science 9

at which the horizontal component of earth’s magnetic field is equal and
opposite to the magnetic field of the magnet.
10. The angle between the magnetic meridian and geographical meridian at
a place is called angle of declination of that place.
11. The angle between the direction of earth’s magnetic field at a place and
the horizontal line is called the angle of dip at that place.

Exercise

1. Choose the best alternative in each case.

a. The presence of electric current in an electric circuit is detected by

i. Ammeter ii. Voltmeter iii. Galvanometer iv. Potentiometer

b. What is the SI unit of potential difference?

i. Ohm ii. Volt iii. Ampere iv. Ohm x m

c. A battery of 9 volts and a bulb of 6 Watts have been used in a torch

light. What will be the amount of current in the circuit?

i. 1.5A ii. 2.0A iii. 2.25A iv. 0.67A

d. The value of angle of dip at the magnetic pole is

i. 0° ii. 90° iii. 45° iv. 72.5°

e. The value of the angle of dip at the Kathmandu valley is

i. 90° ii. 50° iii. 42° iv. 35°

2. Answer these questions in very short.
a. What is a cell?
b. Write down differences between a cell and a generator.
c. Draw the symbols for battery and galvanometer.
d. Why do we use conventional direction of current although it
is not correct ?
e. What is magnetism ?
f. Write down any two advantages of hydroelectricity.

3. Define: b. Ammeter
a. Galvanometer d. Neutral points
c. Voltmeter f. Geographic meridian
e. Magnetic meridian

4. Differentiate between:
a. Potential difference and electromotive force
b. Angle of declination and angle of inclination
c. Resistance and resistivity
d. Ammeter and voltmeter

5. Give reasons:
a. An ammeter is always connected in series in an electric circuit.

147 Times' Crucial Science Book - 9

b. A voltmeter is connected parallel to load in an electric circuit.
c. A bulb contains coiled wire.
d. The value of emf is more than the value of Pd.
e. Bulbs are connected in parallel combination.
f. A freely suspended magnet rests pointing N-S direction.
g. Angle of dip is 90° at magnetic poles.
h. A magnetic compass does not show any particular direction

at neutral point.

6. Solve the following numerical problems.
a. How much charge flows through a wire carrying 3A of
current in 10 minutes?
b. What is the resistance of a material that draws 200mA of
current at 20V potential difference?
c. Calculate:
i. Total resistance from the given diagram.

Fig:(a) 6W Fig:(b) 6W
4W 6W

12v 12v

d. What is the potential difference across a resistor having

resistance of 4W which draws a current 3A.

e. What is the total resistance when two resistors having
resistance 6W and 9W are connected (i) in series and (ii)

parallel combination?

Answers 6. a. 1800 Coulomb b. 100W c.i)Fig:(a)10W,

Fig:(b)3W d. 12V e. 15W, 3.6W

Project Work

Take a conducting wire having 10 feet length, a
torchlight bulb, two dry cells, etc. Then, connect
two cells in series combination. Now join the cells
to each end of the bulb with conducting wires. Put
a switch in the conducting wire. Now the electric
circuit is ready. Observe the brightness of bulb
by changing the number of cells. Again connect the cells in parallel
and observe brightness. Write your observations with reasons.

Times' Crucial Science Book - 9 148

Chapter Valency

8 and
Molecular Formula
Antoine Lavoisier

He is known for studying the properties of
Hydrogen and oxygen. He is sometimes known

as the father of chemiEstsrtyi.mated Periods :6

Objectives

At the end of the lesson, students will be able to:
• define matter;
• explain element, compound and their constituents;
• define octet and duplet rule;
• explain radicals, valency and molecular formula.

Mind Openers

• What is chemistry? Does chemistry have branches?
• What are the differences between atom and molecule?
• Can you define symbol? What is its importance? Discuss.

Introduction

Matter is anything that has mass and occupies space. The matter
exists in several forms and each form of matter has its own
properties. One form of matter can be changed into another form.
There is a separate branch of science to study about the properties
and composition of matter. It is called chemistry.

Chemistry is the branch of science which deals with the study of
properties, composition and transformation of matter.

Chemistry can be divided into three main branches. They are as
follows:

1. Inorganic Chemistry
2. Organic Chemistry
3. Physical Chemistry

1. Inorganic Chemistry

The branch of chemistry that deals with the study of inorganic
compounds is called inorganic chemistry. The compounds such as
H2O, NaCl, NH3, NaOH, H2SO4, etc are inorganic compounds. The
inorganic compounds can be formed by any elements but do not
contain carbon and hydrogen together.

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2. Organic Chemistry

The branch of chemistry in which we study about hydrocarbons
and their derivatives is called organic chemistry. The compounds
which contain direct covalent bond between carbon and hydrogen
are called hydrocarbons. For example, CH4, C2H6, C3H8, C2H5OH,
etc are the organic compounds.

3. Physical Chemistry

The branch of chemistry which deals with the energy changes in the
chemical reactions is called physical chemistry.

In this chapter, we will study about the fundamental terms of
chemistry and the ways to derive molecular formula.

Elements, compounds and their mixtures are regarded as matter.
The elements and compounds have definite composition and
particular properties. So, they are called pure substances. But a
mixture can be formed if two or more substances are placed together
in any ratio. So, a mixture does not have definite composition and
properties and is known as impure substance. A brief classification
of matter is as follows:

Matter

Pure Substances Impure Substances

Elements Compounds Heterogeneous Homogeneous

Mixture Mixture

A short description of pure substances is as follows:

1. Elements
An element is the simplest pure form of substance which is made of
similar kinds of atoms. There are about 118 elements discovered
so far. Among them 92 elements occur naturally and 26 elements
are obtained in laboratories artificially. Hydrogen, carbon, oxygen,
sodium, iron, cobalt, nickel, uranium, etc are some examples of
elements.

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The elements are represented in short form with the help of their
symbols. Symbols are the short and abbreviated forms for the full
name of elements. Generally, the symbols are expressed by taking
first letter of the English or Latin name of elements. In case the
name of more than one element begins with the same letter, two or
three letters are also used as symbol.

Symbols of some common elements are as follows:

SN Element Symbol Atomic Number SN Element Symbol
1 Hydrogen H1
2 Helium He 2 21. Chromium Cr
3 Lithium Li 3
4 Beryllium Be 4 22. Manganese Mn
5 Boron B5
6 Carbon C6 23. Iron (Ferrum) Fe
7 Nitrogen N7
8 Oxygen O8 24. Copper (Cuprum) Cu
9 Fluorine F9
10 Neon Ne 10 25. Zinc Zn
11 Sodium (Natrium) Na 11
12 Magnesium Mg 12 26. Molybdenum Mo
13 Aluminium Al 13
14 Silicon Si 14 27. Silver (Argentum) Ag
15 Phosphorus P 15
16 Sulphur S 16 28. Tin (Stannum) Sn
17 Chlorine Cl 17
18 Argon Ar 18 29. Iodine I
19 Potassium K 19
20 Calcium Ca 20 30. Platinum Pt

31. Gold (Aurum) Au

32. Mercury (Hydrargyrum) Hg

33. Lead (Plumbum) Pb

34. Tungsten (Wolfam) W

35. Lanthanum La

36. Actinium Ac

37. Uranium U

38. Curium Cm

39. Californium Cf

40. Einsteinium Es

2. Compounds

Compounds are the pure substances which are formed by the chemical
combination of two or more elements in a definite proportion by
weight. The properties of compounds are different from those of
elements from which they are formed. For example, water, common
salt, etc are compounds.

A compound is composed of molecules. A molecule is the smallest

151 Times' Crucial Science Book - 9

particle of a compound or an element which has independent
existence. A molecule of a compound is different from that of an
element.

Differences between the molecule of a compound and that of an
element

Molecule of a compound Molecule of an element

1. It is made up of two or more 1. It is made up of similar types of

different types of atoms. atoms.

2. The atoms of different elements 2. It contains the atoms of same
form a molecule of a compound. element.

3. Example: H2O is the molecule 3. Example: H2 is molecule of
of water, H2SO4 is a molecule of hydrogen, O2 is the molecule of
Sulphuric acid, etc. oxygen, etc.

Atom

Atom is the smallest particle of an element that can take part in
a chemical reaction. An atom is represented by the symbol of an
element. For example, H represents an atom of hydrogen, Na
represents an atom of sodium, etc. An atom is an electrically neutral
particle because it has equal number of positively charged protons
and negatively charged electrons.

Structure of atom

An atom is extremely small. The diameter of an atom is approximately
10−10 m. But it is made up of several particles called sub-atomic
particles. Electron, proton and neutron are the sub-atomic particles
present in an atom. Protons and neutrons are located in the nucleus
of an atom but the electrons revolve round the nucleus. The electrons
revolve around the nucleus in a fixed path which is known as orbit
or shell.

Electron

Protons Nucleus
Neutrons

Shell

Structure of an atom

The mass of an atom is concentrated in the nucleus. ln nucleus,

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protons and neutrons are present. The weight of a proton is equal
to the weight of one hydrogen atom. The weight of neutron is nearly
equal to the weight of a proton, i.e. one atom of hydrogen. The
weight of an electron is equal to 1/1837 th the weight of one hydrogen
atom. Since the weight of electron is very less as compared to that
of proton or neutron, it is often neglected while calculating atomic
weight.

