science 10

Differences between s-block and p-block elements:

s- block elements p-block element

1. The last electron of these 1. The last electron of these elements
elements enters in s-sub enterers in p-sub shell.
shell.

2. They form only 2 columns 2. They form 6 columns in the

in the periodic table. periodic table.

3. The chemical reactivity 3. The chemical reactivity decreases
of elements increases on on moving downwards in a group.
moving from top to bottom of
a group.

3. d-block elements

The elements in which the last electron enters the d–sub shell of the
outermost shell are called d-block elements. The elements of group 3
to 12 are d-block elements. The d–block elements are also known as
transition metals.

Examples:

Scandium (Sc) → 1s2, 2s2, 2p6, 3s2, 3p6, 4s2, 3d1
Titanium (Ti) → 1s2, 2s2, 2p6, 3s2, 3p6, 4s2, 3d2
Vanadium (V) → 1s2, 2s2, 2p6, 3s2, 3p6, 4s2, 3d3
Chromium (Cr) →1s2, 2s2, 2p6, 3s2, 3p6, 4s1, 3d5
Manganese (Mn) → 1s2, 2s2, 2p6, 3s2, 3p6, 4s2, 3d5
Iron (Fe) → 1s2, 2s2, 2p6, 3s2, 3p6, 4s2, 3d6
Cobalt (Co) → 1s2, 2s2, 2p6, 3s2, 3p6, 4s2, 3d7
Nickel (Ni) → 1s2, 2s2, 2p6, 3s2, 3p6, 4s2, 3d8
Copper (Cu) → 1s2, 2s2, 2p6, 3s2, 3p6, 4s1, 3d10
Zinc (Zn) → 1s2, 2s2, 2p6, 3s2, 3p6, 4s2, 3d10

4. f -block elements

The lanthanides and actinides are f-block elements. The f–sub shell
of outermost shell of these elements is gradually filled up. These
elements are called inner transition elements because two series of
these elements act as transit between the groups 3 and 4.

The series of 14 elements each are placed at the bottom of the periodic
table.

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Advantages of modern periodic table

Modern periodic table corrected the drawbacks (demerits) of
Mendeleev's periodic table and made the position of elements
clear. There are several advantages of modern periodic table over
Mendeleev's periodic table. Some of them are:

1. Position of metal, metalloid and non-metal

There are separate positions for metal, metalloid and non-metals
in the modern periodic table. The metals are placed on the left
side of periodic table. The non metals lie on the right side of
periodic table which include inert gases, halogens and other
elements. The elements B, C, Si, Ge, As and Te are metalloid
and lie between metals and non-metals.

2. Position of hydrogen

Hydrogen is an element with one valence electron. It can lose
that electron to acquire a positive charge similar to that of alkali
metals. So, it is placed in group IA (group 1). It is the explanation
for placing hydrogen in group IA. Even though the justification
has been given for keeping hydrogen in group IA, the position
of hydrogen is still controversial due to its similarities with
halogens in some other properties.

3. Improvement in the position of elements

The faulty position of elements like argon, potassium, etc is
solved itself while arranging elements as a function of atomic
number.

4. Position of lanthanides and actinides

All the lanthanides have similar properties so they belongs to
group 3 and 6th period. Similarly, all the actinides belong to
group 3 and 7th period. But two series of each of these elements
are placed at the bottom of the periodic table to avoid the undue
sidewise expansion of the periodic table. All lanthanides have
similar chemical properties. Similarly, all actinides have similar
chemical properties.

5. Position of isotopes

The isotopes are the atoms of same element having same atomic
number but different mass number. The isotopes of elements are
due to the difference in number of neutrons in the atoms of same

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element. For example: Protium, deuterium and tritium are three
isotopes of hydrogen.

1n 1p 1p
1n 2n

Protium Deuterium Tritium

Isotopes of hydrogen

Since all the isotopes have same number of protons, they have same
atomic number. So all isotopes occupy same position in the modern
periodic table. Thus, the modern periodic table solved the problem of
position of isotopes.

Valence electrons and valency of elements

The electrons that are present in the outermost shell of an atom are
called valence electrons because they determine the valency of an
element. The valence electrons take part in chemical reactions and
the reactivity of elements depends upon them. A part of periodic
table is given below in which the valence electrons of some elements
are shown:

Group/ 1 2 13 14 15 16 17 0
Period
1 H Be B C N O F He
1 2, 2 2, 3 2, 4 2, 5 2, 6 2, 7 2
2 Li Mg Al Si P S Cl Ne
2, 1 2, 8, 2 2, 8, 3 2, 8, 4 2, 8, 5 2, 8, 6 2, 8, 7 2, 8
3 Na Ca Ar
2, 8, 1 2, 8, 8, 2 2, 8, 8
4 K
2, 8, 8, 1

Valency 1 2 3 4 3 2 1 0

Note: The bold face numbers represent the valence electrons.

According to octet rule, a group of eight electrons in the outermost
shell is always stable. The elements try to gain octet either by losing,
gaining or mutual sharing of their valence electrons. The number
of electrons lost, gained or shared by an atom during a chemical

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reaction is the valency of an element. Hence, the valence electrons
play an important role to determine the valency of an element.

The elements of group IA, IIA and IIIA lose all electrons of their
valence shell to attain octet. So, the valency of Na is 1, Mg is 2 and
Al is 3. The elements of group IVA (i.e, C, Si, etc) mostly share
electrons to attain octet. The elements of group VA to VIIA gain
electrons from other atoms to form octet. Hence, these elements are
the electronegative elements. But the elements of group '0' neither
lose, nor gain nor share electrons because they have their own octet
in their valence shell.

The valency of an element is determined with the help of valence
electrons. For example, Valency of elements of group IA to IVA = No.
of valence electrons.

Valency of group VA to VIIA and 'O' = 8 - No. of valence electrons.

Reactivity of the elements

The elements of a particular group have similar chemical properties
but they do not have the same properties. There is regular change
in reactivity of elements in a group. The reactivity of the elements
is largely affected by the size (radius) of their atoms. As we move
from left to right of a period, the nuclear charge (i.e, number of
protons) increases. But the addition of electrons takes place in the
same shell. As a result, the electrons are pulled closer to the nucleus.
This decreases the size of the atom. Hence, the atomic size of Mg is
smaller than that of Na, the atomic size of Al is smaller than that
Mg and so on.

1. The reactivity of metals: The reactivity of the metals (elements
of group IA, IIA and IIIA) increases along a group on going from
top to bottom of the periodic table. As we go down from top to
bottom of a group, the size of the atom increases due to the
addition of one new shell in each successive member. Thus, the
valence electrons become farther from the nucleus and the force
of attraction between the nucleus and valence shell decreases.
This enables an atom to lose valence electrons easily. Hence, the
chemical reactivity of metals increases on moving from top to
bottom of a group in the periodic table. For example, Na is more
reactive than Li, K is more reactive than Na and so on.

Potassium (K) is more reactive than sodium (Na) because

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potassium has larger size and thus less attractive force exists
between nucleus and valence electrons than that in sodium. So,
potassium loses electron easily than sodium.

2. The reactivity of non-metals: The reactivity of non-metals
(elements of group VA, VIA and VIIA) decreases from top to
bottom of a group. As we go down the group, the size of the
elements increases. Thus, the valence electrons become farther
from the nucleus and force of attraction between the nucleus and
valence shell decreases. This decreases the tendency of an atom
to gain (attract) electrons. Hence, the chemical reactivity of non-
metals decreases on moving down a group in the periodic table.
For example, fluorine is more reactive than chlorine; chlorine
is more reactive than bromine and so on. Among the halogens,
fluorine is the most reactive and iodine is the least reactive.

Fluorine is more reactive than chlorine because fluorine (F) has
smaller size and more electronegativity. Due to this, fluorine
attracts electron more strongly than chlorine.

As we move from left to right in a period, the chemical reactivity
of the elements first decreases and then increases after reaching
the least reactive element.

Learn and Write

1. Elements of group 17 are highly reactive electronegative
elements. Why?

Elements of group 17 have seven valence electrons. Therefore,
they readily take one electron from other atoms to make their
outermost orbit complete and become electronegative. Thus,
elements of group 17 are highly electronegative elements.

2. Potassium is more reactive than sodium. Why?

Potassium has bigger atomic size than sodium. Hence the
valence electron of potassium is farther from the nucleus than
in sodium and is attracted with less force by the nucleus. Thus,
valence electron can be thrown easily in potassium than in
sodium. Therefore, potassium is more reactive than sodium.

3. Sodium is s-block element. Why?

According to electronic configuration of sodium i.e. 1s2, 2s22p63s1,
its last electron is present in s-sub-shell. So, it is s-block element.

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Main points to remember

1. The scientific and systematic classification of elements in
definite rows and columns based on certain periodic law is
called periodic table.

2. The horizontal rows of elements in the periodic table are called
periods

3. The vertical columns of the elements having similar properties
in the periodic table are called groups.

4. Mendeleev's periodic law states that the physical and chemical
properties of elements are periodic function of their atomic
weights.

5. Modern periodic law states that the physical and chemical properties of
elements are periodic function of their atomic number.

6. There are seven periods and nine groups in the modern periodic
table.

7. The elements of group IA are called alkali metals because these
elements readily react with water to form alkali.

8. The elements of group IIA are known as alkaline earth metals because
they form basic hydroxides which are less soluble in water.

