Oasis Science and Technology 9

11p+ 12p+ 13p+ 14p+
12n0 12n0 14n0 14n0

Sodium Magnesium Aluminium Silicon

15p+ 16p+ 17p+ 18p+
16n0 16n0 18n0 22n0

Phosphorus Sulphur Chlorine Argon

19p+ 20p+
20n0 20n0

Potassium Calcium

8.11 Valence Shell and Valence Electrons

The outermost shell of an atom from where loss or gain of electrons takes place is called
valence shell and the total number of electrons which are present in valence shell (outer
shell) are called valence electrons. For example, valence electrons in sodium, magnesium and
chlorine are 1, 2 and 7 respectively. Valence electrons determine the valency of an atom.

8.12 Valency

The combining capacity of an element or a radical with another element or radical to form a
compound or molecule is called valency. In the past, valency was defined as the total number of
hydrogen atoms which are combined with an element during chemical combination. For example:

In HCl, the valency number of chlorine is one.
In H2O, the valency number of oxygen is two.
In NH3, the valency number of nitrogen is three.
In CH4, the valency number of carbon is four.
But all compounds do not have hydrogen atoms. So, this concept has been modified and a new
concept is put forth. According to the new concept, "The total number of electrons lost, gained

valency /ˈveɪl(ə)nsi/ - the combining capacity of an atom with another atom to form a molecule

CHEMISTRY Oasis School Science - 9 143

or shared by an atom during chemical combination is called valency." Valency of sodium is
one because it loses one electron and valency of oxygen is two because it gains two electrons.
The valency of chlorine is one as it gains one electron from other elements. Similarly, the
valency of carbon is four as it shares four electrons during chemical combination.

Examples,

1. Find out the valency of aluminium in AlCl3.

In AlCl3, three atoms of chlorine combine with one atom of aluminium. So, valency
of aluminium is three.

2. Find out the valency of hydrogen and phosphate in H3PO4.
In H3PO4, three atoms of hydrogen combine with one phosphate radical. So, the

valency of phosphate is three and that of hydrogen is one.

8.13 General Idea to Find Out Valency of Some Elements

1. On the basis of the modern periodic table, the valency of an element is equal to the
number of group from the first to the fourth group. For group first, second, third and
fourth, the valency is one, two, three and four respectively.

Group I II III IV
Valency 1234

2. Valency of elements in group fifth, sixth, seventh and eighth is three, two, one and zero
respectively.

Group V VI VII VIII
Valency 3210

3. Valency of zero group elements like He, Ne, Ar, Kr, Xe and Rn is zero.

4. Valency of a radical is equal to the number of charges present in it. For example:

Valency of Na+ = 1

Valency of CO3- - = 2

Valency of Mg++ = 2

Valency of PO - - - = 3
4

5. In case of transition elements, they have two incomplete outer shells. So, the electrons of
these two shells participate in bonding by showing variable valencies. For example:

Cuprous Cu+ = 1 Cupric Cu++ = 2
Stannous Sn++ = 2 Stannic Sn++++ = 4
Mercurous Hg+ = 1 Mercuric Hg++ = 2

Aurous Au+ = 1 Auric Au+++ = 3

Note: The lower valency is ended by suffix –'ous' and higher valency is ended by suffix – 'ic'.

transition elements /trænˈzɪʃ(ə)nˈelɪm(ə)nts/ - one of the group of elements located in the centre of the periodic table

144 Oasis School Science - 9 CHEMISTRY

8.14 Sub-shell
Each and every main shell contains one or more Fact File-2

than one sub-shells which are denoted by s, p, In the compound MgCl2, 2 atoms of

d and f. Chlorine combine with one atom of

The main shells along with their sub-shells are Magnesium. Hence, the valency of
listed in the given box. Magnesium is two.

Main Shells Sub-shells (orbitals)

K (n=1) s
L (n=2) s and p
M (n=3) s, p and d
N (n=4) s, p, d and f

The maximum number of electrons that can be accommodated by each sub-shell is given
below:

Sub-shells (orbitals) Maximum number of electrons
s (sharp) 2
p (principal) 6
d (diffuse) 10
f (fundamental) 14

The above data shows that the K shell (n = 1) contains only one sub-shell (1s) with maximum
two electrons. L shell (n = 2) contains two sub-shells (2s and 2p) with maximum eight electrons.
M shell (n = 3) contains three sub-shells (3s, 3p and 3d) with maximum eighteen electrons. N
shell (n = 4) contains four sub-shells (4s, 4p, 4d and 4f) with maximum thirty two electrons.
The last electron present in a sub-shell determines the block of an element.

Aufbau principle

This principle was given by Wolfgang
Pauli and Niels Bohr in the early 1920s. The
different sub-shells of an atom have different
energy. Electrons always try to enter into
the sub-shell which has less energy.

Aufbau principle states that, "The electrons
in an atom are so distributed that they
occupy shells in the order of their increasing
energy." It means that the shells having low
energy are filled faster than the shells having
high energy in the following sequence.

1s<2s<2p<3s<3p<4s<3d<4p<5s<4d<p
<6s<4f<5d<6p<7s<5f<6d<7p.

CHEMISTRY Oasis School Science - 9 145

Table: Electronics configurations of some elements on the basis of sub-shells (s, p, d, and f)
are given below:

Elements Atomic K L M N O P
6s 6p 6d 6f
No. 1s 2s 2p 3s 3p 3d 4s 4 p 4d 4f 5s 5p 5d 5f

H 11

He 2 2

Li 3 2 1

Be 4 2 2

B 5 22 1

C 6 22 2

N 7 22 3

O 8 22 4

F 9 22 5

Ne 10 2 2 6

Na 11 2 2 6 1

Mg 12 2 2 6 2

Al 13 2 2 6 2 1

Si 14 2 2 6 2 2

P 15 2 2 6 2 3

S 16 2 2 6 2 4

Cl 17 2 2 6 2 5

Ar 18 2 2 6 2 6

K 19 2 2 6 2 6 1

Ca 20 2 2 6 2 6 2

Sc 21 2 2 6 2 6 1 2

Cr 24 2 2 6 2 6 5 1

Fe 26 2 2 6 2 6 6 2

Ni 28 2 2 6 2 6 8 2

Cu 29 2 2 6 2 6 10 1

Ag 47 2 2 6 2 6 10 2 6 10 1

8.15 Radicals

Radicals are charged atoms or group of atoms having a common charge which act as a single
unit during a chemical reaction. They have either positive charge or negative charge. Radicals
are charged particles. So, they are highly reactive and least stable. Hence, they do not occur
in free form and make different types of compounds. On the basis of electric charge, radicals
are of two types:

1.r adicaEl l/ˈeræctdrɪokpl/ osit-ivea crhaadrgiecdaaltsomoorrbgarosuipcorfaatdoimcsahlasving a common charge

146 Oasis School Science - 9 CHEMISTRY

The atoms or group of atoms which have positive charge in them are called electropositive
radicals or basic radicals. Some examples of electropositive radicals with their valencies
are given below:

Radicals having Radicals having Radicals having Radicals having
valency 1 valency 2 valency 3 valency 4
(Monovalent) (Bivalent) (Trivalent) (Tetravalent)
Hydrogen (H+) Beryllium (Be++) Boron (B+++)
Lithium (Li+) Magnesium (Mg++) Aluminium (Al+++) Stannic (Sn++++)
Sodium (Na+) Calcium (Ca++) Ferric (Fe+++)
Potassium (K+) Strontium (Sr++) Auric (Au+++) Plumbic (Pb++++)

Rubedium (Rb+) Barium (Ba++) Chromium (Cr+++) Silicon (Si++++)
Caesium (Cs+) Cupric (Cu++) Manganic (Mn+++)
Cuprous (Cu+) Mercuric (Hg++)
Mercurous (Hg+) Stannous (Sn++)
Ammonium (NH4+) Zinc (Zn++)
Aurous (Au+) Nickel (Ni++)
Manganous (Mn++)

2. Electronegative radicals or acidic radicals

The atom or group of atoms which have negative charge in them are called electronegative
radicals or acidic radicals. Some examples of electronegative radicals are given below:

Radicals having valency 1 Radicals having valency 2 Radicals having valency 3
(Monovalent) (Bivalent) (Trivalent)

_ __ ___

Fluoride (F ) Oxide (O ) Nitride (N )

_ __ ___

Chloride (Cl ) Sulphide (S ) Phosphate (PO4 )

_ __ ___

Bromide (Br ) Sulphite (SO3 ) Phosphide (P )
__
Iodide (I –) ___
Carbonate (CO3 )
_ __ Phosphite (PO3 )

Nitrate (NO3 ) Sulphate (SO4 )
_ __

Nitrite (NO2 ) Zincate (ZnO2 )
_ __ __

Cyanide (CN ) Silicate (SiO3 )
__
_
Peroxide (O2 )
Hydroxide (OH ) __

_ Dichromate (Cr2O7 )
__ __
Chlorate (ClO3 )
_ Thiosulphate (S2O3 )

Bisulphate (HSO4 )
_

Bicarbonate (HCO3 )
_

Metaluminate (AlO2 )

bond /bɒnd/ - the force by which atoms are held together in a compound Oasis School Science - 9 147
CHEMISTRY

Reasonable fact-1

Radicals contain charge, why?
Radicals are atoms or group of atoms having a common charge which are formed by
losing (donating) or gaining of electron/s to attain the duplet or octet state. Hence, they
contain positive charge due to loss of electron/s and negative charge due to gain of elec-
tron/s. Therefore, radicals contain charge.

8.16 Inert Gases

The elements which have eight electrons in their valence shell (except helium) and do not take
part in the chemical reactions are called inert gases. They are kept in the zero group or VIIIA
of the modern periodic table. Inert gases have complete octet or duplet and show zero valency.
Therefore, they are chemically inert and occur in atomic form in gaseous state. The inert gases
with their symbol, atomic number and electronic configuration are given below:

S.N. Name of inert Symbol Atomic Electronic configuration
L MNO
gases number K P

1. Helium He 22
2. Neon Ne
3. Argon Ar 10 2 8
4. Krypton Kr
5. Xenon Xe 18 28 8
6. Radon Rn
36 28 18 8

54 28 18 18 8

86 28 18 32 18 8

On the basis of the above electronic configuration, atomic structures of some inert gases are
given below.

