Modern Graded Science - 9

7. Atoms of noble gasses do not form molecules and they have zero valency because these
atoms have a stable electronic configuration.

Things To Know

1. An element is the simplest form of a pure substance which cannot be split up into two
or more simpler substances by physical and chemical methods.

2. A compound is a substance which consists of the atoms of two or more elements
combined in a definite proportion by weight.

3. The smallest particle of a substance which can exit freely is called a molecule.
4. The smallest particles of elements which can take part in chemical change are

called atoms.
5. Anatom or agroup of atoms that is electrically charged due to the loss or gain of

electrons is called ion.
6. A positively charged ion and a negatively charged ion are called cation and anion

respectively.
7. A radical is defined as a charged atom or group of atoms having a positive or negative

charged which acts as a single unit during a chemical reaction.
8. The combining capacity of elements or radicals with the other elements or radicals is

called valency of those elements or radicals.
9. The outermost energy level or shell of an atom is termed a valence shell and the

electrons present in the valence shell are called valence electrons.
10. The shell is made up s, p, d, and f sub-shells and the sub-shells are made up of orbitals.
11. The maximum number of electrons the s-sub-shell can contain is 2, p-subshell is 6,

d-subshell is 10 and f-subshell is 14.
12. Elements like He, Ne, Ar, Kr, Xe and Rn are inert elements because they do not take

part in chemical reactions due to their valency being zero.
13. Helium has a single shell with two electrons, so it is called duplet.
14. Inert elements (except helium) contain 8 electrons in their outermost shell, so they are

called octet.
15. Octet rule is a process by which the atoms of the elements try to maintain 8 electrons

in their valence shell during the formation of a molecule.
16. The tendency of atoms to attain the stable electronic configuration is the real cause of

a chemical reaction.
17. The chemical bond formed by the transfer of electrons from the valence shell of one

atom to the outermost shell of another atom is called an electrovalent bond.
Classification of Elements 147

Things To Do

1. Prepare the atomic models of twenty elements from the atomic number 1 to 20 by
using beads, tape, thread on a cardboard.

2. Prepare the models of molecular structures of water and methane by using beads, wire
tape and wooden frames of different sizes.

Test Yourself

1. Multiple Choice Questions (MCQs)

a. Water is …………………… .

A. electrovalent compound B. covalent compound

C. ionic compound D. coordinate

b. The electronic configuration of oxygen is 2, 6 and its valency is ………. .

A. 1 B. 2 C. 4 D. 0

c. In d subshell of an atom, the maximum number of electrons that can be adjusted is...

A. 14 B. 10 C. 6 D. 2

d. A covalent bond is formed by …………….. of electrons.

A. sharing B. losing C. gaining D. all of the above

e. The noble gasses like Ne, Ar, etc. are ………… gases.

A. diatomic B. monoatomic C. metallic D. reactive

f. ……. compounds conduct electricity in a solution.

A. ionic B. covalent C. co-ordinate D. pure water

g. …….. is in the shape of a dumb bell..

A. s-orbital B. p-orbital C. d-orbital D. f-orbital

2. Answer the following questions.

a. Define valency.

b. What is the valency of aluminium and nitrate in Al (NO3)3?
c. How many electron pairs are shared by an oxygen molecule?

d. The valency of fluorine is 1. What is its valence electron?

e. What are radical and ions?

f. Define

i. element ii. compound iii. symbol iv. molecule

v. atom vi. metals vii. electrovalent compound viii. covalent compound

ix. chemical bond x. metalloids

g. What are radicals? Mention their types.

h. What is valency? Write the valency of the following radicals and elements.

CO3, PO4, Na, H, O, OH, NH4, Cl, Ca, Al, Si, Hg (ic), NO3
148 Modern Graded Science Class 9

i. In two atoms, atom X loses two electrons and atom Y gains these two electrons to
from a molecule XY. Write any two properties of the compound XY.

j. What do you understand by valence electrons and valence shell?
k. What is octet rule?
l. What are duplet and octet elements? Differentiate.
m. What is an electrovalent bond? Explain with an example.
n. What is a covalent bond? How is it formed? Explain with an example.
o. How many sub-shells are there in L-shell? Name them.
p. Write the electronic configuration of sodium and chlorine.
q. What do you mean by a molecular formula? What information is conveyed by a

molecular formula?
r. Enlist the characteristics of the following:

i. Electrovalent compound ii. Covalent compound

3. Give reasons. b. Neon is a inert element.
a. Oxygen is a diatomic element. d. Some elements have variable valencies.
c. The valency of sodium is one.

4. Differentiate between:
a. electrovalent bond and covalent bond
b. acidic radical and basic radical
c. Ion and radial

5. Diagrammatic questions

a. Write the atomic structure of the following

i. sodium ii. potassium iii. oxygen iv. argon

b. Answer the following questions on the basis of the given
atomic structure:
Element Electronic configurations
i. Is it a metal or non-metal?
A 1s2, 2s2, 2p6, 3s1

ii. What will be its valency? B 1s2, 2s2, 2p6, 3s2, 3p5

iii. It combines with chlorine to form C 1s2, 2s2, 2p6, 3s2, 3p6, 4s2
potassium chloride. Write the 1s2, 2s2, 2p6
molecular formula and the molecular D

structure of potassium chloride.

c. Study the given table and answer the following questions.

i. Name metals and non-metals.

ii. What is the valency A and why?

iii. Write the name and symbol of all elements.

Classification of Elements 149

iv. Write down the molecular farmula of the compound formed by the combination of
A and B; and C and B.

v. Which element is chemically active in between A and D?

d. Sketch the molecular structures of the following:

i. Cl2 ii. H2 iii. NH3

iv. MgO v. Na2S vii. AlCl3

viii. N2 ix. CO2 x. K2O

6. Write the molecular formulae of the following compounds by criss-cross method

a. Ammonium sulphate b. Zinc nitrate

c. Sodium peroxide d. Carbon tetrachloride

e. Sliver nitrate f. Ferric hydroxide

g. Ammonium chloride h. Aluminum sulphate

i. Mercuric oxide j. Calcium carbonate

k. Potassium cyanide l. Auric chloride

m. Aluminium hydroxide n. Potassium thiosulphate

o. Magnesium bicarbonate p. Magnesium sulphate

7. Write the molecular formulae of the following compounds.

a. Sodium oxide b. Cuprous oxide

c. Mercuric oxide d. Calcium nitrate

e. Aluminium oxide f. Ferric chloride

g. Ammonium chloride h. Silver chloride

i. Hydrogen bromide j. Magnesium nitride

k. Hydrogen peroxide l. Sodium sulphate

m. Ammonium sulphate n. Ferrous sulphate

o. Sodium carbonate p. Zinc carbonate

q. Calcium silicate r. Potassium bicarbonate

s. Potassium chlorate t. Aluminium hydroxide

u. Ammonium hydroxide v. Sodium silicate

Amphoteric : Having the characteristic of both acid and base both.

Transition : changing from one state or condition to another

Oxide : a compound made up of metals or nonmetals and oxygen

Electronegativity : the tendency of an atom or radical to attract electrons in the formation
of an ionic bond

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150 Modern Graded Science Class 9

Chapter Chemical Reaction

9

Total estimated PDS: 5 (5T/OP

Competencies

On completion of this lesson, the students will be competent to:

define physical and chemical changes.

define a chemical equation and its components.

represent chemical reaction in the form of chemical equations and methods of
their balancing.

In our daily life, we observe many changes around us. The changes occur in all the substances.
For example, germination of seeds, digestion of food, burning of coal, melting of ice, rusting
of iron and making salt solution etc. When changes take place in substances, either energy is
given or energy is released from them. All the changes can be classified into physical change
and chemical change beside nuclear reaction.

Physical change

Physical change is a temporary and usually reversible change in which no new substances are
formed. In this change, the chemical composition of the substance remains the same but only
the physical properties like state, colour, speed, motion, shape, size, etc. are changed.

Some of the examples of physical change are glowing of an electric bulb, magnetizing an iron
nail, preparing a solution, changing of ice into water and water vapour etc. Physical change is
a reversible change as it occurs in forward as well as backward directions.

Chemical change ice oonnhcoeaotliinngg water.

Chemical change is a permanent and it is usually an irreversible change in which new

substances with different properties are formed. In this change, the physical properties, the

chemical properties and the chemical composition of the substances are changed.

Some common examples of chemical change are changing milk into curd, rusting of iron,

burning magnesium ribbon, decaying grasses, cooking rice, etc. A chemical change is an
irreversible change as it cannot be reversed back from products to reactants. Chemical changes
are called chemical reactions.
Magnisium + Oxygen ∆ Magnesium oxide

reactants product

Chemicaql Reaction 151

Chemical reaction

The process by which a chemical change takes place is called a chemical reaction. It occurs
by addition, decomposition or displacement of the atoms or molecules of the matter into new
substances. For example, when nitrogen reacts with hydrogen to form ammonia.

Nitrogen + Hydrogen Ammonia

N2 + 3H2 2NH3

Those substances that take part in a chemical reaction are called reactants. Likewise, those

substances which are formed after a chemical reaction are called products. Presentation of a
chemical reaction in the form of equation is a called chemical equation. Chemical equations
may be word equations or formula equations.

a. Word equation

The chemical reaction which is expressed in terms of full names of reactants and product is

called a word equation. For example, Hydrogen + Oxygen → Water

Reactants Product

Reactants are written at first and then an arrow is given which means "forms" or "gives" and
then products are written. If reactants or products are two or more than two, they are joined
by the sign of '+'.

The sign of plus (+) on the reactant side indicates 'reacts with' while the sign of plus (+) on
product side indicates 'and' or 'in addition to'. In some cases, two arrows are written between

the reactants and the products by pointing to two ways. These two way arrows () are

generally used for showing the reversible reaction.

The required conditions for the chemical change can be written above the arrow to
make the equation more informative. For example, when water is electrolysed, it gives
hydrogen and oxygen.
Water electricity Hydrogen + Oxygen

In the chemical reaction, the state of reactants and products can be expressed by the following
signs: 's' for solid, '' for liquid, 'g' for gas and 'ag' (aquous) for a solution in water.

For examples,
Water () electricity Hydrogen(g) + Oxygen (g)

b. Formula equation

The chemical reaction which is expressed by using molecular formulae of reactants and
products is called a formula equation. For example,

152 Modern Graded Science Class 9

Hydrogen + Oxygen water

2H2 + O2 2H2O

Skeletal chemical equation The molecules of gaseous
elements like H, N, O, F,
In order to write a chemical reaction, we write a formula Cl, Br and I are diatomic.
equation. If a formula equation holds an unequal number Hence, their molecules are
of atoms of one or more elements on its reactants, the written as H2, N2, O2, F2,
products are called an imbalanced or skeletal formula Cl2, Br2, and I2,.
equation. Thus, the chemical equation in which the total

number of atoms of each element in reactant and product is not equal is called an unbalanced

chemical equation or skeletal chemical equation. Some examples of such chemical equations
are as follows:

1. Sodium + Chlorine Sodium chloride

Na + Cl2 NaCl

2. Nitrogen + Hydrogen Ammonia

N2 + H2 NH3
Zinc chloride + Hydrogen
3. Zinc + Hydrochloric acid

Zn + HCl ZnCl2 + H2

Drawbacks of a skeletal chemical equation

1. A skeletal chemical equation does not follow the law of conservation of mass.
2. It does not tell about the ratio of reactant and product molecules.
3. It does not give idea about the number of atoms of each element in reactants and

products.