9. All the elements of sub-group A in the periodic table are known
as representative elements.

10. The isotopes are the atoms of same element having same atomic
number but different mass number.

11. The region around the nucleus where the probability of finding
electron is maximum is called orbital.

Exercise

A. Choose the best alternative.

1. How many elements were classified in Mendeleev's original
Periodic Table?

a.92 b. 109 c. 36 d. 63

2. The atoms of the same element having same atomic number
but different mass number are called

a. Isotopes b. Isobars

c. Anomalous pair d. Lanthanides

3. Who proposed Modern Periodic Law?

a. Dmitri Mendeleev b. John Dalton

c. Henery Moseley d. None

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4. The elements of group 17 are known as

a. Alkali metals b. Alkaline earth metals

c. Noble gases d. Halogens

5. The elements of group 3 to 12 are called

a. Transition metals b. d-block elements

c. Sub-group B elements d. All of the above

B. Answer these questions in brief:

1. What is periodic table?

2. Define the term group and period.

3. State Mendeleev's periodic law.

4. State Modern periodic law.

5. What is the common name given to the elements of group 1? Why?

6. Write down any four advantages of modern periodic table
over Mendeleev's periodic table.

7. Write down any four drawbacks of Mendeleev's periodic table.

8. Write down any two features of modern periodic table.

9. Write down the maximum number of electrons that can be
accommodated in p and d sub shells.

10. What is the relation between atomic size and reactivity of
elements in a group of metals?

11. In which groups of modern periodic table are highly reactive
metals and halogens placed?

12. Explain why does the reactivity elements of group 1 increase as
we go down the group.

13. Which one is more reactive between fluorine and chlorine? Why?

14. Why is potassium more reactive than calcium although both
lie in the same period?

15. What is p-block element? Give examples.

16. What do you mean by isotopes? Give examples.

17. What do you mean by lanthanides? Why are they kept
separately in the periodic table?

18. What is lanthanide series? What is the position (group and
period) of lanthanides in the modern periodic table?

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19. What are actinides? What is their position in the modern
periodic table?

20. Write down the electronic configuration of following elements
in terms of sub-shells.

Lithium, Nitrogen, Sodium, Potassium, Neon, Scandium,
Chromium, Copper, Zinc.

21. Electronic configuration of an element is given:

1s2, 2s2, 2p6, 3s2, 3p1

To which block of the modern periodic table does the given
element belong? Write balanced equation for the reaction
between this element and chlorine. Also write two uses of
the element.

C. Conceptual questions:

1. A small portion of modern periodic table (group IA) is given.

(i) Give two reasons for placing hydrogen Group IA
along with metals in this group. H
Li
(ii) What is the common name given to the Na
group of these elements? Why? K
Rb
(iii) Which one is more reactive Na or K? Cs
Why?

(iv) Which element is the most reactive

among these elements? Why?

D. Differentiate between:

1. Group and Period.

2. Mendeleev's periodic table and Modern periodic table.

3. s-block elements and p-block elements.

4. d-block elements and f-block elements.

5. Metals and Metalloids.

6. Alkali metals and Alkaline earth metals.

7. Lanthanides and Actinides.

E. Give reasons:

1. Potassium is more reactive than sodium although they both
belong to same group.

2. Chlorine is more reactive than nitrogen.

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3. Sodium is placed in s-block in the periodic table.
4. Fluorine is placed in p-block of periodic table.
5. Hydrogen is placed in group 1 along with alkali metals.
6. Li, Na and K are placed in the same group of the periodic table.
7. Argon is an inert gas.
8. Hydrogen can also be placed in group 17.
9. The atomic size increases as we move from top to bottom of a group.
10. The atomic size decreases as we move from left to right in a period.
11. The atomic size of aluminium is smaller than that of sodium.
12. Chlorine is less reactive than fluorine.
13. Oxygen is more reactive than nitrogen.

Project Work

Make a periodic table in a chart paper. Mark the position of s, p,
d and f-block elements. Also write the group and period numbers
of each columns and rows. Compare your work with that of your
friend.

Glossary

• Credit : praise or recognition for something done
the part played by somebody or something
• Contribution : in causing a result
forecast
• Predicated : an interval between two notes consisting of
eight notes
• Octaves : reasoning
the second last, just inner to the outermost
• Justification : arrangement
• Penultimate : following in an uninterrupted sequence
• Configuration:
• Successive :

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Chapter

8 Chemical Fritz Haber
Reaction

He discovered Haber's process, Born –
Haber Cycle, Fertilizer, Chemical warfare,

explosEivset,imetca. ted Periods: 7 (5T+2P)

Objectives

At the end of the lesson, students will be able to:
• explain various types of chemical reactions;
• to write chemical equations of the chemical changes;
• to explain the factors affecting rate of chemical reactions.

We observe several changes in our daily life. The conversion of one
form of matter, energy, etc into another is called change. The change
in matter can be categorized into two types: Physical change and
Chemical change.

Physical change

A substance undergoes changes in shape, size, mass, volume, state,
etc due to physical change. For examples, tearing of paper into
pieces, melting of ice into water, dissolving salt in water, making
shirt, pant, etc from raw cloth, etc.

Thus, a physical change is a temporary change in which no new
substance is formed. The physical change can be reversed by simple
physical means.

Chemical Change

A chemical change is a permanent change which involves the change
in molecules of matter. The chemical change cannot be reversed
easily. Examples: turning of milk into curd, burning of fuel, digestion
of food, rusting of iron, growth of baby into adult, etc.

Chemical reaction

A chemical change in which combination, decomposition or exchange
of molecules of matters takes place is called chemical reaction. The
chemical reaction can be represented by word equation and formula
equation.

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All chemical reactions are represented by chemical equations. A
chemical equation is the symbolic representation of actual chemical
reaction in terms of symbols and formulae. For example, sodium
reacts with chlorine to form sodium chloride.

Sodium + Chlorine → Sodium chloride

2Na + Cl2 → 2NaCl
Product
Reactants

A chemical equation consists of two parts: reactants and products,
which are separated by an arrow (→). The substances that undergo
chemical change during a chemical reaction are called reactants.
They are written on the left hand side of the chemical equation.
Similarly, a new substance which is formed as a result of chemical
change in the reaction is called product. The products are written on
the right hand side of the equation.

Sometimes, the other conditions required for the chemical reaction
are written over an arrowhead.

For example:

Potassium Chlorate Heat → Potassium Chloride + Oxygen

Mercuric Oxide Heat → Mercury + Oxygen

Hydrogen + Chlorine Sunlight → Hydrochloric acid

A chemical equation is also known as formula equation. But, the
equation represented by the name of chemical substances in words
instead of molecular formula is called word equation. The formula
equation which holds an unequal number of atoms of one or more
elements on either side of arrow is called unbalanced equation. The
unbalanced equation is also called skeleton equation. For example:

Word Equation:

Hydrogen peroxide → Water + oxygen

Skeleton equation (formula equation):

H2O2 → H2O + O2
Balanced equation:

2H2O2 → 2H2O + O2
In a balanced chemical equation, the number of atoms of each element
is same on both sides of the arrow. In the above chemical equation,
the number of hydrogen and oxygen atoms is same in both sides of
the arrow.

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Balanced chemical equation

The chemical reactions are represented by chemical equations using
molecular formula. But these equations may not contain same
number of atoms on both sides. To make the number of atoms of each
element equal on both sides of the equations, the chemical equation
is balanced. The unbalanced equation can be balanced by comparing
the number of atoms on two sides. Hence, a chemical equation in
which the number of atoms of each element on both sides of the
equation is equal is called a balanced chemical equation.

Balancing of chemical equation is done by writing coefficients before
the symbols or formula in such a way that the total numbers of
atoms of each element on both sides becomes equal. This method of
balancing chemical equation is called hit and trial method. The hit
and trial method is also known as hit and success method.

Method of writing balanced chemical equation

Use the following rules for writing balanced chemical equations:

(a) Write down the chemical equation in the form of word
equation. For example;

Potassium Chlorate Heat → Potassium Chloride + Oxygen

(b) Write down the symbol or molecular formula of every element
or compound involved as reactant or product. For example;

2KClO3 ∆ → KCl + O2
In case of elements, symbols are written. But in case of
compounds, molecular formula is used. However, molecular
formulae are used in case of diatomic elements such as H2,
N2, O2, F2, Cl2, Br2 and I2.
(c) Write a suitable coefficient (number) to the atom or molecule
to make the number of atoms either side equal.

2KClO3 ∆ → 2KCl + 3O2
The digit used for multiplying a molecule indicates its number.
For example, 2KClO3 means two molecules of KClO3, where
there are 2 atoms of potassium, 2 atoms of chlorine and 6
atoms of oxygen.

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Examples:

1. Chemical reaction: When silver bromide is exposed to light,
silver and bromine are produced.

Word equation:

Silver bromide light → Silver + Bromine
Formula equation:

AgBr light → Ag + Br2
Balanced equation:

2AgBr light → 2Ag + Br2

2. Chemical reaction: When a mixture of nitrogen and hydrogen
under high pressure is heated in the presence of catalyst,
ammonia is produced.

Word equation:

Nitrogen + Hydrogen Heat, pressure Ammonia
Formula equation: Catalyst, promoter

N2 + H2 Heat, pressure NH3
Catalyst, promoter

Balanced equation:

N2 + 3H2 Heat, pressure 2NH3
Catalyst, promoter

3. Chemical reaction: When calcium carbonate is heated, calcium
oxide and carbon dioxide are produced.

Word equation:

Calcium carbonate ∆ → Calcium oxide + Carbon dioxide

Formula equation:

CaCO3 → CaO + CO2
The equation is balanced itself.