2p+ 10p+ 18p+
2n0 10n0 22n0

Helium Neon Argon

Fig. 8.5

8.17 Ions

Atoms are electrically neutral as they have equal number of protons and electrons having
opposite charges. When an atom loses or gains electron/s, it becomes positively or negatively
charged. Such type of charged atoms are called ions. So ions are positively or negatively
charged atoms. They are formed by loss or gain of electron/s. Examples: Na+, Mg++, Al+++, K+, Ca+ +,

148 Oasis School Science - 9 CHEMISTRY

Fe++, Cu+, Zn++, F–, Cl-, O– –, N– – –, Br–, I–, etc. When an atom loses electron/s from the valence shell,
it becomes positively charged and when an atom gains electron/s, it becomes negatively charged.

8.18 Duplet and Octet State

Helium is the first member of inert gases. It has only two electrons and according to the 2n²
rule these two electrons are present in the first shell (K-shell). So, the arrangement of two
electrons in the K-shell of an atom is called duplet state. Since helium atom has duplet state, it
does not take part in chemical reaction and remains in atomic form.

Except helium, the other five elements have eight electrons in their valence shell. In other
words, they are in stable electronic configuration. The state of having eight electrons in valence
shell (last shell) of an atom is called octet state. The presence of two electrons in helium (He)
and eight electrons in Ne, Ar, Kr, Xe, and Rn is the main cause of stability of these elements.
Hence, they have zero combining capacity (valency).

The elements which have more or less than eight electrons in their last shell are chemically
unstable and they always try to achieve this condition which is called octet rule. So, the
tendency of elements by which they try to maintain eight electrons in their valence shell (last
shell) either by transferring or sharing of electrons is called octet rule. Similarly, some elements
like H, Li, Be try to maintain two electrons in the shell K (last shell) either by transferring or
sharing of electron which is called duplet rule.

8.19 Chemical Bond

Inert gases have eight electrons in their valence shell except helium. Due to stable electronic
configuration, inert gases are stable. Other elements which do not have duplet or octet state
are unstable. Metals have one, two or three electrons in their valence shell whereas active
non-metals have five, six or seven electrons in their valence shell. Metals try to lose electrons
from the valence shell/s to be stable whereas non-metals try to gain one, two or three electrons
to attain stable electronic configuration. However, some elements share electron/s to obtain
stable electronic configuration. Thus, losing, gaining or sharing of electrons by an atom to
obtain stable electronic configuration is the main cause of chemical reaction.

When an atom loses electrons, it gains positive charge and when it gains electrons it acquires
negative charge. In between these opposite charges, there is a force of attraction which is called
chemical bond. So, the force of attraction by which atoms are held together in a molecule is
called chemical bond. For example, in CH4 one atom of carbon and four atoms of hydrogen
are held together by the bond.

H

Chemical bond

CH
H

H

Fig 8.6 Methane molecule

CHEMISTRY Oasis School Science - 9 149

There are various types of chemical bonds but in this unit we will discuss electrovalent and
covalent bonds only.

a. Electrovalent bond or Ionic bond
The chemical bond which is formed by the transfer of electron/s from the valence shell

of metal to the valence shell of non-metal is called electrovalent bond. The compounds
which are formed by the transfer of electron/s from metal to non-metal are called
electrovalent compounds. They contain electrovalent bond/s. For example, NaCl, KCl,
CaCl2, MgCl2, CaO, etc.
During electrovalent bonding, metals lose their electrons and acquire positive charge, i.e.
cation. Similarly, non-metals gain electrons and acquire negative charge, i.e. anion. In
between these opposite charges, there is a force of attraction which is called electrostatic
force of attraction.
Characteristics of electrovalent or ionic compounds

1. Electrovalent or ionic compounds are generally found in solid state.

2. They have high melting and boiling points.

3. They conduct electricity in molten state or aqueous solution.

4. They dissolve in water.

5. They contain metal atom/s in their molecule.

Formation of Sodium chloride (NaCl)
Sodium chloride is an electrovalent or ionic compound. It is formed by the transfer of one
electron from sodium atom to chlorine atom. In sodium chloride, sodium is a metal and
chlorine is a non-metal. Sodium atom has one electron in its valence shell, whereas chlorine
has seven electrons in its valence shell. Sodium atom donates its one electron to the valence
shell of chlorine. Hence, sodium gains positive charge and chlorine gains negative charge.
There is a force of attraction between two opposite charges which is called electrovalent bond.
This bond keeps Na+ and Cl– together in the form of a molecule, i.e. NaCl.

Sodium atom Chlorine atom Na++ Cl– → NaCl
Sodium chloride molecule

Fig. 8.7 Formation of Sodium chloride molecule

Formation of Magnesium chloride (MgCl2)

Magnesium chloride is an electrovalent compound. It is formed by the transfer of two electrons
from one magnesium atom to two chlorine atoms. In magnesium chloride, magnesium is a

150 Oasis School Science - 9 CHEMISTRY

metal and chlorine is a non-metal. Magnesium atom has two electrons in its valence shell,
whereas chlorine atom has seven electrons in its valence shell. During chemical combination,
magnesium donates its two electrons to each of chlorine atom. As a result, magnesium acquires
two positive charges and two chlorine atoms gain one negative charge each. Now, a chemical
bond is formed in between one magnesium and two chlorine atoms.



17p+ 12p+ 17p+
18n0 12n0 18n0

Chlorine atom Magnesium atom Chlorine atom

Mg+++ 2Cl– → MgCl2

Fig. 8.8 Formation of Magnesium chloride molecule

b. Covalent bond

The chemical bond formed by the sharing of electron pair/s in between two or more than

two non-metal atoms is called covalent bond. It is represented by a line (–) in between the

bonded atoms. When one pair of electrons is shared, it is called single covalent bond and

represented by only one line (–). Similarly, when two pairs and three pairs of electrons

are shared, they are called double and triple covalent bonds respectively. Double covalent

bond is represented by two lines (=) and triple covalent bond is represented by three lines (≡).

For example

HH HH

H CCH CC H CC H

HH HH Ethyne
Ethane Ethene (Triple bond)
(Single bond)
(Double bond)

The compounds which have covalent bonds are called covalent compounds. These
compounds are formed by sharing of electron pair/s between non-metallic atoms.
Examples: Carbon dioxide (CO2), Ammonia (NH3), Water (H2O), Methane (CH4), etc.
They have generally low melting and boiling point. There are two types of molecules in
covalent compounds. They are:

i. Homonuclear molecules
The molecules which have the same type of sharing atoms are called homonuclear

molecules. For example, H2, N2, O2, Cl2, Br2, etc.

ii. Heteronuclear molecules
The molecules which have different types of sharing atoms are called heteronuclear

molecules. For example, CO2, H2O, NH3, CH4, HCl, etc.

CHEMISTRY Oasis School Science - 9 151

Characteristics of covalent compounds
1. Covalent compounds are found in solid, liquid and gaseous state.
2. They have low melting and boiling points.
3. They do not conduct electricity.
4. They are insoluble in water.
5. They do not contain metal atom in their molecules.

Formation of Hydrochloric acid (HCl)

Hydrochloric acid is formed by sharing of one pair of electrons between hydrogen atom
and chlorine atom. In hydrochloric acid, there is one hydrogen atom and one chlorine atom.
Hydrogen atom has only one electron in its shell. So, it requires one more electron to get duplet
state. Similarly, chlorine has seven electrons in its valence shell. So it requires one more electron
to get octet state. Therefore, hydrogen atom and chlorine atom share one pair of electrons to get
stable electronic configuration. As a result, hydrogen chloride (HCl) molecule is formed.

1 p + 17p+ 17p+ 1p+
18n0 18n0 0n0
0n0
+

Hydrogen atom Chlorine atom Hydrogen chloride molecule

Fig. 8.9 Formation of Hydrogen chloride molecule

Formation of Methane (CH4)

In methane molecule, there is one carbon atom and four hydrogen atoms. Carbon atom has
four electrons in its valence shell. So it shares with four electrons to get octet state. Similarly,
hydrogen atom has only one electron in its valence shell. So, it shares with one electron to
get duplet state. Here, four electrons of carbon combine with one electron each from four
hydrogen atoms to form CH4.

Methane molecule is formed by sharing of four pairs of electrons between one carbon atom
and four hydrogen atoms.

101npp+0+ 101npp0++ 101npp0++ H
I
6p+ H–C–H
6n0 Or, I
H

(Methane)

1p+
0n0

152 Oasis School Science - 9 Fig. 8.10 Molecular structure of Methane
CHEMISTRY

Differences between Electrovalent and Covalent Compounds

S.N. Electrovalent Compounds S.N. Covalent Compounds

1. The compounds formed by 1. The compounds formed by sharing of

transfer of electrons between electron pairs between atoms are called

atoms are called electrovalent covalent compounds.

compounds.

2. They can conduct electricity in 2. They cannot conduct electricity.
molten/solution state.

3. They have high melting and 3. They have low melting and boiling point.
boiling point.

4. They contain metal atoms in 4. They do not contain metal atoms in their
their molecules. Examples: NaCl, molecules. Examples: H2O, NH3, CH4, etc.
MgCl2, AlCl3, etc.

8.20 Molecular Formula

The molecular formula of a molecule is the symbolic representation of the molecule of an element
or a compound in molecular form. It represents the actual number of atoms of different elements
in a molecule. For example, the molecular formula of sodium chloride is NaCl and that of water
is H2O. It shows that one molecule of sodium chloride (NaCl) consists of one atom of sodium
(Na) and one atom of chlorine (Cl). Similarly, one molecule of water (H2O) consists of two atoms
of hydrogen (2H) and one atom of oxygen. (O) In case of inert gases, i.e. He, Ne, Ar, Kr, Xe and
Rn, the single atom represents atom as well as molecule because they are monoatomic molecules.
Elements like hydrogen, nitrogen, oxygen, chlorine, bromine and iodine have two atoms in their
molecule, viz. H2, N2, O2, Cl2, Br2 and I2 respectively. So, they are called diatomic molecules.

S.N. Symbol S.N. Molecular Formula

1. The symbol of an element is the 1. The molecular formula is the symbolic

abbreviation of the full name of representation of a molecule of an

that element. element or a compound.

2. It represents one atom of an 2. It represents one molecule of an element

element. Examples: H, Na, K, etc. or a compound. Example: H2, NaCl,
H2SO4, etc.

Note
• Cl2 indicates one molecule of chlorine and 2Cl represents two atoms of chlorine.
• NH3 indicates one molecule of ammonia whereas 2NH3 indicates 2 molecules of ammonia.

8.21 Molecular Weight

The molecular weight of a molecule is the sum of atomic weight of all atoms of the molecule.
It is calculated by adding the atomic weight of the atoms present in a molecule.