Each orbit or shell is composed of

sub-shells. There are four kinds of

sub-shells. They are s,p,d and f sub-

shells. The s-subshells are spherical

in shape where as the p-subshells

have dumb-bell shape. S-subshell P-subshell
(dumb-bell shape)
(spherical shape)

Each cell has a particular number of

sub-shells. Each sub-shell can accommodate a definite number of

electrons. It is given in the following table:

Shells KL M N
Sub-shells
s sp s p d s p d f

No. of electrons 2 2 6 2 6 10 2 6 10 14

The table shows that,

The s-sub-shell can hold a maximum of 2 electrons.

The p-sub-shell can hold a maximum of 6 electrons.

The d-sub-shell can hold a maximum of 10 electrons.

The f-sub-shell can hold a maximum of 14 electrons.

Atomic number

The total number of protons present in the nucleus of an atom is
called atomic number. In an electrically neutral atom, the number
of protons is equal to the number of electrons. Thus, atomic number
may also be defined as the total number of protons or electrons
present in an electrically neutral atom. The atomic number is
denoted by Z.

Atomic number = No. of protons

= No. of electrons in an electrically neutral atom

∴ Z = No. of p+ or e−

Atomic weight

The sum of number of protons and neutrons in the nucleus of an

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atom is called atomic weight. It is also known as atomic mass. The
atomic weight of an element is denoted by A.

A = No. of protons + No. of neutrons

∴ A = p+ + n

Electronic configuration of atoms

The systematic distribution of electrons in different energy levels in
an atom is called electronic configuration. Electrons are distributed
in different shells or orbits. A shell or orbit is a well defined fine
circular path in which the electrons move around the nucleus. The
shells can be represented as K, L, M, N,................. or 1, 2, 3, 4,.........
etc.

Electronic configuration of some atoms is given below:

SN Element Symbol No. of No. of No. of At. Electronic
electrons protons neutrons Mass
1. Hydrogen configuration
2. Helium H 1 1 01 K L MN
3. Lithium He 2 2 2 4 1 - --
4. Beryllium Li 3 3 4 7 2 - --
5. Boron Be 4 4 5 9 21--
6. Carbon B 5 5 6 11 22--
7. Nitrogen C 6 6 6 12 23--
8. Oxygen N 7 7 7 14 24--
9. Fluorine O 8 8 8 16 25--
10. Neon F 9 9 10 19 26--
Ne 10 10 10 20 27--
28--

11. Sodium Na 11 11 12 23 2 8 1 -

12. Magnesium Mg 12 12 12 24 2 8 2 -

13. Aluminium Al 13 13 14 27 2 8 3 -

14. Silicon Si 14 14 14 28 2 8 4 -

15. Phosphorus P 15 15 16 31 2 8 5 -

16. Sulphur S 16 16 16 32 2 8 6 -

17. Chlorine Cl 17 17 18 35 2 8 7 -

18. Argon Ar 18 18 22 40 2 8 8 -

19. Potassium K 19 19 20 39 2 8 8 1

20. Calcium Ca 20 20 20 40 2 8 8 2

The distribution of electrons in different shells is governed by
following rules.

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1. The maximum number of electrons present in each shell is given by
2n2 rule. The rule that predicts the maximum number of electrons
that can be accommodated in different shells of an atom is called 2n2
rule. For example,

Shell Number Maximum number of elec-trons
K n = 1 2n2 = 2 × 12 = 2
L n = 2 2n2 = 2 × 22 = 8
M n = 3 2n2 = 2 × 32 = 18
N n = 4 2n2 = 2 × 42 = 32 and so on.

2. The outermost shell of an atom cannot have more than 8 electrons
and the shell just inner to it cannot contain more than 18 electrons.

3. It is not necessary for a shell to be completely filled before a new shell
starts to be filled.

Electronic configuration of some elements is as follows:

155 Times' Crucial Science Book - 9

Duplet and octet

The element helium has two electrons in K-shell. The K-shell is
completely filled with two electrons and it cannot hold any more
electrons. The state in which the outermost shell of an atom has the
capacity to accommodate only two electrons is called duplet. Helium
is only one element in the periodic table which has duplet. It is an
inert gas because it does not undergo any chemical reaction due to
completely filled valence shell.

Some elements have 8 electrons in their valence shell. A stable set
of eight electrons in the outermost shell of an atom is called octet.
The elements like neon, argon, krypton, xenon, etc have octet state.
They are also the inert gases.

Duplet and octet rule

The tendency of an element to maintain two electrons in outermost
shell is called duplet rule. Only hydrogen, lithium, beryllium and
boron have tendency to gain duplet by gaining or losing electrons.

For example, hydrogen attains duplet by gaining one electron. But
lithium attains duplet by losing one electron.

H +1e− → H−
Li −1e− → Li+

The tendency of an element to acquire eight electrons in its outermost

orbit is called octet rule. The atoms can acquire octet state by gaining,

losing or mutual sharing of electrons. The tendency of elements to

acquire octet or duplet is the cause of chemical reactions.

Na − e− → Na+ (octet state)

(2, 8, 1) (2, 8)

Cl + e−→ Cl− (octet state)

(2, 8, 7) (2, 8, 8)

Valency

The shells of an atom can be shown as:

Nucleus + Valence shell
Penultimate shell
Innermost (first) shell

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Valency of an element depends upon the number of electrons
present in its outermost shell.

The outermost shell or orbit of an atom is called valence shell. The
electrons present in the valence shell are called valence electrons.
The valency of an element is determined by the number of its
valence electrons. The combining capacity of an element is called
valency.

All the inert gases (Helium, Neon, Argon, Krypton, Xenon and
Radon) have zero valency because their outermost shells are
completely filled and they have no combining capacity.

Valencies of some common elements are as follows:

SN Name of Symbol Valency SN Name of Symbol Valency
element element
H1 Zn 2
1. Hydrogen He 0 1. Zinc Cr 3
Li 1 Br 1
2. Helium Be 2 2. Chromium Au 1, 3
B3 I1
3. Lithium C4 3. Bromine La 3
N3 Ra 0
4. Beryllium O2 4. Gold U4
F1 Mn 2, 4
5. Boron Ne 0 5. Iodine As 3,5

6. Carbon 6. Lanthanum

7. Nitrogen 7. Radon

8. Oxygen 8. Uranium

9. Fluorine 9. Manganese

10. Neon 10. Arsenic

11. Sodium Na 1 11. Krypton Kr 0
12. Magnesium Mg 2
13. Aluminium Al 3 12. Silver Ag 1
14. Silicon Si 4
15. Phosphorus P 3 13. Xenon Xe 0
16. Sulphur S 2
17. Chlorine Cl 1 14. Tungsten W 2,4
18. Argon Ar 0
19. Potassium K 1 15. Actinium Ac 3
20. Calcium Ca 2
16. Barium Ba 2

17. Vanadium V 4

18. Nickel Ni 2,3

19. Cobalt Co 2

20. Iron Fe 2, 3

The number of electrons gained, lost or shared during a chemical
reaction gives the valency of an element. For example, sodium loses
one electron from its valence shell and attains octet. It is the stable
state of sodium. Hence, its valency is 1. Similarly, the valency of

157 Times' Crucial Science Book - 9

chlorine is also 1 because it gains one electron to be in stable form.

Variable valency

Some of the elements may have two or more valencies. For example,
the valencies of iron are 2 and 3, the valencies of copper are 1 and
2 and so on. This property of elements is called variable valency.
The variable valency of elements is due to the incompletely filled
penultimate shell.

Variable valency of some common elements are as follows:

Element Variable

Sulphur (S) valency
Nitrogen (N) 2, 4, 6
Phosphorus (P) 3, 5
Manganese (Mn) 3, 5
Copper (Cu) 2, 4
Mercury (Hg) 1, 2
1, 2
Gold (Au) 1, 3
Iron (Fe) 2, 3
Lead (Pb) 2, 4
Tin (Sn) 2, 4

The valency of elements of group I, II, III are +1, +2, +3 respectively.
Likewise, the valency of element of group IV is either +4 or -4.
Similarly, the valency of group O is o. The valency of elements
of group V, VI, VII are -3, -2, and -1 respectively. This rule is not
applicable in every element.

Radicals

An atom or group of atoms carrying either positive or negative
charge which acts as a single unit in chemical reactions is called
radical. Radicals do not exist freely in nature because they react
readily with other radicals or elements of their surroundings and
form compounds.