4. Chemical reaction: The strong heating of silver nitrate gives
silver, nitrogen dioxide and oxygen.

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Word equation:

Silver nitrate ∆ → Silver + Nitrogen dioxide + Oxygen

Formula equation:

AgNO3 ∆ → Ag + NO2 + O2
Balanced equation:

2AgNO3 ∆ → 2Ag + 2NO2 + O2
5. Chemical reaction: When a mixture of ammonium chloride

and calcium hydroxide is heated, ammonia gas, calcium chloride
and water are produced.

Word equation:

Ammonium Chloride + Calcium hydroxide → Calcium Chloride +
Ammonia + Water

Skeleton equation:
NH4Cl + Ca(OH)2 → CaCl2 + NH3 + H2O
Balanced chemical equation:
2NH4Cl + Ca(OH)2 → CaCl2 + 2NH3 + 2H2O

Endothermic and exothermic reactions

The reactions which require heat are called endothermic reactions.
The compounds formed by the absorption of heat are called
endothermic compounds. For example:

CaCO3 (s) ∆ → CaO (s) + CO2 (g)
2KClO3 (s) ∆ → 2KCl (s) + 3O2 (g)
N2 + O2 3000°C → 2NO (Nitric oxide)
The reactions which produce heat are called exothermic reactions.
The compounds formed by exothermic reactions are called exothermic
compounds. For example,

C (s) + O2 (g) → CO2 (g) + heat
CaO (s) + H2O → Ca (OH)2 (aq) + heat
CH4 (g) + 2O2 (g) → CO2 (g) +2H2O (g) + heat

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Reversible and irreversible reactions

The reaction which occurs in forward as well as backward direction
is called reversible reaction. For example,

N2 (g) + H2 (g) 4500C,200–900atm 2NH3 (g)
Fe, MO

In the above equation, the forward reaction means the combination

of nitrogen and hydrogen to give ammonia. But the decomposition of

ammonia back to nitrogen and hydrogen is the backward reaction.

A reaction which occurs in forward direction only is called irreversible
reaction. For example,

2Na + 2H2O →2NaOH + H2↑
C (s) + O2 (g) → CO2 (g)

Factors affecting chemical reaction

There are many factors which affect the rate of chemical reaction.
Some of the important factors are given below:

1. Simple contact: Some substances react with each other when
brought in contact with each other. For example: when sodium
is added to the jar containing chlorine, sodium chloride (a salt) is
formed. Similarly, when hydrogen gas is mixed with fluorine gas,
hydrogen fluoride is obtained.

2Na + Cl2 → 2NaCl

H2 + F2 →2HF

2. Contact by solution: Many reactions occur in the form of their
aqueous solution. For example: Sodium chloride (NaCl) and
silver nitrate (AgNO3) do not react in their solid state. But when
aqueous solutions of these substances are mixed, they react to
from precipitate of silver chloride (AgCl).

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

3. Heat: Heat increases the rate of chemical reaction. Heat gives

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energy to the molecules of the reactants and makes them collide
with each other. The collision between the molecules brings
about chemical reaction. For example: When calcium carbonate
is heated, calcium oxide and carbon dioxide are formed.

CaCO3 ∆ → CaO + CO2

4. Light: Some of the reactions take place in the presence of
light. For example: When silver bromide is exposed to sunlight,
it decomposes to from silver and liquid bromine. Similarly,
hydrogen and chlorine combine to from HCl in the presence of
sunlight.

2AgBr sunlight → 2Ag + Br2

H2 + Cl2 sunlight → 2HCl

5. Pressure: High pressure is a necessary condition for many
chemical reactions. For example: Nitrogen reacts with Hydrogen
under high pressure to give ammonia.

N2 + 3H2 4500C,200–900atm 2NH3
Fe, MO

6. Electricity: Electricity brings about both combination and
decomposition reaction. When electricity is passed through

acidified water, the water decomposes to from hydrogen and
oxygen gas.

2H2O H2SO4 2H2 + O2
electricity

Similarly, when electric spark is passed through a mixture of
hydrogen and oxygen, they unite to form water.

2H2 + O2 electric−spark → 2H2O

7. Catalyst: A catalyst is a chemical substance which increases or
decreases the rate of chemical reaction without being consumed
itself. The process of changing the rate of chemical reaction by
the use of catalyst is called catalysis. Many chemical reactions
take place in the presence of catalyst. The catalysts are classified

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into two types:

a. Positive catalyst

b. Negative catalyst

a. Positive Catalyst: The catalyst which increases the rate of
chemical reaction without being consumed itself is called positive
catalyst. For example, when potassium chlorate is heated
strongly,it decomposes to produce potassium chloride and oxygen.

2KClO3 360°C → 2KCl + 3O2
(No use of catalyst, so the reaction occurs at high temperature)

But when manganese dioxide (MnO2) is added, the reaction takes
place at lower temperature.

2KClO3 240°C 2KCl + 3O2
MnO2

(Use of Catalyst, so the reaction occurs at reduced temperature)

MnO2 also acts as positive catalysts in the decomposition of
hydrogen peroxide (H2O2).

2H2O2 MnO2 → 2H2O + O2↑

b. Negative catalysts: A catalyst which decreases the rate
of chemical reaction without being consumed itself is called
negative catalyst. For example, glycerine or phosphoric acid acts
as a negative catalyst in the decomposition of hydrogen peroxide.

2H2O2 glycerine 2H2O + O2↑
(slowed rate)

Types of chemical reactions

The chemical reactions can be classified into following types on the
basis of nature of the reactions.

1. Combination or addition reaction: The chemical reaction in
which two or more substances combine to form a new substance
is called combination reaction. For example:

H2 (g) + O2 (g) → H2O (l)

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2Na (s) + Cl2 (g) →2NaCl (s)

N2 (g) + 3H2 (g) 4500C, 200-900 atm 2NH3 (g) + heat
Fe+MO
Fe(s) + S (s)→ FeS(s)

2. Decomposition or dissociation reaction: The chemical
reaction in which a compound decomposes into two or more
simpler substances is called decomposition reaction. It is just
the opposite of a combination reaction. A decomposition reaction
occurs due to heat, light, electricity, catalyst, etc. Examples:

a. By heat: Heat decomposes oxides, carbonates, chlorates,
nitrates, etc of metals. Sometimes, heat as well as catalyst are
needed.

2PbO2 (s) ∆ → 2PbO (s) + O2 (g)

CaCO3(s) ∆ → CaO (s) + CO2 (g)

2KClO3(s) 240°C 2KCl (s) + 3O2 (g)
MnO2

2Pb(NO3)2 ∆ → 2PbO + 4NO2 + O2 ↑

b. By electricity: Electricity decomposes liquid or molten states

of the compounds.

2H2O (l) Electricity 2H2 (g) + O2 (g)
acid

2NaCl (molten) Electricity → 2Na (s) + Cl2 (g)

c. By light: Silver bromide is decomposed to silver and bromine
by light.

2AgBr (s) light → 2Ag (s) + Br2 (l)
This reaction is useful in photography.

d. By Catalyst: Manganese dioxide acts as a positive catalyst in
the decomposition of hydrogen peroxide.

2H2O2 (aq) MnO2 → 2H2O (l) + O2 (g)

3. Displacement or replacement reaction: The chemical
reaction in which an element or radical of a compound is

Times' Crucial Science Book - 10 164

replaced by another element or radical to form new compounds
is called displacement reaction. Generally, a more reactive
element displaces a less reactive element from its compound.
Displacement reaction can be divided into two types:

a. Single displacement reaction

b. Double displacement reaction

a. Single displacement reaction: The chemical reaction
in which an element from a compound is replaced by another
element to form a new compound is called single displacement
reaction. For example,

Zn(s) + H2SO4 (aq) → ZnSO4 (aq) + H2 (g)
Here, zinc displaces hydrogen from sulphuric acid.

b. Double displacement reaction: A chemical reaction
in which two reacting compounds exchange corresponding
elements or radicals to form two new compounds is called
double displacement reaction. Such reactions are shown by ionic
compounds in solution state. Examples:

AgNO3 (aq) + NaCl (aq) → NaNO3 (aq) + AgCl ↓
2AgNO3 (aq) + CaCl2 (aq) → Ca(NO3)2 (aq) + 2AgCl ↓

4. Acid-base or Neutralization reaction: A chemical reaction in
which an acid reacts with a base to form salt and water is called
acid base reaction. This reaction is also called neutralization
reaction because acid and base neutralize each other to form salt.
A salt in general is a neutral compound. Examples:

HCl (aq) + NaOH (aq) → NaCl (aq) + H2O (l)

acid base salt water

H2SO4 (aq) + 2NaOH (aq) → Na2SO4 (aq) + 2H2O (l)
HNO3 (aq) + KOH (aq) → KNO3 (aq) + H2O (l)

Rate of chemical reaction

The decrease in the concentration of reactants or the increase in
concentration of products per unit time is called rate of chemical
reaction or simply reaction rate. It shows the speed at which

165 Times' Crucial Science Book - 10

reactants are transformed into products. The rate of reaction provides
information whether a particular reaction is slow or fast. The rate of
reaction is affected four main factors, which are discussed below:

1. Temperature : In general, the increase in temperature increases
the rate of chemical reaction. The increased temperature
vibrates or moves the molecules of the reactants. This makes
the reacting molecules come close to each other rapidly. If the
molecules come closer, they collide with each other quickly and
the reaction becomes faster.