CHEMISTRY Oasis School Science - 9 153

i) Molecular weight of Water (H2O) = H × 2 + O × 1
= 1 × 2 + 16 × 1 = 18

ii) Molecular weight of Calcium carbonate (CaCO3) = Ca × 1 + C × 1 + O × 3
= 40 × 1 + 12 × 1 + 16 × 3 = 100

iii) Molecular weight of Nitric acid (HNO3) = H × 1 + N × 1 + O × 3
= 1 × 1 + 14 × 1 + 16 × 3

= 1 + 14 + 48 = 63

iv) Molecular weight of Aluminium sulphate [ Al2(SO4)3] Al × 2 + 3 (S × 1 + O × 4)
=

= 26 × 2 + 3 (32 × 1 + 16 × 4)
= 54 + 3(96) = 342

8.22 Methods of Writing Molecular Formula

We should follow the following steps to write the correct molecular formula of a molecule.

1. Write the symbol of basic (positive) and acidic (negative) radicals side by side.

2. Write the valency of each radical on upper right corner of each.

3. Exchange the valency of these radicals. Take HCF if it is necessary.

4. Combine radicals with exchanged valency.

5. If radicals have different atoms, it is enclosed within brackets.

For example,

1. Sodium chloride 2. Calcium chloride 3. Aluminium chloride
Na Cl Ca Cl
21 Al Cl
11 31

Na 1 Cl 1 Ca1 Cl 2 Al 1 Cl 3
CaCl2 AlCl3
NaCl
5. Boron oxide
4. Carbon tetrachloride BO 6. Ammonium sulphate
C Cl 32
41 NH 4 SO 4

1 2

C1 Cl4 B2 O 3 NH 4 SO 4
CCl4 B2O3
1 2
7. Calcium sulphate
Ca SO4 (NH4)2SO4
22
8. Aluminium nitrate 9. Calcium bicarbonate
Ca HCO 3
Al NO 3
21
3 1

Ca2CaSO4SO2 4 Al NO 3 Ca HCO 3
154 Oasis School Science - 9
3 1 2 1

Al(NO3)3 Ca(HCO3)2

CHEMISTRY

8.23 Information Obtained form Molecular Formula

1. Molecular formula represents one molecule of a substance.
2. It indicates the total number of atoms of the same or different element/s in each molecule.
3. It indicates percentage composition of each element present in the compound.
4. The valency or combining capacity of each element can be found from the molecular

formula.
Example: In ammonia molecule (NH3), the valency of nitrogen is three and that of

hydrogen is one.

5. We can calculate molecular weight from the molecular formula. For example,

The molecular weight of ammonia (NH3) = N × 1 + H × 3
= 14 × 1 + 1 × 3 = 14 + 3 = 17

SUMMARY

• The branch of science which deals with the study of matter, its composition and properties
is called Chemistry.

• The smallest particle of an element that can take part in chemical reaction without division
is called atom.

• Compound is a chemical substance formed by the combination of two or more elements
in a definite proportion by weight.

• The smallest particle of an element or a compound which is capable of independent
existence is called molecule.

• The total number of protons present in the nucleus of an atom is called atomic number. It
is denoted by Z.

• The systematic distribution of electrons in different shells of an atom is called electronic
configuration.

• The combining capacity of an element or a radical with another element or a radical to
form a compound or molecule is called valency.

• The arrangement of two electrons in the K-shell of an atom is called duplet state.

• The state of having eight electrons in valence shell (last shell) of an atom is called octet state.

• The tendency of elements by which they try to maintain eight electrons in their valence
shell (last shell) either by transferring or sharing of electrons is called octet rule.

• The force of attraction by which atoms are held together in a molecule is called chemical bond.

• The chemical bond which is formed by the transfer of electron/s from the valence shell of
metal to the valence shell of non-metal is called electrovalent bond.

• The chemical bond formed by the sharing of electron pair/s in between two or more than
two non-metal atoms is called covalent bond.

• The molecular formula of a molecule is the symbolic representation of the molecule of an
element or a compound in molecular form.

CHEMISTRY Oasis School Science - 9 155

Exercise

Group-A

1. What is an element?

2. What is a compound?

3. What is electronic configuration?

4. Write down the maximum number of electrons that can accommodate in the given
sub-shells.

i) s ii) p iii) d iv) f

5. What are sub-atomic particles?

6. What are proton and electron?

7. What is an atom? Give one example.

8. What is a molecule? Give one example.

9. What is a diatomic molecule? Give one example.

10. What is a radical? How many types of radicals are there?

11. What are electronegative radicals?

12. Name one electropositive and electronegative ion each.

13. What is a chemical bond?

14. What are covalent compounds?

15. What type of compounds are called electrovalent compounds?

16. What are valence electrons? How many valence electrons are found in an oxygen atom?

17. What is molecular formula? Give one example.

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

19. What is octet state? Give one example.

20. What is octet rule?

21. What is 2n2 rule?

22. What type of gases are called inert gases?

23. Write down the valency of the given radicals.

i) Nitrate ii) Phosphate iii) Chloride

iv) Silicate v) Sulphate vi) Carbonate

vii) Hydroxide viii) Ammonimum ix) Bisulphate

x) Bromide xi) Nitrite xii) Sulphite

xiii) Iodide xiv) Chlorate xv) Oxide

156 Oasis School Science - 9 CHEMISTRY

Group-B

1. `Differentiate between atom and molecule in any two points.

2. Sodium is called an element and water is called a compound, why?

3. Write any two differences between elements and compounds.

4. Differentiate between protons and electrons in any two points.

5. What is meant by duplet state and octet state?

6. Differentiate between atom and radical with one example of each.

7. Sodium atom cannot exist freely in nature but argon atom can. Give reason.

8. Write any two differences between electrovalent bond and covalent bond.

9. Differentiate between electropositive radicals and electronegative radicals with one
example of each.

10. Atoms are electrically neutral. Justify this statement.

11. Write any two differences between electrovalent compounds and covalent compounds.

12. Sodium chloride is called electrovalent compound but ammonia is called covalent
compound, why?

13. The valency of neon is zero but that of sodium is one, why?

14. The elements of group 1 are reactive but those of group 18 are inert, why?

15. Electrovalent compounds are electrolysed but covalent compounds are not electrolysed,
why?

16. Identify electropositive and electronegative radicals from the following:

i) Sulphate ii) Ammonium iii) Nitrate iv) Carbonate

Group-C
1. What is a chemical bond? How is it formed? Write with an example.
2. How is covalent bond formed? Write in brief.
3. Calculate the maximum number of electrons present in shell L and N by using 2n2 rule.
4. Describe the formation of sodium diloride with a neat figure.
5. Describe the formation of carbon dioxide with a neat figure.

Group-D

1. Draw the molecular structure of the following molecules:

i) Sodium chloride ii) Carbon dioxide

2. In a part of the periodic table given below, some elements are denoted by symbols A, B,
C, D and E. Answer the following questions on this basis:

CHEMISTRY Oasis School Science - 9 157

Group IA IIA IIIA IVA VA VIA VIIA 0

Period (1) (2) (13) (14) (15) (16) (17) (18)

2 Beryllium D

3 B A Sulphur C Argon

4 Potassium E

i) Name the elements 'B' and 'D' and state their valency also.

ii) Name the compound formed by the combination of elements 'A' and 'C'. Also
write down the type of bond present between them.

3. Answer the following questions on the basis of given figures.

8p+ 11p+ 17p+
8n0 12n0 18n0

AB C

i) What is the valency of B and C? Why?

ii) Name the compound formed by the combination of the elements and A and B.
Also write down the type of the bond with reason.

4. Name any four elements having variable valency. Also, write their valencies.

5. Explain the method of writing molecular formula with examples.

6. Calculate the molecular weight of the given compounds.

i) Calcium carbonate ii) Ammonium nitrate.



158 Oasis School Science - 9 CHEMISTRY

9UNIT Estimated teaching periods

Theory 5

Practical 0

Chemical

Reaction Avogadro

Objectives

After completing the study of this unit, students will be able to:

• define physical change and chemical change and differentiate between them.
• describe the method of writing chemical equation and write chemical

equation.

9.1 Introduction

Various types of changes occur in our surroundings as well as within the body of living beings.
For example, rusting of iron, burning of wood, formation of cloud, digestion of food, burning
of fuel, growing of plants and animals, melting of ice, etc. Some of these changes are physical
changes whereas others are chemical changes. In a physical change, new substances are not
formed but new substances having different properties are formed in a chemical change.
Chemical change is a permanent change that occurs as a result of chemical reaction.

9.2 Physical Change

The temporary and reversible change in which no new substance is formed is called a physical
change. In a physical change, the chemical composition of the substance does not change but
the physical properties like colour, odour, taste, shape, size, etc. are changed. Physical change
is called a reversible change because it occurs in forward as well as backward directions.

Ice on heating Water on heating Steam
on cooling on cooling

Characteristics of physical change
i. Physical change is a temporary change.
ii. In this change, no new substance is formed.
iii. It is usually a reversible change.
iv. In this change, only physical properties like colour, odour, taste, physical state, etc.

are changed.

reversible /rɪˈvɜːsəbl/ - that can be changed to its original form
CHEMISTRY
Oasis School Science - 9 159

Some examples of physical change are given below:

i. Melting of ice or wax

ii. Magnetizing and demagnetizing of iron

iii. Vaporization of water

iv. Making different objects from wood, soil and paper

v. Formation of a solution

When water is heated, it changes into vapour. It is a physical change. In this change,
new substance is not formed. In this change, water, i.e. liquid state changes into gaseous
state but water and water vapour have similar chemical properties as both contain H2O
molecules as shown in the given figure.

H HH H heat HH HH
OO O O

Water (liquid state) Water vapour

9.3 Chemical Change Fig. 9.1

The permanent change in which new substance having different properties is formed is called
a chemical change. In a chemical change, the physical properties as well as the chemical
composition of the substance are changed. Generally, the product formed during this change
cannot be reversed back to get reactants. So chemical change is also called an irreversible change.

Properties of chemical change
i. Chemical change is a permanent change.
ii. In a chemical change, new substance is formed.
iii. It is an irreversible change.
iv. In this change, physical properties and chemical composition are changed.
Some examples of chemical change are given below:
i. Burning of paper and fuel
ii. Rusting of iron
iii. Electrolysis of water
iv. Burning of candle
v. Digestion of food

When water (H2O) is electrolysed, it changes into hydrogen (H2) and oxygen (O2) gas. It
is a chemical change because two new substances having different chemical properties
are formed. This process is shown in the given figure.

