives hydrogen ions when dissolved in water

CHEMISTRY Oasis School Science - 10 165

hydrogen ion. An acid cannot show its properties in the absence of water. Dry or pure
concentrated acid does not turn blue litmus paper into red. Thus, in the presence of water,
hydrochloric acid gives a proton to water forming the hydronium ion and chloride ion.

HCl H+ + Cl –
H3O+ (Hydronium ion)
H+ + H2O
Or, H3O+ + Cl–

HCl + H2O

9.3 Classification of Acids

1. Classification of acids on the basis of strength
There are two types of acids on the basis of strength. They are strong acids and weak acids.

a. Strong acids: Acids which undergo almost complete dissociation in aqueous solution
and produce high concentration of hydrogen ions are called strong acids. Due to more
hydrogen ion concentration, they are good conductors of electricity and have low pH
value. Examples: Hydrochloric acid (HCl), sulphuric acid (H2SO4), nitric acid (HNO3),
etc.

Note: The degree of dissociation of nitric acid is 92%. This means that out of
100 molecules of nitric acid, 92 molecules dissociate into ions.

b. Weak acids: Acids which undergo Reasonable Fact-1
partial dissociation in aqueous solution
and produce low concentration of Acetic acid is called a weak acid, why?
hydrogen ions are called weak acids.
Due to low concentration of hydrogen Ans: Acetic acid is an acid having low degree
ions, they do not conduct electricity of ionization. It produces low concentration of
easily and have high pH value. hydrogen ions in its solution. So, acetic acid is
called a weak acid.

Examples: Acetic acid (CH3COOH), carbonic acid (H2CO3), formic acid (HCOOH),
ascorbic acid, (H2C6H6O6), etc.

Note: Phosphoric acid (H3PO4) is a weak acid. Its degree of dissociation is about 26%.
This means that out of 100 molecules of phosphoric acid, about 26 molecules

undergo dissociation to give ions. When the degree of dissociation is more
than 30%, it is considered a strong acid.

2. Classification of acids on the basis of chemical nature

There are two types of acids on the basis of chemical nature or source. They are organic acids
and inorganic acids.

a. Organic acids: Acids which are obtained from living Fact File - 1
organisms and have hydrocarbons are called organic

acids. Examples: Acetic acid (CH3COOH), formic Hydrochloric acid (HCl) is an
acid (HCOOH), ascorbic acid (H2C6H6O6), maleic acid inorganic acid which is found
(H2C4H2O4), etc. These are weak acids. So, they in gastric juice.

166 Oasis School Science - 10 CHEMISTRY

produce less concentration of hydrogen ions in aqueous solution and conduct very low
amount of electricity.

b. Inorganic acids (Mineral acids): Acids which Fact File - 2
are obtained from minerals and do not have

hydrocarbons are called inorganic acids or Why can't pickle be stored in metal
mineral acids. Inorganic acids may be strong vessels for a long time?

or weak. Inorganic acids are commonly used Ans: Pickles contain weak acids
which slowly react with metal vessels
in laboratories.
and corrode them. Therefore, pickles
Examples: Hydrochloric acid (HCl), nitric cannot be stored in metal vessels for a
long time.
acid (HNO3), sulphuric acid (H2SO4), carbonic
acid (H2CO3), etc.

9.4 Properties of Acids

a. Physical properties of acids Fact File - 3

1. Acids possess sour taste due to the presence Lemon juice consists of citric acid,
of hydrogen ions (H+). Most of the fruits are vinegar consists of acetic acid, apple
sour in taste. However, it is dangerous to juice consists of malic acid and sour
touch and taste acids in the laboratory as milk consists of lactic acid.
they burn our skin, tongue, etc.

2. Acids change blue litmus paper into red and
methyl orange into red.

3. Strong acids are corrosive in nature.

b. Chemical properties of acids

1. Acids react with bases/ alkalis and produce salt and water.

Acid + Base Salt + Water

HCl + NaOH NaCl + H2O

H2SO4 + CaO CaSO4 + H2O

Reasonable Fact-2

The taste of acid is sour, Why?
Ans: Acid contains hydrogen ions (H+ ions). These ions stimulate the taste buds of our
tongue that detect sour taste. Therefore, the taste of acid is sour.

`Reasonable Fact-3

Why is sulphuric acid called the king of chemicals, or kingly water?
Ans: Sulphuric acid (H2SO4) is widely used in industries and laboratories. Therefore,
it is called the king of chemicals, or kingly water.

corrosive /kəˈrəʊsɪv/ - tending to destroy something slowly by chemical action
CHEMISTRY Oasis School Science - 10 167

2. Dilute acids react with reactive metals and produce salt and hydrogen gas.

Metal + Dilute acid Salt + Hydrogen

Mg + 2HCl MgCl2 + H2↑
Zn + 2HCl ZnCl2 + H2↑
Zn + H2SO4 ZnSO4 + H2↑

3. Acids react with carbonates and bicarbonates and produce salt, water and carbon
dioxide.

Carbonate + Acid Salt + Water + Carbon dioxide

MgCO3 + H2SO4 MgSO4 + H2O + CO2↑
CaCl2 + H2O + CO2↑
CaCO3 + 2HCl 2NaCl + H2O + CO2↑
Water + Carbon dioxide
Na2CO3 + 2HCl Salt +

Bicarbonate + Acid

NaHCO3 + HCl NaCl + H2O + CO2↑
Ca(HCO3)2 + 2HCl CaCl2 + 2H2O + 2CO2↑

4. Acids react with sulphites and bisulphites to liberate sulphur dioxide.

Sulphite + Acid Salt + Water + Sulphur dioxide

CaSO3 + 2HCl CaCl2 + H2O + SO2↑
Salt + Water + Sulphur dioxide
Bisulphite + Acid

NaHSO3 + HCl NaCl + H2O + SO2↑

5. Acids react with sulphides to liberate hydrogen sulphide gas along with salt.

Sulphide + Acid Salt + Hydrogen sulphide

ZnS + 2HCl ZnCl2 + H2S

6. Acids dissolve in water and give hydrogen ions.

HCl H+ + Cl–
2H+ +
H2SO4 H+ + SO4– –
HNO3 NO3–

9.5 Uses of Acids

1. Sulphuric acid is used in industries for making drugs, detergents and chemical fertilizers.
It is widely used in laboratories and industries. Therefore, sulphuric acid is also called
the king of chemicals, or kingly water.