On the basis of type of electric charge contained by radicals, they
can be classified into two types:

Times' Crucial Science Book - 9 158

1. Electropositive radicals

Electropositive radicals are those radicals which possess positive
charge. They possess positive charge because the number of positive
charge exceeds the number of electrons. Electropositive radicals are
also known as basic radicals or metallic radicals. Some examples of
electropositive radicals are as follows:

Valency'1' Valency'2' Valency'3' Valency'4'
Monovalent radical Divalent radical Trivalent radical Tetravalent radical
Hydrogen (H+) Magnesium (Mg++) Boron (B+++) Plumbic (Pb++++)
Sodium (Na+) Calcium (Ca++) Aluminium (Al+++) Stannic (Sn++++)
Potassium (K+) Zinc (Zn++) Chromium (Cr+++) Manganic (Mn++++)
Silver (Ag+) Cupric (Cu++) Auric (Au+++)
Cuprous (Cu+) Barium (Ba++) Arsenous (As+++) -
Mercurous (Hg+) Mercuric (Hg++) Ferric (Fe+++) -
Aurous (Au+) Stannous (Sn++) -
Ammonium(NH4+) Ferrous (Fe++) - -
- -

2. Electronegative radicals

Electronegative radicals are those radicals which carry negative
charge. They posses negative charge because the number of
electrons exceeds the number of protons present in the atoms.
Electronegative radicals are also known as acidic or non-metallic
radicals. Some examples of electronegative radicals are as follows:

Monovalent (’1’) Bivalent (’2’) Trivalent (’3’) Tetravalent (’4’)

Fluoride (F—) Oxide (O— —) Nitride (N— — —) Carbide (C4—)
Phosphate(PO4— — —) -
Chloride (Cl—) Sulphide (S— —) Phosphide (P— — —) -
-
Bromide (Br—) Carbonate (CO3— —) - -
Silicate (SiO3— —) - -
Iodide (I—) Peroxide (O2— —) - -
Sulphite (SO3— —) - -
Hydroxide (OH—) Sulphate (SO4— —) - -
Zincate (ZnO2— —) - -
Chlorate (ClO3—) -
Nitrite (NO2—) -
Nitrate (NO3—)
Bicarbonate (HCO3—) -
Cyanide (CN—)

Permanganate (MnO4—) - --
Bisulphate (HSO4—) - --

159 Times' Crucial Science Book - 9

Way to find out the valency of radicals:

SN Radicals Number of atoms and their Residual
1. Hydroxide (OH—)
2. Carbonate (CO3— —) valencies valency

3. Bicarbonate(HCO3—) Valency of O = −2 −2 + 1 = −1
4. Ammonium (NH4+) Valency of H = +1
5. Phosphate (PO4— — —)
Valency of C = +4 +4 −6 = −2
Valency of O3 = (−2) × 3 = −6

Valency of H = +1

Valency of C = +4 +5 −6= −1

Valency of O3 =−6

Valency of N = −3 −3 + 4 = 1
Valency of H4 = (+1 × 4) = +4

Valency of P = +5 −8 + 5= −3
Valency of O4 = (−2) × 4 = −8

Ions
The electrically charged atoms or particles which carry positive or
negative charge are called ions. On the basis of charge carried by
ions, they can classified into two types: cation and anion.
a. Cations: The ions which carry positive charge are called cations.
An atom becomes cation by losing electron during a chemical
reaction. For example, H+ , Na+, Mg++, K+, Ca++, etc.
b. Anions: The ions which carry negative charge are called anions.
An atom becomes anion when it gains electrons during chemical
reaction. For example, Cl−, ClO3−, SO4−−, NO3−, PO4−−− etc are some
examples of anion.
Difference betweeen cations and anions

Cations Anions

1. They carry positive charge. 1.They carry negative charge.

2. They are formed by losing electrons. 2. They are formed by gaining
electrons.

3. Mostly metals form cations. 3. Mostly non-metals form
anions.

4. They are discharged at cathode 4. They are discharged at

during electrolysis. anode during the process of

electrolysis.

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Chemical bond
All elements except inert gases try to become stable by gaining,
losing or sharing electrons. This is the main cause of chemical
reaction. Metals lose electrons and become positively charged
while non-metals gain electrons and become negatively charged.
When positively charged particles come near to negatively charged
particles, they combine to form compound. In such compounds,
the atoms are held together by a force called chemical bond. In
other words, the force of attraction between two or more atoms in a
compound is called chemical bond. A chemical bond can be formed
by gaining, losing or mutual sharing of electrons.
There are three types of bonds. They are electrovalent bond,
covalent bond and co-ordinate covalent bond. In this chapter, we
shall discuss about electrovalent and covalent bond.
1. Electrovalent bond
The bond formed by the complete transfer of electrons from one atom
to another during a chemical reaction is called electrovalent bond. It
is also known as ionic bond. The compounds which have electrovalent
bond are called electrovalent compounds. Electrovalent compounds
are generally solid in nature. They have high melting and boiling
point. They undergo electrolysis. Calcium oxide, sodium chloride,
magnesium chloride, magnesium oxide, calcium chloride, etc are
some examples of electrovalent compounds.
Formation of sodium chloride
The electronic configuration of sodium is (2, 8, 1) and that of chlorine
is (2, 8, 7). One atom of sodium and an atom of chlorine combine
with each other to form sodium chloride (NaCl). This combination
is called chemical reaction in which sodium atom loses one electron
and chlorine atom gains that electron. This ‘loss and gain’ of electron
results in a force of attraction between these two atoms. This force
of attraction is called electrovalent bond.

p=11 p=17 p=11 p=17
n=12 n=18 n=12 n=18

(Na:2,8,1) (Cl:2,8,7) (Na−:2,8) (Cl+:2,8,8)

Attractive force

(electrovalent bond)

Formation of sodium chloride (NaCl)

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2. Covalent bond

The bond which is formed by the mutual sharing of electrons
between the combining atoms during chemical reaction is called
covalent bond. The compounds which have covalent bond are called
covalent compounds. Ammonia, water, methane and carbon dioxide
are some examples of covalent compounds. The covalent bond is
also present in the molecule of some elements such as hydrogen,
oxygen, nitrogen, chlorine etc. Covalent compounds are generally
found in liquid or gaseous state. The solid covalent compounds have
low melting and boiling points. They cannot work as electrolyte.
They can easily dissolve in organic solvents.

Formation of water molecule

The electronic configuration of oxygen is (2, 6) and that of hydrogen
is 1. Oxygen and hydrogen atoms share 2 pairs of electrons and
form a molecule of water (H2O). In this combination, two hydrogen
atoms contribute one electron each for sharing with two electrons
from oxygen.

p=1 p=1

p=8 p=8
n=8 n=8

p=1 p=1

Formation of Water (H2O)

Differences between ionic and covalent compounds

Electrovalent compounds Covalent compounds

1. They are formed by the complete 1. They are formed due to the mutual

transfer of electrons from one atom sharing of electrons between the

to another. combining atoms.

2. They are generally solids. 2. These compounds may be gas,
liquid or solid. The solids are soft
and brittle.

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

points. boiling points.

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4. They conduct electricity in their 4. They generally do not conduct
molten or aqueous solution state. electricity.

5. They are usually insoluble in 5. They are usually soluble in water.
water.

Molecular formula

The symbolic representation of a molecule of a compound or an
element which shows the actual number of atoms in a molecule is
called molecular formula. For example, the molecular formula of
ammonia is NH3. Here, NH3 represents a molecule of ammonia
which consists of four atoms. Among four atoms, one is nitrogen
and three are hydrogen atoms.

Information from molecular formula
1. Molecular formula represents the ratio of different elements present in
the given compound.
2. We can find the valency of combining atoms from the molecular formula .
3. It gives the actual number of atoms present in a molecule of a compound.

4. We can calculate the molecular weight of a molecule of the compound.

Molecular weight

The sum of atomic weights of all the atoms present in a molecule is
called molecular weight. lt is also known as molecular mass.

For example: =2×H+1×O
Molecular weight of H2O = 2 × 1 + 1 × 16
= 18
Molecular weight of CO2 =1×C+2×O
= 12 + 32
Molecular weight of NH3 = 44
=1×C+3×H
= 1 × 12 + 3 × 1
=17

Method to write molecular formula of compounds

1. Write the symbol of elements or radicals side by side. Write
metal first and then non-metal. For example, to write molecular
formula of calcium oxide.

Calcium Oxygen

Ca O

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2. Write the valency of each element or radical on the top of symbol.

Calcium Oxygen

22

Ca O

3. Interchange the valency of the elements or radicals.

2 2 = Ca2O2
Ca O

4. Remove the common valency if present.

Ca2O2 = CaO (Here, 2 is common and is removed)
2. Ammonia
Examples:
1. Water

Hydrogen Oxygen Nitrogen Hydrogen
1 3
H 2 = H2O N 1 = NH3
O H

3. Common salt (sodium chloride) 4. Calcium chloride
Calcium
Sodium Chloride = NaCl 2 Chloride
11 Ca 1 =CaCl2
Na Cl Cl

5. Magnesium Oxide 6. Sodium hydroxide

Magnesium Oxygen Sodium Hydroxide
2 1 =NaOH
2 2 = Mg2O2 Na OH
Mg O = MgO

7. Copper sulphate 8. Ammonium Sulphate

Copper Sulphate Ammonium Sulphate
22
Cu SO4 = Cu2(SO4)2 1 2 =(NH4)2SO4
= CuSO4 NH4 SO

9. Ferric oxide 10. Potassium Cyanide

Ferric oxide Potassium Cyanide
1 =KCN
3 2 = Fe2O3 1 CN
Fe O K

Times' Crucial Science Book - 9 164

SN Name of Compound Chemical Name Molecular Formula
1. Water Hydrogen Oxide H2O
2. Common salt Sodium Chloride NaCl
3. Sulphuric acid Hydrogen Sulphate H2SO4
4. Ammonia Hydrogen Nitride NH3
5. Lime water Calcium Hydroxide
6. Baking soda Sodium Bicarbonate Ca(OH)2
7. Blue vitriol Copper Sulphate NaHCO3
8. Nitric acid Hydrogen Nitrate CuSo4.5H2O
9. Hydrochloric acid Hydrogen Chloride
10. Gypsum Calcium Sulpahte HNO3
HCl
11. Chalk, limestone or marble Calcium Carbonate CaSO4.2H2O