H HH H electrolysis HH HH +OO
OO Fig. 9.2 Oxygen (O2)
Hydrogen (H2)
Water (H2O)

irreversible /ɪrɪˈvɜːsəbl/ - that cannot be changed back into its original form

160 Oasis School Science - 9 CHEMISTRY

Differences between Physical and Chemical change

S.N. Physical change S.N. Chemical change
1. It is a temporary change. 1. It is a permanent change.
2. In this change, no new substance 2. In this change, new substance is formed.

is formed. 3. It is an irreversible change.
3. It is a reversible change.

Examples: Melting of ice, Examples: Rusting of iron, burning of
magnetizing of iron, etc. paper, etc.

9.4 Chemical Reaction

The combination, decomposition or exchange that takes place in the molecules of matter
during a chemical change is called chemical reaction. For example, when hydrogen gas burns
in oxygen, it forms water. It can be expressed as follows:

Hydrogen + Oxygen Burn Water

2H2 + O2 Burn 2H2O

Similarly, when calcium carbonate is heated, it decomposes into calcium oxide and carbon
dioxide gas. It can be expressed as follows:

Calcium carbonate ∆ Calcium oxide + Carbon dioxide

(Reactant) (Products)

CaCO3 ∆ CaO + CO2

A chemical reaction is expressed in word equation and chemical equation.

9.5 Word Equation

The chemical reaction expressed by writing the full names of reactants and products is called
a word equation. In a word equation, there are some demerits which are given below:

1. It takes more space and more time.
2. We cannot count the total number of atoms and molecules of reactants and products.

3. We cannot balance the equation.

Examples:

1. Hydrogen + Oxygen burn Water
2. Sodium + Chlorine contact Sodium chloride
3. Calcium carbonate Calcium oxide + Carbon dioxide
heat

4. Potassium chlorate heat Potassium chloride + Oxygen

9.6 Chemical Equation

The chemical reaction expressed by writing symbols and molecular formulae of reactants and
products is called chemical equation. A chemical equation is more informative than a word
equation.

CHEMISTRY Oasis School Science - 9 161

Examples: ∆ 2H2O
1. 2H2 + O2 2NaCl
∆ CaO + CO2
2. 2Na + Cl2 ∆ 2KCl + 3O2

3. CaCO3

4. 2KClO3

9.7 Reactants and Products

The chemical substances which take part in a chemical reaction are called reactants. The chemical
substances which are produced after chemical reaction are called products. Reactants are written
on the left side of the arrow whereas products are written on the right side of the arrow.

For example,

HCl + NaOH NaCl + H2O
Reactants Products

9.8 Unbalanced or Skeleton Chemical Equation

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

For example: ∆ KCl + O2

KClO3

In above equation, the number of oxygen atoms in reactant and product sides is not equal
so it is called unbalanced chemical equation. Some more examples of unbalanced chemical
equations are as follows:

Zn + HCl ZnCl2 + H2

KOH + H2SO4 MnO2 K2SO4 + H2O
H2O2 H2O + O2

HCl + Ca(OH)2 CaCl2 + H2O

9.9 Balanced Chemical Equation

The chemical equation written by balancing the total number of atoms of each element in
reactants and products is called balanced chemical equation. In this chemical equation, the
number of atoms of each element is equal in reactants and products. It gives more information
than the unbalanced chemical equation. Balanced chemical equation proves the law of
conservation of mass. Above unbalanced chemical equations can be balanced as follows:

2KClO3 ∆ 2KCl + 3O2
Zn + 2HCl MnO2 ZnCl2 + H2
2KOH + H2SO4 K2SO4 + 2H2O
MnO2
2H2O 2 2H2O + O2

2HCl + Ca(OH)2 CaCl2 + 2H2O

162 Oasis School Science - 9 CHEMISTRY

In above equations, the number of atoms of different elements in reactant and product sides is
equal. So, they are called balanced chemical equations.

Fact File-1

According to Dalton's atomic theory, atoms can neither be created nor be destroyed and
remain intact before and after the chemical reaction. So, every chemical equation should
be balanced.

9.10 Methods of Writing Balanced Chemical Equations

Following points should be remembered while balancing the chemical equations:

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

For example: Nitrogen + Hydrogen Ammonia

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

For example: N2 + H2 NH3

3. The number of atoms of each element are balanced by using suitable coefficient without
changing the molecular formulae of reactants and products.

N2 + 3H2 2NH3

4. The number of atoms in the biggest molecule should be balanced before balancing the
number of hydrogen and oxygen atoms.

This method of balancing chemical equation is called hit and trial method.

9.11 Some More Examples of Balanced Chemical Equation

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

Unbalanced chemical equation: KClO3 ∆ KCl + O2

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

2. Word equation : Sodium + Chlorine Sodium chloride

Unbalanced formula equation : Na + Cl2 NaCl
Balanced formula equation: 2Na + Cl2 2 NaCl
3. Word equation : Potassium + Oxygen Potassium oxide

Unbalanced formula equation : K + O2 K2O
Balanced formula equation: 4K + O2 2K2O
4. Word equation : Magnesium + Oxygen
Magnesium oxide

Unbalanced formula equation : Mg + O2 MgO
2 MgO
Balanced formula equation: 2 Mg + O2 Calcium oxide + Carbon dioxide
5. Word equation: Calcium carbonate heat CaO + CO2
Balanced formula equation: CaCO3 ∆

CHEMISTRY Oasis School Science - 9 163

6. Word equation : Zinc + Hydrochloric acid Zinc chloride + Hydrogen

Unbalanced formula equation: Zn + HCl ZnCl2 + H2

Balanced formula equation: Zn + 2HCl ZnCl2 + H2

7. Word equation : Calcium chloride + Silver nitrate Calcium nitrate + Silver chloride

Unbalanced formula equation: CaCl2 + AgNO3 Ca (NO3)2 + AgCl

Balanced formula equation: CaCl2 + 2AgNO3 Ca(NO3)2 + 2 AgCl
8. Word equation : Hydrogen peroxide catalyst Water + Oxygen

Unbalanced formula equation: H2O2 MnO2 H2O + O2

Balanced formula equation: 2H2O2 MnO2 2H2O + O2

9.12 Information Obtained from a Balanced Chemical Equation

Following pieces of information can be obtained from a balanced chemical equation.
1. The names and symbols of reactants and products
2. The total number of atoms or molecules of reactants and products
3. The ratio of molecular weight of reactant and product molecules
4. The type of chemical reaction

9.13 Limitation of a Balanced Chemical Equation

A balanced chemical equation cannot provid information about

1. The physical state of reactants and products
2. Concentration of reactants
3. Conditions required for the reaction like heat, light, pressure, catalyst, etc.

4. The duration of the chemical reaction
5. The rate of chemical reaction.

9.14 Modification of Chemical Equation

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

for gas and 'aq' for aqueous solution.

2. Concentration of reactants are denoted by dil. for dilute and conc. for concentrated.
3. The conditions like temperature, pressure, light, catalyst, etc. are written above or below

the arrow.
4. A double ways arrow ( ⇋ ) is used to represent a reversible reaction and a single way

arrow (→ ) is used to represent an irreversible reaction.
5. For endothermic reaction, positive sign (+ ∆) and for exothermic reaction, negative sign

(–∆) is used over the arrow.

catalyst /ˈkætəlɪst/ - a chemical substance that increases or decreases the rate of chemical reaction

164 Oasis School Science - 9 CHEMISTRY

9.15 Exothermic Reaction

The chemical reaction which releases heat during the chemical change is called exothermic
reaction. For example,

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

CH4 + 2O2

Zn + 2HCl

9.16 Endothermic Reaction

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

For example,

N2 + O2 Heat 2NO
Heat
2KClO3 2KCl + 3O2↑
Heat CaO + CO2↑
CaCO3 Heat NaCl + 2H2O + N2↑
Heat 2Na + Cl2 ↑
NaNO2 + NH4Cl

2NaCl

9.17 Reversible and Irreversible Reaction

The chemical reaction which occurs in forward as well as backward directions is called a
reversible reaction. Generally, a double ways arrow (⇋) is used to show reversible reaction. For
example, when nitrogen combines with hydrogen, it gives ammonia. But after the application
of low pressure and high temperature, ammonia gets decomposed into hydrogen and nitrogen.

Nitrogen + Hydrogen Ammonia

N2 + 3H2 2NH3

The chemical reaction which occurs only in one direction (forward direction) is called
irreversible reaction. Generally, a single way arrow (→) is used to show such reactions. For
example, when calcium carbonate is decomposed, it gives calcium oxide and carbon dioxide.
But calcium carbonate cannot be obtained by combining calcium oxide and carbon dioxide.

Calcium carbonate Heat Calcium oxide + Carbon dioxide

CaCO3 ∆ CaO + CO2

9.18 Catalyst

A catalyst is a chemical substance which increases or decreases the rate of chemical
reaction remaining itself chemically unchanged. There are two types of catalysts.

CHEMISTRY Oasis School Science - 9 165

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

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

2H2O2 MnO2 2H2O + O2↑
(catalyst)
b. Negative catalyst

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

2H2O2 Glycerine 2H2O + O2↑

Characteristics of a catalyst

i) The mass and chemical nature of a catalyst does not change during a chemical
change.

ii) A catalyst does not initiate a chemical change.

iii) A catalyst increases or decreases the rate of a chemical reaction.

SUMMARY

• The temporary and reversible change in which no new substance is formed
is called a physical change.

• The permanent change in which new substance is formed is called a chemical
change. In a chemical change, the physical properties as well as the chemical
composition of the substance are changed.

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

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

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

• The chemical substances which take part in a chemical reaction are called
reactants.

• The chemical substances that are formed after a chemical change are called products.

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

• The chemical equation written by balancing the total number of atoms of
each element in reactants and products is called balanced chemical equation.

• The chemical reaction which occurs in forward as well as backward directions
is called reversible reaction.

• The chemical reaction which occurs only in one direction (forward direction)
is called irreversible reaction.

166 Oasis School Science - 9 CHEMISTRY

Exercise

Group-A

1. What is a physical change?

2. What is a chemical change?

3. What is a chemical reaction?

4. What are 'reactants' and 'products'?

5. What is a word equation?

6. What is a chemical (or formula) equation?

7. Write any two pieces information which can be obtained from a balanced chemical
equation.

8. Write any two limitations of a balanced chemical equation?

9. What is a catalyst?

10. What type of catalyst is called a positive catalyst?

11. Define negative catalyst with one example.

12. Write any two characteristics of catalysts.

Group-B

1. `Melting of ice is called physical change but rusting of iron is called chemical change,why?