2. Hydrochloric acid is used in laboratories and in tanning and printing industries.

3. Nitric acid is used for making explosives, plastics and dyes.

4. Carbolic acid (phenol) is used to kill germs.

168 Oasis School Science - 10 CHEMISTRY

5. Boric acid is used for washing eyes and wounds.
6. Acetic acid (vinegar) is used for preserving and flavoring foods.
7. Citric acid is used in medicines, as a source of vitamin C and flavoring drinks.
8. Carbonic acid is used in soft drinks and soda water.
9. Oxalic acid is used to remove ink-stain.

10. Tartaric acid is used in baking powder.

9.6 Base

Metallic oxides and hydroxides are called bases. Fact File - 4
Most bases or metal oxides dissolve in water and
give hydroxyl ions. The bases that dissolve in Bases like PbO, HgO, CuO do not
water and produce hydroxyl (OH-) ions are called dissolve in water.

alkalis, e.g. NaOH, KOH, Ca(OH)2, Mg(OH)2, etc.
Some bases like PbO, BaO, HgO, CuO, etc. do not dissolve in water. So these compounds are

bases but not alkalis. Therefore, all alkalis are bases, but all bases are not alkalis.

Most bases, or metallic oxides, dissolve in water and form metal hydroxides.

Na2O + H2O 2NaOH

MgO + H2O Mg(OH)2

K2O + H2O 2KOH

CaO + H2O Ca(OH)2

Alkalis give hydroxyl (OH-) ions when dissolved in water.

NaOH +H2O Na+ + OH_

Mg(OH)2 +H2O Mg++ + 2OH_

KOH +H2O K+ + OH_

Ca(OH)2 +H2O Ca++ + 2OH_

9.7 Preparation of Bases

There are different methods by which we can prepare bases. They are:

1. By direct combination of metal with oxygen

2Ca + O2 2CaO
2Mg + O2 2MgO
4Na + O2 2Na2O
4Al + 3O2 2Al2O3
2Hg + O2 2HgO

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2. By thermal decomposition of metallic carbonates

CaCO3 ∆ CaO + CO2
ZnCO3 ∆ ZnO + CO2
CuCO3 ∆ CuO + CO2

3. By thermal decomposition of metallic nitrates

2 Ca (NO3)2 ∆ 2CaO + O2 + 4NO2
2 Cu (NO3)2 ∆ 2CuO + O2 + 4NO2
2 Pb (NO3)2 ∆ 2PbO + O2 + 4NO2

Differences between Bases and Alkalis

S.N. Bases S.N. Alkalis
1. Alkalis are metallic hydroxides.
2. Bases are metallic oxides. 1. All alkalis give hydroxyl ions when
dissolved in water.
3. All bases do not give hydroxyl ions 2. All alkalis are bases.

when dissolved in water. Examples: NaOH, KOH, Mg(OH)2,
Ca(OH)2 , etc.
All bases are not alkalis. 3.

Examples: MgO, CaO, Na2O, HgO,
Fe2O3, etc.

9.8 Strong and Weak Bases

Bases (alkalis) which give more amount of hydroxyl ions in aqueous solution are called strong
bases, e.g., Sodium hydroxide (NaOH), potassium hydroxide (KOH), calcium hydroxide
[Ca(OH)2]. They undergo almost complete dissociation and have a high pH value.

Bases (alkalis) which give less amount of hydroxyl ions in aqueous solution are called weak
bases, e.g., Ferric hydroxide [Fe(OH)3], copper hydroxide [Cu(OH)2], etc. They have less
degree of ionization and have a low pH value.

Reasonable Fact-4

All alkalis are bases, but all bases are not alkalis. Justify.

Ans: Bases are substances that may or may not dissolve in water but alkalis are bases
that dissolve in water. Insoluble bases are not alkalis. Therefore, we can say that all
alkalis are bases, but all bases are not alkalis.

Reasonable Fact-5

Ammonium hydroxide is called a weak alkali, why?
Ans: Ammonium hydroxide is an alkali which produces low concentration of hydroxyl
ions in a solution. So it is called a weak alkali.

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9.9 Properties of Bases or Alkalis

a. Physical properties of Bases or Alkalis Reasonable Fact-6

1. Bases are soapy in touch and bitter in taste. Why are bases (alkalis) bitter

2. Bases turn red litmus paper to blue, methyl in taste?
orange to yellow and phenolphthalein to pink. Ans: Bases (alkalis) contain

3. Strong bases or alkalis like NaOH, KOH OH- -ions which stimulate the

dissolve oil and grease. taste buds that detect bitter

4. Strong bases or alkalis burn our skin. taste. Therefore, bases (alkalis)
are bitter in taste.

b. Chemical properties of Bases or Alkalis

1. Bases or alkalis react with acids to form salt and water.

Base + Acid Salt + Water

2KOH + H2SO4 K2SO4 + H2O

NaOH + HCl NaCl + H2O

NaOH + CH3COOH CH3COONa + H2O

2. Bases or alkalis react with carbon dioxide and form corresponding carbonate and
water.

Alkali + Carbon dioxide Carbonate + Water

2NaOH + CO2 Na2CO3 + H2O
2KOH + CO2 K2CO3 + H2O
Ca(OH)2 + CO2 CaCO3 + H2O

3. Alkalis liberate ammonia from ammonium salts.

Alkali + Ammonium salt Salt + Water + Ammonia

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

4. Alkalis produce insoluble metal hydroxides when they react with heavy metal salts.

Alkali + Heavy metal salt Insoluble hydroxide + Salt

2NaOH + CuCl2 Cu(OH)2 + 2NaCl
2NaOH + CuSO4 Cu(OH)2 + Na2SO4
2NaOH + FeCl2 Fe(OH)2 + 2NaCl
3NaOH + FeCl3 Fe(OH)3 + 3NaCl
3NH4OH + FeCl3 Fe(OH)3 + 3NH4Cl
2NaOH + ZnCl2 Zn(OH)2 + 2NaCl

CHEMISTRY Oasis School Science - 10 171

9.10 Uses of Bases

1. Sodium hydroxide (NaOH) is used to make soaps, detergents, paper, etc. and in
purification of petroleum products.

2. Calcium hydroxide [Ca(OH)2], or slaked lime, is used for making mortar and bleaching
powder, to reduce hardness of water and to neutralize acidity of soil.