12. Plaster of paris Calcium Sulphate CaCO3
CaSO4.1/2H2O
13. Washing soda Sodium Carbonate Na2CO3.10H2O
14. Quick lime Calcium Oxide
15. Marsh gas Methane CaO
CH4
16. Calcium chloride Calcium chloride
CaCl2
17. Magnesium carbonate Magnesium carbonate
MgCO3
18. Potassium sulphate Potassium sulphate
K2SO4
19. Potassium carbonate Potassium carbonate
K2CO3
20. Lunar caustic Silver nitrate
AgNO3

Main points to remember
1. The simplest and pure form of substance which is made of similar kinds

of atoms is called element.
2. Symbols are the short and abbreviated form for the full name of elements.
3. The pure substances which are formed by the chemical combination

of two or more elements in a definite proportion by weight are called
compounds.
4. An atom is the smallest particle of an element that can take part in a
chemical reaction.
5. A molecule is the smallest particle of a compound or an element which
has independent existence.
6. The state in which the outermost shell of an atom has the capacity to
accommodate no more than two electrons is called duplet.
7. A stable set of eight electrons in the outermost shell of an atom is called octet.
8. The number of electrons in the valence shell of an atom that determines

165 Times' Crucial Science Book - 9

the combining capacity of an element is called valency.
9. An atom or group of atoms carrying either positive or negative charge

that behaves as a single unit in a chemical reaction is called radical.
10. The radicals that carry positive charge due to the loss of electron are

called electropositive radicals.
11. Electronegative radicals are the radicals which carry negative charge

due to the gain of electrons.
12. The attractive force which holds two or more atoms in a molecule of a

compound is called chemical bond.
13. The bond which is formed by the transfer of electrons from one atom to

another during chemical reaction is called electrovalent bond.
14. The bond which is formed by the mutual sharing of electrons among

atoms during chemical reaction is called covalent bond.
15. The symbolic representation of the molecule of a compound or element,

which shows the actual number of atoms present in a molecule is called

molecular formula.
Learn and Write

1. What is the difference between H2 and 2H?
H2 represents a molecule of hydrogen whereas 2H represents
two atoms of hydrogen.

2. What is the difference between Co and CO?
Co represents the element Cobalt but CO represents the
compound carbon monoxide.

3. Neon is an inert gas. Why?
Neon atom has 8 electrons in the outermost orbit and thus it is in
octet state. Hence, its outermost orbit is already complete and it
does not take part in chemical reactions and is an inert gas.

4. Valency of chlorine is 1. Why?
Electronic configuration of chlorine is 2, 8, 7. Therefore, it needs
one electron to make its outermost orbit complete. During its
combination with other elements, it takes one electron from
them. Thus, its valency is 1.

5. Sodium is a reactive metal. Why?
Sodium has only one electron in the outermost orbit. Therefore,
it always tries to lose that one electron to make the outermost
orbit complete. Therefore, it is a reactive metal.

6. What is variable valency? Why do some elements have
variable valency?
The property of an elemement to have more than one kind of
valency is called variable valency. An element has a single

Times' Crucial Science Book - 9 166

kind of valency if only the valence electrons take part in bond
formation. If the valence electrons as well as the electrons of
penultimate shell of an element take part in bond formation,
such element has variable valency.
7. What changes occur in the atoms if they combine to
form a compound?
Each atom attains a stable configuration of the nearest inert
gas after it undergoes chemical combination. The combining
atoms completely lose their individual property and a molecule
with different properties is formed.
8. What are the characteristics of covalent compounds?
The characteristics of covalent compounds are as follows:

a. Covalent compounds are generally found in liquid or gas state.
b. The solid covalent compounds have low melting and boiling points.
c. They are non-electrolyte and hence they do not undergo electrolysis.
d. Most of them are insoluble in water. They are soluble in organic solvents.
9. Why don’t the covalent compounds undergo electrolysis?
The covalent compounds contain covalent bonds. The covalent
bonds are formed by the mutual sharing of electrons. Hence,
these compounds do not dissociate in to ions in their aqueous
solution or molten state. As these compounds have no ions,
they cannot undergo electrolysis

Exercise

1. Choose the best alternative in each case.

a. The hydrogen nucleus is equivalent to

i. Electron ii. Proton iii. Neutron iv. All of these

b. An element having three shells consists of one electron in its

valence shell; the element is

i. Fluorine ii. Magnesium iii. Sodium iv. Potassium

c. Which of the following elements has a variable valency?

i. Calcium ii. Silver iii. Chlorine iv. Copper

d. What is the molecular formula of sulphuric acid?

i. HCl ii. H2SO4 iii. HNO3 iv. NaOH

e. Which of the following is the most reactive non-metal?

i. Fluorine ii. Chlorine iii. Bromine iv. Iodine

2. Answer these questions in very short.
a What is chemistry?
b. How many elements are discovered so far?
c. What is a molecule? Which is a molecule of hydrogen either

167 Times' Crucial Science Book - 9

H2 or 2H?
d. Define octet with examples.
e. What is a valence shell?
f. What do you mean by Co and CO?
g. What is duplet rule?
h. Calculate the molecular weight of CaCO3.
i. Give any four examples of electronegative radicals.

3. Define:

a. Atom b. Duplet c. Octet rule d. Symbol

e. Electronic configuration f. Chemical bond

4. Give reason:
a. The valency of sodium is 1.
b. The valency of argon is 0.
c. Sodium chloride is an electrovalent compound.
d. Carbon dioxide is a covalent compound.
e. An atom is electrically neutral.

5. Differentiate between:
a. Elements and compounds
b. Electrovalent and covalent bond
c. Molecule and atom
d. Symbol and molecular formula
e. Molecule of a compound and molecule of an element
f. Atomic number and atomic weight

6. Write the molecular formulae of following:

a. Water b. Ammonia

c. Sodium chloride d. Calcium oxide

e. Carbon dioxide f. Magnesium hydroxide

g. Potassium oxide h. Methane

i. Potassium chloride j. Nitric acid

k. Sulphuric acid l. Hydrochloric acid

m. Limestone n. Sulphur trioxide

o. Nitrous acid p. Ferrous chloride

q. Hydrogen peroxide s. Auric chloride

r. Sodium peroxide t. Phosphoric acid

u. Potassium bicarbonate v. Ammonium sulphate

w. Potassium chlorate x. Silver chloride

y. Washing soda z. Ammonium nitrate

Times' Crucial Science Book - 9 168

7. Write down the names of following compounds.

a. Ca(OH)2 b. H2SO4 c. CH4 d. MgCO3
e. Ca(HCO3)2
i. KCN f. Ca3(PO4)2 g. HNO3 h. KClO3

m. HCl j. Na2O k. HgO l. CuSO4

n. (NH4)2SO4

Project Work

Observe different elements or compounds that are being used
in your home and school. Write down their names, symbols or
molecular formulae. Also explain their uses.

Glossary
Organic: related to the parts of organisms, containing carbon and

hydrogen
Existence: state of being present
Approximately: nearly, roughly, about
Accommodate: hold, suit, have room for
Variable: changing, not fixed
Penultimate: just inner to the outermost
Discharge: to lose charge, to remove charge
Brittle: ability to break down in to pieces

169 Times' Crucial Science Book - 9

Chapter Fritz Haber

9 Chemical
Reaction

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

explosive, etc. Estimated Periods :6

Objectives

At the end of the lesson, students will be able to:

• define physical and chemical changes with examples;
• differentiate between physical and chemical change;
• define chemical change and represent chemical change in terms of word and formula

equations;
• define reversible and irreversible reaction.

Mind Openers

• What is a change?
• What are physical and chemical changes? Can you tell their examples?
• What is a chemical reaction? Can you represent chemical reaction in terms of

symbols and molecular formula? Discuss.

Introduction

Several types of changes occur in our surrounding continuously.
Some of these changes take place slowly while some take place
faster. The conversion of one form of things into another is called
change. The change can be categorized into two types physical
change and chemical change.

1. Physical change

A physical change is a temporary change in which no new substance
is formed. The shape, size, volume, state, etc of the given substance
may change in this change. For example, if ice is heated, it changes
into water. If water is again heated, it changes into vapour. In
these cases, there is change in the physical state of matter. If water
vapour is cooled, it changes into water and the further cooling of
water produces ice. Hence, such changes can be reversed. It means
that the physical change is temporary and reversible.

Some common examples of physical change are as follows:
1. Melting of ice into water

Times' Crucial Science Book - 9 170

2. Evaporation of water into vapour
3. Dissolving salt into water
4. Tearing paper into pieces.
5. Making sickle from iron rod.
6. Magnetizing an iron nail.
7. Making shirt from raw cloth

8. Melting ghee into liquid state

In physical change, there is no change in the molecular level of
substances.

2. Chemical change

A chemical change is a permanent change in which an entirely new
substance with new properties is formed. There is change in the
molecular level of matter. Chemical change cannot be reversed
easily. For example, if boiled milk is cooled and acted by a specific
type of bacteria for some hours, it gets converted to curd. But the
curd cannot be reversed back to milk. Hence, chemical change is a
permanent and irreversible change.

Some common examples of chemical change are as follows:
1. Burning of firewood
2. Rusting of iron
3. Turning of milk into curd
4. Digestion of food in our body
5. Growth of baby into adult
6. Electrolysis of water

Some of the chemical changes can be reversed by chemical processes.

For an example, electrolysis of water decomposes it into hydrogen
and oxygen.