2. Chemical equation should be balanced, why?

3. Formula equation is more informative than word equation. Give reason.

4. Write any two differences between physical change and chemical change.

5. Differentiate between reactants and products in any two points.

6. Write any two differences between word equation and formula equation.

7. Differentiate between exothermic reaction and endothermic reaction in any two points.

Group-C

1. What do you mean by 'reactants'? Write with examples. Write down the balanced chem-
ical equation of the reaction between sulphuric acid and ammonium hydroxide.

2. Write down the information which can be obtained from a balanced chemical equa-
tion. Write a balanced equation of the chemical reaction between iron (III) chloride and
Ammonium hydroxide.

3. What are the limitations of a balanced chemical equation? Mention any three points.

4. Write down the method of writing a balanced chemical equation.

5. We can get more information from a formula equation than from a word equation. Justify
this statement. Write down the balanced chemical equation of the given equation.

Aluminium + Oxygen Aluminium oxide

CHEMISTRY Oasis School Science - 9 167

Group-D

1. Study the given word equation and answer the following equations.
Calcium carbonate Heat Calcium oxide + Carbon oxide

i) Express the above word equation in the form of formula equation.

ii) Mention any two pieces of information that can be obtained from the given
equation.

iii) Name the factor that brings out chemical change in the above reaction.

2. Change the given word equations into balanced chemical equation. Different between
reactants and products.

i) Sulphuric acid + Zinc Zinc sulphate + Hydrogen

ii) Iron + Oxygen Ferric oxide

3. Complete and balance the given chemical equations:

i) Na2CO3 + .............. NaCl + .....................
ii) HCl + ............ CaCl2 + H2O

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

i) When magnesium reacts with nitric acid

ii) When silver nitrate reacts with sodium chloride

5. Correct and rewrite the given chemical equations:

i) NaCO3 + CaCl2 NaCl2 + CaCO3
ii) HNO2 + NaHCO2 NaNO3 + H2O + CaCO2
iii) NaOH + NH3Cl NaCl + NH2 + H2O
iv) CaCl2 + AgNO3 CaCO3 + AgCl

168 Oasis School Science - 9 CHEMISTRY

UNIT 10 Estimated teaching periods

Theory 5

Practical 3

SOLUBILITY

Solution

Objectives

After completing the study of this unit, students will be able to:

• define unsaturated, saturated and super saturated solution and prepare such
solutions.

• define solubility of a substance and explain the relation between solubility
and temperature.

• explain crystallization in brief.

10.1 Introduction

In our surroundings, there are many substances which are intermingled with one another.
In atmosphere, different types of gases and dust particles are mixed together. We prepare
'sherbet' by putting sugar into water. Sugar gradually dissolves into the water. Many solutions
of similar types are used in our daily life, such as tea, medicine, drinks, etc. Not only solutions
but also different types of mixtures are also familiar to us. Rice coats in rice, sand in rice, sand
in lime, dust in atmosphere, etc. are some examples of mixture. So, a mixture is a mass which
is obtained by mixing two or more substances in any proportion by weight. In a mixture, the
mixing components do not lose their identity. The components of a mixture are chemically
unreactive and they may be present in solid, liquid or gaseous state. However the mixture of
solid and gas is not found in nature.

10.2 Types of Mixture

Depending on the nature, there are two types of mixtures. They are:

i) Homogeneous mixture

ii) Heterogeneous mixture

i) Homogeneous mixture

The mixture in which mixing components are distributed uniformly and they cannot be
identified by our naked eyes is called homogeneous mixture. In sugar solution, sugar
particles are distributed throughout the water. In brass, metals are mixed uniformly.
Sugar solution, salt solution, soda water, alcohol in water, water in milk, brass, etc. are
some examples of homogeneous mixture.

solubility /ˈsɒljʊbɪlƏti/- the amount of solute required to form a saturated solution in 100 g of a solvent at a certain temperature

CHEMISTRY Oasis School Science - 9 169

ii) Heterogeneous mixture

The mixture in which mixing components are not distributed uniformly and they can be
identified with our naked eyes is called heterogeneous mixture. Husk in rice, sand in rice,
muddy water, dust in air, smoke in air, etc. are some examples of heterogeneous mixtures.

Homogeneous mixture Fig. 10.1 Heterogeneous mixture

Differences between Homogeneous and Heterogeneous mixture

S.N. Homogeneous mixture S.N. Heterogeneous mixture

1. In this mixture, mixing components 1. In this mixture, mixing components

are distributed uniformly. are not distributed uniformly.

2. We cannot identify all the mixing 2. We can identify all the mixing
components. components.

Examples: Salt solution, brass, etc. Examples: Muddy water, stones in
rice, etc.

10.3 Classification of Mixture

A mixture consists of component particles of different size. On the basis of the size of the
component particles, there are three types of mixture. They are as follows:

i) Solution ii) Colloids iii) Suspension

S.N. Mixture Size of particles
1. Solution 10-7 cm or less
2. Colloid 10–7 to 10–5 cm
3. Suspension 10–5 cm or large

i) Solution Fig. 10.2

Take a glass of water and add some sugar in it. Stir the mixture till the
sugar disappears. Now, we will get a homogeneous mixture of sugar
and water which is called a solution. So, a solution is a homogeneous
mixture of two or more substances which are mixed up to a certain
limit. In a solution, the size of particles is 10–7 cm or less in diameter.
They are so tiny particles that they cannot be seen with our naked
eyes. Solution contains two components, i.e. solute and solvent. Thus,
solution is a homogeneous mixture of solute and solvent.

Solution = Solute + Solvent

170 Oasis School Science - 9 CHEMISTRY

Solute and Solvent

The component of a solution which is present in relatively large amount is called solvent.
Solvent dissolves the solute to make a solution. Water, alcohol, ether, phenol, petrol, etc.
are the examples of solvent. So, a solvent is a substance which dissolves the solute into it
to form a solution. Similarly, the component of a solution which is present in relatively
less amount and gets dissolved into a solvent is called solute. Sugar, salt, copper sulphate,
etc. are the examples of solute. So, a solute is a substance which dissolves into a solvent.

Sugar + Water Sugar solution
(Solute) (Solvent) (Solution)


Differences between Solute and Solvent

S.N. Solute S.N. Solvent
1. It is the component of a solution
It is the component of a solution 1. which is present relatively more in
2. which is present relatively less quantity than solute.
in quantity than a solvent. It dissolves the solute.

It gets dissolved into a solvent. 2. Examples: Water, alcohol, petrol, ether,
etc.
Examples: Sugar, salt, etc.

ii) Colloid

The mixture in which the size of component particles of the mixture Fig. 10.3
ranges between 10–7cm to 10–5cm in diameter is called colloid. When
sulphur is dissolved in alcohol and poured into water, we get turbid
solution. This turbid solution will not settle even if it is kept for
several days. If we filter the turbid solution, filtrate again gives turbid.
This type of solution is called colloidal solution or sol. The particles
of a colloid can pass through the filter paper and cannot be seen
under the simple microscope. Milk, blood, gum, etc. are some
examples of colloids.

iii) Suspension

The heterogeneous mixture in which the diameter of mixing particles
is 10–5cm or large is called suspension. The components of a suspension
contain larger particles so they are visible with naked eyes. The
particles of a suspension can be separated by simple physical methods
like filtration, sedimentation, etc. Sand in water, mud in water, dust
particles in air are some examples of suspensions.

10-7cm or small 10-5cm to 10-7 cm 10-5cm or large Fig. 10.4

(Solution) (Colloid) (Suspension)
Fig. 10.5

CHEMISTRY Oasis School Science - 9 171

11.4 Concentration of a Solution

The amount of a solute which is present in a given amount of a solvent is called concentration
of the solution. If less amount of a solute is present in more amount of a solvent, it is called
less concentrated or dilute solution. Similarly, if the amount of the solute is relatively more, it
is called concentrated solution. Here, concentrated solution is denoted by "conc." and dilute
solution by "dil."

Fig. 10.6(a) Dilute solution of CuSO4 (b) Concentrated solution of CuSO4

"Dilute" and "concentrated" are relative terms and they do not specify the exact concentration
of the solution. If the amount of a solute in a solution is relatively less, it is called dilute
solution and if the amount of a solute in a solution is relatively more, it is called concentrated
solution. In a dilute solution, there is less solute and less density whereas in a concentrated
solution, there is more solute and more density.

Differences between Dilute and Concentrated solution

S.N. Dilute solution S.N. Concentrated solution

1. It contains relatively less amount of a 1. It contains relatively more amount of

solute in given amount of a solvent. a solute in given amount of a solvent.

2. It has less density. 2. It has more density.

10.5 Unsaturated, Saturated and Supersaturated Solution

Take some amount of water and add small amount of copper sulphate in it. On constant
stirring, it gives a solution. In this solution, more amount of copper sulphate is added and
if the added extra solute also dissolves, it is known as unsaturated solution. So, the type of
solution in which further amount of solute can dissolve easily at a certain temperature is
called unsaturated solution.

If we go on adding more amount of copper sulphate to the solution, a condition will appear
at which no more copper sulphate dissolves at that temperature. This solution is called
saturated solution. So, the solution which cannot dissolve any more amount of solute at a
fixed temperature is called saturated solution.

A saturated solution does not show any change in concentration when it is brought in contact
of a solute. This is because concentrated solution remains in equilibrium state. If we increase

172 Oasis School Science - 9 CHEMISTRY

the temperature of the saturated solution, more amount of solute can be dissolved. But when
the solution is cooled down, the extra amount of solute which was dissolved comes out and
sediments at the bottom of the container in the form of crystals. So, the saturated solution
prepared at high temperature which throws out excess solute as a solid on cooling is called
supersaturated solution.

17 g CuSO4 21 g of CuSo4 30 0C 24 g of CuSO4
100 ml water 100 ml water
20 0C 20 0C

A BC

Fig. 10.7

In fig. A, the type of solution is unsaturated because 21 g of CuSO4 can be dissolved in 100 ml
of water at 200C but the solution has only 17g of CuSO4.

In fig. B, the type of solution is saturated as it cannot dissolve any more amount of CuSO4 at
200C. If we add a crystal of CuSO4 in this solution, it does not dissolve at that temperature but
dissolves only when the temperature is increased.

In fig. C, the type of solution is supersaturated. If a crystal of CuSO4 is added to it, the size
of the crystal increases. When the solution is cooled from 300C to 200C, it throws 3 gram, i.e.
(24g–21g) of CuSO4 as solid in the solution.