3. Potassium hydroxide (KOH) is used in batteries and to make soft soap.

4. Aluminium hydroxide [Al(OH)3] and magnesium hydroxide [Mg(OH)2] are used to
reduce hyperacidity of stomach.

5. Ammonium hydroxide (NH4OH) is used to remove grease and stains from clothes and
to make fertilizers.

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

9.11 Salt

Salt is a chemical substance which is formed by Reasonable Fact-7
partial or complete replacement of hydrogen ions of
an acid by a metal or ammonium radical. In general, Why is sodium chloride called a salt?
salts are neutral compounds, but some may be acidic
or basic in nature. The process by which acid and Ans: Sodium chloride is called a salt
base react together to give salt and water is called because it is formed by complete
neutralization reaction. In this reaction, H+ ions of replacement of hydrogen atoms of
the acid are completely replaced by a metal. hydrochloric acid by sodium metal.

Acid + Base Salt + Water

HCl + NaOH NaCl + H2O

H2SO4 + 2KOH K2SO4 + 2H2O

In acid-base reaction, there may be complete or partial replacement of hydrogen atom(s) of

acid by a metal or ammonium radical. For example,

NaOH + H2SO4 NaHSO4 + H2O

In this example, only one hydrogen atom from sulphuric acid is replaced by sodium atom.

2NaOH + H2SO4 Na2SO4 + 2H2O

Here, both the hydrogen atoms are replaced by sodium atoms.

A salt contains two types of radicals: basic radical and acidic radical. During the formation of
a salt, the radical derived from a base is called a basic radical and the radical derived from an
acid is called an acidic radical. For example,

NaOH + HCl NaCl + H2O

amorphous /əˈmɔːfəs/ - having no definite shape, form or structure
172 Oasis School Science - 10 CHEMISTRY

In the salt NaCl, the sodium radical (Na+) is derived from the base (NaOH). So, it is a basic
radical. Similarly, the chloride radical (Cl-) is derived from an acid (HCl). So, it is an acidic radical.

9.12 Types of Salts

On the basis of chemical nature, there are five types of salts. They are as follows:

i) Normal salts: Salts formed by the chemical

reaction between a strong acid and a strong Reasonable Fact-8

base or a weak acid and a weak base are Sodium bicarbonate is called an acid
called normal salts. Examples: Sodium salt, why?

chloride (NaCl), potassium chloride (KCl), Ans: Sodium bicarbonate is formed

potassium nitrate (KNO3), etc. Normal salts by the partial displacement of

are also called neutral salts. hydrogen atoms of the acid. So sodium

ii) Acid salts: Salts formed by partial bicarbonate is called an acid salt.

replacement of hydrogen ions of an acid by

metals are called acid salts. Examples : Sodium bisulphate (NaHSO4), calcium bicarbonate
[Ca(HCO3)2], etc.

iii) Base salts: Salts formed by partial replacement of hydroxyl ions of a base by an acidic
radical are called base salts. Examples : Basic zinc chloride or zinc hydroxy chloride
[Zn (OH)Cl], basic lead chloride or lead hydroxy chloride [Pb(OH)Cl], etc.

iv) Acidic salts: Salts formed by the chemical reaction between a strong acid and a weak

base are called acidic salts. For example, CuSO4, PbCl2, etc.

Strong acid + Weak base Acidic salt + Water

H2SO4 + CuO CuSO4 + H2O
2HCl + PbO PbCl2 + H2O

v) Basic salts: Salts formed by the chemical reaction between a weak acid and a strong base

(alkali) are called basic salts. For example, Na2CO3, CH3COONa, K2CO3, etc.

Weak acid + Strong base Basic salt + Water

H2CO3 + 2NaOH Na2CO3 + 2H2O

CH3COOH + NaOH CH3COONa + H2O

vi) Hydrated salts: The salts containing certain Reasonable Fact-9
molecules of water are called hydrated salts.
Example: Sodium carbonate (Na2CO3. 10H2O) Sodium chloride is salty in taste, why?

Calcium sulphate (CaSO4. 7H2O) Ans: Sodium chloride stimulates the
Copper sulphate (CuSO4.5H2O), etc. taste buds of our tongue that detect
salty taste. So, sodium chloride is
9.13 Properties of Salts satly in taste.

1. Generally, salts are neutral, but some may be acidic or basic in nature.

2. Most of the salts are water soluble, but chloride salts of silver and lead and sulphate salts
of lead and barium are insoluble.

CHEMISTRY Oasis School Science - 10 173

3. Salts of metals, like Na, K, Mg, Ca, Al and Ba are white or colorless whereas salts of Cu,
Co, Mn, Ni, Fe and Cr have color.

4. Salts which are formed by the neutralization of a strong acid and a strong base or a weak
acid and a weak base are neutral to indicators and salts of a strong acid and a weak base
or a weak acid and a strong base change the color of the indicators.

5. Salts are electrovalent compounds. They conduct electricity in a molten or solution state.
6. Some salts are amorphous whereas some are crystalline.

9.14 Preparation of Salts

Salts can be prepared by different methods. Some of them are given below:

1. By direct combination of metals and non-metals

Metal + Non-metal Salt

2Na + Cl2 2NaCl
2Fe + 3Cl2 2FeCl3
Fe FeS
+ S

2. By the reaction of metals with acid

Metal + Acid Salt + Hydrogen

Zn + H2SO4 ZnSO4 + H2
Mg + 2HCl MgCl2 + H2
CaSO4 + H2
Ca + H2SO4

Fe + H2SO4 FeSO4 + H2

3. By the reaction of acids with alkalis

Acid + Base Salt + Water

HCl + NaOH NaCl + H2O
+ 2H2O
(Strong acid) (Strong base) (Neutral salt)

H2SO4 + Zn(OH)2 ZnSO4

(Strong acid) (Weak base) (Acidic salt)

H2CO3 + NaOH Na2CO3 + 2H2O

(Weak acid) (Strong base) (Basic salt)

CH3COOH + NH4OH CH3COONH4 + H2O

(Weak acid) (Weak alkali) (Neutral salt)

4. By the reaction of carbonates and bicarbonates with acids.

Carbonates + Acid Salt + Water + Carbon dioxide

CaCO3 + 2HCl CaCl2 + H2O + CO2
H2O + CO2
ZnCO3 + 2HCl ZnCl2 +

anaemic / ə ˈ n i ː m ɪ k / - suffering from anaemia, i.e., deficiency of RBCs in blood

174 Oasis School Science - 10 CHEMISTRY

Bicarbonates + Acid Salt + Water + Carbon dioxide
2H2O + 2CO2
2NaHCO3 + H2SO4 Na2SO4 + 2H2O + 2CO2
Mg(HCO3)2 + H2SO4 MgSO4 +

5. By the reaction of metallic oxides with acids

Metallic oxides + Acid Salt + Water
H2O
MgO + H2SO4 MgSO4 + H2O
FeO + 2HCl FeCl2 +

9.15 Uses of Salts

1. Table salt (NaCl) is used in our foods and also as a preservative.
2. Sodium carbonate is used in the manufacture of soaps, detergents and glasses. It is also

used to reduce the hardness of water.
3. Calcium sulphate is used for plastering of fractured bones.
4. Sodium bicarbonate is used in baking powder, for reducing hyperacidity and in a fire

extinguisher.
5. Copper sulphate is used for making fungicides and in copper plating.
6. Aluminium chloride is used in a dry cells as a electrolyte.
7. Silver bromide is used in photography.
8. Anhydrous ferrous sulphate is used as a medicine to treat anaemic patients.
9. Ammonium chloride is used in a dry cell as an electrolyte.
10. Ammonium sulphate is used as a chemical fertilizer.
11. Zinc sulphate is used to make the white pigment of the eye.