2. Pressure : Pressure mainly affects the chemical reaction of
gaseous reactants. The increased pressure improves the rate
of reaction of gaseous reactants. When gaseous reactants are
mixed for reaction, their molecules are far apart from each other.
If pressure is increased in the vessel containing gas molecules,
their volume decreases. This makes molecules come closer and
collide with each other rapidly. The effective collision of such
molecules results in products and the reaction becomes faster.

3. Catalyst : A catalyst has a special property to increase or decrease
the rate of chemical reaction without being consumed itself. Its
concentration and composition remains unchanged at the end of
the reaction. The use of a positive catalyst increases the rate of
reaction. A catalyst lowers the energy required to break the bond
between the molecules of reactants. Thus, even the lower energy
causes the reaction and it occurs at faster rate.

4. Surface area : The greater surface area of reactants increases the
rate of reaction. The surface area increases if the solid reactants
are crushed into their powder. In powder form, more molecules are
exposed to the other reacting molecules. The exposed molecules
collide with each other and the reaction becomes faster. Further,
the reaction rate is enhanced if is it possible to use reactants in
their solution phase.

Times' Crucial Science Book - 10 166

Learn and Write

1. Heat increases the rate of chemical reactions. Why?

Heat gives energy to the molecules of the reactants and makes
them collide with each other. The collision between the molecules
brings the chemical change faster.

2. Acid-base reaction is called neutralization reaction. Why?

As a result of acid-base reaction, salt and water are produced.
Generally salts are neutral. Therefore, acid-base reaction is
called neutralization reaction.

3. What happens when magnesium ribbon is burnt in air?

When magnesium ribbon is burnt in air, it gives magnesium oxide.
2Mg + O2 → 2MgO (magnesium oxide).

Main points to remember

1. A chemical reaction is a chemical change in which exchange,
combination or decomposition occurs in the molecule of substances
to produce new substance.

2. A catalyst is a chemical substance which alters the rate of
chemical reaction but itself remains unchanged at the end of the
reaction. The process of changing the rate of chemical reaction by
the use of catalyst is called catalysis.

3. There are different types of chemical reactions. They are:
combination reaction, decomposition reaction, displacement
reaction, neutralization reaction etc.

4. Areactioninwhichtwoormoresubstancescombinetogethertoform
a new substance is called combination reaction.

5. Areactioninwhichacompounddecomposesintotwoormoresimpler
substance is called decomposition reaction.

6. A chemical reaction in which an element or radical in a
compound is displaced by another element or radical to form
new compounds is called displacement reaction.

7. A chemical reaction in which an acid reacts with a base to form
salt and water is called neutralization reaction.

167 Times' Crucial Science Book - 10

Exercise

A. Choose the best alternative.

1. The substance that undergoes chemical change during a
chemical reaction is called

a. Chemical equation b. Product c. Reactant d. Reaction

2. If the number of atoms of each element is same on both sides
of the arrow, the equation is called

a. Balanced chemical equation b. Skeleton equation
c. Unbalanced equation d. Word equation

3. A simple chemical equation cannot tell

a. Concentration of reactants and products

b. Physical state of reactants and products

c. Speed of the reaction

d. All of the above

4. In the thermal decomposition of potassium chlorate, MnO2 acts as a
a. Catalyst b. Negative catalyst c. Positive catalyst d. All

5. When ammonium cyanate is heated, it turns into urea. Such
reaction is called

a. Decomposition reaction b. Single displacement reaction

c. Hydrolysis reaction d. Rerarrangement reaction

B. Answer these questions in brief.
1. What do you mean by chemical reaction?
2. What is a balanced chemical equation? Give an example.
3. What conditions are required for a chemical change? Explain
any two with examples.
4. What is a catalyst? Give two examples each of positive and
negative catalyst.
5. Write down any three informations that can be obtained from
a balanced chemical equation.
6. Give an example of chemical reaction that takes place in the
presence of light.
7. Give an example of reaction that takes place in the presence of heat.
8. Write down any three limitations of a chemical reaction.
9. Define exothermic and endothermic reaction with one
example of each.
10. Give an example of electrolytic decomposition reaction.

Times' Crucial Science Book - 10 168

11. Give two examples of reactions that are made either faster or
slower by catalyst.

12. Explain the effect of temperature on the rate of reaction.
13. Write down any three methods of increasing the rate of

chemical reaction.
14. Define the following with one example of each:

a. Combination reaction b. Decomposition reaction
c. Neutralization reaction d. Single displacement reaction
e. Double displacement reaction f. Rearrangement reaction

C. Differentiate between:
1. Combination and decomposition reaction.
2. Physical change and chemical change.

D. Give reasons.
1. It is necessary to balance a chemical equation.
2. The acid base reaction is also known as neutralization reaction.
3. The increase in concentration of reactants increases the rate
of reaction.

E. What happens when (Write with balanced chemical equation):
1. Potassium chlorate is heated?
2. Lead nitrate is heated?
3. Calcium reacts with water?
4. Phosphorus is burnt in air?
5. Mercuric oxide is heated?
6. Potassium bromide and chlorine gas are mixed together?
7. Sodium metal comes in contact with water?
8. Lime water is added to dilute solution of hydrochloric acid?
9. Silver bromide is exposed to sunlight?

F. Write down the balanced chemical equations for the following reactions:
1. Aluminium + Oxygen → Aluminium Oxide.

2. Calcium bicarbonate ∆ → Calcium Carbonate + Water + Carbon
dioxide.

3. Copper Carbonate ∆ → Copper oxide + Carbon dioxide.
4. Calcium Hydroxide + Carbon dioxide → Calcium carbonate +

Water.
5. Nitric acid + Sodium hydroxide → Sodium nitrate + Water.
6. Sodium carbonate + Hydrochloric acid → Sodium chloride + Water

+ Carbon dioxide.
7. Ammonium chloride + Sodium nitrite → Sodium chloride + Water +

Nitrogen.

169 Times' Crucial Science Book - 10

8. Calciumchloride+Silvernitrate→Calciumnitrate+Silverchloride.

G. Solve the following numerical problems.
1. How many helium atoms are present in 0.2 kg of helium?
2. How many water molecules are present in 54g of water?
3. Calculate the mass of MgO formed if 10gm of magnesium is
burnt in air.
4. Calculate the weight (in gram) of 2 mole of methane gas.
5. If 10gm of KOH reacts with H2SO4 to from K2SO4, find the
mass of K2SO4 formed.
6. Calculate the mass of CaO formed if 0.5 kg of CaCO3 is
heated.

Project Work

Take some dilute sulphuric acid in a test tube. Add a granule of
zinc into it. What will you observe? Cover the mouth of the test
tube with the mouth of another test tube and hold them together
in vertical position for about 2 minutes. Introduce a burning
candle to the mouth of the upper test tube. What will you observe?
Why is it so? Explain your observations and findings.

Glossary : use of symbol and formula

• Symbolic : substances that undergo chemical change during
• Reactant the chemical reaction

• Product : the new substances which are formed as a result
of chemical change in the reactants
• Coefficient
: the number placed before a letter that represents
• Catalyst a variable in algebra

• Precipitate : a substance that increases the rate of a chemical
reaction without itself undergoing any change

: to cause a solid to separate out from a solution as
a result of a chemical reaction, or separate out in
this way

Times' Crucial Science Book - 10 170

Chapter

9 Acid, Base

and Salt
Soren Sorensen
The potential of Hydrogen, or pH, scale was invented by the Danish biochemist Soren
Peter Lauritz Sorensen. He introduced this scale in 1909 to measure the acidity and
basicity oEf susbsttainmces ated Periods: 8 (6T+2P)

ObAjetctthieveesnd of the lesson, students will be able to:

• define acid, base and salt;

• explain properties of acid, base and salt;

• differentiate between base and alkali;

• explain uses of acid, base and salt;

• write and balance simple acid-base reactions.

There are various types of compounds in our surrounding. They are
categorised into three categories. They are: acid, base and salt.
These compounds are obtained from natural as well as man-made
sources. For example, the acids are obtained from fruits and vegetables
such as lemon, orange, mango, amala, etc. Acetic acid (vinegar) is the
first known acid to human and is used to make pickles. Ascorbic acid
(vitamin C), found in amala is used as medicine. Similarly, caustic
soda (sodium hydroxide) is a base which is used to manufacture
soaps. Common salt and copper sulphate are salts which are very
useful. Common salt is our kitchen salt whereas copper sulphate is
used as fungicide, electrolyte, etc.

Acids

The word acid has been derived from the Latin word acidus which

means sour in taste. Initially, the sour tasting substances were

regarded as acid. But this definition could not include all the acids.

Hence, an acid is defined as a substance which produces hydrogen

ions (H+) when dissolved in water. For example, hydrochloric acid,

sulphuric acid, nitric acid, carbonic acid, acetic acid, etc.
HCl →H+ + Cl–

H2SO4 → 2H+ + SO4– – H+ + CH3COO–
HNO3 → H+ + NO3–
CH3COOH

171 Times' Crucial Science Book - 10

In fact, hydrogen ions do not exist in solution. They immediately

combine with H2O to form H3O+. The hydrated hydrogen ion that
exists in the solution of acids is called a hydronium ion (H3O+).