Oxygen gas (O2) Half cut plastic bottle
Graphite anode
Hydrogen gas (H2)
Water with Sulphuric acid

Graphite cathode

Rubber cork

Battery Times' Crucial Science Book - 9
Electrolysis of water

171

A molecule of water is made by the chemical combination of
hydrogen and oxygen atoms. By electrolysis, the water molecule
can be decomposed into totally new substances, i.e. oxygen and
hydrogen.

Differences between physical and chemical change

Physical change Chemical change

1. Physical change is a temporary 1. Chemical change is a permanent

change in which no new substance change in which a new substance is

is formed. formed.

2. There is change in physical 2. There is change in both physical and

properties only. chemical properties.

3. There is no change in energy. 3. There is change in energy and mass.
Energy is either absorbed or evolved.

4. It is usually reversible. 4. It is usually irreversible

5. It involves change in shape, size, 5. It involves change in molecular level

volume, state of matter, etc. of matter.

Chemical reaction

A chemical reaction is a chemical change in which combination,
decomposition or exchange of molecules of matter takes place. The
chemical reaction can be represented by using symbols of atoms
and molecular formulae of molecules.

Chemical equation

The symbolic representation of actual chemical reaction in terms of
symbols and formulae is called chemical equation. There are two
types of chemical equations on the basis of the way of representation.
They are:

(i) Word equation
(ii) Symbolic equation or formula equation

i. Word equations

Word equation is a form of chemical equation which is represented
by writing the full names of reactants and products. The substances
which undergo the chemical change are called reactants. They are
written on the left hand side of chemical equation. Similarly, the
substances which are produced as a result of chemical change in the
reactants are called products. They are written on the right hand
side of chemical equation. Some examples of word equations are as

Times' Crucial Science Book - 9 172

follows:

Hydrogen + oxygen → Water

Carbon + hydrogen → Methane

Carbon + oxygen → Carbon dioxide

Calcium carbonate → Calcium oxide + Carbon dioxide

(Reactant) (Products)

ii. Symbolic equation or formula equation

Symbolic equation is a form of chemical equation which is

represented by writing the symbols and formulae of reactants and

products. Some examples of symbolic equation are as follows:

2H2 + O2 → 2H2O
Reactants Product

C + 2H2 → CH4
Reactants Product

2KClO3 → 2KCl+ 3O2
Reactant Products

The physical states of the reactants and products can be represented
by using symbols.

For example:
Calcium carbonate (s)→heat Calcium oxide (s) + Carbon dioxide (g)
CaCO3 (s)→∆ CaO(s) + CO2(g)

Here, s = solid, g = gas

Heat required for the reaction is represented by ∆ (delta) above the
arrow. When calcium carbonate is heated, it decomposes to give
calcium oxide and carbon dioxide.

Endothermic and exothermic reactions

A chemical reaction which needs heat is called endothermic reaction.

For example,
2KClO3→∆ 2KCl + 3O2

A chemical reaction in which heat is produced is called exothermic
reaction.

CH4 + 2O2 → CO2 + 2H2O + heat
Reversible and irreversible reactions

A chemical reaction in which reactants are changed into products

173 Times' Crucial Science Book - 9

and the products are changed back into reactants is called reversible
reaction. Reversible reaction is denoted by arrows that point forward
and backward (l).

For example:

N2 + 3H2 l 2NH3

A chemical reaction which occurs in only one direction, i.e., reactants
to product is called irreversible reaction. Irreversible reactions are
represented by (→).

2H2 + O2 → 2H2O
CaCO3 →∆ CaO + CO2

Unbalanced and balanced chemical equations
The chemical equation which has unequal number of atoms in the
reactant and product side is called unbalanced chemical equation.
Such equations are also called Skeleton chemical equations.

For example:

H2 + O2 → H2O

Na + Cl2 → NaCl
The chemical equation which has equal number of atoms in reactant
and product side is called balanced chemical equation.

For example
2H2 + O2→∆ 2H2O

CH4 + 2O2 → CO2 + 2H2O

Diatomic elements
The elements whose molecules are made up of two atoms are
called diatomic elements. When diatomic elements take part in
reaction, they are always represented in molecular form. Hydrogen,
nitrogen, oxygen, fluorine, chlorine, iodine and bromine are the
diatomic elements. They are written as H2, N2, O2, F2, Cl2, I2 and Br2
respectively.

Examples:

2H2 + O2 →∆ 2H2 O
H2 + F2 →∆ 2HF
2Na + Cl2→∆ 2NaCl

Times' Crucial Science Book - 9 174

Methods of writing chemical equations

1. Write the word equation. Write down the names of reactants
on left side of arrow and that of products on the right side.

Hydrogen + Chlorine → Hydrochloric acid

2. Rewrite the equation in symbolic form.

H2 + Cl2 → HCl

3. If the equation obtained is not balanced equation, balance it
y counting the number of atoms in each side and writing numerical
coefficient where it is necessary.

H2 + Cl2 → 2HCl

4. If any conditions are necessary for the reaction, write it
above the arrow. Thus, a required equation is obtained.

H + Cl Sunlight 2HCl
22

Some other examples:

1. Word equation:

Zinc + dil. sulphuric acid→ Zinc sulphate + Hydrogen

Chemical equation: Zn + dil. H2SO4 → ZnSO4 + H2 ↑
2. Word equation:

Potassium chlorate → Potassium chloride + Oxygen

Skeleton equation: KClO3 → KCl + O2 ↑
Balanced equation: 2KClO3 → 2KCl+ 3O2 ↑
3. Word equation:
Hydrogen peroxide MnO2 Water + Oxygen

Skeleton equation: H2O2 MnO2 H2O + O2↑
Balanced equation: 2H2O2 MnO2 2H2O + O2↑
4. Word equation:

Sodium chloride + Silver nitrate → Sodium nitrate + Silver chloride

Chemical equation: NaCl + AgNO3 → NaNO3 + AgCl
5. Word equation: Hydrogen + Nitrogen l Ammonia

Skeleton equation: H2 + N2 l NH3
Balanced equation: 3H2 + N2 l 2NH3

6. Word equation:

Balanced equation: ZnO + C→∆ Zn + CO

175 Times' Crucial Science Book - 9

7. Word equation:

Phosphorus + Oxygen → Phosphorus pentoxide

Skeleton equation: P + O2 → P2O5
Balanced equation: 4P + 5O2 → 2P2O5
8. Word equation:
Skeleton equation: Mg+ N2 →∆ Mg3N2
Balanced equation: 3Mg + N2 →∆ Mg3N2

9. Word equation:
Lime water + Carbon dioxide → Calcium carbonate + water

Chemical equation: Ca(OH)2 + CO2 → CaCO3 +H2O
10. Word equation:

Sodium hydroxide + Hydrochloric acid → Sodium chloride + water

Chemical equation: NaOH + HCl → NaCl + H2O
Information conveyed by chemical equation

A chemical equation gives full meaning if it is balanced. A balanced
chemical equation gives the following information:

1. It represents the name of the reactants and products.
2. It shows the exact amount of reactants and products.
3. It represents the ratio of molecular weight of the reactant and product

separately.
4. It can show the conditions required for a chemical reaction like heat,

light, catalyst, etc.

5. It can show the type of chemical reaction.

Catalyst

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

a. Positive catalyst b. Negative catalyst

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

Times' Crucial Science Book - 9 176

2KClO3 360°C 2KCl + 3O2
(No Catalyst, so high temperature is required)

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

2KClO 240°C 2KCl + 3O2
3 MnO2

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

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

2H2O2 MnO2 2H2O + O2↑

b. Negative catalysts: A catalyst which decreases the rate of chemical

reaction without being consumed itself is called negative catalyst.

For an example, glycerine or phosphoric acid acts as a negative

catalyst in the decomposition of hydrogen peroxide.

2H O 2H O + O ↑Glycerine22
2 2 (Slowed rate)

Learn and Write

1. What happens when magnesium burns in oxygen?

When magnesium burns in oxygen, it gives magnesium oxide.

2Mg + O2 → 2MgO

Magnesium + Oxygen → Magnesium oxide

2. Melting of ice is a physical change. Why?

When ice is heated, it melts. During this change, there occurs
only the change in its state. There is no change in the molecular
level. It means molecular structure of ice and water is not
different. Therefore, melting of ice is a physical change.

3. Dissociation of water into hydrogen and oxygen on the
passage of current is a chemical change. Why?

When current passes through acidified water, the water
dissociates into oxygen and hydrogen. In this change, the
change occurs in the molecular level. In one molecule of water,
there are 3 atoms in which two are hydrogen atoms and one is
oxygen atom. When water dissociates, the hydrogen and oxygen
atoms get separated forming separate hydrogen and oxygen
molecules. Therefore, dissociation of water into hydrogen and
oxygen is a chemical change.

177 Times' Crucial Science Book - 9

Main points to remember

1. A physical change is a temporary change in which no new substance is
formed.

2. A chemical change is a permanent change in which a new substance
with entirely new properties is formed.

3. A chemical reaction is a chemical change in which combination,
decomposition or exchange of molecules of matter takes place.

4. A chemical equation is the symbolic representation of actual chemical
reaction in terms of symbols and formulae.

5. The substances which undergo chemical change are called reactants
whereas the substances which are formed as a result of chemical
change in reactants are called products.

6. The chemical reaction which has unequal number of atoms in reactant
and product side is called unbalanced chemical reaction.