Neutralization Reaction

When an acid and a base are mixed, both acid and base lose their properties and form two neu-
tral substances, i.e., salt and water. Such a chemical reaction is called a neutralization reaction.

Examples:

Strong acid + Strong base Neutral salt + Water

HCl + NaOH NaCl + H2O

Weak acid + Weak base Neutral salt + Water

H2CO3 + 2NH4OH (NH4)2CO3 + 2H2O

Applications of neutralization reaction
1. Farmers use lime to neutralize the acidity of soil.

2. Human beings use magnesium hydroxide to reduce hyperacidity.
3. Honeybees and ants inject formic acid into our body. It can be neutralized by using soap.
4. Acetic acid is used to neutralize the acid injected by the bumble bee in our skin.

CHEMISTRY Oasis School Science - 10 175

SUMMARY

• Acids are those chemical substances that give hydrogen ions when dissolved in
water.

• Acids which undergo almost complete dissociation in aqueous solution and
produce high concentration of hydrogen ions are called strong acids.

• Acids which undergo partial dissociation in aqueous solution and produce low
concentration of hydrogen ions are called weak acids.

• Acids which are obtained from living organisms and have hydrocarbons are called
organic acids.

• Acids which are obtained from minerals and do not have hydrocarbons are called
inorganic acids.

• Metallic oxides and hydroxides are called bases. The bases that dissolve in water
and produce hydroxyl (OH–) ions are called alkalis.

• Bases (alkalis) which give more amount of hydroxyl ions in aqueous solution are
called strong bases.

• Bases (alkalis) which give less amount of hydroxyl ions in aqueous solution are
called weak bases.

• Salt is a chemical substance which is formed by partial or complete replacement of
hydrogen ions of an acid by a metal or ammonium radical.

• A salt contains two types of radicals: basic radical and acidic radical.

• Salts formed by partial replacement of hydrogen ions of an acid by metals are
called acid salts.

• Salts formed by the chemical reaction between a strong acid and a strong base or a
weak acid and a weak base are called normal salts.

• Salts formed by partial replacement of hydroxyl ions of a base by an acidic radical
are called base salts.

• Most of the salts are water soluble, but chloride salts of silver and lead and sulphate
salts of lead and barium are insoluble.

• Salts of metals, like Na, K, Mg, Ca, Al and Ba are white or colorless whereas salts of
Cu, Co, Mn, Ni, Fe and Cr have color.

• Salts which are formed by the neutralization of a strong acid and a strong base or a
weak acid and a weak base are neutral to indicators and salts of a strong acid and
a weak base or a weak acid and a strong base change the color of indicators.

• Sodium carbonate is used in the manufacture of soaps, detergents and glasses. It is
also used to reduce the hardness of water.

176 Oasis School Science - 10 CHEMISTRY

Exercise

Group-A
1. Define acid with any two examples.
2. What is an inorganic acid? Give any two examples.
3. Define organic acids with any two examples.
4. Name the acid found in each of the given substances:
i) Juice of lemon ii) Vinegar
5. Mention any two physical properties of acids.
6. Define base and give any two examples.
7. What is an alkali? Give any two examples.
8. Write any two physical properties of bases.
9. Name any two alkalis that react with skin.
10. What is a salt? Give any two examples.
11. Write any two physical properties of salts.
12. Define basic salt with one example.
13. Define neutral salt with one example.
14. Name the bases which are used for given activities:
a) To soften hard water b) To make soft soap
15. Name the bases which are used for given activities:
a) To make plastics and chemical fertilizers
b) To make bleaching powder
16. Name the bases which are used for given activities.

a) To purify sugar b) To purify petroleum products

Group-B
1. What are strong acids and weak acids? Give any two examples of each.
2. Why is sodium chloride salty in taste?
3. What type of salts are neutral? Explain with examples.
4. Write any two differences between acid and base.
5. Write any two differences between alkali and base.
6. Write any two differences between strong acid and weak acid.
7. Write any two differences between concentrated acid and dilute acid.
8. H2SO4 is called an acid, why?
9. All alkalis are bases but all bases are not alkalis, why?
10. Citric acid is called a weak acid, why?
11. Sodium hydroxide is called a base but sodium chloride is called a salt. Why?
12. Hydorchloric acid is kept in plastic or glass bottle, why?

CHEMISTRY Oasis School Science - 10 177

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

Group-C

1. What happens when acid is reacted with metallic carbonate? Write with the balanced
chemical equation.

2. Name the compounds which are formed due to the chemical reaction between Nitric
acid and Calcium bicarbonate.

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

4. Write any three uses of acids.

5. Write any three uses of salts.

6. Give an application of each of the given compounds:

a) Hydrochloric acid b) Nitric acid c) Carbonic acid

7. Give an application of each of the given compounds:

a) Sodium chloride b) Citric acid c) Magnesium hydroxide

8. Give an application of each of the given compounds:

a) Calcium sulphate b) Sodium hydroxide c) Sulphuric Acid

9. Give an application of each of the given compounds:

a) Copper sulphate b) Sodium carbonate c) Sodium bicarbonate

10. Give an application of each of the given compounds:

a) Carbolic acid b) Quick lime c) Ferrous sulphate

11. Give an application of each of the given compounds:

a) Ammonium chloride b) Ammonium sulphate c) Vinegar

12. Write down the following chemical changes into balanced chemical equations.

a) Metal oxide + Acid Metal salt + Water

b) Acid + Metal Salt + Water

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

a) Alkali + Ammonium salt Salt + Gas + Water

b) Acid + Base Salt + Water

Group-D

1. Write any two chemical properties of acids with balanced chemical equation.

2. Write any two chemical properties of bases with balanced chemical equation.

3. What happens in the given conditions? Write with balanced chemical equation.

i) When a metal reacts with dilute sulphuric acid.

ii) When an alkali reacts with ammonium salt.