Balanced chemical equation

An unbalanced chemical equation is necessary to be balanced by comparing the numbers of
atoms in reactant and product molecules. The chemical equation, in which total number of
atoms in reactants and products is equal is called a balanced chemical equation. A balanced
chemical equation is more informative than an unbalanced chemical equation. It is also based
on the law of conservation of mass. For examples,

Sodium reacts with chlorine to produce sodium chloride.

Sodium + Chlorine Sodium chloride [word equation]

Na + Cl2 NaCl [skeletal equation]

2Na + Cl2 2NaCl [balanced equation]

In this balanced chemical equation, the number of sodium and chloride atoms is equal on
each side of the chemical equation.

Chemicaql Reaction 153

Writing and balancing chemical equations

There are various ways to balance a chemical equation. Hit and trial method is the simplest
one. Hit and trial method is also termed as hit and success method. The following points
should be kept in mind to write and balance a chemical equation using this method.

1. Write the molecular formulae of all the reactants and products correctly.

a. To write molecular formulae of elements, write the symbols only. For example, Na
for sodium, Ca for calcium, Au for gold, S for sulphate, etc. But few elements are
d(Nia2t)o, mOixcyig.ee.nt(hOei2r),mFoluleocruinlees(hFa2)v,eCthwlooraintoem(Cs.l2E),xBamropmleisnear(eBHr2)ydarnodgIeond(iHne2)(,In2)i.trogen

b. To write molecular formulae of compounds, use the symbols and valences of the
radicals involved in that molecule.

2. Separate reactants and products by a sign of arrow. If reactants or products are more than
one, connect them by a sign of plus.

3. Balance the atoms of O and H at last (the atoms used at many places in an equation should
be balanced at last).

4. For balancing, the number should be added as co-efficient i.e. in the front of the molecules.
While balancing count the atoms in following ways.

a. The number written at the right lower corner of an atom is counted for that atom only
For example, in MgSO4, there are one 'Mg', one 'S' and four 'O'.

b. The number written at the right lower corner of a bracket is for all the atoms enclosed
in the bracket. For example, in Al2 (SiO3) has two Al, three 'S' and nine 'O'.

c. The coefficient number is for all the atoms of the molecule. For example, in 2Al2
(SiO3)3, there are four Al, six Al and eighteen 'O'.

Example: ammonia gas is prepared by the reaction of ammonium chloride and calcium
hydroxide. The reaction is given below.

1. Ammonium choride + Calcium hydroxide Calcium chloride + Ammonia +
water [Word equation]

2. NH4Cl + Ca(OH)2 CaCl2 + NH3 + H2O [Skeletal equation]

3. Here NH4Cl is multiplied by 2 to equalize with the Cl atoms.

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

4. Now, to equalize N atoms, NH3 is multiplied by 2.

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

5. To equalize oxygen and hydrogen, H2O is multiplied by 2.

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

In this way, we get the balanced chemical equation.

154 Modern Graded Science Class 9

Some more examples
1. When ethane is burnt in oxygen, carbon dioxide gas and water are produced.
Ethane + Oxygen heat Carbon dioxide + Water
[Word equation]

C2H6 + O2 ∆ CO2 + H2 [Skeletal equation]

2C2H6 + 7O2 ∆ 4CO2 + 6H2O [Balanced equation]

2. When potassium burns with oxygen to form potassium oxide.

Potassium + Oxygen ∆ Potassium oxide [Word equation]

K + O2 ∆ K2O [Skeletal equation]
4K + O2 ∆ 2K2O [Balanced equation]

3. Copper reacts with concentrated nitric acid to form copper nitrate, nitrogen dioxide and
water.

Copper + conc. Nitric acid Copper + Nitrogen + water [Word equation]

nitrate dioxide

Cu + HNO3 Cu(NO3)2 + NO2 + H2O [Skeletal equation]

Cu + 4HNO3 Cu(NO3)2 + 2NO2 + 2H2O [Balanced equation]

4. When calcium carbonate is heated, it forms calcium oxide and carbon dioxide.

Calcium carbonate heat Calcium oxide + Carbon dioxide [Word equation]

CaCO3 heat CaO + CO2↑ [Skeletal equation]
CaCO3 heat CaO + CO2↑ [Balanced equation]

5. When potassium chlorate is heated in the presence of manganese dioxide, it produces

potassium chloride and oxygen.

Potassium chlorate heat Potassium chloride + oxygen [Word equation]

KClO3 MnO2 KCl + O2↑ [Skeletal equation]

2KClO3 2KCl + 3O2↑ [Balanced equation]

6. When nitrogen reacts with oxygen at high temperature to form nitric oxide.

Nitrogen + Oxygen 3000oC Nitric Oxide [Word equation]

N2 + O2 NO [Skeletal equation]

N2 + O2 2NO [Balanced equation]

7. When zinc reacts with hydrochloric acid to form zinc chloride and hydrogen.

Zinc + Hydrochloric acid Zinc chloride + Hydrogen [Word equation]

Chemicaql Reaction 155

Zn + HCl ZnCl2 + H2 [Skeletal equation]

Zn + 2HCl ZnCl2 +H2 [Balanced equation]

8. When a solution of silver nitrate is treated with a solution of sodium chloride, a white
precipitate of silver chloride is formed.

Silver nitrate + Sodium chloride Silver chloride + Sodium nitrate [Word equation]

AgNO3 + NaCl AgCl + NaNO3 [Skeletal equation]

AgNO3 +NaCl AgCl + NaNO3 [Balanced equation]

Information conveyed by a balanced chemical equation

The following information can be obtained from a balanced chemical equation.
1. Formulae of the reactants and products
2. Total number of atoms or molecules of reactants and products
3. Type of chemical reaction
4. Ratio of masses of reactants and products.
5. Names of the type of reactants and products.

Limitations of chemical equations

1. Chemical equations cannot give the physical state of reactants and products.
2. It does not give any information about the rate of chemical reaction.
3. It does not tell about the time taken by the reaction to complete.
4. It does not explain the actual concentration of the reactants.
5. It does not state whether the reaction occurs by itself or certain conditions like heat,

light, catalyst are required.
6. It does not show that the reaction is either explosive or not.
7. It does not tell that chemical reaction will start immediately or it will take time to be

started.

Endothermic and exothermic reactions

Chemical reactions take place with either production of heat or absorption of heat. Accordingly
chemical reactions are classified into two types: exothermic reaction and endothermic reaction.

Exothermic reaction : The chemical reaction in which heat is released to the surroundings is
called exothermic reaction. The amount of heat energy released is written along with products.
This indicates that heat is a given out.

Examples,

a. C + O2 CO2 + Heat c. Zn + 2HCl ZnCl2 + H2↑ + Heat
b. CH4 + 2O2 Ca(OH)2 + Heat
CO2 + 2H2O + Heat d. CaO+ H2O

156 Modern Graded Science Class 9

Endothermic reaction: The chemical reaction which is occurred by the absorption of heat
from the surroundings is called an endothermic reaction. The amount of heat energy given is
written along with reactants. This indicates that heat is absorbed.

Examples, CaO + CO2↑ c. N2 + O2 + Heat 2NO↑
a. CaCO3 + Heat 2Hg+O2↑ d. 2KClO3 + Heat 2KCl + 3O2↑
b. 2HgO + Heat

Catalyst

A catalyst is a chemical substance that alters the rate of chemical reaction but is not consumed
by the reaction. Hence, a catalyst can be recovered chemically unchanged at the ends of
chemical reaction. Catalyst can be divided into two types, they are positive catalyst and
negative catalyst.

Positive catalyst: The catalyst that increases the rate of a chemical reaction is called a positive
catalyst. Example,

When potassium chlorate is heated with manganese dioxide (MnO2), it increases the rate of
decomposition of potassium chloride and oxygen. Thus, MnO2 is a positive catalyst.

2KClO3 ∆ 2KCl + 3O2↑
MnO3

Negative catalyst: The catalyst that decreases the rate of a chemical reaction is called a negative

catalyst. Example,

When hydrogen peroxide is mixed with glycerol [C3H5(OH)3], the glycerol decreases the rate
of decomposition of hydrogen peroxide into water and oxygen. Thus, glycerol is a negative
catalyst.

2H2O2 Glycerol 2H2O + O2↑

Characteristics of catalyst

1. The mass and chemical composition of a catalyst Sodium hydroxide and
remains unchanged in the chemical reaction. potassium hydroxide are called
caustic soda and caustic potash
2. A small quantity of the catalyst is required for the respectively. This is because
chemical reaction. they produce burning effect on
the skin.
3. The catalyst is specific in nature. It means a specific
catalyst can be used only for a specific reaction.

4. The catalyst does not initiate a reaction.

Chemicaql Reaction 157

S me Reasonable Facts

1. Rusting of iron is a chemical change because it is a permanent change in which a
substance called iron oxide is formed and the mass of iron is increased after rusting.

2. Glycerol is a negative catalyst because it decreases the rate of chemical reaction such as
the decomposition of hydrogen peroxside to water and oxygen.

Things To Know

1. A physical change is a temporary and reversible change. During the change no new
substance is formed.

2. A chemical change is a permanent and irreversible change in which a new substance
with different properties is formed.

3. A chemical equation is representation of a chemical reaction in terms of equation.
4. Reactants are those substances that undergo a chemical change during a chemical

reaction.
5. Products are those substances, which are formed as a result of a chemical change in the

reactants.
6. A chemical equation can be expressed in terms of words which is called a word

equation.
7. The chemical equation which is written in terms of molecular formulae is called a

formula equation.
8. Any formula equation that holds an unequal number of atoms of each element on its

two sides, i.e., reactants and products is called a skeletal equation.
9. If the number of atoms of each element in the reactants and products of a chemical

equation are equal, it is called a balanced chemical equation.
10. Hit and trial method of balancing a chemical equation involves the equalization of

the number of atoms of each side of the equation by multiplying the reactants and
products with a suitable coefficient.
11. The chemical reaction which is occurred by the absorption of heat from the
surroundings is called an endothermic reaction.
12. The chemical reaction in which heat is released to the surroundings is called exothermic
reaction.
13. The chemical substance that alters the rate of chemical reaction is called catalyst.
14. The catalyst that increases the rate of a chemical reaction is called a positive catalyst.
15. The catalyst that decreases the rate of a chemical reaction is called a negative catalyst.

158 Modern Graded Science Class 9

Things To Do

Make two lists of basic radicals and acid radicals with their symbols and volencies. Paste them
on the wall in front of your study table and memorise the symbols and valencies of four basic
and two acid radicals daily.