7. The chemical reaction which has equal number of atoms in reactant
and product side is called balanced chemical reaction.

8. The elements whose molecules are made up of two atoms are called
diatomic elements.

Exercise

1. Define:

a. Chemical equation b. Reversible reaction

c. Balanced chemical equation d. Reactant

e. Unbalanced chemical equation f. Product

g. Irreversible reaction h. Endothermic reaction

2. Answer these questions:
a. What is physical change? Give examples.
b. What is chemical equation? What are the ways to write
chemical equations?
c. What do you mean by word equation and symbolic equation?
Give two examples of each.
d. Define diatomic elements with examples.

3. Differentiate between:
a. Physical and chemical change
b. Endothermic and exothermic reaction
c. Reversible and irreversible reaction
d. Balanced and Skeleton chemical equation

4. Give reasons:
a. Melting of ice into water is a physical change.

Times' Crucial Science Book - 9 178

b. Change of milk into curd is a chemical change.
c. The reaction between hydrogen and nitrogen to give

ammonia is a reversible reaction.

5. Change the following word equations into balanced chemical
equations.

a. Hydrogen + Oxygen →Water.
b. Hydrogen + Nitrogen → Ammonia.
c. Hydrogen + Chlorine → Hydrochloric acid.
d. Carbon + Oxygen → Carbon dioxide.
e. Sodium + Oxygen →Sodium oxide.
f. Sodium + Chlorine → Sodium chloride.
g. Magnesium + Nitrogen → Magnesium nitride.
h. Lime water + Carbon dioxide → Calcium carbonate + Water.
i. Potassium chlorate → Potassium chloride + Oxygen.
j. Calcium hydroxide + Carbon dioxide → Calcium carbonate

+ Water.
k. Calcium hydroxide + hydrochloric acid → Calcium chloride

+ Water.
l. Calcium carbonate + Hydrochloric acid → Calcium chloride

+ Water +Carbon dioxide.
m. Copper carbonate → Copper oxide + Carbon dioxide.
n. Copper + Sulphuric acid →Copper sulphate + water +

Sulphur dioxide.

6. Balance the following skeleton equations.

a. Al + HCl → AlCl3 + H2 c. Ag + Br2 → AgBr
b. CuCO3 → CuO + CO2 d. H2SO4 + NaOH → Na2SO4 + H2O
e. KClO3 → KCl + O2 f. P4 + O2 → P2O5

Project Work

Several changes occur in your home, school or surrounding. Notice the
changes that occur at your home, especially in the kitchen. Classify these
changes as physical change and chemical change with justification.

Glossary
Reversible: that can be reversed
Temporary: lasting for a short time
Dissociation: break down

179 Times' Crucial Science Book - 9

Chapter

10 Solubility

Mikhail Tswett

He is also known by the name Mikhail Tsvet.
He is known for the discovery of adsorption
chromatography.

Estimated Periods :6

Objectives

At the end of the lesson, students will be able to:
• define a mixture and explain its types;
• explain different types of solutions;
• differentiate between saturated and unsaturated solution;
• explain solubility and solubility curve.

Mind Openers

• What is a mixture?
• What do you mean by homogeneous and heterogeneous mixtures?

• Are mixtures useful? Discuss.

Introduction
The substances are generally found in the form of mixtures. The
mixture can be formed among all three states of matter, i.e. solid,
liquid and gas.
A mixture is a mass formed by mixing two or more substances in
any proportion by weight without any change in the individual
property of the components. So, each component of mixture retains
its own identity and properties. The mass obtained by mixing salt
and water, water and sand, sand and salt, sugar and water, etc are
some examples of mixture.
Classification of mixtures
On the basis of the distribution of components, mixture can be
classified into two types. They are: heterogeneous mixture and
homogenous mixture.
1. Heterogeneous mixture
A mixture in which the components are not distributed uniformly
and they can be identified by naked eyes is called heterogeneous
mixture. Mixture of sand and water, sugar and sand, smoke, oily
water, etc are some examples of heterogeneous mixture.
2. Homogeneous mixture
A mixture in which the components are uniformly distributed and

Times' Crucial Science Book - 9 180

cannot be identified by naked eyes is called homogeneous mixture.
Mixture of salt and water, sugar and water, alcohol and water, etc
are some examples of homogeneous mixture.

Differences between homogeneous mixture and
heterogeneous mixture

Homogeneous mixture Heterogeneous mixutre

1.The components are distributed 1. The components are not distributed

uniformly. uniformly.

2. Each component of the mixture cannot be 2. The components of the mixture can

identified by naked eyes. be identified by naked eyes.

3. A solution is a homogeneous mixture. 3. A suspension is a heterogeneous
mixture.

On the basis of size of particles
Mixture can be classified into three types on the basis of size of
particles: suspension, colloids and solution.
1. Suspension
Suspension is a heterogeneous mixture in which the size of the
particles is 10−5cm or larger. The size of particles means the diameter
of the particles. The particles of suspension can easily be identified
with the help of simple microscope as well as naked eye. Muddy
water, sandy water, etc are some examples of suspension.
The components of a suspension can be separated by, sedimentation
decantation and filtration processes.
2. Colloid
A homogeneous mixture in which the size of the particles ranges
between 10−5 to 10−7cm is called colloid. Blood, urine, wax, milk, etc
are some examples of colloid. The components of a colloid can be
separated by using ultrafiltration and centrifugation methods.
3. Solution
A solution is a homogeneous mixture in which the size of particles
is 10−7cm or smaller. The components which are needed to make
solution are solute and solvent.

Solvent + Solute = Solution

A component of a solution which dissolves solute to form a solution
is called solvent. Solvent is present in relatively more amount than
solute. Water, alcohol, kerosene, petrol, benzene, etc are some
examples of solvent.

181 Times' Crucial Science Book - 9

A solute is a component of a solution which gets dissolved in solvent.
The amount of solute present in solution is relatively less than that
of solvent. Sugar, salt, alum, etc are some examples of solute.

Concentration of solution

The amount of solute present in given amount of solvent is called
concentration of solution.

On the basis of relative amount of solute present in solution, it can
be classified into two types: dilute and concentrated solution.

Dilute and concentrated solution

A solution which contains relatively less amount of solute in the
given amount of solvent is called dilute solution. It is represented
by ’dil.’ in chemical equation.

A solution which contains relatively more amount of solute in
the given amount of solvent is called concentrated solution. It is
represented by conc in chemical equation.

The terms dilute and concentrated are relative terms and they do
not tell the exact concentration of the solution.

Saturated, Unsaturated and Supersaturated Solution

A solution can be classified into three types on the basis of its ability
to dissolve more amount of solute. They are: unsaturated, saturated
and super saturated solution.

1. Unsaturated and saturated solution

The solution which can dissolve more amount of solute at a particular
temperature is called an unsaturated solution.

Take a glass of water, two spoonfuls of sugar and a glass rod. Add
a little amount of sugar in the water and stir it with the glass rod.
The sugar dissolves easily in it due to stirring. Again add some more
sugar to the glass and stir it. The sugar dissolves again. Since the
solution dissolves more amount of solute, it is unsaturated solution.
After adding sugar in the solution for some time, the solution at
that temperature is unable to dissolve any more amount of sugar.
This solution is called saturated solution.

Thus, the solution which cannot dissolve any more amount of solute
at a given temperature is called a saturated solution. A saturated
solution becomes unsaturated solution if temperature or the amount
of solvent is increased.

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Differences between unsaturated and saturated solution

Unsaturated solution Saturated solution

1. The solution which can dissolve 1. The solution which cannot

more amount of solute at a particular dissolve any more amount of solute

temperature is called unsaturated at a particular temperature is called

solution. saturated solution.

2. It contains less solute than it can 2. It contains solute as much as it

dissolve. can dissolve.

3. It becomes saturated when solute 3. It becomes unsaturated on

is added. heating.

2. Supersaturated solution
The solution which holds more solute than that is required to
make a saturated solution at a particular temperature is called
supersaturated solution. On heating, the saturated solution becomes
unsaturated and hence it can dissolve more amount of solute in it.
On cooling, the solution throws out excess amount of solute in the
form of crystals. Thus, the solution turns to be supersaturated
solution. This solution contains more solute dissolved than a
saturated solution at that temperature.

Activity 10 .1 To identify the unsaturated, saturated and suer saturated

solution

Materials required:

Three beakers each containing saturated, unsaturated and supersaturated
solution of sugar, glass rods, some crystals of sugar, sugar tongs, etc.

AB C

Procedure:
1. Place the beakers with solution on the table and name them as A, B andC.
2. Drop a crystal of sugar in each beaker and wait for sometime.

Observation:

Suppose you will observe that:
1. Beaker ’A’ easily dissolves the added sugar. Since this solution can

dissolve more solute, it is unsaturated solution.
2. Beaker ’B’ does not dissolve the added sugar even while stirring. Then

it issaturated solution.
3. Beaker ’C’ also does not dissolve the sugar. Instead of dissolving, the

size of crystal increases. Then, it is supersaturated solution.

Conclusion:

Unsaturated, saturated and super saturated solution can be identified with
the help of some crystals of the solute.

183 Times' Crucial Science Book - 9

Importance of solution
1. The plants obtain minerals from soil in the form of solution
2. Digested foods circulate in our body in the form of solution.
3. Most of the medicines, paints, cold drinks, etc are in the form of solution.
4. Oxygen is present as solute in water. Aquatic animals breathe in
dissolved oxygen.