4. Write short note on ‘different types of salt’.

5. Why water can be considered as an acid as well as a base? Write the following reactions
into balanced formula.

i) Sodium hydroxide + Hydrochloric acid Sodium chloride + Water

ii) Magnesium oxide + Sulphuric acid Magnesium sulphate + Water

178 Oasis School Science - 10 CHEMISTRY

UNIT10 Estimated teaching periods

Theory 5

Practical 2

SOME GASES Antoine Lavoisier
(1743–1794 AD)
Objectives

After completing the study of this unit, students will be able to:
• explain the laboratory preparation of ammonia and carbon dioxide.
• explain the properties and uses of ammonia and carbon dioxide.

A. AMMONIA

10.1 Introduction

Ammonia is a gaseous compound made up of one nitrogen atom and three hydrogen atoms.
It was first prepared by Antoine Lavoisier by heating ammonium chloride (sal ammoniac). But
its composition was studied by Berthecol and Davy.

Ammonia occurs in free as well as in combined states. In a free state, it is present in air and soil
whereas in a combined state, it is found in different ammonium salts like ammonium sulphate
[(NH4)2SO4], ammonium phosphate [(NH4)3PO4], ammonium chloride [NH4Cl], etc. Ammonia
gas is also synthesized in nature by nitrifying bacteria.

AMMONIA NH3
Molecular formula : 17
Molecular weight :

p+ = 1
n0 = 0

p+ = 7
n0 = 7

p+ = 1 p+ = 1
n0 = 0 n0 = 0

Fig 10.1: Molecular structure of ammonia (NH3)
CHEMISTRY Oasis School Science - 10 179

10.2 General Methods of Preparation of Ammonia Gas

1. By heating ammonium salts

Ammonia gas can be prepared by heating ammonium salts, like ammonium sulphate

[(NH4)2SO4], ammonium chloride (NH4Cl), ammonium carbonate, (NH4)2CO3], etc.
∆
(NH4)2SO4 2NH3 + H2SO4
∆
NH4Cl NH3 + HCl

(NH4)2CO3 ∆ 2NH3 + H2O + CO2

2. By heating ammonium salts with strong bases

Ammonia gas can be prepared by the reaction of ammonium salts, like ammonium

chloride (NH4Cl), ammonium sulphate [(NH4)2 SO4], etc. with strong bases like sodium
hydroxide (NaOH), potassium hydroxide (KOH), etc.

NH4Cl + KOH ∆ KCl + H2O + NH3↑
(NH4)2SO4 + 2NaOH
∆ Na2SO4 + 2H2O + 2NH3 ↑

10.3 Laboratory Preparation of Ammonia Gas

Principle

In the laboratory, ammonia gas is prepared by heating a mixture of ammonium chloride and
calcium hydroxide (slaked lime) in the ratio 2:1.

2NH4Cl + Ca(OH)2 ∆ CaCl2 + 2H2O + 2NH3 ↑

Apparatus required

Hard glass test tube, delivery Hard glass test tube Ammonia gas
tube, gas jar, Bunsen burner, Stand
stands, red litmus paper, cork

Chemicals required NH4Cl+Ca(OH)2 Gas jar Stand
Delivery tube Moist red
i) Ammonium chloride [NH4Cl] litmus paper
ii) Calcium hydroxide [Ca(OH)2]
Burner

Procedure Fig. 10.2 Laboratory preparation of ammonia gas

1. Make a mixture of ammonium chloride and calcium hydroxide in the ratio 2:1.
2. Keep the mixture in a hard glass test tube, and arrange the apparatus as shown in

fig. 10.2.

3. Supply heat from the Bunsen burner.

4. After heating, ammonia gas is produced. It is collected in the gas jar by downward
displacement of air.

slaked lime /sleɪk laɪm/ - a white powder made by combining lime with water

180 Oasis School Science - 10 CHEMISTRY

Reasonable Fact-1

Ammonia gas cannot be collected in a gas jar either by the downward displacement of
water or upward displacement of air, why?

Ans: Ammonia gas cannot be collected in a gas jar by the downward displacement of water because
this gas is highly soluble in water. It forms ammonium hydroxide (NH4OH) when passed in water.
Ammonia gas is lighter than air. So, it cannot be collected in the gas jar by the upward displacement
of air.

Reasonable Fact-2

A moist red litmus paper is used to test ammonia gas, why?
Ans: Ammonium hydroxide is an alkali. So, a wet red litmus paper is used to test if the gas is
ammonia or not.

Precautions
1. The apparatus should be made air tight.
2. The hard glass test tube should be slightly slanted.
3. Ammonia gas is collected by downward displacement of air.
Test of ammonia gas
1. Ammonia gas turns moist red litmus paper into blue because it is basic in nature.
2. It has a strong pungent odor.
3. When a glass rod dipped in conc. hydrochloric acid (HCl) is brought in contact

with ammonia gas, it forms a white fume of ammonium chloride (NH4Cl).

Fact File - 1

We bring moist red litmus paper near the mouth of the gas jar to test whether the gas jar is filled by
ammonia gas or not. Since ammonia gas is basic in nature, it turns moist red litmus paper blue.

Fact File - 2

The gas jar is inverted because ammonia is lighter than air. So, this gas is collected by the downward
displacement of air.

Fact File - 3

Ammonia gas is not collected by the downward displacement of water similar to hydrogen and
oxygen gas because it is highly soluble in water and gives ammonium hydroxide.

NH3 + H2O NH4OH
(Ammonium hydroxide)


Fact File - 4

A hard glass test tube is slightly inclined during preparation of ammonia gas, otherwise the steam
produced during the reaction exerts pressure, which may crack the test tube. So, the hard glass test
tube is slanted to prevent it from cracking.

Fact File - 5

Lime tower (CaO) is used to obtain dry and pure ammonia gas.

Fact File - 6

A large amount of ammonia gas can be prepared at high pressure. We should not maintain high
pressure to avoid an explosion.

CHEMISTRY Oasis School Science - 10 181

10.4 Industrial Preparation of Ammonia Gas

For commercial purpose, ammonia gas is manufactured by Haber's process. In this process,
a mixture of nitrogen and hydrogen (1:3) is combined directly at 500 0C temperature and 200
to 600 atmospheric pressure. For this reaction, finely powdered iron is taken as a catalyst and
molybdenum as a promoter. If the catalyst is not used, the rate of reaction becomes extremely
slow, and we need to supply high heat and high pressure.