H+ + H2O → H3O+

Hydrogen ion Water molecule Hydronium ion

Hence, acids are defined as compounds which produce hydronium

ions (H3O+) when dissolved in water.
Strong and weak acids

An acid which ionizes completely to give high concentration of hydrogen

ions (H+) in aqueous solution is called strong acid. The aqueous

solution of these acids have low pH value. Generally, inorganic acids

are strong acids. For example, hydrochloric acid (HCl), sulphuric acid

(H2SO4), nitric acid (HNO3), hydroiodic acid (HI) etc.
An acid which gives low concentration of hydrogen ion (H+) in

aqueous solution is called weak acid. These acids ionize partially and

are comparatively poor conductor of electricity. For example, acetic

acid, (CH3COOH), formic acid (HCOOH), carbonic acid (H2CO3), etc.
Organic and inorganic acid

The acids which are obtained from living organisms (plants and

animals) are called organic acids. For example: acetic acid, formic

acid, tartaric acid, lactic acid, citric acid, malic acid, ascorbic acid, etc.

The acids which are obtained from minerals and prepared in

laboratory are called inorganic acids. For example: hydrochloric acid

(HCl), sulphuric acid (H2SO4), nitric acid (HNO3), etc.

Activity 9 .1 To identify the weak and strong acid.

Materials required: Test tubes

Chemicals required:

Pieces of magnesium ribbon, dilute sulphuric acid, acetic acid, etc.

Procedure
1. Fill about one-third of a test tube with dilute hydrochloric acid.
2. Fill about one-third of another test tube with acetic acid.
3. Drop a small piece of magnesium ribbon in both the test tubes.

Observe what happens?

Observation:

You will observe that sulphuric acid reacts with magnesium
faster than acetic acid.

Times' Crucial Science Book - 10 172

Conclusion:

Sulphuric acid is a strong acid and acetic acid is a weak acid. Repeat
the above mentioned activity taking the pieces of granulated zinc
instead of magnesium ribbon. Observe the reactions and note
which one is faster. What will you conclude from the observation?

Differences between organic and inorganic acid:

Organic acid Inorganic acid

1. It is obtained from the 1. It is obtained from minerals and
plant or animal source. can be prepared in laboratory.

2. They are weak acids. 2. They are generally strong acids.

Examples: acetic acid, lactic Examples: hydrochloric acid,
acid, formic acid etc. sulphuric acid, nitric acid, etc.

Sources of organic acid
The organic acids are found naturally in fruits and vegetables. Some
of them also are present in the body of insects, other animals, etc. All
organic acids are sour in taste.
Some of the naturally occurring organic acids and their sources are
given below:

S. N Name of Acid Source

1. Acetic acid Vinegar, sour pickles etc.

2. Citric acid Lemon, orange, mandarin, tomato etc.

3. Lactic acid Milk, curd etc.

4. Formic acid Ant bite, nettle sting etc.

5. Oxalic acid Oxalis (Chariamilo)

6. Tartaric acid Grapes, pomelo etc.

7. Ascorbic acid (Vit. C) Amala, Harro, etc.

Properties of acids
1. Acids are generally sour in taste. But the acids such as boric
acid, salicylic acid, etc do not have sour taste.
2. Aqueous solutions of acids conduct electricity.
3. Acids change the colour of indicator as follows:

Indicators Change of colour

Blue litmus paper Red

Methyl orange (Orange) Red

Phenolphthalein (Colourless) Colourless (no change)

173 Times' Crucial Science Book - 10

4. Acids give hydrogen ions when dissolved in water.
HCl → H+ + Cl–
HNO3 → H+ + NO3–
H2SO4 → 2H+ + SO4– –

5. Acids react with metals to form salt and hydrogen gas.

Acid + metal → Salt + hydrogen gas

2HCl + Zn → ZnCl2 + H2↑
H2SO4 + Mg → MgSO4 + H2↑
6. Acids react with bases (metallic oxides and hydroxides) to
form salts and water.

Acid + Base → Salt + Water

HCl + NaOH → NaCl + H2O
2HCl + CaO → CaCl2 + H2O
HNO3 + KOH → KNO3 + H2O
H2SO4 + 2NaOH → Na2SO4 + 2H2O
7. Acids react with carbonates and bicarbonates to form salt,
water and carbon dioxide gas.

Acid + carbonate or bicarbonate → Salt + water + carbon dioxide

2HCl + Na2CO3 → 2NaCl + H2O + CO2
H2SO4 + MgCO3 → MgSO4 + H2O + CO2
2HCl + Ca(HCO3)2 → CaCl2 + 2H2O + 2CO2

Uses of acids

1. Sulphuric acid and nitric acid are used to make chemical
fertilizers, drugs, explosives, etc.

2. Carbonic acid is used to make soft drinks like soda water and
other drinks.

3. The acids like HCl, H2SO4, HNO3 etc are used in laboratory to
perform different experiments.

4. Carbolic acid or Phenol (C6H5OH) is used as germicide or
disinfectant.

5. Citric acid is a source of vitamin C.

6. Acetic acid (vinegar), tartaric acid, citric acid, etc are used for
giving taste to food.

7. Picric acid, boric acid, etc are used for washing wounds.

Times' Crucial Science Book - 10 174

Precautions to be taken while using acids

It is dangerous to taste or touch strong acids. Hence, following
precautions should be remembered while using acids:

a. We should never touch or taste the strong acids because they
can burn the skin.

b. We should never fill acid in metallic vessels.
c. If any part of our body comes in contact of acid, we should

wash the affected part thoroughly with water or with dilute
solution of sodium bicarbonate. But we should never use
strong alkali like NaOH or KOH because they are also

Base corrosive.

Bases are metallic oxides or hydroxides which react with acids to
give salt and water. For example, sodium oxide, potassium oxide,
sodium hydroxide, calcium hydroxide, iron oxide etc.

Base + Acid → Salt + Water

NaOH + HCl → NaCl + H2O
CaO + H2SO4 → CaSO4 + H2O

The bases which are soluble in water are called alkalis. The alkalis
produce hydroxyl ions (OH–) when dissolved in water. For example,
sodium hydroxide, potassium hydroxide, ammonium hydroxide,
calcium hydroxide, etc.

KOH → K+ + OH¯

Hydroxyl ion
NaOH → Na+ + OH ¯

Activity 9 .2 To prepare a base using magnesium (metal).

Materials required:
Beaker, a burning candle, a glass rod, funnel, filter paper, stand, etc.

Chemicals required:
A piece of magnesium ribbon, water, phenolphthalein (indicator) etc.

Procedure
1. Hold a small piece of magnesium ribbon with the help of tongs.
2. Burn the magnesium ribbon with the help of burning

candle and hold it above a white sheet of paper. Burning of
magnesium in the air produces a white powder of magnesium
oxide which is a base.

175 Times' Crucial Science Book - 10

3. Now, take a beaker with little water and add the white
powder into it. Stir the mixture to dissolve completely. Add a
little more water if needed.

4. Filter the solution and take a small quantity of solution in a
test tube.

5. Add a drop of phenolphthalein into the test tube and
observe the change in colour.

Observation:

The solution of the test tube turns pink.

Conclusion:

Magnesium burns in the oxygen of air to form magnesium oxide.
This oxide dissolves in water to form an alkali, magnesium
hydroxide. An alkali (base) turns the colour of phenolphthalein
into pink.

Reactions involved:

Magnesium + Oxygen ∆ → Magnesium oxide

2Mg + O2 ∆ → 2MgO (a base)
Magnesium oxide + water → Magnesium hydroxide

MgO + H2O → Mg(OH)2 (an alkali)

Preparation of bases

The bases can be prepared in the following ways:

1. Direct combination of metal with oxygen: Sodium directly
combines with the oxygen of atmosphere to form sodium oxide.
But other metals combine with oxygen while heating.

4Na + O2 → 2Na2O

4Al + 3O2 ∆ → 2Al2O3

2Ca + O2 ∆ → 2CaO

2Cu + O2 → 2CuO

2. Action of water on soluble metal oxides: The oxides of very
active metals such as sodium, potassium, calcium, magnesium,
etc react with water to produce alkali.

Times' Crucial Science Book - 10 176

Na2O + H2O →2NaOH
CaO + H2O → Ca(OH)2
K2O + H2O → 2KOH
3. Direct reaction of metal with water: Sodium, potassium and
calcium directly react with water to produce hydrogen gas and
alkali.

2Na + 2H2O → 2NaOH + H2↑
Ca + 2H2O → Ca(OH)2 + H2↑
4. Decomposition of metallic carbonates on heating: When
calcium carbonate is heated, it decomposes to calcium oxide and
carbon dioxide.

CaCO3 ∆ → CaO + CO2

All oxides and hydroxides of metal are bases. Some bases are soluble
in water while most of them are insoluble. The bases like KOH,
Mg(OH)2, NaOH, Ca(OH)2 etc are water soluble bases and are called
alkalis. Cupric oxide, aluminium oxide, ferric oxide, silver oxide
etc. are also the bases because they are the oxides of metal. But
these oxides are insoluble in water and hence they are not alkalis.
Therefore, all alkalis are bases but all bases are not alkalis.

Strong and weak alkalis (bases)

The alkalis which completely dissociate in water to produce high
concentration of hydroxyl (OH–) ions are called strong alkalis. They
have high pH value. For example, NaOH, KOH etc.

The alkalis which dissociate partially in water to produce less
concentration of hydroxyl (OH–) ion are called weak alkalis. For
example, Cu(OH)2, Fe(OH)3, Al(OH)3, etc.
Differences between bases and alkalis:

Bases Alkalis

1. All oxides and hydroxides of 1. Only hydroxides of very
metal are called bases. active metals are alkalis.

2. They may be soluble or 2. They are soluble in water.
insoluble in water.

3. All bases are not alkalis. 3. All alkalis are bases.

Examples: CuO, NaOH, CaO, Examples: NaOH, KOH etc.
KOH, HgO, Ag2O, Fe(OH)2 etc.