Test Yourself

1. Multiple choice questions (MCQs)

a. The ……………… change is the type of change in which new substances with

different properties are formed.

A. physical B. chemical C. organic D. reversible

b. What are the substances which take part in chemical reaction called?

A. products B. chemical substances C. reactants D. physical substances

c. The rate of a reaction is increased by the presence of another substance called a …

A. catalyst B. positive catalyst C. negative catalyst D. promoter

d. The rate of reaction depends on the ……………… of the reactions.

A. colour B. state C. nature D. atomic weight

e. Heat is absorbed in an …………. reaction.

A. exothermic B. reversible C. endothermic D. irreversible

2. Answer the following questions.
a. What is a chemical reaction? In how many ways can chemical reactions take place?
b. Enlist any four examples of chemical change.
c. What do you mean by a diatomic element? Give any four examples.
d. What is a catalyst? Write the characteristics of a catalyst.
e. What information can be obtained from a balanced chemical equation?
f. Enlist the limitation of a balanced chemical reaction.

g. Define
i. reactant ii. product iii. word equation iv. formula equation

3. Differentiate between:
a. physical change and chemical change
b. skeletal chemical equation and balanced chemical equation
c. endothermic reaction and exothermic reaction.

4. Change the following word equations into balanced equations. 159
a. Hydrogen + Chlorine → Hydrogen chloride
b. Hydrogen + Nitrogen → Ammonia
c. Potassium chlorate → Potassium chloride + Oxygen
Chemicaql Reaction

d. Zinc + Hydrochloric acid →Zinc chloride + Hydrogen
e. Copper + Sulphuric acid →Copper sulphate + Water + Sulphur dioxide
f. Aluminium + Sulphuric acid→Aluminium sulphate + Hydrogen
g. Lime water + Carbon dioxide →Calcium carbonate + Water
h. Ammonia + Hydrogen chloride →Ammonium chloride
i. Sodium carbonate + Water →Sodium hydroxide + Carbonic acid
j. Iron + Oxygen → Iron oxide
k. Phosphorus + oxygen → Phosphorus pentoxide

5. Balance the skeletal equations.

a. K + O2→ K2O b. Hg + O2→HgO
c. Na + Cl2→ NaCl d. AgBr→ Ag + Br2
e. KClO3→KCl + O2 f. Fe + CuSO4→ Fe SO4 + Cu

6. Fill in the gaps and balance the following equations.

a. Ca(OH)2 + ............ → CaCO3 + ............
b. H2O2→ H2O + ............
c. Mg + ............ → MgO

d. S + ............ → SO2
f. ............ + S → FeS

g. CaCO3→............ + ............

7. What happens when:

a. Magnesium burns in oxygen

b. Calcium carbonate is heated

c. Hydrogen reacts with nitrogen at necessary conditions

d. Potassium hydroxide reacts with nitric acid

e. Carbon dioxide is treated with limewater

8. Potassium hydroxide reacts with hydrochloric acid to form salt and water.
a. Write the word equation and the balanced formulae equation for the above reaction.
b. What will happen if sodium hydroxide is used instead of potassium hydroxide?

Reversible : the reaction in which the reactants form the products and vice versa
Electrolysis : separation of a substance into its chemical components by passing an

electric current in its solution
Irreversible : the chemical reaction occurs only in one direction (reactants → product)

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160 Modern Graded Science Class 9

Chapter Solubility

10

Total estimated PDS: 6 (3T/3P

Competencies
On completion of this chapter, the students will be competent to:

define solution and its types.
prepare unsaturated and saturated solution.
define solubility of a substance.
define super saturated solution
define the relation between solubility of a substance and temperature.
describe crystallization.

We have different types of substances around us. Some of them are pure substances and others
are mixtures. Mixtures are impure substances which are made by different substances. A
mixture is defined as a mass which is obtained by mixing two or substances in any proportion
by weight. Examples are sand in rice, sand and water, etc. On the basis of the size of the
particles of the components of the mixture, they are divided into suspensions, colloids and
solutions. Study them comparatively in the following table.

Basis Solutions Colloids Suspensions
1. Size of particles Smaller than 10–7 cm Intermediate of 10–7 cm and 10–3 cm Bigger than 10–3cm
2. Visibility of particles Invisible (Visible un- May be visible in beam of light (visi- Visible under naked
der ultramicroscopic) ble under light microscope) eyes
3. Settlement of parti- Do not settle down Does not settle down in an ordinary Settle down easily
cles when kept still condition
for some time Cannot be filtered Can be filtered
sually transparent Cannot be filtered Opaque
4. Filtration of particles Cloudy translucent Heterogeneous
5. Appearance of mixture Intermediate of homogenous and
heterogeneous Sand water, muddy
6. Type of mixture Homogeneous blood, milk, gum, dust in air, etc. water, etc.

7. Examples salt + water, alcohol +
water, air, brass, etc.

Thus, solutions are homogeneous mixtures and suspensions are heterogeneous mixtures but
colloids are the intermediate of homogeneous and heterogeneous mixtures. Due to this reason,
sometimes, they are termed as colloidal solutions and colloidal suspensions too. Sometimes,
colloids are also termed as emulsion which means immiscible liquids.

Solubility 161

Solution

A solution is defined as the homogeneous mixture of two or more than two substances. It
may be formed by any state of matter. The size of the particles in a solution is 10-7 cm or less.
A powerful microscope is required to see the particles in a solution. In other words, Solution
is defined as the homogeneous mixture of a solvent and a solute. For example, sugar and
water forma sugar solution, salt and water form a salt solution, brass is formed from zinc and
copper, etc. Thus, solution = solute + solvent.
a. Solute: A solute is the component of a solution which gets dissolved into solvent to form

a solution. In a solution, if the components are in different states, the solute changes its
state. If the components are in the same state, the amount of the solute is found lesser
than the amount of the solvent.
b. Solvent: A solvent is the substance which dissolves the solute to form a solution. If the two
components of a solution are in different states, the solvent does not change its state.If both
the components of a solution are in the same state, the solvent is in a larger amount.
1. Water: It is a universal solvent found in the liquid state. It has the capacity of dissolving

many substances like common salt, copper sulphate, sugar, etc.
2. Alcohol: It is also found in the form of spirit. It can dissolve resin and iodine. The

medicine named tincture of iodine is prepared by dissolving iodine in alcohol.
3. Petrol and kerosene: These can dissolve ghee, grease, oil, fat, etc. They are used for

removing the stain of these substances in clothes.
4. Ether: It is an organic solvent. It can dissolve fat, oil, resin, etc.

Some examples of solutions

Solution formed in Examples of solution Solute Solvent
A. Solid 1. Brass Copper Zinc
2. Amalgam Mercury Zinc
3. Hydrogen in platinum Hydrogen Platinum
Salt Water
B. Liquid 1. Salt solution Alcohol Water
2. Alcohol in water, (when amount of water is more
Carbon dioxide Water
than that of alcohol) Carbon particles Air
3. Soft drinks Water vapour Air
Oxygen Nitrogen
C. Gas 1. Smoke in air
2. Water vapour in air
3. Air

Unsaturated and saturated solutions

Activity 10.1
To make saturated and unsaturated solutions.

1. Materials needed, water, beaker, sugar crystal, glass rad. take a beaker containing some
water at the room temperature.

162 Modern Graded Science Class 9

2. Add a little amount of sugar crystals in it and stir it with a glass rod. The sugar crystals get
dissolved in it.

3. Add some more crystals of sugar to it and stir again. It also gets dissolved.
4. Repeat this process of dissolving more and more sugar in that fixed amount of water till

some of the sugar crystals remain in the undissolved state in the solution.

From the activity, it is clear that there
is a limit of dissolving sugar by a fixed
amount of water at a definite temperature.
When it reaches the limit, the solution
is said to be a saturated solution at that

temperature. The solution which can

dissolve excess amount of solute in it

at the given temperature is called an Fig. 10.1
unsaturated solution at that temperature.

The solution can be changed into saturated by cooling it or by adding more solute in it. It is

less dense than the saturated solution. The solution which cannot dissolve the excess amount

of solute in it at a particular temperature is called a saturated solution at that temperature. It

is denser due to the presence of maximum amount of solute in it. A saturated solution can be

changed into unsaturated by heating the solution or by adding more solvent in it.

Difference between unsaturated and saturated solutions

Unsaturated solution Saturated solution
1. This solution has capacity of dissolving more 1. This solution is unable to dissolve more solute at a partic-

solute at a particular temperature. ular temperature.
2. On heating, it remains unsaturated. 2. On heating, it becomes unsaturated.
3. Precipitation does not appear on cooling. 3. Precipitation of the solute appears on cooling.
4. It has less density. 4. It has more density.

5. It has less concentration of the solute. 5. It has more concentration of the solute.

Supersaturated solution

Activity 10.2
To prepare a supersaturated solution.

Apparatus required: beaker, water, salt glassrod, tripods stand, wire gauel and spirit lamp.

Procedure:
1. Take some amount of saturated solution in a beaker and heat the beaker gently.
2. Add a little amount of sugar in the beaker containing the saturated solution with
increasing temperature. You will see that the additional sugar gets dissolved.

Solubility 163

3. Continue this process of dissolving the sugar more and more by keeping the definite
volume of solution. After some time, we can see that no more sugar dissolves in
the solution.

4. Cool the saturated solution in cold water. The solution crystallizes less amount of solute
and the solution holds more amount of solute than the required amount of solute for
saturated solution at that temperature. It is a supersaturated solution.

Hence, the saturated solution at higher temperature that holds excessively more solute
than the required amount for a saturated solution at that temperature is called a
supersaturated solution.

Activity 10.3
To identify the unsaturated, saturated and supersaturated solution

Apparatus required: beakers, glass rod, water, crystals of sodium chloride.
Procedure:

1. Take three beakers A, B and C filled with equal amount of unsaturated, saturated and
supersaturated salt solutions respectively and stir them with a glass rod.

2. Put a small salt (crystals) in each beaker.
3. What will you observe in each beaker containing the solution?
In beaker A, the crystals of salt are gradually dissolved, so the solution of beaker A forms an
unsaturated solution. In beaker B, there is no change in the size of salt crystal, when added in
the solution. So, the solution of beaker B is a saturated solution.

Fig. 10.2 (b) unsaturated solution (b) saturated solution (c) supersaturated solution

In beaker C, there is increase in the size of the crystal. Thus, the solution of beaker C is a
supersaturated solution.

Importance of solution

The following points show the importance of solution:
1. Oxygen is found dissolved in water and in air which is used by living
organisms for respiration.
2. Plants take different minerals and salts from the soil in the form of a solution.
164 Modern Graded Science Class 9

3. The digested food circulate in different parts of the body in the form of a solution.
4. Most of the reactions are taking place easily in the form of a solution.
5. Most of the medicines and paints are found in the form of a solution.
6. Solution is essential in many industries.

Solubility of a substance

Different substances have different physical and chemical properties. So, different solutes of
the same weight cannot be dissolved in the same volume of the given solvent at a particular
temperature. The amount of the solute dissolved in a saturated solution at a fixed (given)
temperature is solubility of the solute. Thus, the solubility of a substance (solute) at a given
temperature is defined as the amount of the substance (solute) dissolved in 100 g of solvent to
make a saturated solution at that temperature.