Differences between Saturated and Unsaturated Solution

S.N. Saturated solution S.N. Unsaturated solution

1. In this solution, no more solute can 1. In this solution, more amount of

dissolve at the given temperature. solute can dissolve easily at the given

temperature.

2. It has more saturation and more 2. It has less saturation and less density.
density.

Differences between Saturated and Supersaturated solution

S.N. Saturated solution S.N. Supersaturated solution

1. In this solution, no more solute can 1. It is the saturated solution prepared

dissolve at the given temperature. at high temperature.

2. It does not change concentration 2. Its concentration decreases when it

when it comes in contact of a solute. comes in contact of a solute.

10.6 Importance of Solution

1. Plants absorb food and minerals from the soil in the form of solution.
2. Aquatic organisms take dissolved oxygen from the water in the form of solution.
3. Most of the medicines are prepared in the form of solution.

saturation /sætʃəˈreɪʃn/ - a degree to which sth is absorbed in sth else

CHEMISTRY Oasis School Science - 9 173

4. The digested food is distributed throughout the body in the form of solution.
5. Many chemical reactions occur in the form of solution.

Activity 1

Objective : To prepare unsaturated, saturated and supersaturated solution

Materials required : Copper sulphate, water, beaker, Bunsen burner, tripod stand, glass
rod, matchbox

Procedure

i. Take 50 ml of water in a beaker. Add some amount of copper sulphate into it and stir
it with a glass rod.

ii. Go on adding the extra amount of copper sulphate and stir the solution till the added
solute stops to dissolve.

iii. Increase the temperature of the solution and add more amount of copper sulphate.

Observation : Initially solute dissolves easily but after some time no more solute can
dissolve. However, after increasing temperature, more amount of solute
gets dissolved.

Conclusion : The first condition is called unsaturated solution, second condition is called
saturated solution and the last condition is called supersaturated solution.

Glass rod CuSO4 Glass rod
Beaker
Beaker
CuSO4

Stand Burner
Burner

Fig. 10.8

10.7 Solubility

The solubility of a substance at a fixed temperature is the amount of solute (in gram) required
to form a saturated solution in 100 grams of a solvent at that temperature. Different substances
have different solubility because they have different physical and chemical properties. The
solubility of a substance can be calculated by using the given formula:

Weight of solute (in gram)
Solubility = Weight of solvent (in gram) × 100

174 Oasis School Science - 9 CHEMISTRY

For example, the solubility of copper sulphate is 25 at 300C. It means that 25 grams of copper
sulphate dissolves in 100 grams of water at 300C to form a saturated solution. Solubility has no
unit as it is the percentage ratio of two similar physical quantities.

The solubility of some common substances at 200C and 300C temperature is given below:

S. No. Substances (Solutes) Solubility at 200C Solubility at 300C

1. Sodium chloride 35 37
2. Sodium nitrate 88 95
3. Copper sulphate 21 24
4. Ammonium chloride 37 41
5. Potassium nitrate 27 37
6. Magnesium chloride 32 34

From above table, it becomes clear that different solutes have different solubility at a fixed
temperature and the same solute has different solubility at different temperatures. It also
shows that the solubility of a solute increases with increase in temperature and vice-versa.

Fact File-1

The solubility of copper sulphate at 200C is 21 means that at 200C, 21 grams of copper
sulphate dissolves into 100 grams of solvent (water) to form a saturated solution.

Reasonable fact-2

Temperature is mentioned along with the solubility of a substance, why?
The solubility of a solute depends on the temperature. It may increase or decrease
with the increase or decrease in temperature but the solubility of a solute is fixed at
a certain temperature. Therefore, temperature is mentioned along with the solubility
of a substance.

Worked out Numerical 1

Find out the solubility of a salt when 44g of it is dissolved in 50 g of water at 20°C temperature.
Solution:

Given, At 200C, 44 g
Weight of salt (solute) = 50 g
Weight of water (solvent) =
?
Solubility =

We know, Solubility = Wt. of solute (in gram) × 100
Wt. of solvent (in gram)

= 44 g
× 100
=
50 g
88

Therefore, the solubility of the salt is 88 at 20°C.

CHEMISTRY Oasis School Science - 9 175

Worked out Numerical 2

The solubility of copper sulphate at 30°C is 24. What amount of water is required to make
a saturated solution of 0.5 kg of copper sulphate?

Solution:

Given, At 300C, Solubility = 24


Weight of solute = 0.5 kg

= 500 g [∵ 1 kg = 1000 g]

Weight of solvent = ?

We know, Solubility = Wt. of solute (in gram) × 100
Wt. of solvent (in gram)

or, 24 = 500
Solvent × 100
or, Solvent = 500
24 × 100
= 2083.33g.

Therefore, the weight of solvent (water) is 2083.33 g.

Worked out Numerical 3

If 20 g saturated solution contains 8 g of the solute at 25°C, calculate the solubility of the
solute.

Solution:

Given, At 250 C

Weight of saturated solution = 20 g

Weight of solute = 8 g

Weight of solvent = Weight of solution – Weight of solute

= 20g – 8g = 12g

Solubility = ?

We know, Solubility = Wt. of solute (in gram)
Wt. of solvent (in gram) × 100

= 8
= 12 × 100

66.66

Therefore, the solubility of the solute is 66.66 at 25°C.

176 Oasis School Science - 9 CHEMISTRY

Worked out Numerical 4

If 20 g saturated solution at 80°C is cooled down to 30°C, how many grams of solute will be
crystallized out. The solubility of the solute at 80°C is 100 and at 30°C is 25.

Solution:

At higher temperature, i.e. 80°C

Let, Weight of solute = xg
20 g
Weight of solution = (20 – x) g
100
∴ Weight of solvent =

Solubility =

Now,

Solubility = Wt. of solute (in gram)
× 100
100 =
Wt. of solvent (in gram)

x
20 – x × 100

x = 10 g

So, Weight of solvent = Weight of solution – Weight of solute

= 20 – 10

= 10 g

At lower temperature, i.e. 30°C

Let, Weight of solute = yg
10 g
Weight of solvent = 25

Solubility =

Now, Solubility = Wt. of solute (in gram)
× 100
25 =
y = Wt. of solvent (in gram)
y
10 × 100

2.5 g

Therefore, the amount of solute that crystallizes on cooling = x – y

= 10 – 2.5

= 7.5 g

∴ When the given solution is cooled from 80°C to 30°C, 7.5 gm crystals are obtained.

CHEMISTRY Oasis School Science - 9 177

10.8 Factors Affecting Solubility

Different factors affect the solubility of a solute. Some of them are described below:

i. Temperature

Temperature is one of the major factors that affect the solubility of a solute. When
temperature is increased, solubility of a substance also increases and vice-versa. When
the temperature of solution is increased, the kinetic energy of solute and solvent
molecules also increases. As a result, intermolecular space among the molecules of the
solvent increases. Hence, more amount of solute gets dissolved at high temperature.

Similarly, when temperature is decreased, kinetic energy of the solute and solvent
molecules is also decreased. As a result, spaces of the solvent molecules decrease. Hence,
solution throws out the excess solute in the form of crystals.

ii. Size of particles

Small sized solute particles dissolve more rapidly than the lump. This is because small
size of solute particles increases the total surface area and decreases the surface area of
individual particle. Hence, tiny particles of a solute can enter into the small spaces.

10.9 Solubility Curve YSolubility
Potassium nitrate
Solubility of most of the substances increases 120 Magnesium chloride
with increase in temperature and decreases with 110 X
decrease in temperature. The best way to show 100 Calcium sulphate
this relation is by using simple graph line. So,
the curve obtained by plotting solubility of a 90 6 8 10 12 14 16 18 20
substance at different temperatures along Y-axis 80 Temperature in 0C
and corresponding temperature along X-axis is 70
called a solubility curve. Different substances 60 Fig. 10.9
have different solubility curve. Potassium nitrate, 50
potassium chloride, sodium chloride, etc. show the
continuous solubility curve without break. In these 40
substances, solubility increases with increase in 30
temperature. 20
10

02 4

Sodium sulphate is an example of a solute whose solubility increases with increase in
temperature up to 32.4°C. After this temperature, solubility decreases with increase in
temperature. So, this type of curve line is called discontinuous curve with break.

Information obtained from solubility curves

1. Solubility curve can be used to find out the solubility of a substance at given temperature.

2. Solubility curve shows the variation of solubility with temperature.

3. Solubility curve can be used to compare the solubility of different substances at different
temperatures.

4. The amount of a solute that crystallized out after cooling the solution can be calculated.
5. The substances which undergo crystallization first can be predicted.

178 Oasis School Science - 9 CHEMISTRY

10.10 Crystalline Solids Fig. 10.10 Crystals of sugar

A crystal is a solid substance having a fixed geometrical shape
bounded by smooth surfaces and sharp edges. Crystals are
arranged in a specific three dimensional pattern. Different
types of solids have crystals of different shape and size. The
solid substance which is made by repeating crystals is called
crystalline solid. It has a fixed geometrical shape and fixed
melting point and boiling point. Sodium chloride, copper
sulphate, sugar, etc. are some examples of crystalline solids.

Characteristics of crystals

1. They have a definite geometrical shape.

2. They have sharp melting point.

3. They are held together by their sharp edges.

4. They are present in pure form.

10.11 Crystallization

The process by which fine crystals are obtained on cooling a hot super saturated solution
slowly is called crystallization. When a saturated solution at higher temperature is cooled
down slowly, it throws the extra solute in the form of crystals and sediments at the bottom of
the container. These crystals are observed under the microscope to find out the shape.

Activity 2

Objective : To get crystals of copper sulphate

Materials required : Copper sulphate, water, glass rod, evaporating beaker, tripod
stand, Bunsen burner

Procedure : Prepare the saturated solu-
tion of copper sulphate
at higher temperature by
keeping solid copper sul-
phate in the solution with
constant stirring. Now, the
solution is allowed to cool
down at room temperature
slowly.

Observation : We will get residue (crys- Fig. 10.11 Crystals of CuSO4
tals) at the bottom of evap-

orating dish. Now, these

crystals are observed under the microscope.

Conclusion : Fine crystals of copper sulphate are obtained.

amorphous /əˈmɔːfəs/ - having no definite shape, form or structure Oasis School Science - 9 179
CHEMISTRY

10.12 Amorphous Solids

The solid substances which do not have a fixed geometrical shape are called amorphous
solids. Glass, rubber, plastic, etc. are some examples of amorphous solids. These solids do not
melt on heating but become soft.

SUMMARY

• A mixture is a mass which is obtained by mixing two or more substances in any
proportion by weight.

• The mixture in which mixing components are distributed uniformly and they
cannot be identified with our naked eyes is called homogeneous mixture.