5. Most of the chemical reactions take place in the solution form.

Solubility
Solubility of a solute at a given temperature is the amount of solute
required
to make a saturated solution in 100gm of solvent at that temperature.
Solubility of some substances are as follows:

SN Solute Solubility at
20°C

1. Sodium chloride 35

2. Copper sulphate 21

3. Sodium nitrate 88

4. Potassium nitrate 21

5. Calcium sulphate 0.20

Solved Numerical Problem 10.1

Calculate the solubility of sodium nitrate if 85g of sodium nitrate can be dissolved
in 100g of water at 20°C.

Solution:

Given,

Weight of solute (W1) = 85g
Weight of solvent (W2) = 100g
Solubility (S) = ?

Now,

W1 × 100 or, = 85
W2 × 100 = 85
S= 100

∴ The solubility of sodium nitrate at 20°C is 85.

Times' Crucial Science Book - 9 184

Solved Numerical Problem 10.2

The solubility of sodium chloride at 20°C is 35. What amount of salt is required to
make a saturated solution in 2 kg water at the same temperature?

Solution:

Given,

Solubility (S) = 35

Weight of water (W2) = 2 kg = 2000g
Weight of sodium chloride (W1) = ?

Now,

S = W1 × 100 or, 35 = W1 × 100
W2 2000

or, 35 × 2000 or, W1=700g
W1 = 100

∴ The required amount of sodium chloride is 700g..

Factors affecting solubility

The solubility of a substance is affected by the following factors:

1. Temperature

The solubility of a substance increases with the increase in
temperature. The increase in temperature of a solution increases
the kinetic energy of solute and solvent molecules. The increase in
kinetic energy makes the solvent molecules move apart creating
a large intermolecular space as well as it breaks up the solute
particles. Hence, the increased intermolecular space of the solvent
can hold more solute particles and the solubility of solute increases.

When the temperature is decreased, the kinetic energy of the solute
and solvent molecules decreases. As a result, the solvent molecules
come closer decreasing the intermolecular space. Also, the solute
particles do not break down as rapidly as in higher temperature.
Hence, less amount of solute is dissolved.

2. Size of solute particles

The finely powdered solute dissolves faster than its big lump. It

185 Times' Crucial Science Book - 9

is because small size of solute particles can be held easily by the
intermolecular space of the solvent.

3. Shaking or stirring

A solute dissolves faster if the mixture is shaken or stirred. Stirring
or shaking increases the kinetic energy of the molecules. It also
helps in the uniform distribution of solute particles in the solvent
molecules. Thus, the solute dissolves faster.

Solubility curve 130

The solubility of a substance increases with the 120
rise in temperature and decreases with the fall Solubility (100 gm of water)
of temperature. This variation of solubility can 110
be plotted on a graph. The curved line obtained Potassium
on a graph paper by plotting the solubility of 100 nitrate
a substance at different temperatures is called
solubility curve. In the solubility curve, the 90
temperature is plotted on X-axis and solubility
is plotted on Y-axis. 80

Solved Numerical Problem 10.3 70 MagSnoedsiiuumm chloride
60
50 chloride
40

30

20

10

0

10 20 30 40 50 60 70 80 90 100
Temperature°C

If the solubility of copper sulphate at 70°C and 20°C is 85 and 25 respectively,
what amount of copper sulphate will be crystallized out when its 50g saturated
solution is cooled from 70°C to 20°C?

Solution:
Given,

At 70°C,

Solubility = 85

Amount of solution = 50g

Let the weight of solute = x g

∴ Weight of solvent = (50 —x)g

Now, we have
Solubility = weight of solute × 100
weight of solevnt

x
or, 85 = 50 − x × 100 or, 100x=85(50 − x) or, 100x =4250−85

or, 185x =4250 ∴x =22.97

Times' Crucial Science Book - 9 186

Hence, amount of solvent= 50 − x = 50 − 22.97 = 27.03g

At 20°C,

Solubility = 25

Amount of solvent = 27.03g

Let, the weight of solute = y

Again, y
solubility = weight of solute × 100 or, 25 = 27.03 × 100
weight of solevnt

or, 100y = 27.03 x 25 ∴y = 6.75g
Now, the amount of solute is 6.75g.
∴ The amount of copper sulphate which gets crystallized out

= (22.97 − 6.75)g
= 16.22g

Information from a solubility curve
1. It helps us to find the solubility of a substance at different temperatures.
2. Solubility of different substances at a particular temperature can be
compared.
3. It shows the variation of solubility of a substance with temperature.
4. Solubility curve helps to indicate the temperature of formation of
hydrate.
5. It helps us predict which one will crystallize first if two or more solutes
are present in a solution.
6. It also helps to calculate the amount of solute precipitated during the
cooling of saturated solution.

Crystallization

The process of formation of crystals from a hot saturated solution of
a solute while cooling is called crystallization.

A crystal is a solid substance, which has a definite geometrical shape
with smooth surface and sharp edges. Crystals are pure substances.
They have fixed melting points. The crystals are arranged in a
specific three-dimensional pattern.

187 Times' Crucial Science Book - 9

Activity 10.2 To obtain the crystals of copper sulphate

Materials required:

Solution of copper sulphate, burner, copper sulphate crystals, tripod
stand, porcelain basin, glass rod etc.

Procedure:
1. Take a solution of copper sulphate in

a porcelain basin.
2. Heat the solution and dissolve more

crystals of copper sulphate.
3. Keep on heating to evaporate the

excess of solvent.
4. After heating for some minutes, insert a glass rod into the boiling
solution and draw it back. If tiny rough crystals are seen on the glass
rod, the solution is ready for crystallization. Now, remove the solution
from the burner and place it in a cool and dry place.
5. Leave the solution undisturbed for 1−2 hours. If you need crystals
fast, place the hot solution over a trough containing cold water. But
this process does not produce big crystals.

Observation:

Attractive blue crystals of copper sulphate are seen in the procelain
basin after cooling the solution.

Conclusion:

When a hot saturated solution of a solute is cooled down slowly,
crystals are formed. This process is called crystallization.

Types of solids
On the basis of shape, solid substances can be classified into two
categories. They are: amorphous solid and crystalline solid.
1. Amorphous solids
The solid substances which do not have fixed geometrical shape are
called amorphous solids. Rubber, plastics, soil, glass etc are some
amorphous solids.

2. Crystalline solids

The solid substances which have fixed geometrical shape with
smooth surface and sharp edges are called crystalline solids. They
have a fixed three dimensional arrangement of the constituent

Times' Crucial Science Book - 9 188

particles. They have definite melting points. Sodium chloride
(NaCl), alum, copper sulphate (CuSO4), etc
Learn and Write

1. Temperature is also mentioned while explaining solubility. Why?
Solubility of a solute changes with the change in temperature.
It increases with the rise in temperature and decreases with
the fall in temperature. Therefore, temperature should be
mentioned to specify the solubility.

2. Saturated solution becomes unsaturated on heating. Why?
When saturated solution is heated, there will be increase in
the intermo-lecular space of the solvent molecules. Due to it,
more solute particles can be accommodated. Thus, saturated
solution becomes unsaturated on heating.

3. Muddy water is a suspension. Why?
In muddy water, the size of particles is greater than 10−5cm.
The particles can be seen with naked eyes. Therefore, muddy
water is a suspension.

4. Salt solution is a homogeneous mixture. Why?
In salt solution, the components are mixed uniformly and
the particles cannot be seen with naked eyes. Therefore, salt
solution is a homogeneous mixture.

5. What happens when a supersaturated solution is cooled?
When a supersaturated solution is cooled, the solubility of the
solute decreases. Hence, some amount of solute gets separated
from the solution. The separated solute settles at the bottom of
the containing vessel.

6. Is it possible to prepare a saturated solution from an
unsaturated solution without adding solute?
It is possible to prepare a saturated solution from an
unsaturated solution without adding solute in the following
ways:
a. An unsaturated solution at higher temperature becomes saturated
if it is cooled.
b. If the solution in unsaturated at low temperature, it can be made
saturated by evaporating excess solvent from it.

7. What is a crystal? Write any two characteristics of a crystal.
A crystal is a solid substance, which has a definite geometrical
shape with smooth surface and sharp edges. A crystal has the
following characteristics:
a. A crystal has a sharp melting and boiling point.

189 Times' Crucial Science Book - 9

b. The crystals are arranged in a specific three dimensional pattern.

Main points to remember

1. A mixture is a mass formed by mixing two or more substances in any
proportion by weight without any change in the individual property of
the components.

2. A mixture in which the components are not distributed uniformly and
they can be identified by naked eyes is called heterogeneous mixture.

3. A mixture in which the components are uniformly distributed and
cannot be identified by naked eyes is called homogeneous mixture.

4. Suspension is a heterogeneous mixture in which the size of particles is
10−5cm or larger.

5. Colloid is a homogeneous mixture in which the size of unit particles
ranges from 10−5cm to 10−7cm.

6. Solution is a homogeneous mixture in which size of unit particles is
10−7cm or smaller.

7. A solution which has relatively less amount of solute in the given
amount of solvent is called dilute solution.

8. A solution which contains relatively more amount of solute in the given
amount of solvent is called concentrated solution.

9. The solution which can dissolve other more amount of solute at a given
temperature is called an unsaturated solution.

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

11. The solution which holds more solute than required to make saturated
solution at that temperature is called supersaturated solution.