Mathematically,

The solubility of some solute substances at 30°C is tabled below:

S.N0. Solute substance Temperature °C Solubility g/100 g of water
1. Potassium chlorate 30 9
2. Copper sulphate 30 25
3. Potassium nitrate 30 27
4. Sodium chloride 30 36
5. Sodium nitrate 30 95
6. 30 220
Sugar

The solubility of any chemical substance such as

copper sulphate at a given temperature 30oC is 25

means, 25g of copper sulphate is dissolved in 100g Many salts like barium nitrate and sodium

of water to make a saturated solution at that 30oC biarsenate show a large increase in solubility
with temperature. A few such salts as calcium
temperature. The solubility of different substances sulphate and cerium sulphate become less

at a given or constant temperature is different due soluble in water as temperature increases.

to their different chemical properties.

Solubility 165

Example 1: At 20°C, 5.1 gram of sugar dissolves in 2.5 gram of water to form a saturated sugar
solution. Find the solubility of sugar at that temperature.

Solution
Here, Weight of solute (w1) = 5.1 g.
Weight of solvent (w2) = 2.5 g
Solubility at 20°C (s) = ?

We know that, Weight of solute (w1)
Weight of solvent(w2)
Solubility (s) = × 100

= 5.1 × 100 = 204
2.5

Therefore, the solubility of sugar at 20°C is 204.

Example 2: The solubility of common salt at 30°C is 200. What amount of water is required at
that temperature to prepare a saturated solution of 1kg of salt?

Solution

Here, Weight of solute (w1)= 1kg = 1000 g
Solubility of salt at 30°C (s) = 200

Weight of solvent (w2)= ?

We know that, Weight of solute(w1)
Weight of solvent(w2)
Solubility (s)= × 100

1000
or, 200 = w2 × 100

or, W2 = 100200×0100= 500 g

Therefore, the weight of water in the saturated solution = 500 g.
Example 3: If 75 gram of a saturated solution of sodium nitrate in water at 30°C, contains 15
gram of solute, find the solubility of sodium nitrate at that temperature.

Solution
Here, Weight of saturated solution (w) = 75 g
Weight of solute (w1)= 15 g
Weight of solvent(w2)= Weight of saturated solution (w) - Wt. of solute(w1)

= 75 - 15 = 60
Solubility (s) = ?
166 Modern Graded Science Class 9

We know that,

Solubility (s) = Weight of soluten (w1) × 100
Weight of solvent (w2)

= 15 × 100 = 25
60

Therefore, the solubility of sodium nitrate at 30° C is 25.

Example 4: When 15g of a saturated solution of sodium nitrate at 30oC is cooled down at
10oC, then how much sodium nitrate will be precipitated if the solubility of sodium nitrate at

30oC and 10oC is 95 and 30 respectively?

Solution

According to the statement:

At 30°C, 95 g of NaNO3 forms 195 g of saturated solution (95 g solute + 100 g solvent)
At 10°C, 30 g of NaNO3 forms 130 g of saturated solution (30 g solute + 100 g of solvent)

The difference in the weight of solution is: 195 – 130 = 65 g

∴ 195g of saturated solution when cooled from 30° C to 10° C separates 65 g NaNO3

∴ 1 g of saturated solution when cooled from 30° C to 10°C separates 65 g of NaNO3
195
65
∴ 15 g of saturated solution is when cooled from 30°C to 10°C separates 195 × 15 g of NaNO3

= 5 g.

Thus, 5 g sodium nitrate is separated by cooling the solution from 30° to 10°C.

Effect of heat on solubility

Prior to studying solubility, it is better to consider about solution. How is a solution prepared?
How is the status (position) of the molecules of the solute, the solvent and the solution?

The speed of molecules in a solution becomes faster or slower on heating. What does the
statement say?

The molecules of a solute in the solid state are held together by intermolecular force of
attraction. So, the molecules of solid substances are compactly packed. When the solid
substance is heated, the molecules begin to vibrate about their mean position. The vibrating
molecules possess kinetic energy. With the rise in temperature, the kinetic energy of molecules
increases. This helps the molecule to move in high speed.

The molecules of a solvent in the liquid state are loosely held together due to less intermolecular
force of attraction in comparison to solid substances. So, the molecules of liquids are loosely
packed. The molecules of a solvent in the liquid state are constantly moving around with
different speeds. If we heat the solvent, the kinetic energy of the molecules of a solvent increases
more and they move faster. Thus, on heating the solution, the kinetic energy of the molecules

Solubility 167

increases more and the molecules move in a high speed.
Similarly, on shaking, the kinetic
energy of molecules of a solution
increases and the molecules catch
up more speed. Thus, on heating
and shaking the kinetic energy of
the molecules of a solute increases.
As a result, the molecules of the
solute strike each other. Then they
are separated from each other. The
separated molecules of the solute get mixed with the molecules of the solvent to make a
solution.
When common salt is mixed in water, the negatively charged oxygen of water pulls the
positively charged sodium of the salt. Similarly, positively charged hydrogen atoms of water
pull the negatively charged chlorine. It helps to dissolve salt in water.

When molecules are heated at a high
temperature, the intermolecular
space increases. This is because the
inter-molecular force becomes
weaker. It helps to adjust more
molecules of the solute in the
increased space. For this reason, the
solubility of substances increases
with the rise in temperature.

On the other hand, if the solute substances in the solid state are finely crushed, the surface
area increases. It makes the molecules of the solvent come
in contact with the molecules of the solute at a faster rate.
This also increases the solubility of the substances.

Solubility curve

The solubility of the substances is different at different
temperatures. In general, the solubility of different
substances increases with the rise in temperature.

A curve obtained by plotting the solubility of a substance
against their respective temperatures is known as a
solubility curve. The solubility of copper sulphate at
different temperatures is given in the table below:

168 Modern Graded Science Class 9

Temperature (°C) 0 10 20 30 40 50 60 70
Solubility of copper sulphate 14 17 21 24 29 34 40 47

While drawing a solubility curve of a substance, temperature is taken in the x-axis and the
solubility of substances is shown at y-axis. Some solubility curves of different substances are
shown in the given diagram:

Information obtained from solubility curve

After studying the solubility curve, the following
information can be obtained.

1. The solubility of a substance at a particular
temperature can be known.

2. The solubility it helps us to compare the
solubility of different substances at the same
temperature.

3. It gives a clear idea that solubility of a substance
changes with the temperature.

4. The solubility curve helps us to predict what
ammount of solute will crystallize first from a
hot saturated solution when cooled.

Crystals

The saturated solution of a solid like copper sulphate is prepared at high temperature and it is
allowed to cool down. It is found that the excess solid is deposited spontaneously in the form
of a granular substance in geometrical shapes and is made of sharp edges and smooth surfaces.
Such granular substances are called crystals.
A crystal is a piece of solid particles which has a regular geometrical shape, smooth surfaces
and sharp edges.

Characteristics of crystals

1. They are solid particles with a definite or regular shape.

2. They are pure substances with a fixed melting point.

3. They are arranged in three dimensional patterns.

Crystallization

Crystallization is the process in which crystals are formed by cooling a hot saturated solution
of solid materials.

Solubility 169

Activity 10.4
To obtain crystals from copper sulphate solution
Apparatus required: Procelain basin, Glass rad, burner, tripot stand, wire gauze, copper
sulphate solution, filter paper, copper sulphate crystal.

Procedure:

1. Take a solution of copper sulphate in a porcelain basin.

2. Make the solution hot saturated either by dissolving
more copper sulphate in the solution or by evaporating
the solvent.

3. Continue heating the solution until the small crystals
are seen on the inner walls or bottom of the basin.

4. Now cut off the source of heat allow it to cool in air.
After a while, the crystals are settled at the bottom of
the basin.

5. Finally, separate the crystals from the mother liquor and place them on the filter paper

to dry. In this way, the crystals of CuSO4 are thread
obtained.

Sugar candy is a crystalline pure sugar formed by beaker
repeated boiling and slow evaporation of

concentrated solution of sugar. To prepare a ssuatguarrasteodlution
sugar candy, tie a small crystal of sugar candy

with a thread and dip in the concentration or small crystal
saturated sugar solution as shown in the figure. of sugar candy

After a few days, a large sized crystal of sugar in Fig. 10.8 formation of suger candy by crystillization

prepared. In this way, a sugar candy is formed by crystallization process.

Based on the shape, there are two kinds of solid substances. They are amorphous and crystalline.

1. Amorphous solids: Those solid substances which

do not have any definite geometrical forms or

shapes are called amorphous solids. On heating, If the solution is allowed to cool fast,
they do not melt but soften and become less viscous. small sized crystals are formed. If the
Lime, glass, rubber, plastics, etc. are some examples solution is allowed to cool slowly,

of amorphous solids. large sized crystals are formed.

2. Crystalline solids: Crystalline solids have typical

geometrical shapes along with definite and rigid morphology. The crystalline substances

have a sharp melting point, definite three-dimensional arrangement of the constituent

particles. Most of the solid, like NaCl, CuSO4, etc. are crystalline solids.

170 Modern Graded Science Class 9

S me Reasonable Facts

1. Temperature is also mentioned with solubility because solubility of a substance is different
at different temperatures.

2. When a saturated solution is heated, it becomes unsaturated. It is because when a saturated
solution is heated, the intermolecular spaces increase and more particles of solute can be
adjusted there.

3. Water is a universal solvent: It is because water can dissolve many types of solutes in it.
4. Soap does not form lather properly in winter. It is because in winter due to low temperature

of water, it cannot dissolve enough soap in it to form proper lather.
5. Brass is a solution: It is because brass is an alloy that is made by mixing two metals i. e.

copper and zinc in a proper ratio. It is a special type of solution of two solids and the
particle of any solid of the component is not visible in it.

Things To Know

1. There are two types of mixtures: homogeneous mixture and heterogeneous mixture.

2. On the basis of the size of the particles, mixtures are of three types: solution, colloids
and suspension.

3. Solution is a type of the homogeneous mixture of two or more than two substances.

4. A colloid is an intermediate of also a homogeneous mixture and heterogeneous mixture,
the particle size of which is in between that of particles of solution and suspension. It is
also termed as colloidal solution as well as colloidal suspension.

5. Suspension is a heterogeneous mixture which is formed when solid substances are
mixed with the liquid.

6. The solution which can dissolve excess amount of a solute in it at the given temperature
is called an unsaturated solution at that temperature.

7. The solution which is unable to dissolve excess amount of a solute in it at a particular
temperature is called a saturated solution at that temperature.

8. The solution at a particular temperature that has dissolved more than required amount
of solute for a saturated solution at that temperature is called a supersaturated solution.

9. The maximum amount of solute substance dissolved in 100g of solvent at a definite

temperature is called solubility of that solute at that temperature.
c=urwwv2e1
S × 100 by plotting the solubility of a substance at different temperatures
10. A obtained

against the respective temperatures is called a solubility curve.