N2 + 3H2 500 0C/200 to 600 atm. 2NH3
Fe (catalyst) and Mo (promoter)

Note:
A promoter is a chemical substance that makes a catalyst more active during a chemical
reaction, e.g., Molybdenum (Mo).

Conditions required for Haber's process

A German chemist Fritz Haber, in 1913 AD, discovered the following conditions for the
industrial preparation of ammonia gas.

1. Temperature 500 0C
2. Pressure 200 to 600 atm.
3. Catalyst Iron powder (Fe)

4. Promoter Molybdenum (Mo)

10.5 Properties of Ammonia Gas

a. Physical properties

1. It is a colorless and tasteless gas having a strong pungent odor.

2. It is lighter than air.

3. It is highly soluble in water.

4. It turns moist red litmus paper blue.

5. It is a non-combustible gas.

b. Chemical properties
1. Ammonia dissolves in water and forms an alkali, i.e., ammonium hydroxide.

NH3 + H2O NH4OH

2. Ammonia is a basic gas, so it reacts with acid to give salt.

NH3 + HCl NH4Cl

2NH3 + H2SO4 (NH4)2SO4

NH3 + HNO3 NH4NO3

3. Ammonium hydroxide, i.e., ammonia solution, reacts with acid to give salt and water.

2NH4OH + H2SO4 (NH4)2 SO4 + 2H2O

combustible / k ə m ˈ b ʌ s t ə b l / - able to begin burning easily

182 Oasis School Science - 10 CHEMISTRY

NH4OH + HCl NH4Cl + H2O

NH4OH + HNO3 NH4NO3 + H2O

4. Ammonia and carbon dioxide react together at about 1500 °C and under certain pres-
sure (30 atm.) to give urea. It is a very useful chemical fertilizer rich in nitrogen.

2NH3 + CO2 1500 0C NH2 - CO - NH2 + H2O
Pressure (urea)

(30 atm.)

5. Ammonia reacts with conc. hydrochloric acid and forms solid particles of ammonium chloride.

NH3 + conc. HCl NH4Cl

6. Ammonia reacts with oxygen and forms nitrogen gas and water.

4NH3 + 3O2 2N2 + 6H2O

10.6 Uses of Ammonia Gas

1. It is used in the manufacture of chemical fertilizers, like ammonium nitrate, ammonium
phosphate, urea, etc.

2. Liquid ammonia is used in refrigerators and cold storages.

3. It is used in the manufacture of different types of industrial products, like nitric acid,
plastics, dyes, nylon, rayon, washing soda, chemical explosives, etc.

4. It is used to develop the blue print of maps.

5. It is used to make medicines like ammonium carbonate, ammonium chloride, etc.
6. Liquid ammonia and ammonia solution are used as a laboratory reagent.

Activity 1

To demonstrate that ammonia is highly soluble in water and basic in nature

A round bottom flask containing ammonia gas is NH3
taken, and a delivery tube is fitted on it as shown Round bottom flask

in the diagram. The other end of the delivery tube

is kept within the water containing some drops Stand
of phenolphthalein. The delivery tube carries a

jet inside the round bottom flask. When ammonia

gas cools, it creates low pressure inside the flask, Water
so some water drops rise up in the flask and form

ammonium hydroxide. As a result, more vacuum Fig. 10.3 Fountain experiment
is created inside the flask. Hence, water rushes up

with high pressure to fill up the vacuum, which looks like a fountain. Afterwards the

entire flask appears pink. This experiment proves that ammonia is highly soluble in water

and basic in nature.

Note: If we take kerosene in place of water, it does not form ammonium hydroxide
and does not give a pink color.

CHEMISTRY Oasis School Science - 10 183

SUMMARY

• The molecular formula of ammonia is NH3, and its molecular weight is 17.

• In the laboratory, ammonia gas is prepared by heating a mixture of ammonium
chloride and calcium hydroxide.

2NH4Cl + Ca(OH)2 CaCl2 + 2H2O + 2NH3↑

• Ammonia is a basic gas, so it turns moist red litmus paper blue.

• Ammonia is collected by the downward displacement of air because it is lighter
than air.

• For commercial purpose, ammonia is manufactured by Haber's process. In
this process, a mixture of hydrogen and nitrogen in the ratio 3:1 is heated at
5000C and 200 to 600 atm. pressure in the presence of an iron catalyst and a
molybdenum promoter.

N + 3H 500 0C, 200 to 600 atm. 2NH3
2 2 Fe and Mo

• Ammonia is highly soluble in water. It forms ammonium hydroxide, i.e.,
ammonia solution when dissolved in water.

NH3 + H2O NH4OH

• Ammonia reacts with acid to give ammonium salt.

NH3 + HCl NH4Cl

• Ammonia solution reacts with acid to give salt and water.

NH4OH + HCl NH4Cl + H2O

• Ammonia reacts with carbon dioxide at about 15000C and under certain pressure
(30 atm.) to give urea.

2NH3 + CO2 1500 0C NH2-CO-NH2 + H2O
Pressure

(30 atm.)

• Ammonia is used in the production of different types of industrial products,
like plastics, washing soda, dyes, nylon, rayon, chemical fertilizers, etc.

• Ammonia is used to develop a blue print, to make medicines and as a laboratory
reagent.

184 Oasis School Science - 10 CHEMISTRY

B. CARBON DIOXIDE

10.7 Introduction

Carbon dioxide is a gaseous compound made of one atom of carbon and two atoms of oxygen.
Carbon dioxide occurs in a free or combined state in nature. In the free state, it is present in
the atmosphere, about 0.03% by volume. In the atmosphere, carbon dioxide is reached by the
burning of fuels like diesel, petrol, kerosene, wood, coal, etc. and by the respiration of living
beings. In the combined state, it is present in different types of carbonates, bicarbonates, etc.
Since carbon dioxide is heavier than air, it is present at a lower level of the atmosphere and
occurs in deep wells, mines, caves, etc. Carbon dioxide has a great role to play in the living
world because it is used by green plants for photosynthesis to convert solar energy into chem-
ical energy. This chemical energy is supplied to all animals for their existence. Carbon dioxide
is produced by the reaction of limestone (CaCO3), dolomite (CaCO3.MgCO3) and magnesite
(MgCO3) with dilute acids.

Carbon Dioxide CO2
44
Molecular formula –
Molecular weight –




p+= 8 p+= 6 p+= 8
n0= 8 n0= 6 n0= 8

Fig 10.4: Molecular structure of CO2

10.8 General Methods of Preparation of Carbon Dioxide Gas

1. By the combustion of hydrocarbons like methane, ethane, etc.

When hydrocarbons like methane, ethane, etc. are burnt, carbon dioxide gas is produced.