177 Times' Crucial Science Book - 10

Properties of bases
1. The bases are soapy (slippery) to touch.
2. They have bitter taste.
3. They turn the colour of indicators as follows:

Indicators Change of colour

Red litmus (red) Blue

Methyl orange (Orange) Yellow

Phenolphthalein (Colourless) Pink

4. Water soluble bases produce hydroxyl ion when dissolved in

water. NaOH → Na+ + OH -

Ca(OH)2 H2O Ca+ + + 2OH -

5. They react with acids to form salt and water. The reaction

of base with acid to form neutral salt and water is called

neutralization reaction.

Base + Acid → Salt + Water

Examples:

NaOH + HCl → NaCl + H2O
Ca(OH)2 + 2HCl → CaCl2 + 2H2O

6. Alkali solution reacts with carbon dioxide to form metallic
carbonate and water.

Examples:

2NaOH + CO2 → Na2CO3 + H2O
2KOH + CO2 → K2CO3 + H2O
Mg(OH)2 + CO2 → MgCO3 + H2O

7. They react with some salt to form insoluble metallic hydroxide.

Examples:

2NaOH + CuSO4 → Cu(OH)2 + Na2SO4
FeCl3 + 3NH4OH → Fe(OH)3 + 3NH4Cl

8. They react with ammonium salt on heating and give ammonia gas.

Examples:

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

Times' Crucial Science Book - 10 178

Uses of bases
1. Sodium hydroxide is used in the manufacture of soap, paper,
rayon etc. It is also used in refining of petroleum products.
2. Quick lime (CaO) is used for softening hard water, production
or purification of sugar, manufacture of washing soda etc.
3. Calcium hydroxide is used for white washing of buildings,
softening hard water etc. Calcium hydroxide is also used to
manufacture mortar and bleaching powder.
4. Ammonium hydroxide is used for the manufacture of
fertilizers, plastics, dyes etc.
5. Potassium hydroxide is used in the manufacture of soft soaps.
6. Magnesium hydroxide and aluminium hydroxide are is used
as antacid to control acidity in stomach.

7. Sodium hydroxide, potassium hydroxide and ammonium
hydroxide are very important lab reagents.

Difference between acid and base

Acid Base

1. It gives hydrogen ion (H+) 1. Water soluble base gives
when dissolved in water. hydroxyl ion (OH–)ion when
dissolved in water.

2. It turns blue litmus in to red. 2. It turns red litmus in to blue.

3. It does not change the colour 3. It changes the colourless

of Phenolphthalein. solution of Phenolphthalein

to pink.

4. It reacts with a base to form 4. It reacts with an acid to

salt and water. form salt and water.

5. It has sour taste. 5. It has bitter taste.

6. Its pH is less than 7. 6. Its pH is more than 7.

Salt

A salt is formed by the neutralization of an acid by a base. The salts
are generally neutral to indicators. A salt is a compound which
is formed by the partial or complete replacement of one or more
hydrogen atoms of an acid by one or more metal atoms or positive
radicals.

For example, sodium chloride is a salt which is formed by the
replacement of one hydrogen atom of an acid by one sodium atom.

NaOH + HCl → NaCl + H2O

179 Times' Crucial Science Book - 10

If sodium replaces one hydrogen atom from sulphuric acid (partial
replacement) sodium bisulphate is formed. But the replacement of
both hydrogen atoms from sulphuric acid by sodium produces sodium
sulphate (complete replacement).

NaOH + H2SO4 → NaHSO4 + H2O (partial replacement)

Sodium bisulPHate

2NaOH + H2SO4 → Na2SO4 + H2O (complete replacement)
Sodium sulphate

Each salt molecule contains two radicals: basic radical (electropositive
radical) and acidic radical (electronegative radical). The radical that
comes from a base and carries positive charge is called basic radical
whereas the radical that comes from an acid and carries negative
charge is called acidic radical. Therefore, the basic radical is known
as metallic radical and the acidic radical is known as non-metallic
radical. For example, in the salt NaCl, Na+ is basic radical and Cl– is
acidic radical.

Preparation of salts

Salts can be prepared by the following methods:

1. Direct combination of metals with non-metals:

Some active metals such as sodium, calcium, iron, etc react
directly with non-metals to form salts.

2Na + Cl2 → 2NaCl (Sodium chloride)
2Fe + 3Cl2 → 2FeCl3 (Ferric chloride)
2. Action of acids on metals:

Metals react with acids to produce salt and hydrogen gas.

Zn + H2SO4 → ZnSO4 + H2↑
Mg + 2HCl → MgCl2 + H2↑
Cu + H2SO4 → CuSO4 + SO2 + H2O
3. Action of acids on metallic oxide: Metal oxides react with
acid to form salt and water.

Na2O + 2HCl → 2NaCl + H2O
CuO + H2SO4 → CuSO4 + H2O
4. Neutralization of an acid by an alkali (hydroxide base):
The reaction of an alkali with an acid produces salt and water.

NaOH + HCl → NaCl + H2O

Strong alkali strong acid Neutral salt

Times' Crucial Science Book - 10 180

H2SO4 + Zn(OH)2 → ZnSO4 + H2O
Acidic salt
Strong acid Weak alkali
Na2CO3 + 2H2O
H2CO3 + 2NaOH → Basic salt

Weak acid Strong alkali CH3COONH4 + H2O
Neutral salt
CH3COOH + NH4OH →
Weak acid (Ammonium ethanoate)
weak alkali

Salts formed by the reaction of strong acid and strong base
are generally neutral. If the salts are formed by the reaction
of strong acid and weak base, they are acidic. If the salts are
formed by the reaction of weak acid with strong base, they are
alkaline. Similarly, salts formed from weak acid and weak base
are neutral.

5. Action of acid on metallic carbonates: The action of acid on
metallic carbonates also produces salts.

CaCO3 + 2HCl → CaCl2 + H2O + CO2↑

Types of salts

There are three types of salts. They are:

(a) Normal salts: It is a salt formed by complete replacement
of hydrogen atoms of an acid by a metal. A normal salt is
generally neutral to indicators. For example

NaCl, KCl, BaSO4, Na2SO4, NH4Cl etc,

(b) Acidic salts: It is a salt formed by partial displacement of
hydrogen of an acid by a metal or electropositive radical.
Some of the acidic salts are acidic to indicators.

Example: NaHCO3, NaHSO4, Ca(HCO3)2, KHSO4 etc.

(c) Basic salts: The salt which is formed by the neutralization

reaction between weak acid and strong base is called basic

salt. Such salts are alkaline in nature. Examples: Na2CO3,
CH3COONa, CH3COOK, etc.

Strong base + Weak acid Basic salt + Water

NaOH + H2CO3 Na2CO3 + H2O

NaOH + CH3COOH CH3COONa + H2O

(d) Hydrated salts: The salts which contain water molecules in

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their crystals are called hydrated salts. The water molecules

present in the crystals of salts are called water of crystallization

or water of hydration. Each salt contains a specific number

of water of crystallization. Examples of hydrated salts: Blue

vitriol (copper II sulphate pentahydrate, CuSO4.5H2O), white
vitriol (zinc sulphate heptahydrate, ZnSO4.7H2O), Epsom salt
or magnesium sulphate heptahydrate (MgSO4.7H2O), sodium
sulphate decahydrate (Na2SO4.10H2O), etc.

Although CuSO4 is a normal salt formed by the complete replacement
of hydrogen atoms of sulphuric acid by copper, it shows acidic nature

in aqueous solution due to hydrolysis. Copper sulphate ionizes into
Cu++ and SO4– –
dissociates into in its aqueous solution. At the same time, water also
H+ and OH– ions. Thus H+ ions react with SO4– – ions

to form sulphuric acid, which is a strong acid. But Cu++ ions combine
with OH– ions to form a weak base, Cu(OH)2. As a result the solution
contains more concentration of H+ ions and becomes acidic.

CuSO4 + 2H2O → H2SO4 + Cu(OH)2

Similarly, the aqueous solution of Na2CO3 is alkaline due to hydrolysis
reaction.

Na2CO3 + H2O → H2CO3 + NaOH

Properties of salts

The salts have a variety of properties, some of which are given below:

(1) Sodium chloride has a characteristic salty taste. But most of
the salts taste bitter. Some salts are tasteless too.

(2) Salts are generally neutral to indicators (e.g. NaCl, KCl,
KNO3 etc.). Some salts are acidic (CuSO4) whereas some are
alkaline (basic), e.g. Na2CO3.

(3) Salts of metals like Na, K, Al, Mg, Ca, and Ba are colourless
(white) whereas the salts of the transition metals such as Fe,
Co, Ni, Cu, Mn, Cr, etc, are coloured.

(4) Most of the salts dissolve in water. All sodium, potassium and
ammonium salts, all bicarbonates and nitrates, all chlorides
(except those of silver and lead) and all sulphates (except
those of lead and barium) are soluble in water.

(5) The salts are electro valent or ionic compounds. They conduct
electricity in their fused or aqueous solution state due to
ionization.