Solubility 171

11. The solid particle having a regular geometrical shape and bounded by sharp edges is
called a crystal.

12. The process in which crystals are formed by cooling a hot saturated solution substance
is called crystallization.

Things To Do

Make large-sized crystals of sugar by crystallization process.
Test Yourself

1. Multiple choice questions (MCQs)

a. A mixture in which the components can be seen under naked eyes is:

A. solution B. colloid C. suspension D. emulsion

b. In a solution containing 1 liter alcohol and 2 liters water:

A. solute is alcohol and solvent is water

B. solute is water and solvent is alcohol
C. any one of them is a solute and another is a solvent

D. both of them are solutes and both are solvents

c. Which of the following does not affect on solubility?
A. temperature B. pressure C. nature of solute and solvent D. state of matter
4. Which component of a solution changes its state?

A. solute B. solvent C. both solute and solvent D. solution

d. Which character of the following is found in a suspension only, not in a solution and
a colloid?

A. homogeneous mixture B. visible particles

C. particles cannot pass through filter paper

D. particles move here and there in the mixture.
e. Which of the following is a heterogeneous mixture?

A. salt solution B. tincture of iodine C. water and alcohol D. muddy water

f. Which of the following cannot be known by using a solubility curve?

A. solubility of a substance at a particular temperature
B. comparison of solubility of two solute in the same solution

C. state of solute and solvent

D. relation of temperature and solubility of a substance

g. Which formula is used to calculate solubility of a substance?
100
A. S = W1 × W2 B. S = WW12 × 100
D. S = W1 × 100
C. S = W2 × W1
100 W2

172 Modern Graded Science Class 9

2. Answer the following questions

a. What is meant by a solution? Give any four examples of a solution.

b. What is a saturated solution? Give an example.

c. Define:

i. mixture ii. supersaturated solution

iii. solubility of a substance iv. homogeneous mixture

d. What is a solubility curve? Write the information the can be obtained from a solubility curve.

e. "Solution is a homogeneous mixture". Justify.
f. How can you identify unsaturated, saturated and supersaturated solutions?
g. What is the importance of solution?

h. Explain how the crystals of copper sulphate can be formed.

i. What are crystals? Write the characteristics of crystals. How are crystals made?

j. The solubility of sugar at 30°C is 220. What does it mean?

3. Differentiate between:
a. solution and suspension

b. saturated and unsaturated solutions

c. amorphous solids and crystalline solids

d. solute and solvent

e. solution and colloid

4. Give reasons.
a. In winter, soap does not dissolve properly in water.

b. Colloid is an intermediate state of solution and suspension.

c. A saturated solution becomes unsaturated when it is heated.

d. In the atmosphere, water is a solute.

e. Big crystals are formed when a hot saturated solution is cooled slowly.

5. Diagrammatic questions.

Temperature °C 0 10 20 30 40 50 60 70
Sodium chloride in 100 g water/g 25 30 35 37 39 40 41 41
Ammonium chloride in 100 g/g 28.4 32.8 37.3 41.3 46.2 50.6 55 64

a. Draw solubility curves of sodium chloride and ammonium chloride from the data above.

b. Answer the following questions by studying the solubility curves of the sodium
chloride and ammonium chloride given above.

i. What is the solubility of sodium chloride at 25°C. Solubility 173

ii. What is the relation between increase in temperature and solubility?

iii. Find out the amount of ammonium chloride precipitated when the saturated
solution prepared at 50°C is cooled to 20°C.

c. Answer the questions asked on the basis of given
diagram.

i. What does the demonstration show?

ii. What type of solution is required for the
process shown in the diagram?

iii. What will happen if the solution is cooled
fast?

iv. What will happen if the solution is cooled slowly?

6. Numerical problems.

a. At 40°C, 45 g of potassium chloride dissolves in 30g of water to from a saturated
solution. Find the solubility of potassium chloride.
(Ans: 150 )

b. At 30°C, 103g of saturated solution of sodium chloride contains 30.9g of sodium
chloride. Find the solubility of sodium chloride at 30°C.
(Ans: 42.85)

c. The solubility of sugar in water at 20°C is 204. Calculate the amount of sugar that
should be dissolved in 22g of water at 20°C to make a saturated solution.
(Ans: 44.88 g)

d. Five grams of water dissolves 60 gram of ammonium nitrate in it to form a saturated
solution at 30°C. Calculate the solubility of ammonium nitrate at that temperature.
(Ans: 1200)

e. How much of copper sulphate crystals will be separated if 25g saturated solution of
copper sulphate at 600 C is cooled down to 300C if solubility of copper sulphate at 600
C is 50 and at 300 C is 30 respectively.
(Ans: 3.33g)

Stain : a substance used for staining wood, fabric, dye, etc.

Inter molecular space : space between moleculs

Precipitation : Setteling down of undissolved paste like substance in a mixture

∆∆∆

174 Modern Graded Science Class 9

Chapter Some Gases

11

Total estimated Pds: 8/ (6T/2P

Competencies
On completion of this chapter, the students will be competent to:

describe the laboratory preparation of hydrogen, oxygen and nitrogen.
enlist the properties and uses of hydrogen, oxygen and nitrogen.

There are eleven elements which are found in the gaseous state. Some of them are hydrogen,
nitrogen, oxygen, inert gases like helium, neon, argon, etc. There are some compounds such as
carbon dioxide, water vapour, ammonia, etc. which are found in the gaseous state in nature.
Gases like nitrogen, oxygen, inert gases and water vapour mix together in different volumes to
form air. The thick layer of air which surrounds the earth is called atmosphere. The main gases
in the air, which occupy about 99% volume of the air, are nitrogen and oxygen. The composition
of different gases in the air with there percentage volume is tabled below:

The human body is composed of mainly four S.N. Gases in air % by volume
elements: hydrogen, oxygen, nitrogen and
carbon. It is estimated that about 96.2% of the 1. Nitrogen (N2) 78.08%
total body weight of human beings is made by
these elements. These elements are combined 2. Oxygen (O2) 20.95%
to other elements to form many other essential
substances. Such substances are essential to 3. Carbon dioxide (CO2) 0.039%
sustain the life of human beings.
4. Argon (Ar) 0.93%

5. Hydrogen (H2) 0.00005%

6. Ozone (O3) 0.000004%

7. Other gases (Ne, He, 0.3768%
CH4, Ar, Kr, etc.)
All living beings use oxygen for respiration.
(Source: www.physical geography.net)
Nitrogen is used by plants in the form of fertilizers. Plants take carbon dioxide and give out

oxygen during photosynthesis. Carbon dioxide is formed in the air also by the combustion of

fuel. In this way, these different gases are interchanged between air and living things and their

content remains in a balanced form in the atmosphere. The content of these gases is always

constant in the air. But the amount of moisture and carbon dioxide may vary from place to

place. Gases like hydrogen, nitrogen and oxygen can be obtained either from the atmosphere
or from manufacturing in industries.

Some Gases 175

A. Hydrogen

Hydrogen is found in the sun, stars and spaces between the stars. But it constitutes 0.00005%
of the earth's atmosphere by volume. It is an inflammable gas which reacts with other elements
to form various compounds like acid, hydrocarbon and carbohydrate, etc.

Facts about hydrogen

Symbol H 1p+ 1p+
Atomic number 1
Atomic weight 1.008 amu Fig. 11.1 molecular structure of hydrogen
Valency
Molecular formula 1
Molecular weight
Position in periodic table : Group H2
2.016 amu
: Period IA
Electronic configuration 1
(1s1 )
Boiling point -253°C
-259°C
Freezing point

Discovery

Hydrogen was first proved to be a distinct element by an English scientist, called Henry
Cavendish in 1766 who named it inflammable gas. In 1783, Lavoisier proposed the name
hydrogen to it, as it has capacity of forming water on burning with oxygen. The name hydrogen
means water producer.

Occurrence

Hydrogen is a reactive element and thus does not occur much in the free state. It is found in
volcanic gases and natural gases in trace amount. But it occurs abundantly in a combined
state. In the combined form, it is an important constituent of water, acid, alkali and many
organic compounds of vegetables and animal products. The chief source of hydrogen is water
in which its amount is double in comparison to oxygen.

General methods for preparation of hydrogen gas

Hydrogen gas can be prepared by the following methods.
1. From acids: Metals like zinc, iron, magnesium, etc. are more electropositive than

hydrogen and react with dilute acid to produce hydrogen gas.

Zn + dil. H2 SO4 → ZnSO4 + H2↑
Mg + H2 SO4 → MgSO4 + H2↑

176 Modern Graded Science Class 9

Metals like silver, copper, etc. are less electropositive than hydrogen. So these metals do not
produce hydrogen gas from dilute acids.

2. From alkalis: Hydrogen gas can be obtained from the action of metals like zinc, aluminum,
etc. with boiling caustic soda.

Zn + 2NaOH → Na2ZnO2 + H2↑

(sodium zincate)

2Al + 2NaOH + 2H2O → 2NaAlO2 + 3H2↑
(sodium aluminate)

3. From water: At ordinary temperature, highly active metals like sodium, potassium,
calcium, etc. react with water to liberate hydrogen gas.

2Na + 2H2O → 2NaOH + H2↑
2K + 2H2O → 2KOH + H2↑
B. Metals like magnesium, zinc, alumonium, etc. react with boiling water (steam); they react

together to produce hydrogen gas.

Mg + H2O → MgO + H2↑

Laboratory preparation of hydrogen gas

Principle: When granulated zinc reacts with dilute sulphuric acid or dilute hydrochloric
acid, they react together to form hydrogen gas. The principle reaction is as follows.

Zinc + Dilute Sulphuric acid → Zinc sulphate + Hydrogen.

Zn + dil. H2SO4 → ZnSO4 + H2↑
OR

Zinc + Dilute Hydrochloric acid → Zinc Chloride + Hydrogen

Zn + dil. 2H Cl →ZnCl2 + H2↑

Requirements:

Apparatus: Woulfe's bottle, delivery tube, cork, beehive shelf, gas jar, thistle funnel,
pneumatic trough.

Chemicals: dilute sulphuric acid(H2SO4)/dilute hydrochloric acid(HCl), granulated zinc
(Zn).

Take a few grains of granulated zinc in Woulfe's bottle fitted with a thistle funnel and a delivery
tube with corks. Put the next end of the delivery tube under water in a pneumatic trough
having a bee-hive shelf. Invert a gas jar completely filled with water over the bee-hive shelf
and let the end of the delivery tube into it. Now, pour dilute sulphuric acid through the thistle
funnel till it covers the pieces of zinc and the lower end of the funnel dips in the acid.

Some Gases 177

Fig. 11.2 laboratory preparation of hydrogen gas

A brisk action sets in and a gas evolves. The evolved gas is collected in the gas jar by the
downward displacement of water. The zinc sulphate is left in the Woulfe's bottle in the form of
solution. The collected gas is hydrogen.