• The mixture in which mixing components are not distributed uniformly and they
can be identified with our naked eyes is called heterogeneous mixture.

• A solvent is a substance which dissolves the solute into it to form a solution.

• The amount of a solute which is present in a given amount of a solvent is called
concentration of the solution.

• The type of solution in which additional amount of solute can dissolve easily at a
certain temperature is called unsaturated solution.

• The solution which cannot dissolve any more amount of solute at a fixed temperature
is called saturated solution.

• The saturated solution prepared at high temperature which throws out excess
solute as a solid on cooling is called supersaturated solution.

• The solubility of a substance at a fixed temperature is the amount of solute (in gram)
required to form a saturated solution in 100 grams of a solvent at that temperature.

• The curve obtained by plotting solubility of a substance at different temperatures
along Y-axis and corresponding temperature along X-axis is called a solubility
curve.

• A crystal is a solid substance having a fixed geometrical shape bounded by smooth
surfaces and sharp edges.

• The solid substance which is made by repeating crystals is called crystalline solid.

• The process by which fine crystals are obtained on cooling a hot super-saturated
solution slowly is called crystallization.

• The solid substances which do not have a fixed geometrical shape are called
amorphous solids.

180 Oasis School Science - 9 CHEMISTRY

Exercise

Group-A

1. What is a mixture? Give one example.

2. Define homogeneous mixture with one example.

3. What is heterogeneous mixture? Give one example.

4. Separate homogeneous mixture and heterogeneous mixture from the given mixture.

i) Sugar solution ii) Sand and sugar

iii) Alcohol and water iv) Turbid water

5. Define solute and solvent.

6. What is solubility? Write down the formula to calculate the solubility of a substance.

7. What is a solution? Give one example.

8 Define: unsaturated solution and saturated solution.

9. What are suspension and colloid? Give any two examples of each.

10. What is a solubility curve? Write.

11. What is a crystal? Write one method of making big crystals.

12. Define dilute solution and concentrated solution.

13. If some more solute can be dissolved in a solution at 250C, write down the type of
solution.

14. If concentration of solution reduces and crystals appear when solute is added to the
solution at 400C, name the type of solution.

15. What happens when the saturated solution of copper sulphate is heated?

16. What happens when supersaturated solution of sodium chloride is cooled?

17. What type of substance is called a crystalline substance?

18. What type of substance is called an amorphous substance?

19. What is the size of particles dissolved in a solution?

Group-B

1. The mixture of sugar and water is called homogeneous but the mixture of salt and sand
is called heterogeneous, why?

2. Write any two differences between homogeneous mixture and heterogeneous mixture.

3. Differentiate between solute and solvent in any two points.

4. Write any two differences between dilute solution and concentrated solution.

5. Write any two differences between unsaturated and saturated solution.

6. More amount of solute can be dissolved by heating a saturated solution. Give reason.

CHEMISTRY Oasis School Science - 9 181

7. Temperature is mentioned along with the solubility of a substance. Give reason.
8. The solubility of a solute increases on increasing the temperature of a solvent, why?
9. Differentiate between crystalline solids and amorphous solids.

Group-C

1. What is a crystallization? Write down two methods of making big crystals.

2. How is saturated solution formed? Write. What type of substance is called an amorphous
substance? Write with examples.

3. How is the solubility of a solute calculated? Write. What is the size of the particles
dissolved in a solution? Write.

4. Describe in brief the utility of solution in our daily life. What happens when the saturated
solution of copper sulphate is heated?

5. Describe in brief the method of finding the solubility of common salt.

Group-D

1. Study the given solubility curve and answer the following Y
120

questions. 110

Solubility100
Potassium nitrate
i) What is the solubility of sodium chloride at 250C? 90

80
70

ii) What is the relationship between increase in 60 Magnesium chloride
temperature and solubility? 50

40
30

iii) What happens when a saturated solution of 20
sodium chloride at 500C is cooled to 200C? 10 Calcium sulphate

0 2 4 6 8 10 12 14 16 18 20 X
Temperature in 0C

Fig. 10.9

iv) What is the solubility of ammonium chloride at

240C and 650C from the solubility curve?

2. At 200C, the solubility of sodium chloride is 35. How much gram of sodium chloride
is required at that temperature to make a saturated solution in 20 gram of water?
Differentiate between crystal and crystallization.

3. Cooper sulphate is called crystalline substance. Give reason. Differentiate between
suspension and colloids.

4. The solubility of a salt is 36 at 350C. How much salt is required to make a saturated
solution in 144 gram of water at that temperature? Differentiate between solubility and
solubility curve.

182 Oasis School Science - 9 CHEMISTRY

UNIT11 Estimated teaching periods

Theory 5

Practical 3

Some Gases

Henry Cavendish

Objectives

After completing the study of this unit, students will be able to:

• prepare hydrogen, oxygen and nitrogen gases in laboratory.
• explain physical properties, chemical properties and uses of hydrogen,

oxygen and nitrogen gases.

11.1 Introduction

The earth is surrounded by a thick layer of air which is called atmosphere. The air consists
of various gases like nitrogen, oxygen, carbon dioxide, neon, argon, water vapour, etc. About
99% volume of the air is occupied by nitrogen and oxygen. The different gases present in air
with their percentage volume are given below:

S.N. Gases in air Volume (%)
1. Nitrogen (N2) 78.08

2. Oxygen (O2) 20.95
3. Carbon dioxide (CO2) 0.0360
4. Argon (Ar) 0.93
0.002
5. Neon (Ne) 0.000004
0.003945
6. Hydrogen (H2) 0.00005
7. Ozone (O3)
8. Other gases (Ne, He, CH4, N2O)

Many elements exist in gaseous state. Some of them are hydrogen, helium, nitrogen, oxygen,
neon, chlorine, argon, krypton, etc. Gases are very useful for human beings as well as other
organisms. Hydrogen gas is used to fill balloons and prepare vegetable ghee. Oxygen gas is
used by living organisms for breathing. Nitrogen gas is used for making chemical fertilizers.
Carbon dioxide is used by green plants for photosynthesis. It is also used in soft drinks like
coca-cola, beer, soda water, etc.

In this unit, we will study the occurrence, laboratory preparation, industrial preparation,
properties and uses of hydrogen, oxygen and nitrogen.

CHEMISTRY Oasis School Science - 9 183

A. HYDROGEN GAS H 1p+ 1 p +
1
Symbol 1 Fig. 11.1 Molecular structure of
Atomic number 1 hydrogen gas
Atomic mass H2
Valency 2
Molecular formula 1s1
Molecular weight
Electronic configuration

11.2 Discovery

Henry Cavendish, a British scientist, discovered hydrogen gas in 1766 AD and named it
"inflammable gas". Later in 1783 AD, Antony Lavoisier proposed the name hydrogen (Greek
word, hydro-water, genas – produce) because it produces water on burning with oxygen.

11.3 Occurrence

Hydrogen is the lightest, simplest and reactive element. Thus, it does not occur in free state in
nature. In combined state, it is present in different compounds like water, acids, hydrocarbons,
carbohydrates, etc. Large amount of hydrogen is present in the sun and its trace is present in
volcanic gases. The chief source of hydrogen on the earth is water.

11.4 General Methods of Preparing Hydrogen Gas

Hydrogen gas can be prepared by either of the following methods.

i) From water

Hydrogen gas can be prepared by the electrolysis of water.

2H2O Electricity 2H2 + O2

Hydrogen gas can also be prepared by the reaction of active metals like sodium,
potassium with water at ordinary temperature.

2Na + 2H2O 2NaOH + H2
2KOH + H2
2K + 2H2O

ii) From acids
Electropositive metals like iron, zinc, magnesium, etc. react with acids to give hydrogen gas.

Zn + dil. H2SO4 ZnSO4 + H2
Mg + dil. 2HCl MgCl2 + H2

184 Oasis School Science - 9 CHEMISTRY

iii) From alkalis
Metals react with alkalis to give hydrogen gas.

Zn + 2 NaOH Na2ZnO2 + H2
(Sodium zincate)

2 Al + 2 NaOH + 2H2O 2 NaAlO2 + H2
(Sodium aluminate)

11.5 Laboratory Preparation of Hydrogen Gas

Principle

In laboratory, hydrogen gas can be prepared by the reaction of zinc granules with dilute
hydrochloric acid or dilute sulphuric acid.

Zinc + Hydrochloric acid Zinc chloride + Hydrogen

Zn + dil. 2 HCl ZnCl2 + H2

Zinc sulphate + Hydrogen

Zinc + Sulphuric acid ZnSO4 + H2

Zn + dil. H2SO4


Apparatus required

i. Woulfe's bottle ii. Thistle funnel

iii. Gas jar iv. Delivery tube

v. Water trough vi. Beehive shelf vii. Cork

Chemicals required

i. Dilute hydrochloric acid (HCl) or dilute sulphuric acid (H2SO4)
ii. Zinc granules (Zn)



Dilute H2SO4



Delivery tube

Thistle funnel

Cork Hydrogen gas
Gas jar

Woulfe's bottle Beehive shelf
Granulated zinc
Water trough
Water

Fig. 11.2 Laboratory preparation of hydrogen gas (H2)

Procedure
i. Keep some granules of zinc in the Woulfe's bottle and arrange the apparatus as
shown in the figure.

alkali / ˈ æ l k ə l ə ɪ / - the base that dissolves in water

apparatus /ˌæpəˈreɪtəs/ - the tools or other pieces of equipment that are needed for a particular activity

CHEMISTRY Oasis School Science - 9 185

ii. Pour dilute hydrochloric acid or dilute sulphuric acid through the thistle funnel.

When zinc reacts with dilute acid, it gives off hydrogen gas. This gas passes through
delivery tube. The gas is collected in the gas jar by downward displacement of water.

Precautions Reasonable fact-1
i. The apparatus should be made airtight.
ii. We should take granulated impure zinc. Why is it difficult to get hydrogen
iii. We should not use concentrated in air? Give reason.
sulphuric acid.
iv. The lower end of the thistle funnel must It is difficult to get hydrogen gas
be dipped in the acid solution. freely in the air because:
v. The experiment must be conducted a) It is highly reactive element. So,
away from the fire.
it reacts with other elements to
11.6 Test of Hydrogen Gas form compounds.
b) It is the lightest gas known and
flies upward. So, it cannot be
found in the air.

When a burning match stick is introduced into the gas jar containing the gas produced, the gas
burns with pale-blue flame and the matchstick gets extinguished with a pop sound. It proves
that the gas is hydrogen.