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Uses of salts

The salts are very useful in our daily life. The uses of some of the
salts are as follows:

S Name of Mol. formula Common Uses
No. Salt name
NaCl As edible salt. In
1. Sodium Common the manufacture of
chloride salt sodium hydroxide,
washing soda, etc.
2. Sodium Na2CO3.10H2O Washing
In the manufacture
carbonate soda of glass, soap and
detergent. It is also
3. Sodium NaHCO3 Baking used to remove
soda hardness of water.
bicarbonate
It is used as
4. Copper CuSO4.5H2O Blue baking powder,
CaSO4.2H2O vitriol fire extinguisher,
sulphate to reduce acidity of
Gypsum stomach, etc.
5. Calcium
sulphate For electroplating,
preserving wood, as
6. Ferrous FeSO4.7H2O Green fungicide etc.
sulphate vitriol
White To manufacture
7. Zinc ZnSO4.7H2O vitriol chalk, in the
sulphate - manufacture of
cement, etc.
8. Ammonium (NH4)2SO4 Plaster of
sulphate paris Used as medicine, in
factories, etc.
9. Calcium CaSO4.21 H2O
sulphate Used in eye medicine,
for white pigment etc.
10. Magnesium MgSO4.7H2O Epsom
sulphate salt As a chemical
fertilizer.

For plastering
fractured bones of
body, for making
ceramics and statues,
etc.

Used in medicine, dye
industry etc.

183 Times' Crucial Science Book - 10

11. Ammonium NH4Cl Sal As a good electrolyte
chloride ammoniac in dry cell, as a lab
reagent, etc.
12. Zinc ZnCO3 Calamine
For making ointment
carbonate for curing skin
diseases.

Indicator

An indicator is a chemical compound which indicates the acidic, basic
or neutral nature of a substance by changing its colour. The common
indicators are litmus paper, methyl orange and phenolphthalein.
Litmus paper is the most commonly used indicator in a laboratory. It
is prepared from a plant called lichen.
The change of colour of common indicator in acidic, basic and salt
solutions is given below:

S. Indicator and Colour in Colour in Colour in
N. its colour acidic soln alkaline salt soln

soln

1. Blue litmus paper (blue) Red No change No change

2. Red litmus paper (red) No change Blue No change

3. Methyl orange (orange) Red Yellow Faint
orange

4. Phenolphthalein Colourless Pink (red) No change
(colourless) (No change)

5. Red cabbage juice (red) Red Green Faint
purple

Indicators such as litmus, methyl orange and Phenolphthalein
indicate whether a substance is acidic, alkaline or neutral but they
cannot measure the strength of acid and alkali.

Universal indicator

An indicator prepared by mixing several ordinary indicators of
different colours is called universal indicator. It shows different
colours in the solutions of acids or alkalis of different concentrations.
It indicates the acidic or basic character of a substance as well as the
strength of acid and base.
A red or deep red colour in the universal indicator indicates acidity
whereas blue or deep blue colour denotes alkalinity. The strength of
acid or base can be determined by matching the change in colour of
universal indicator with the colour in pH chart.

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Application of neutralization reaction

The treatment of insect bites and acidity of stomach involves the use
of neutralization reaction. The bite of an ant or stinging of nettle
injects an acid called formic acid into our skin. The acid causes pain
with itching and swelling in the affected part. The effect of this acid is
neutralized by using calamine lotion (ZnCO3) or sodium bicarbonate.
Zinc carbonate or sodium bicarbonate reacts with the acid and
nullifies its effects. Thus, the pain is relieved. Similarly, the excess
of hydrochloric acid secreted in stomach causes hyper-acidity. To
neutralize the effect of acid, mild alkali, such as sodium carbonate is
taken. A mixture of magnesium hydroxide and aluminium hydroxide
is also used to get relief from hyper-acidity. This mixture is commonly
called antacid.

pH of soil

Optimum pH of soil is necessary for the growth of crops, fruits and
other plants. Some plants grow well in acidic soil, some in basic
whereas some grow well in the neutral pH.

If the acidity of the soil is increased beyond limit, it can cause harms
in the growth of plants. The increased acidity of soil is removed by
adding necessary amount of calcium hydroxide in the soil.

Main points to remember

1. An acid is a substance which gives hydrogen ion when dissolved
in water.

2. The hydrated hydrogen ion that exists in the solution of acids is
called hydronium ion.

3. Bases are the metallic oxides or hydroxides which react with
acid to give salt and water.

4. The bases which are soluble in water are called alkalis.
5. A salt is a compound which is formed by partial or complete

replacement of one or more hydrogen atoms of an acid by
equivalent number of metal atoms or electropositive radical.
6. An indicator is a chemical compound which indicates the acidic,
basic or neutral nature of a substance by changing its colour.
7. An indicator prepared by mixing several organic indicators of
different colours is called universal indicator.
8. The pH of a solution is defined as the measure of hydrogen ions
present in the given solution.

185 Times' Crucial Science Book - 10

Exercise

A. Choose the best alternative.

1. Which of the following is an inorganic acid?

a. HCl b. H2SO4 c. HNO3 d. All of these

2. What is the chemical name of vinegar?

a. Acetic acid b. Ascorbic acid

c. Vitamin C d. Calcium hydroxide

3. Which of the following is an organic acid?

a. Acetic acid b. Lactic acid

c. Citric acid d. All of these

4. Which of the following is a base but not alkali?

a. KOH b. NaOH c. Fe(OH)3 d. Ca(OH)2

5. Quick lime

a.CaO b. Na2O c. Ca(OH)2 d. KOH

B. Answer these questions in brief.

1. What is an acid? Name the ions present in aqueous nitric acid.

2. Why is H2SO4 called an acid?

3. What happens when copper is kept in hot concentrated
sulphuric acid?

4. Write down any four chemical properties of acids with
balanced chemical equations.

5. What are alkalis? Name the ions present in aqueous sodium
hydroxide.

6. Name any four compounds which can neutralize acids.

7. Write down any four chemical properties of bases with
balanced chemical equations.

8. What is meant by organic acid? Give examples.

9. What is a salt? Give some examples of water-soluble salts.

10. Mention any four ways of preparing salts.

11. What is an acidic salt? Give examples.

12. What is the pH value of a neutral solution?

13. Define basic salts with examples.

Times' Crucial Science Book - 10 186

14. Write down any two uses of sodium bicarbonate.

15. Name the basic and acidic radical in the compound NaCl.

16. Why is Phenolphthalein not much appropriate to identify
whether the given solution is acid, base or salt?

17. What is the use of finding out the pH of soil?

18. What type of salt is formed by the reaction between strong
acid and weak alkali?

19. What is meant by an universal indicator?
C. Give reasons:

1. Acetic acid is a weak acid.
2. Sodium hydroxide is a strong base.
3. All alkalis are bases but all bases are not alkalis.
4. When copper sulphate is dissolved in water, the solution

becomes acidic.
5. We use sodium bicarbonate to relieve hyper-acidity.
6. Calamine lotion is used to kill the pain of ant bite or nettle sting.
7. It is dangerous to touch or taste strong acids as well as alkalis.
D. Write down differences between:
1. Strong and Weak acid,

2. Inorganic and Organic acid,

3. Acid and Base.

4. Base and Alkali.

5. Indicator and Universal indicator.

6. Basic radical and Acid radical.
E. What happens when

1. Calcium carbonate is treated with dilute hydrochloric acid?

2. Carbon dioxide gas is passed through NaOH solution?

3. Ammonium salt is treated with strong alkali?

4. Hydrochloric acid reacts with sodium hydroxide?
F. Complete the following equations and balance them if necessary.

1. Zn + ............ → ZnSO4 + H2

2. Ca(OH)2 + ………… → CaCO3 + ……………

3. Cu + Conc. H2SO4 → …………….. + ………….. + ……………

187 Times' Crucial Science Book - 10

4. H+ + H2O → ……………
5. Na2O + → …………. → NaOH
6. CH3COOH + NH4OH → …………. + …………….
7. H2CO3 + NaOH → ...................... + .....................
G. Complete the following table:

SN. Indicator and Colour Colour Colour in
its colour in HCl in NaOH NaCl
solution solution
solution

1. Red litmus

2. Blue litmus

3. Methyl orange

4. Phenolphthalein

H. Conceptual questions:
1. Write down the name and molecular formula of the compound
which gives hydrogen and chloride ions when dissolved in
water. If few drops of methyl orange are added to this solution,
what will be its colour? Why? What will happen if the above
mentioned compound is treated with calcium oxide? Write
with balanced chemical equation.
2. Three different solutions such as caustic potash, hydrochloric
acid and salt solution are placed in three different beakers. How
would you identify each solution with the help of methyl orange?
3. The production of gases like CO2, SO2 etc. from industries can
affect our homes and historical monuments. Explain with reasons.
4. Write down the name and molecular formula of the compound
which releases sodium and hydroxyl ions in water. What will
be the colour of Phenolphthalein in this solution?

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Project Work

Take solutions of acid, alkali and salt in three different beakers.
Study the change in colour of the indicators in these solutions by
taking a little of solution in different test tubes and adding the
indicators.

Glossary : corrosive or burning by chemical action
: quantity of the solute present in the solution
• Caustic : to a degree but not completely
• Concentration : in the form of small grains or particles
• Partially : a balm, cream, gel, etc that is used as a medicine
• Granulated : a drug that reduces or neutralizes stomach acid
• Ointment
• Antacid

189 Times' Crucial Science Book - 10

Chapter

10 Some Gases
Henry Cavendish
He is known for the discovery of hydrogen and
measuring the earth's density, etc.
Estimated Periods: 7 (5T+2P)

Objectives

At the end of the lesson, students will be able to:
• explain the lab preparation of carbon dioxide and ammonia gas;
• explain the properties and uses of carbon dioxide and ammonia gas.

Carbon dioxide

Molecular formula - CO2
Molecular weight - 44 amu
Nature of bonding: Covalent bonds
Volume in atmosphere: 0.03%
Freezing point: 78°C

8p 12p 8p

8n 12n 8n

Molecular structure of CO2
Carbon dioxide is a compound gas. It was first prepared by Van
Helmond in 1630 AD and its properties were studied in detail by
Lavoisier.