Precautions

1. Impure zinc should be used If the reaction between zinc and acid is very slow, we
instead of pure zinc because the should add a little copper sulphate solution to accelerate
reaction between the pure zinc the chemical reaction between them. Here, copper
and the dilute sulphuric acid is sulphate acts as a catalyst.
very slow whereas impurities
present in the zinc increase the
rate of reaction.

2. Concentrated sulphuric acid
should not be used because it produces sulphur dioxide instead of hydrogen gas.

Zn + conc. 2H2SO4 → ZnSO4 + 2H2O + SO2↑

3. The end of the thistle funnel must be under the acid in the Woulfe's bottle.

4. The apparatus should be made airtight.

5. The bee-hive shelf should be under the water level in the pneumatic trough.

6. The experiment should be conducted far away from fire.

Test of hydrogen

Hydrogen burns in air with a faint pale-blue flame. When a lighted splinter is introduced to
the mouth of the gas jar, the gas burns with pop sound and the splinter gets extinguished.
Thus, we can conclude that the produced gas is hydrogen.

Manufacture of hydrogen

Hydrogen gas is used for many purposes. So, it is manufactured in large scale. Usually, the
following two methods are used for the manufacture of hydrogen:

178 Modern Graded Science Class 9

1. By the electrolysis of water: For electrolysis, a small amount of dilute acid is poured into
a voltameter containing water to make a
strong electrolyte. In the voltameter, iron is
used as cathode while the nickel-plated iron
acts as anode. An asbestos diaphragm
separates these two electrodes from each
other. This diaphragm prevents the mixing of iron cathode
hydrogen gas and oxygen gas. When an
electric current is passed, hydrogen is formed
at the cathode and oxygen at the anode. voltameter
2H2O Electricity 2H2+ O2
Fig. 11.3 electrolysis of water

2. From methane-steam process: When a mixture of steam and methane is passed over a
heated nickel catalyst at 700°C-1000C°, and compressed to 3-25 atmosphere in reactor,
hydrogen gas is manufactured. Methane for this process is obtained as a by-product of
petroleum industry.

This is the safest and cheapest method of Hydrogen is the only element whose isotopes have
manufacture of hydrogen gas. different names. The types of hydrogen are Protium
(P), Deuterium (D) and Tritium (T).
Nascent hydrogen
The ordinary isotope of hydrogen with no neutron is
The atomic form of hydrogen produced at the called protium. Its nuclear symbol is 11H. Deuterium
time of a chemical reaction is known as nascent has two neutrons with its mass number 2. Its nuclear
hydrogen. The name nascent means newly symbol is 12H. Similarly, tritium has three neutrons
born. Nascent hydrogen is very unstable. So, with its mass number 3. Its nuclear symbol is 13H. The
the produced nascent hydrogen immediately atoms of these isotopes of hydrogen have one electron to
combine themselves to produce hydrogen balance the charge of one proton.
molecule which is less reactive.
Heavy water, D2O, also called deuterium oxide, in
Zn + H2SO4 → ZnSO4 + 2H which both common hydrogen isotope, protium atoms
H + H → H2↑ (nascent hydrogen) have been replaced with hydrogen isotope deuterium. It
is present naturally in water, but in only small amount
or less than 1 part in 5,000. Heavy water is used in
certain types of nuclear reactors, where it acts as a
neutron moderator.

Physical properties of hydrogen

1. Hydrogen is a colourless, tasteless and odourless gas. 179
2. It is almost insoluble in water.
3. It is the lightest gas known. 22.4 litre of this gas weighs only 2 grams.
4. It is a combustible gas but it does not support burning.
5. It is neutral to litmus.
6. It liquefies at -253°C and solidifies at -259°C.

Some Gases

Chemical properties of hydrogen

1. Hydrogen burns in air or in oxygen with a blue flame and it produces water.
2H2 + O2 ∆ 2H2O
2. Metals like sodium, potassium and calcium burn in hydrogen forming corresponding

unstable metallic hydrides.
2Na + H2 300oC 2NaH (Sodium hydride)
Ca + H2 300oC CaH2 (Calcium hydride)
3. When dry hydrogen is passed over heated oxides of iron, copper, lead, etc. it reduces the

metallic oxides to their metals.
Fe2O3 + 3H2 ∆ 2Fe + 3H2O

CuO + H2 ∆ Cu + H2O

PbO + H2 ∆ Pb + H2O

The removal of oxygen from the compound during reaction is called reduction reaction

and hydrogen which brings about reduction is called reducing agent. The compound

from which oxygen is removed is said to have reduced. But the oxides of calcium, zinc
and magnesium are not reduced by hydrogen.

4. Hydrogen reacts with chlorine, bromine, iodine and fluorine (halogens) in different
conditions to form their acids.
darkness
H2 + F2 sunlight 2HF (Hydrofluoric acid)
H2 + Cl2 400°C 2HCl (Hydrochloric acid)
H2 + Br2 400°C 2HBr (Hydrobromic acid)
H2 + I2 2HI (Hydroidic acid)

5. Hydrogen combines with nitrogen under 200-900 atmospheric pressure at 500°C
temperature and in the presence of a catalyst, iron with molybdenum as promotor to give
ammonia.
N2 + 3H2 500ºC, Fe, Mo 2NH3 (Ammonia)
200 - 900 atm

6. Hydrogen reacts with burning coal and forms methane. Similarly hydrogen reacts with
sulphur on heating to produce hydrogen sulphide.

2H2 + C → CH4 (Methane)
H2 + S ∆ H2S (Hydrogen sulphide)
7. When hydrogen is passed in vegetable oil in the presence of catalyst nickel (Ni) at 200°C

temperature and at 8-10 atmospheric pressure, the oil solidifies and vegetable ghee is

180 Modern Graded Science Class 9

formed. The process is called hydrogenation.

H2 + Vegetable oil 200oC, Ni Vegetable ghee.
8-10 atm
Uses of hydrogen

1. It is used in manufacturing of ammonia and fertilizers.
2. It is used in manufacturing of vanaspati ghee by the process of hydrogenation.
3. It is used as a reducing agent as it reduces metallic oxides into corresponding metals.
4. Hydrogen is used for the production of oxy-hydrogen for welding and cutting metals.
5. The liquid form of hydrogen is used as fuel in rockets.

S me Reasonable Facts

1. Zinc is mostly used for the laboratory preparation of hydrogen gas rather than other
metals because of the following facts:

a. Metals like sodium and potassium react violently with acid.

b. Calcium and magnesium are very expensive in comparison to zinc.

c. Aluminium forms a protective coating of Al2O3 whereas iron reacts with acid
very slowly and requires more heat.

2. Many metals displace the hydrogen from acids because metals are more electropositive
than hydrogen.

3. Hydrogen gas is not found in air because it is more reactive and the lightest gas known.

Things To Know

1. Air is a mixture of different gases like nitrogen, oxygen, carbon dioxide, inert gases,
water vapour, etc.

2. Hydrogen is a reactive and lightest gas known, so it is found only in a combined state
and not in air.

3. Hydrogen gas is prepared in the laboratory by treating zinc with dilute sulphuric acid.
4. Hydrogen gas is collected by the downward displacement of water in the laboratory as

it is the lightest gas known and insoluble in water.
5. Hydrogen is a combustible gas but not a supporter of combustion.
6. Hydrogen reacts with nitrogen at suitable conditions to produce ammonia.
7. Hydrogenation is a process in which unsaturated compounds combine with hydrogen

in the presence of a catalyst and saturated compounds are produced.

Some Gases 181

Things To Do

1. Take a balloon filled with H2 gas and transfer it to a gas jar. Test the properties of hydrogen.
2. Draw a labeled diagram of laboratory preparation of hydrogen gas on a sheet of chart

paper and paste on your classroom/study room.

Test Yourself

1. Multiple choice questions (MCQS)
a. Who discovered hydrogen?

A. Lovoiser B. Priestley C. Sheele D. Covendis

b. At what temperature does hydrogen boil?

A. 253o C B. 219o C C. 210o C D. 253o C

c. Which chemicals are used in lab preparation of hydragen?

A. Zn and Ca (OH)2 B. Zn and H2SO4 C. Zn and Mg (OH)2 D. Zn and H2CO3
d. Hydragen gas is collected by

A. downwards displacement of water B. upward displacement of water

C. upward displacement of air D. Downward displacement of air

e. How much gas is found in 2 g of hydrogen at normal temperature and pressure?

A. 2.4 litre B. 12.4 litre C. 22.4 litre D. 32.4 litre

2. Answer the following questions.

a. Why is hydrogen gas collected by downward displacement of water?

b. Give any four precautions to be taken during the laboratory preparation of hydrogen gas.

c. How do you test hydrogen gas?

d. How is hydrogen manufactured from methane?

e. What are the physical properties of hydrogen gas?

f. What is hydrogenation?

g. What do you mean by a reduction reaction and a reducing agent?

h. How is hydrogen manufactured from electrolysis of water?

i. Write the composition of air.

3. Give reasons.

a. Concentrated H2SO4 cannot be used instead of dilute H2SO4 in laboratory preparation
of hydrogen gas.

b. Pure hydrogen gas is not used in balloons.

182 Modern Graded Science Class 9

c. Hydrogen gas is collected by downward displacement of water.
d. Pure zinc is not used in laboratory preparation of hydrogen gas.

4. What happens when:

a. Hydrogen gas is passed over hot ferric oxide.
b. Hydrogen gas combines with chlorine.
c. Iron reacts with hydrochloric acid.
d. Hydrogen burns with oxygen.

5. Complete and balance the chemical equations given below.

a. Zn + NaOH → …………. + H2

b. Mg + ……… → MgO + H2↑
c. Na + H2 ∆
…………………….

d. Fe2O3 + 3H2 → 2Fe + …………………..

6. Diagammatic questions

a. Answer the following questions after observing the figure.

i. Which gas is going to be prepared in the

figure?

ii. What are the mistakes in the arrangement?

Correct the figure.

iii. Write the formula equation for the preparation

of the gas.

iv. What happens when a burning match stick is

held near the gas jar?

v. Give any three uses of the gas.

vi. What happens when concentrated sulphuric acid is used instead of dilute sulphuric

acid in it?

b. Sketch a neat and labelled diagram of the apparatus set for the preparation of that gas

in laboratory that can change wet blue litmus into red.

Isotope : elements having the same atomic number but different atomic masses
Contaminate : make impure
Precaution : the trait of practicing caution in advance
Manufacture : produce something in large amount from raw materials naturally or

artificially

Some Gases 183

B. Oxygen

Facts about oxygen

Symbol O Fig. 11.4 molecular structure of
Atomic number oxygen molecule
Atomic weight 8
Valency 16
Molecular formula
Molecular weight 2
Position in periodic table : Group
O2
: Period 32 amu
Electronic configuration VI A
2
Boiling point
(1s2, 2s2, 2p4)
Freezing point -183° C
-219° C

Discovery

Oxygen was discovered by Scheele in 1773 and he called it 'fire air' or 'vital air'. In 1774,
Priestley discovered the gas by concentrating the sun's rays on red-oxide of mercury with the
help of powerful lens. He called it 'perfect gas' or 'very active gas'. In 1776, the name 'oxygen'
was given by Antony Lavoisier after studying its properties. Althougn Sheele and Priestly
discovered oxygen independently, the credit of discovering of oxygen goes to Priestly.