12. The amount of solute required to make a saturated solution in 100g of a
solvent at a particular temperature is called solubility.

13. The curved line obtained on a graph paper by plotting the solubility of
a substance at different temperature is called solubility curve.

14. The process of formation of crystal from a hot saturated solution of a
substance while cooling is called crystallization.

15. A crystal is a solid substance which has a definite geometrical shape

with smooth surface and sharp edges.

Exercise

1. Choose the best alternative in each case.

a. Which of the following is a pure substance?

i. Air ii. Tea iii. Syrup iv. Oxygen

b. The size of particles in a solution is
i. Greater than 10-7cm in diameter

Times' Crucial Science Book - 9 190

ii. Equal to 10-7cm in diameter or less

iii. Equal to 10-7 cm in diameter

iv. All of these

c. Which of the following is an amorphous substance?

i. Sodium Chloride ii. Plastics

iii. Copper sulphate iv. Alum

d. Which of the following is a homogeneous mixture?

i. Muddy water ii. Oily water

ii. Sea water iv. Sandy water

e. Which of the following is a colloid?

i. Blood ii. Milk iii. Urine iv. All of these

2. Answer these questions in very short:

a. What is a mixture?

b. What is the size of the particles in a colloid?

c. Give any two examples of a colloid.

d. What do you understand by saturated and unsaturated

solution?

e. What is a supersaturated solution?

f. What do you mean by solute?

g. What is solubility?

h. What is the size of particles in a solution?

i. Define (with examples)

(i) Amorphous solids (ii) Crystalline solids

(iii) Crystals (iv) Solvent

(v) Colloid (vi) Suspension

3. Differentiate between:
a. Heterogeneous and homogeneous mixture.
b. Dilute and concentrated solution.
c. Unsaturated and saturated solution.
d. Solute and solvent
e. Crystalline and amorphous solids.

4. Give reasons:
a. Solubility of a substance increases with rise in temperature.
b. When a saturated solution is heated, it becomes unsaturated.
c. A solute dissolves faster if the mixture is stirred.

5. Answer these questions:
a. Explain homogeneous and heterogeneous mixtures with
examples.
b. Define dilute and concentrated solution with examples.

191 Times' Crucial Science Book - 9

c. What is solubility curve? Write down the information that
can be obtained from a solubility curve.

d. What is a crystal? Write down its properties.
e. How do you obtain crystals from a solution of copper

sulphate? Explain the process with diagram.

6. Solve the following numerical problems.
a. What is the solubility of sugar if 60 gram of sugar can be
dissolved in 200 gram of water at 50°C?

b. Calculate the solubility of common salt at 30°C if 80 gram

of salt dissolves in 250 gram of water to make saturated
solution.
c. If the solubility of copper sulphate at 50°C is 34, what

amount of copper sulphate will be required to make its

saturated solution in 80 gram water?

d. The solubility of sodium nitrate at 20°C and 30°C is 88
and 95 respectively. What amount of sodium nitrate will
crystallize out if 130 gram saturated solution is cooled from

30°C to 20°C?

Answers 6. a. 30 b. 32 c. 27.2g d. 4.67g

Project Work

Take suspension, colloid and solution in three different transparent
glasses. Find out differences among these three mixtures and write them
down.

Glossary
Proportion: ratio
Naked eyes: eyes without spectacles or contact lens
Ultrafiltration: filtration using powerful filter
Variation: change, difference
Precipitate: a solid particle that appears from solution

Times' Crucial Science Book - 9 192

Chapter

11 Some Gases

Henry Cavendish

He is known for the discovery of hydrogen and
measuring the earth's density, etc.

Estimated Periods :12

Objectives

At the end of the lesson, students will be able to:
• explain electronic configuration, laboratory preparation and properties of some

gases (hydrogen, oxygen and nitrogen);

• explain the uses of different gases.

Mind Openers

• Can you say how hydrogen gas is prepared?

• How are hydrogen, oxygen and nitrogen gases useful?

• Which gas is used as fire extinguisher? Why? Discuss.

Introduction

The earth is surrounded by atmosphere. The atmosphere consists
of a mixture of different gases such as nitrogen, oxygen, carbon
dioxide, etc. The gases can be elements or compounds. There
are 11 elements in the periodic table which are gases at normal
temperature and pressure. For example, hydrogen, helium, nitrogen,
oxygen, fluorine, neon, chlorine, argon, krypton, xenon and radon
are gaseous elements. Some compounds that are found in gas state
are carbon dioxide, methane, oxides of nitrogen, oxides of sulphur,
ammonia, etc. Among all gaseous state of matter, hydrogen is the
lightest and simplest gas.

Composition of air

Nitrogen and oxygen are the major components of air. They
constitute 99.03% of the total volume of air. Nitrogen alone occupies
78.08% of total volume of air and oxygen occupies 20.95% of total
volume of air. water vapour is pressent in variable amount in air.

193 Times' Crucial Science Book - 9

SN Gases in air Percentage by volume

1. Nitrogen (N2) 78.08%
2. Oxygen (O2) 20.95%
3. Carbon dioxide (CO2) 0.03%
0.93%
4. Argon (Ar)

5. Hydrogen (H2) 0.00005%
6. Ozone (O3) 0.000004%
7. Other gases 0.003945%

Carbon dioxide is the fourth most abundant gas in the atmosphere.
It occupies about 0.03% to 0.04% of the total volume of air. The
amount of carbon dioxide in the atmosphere is roughly balanced by
the processes of respiration and photosynthesis by plants. But the
amount of CO2 is increasing day by day in the atmosphere due to
human activities. The amount of other gases is somehow constant
in the atmosphere.

Hydrogen

Symbol: H Atomic number : 1

Molecular formula : H2 Atomic weight : 1.008 amu
Molecular weight : 2.016 amu Valency : 1

Position in the periodic table: Period - 1, Group - IA

p=1 p=1 p=1
n=0 n=0 n=0

Hydrogen atom Hydrogen molecule

Hydrogen was discovered by an English scientist Henry Cavendish
in 1766 AD. He named it as inflammable gas because of its property
to burn in air. The name hydrogen was given by Lavoisier in 1783
AD after studying its properties.
Hydrogen is the simplest and lightest gas. It is very reactive gas.
So, it is mostly found in the form of compound. It is abundant in
the form of water in rivers, lakes, ponds, oceans, seas, etc. Free
hydrogen gas is very rare. It forms less than 0.01% of total volume
of air in free state and is mostly present in the uppermost layer of
atmosphere.

Times' Crucial Science Book - 9 194

General methods of preparation of hydrogen

1. By the action of acids on metals

The metals such as zinc, magnesium, etc react with dilute acids to
produce hydrogen gas and metal salts.

Zn + H2SO4 → ZnSO4 + H2↑
Zn + 2HCl → ZnCl2 + H2↑
2. By the reaction of metals with alkalis

Zinc and aluminium react with alkalis such as sodium hydroxide to
produce hydrogen gas.

Zn + 2NaOH → Na2ZnO2 + H2↑
Sodium zincate

2Al + 2NaOH + 2H2O → 2NaAlO2 + 3H2↑
Sodium meta aluminate

3. From water

Hydrogen gas can be produced by electrolysis of water in the presence
of sulphuric acid as electrolyte. Oxygen gas is also produced in this
reaction.

2H2O electrolysis 2H2↑ + O2↑

4. Hydrogen gas can also be prepared by the action of reactive
metals like sodium, calcium, potassium, etc on water.

2K + 2H2O → 2KOH + H2↑
2Na + 2H2O → 2NaOH + H2↑
Laboratory preparation of hydrogen gas

Principle:

Hydrogen gas is prepared in the laboratory by the action of dilute
sulphuric acid on granulated zinc. Granulated zinc is the impure
zinc in the form of small pieces.

Zn + H2SO4 → ZnSO4 + H2↑
In this process, impure zinc should be used instead of pure zinc
because the action of pure zinc on the dil. H2SO4 is very slow. The
impurity in zinc makes the reaction faster.

195 Times' Crucial Science Book - 9

Apparatus required

Woulfe’s bottle, thistle funnel, delivery tube, gas jar, water trough,
beehive self etc.

Chemicals required:
1. Granulated zinc (Zn)
2. Dilute sulphuric acid (dil. H2SO4)
3. Water (H2O)

Dilute sulphuric acid

Thistle funnel Hydrogen gas
Gas jar

Delivery tube

Woulfe's bottle Trough
Granulated zinc
Water
Beehive shelf

Laboratory preparation of hydrogen gas

Procedure
1. Put some amount of granulated zinc in Woulfe’s bottle.
2. Connect a thistle funnel and delivery tube to two openings of
Woulfe’s bottle with the help of cork.
3. Make the open end of the delivery tube lie below the gas jar
through the beehive shelf.
4. Pour dilute sulphuric acid into the Woulfe’s bottle through
thistle funnel.

Now, dilute sulphuric acid reacts with granulated zinc and
hydrogen gas is produced. Hydrogen gas travels through delivery
tube and is collected in the gas jar by downward displacement
of water.

Hydrogen gas is collected in the gas jar by downward displacement
of water because it is lighter than water and it does not react
with water. Also, it is the least soluble in water.

Precautions
1. The apparatus should be airtight.
2. Impure granulated form of zinc should be used.
3. The end of the thistle funnel should be dipped in the solution.
4. The end of the delivery tube should not touch the solution
within Woulfe’s bottle.

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