Note

• We should use impure granulated zinc to increase the rate of chemical reaction because
pure zinc reacts very slowly with dilute acid.

• Concentrated sulphuric acid produces sulphur dioxide gas (SO2) in place of hydrogen gas.
So, we should use dilute sulphuric acid.

• The experiment should be performed away from the fire; otherwise hydrogen itself burns
and may cause explosion.

• Hydrogen is the lightest known gas. So, it is collected by downward displacement of
water.

• If the reaction occurs slowly, we should use little amount of copper sulphate to accelerate
the reaction.

• Impurities present in the zinc also play a role of positive catalyst in the reaction.

• The lower end of the thistle funnel must be dipped in the solution to protect the hydrogen
from escaping.

• We pass hydrogen gas through water because it is slightly soluble in water.

11.7 Manufacture of Hydrogen Gas

For commercial purpose, hydrogen gas is prepared in the following two ways.

i. From methane – steam reaction
When a mixture of methane gas and water vapour is heated strongly at about 1200°C

temperature and 30 atm. pressure with nickel as catalyst, we get hydrogen gas.

CH4 + H2O 1200ºC/30 atm. CO + 3H2
Ni


186 Oasis School Science - 9 CHEMISTRY

Note

• Catalyst is a chemical substance which either increases or decreases the rate of
chemical reaction, itself remaining chemically unchanged.

ii. By the electrolysis of water

Hydrogen gas can also be manufactured by passing electricity in the acidified water.

Electricity 2H2 + O2
+ H2SO4
2H2O



11.8 Physical Properties of Hydrogen Gas

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

ii. It is the lightest gas.

iii. It is neutral to indicators.

iv. It is slightly soluble in water.

v. It becomes liquid at –253°C and becomes solid at –259°C.

11.9 Chemical Properties of Hydrogen Gas

i. Reaction with air

Hydrogen gas burns with air or oxygen and forms water.

2H2 + O2 ∆ 2H2 O

ii. Reaction with metals

Metals like sodium, potassium, etc. react with hydrogen gas to give corresponding
metallic hydrides.

2 Na + H2 300º C 2NaH (Sodium hydride)

Ca + H2 300º C Ca H2 (Calcium hydride)

iii. Reaction with halogens

Halogens like fluorine, chlorine, bromine react with hydrogen to give corresponding

acids. In dark

H2 + F2 2HF (Hydrofluoric acid)

H2 + Cl2 Light 2HCl (Hydrochloric acid)

H2 + Br2 400º C 2HBr (Hydrobromic acid)

H2 + l2 440ºC 2Hl (Hydrogen iodide)

iv. Reaction with metallic oxides

Hydrogen reduces the metallic oxides into free metals when it is passed through
heated metallic oxides. This property of hydrogen is called reducing property of
hydrogen.

CHEMISTRY Oasis School Science - 9 187

CuO + H2 ∆ Cu + H2O
PbO + H2 ∆ Pb + H2O
Fe3O4 + 4H2 ∆
ZnO + H2 ∆ 3Fe + 4H2O
Zn + H2O

Hydrogen gas Lead oxide Excess hydrogen
(PbO) and water vapour

Burner
Stand

Fig. 11.3

v. Reaction with nitrogen

Hydrogen reacts with nitrogen at 5000C temperature and 200 atm. pressure in the
presence of iron catalyst and molybdenum promoter to give ammonia gas.

3H2 + N2 500ºC, 200 atm. 2NH3

Fe/Mo

vi. Reaction with unsaturated hydrocarbons

Hydrogen combines with unsaturated hydrocarbons to give saturated
hydrocarbons.

C2H4 + H2 Ni
C2H6
(Ethene)
(Ethane)

vii. Hydrogenetaion: When hydrogen gas is passed into vegetable oil at about 2000C
and 8-10-atmospheric pressure in the presence of nickel it forms vegetable ghee.
This process is called hydrogenation.

H2 + vegetable oil 8 – 10 atm. Vegetable ghee
Ni / 200ºC

11.10 Uses of Hydrogen Gas

i. In the past, hydrogen gas was used to fill the balloons and ships because it is
the lightest gas.

ii. It is used in the manufacturing of vegetable ghee by the process of hydrogenation.
iii. It is used as a fuel in rockets.
iv. It is used in the manufacturing of ammonia gas and chemical fertilizers.
v. It is used as a reducing agent.
vi. It is used for making oxyhydrogen flame for cutting and welding of metals.

188 Oasis School Science - 9 CHEMISTRY

B. OXYGEN GAS

Symbol O 8p+ 8p+
Atomic number 8 8no 8n0
Atomic mass 16
Valency 2 Fig. 11.4 Molecular structure of oxygen
Molecular formula O2
Molecular weight 32
Electronic configuration 1s2, 2s2, 2p4

11.11 Discovery

Oxygen was discovered by a Swedish chemist Scheele in 1773 AD and named as fire air or
vital air. An English scientist Joseph Priestly called it perfect gas or very active gas. In 1776 AD,
French scientist Antony Lavoisier gave the name "oxygen".

11.12 Occurrence

Oxygen occurs free as well as in combined state in nature. It constitutes about 48% of the
earth's crust and about 21% of the atmosphere. In combined state, it is present as a constituent
of different compounds like acid, base, salt, fat, carbohydrate, protein, sand, etc. About 72% of
the human body is occupied by oxygen. More than 49% of the materials found on the earth's
surface are the compounds of oxygen.

11.13 General Methods for the Preparation of Oxygen Gas

i. From metallic oxides

Metallic oxides like mercuric oxide, lead oxide, etc. give oxygen gas after heating.

2 HgO ∆ 2 Hg + O2
2 PbO ∆ 2 Pb + O2

ii. From metal peroxides

Oxygen gas can be prepared when metal peroxides like sodium peroxide,
potassium peroxide, etc. are treated with water.

2Na2O2 + 2H2O 4 NaOH + O2

2K2O2 + 2H2O 4 KOH + O2

11.14 Laboratory Preparation of Oxygen Gas

There are two methods of preparing oxygen gas in the laboratory. They are:

CHEMISTRY Oasis School Science - 9 189

a. By using heat
b. Without using heat

a. By using heat

Principle

When 3:1 ratio of potassium chlorate and manganese dioxide is heated at about 250 0C,
we get oxygen gas.

Potassium chlorate 2500 C Potassium chloride + Oxygen
MnO2

2KClO3 2500 C 2KCl + 3O2
MnO2

Apparatus required

i. Hard glass test tube ii. Delivery tube

iii. Water trough iv. Beehive shelf

v. Gas jar vi. Bunsen burner

vii. Cork viii. Stand

Chemicals required

i. Potassium chlorate (KClO3)
ii. Manganese dioxide (MnO2)

Hard glass test tube Cork Oxygen gas
Potassium chlorate Delivery tube Gas jar

+
Manganese dioxide

Stand Water trough
Beehive shelf
Bunsen burner

Fig. 11.5 Laboratory preparation of oxygen gas (O2) by using heat

Procedure
i. The 3:1 ratio of potassium chlorate and manganese dioxide is kept in the hard

glass test tube and the apparatus is set as shown in the figure.
ii. Heat is supplied by Bunsen burner.
iii. When potassium chlorate is decomposed, it gives off oxygen gas.
iv. The oxygen gas produced in the hard glass test tube is collected in the gas jar by

downward displacement of water.
Precautions
i. We should take pure manganese dioxide (MnO2).

190 Oasis School Science - 9 CHEMISTRY

ii. The hard glass test tube should be made inclined.
iii. Heat should be supplied constantly.
iv. The apparatus should be made airtight.

b. Without using heat

Principle

Oxygen gas can be prepared in the laboratory by the decomposition of hydrogen
peroxide in the presence of manganese dioxide.

Hydrogen peroxide MnO2 Water + Oxygen
2H2O2 MnO2 2H2O + O2

Apparatus required

i. Conical flask ii. Thistle funnel

iii. Delivery tube iv. Gas jar

v. Beehive shelf vi. Water trough

vii. Cork viii. Stopper

Chemicals required

i. Hydrogen peroxide (H2O2)
ii. Manganese dioxide (MnO2)

Hydrogen peroxide

Stopper Delivery tube

Cork Oxygen gas
Conical flask Gas jar

Mixture of H2O2 + MnO2 Beehive shelf Water
Water trough

Fig. 11.6 Laboratory preparation of oxygen without using heat

Procedure
i. Keep some manganese dioxide in the conical flask and arrange the apparatus as

shown in the figure.

ii. Pour hydrogen peroxide drop by drop with the help of the stopper.

iii. When hydrogen peroxide decomposes in the presence of manganese dioxide, it
gives off oxygen gas.

CHEMISTRY Oasis School Science - 9 191

iv. Oxygen gas produced in the conical flask is passed through the delivery tube and
it is collected in the gas jar by downward displacement of water.

11.15 Test of Oxygen Gas

i. When a burning match stick is brought near the mouth of the gas jar, the flame becomes
brighter indicating that the gas jar is filled with oxygen gas.

ii. When a burning magnesium ribbon is kept inside the gas jar, it burns and gives white
ash of magnesium oxide (MgO). It also proves that the gas is oxygen.

Note
• Oxygen is not inflammable but a very good supporter of combustion.
• Oxygen is less soluble in water. So, it is collected by passing through water.
• Oxygen is a light gas. So, it is collected by downward displacement of water.
• Nascent oxygen is an atomic form of oxygen which is produced during chemical

reaction. It is more reactive and it has more internal energy.
• Manganese dioxide (MnO2) is a positive catalyst which increases the rate of chemical

reaction.

Activity 1

• Take some potassium permanganate in a hard glass test tube.

• Heat the test tube gently by using a Bunsen burner. A gas is produced during this
process.

• Collect the gas in a test tube.

• Identify the gas produced by this process.

2KMnO4 ∆ K2MnO4 + MnO2 + O2

11.16 Manufacture of Oxygen

For commercial purpose, oxygen gas is manufactured in the following ways:
a. From air
b. From electrolysis of water

a. From air

The mixture of different types of gases is collected and compressed into low volume
with a high pressure. This mixture of gases now changes into liquid state. In this liquid
mixture, there are only nitrogen and oxygen gases. The boiling point of nitrogen (–1960C)
is lower than that of oxygen (–1830C). So, nitrogen gas escapes out leaving only oxygen
gas. Then the oxygen gas is collected in cylinders.

nascent / ˈ n æ s n t / - newly formed atomic form and highly reactive
combustion /kəmˈbʌstʃn/ - the process of burning

192 Oasis School Science - 9 CHEMISTRY