Occurrence Ozone is a form of oxygen formed
by three oxygen atoms. It is pale
Oxygen is an abundantly occurring element on the earth. blue in its pure state. It is formed
It constitutes about 49% in the form of various oxides of during electrolysis of water by the
the earth's crust. It is occurred in a free state as well as in use of electrical appliances and
a combined state. The atmosphere contains about 20.95% lighting in the sky. It is present
by the volume of free oxygen. In the combined state, it is as a thick layer in the stratosphere
richly found in combination with other elements in water, layer of the atmosphere. It absorbs
sand, acids, alkalis, carbohydrate, fat, wood, etc. In the the harmful UV rays coming from
human body, it occupies about 72%. the sun and prevents to reach the
surface of the earth protecting the
General methods for the preparation of oxygen gas human beings from cancer and
genetic disorders.
Oxygen gas can be prepared by the following methods:

1. From metallic oxides: Metallic oxides such as
mercuric axide silver oxide, etc. give oxygen when
heated.
2HgO Heat
2Hg + O2↑

2Ag2O ∆ 4Ag + O2↑

184 Modern Graded Science Class 9

2. From metal peroxide: When sodium peroxide is treated with water, oxygen is liberated.

2Na2O2 + 2H2O → 4NaOH + O2↑
2K2O2 + 2H2O ∆ 4KOH + H2↑

Laboratory preparation of oxygen gas

Oxygen gas is prepared in a laboratory by the following two methods:

1. Using heat

Principle: When potassium chlorate is heated in the presence of manganese dioxide in
the ratio of 3:1, it decomposes at 250°C into potassium chloride and oxygen.

Potassium chlorate 250oC Potassium chloride + Oxygen
MnO2
2KClO3 250°C 2KCl + 3O2↑
MnO2
Requirements:

Apparatus: hard glass test tube, delivery tube, cork, beehive shelf, gas jar, thistle funnel,

pneumatic trough, bunsen burner, retort stand.

Chemicals: Potassium chlorate (KClO3), Manganese dioxide (MnO2)

Take a mixture of powdered potassium chlorate and manganese dioxide in the ratio of 3:1 in
a hard glass test tube. Fit it with a
cork and a delivery tube and
adjust the test-tube in a stand as
shown in the figure. Insert the
other end of the delivery tube
into a bee-hive shelf. A gas jar is
filled with water and inverted Fig 11.5 laboratory preparation of oxygen gas (by using heat)
over the bee-hive shelf in the pneumatic trough. Now, heat the mixture carefully. When
potassium chlorate is decomposed, oxygen is liberated. The first formed oxygen gas is
contaminated with the air inside the hard glass test tube and the delivery tube and it is allowed
to escape. Then the oxygen is collected in the gas jar by downward displacement of water.

Precautions

1. The manganese dioxide (MnO2) should be pure, otherwise, the particles of carbon along
with MnO2 burn inside the tube and it may cause an explosion.

2. The test-tube should be fixed in its mouth in a slightly slanting position.
3. The apparatus should be made airtight.
4. Heat should be given uniformly and should not be removed suddenly because it causes
low pressure inside the tube and water comes up into the tube whereby an explosion may
take place.
Some Gases 185

2. Without using heat

Principle: Oxygen gas can be prepared in the laboratory by hydrogen peroxide and
manganese dioxide. In this reaction, manganese dioxide is used as a catalyst.

Hydrogen peroxide MnO2 Water+Oxygen

2H2O2 MnO2 2H2O + O2↑

Requirements:
Apparatus: conical flask, thistle funnel, delivery tube, cork, beehive shelf, gas jar, pneumatic
water trough.
Chemicals: Hydrogen peroxide (H2O2), Manganese dioxide (MnO2).

The apparatus is set up as shown
in the figure. Hydrogen peroxide
is poured in manganese dioxide
and water having a conical flask
with the help of a thistle funnel.
Inside the flask, the decomposition
of hydrogen peroxide is taking
place in the presence of manganese
dioxide and oxygen gas is liberated. Fig. 11.6 laboratory preparation of oxygen gas (without using heat)

The produced oxygen gas is collected in the gas jar by downward displacement of water as it is
less soluble in water and light. Here, manganese dioxide acts as a catalyst. It increases the speed
of decomposition of hydrogen peroxide. In this way, oxygen gas is produced without heat.

Test of oxygen gas

1. To test whether the produced gas is oxygen or not, introduce a glowing matchstick in the jar
containing gas. This burns with bright light. It proves that the gas in the gas jar is oxygen.

2. When a burning magnesium ribbon is introduced inside the gas jar, it burns with dazzling
light and produces white ash of magnesium oxide. This proves that the gas jar contains oxygen.

Manufacture of oxygen

Usually, oxygen gas is manufactured by the following two methods:
1. From air: By compression, cooling and sudden expansion of air, the air is changed into its

liquid state and is freed from moisture and carbon dioxide. The liquid air has nitrogen and
oxygen only. The liquid air is allowed to vaporize. Being the boiling point of nitrogen (-196°
C) lower, it escapes first and is separately collected. The left liquid is nearly pure oxygen. It
is then allowed to vaporize up to the boiling point (-183° C), and collected in cylinders.

186 Modern Graded Science Class 9

2. By the electrolysis of water: Oxygen can also be manufactured by the electrolysis of water
as explained under hydrogen gas.

electrolysis

H2SO4
 2H2O 2H2 + O2

Physical properties of oxygen gas

1. Oxygen is a colourless, odourless and tasteless gas.
2. It is neutral to an indicator.

3. It is slightly soluble in water (about 3%), hence fish and many other aquatic animals
can live in water.

4. It is a non-combustible gas but it is a combustion supporter.
5. It is slightly heavier than air.

6. It liquefies at -183° C and solidifies at -219° C.

Chemical properties of oxygen gas

1. Some metals like potassium, sodium, calcium and magnesium burn with a bright flame
in oxygen to produce their oxides.

4K + O2 ∆ 2K2O (Potassium oxide)

2Ca + O2 ∆ 2CaO (Calcium oxide)

2Mg + O2 ∆ 2MgO (Magnesium oxide)

4Na + O2 ∆ 2Na2O (Sodium oxide)

2. Most of the non-metals like carbon, phosphorus and sulphur burn in oxygen to form
their oxides.

4P + 5O2 ∆ 2P2O5 (phosphorus pentoxide)
C + O2 CO2 (Carbon dioxide)
S + O2 ∆
∆ SO2 (Sulphur dioxide)

3. Iron changes into ferric oxide when it is strongly heated in the presence of oxygen.

3Fe + 2O2 ∆ Fe3O4 (Ferric oxide)

But in the presence of moisture, iron reacts with oxygen to form rust. The process of
formation of rust is called rusting.

4Fe + 3O2 + XH2O ∆ 2Fe2O3. XH2O (Rust)

Some Gases 187

4. Organic compounds like methane, ethyl alcohol (C2H5OH), carbohydrate, oil, etc. burn
with oxygen to give carbon dioxide.
CH4 + 2O2 ∆ CO2 + 2H2O + Heat
(methane)
C2H6 + 3O2 ∆ 2CO2 + 2H2O + Heat
(ethane)

5. All the living organisms require oxygen for the oxidation of glucose inside the cell. Oxygen
oxidizes the food or glucose and releases energy during respiration.
C6H12O6 + 6O2 biocatalyst enzyme 6CO2 +6H2O + Energy

Carbon dioxide and water are produced as by-products of respiration.

Uses of oxygen

1. Oxygen is used by living beings for the oxidation of food or glucose to produce energy
during respiration.

2. It is used in hospitals for the artificial respiration of pneumatic patients.
3. Miners, mountaineers, astronauts and divers use oxygen for artificial respiration.
4. Liquid oxygen is used as a fuel for rockets and missiles.
5. For burning petrol, kerosene, coal, wood, etc. oxygen is essential.
6. It is used in oxy-acetylene flame or oxy-hydrogen flame to cut and weld metals. This

flame is made by burning acetylene with oxygen.
7. Oxygen is used to remove carbon or other non-metallic impurities present in iron

during the extraction of steel.

S me Reasonable Facts

1. In the laboratory preparation of oxygen gas using heat the mouth of the hard glass
test tube is slanted down. It is because, if the test tube is not heated after heating once,
the water of the trough rises up through the delivery tube. In this condition, the water
does not fall directly at the botton of the test tube. Thus, it is protected against cracking
due to irregular expansion of the glass.

2. To test oxygen gas, alkaline pyrogallol solution is used because in it oxygen gets
dissolved and the colour of the solution changes into dark brown.

3. When an alkaline pyrogallol solution is added to an oxygen containing jar, the oxygen
gets dissolved in it and the colour of the solution changes into dark brown.

188 Modern Graded Science Class 9

Things To Know

1. In a laboratory, oxygen gas is prepared by heating the potassium chlorate and manganese
dioxide in the ratio 3:1.

2KClO3 ∆ 2KCl + 3O2↑

MnO2

2. Oxygen gas can be prepared in a laboratory without heat. Hydrogen peroxide
decomposes into water and oxygen in the presence of manganese dioxide. Here
manganese dioxide acts as a catalyst.

2H2O2 MnO2 2H2O + O2↑

3. Oxygen is a colourless, odourless and tasteless gas. It is slightly soluble in water. It is a
supporter of combustion. It is neutral to an indicator.

4. Metals burn with oxygen and produce their oxides.

5. Oxygen gas is tested by introducing a magnesium ribbon into the jar containing
oxygen. It gives dazzling light.

Things To Do

Take a few grains of potassium permanganate in a test tube and supply heat to it. During the
course of heating, a gas is produced. Identify the gas produced and study the property of the gas.
2KMnO4 ∆ K2MnO4+MnO2 + O2

Test Yourself

1. Multiple choice questions (MCQS)

a. Which chemicals are used to prepare oxygen gas without heating process?

A. KClO3 B. KClO3 and MnO2 C. H2O2 and MnO2 D. H2O2 and KClO3

b. In the presence of less oxygen coal forms ................ by burning in oxygen.

A. CO2 B. CO C. CO3 D. SO2

c. At what temperature does oxygen liquify ?

A. 210o C B. 183o C C. 219o C D. 190o C

d. What ratio of KClO3 and MnO2 is required in preparation of oxygen?

A. 4 : 1 B. 3 : 1 C. 2 :1 D. 2 : 2

2. Answare the following questions.

a. What is the role of manganese dioxide in the preparation of oxygen gas in
the laboratory?

b. Enlist the physical properties of oxygen?