CH4 + 2O2 CO2 + 2H2O
4CO2 + 6H2O
2C2H6 + 7O2

2. By the reaction of acid with carbonates and bicarbonates

When carbonates and bicarbonates of different metals, like calcium, magnesium, etc. are
treated with acid, carbon dioxide is produced.

CaCO3 + 2HCl CaCl2 + H2O + CO2
2NaCl + H2O + CO2
Na2CO3 + 2HCl CaCl2 + 2H2O + 2CO2

Ca (HCO3)2 + 2HCl

CHEMISTRY Oasis School Science - 10 185

3. By burning carbon in plenty of oxygen

When carbon burns in plenty of oxygen, carbon dioxide is formed.

C + O2 CO2

Fact File - 7

Dilute sulphuric acid cannot be used in place of dilute hydrochloric acid because
calcium sulphate is formed after the reaction, which covers the remaining part of
the marble and stops further reaction.

4. By heating limestone vigorously

When limestone or marble is heated
vigorously, carbon dioxide is formed.
CaCO3 ∆ CaO + CO2

10.9 Laboratory Preparation of Carbon Dioxide Gas

Principle
In the laboratory, carbon dioxide gas is prepared by the reaction of dilute hydrochloric acid
with calcium carbonate (limestone or marble).

CaCO3 + 2HCl CaCl2 + H2O + CO2↑

Apparatus required

Woulfe’s bottle, thistle funnel, gas jar, delivery tube, corks, matchbox

Chemicals required

1. Calcium carbonate or marble or limestone pieces (CaCO3)
2. Dilute hydrochloric acid (HCl)

Dilute hydrochloric acid

Thistle funnel Delivery tube
Cork Moist blue
litmus paper
Woulfe's bottle Carbon dioxide gas

Gas jar

Lime stone
pieces

Fig. 10.5 Laboratory preparation of carbon dioxide gas

186 Oasis School Science - 10 CHEMISTRY

Procedure

1. Keep some pieces of marble in a Woulfe’s bottle and arrange the apparatus as
shown in the figure.

2. Pour some dilute hydrochloric acid through the thistle funnel until it covers the
lower end of the thistle funnel and the marble pieces.

3. Chemical reaction takes place between the marble and hydrochloric acid, and it
produces carbon dioxide gas.

4. The gas is passed through the delivery tube and collected in the gas jar by the up-
ward displacement of air.

Fact File - 8

Carbon dioxide gas is collected in the gas jar by the upward displacement of air
because it is heavier than air.

Fact File - 9

The lower end of the thistle funnel should be dipped in the solution otherwise carbon
dioxide gas can escape.

Fact File - 10

Carbon dioxide gas is soluble in water. So it is not collected in the gas jar by passing
through water.

Test of carbon dioxide gas

1. A burning match stick is kept near the mouth of the gas jar. If the burning match
stick extinguishes, it proves that the gas jar is filled with carbon dioxide gas.

2. When carbon dioxide is passed through lime water, i.e. calcium hydroxide, it turns

milky white due to the formation of insoluble calcium carbonate (CaCO3).

CO2 + Ca(OH)2 CaCO3 + H2O

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

color disappears slowly due to the formation of water soluble calcium

bicarbonate [Ca(HCO3)2].

CaCO3 + H2O + CO2 Ca(HCO3)2

3. Carbon dioxide is acidic in nature so it turns moist blue litmus paper red.

Precautions
1. The apparatus should be made air tight.

2. The lower end of the delivery tube should not touch the solution within the
Woulfe's bottle.

3. The lower end of the thistle funnel should be dipped in the solution.
4. Carbon dioxide is collected in the gas jar by the upward displacement of air.

combustion / k ə m ˈ b ʌ s t ʃ ə n / - the process of burning

CHEMISTRY Oasis School Science - 10 187

10.10 Industrial Preparation of Carbon Dioxide Gas

For commercial purpose, carbon dioxide gas is manufactured by heating calcium carbonate (lime-
stone or marble). During this process, calcium oxide, i.e., lime, or quick lime, is also produced.

CaCO3 CaO + CO2↑

When calcium oxide reacts with water, it gives calcium hydroxide, which is also known as

slaked lime.

CaO + H2O Ca(OH)2

10.11 Properties of Carbon Dioxide Gas

a. Physical properties
1. It is a colorless, odourless and tasteless gas.
2. It is soluble in water and produces carbonic acid when dissolved in water.
3. It is heavier than air.
4. It is an acidic gas. So, it turns moist blue litmus paper red.
5. It is neither combustible nor a supporter of combustion.
6. When carbon dioxide gas is cooled to below -78°C, it is converted into solid form, which

is known as dry ice. It is known as dry ice because it melts without wetting the surface.

b. Chemical properties

1. Reaction with water

Carbon dioxide reacts with water and gives carbonic acid.

CO2 + H2O H2CO3 (Carbonic acid)
Fact File - 11

Dry ice is used in refrigerators to preserve foods, fruits, vegetables and meat.

2. Reaction with lime water

Carbon dioxide reacts with lime water and gives insoluble calcium carbonate (CaCO3),
which makes the solution milky-white.

CO2 + Ca(OH)2 CaCO3 + H2O

When carbon dioxide is passed continuously into lime water for a long time, the milky color
disappears slowly due to the formation of water soluble calcium bicarbonate Ca(HCO3)2.

CaCO3 + H2O + CO2 Ca(HCO3)2

3. Green plants convert solar energy into chemical energy by photosynthesis. In this
process, carbon dioxide and water react together to give starch and oxygen.

6CO2 + 6H2O sunlight C6H12O6 + 6O2
chlorophyll

188 Oasis School Science - 10 CHEMISTRY

4. Carbon dioxide is neither combustible nor a supporter of combustion. But a burning
magnesium strip burns in carbon dioxide gas with dazzling light. During this process,
white powder of magnesium oxide and black particles of carbon are produced.

2Mg + CO2 2MgO + C

5. Carbon monoxide is formed when carbon dioxide gas is heated with red hot coke at
about 9000C.

CO2 + C 900ºC 2CO

6. Ammonia reacts with carbon dioxide at about 15000C and certain pressure (30 atm.) to
forms urea and water.

2NH3 + CO2 1500ºC NH2 – CO – NH2 + H2O
pressure (30atm.)