Occurrence

Carbon dioxide occurs free as well as in combined state in nature. In
free state, it is present in atmosphere in about 0.03% by volume. It
is released into the atmosphere through burning of fuels, respiration
by plants and animals, forest fire, etc. It is found in the form of
compounds such as carbonates, bicarbonate, etc. More concentration
of CO2 is found in caves, mines, deep well, etc because it is heavier
than air.

Times' Crucial Science Book - 10 190

General methods of preparation

1. Carbon dioxide is produced by burning hydrocarbons with excess of oxygen.

CH4 + 2O2 burning → CO2 + 2H2O + heat

2C2H6 + 7O2 burning → 4CO2 + 6H2O + heat
2. Carbon dioxide is formed when carbon burns in air with excess of oxygen.

C + O2 → CO2 + heat
3. Carbonates and bicarbonates of metals such as calcium,

magnesium, etc, react with dilute acid to produce carbon dioxide.

CaCO3 + 2HCl → CaCl2 + H2O + CO2
Na2CO3 + 2HCl → 2NaCl + H2O + CO2
Mg(HCO3)2 + 2HCl → MgCl2 + 2H2O + 2CO2
4. Carbon dioxide is produced in large amount by heating limestone.

CaCO3 → CaO + CO2

Laboratory preparation of carbon dioxide

Principle

Carbon dioxide gas is prepared in the laboratory by the action of
dilute hydrochloric acid on calcium carbonate.

CaCO3 + 2HCl → CaCl2 + H2O + CO2
In this reaction, dilute sulphuric acid cannot be used in place of dilute
hydrochloric acid because calcium sulphate formed after reaction
covers the remaining part of marble and prevents the further reaction.

Apparatus required

Round bottom flask, thistle funnel, delivery tube, gas jar, cork, match
stick, litmus paper etc.

Chemicals required
1. Calcium carbonate (CaCO3)
2. Dilute hydrochloric acid (HCl)

Procedure
1. Place some pieces of marble in round bottom flask.
2. Connect a thistle funnel and delivery tube with the round
bottom flask using a cork as shown in the figure.

191 Times' Crucial Science Book - 10

Thistle funnel

Delivery tube

Round bottom flask Gas jar

HCl CO2 gas

CaCO3

Laboratory preparation of carbon dioxide gas

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

4. The chemical reaction between hydrochloric acid and marble
pieces produces carbon dioxide gas.

5. This gas passes through delivery tube and is collected in the
gas jar by upward displacement of air.

Carbon dioxide gas is collected in the gas jar by upward displacement
of air because it is heavier than air. In this process, the cork should
be tightly fitted with the round bottom flask otherwise the gas can
escape and no gas is collected in the gas jar.

Precautions

1. The apparatus should be air tight.
2. The end of thistle funnel should be dipped in the solution

otherwise the CO2 gas can escape through the thistle funnel.
3. The end of delivery tube should not touch the solution within

Woulfe's bottle. If the delivery tube is dipped in the solution,
no gas can travel through the delivery tube. So, the collection
of excess gas in the round bottom flask can cause explosion.

Test of carbon dioxide gas

1. When a moist blue litmus paper is inserted into the jar
containing the gas, the blue litmus paper turns into red. It
is because CO2 combines with H2O of moist litmus paper to
form an acid, which turns blue litmus into red.

2. When a burning match stick is brought near the mouth of
gas jar, the burning match stick extinguishes. It proves that
the gas jar contains carbon dioxide gas.

3. When carbon dioxide is passed through lime water (calcium

Times' Crucial Science Book - 10 192

hydroxide), the lime water turns milky white due to the
formation of insoluble calcium carbonate (CaCO3). This
proves that the gas is CO2.
4. When a burning magnesium ribbon is lowered into a jar
containing CO2 gas, the magnesium-ribbon continues to burn
forming magnesium oxide (grey dust) and carbon (black).

Manufacture of carbon dioxides gas

Carbon dioxide gas is manufactured in industrial scale by heating
calcium carbonate (limestone or marble) in a furnace. Calcium oxide
(quick lime) is another useful substance produced in this reaction.

CaCO3 ∆ → CaO + CO2 ↑

(limestone) (quicklime)

When calcium oxide reacts with water, it produces calcium hydroxide,
which is also known as slaked lime or lime water.

CaO + H2O → Ca (OH)2
(Slaked lime or lime water)

Properties of carbon dioxide

Physical properties
1. It is colourless, odourless and tasteless gas.
2. It dissolves sparingly in water.
3. It turns moist blue litmus paper into red.
4. It is heavier than air.
5. Carbon dioxide can be liquefied to a colourless liquid at 0°C
under the pressure of 40 atmospheres.
6. Carbon dioxide can be solidified by cooling it below -78°C or
applying pressure of 70 atmospheres at room temperature.
The solid form of CO2 is known as dry ice. It is called dry
ice because it melts without wetting the other articles like
paper, clothes etc. Dry ice is used in refrigerators to preserve
foods.
7. It is neither combustible nor the supporter of combustion.

Chemical properties

1. When carbon dioxide is dissolved in water, it produces
carbonic acid. lt is a weak acid.

193 Times' Crucial Science Book - 10

CO2 + H2O → H2CO3
Carbonic acid

2. When carbon dioxide is passed through lime water for a short
time, the lime water turns milky. It is due to the formation of
insoluble calcium carbonate.

CO2 + Ca (OH)2 → CaCO3 + H2O
Calcium carbonate (milky white)

If carbon dioxide is passed through the mixture for a long time, the
milky white colour disappears due to the formation of soluble calcium
bicarbonate.

CaCO3 + H2O + CO2 → Ca(HCO3)2
Calcium bicarbonate

3. Carbon dioxide is neither combustible nor the supporter of
combustion. But a burning magnesium ribbon continues to
burn in the atmosphere of carbon dioxide gas with dazzling
light. This reaction produces white powder of magnesium
oxide and black particles of carbon.

2Mg + CO2 → 2MgO + C

Magnesium oxide Carbon black

4. Ammonia reacts with carbon dioxide at about 150°C under

high pressure to produce urea.

2NH3 + CO2 1500C NH2 — CO — NH2 + H2O
Pressure Urea

5. When carbon dioxide is heated with red hot coke at 900°C,
carbon monoxide is produced.

CO2 + C 900°C → 2CO

6. Green plants convert solar energy into chemical energy in
the form glucose and oxygen.

6CO2 + 6H2O Sunlight C6H12O6 + 6O2
Chlorophyll

Glucose

7. Carbon dioxide reacts with NaOH or KOH solution to produce
carbonate salt.

Times' Crucial Science Book - 10 194

2NaOH + CO2 → Na2CO3 + H2O

Sodium carbonate

Uses of carbon dioxide

1. It is used in the manufacture of soft drinks like soda water,
coca cola, beer etc,

2. Solid carbon dioxide (dry ice) is used as refrigerant to preserve
fruits, fish, meat etc.

3. It is used in the manufacture of fertilizers (like urea), washing
soda (sodium carbonate).

4. It is used in fire extinguishers.
5. It is used in the purification of sugarcane juice. The process

of purifying sugarcane juice using CO2 is called carbonation.
6. Green plants use carbon dioxide for making food during

photosynthesis.
7. A mixture of gases containing 10 – 15 per cent oxygen and

carbon dioxide is called carbogen. It is used in artificial
respiration of pneumonia patients.

Fire extinguisher

Fire extinguisher is a protective Knob
device which produces carbon

dioxide to extinguish fire in

emergency conditions. Glass

The fire extinguisher consists of a vessel

metallic vessel containing sodium Conc.
bicarbonate or sodium carbonate NaHCO3 H2SO4
and a bottle of sulphuric acid. The solution
Cylinder
glass bottle of sulphuric acid is Nozzle
fitted inside the vessel in such a

way that it can be broken easily

when the knob is pressed down. Fire extinguisher
Sodium bicarbonate is filled in the

cylindrical glass bottle.

When the knob is pressed down or the cylinder is dropped on the

ground, the glass bottle breaks, whereby conc. H2SO4 and NaHCO3
solution get mixed with each other. The acid reacts with NaHCO3
and produces CO2 which comes out from the nozzle.

2NaHCO3 + H2SO4 → Na2SO4 + 2H2O + 2CO2

195 Times' Crucial Science Book - 10

Na2CO3 + H2SO4 → Na2SO4 + H2O + CO2
Thus produced CO2 extinguishes the fire.
Carbon dioxide gas is heavier than air and hence occupies the lower
most layer in the atmosphere. When carbon dioxide is supplied over
the burning objects, it displaces air (oxygen) and covers the burning
objects as that by a blanket. Thus, the fire gets extinguished due to
the lack of oxygen.

Ammonia

Molecular formula : NH3
Molecular weight : 17

Nature of bonding : Covalent bonds

Ammonia is a compound gas which consists of one atom of nitrogen
and three atoms of hydrogen. Lavoisier prepared ammonia for
the first time by heating ammonium chloride (Sal ammoniac) and
calcium hydroxide. But its composition was established by Berthecol
and Davy.

1p

7p
7n

1p 1p

Molecular structure of ammonia

Occurrence

Ammonia occurs free as well as in combined state. In free state, it is
present in air near the places where nitrogenous organic compounds
are decaying. In combined state, it is found as different ammonium
salts like, ammonium phosphate, ammonium chloride, ammonium
sulphate, etc.

General methods of preparation
1. By heating ammonium salts

Ammonium sulphate, ammonium carbonate, etc decompose on

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