Some Gases 189

c. What happens when burning charcoal is kept in a gas jar containing oxygen?
d. What is rusting of iron?
e. Define oxy-acetylene flame.
f. What are the uses of oxygen?

g. What happens when the following elements and compounds burn in oxygen? Write
the chemical reaction.

i. Calcium ii. Phosphorus iii. Magnesium iv. Glucose

h. Complete the following chemical reactions:

i. .................... ∆ 2KCl + +3O2↑

ii. 2Ag2O ∆ .................... + O2↑
iii. C6H12O6 + 6O2 biocatalyst .................... + .................... + Energy
iv. .................... electricity
2H2↑ + O2↑

3. Diagrammatice questions
a. How is oxygen gas prepared in a laboratory by the action of heat? Draw a necessary

labeled diagram and mention the chemical equation.

b. An apparatus is fitted for preparing oxygen gas
in the laboratory as shown in the figure. Answer the
questions on the basis of the diagram.

i. Name the chemicals 'A' and 'B'.

ii. Write a balanced chemical equation for the
reaction that takes place in the conical flask.

iii. Write one sure test for the gas collected in the
gas jar.

iv. Why has the gas been collected by downward displacement of water? \

Rust : a red oxide coating on iron caused by the action of oxygen and moisture
Biocatalyst : a biochemical substance like enzyme
Missile : a weapon that is forcibly projected at a targets but is not self-propelled
Weld : a metal joint formed by softening with heat and fusing together
Acetylene : a colourless flammable hydrocarbon gas used chiefly in welding purpose
Dazzling : shining intensely

190 Modern Graded Science Class 9

C. Nitrogen N Fig. 11.7 molecular structure of nitrogen
7
Facts about nitrogen 14

Symbol 3
Atomic number
Atomic weight N2
Valency
Molecular formula 28 amu
Molecular weight VA
Position in periodic table : Group 2

: Period (1s2, 2s2 2p3)
Electronic configuration - 196°C
Boiling point - 210° C
Freezing point

Discovery

Nitrogen was discovered in 1772 by Daniel Rutherford. He named it mephitic air (poisonous
air). In 1776, Lavoisier studied the properties of nitrogen and proved that it does not take part
in combustion and respiration. Due to this reason, he named it 'azota' (Greek a = no; zoe =
life). The name nitrogen was suggested by Chaptal in 1790 because nitrogen is a constituent
of 'nitre' (KNO3/NaNO3).

Occurrence

Nitrogen is one of the inactive gases, however, it has a significant role to play in the life of
plants, animals and used for industrial purpose. It is found in both free and combined states.
It occurs in the free state in air and occupies 78.07% by volume. In combination, it is found
in ammonia, saltpeter [potassium nitrate or sodium nitrate (nitre)], etc. It is also found in
proteins, enzymes, RNA (Ribo Nucleic Acid), DNA (Deoxyribo Nucleic Acid), amino acids
and the products of vegetables and animals.

General method of preparation of nitrogen gas

Nitrogen gas can be prepared by burning Fig. 11.8 preparation of nitrogen gas by
phosphrus in air. A crucible with phosphorus burning phosphrus
is taken and allowed to float on water. Then
phosphrus is burnt and the crucible is covered
with a gas jar as shown in the figure. The oxygen
of the gas jar is consumed during the burning
of phosphorus and phosphorus pentoxide is
formed. This oxide dissolves in water and the

Some Gases 191

level of water rises in the jar. The gas left in the jar is nitrogen.

4P + 5O2 ∆ 2P2O5 (Phosphorus pentaoxide)

Laboratory preparation of nitrogen gas

Principle: When the solution of sodium nitrite and ammonium chloride is heated they react
together and liberate nitrogen gas.

sodium nitrate + ammonium chloride sodium chloride + water + nitrogen

NaNO2 + NH4Cl ∆ NaCl + 2H2O + N2↑

Requirements:

Apparatus: round bottom flask, delivery tube, bunsen burner, cork, beehive shelf, gas jar,
pneumatic trough, wire gauze, tripod stand.

Chemicals: ammonium nitrite (NaNO2), ammonium chloride (NH4Cl).

Take a mixture of ammonium chloride
and sodium nitrite in equivalent amount
in a round bottom flask. Some water
is added to make the solution of the
mixture. Then the delivery tube is fitted
with a round bottom flask with the help
of a cork. The apparatus is fitted as shown
Fig. 11.9 laboratory preparation of nitrogen gas
in the diagram. The flask is heated gently with a source of heat. After that the gas is produced
by the reaction of the chemicals and is collected in the gas jar by downward displacement
of water. The nitrogen obtained from nitrogenous compounds (the sodium nitrite and

ammonium chloride) is about pure which is also called chemical nitrogen.

Precautions

1. Ammonium chloride and sodium nitrite should not be used in their solid forms because
ammonium chloride, being a sublimate, sublimes and escapes in the gases form on
heating.

2. The fitted apparatus should be made airtight.
3. Heat should be provided uniformly.
4. Ammonium chloride and sodium nitrite should be taken in equivalent amount.

Test of nitrogen

Nitrogen gas can be tested in the following ways:

1. To test whether the produced gas is nitrogen or not, introduce a burning splinter in
a jar containing nitrogen gas. The splinter gets extinguished. This proves that the jar

192 Modern Graded Science Class 9

contains nitrogen gas because nitrogen is non-combustible and non-supporter of
combustion.
2. While a burning magnesium ribbon is inserted in a jar containing this gas, if it burns
continuously and forms magnesium nitride (light yellow coloured ash), it proves that the
gas is nitrogen.

3Mg + N2 → Mg3N2 (Magnesium nitride)

Manufacture of nitrogen

From liquid air: Nitrogen gas is manufactured in a large scale by the fractional distillation of
liquid air (for details, see manufacture of oxygen gas).

Physical properties of nitrogen gas

1. Nitrogen is a colourless, odourless and tasteless gas.
2. It is slightly lighter than air. Its density is 14 whereas the density of air is 14.4.
3. It is very slightly soluble in water.
4. It liquefies at – 196°C and solidifies at –210°C.
5. It is neither combustible nor a supporter of combustion.
6. It is neutral to litmus.

Chemical properties of nitrogen gas

1. Reactive metals like magnesium, calcium, aluminium, etc. unite with the gas, when they
are burnt in nitrogen and form the respective nitrides.

3Mg +N2 ∆ Mg3 N2 (Magnesium nitride)
Ca3N2 (Calcium nitride)
3Ca +N2 ∆ 2AlN (Aluminium nitride)

2Al + N2 ∆

2. When a mixture of nitrogen and hydrogen is heated at about 500°C, at the atmospheric
pressure of 200-900 atmosphere and in the presence of a catalyst, finely divided iron and
molybdenum, ammonia gas is liberated.

N2 + 3H2 520000°-C90/F0ea, tMmo. 2NH3↑

3. When a mixture of eqaul volume of nitrogen and oxygen is heated at 2000°C-3000°C in
an electric arc, they combine together and give nitric oxide. Nitric oxide is also produced
during lightning.

N2 + O2 electric spark 2NO (Nitric oxide)

Uses of nitrogen gas

1. Nitrogen gas is used for manufacturing nitric acid, ammonia and calcium cyanamides.
2. It is used in filling electric bulb to produce an inert atmosphere in order to prevent the
oxidation of tungsten filament.
Some Gases 193

3. It is used for manufacturing of a nitrogen rich fertilizer such as ammonium sulphate and
ammonium nitrate.

4. It is used for replacing fuels in fuel tanks of aeroplanes to prevent the formation of an
explosive mixture of fuel and air.

5. It is used for manufacturing explosives like nitroglycerine.
6. Liquid nitrogen is used as a cooling agent.
7. It is used in high temperature thermometer to reduce evaporation.

S me Reasonable Facts

1. Nitrogen gas is an inactive gas because it is a diatomic. There exists a triple covalent bond
between the two nitrogen atoms at ordinary conditions.

2. Nitrogen gas can be collected by downward displacement of water because it does not
dissolve in water.

3. Nitrogen gas is used for replacing fuels in fuel tanks of aeroplanes to prevent the formation
of an explosive mixture of fuel and air.

Things To Know

1. Nitrogen gas is prepared in the laboratory by heating the mixture of ammonium chloride
and sodium nitrite in the ratio 1:1.

NH4Cl + NaNO2 ∆ NaCl + N2↑ + 2H2O

2. Nitrogen is a colourless, odourless and tasteless gas.

3. It is neither combustible nor a supporter of combustion. But magnesium burns with
nitrogen to give magnesium nitride.

3Mg + N2 ∆ Mg3N2

4. Nitrogen combines with hydrogen in the ratio of 1:3 to form ammonia gas.
N2 + 3H2 ∆ 2NH3

5. Nitrogen combines with oxygen to form oxide of nitrogen.

N2 + O2 ∆ 2NO

6. Nitrogen gas is used for manufacturing ammonia, nitric acid, fertilizers, etc.

7. Nitrogen is used in electric bulbs and aeroplanes.

194 Modern Graded Science Class 9

Things To Do

Draw a colourful chart of laboratory preparation of nitrogen gas with the equation and paste
it on your classroom.

Test Yourself

1. Multiple choice questions (MCQS)

a. Who discovered nitrogen gas?

A. I.A. Chaptal B. D. Rutherford C. Lavoisier D. Sheele

b. What is the purpose of burning phosphorus in the preparation of nitrogen in air?

A. to absorb oxygen B. to absorb water

C. to absorb carbondioxide D. to absorb neon

c. The chemicals needed to form nitrogen in a laboratory are...

A. NaNO2 and NH4Cl B. NaNO3 and NH4Cl

C. NaNO2 and NH4NO3 D. NaNO3 and NH4NO3

d. There is .......... bond in betweem N and N atom in the molecule of nitrogen.

A. single B. double C. triple D. ionic

2. Answer the following questions.

a. What is the position of nitrogen in the periodic table?

b. Which chemical is used for reacting with NH4Cl to produce nitrogen gas in the
laboratory?

c. Give any three precautions which should be noted for the preparation of nitrogen in
the laboratory.

d. How will you prepare nitrogen gas in industrial scale?

e. What are the physical properties of nitrogen?

f. What happens when a burning magnesium ribbon is inserted in a jar containing
nitrogen gas? Also write the chemical equation.

g. What is chemical nitrogen?

h. What happens when nitrogen combines with oxygen at a high temperature?

i. What are the uses of nitrogen?

j. Complete the following chemical reactions.

i. 3Mg + N2 ∆ …………….. .

ii. 4P + 5O2 ∆ ……………….. .

iii. ……. + ………… Ca3N2 Some Gases 195

iv. 2Cu + ……….. → CuO A
v. N2 + ……………. → 2NH3
3. Diagrammatic questions
a. Draw a neat and labeled diagram for the preparation of

nitrogen gas in the laboratory. Mention the chemical
equation.
b. The apparatus is set for the preparation of nitrogen
gas in a laboratary. In it two major mistakes are made.
i. Identify the mistakes made in the setting.
ii. Name A.
iii. What will happen if burning magnesium is taken
in the gas formed?

iv. Resketch the diagram correctly and label its major parts.

Inflammable : that can be set on fire easily
Apparatus : set of instruments used especially in scientific experiments
Miner : a person who works in an underground mine
Crucible : a device made of a mixture of graphite and clay usually that is used to heat

things in it at a very high temperature.
Electric arc : electric pressure

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196 Modern Graded Science Class 9