(Urea)

10.12 Uses of Carbon Dioxide Gas

1. Green plants use carbon dioxide to prepare Fact File - 12
food during photosynthesis.
The mixture of 10-15% oxygen and

2. It is used for making soft as well as hard drinks carbon dioxide is called carbogen. It

like soda water, coca cola, beer, etc. is used in the artificial respiration of

3. It is used in the manufacture of fertilizers like pneumonic patients.
urea and washing soda, i.e., sodium carbonate.

4. Solid carbon dioxide, i.e., dry ice is used as a refrigerant to preserve foods, fruits, meat, etc.

5. It is used in fire extinguishers.

6. It is used in carbonation process to purify sugarcane juice.

10.13 Working Mechanism of a Fire Extinguisher Knob

A fire extinguisher is a protective device Nozzle for gas
which is used to extinguish a fire. It has
a red metallic vessel containing a sodium
bicarbonate solution and a bottle of
sulphuric acid.

When we need carbon dioxide gas to Concentrated H2SO4
extinguish a fire, the vessel is inverted.

Afterwards, the bottle of concentrated Saturated sodium bicarbonate
sulphuric acid strikes against the floor. solution
Metallic cylinder

The plug of the acid bottle falls down Fig. 10.6(a) Fire extinguisher Fig. 10.6 (b) Internal structure of a fire extinguisher
and acid comes in contact with the

sodium bicarbonate solution. When

those chemicals react together, they give carbon dioxide gas.

2NaHCO3 + H2SO4 (conc.) Na2SO4 + 2H2O + 2CO2↑

The gas produced in the above reaction comes out through the nozzle at high pressure and
puts off the fire, just like putting a blanket over the surface of the flame.

CHEMISTRY Oasis School Science - 10 189

SUMMARY

• The molecular formula of carbon dioxide is CO2 and its molecular weight is 44.
• Carbon dioxide is present in the atmosphere about 0.03% by volume.

• In the laboratory, carbon dioxide gas is prepared by the reaction of calcium
carbonate with dilute hydrochloric acid.

CaCO3 + 2HCl CaCl2 + H2O + CO2↑

• For industrial purpose, carbon dioxide is prepared by heating calcium carbonate.

CaCO3 ∆ CaO + CO2

• Carbon dioxide can be prepared by the combustion of hydrocarbons like
methane, ethane, etc.

CH4 + 2O2 CO2 + 2H2O

• Carbon dioxide is prepared from carbonates when they react with acid.

Na2CO3 + 2HCl 2NaCl + H2O + CO2↑

• Carbon dioxide dissolves in water and forms carbonic acid.

H2O + CO2 H2CO3

• CO2 is a colorless, odorless and tasteless gas of an acidic nature.

• CO2 is used for photosynthesis and to make soft drinks, dry ice, urea, etc.

• When carbon dioxide is cooled to below -78°C, it converts into solid form, which
is known as dry ice. It is used as a refrigerant to preserve foods.

• Carbon dioxide reacts with lime water to give insoluble calcium carbonate.

CO2 + Ca(OH)2 CaCO3 + H2O

• Plants convert solar energy into chemical energy by photosynthesis. In this

process, carbon dioxide and water react together to give starch and oxygen.

6CO2 + 6H2O sunlight C6H12O6 + 6O2↑
chlorophyll

• Solid carbon dioxide, i.e., dry ice is used as a refrigerant to preserve foods, fruits,
meat, etc.

• A fire extinguisher is a protective device which is used to extinguish a fire. It has a
red metallic vessel containing sodium bicarbonate and a bottle of sulphuric acid.

190 Oasis School Science - 10 CHEMISTRY

Exercise

Group-A
1. Write down the molecular formula and molecular weight of carbon dioxide.
2. Where is carbon dioxide found in nature?
3. What happens when carbon dioxide gas is cooled to 780C?
4. Write any two physical properties of carbon dioxide gas.
5. Write any two chemical properties of carbon dioxide gas.
6. Write down the molecular formula and molecular weight of ammonia.
7. Write any two physical properties of ammonia gas.
8. Name two chemicals which are required to prepare carbon dioxide gas in laboratory.
9. What is dry ice?
10. Which gas is obtained by reacting slaked lime and ammonium chloride?
11. What is Haber’s process?
12. What is promoter? Give an example of it.
13. Name two chemicals that are required to prepare ammonia gas in laboratory.
14. Which gas is obtained by reacting limestone and dilute hydrochloric acid?
15. Which gas is produced by the chemical reaction of dilute hydrochloric acid and marble

pieces?
16. Name the gas produced when a mixture of ammonium chloride and calcium hydroxide

is heated.
Group-B

1. Carbon dioxide gas is not collected over water, why?
2. A moist red litmus paper is used for testing ammonia gas, why?
3. Ammonia gas cannot be collected in an erect gas jar, why?
4. Write any two differences between carbon dioxide and ammonia gas.
5. Carbon dioxide gas is important for green plants, why?
6. Carbon dioxide can be tested with the help of moist blue litmus paper, why?
7. Why is carbon dioxide gas used in fire extinguisher?
8. It is dangerous to clean well remaining closed for a long time, why?
9. Lime water turns to milky when carbon dioxide gas is passed through it, why?
10. What happens when a burning magnesium is inserted in the jar containing carbon diox-

ide? Write with the chemical equation.

CHEMISTRY Oasis School Science - 10 191

Group-C
1. How is carbon dioxide gas prepared in laboratory? Write with the balanced chemical

equation.
2. Draw a neat and labeled diagram showing laboratory preparation of carbon dioxide gas.
3. What are three precautions that should be adopted during the preparation of carbon

dioxide in laboratory?
4. Write any three methods for testing carbon dioxide gas.
5. Write any three uses of carbon dioxide gas.
6. Draw a neat and labeled diagram showing laboratory preparation of ammonia gas.
7. Write any three precautions that should be adopted during laboratory preparation of

ammonia gas.
8. Mention any three methods for testing ammonia gas in laboratory.
9. Write any three uses of ammonia gas.
10. How is urea prepared from ammonia gas? Describe in brief.

Group-D

1. How is carbon dioxide manufactured in industrial scale? Write with chemical equation.

2. Study the given figure and answer the following questions.

Delivery
tube

dil.HCl

Woulfe’s
bottle

Pieces of
CuCO3

i) Name the gas which is being produced in the diagram.
ii) Why is the gas jar kept erect?
iii) How can we test this gas?

3. How is ammonia gas prepared in industrial scale? Describe.

4. Laboratory preparation of ammonia gas is given in the figure. What is the method to
test whether the gas jar is full or not? Why cannot this gas be collected in the gas jar by
downward displacement of water?

MixtuarnedoNf HC2aC(Ol H)2 Ammonia gas

Lime tower
CaO

192 Oasis School Science - 10 CHEMISTRY