Optional Science Book Class 10

4Step

1. Describe the mole concept.
2. What are quantum numbers? Write very short about all quantum numbers.
3. Describe about the principal quantum number.
4. What is azimuthal quantum number? Describe the information obtained from the

azimuthal quantum number.
5. Describe the information which are obtained from the magnetic quantum number.
6. Explain in short about the spin quantum number.
Multiple choice questions

1. What is the use of amu?

a) To measure the length of very small distance

b) To measure the time of very small duration

c) To measure the mass of very small particles

d) To observe the very small particles

2. Which of the following is a Avogadro’s number?

a) 6.022 × 1023 atoms b) 6.022 × 1023 molecules

c) 6.022 × 1023 ions d) All of above a, b and c

3. How many quantum numbers are there?

a) 1 b) 2 c) 3 d) 4

4. Which symbol is used to denote principal quantum number?

a) n b) l c) m d) s

5. Which quantum number explains the orientation of the orbitals?

a) Principal quantum number b) Azimuthal quantum number

c) Magnetic quantum number d)Spin quantum number

6. What is the value of “l” for n=1

a) 3 b) 2 c) 1 d) 0

7. What is the normality for one molar HCl?

a) 1 b) 2 c) 3 d) 4

8. What does it mean by N/10?

a) Decimolar solution b) Decinormal solution

c) Normal solution d) Molar solution

***

CHEMISTRY 0Optional Science - 10 151

UNIT PERIODIC
TABLE AND
8 PERIODIC LAWS

..............,......................,.........................................................

About the inspiring personality

Wolfgang Pauli born in 25 April 1900 AD in Vienna, Austria and died in 15 December
1958 AD in Switz. Pauli propounded the Pauli Exclusion Principle. For this great work he
was awarded with a Nobel Prize in 1945 AD. Pauli along with Neils Bohr also formulated
the Aufbau principle. Not only this, Pauli also made a major contribution in the quantum
mechanics.

Syllabus issued by CDC Learning objectives:

• Mendeleev's and Modern periodic After completing the study of this unit, students

Table (s, p, d and f concept) will be able to:

• Application of periodic table • Show the electronic configuration based on s,

• Aufbau principle p, d and f.

• Electronic configuration • Introduce s-block, p-block, d-block and f-block

• Valency and variable valency elements along with their characteristics.

• Periodic variation • Define valency, variable valency, and radical.

• Atomic size, Ionization • Explain variation of atomic size,
ionization potential, electron affinity and
potential, Electron affinity and
electronegativity in the period and group.
Electronegativity

Key terms and terminologies of the unit

1. Classification of elements: The grouping of elements according to their similarities
and differences is called the classification of elements.

2. Mendeleev's periodic law: According to Mendeleev's periodic law, the physical
and the chemical properties of elements are the periodic functions of their atomic
weights.

3. Mendeleev's periodic table: The table or chart which is obtained after arranging
elements on the basis of increasing atomic weight is called Mendeleev's periodic table.

4. Modern periodic law: According to Modern periodic law, the physical and the
chemical properties of elements are the periodic functions of their atomic numbers.

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5. Modern periodic table: The table or the chart which is obtained after arranging
elements on the basis of their increasing atomic numbers is called the modern
periodic table.

6. Periods: The horizontal rows of the periodic table are called periods. In periods,
elements with gradual change in characters and atomic numbers are kept.

7. Groups: Vertical columns of the periodic table where elements with similar
characteristics are kept are called groups.

8. Periodic variation: The variation in the properties of elements across a period or down
the group of the modern periodic table in a periodic manner is called periodic variation.

9. Atomic size: The distance between nucleus and the outermost shell of an isolated
gaseous atom is called atomic size.

10. Ionization potential: The amount of energy required to remove the loosely bonded
outermost electron from an isolated gaseous atom is called ionization potential.

11. Valence electrons: The total numbers of electrons present in the outermost shell of
an atom are called valence electrons.

12. Valency: The total number of electrons gained or lost or shared during the chemical
combination is called valency.

13. Electron affinity: The amount of energy released by an atom when an electron is
added to its valence shell is called electron affinity.

14. Electronegativity: The amount of energy required for an atom to attract foreign
electrons towards itself is called electronegativity.

15. Aufbau principle: According to Aufbau principle, “filling of electrons in the sub-
shells always occurs from the lower energy level to the higher energy level.”

16. Electronic configuration: The systematic distribution of electrons in different shells
and sub-shells is called electronic configuration.

17. s-block: The block of alkali metals (IA) and alkaline earth metals (IIA) which is
present at the left-hand side of the periodic table is called s-block.

18. s-block elements: The elements which have the last electron in "s" sub-shell are
called s-block elements.

19. p-block: The block of metals, non-metals, metalloids and inert gases which are
present in IIIA, IVA, VA, VIA, VIIA and zero groups is called p-block.

20. p-block elements: The elements which have the last electron in "p" sub-shell are
called p-block elements.

21. d-block: The block of transitional metals which is present in between s-block and
p-block of the modern periodic table is called d-block.

22. d-block elements: The elements which have the last electron in d-sub shell are
called d-block elements.

23. f-block: The block of lanthanides and actinides which is present separately below
the main periodic table is called f-block.

24. Valence shell: The outermost shell of an atom is called valence shell.

25. Valence electrons: The total number of electrons present in the valence shell of an
atom are called valence electrons.

26. Valency: The number of electrons gained, lost or shared by an element during the
chemical reaction is called valency.

27. Variable valency: More than one valency of an element is called variable valency.

CHEMISTRY 0Optional Science - 10 153

Introduction

In the initial stage of the study of science, very few elements were known to us. These elements
could be studied individually. With the rapid development of scientific research, more and
more new elements with their different physical and chemical properties are discovered. Till
now, about one hundred eighteen elements with different atomic masses and numbers are
known to us. So, it is quite difficult to study these elements separately. To solve this problem,
chemists realized that elements with similar and dissimilar properties should be divided into
different groups. So, classification of elements refers to the process of grouping of elements on
the basis of their similarities and dissimilarities. The grouping of elements according to their
similarities and differences is called the classification of elements.

Classification is a scientific process. It makes the study of elements easy, fast and clear. At
the beginning of the classification, elements were divided into three groups on the basis of
their physical states like solid, liquid and gas. After physical state, elements were classified
according to the metallic characters such as metals, non-metals and metalloids. In the field
of classification, Antony Lavoisier was the first chemist, who studied elements on the basis
of their metallic nature and categorized them into two groups, i.e. metals and non-metals. At
that time, chemists knew only thirty-two elements. Lavoisier's classification could not last
long, as some elements having both the characters of metal and non-metal were discovered
later. There is no doubt that different chemists have continuously tried to classify elements
systematically for the better and advanced study of elements.

Mendeleev's Period Law

Russian chemist Dmitri Mendeleev followed the idea which was given by John Newland in
1864 AD. At this time, about 63 elements were discovered. Mendeleev studied the physical and
chemical properties of these elements along with their compounds. After the complete study
of these elements, he made a sequence of elements in an ascending order of atomic weights.
similar to the sequence which was given by John Newland in 1864 AD. In this sequence, he
found that the elements with similar properties occur at a regular interval. This observation
led Mendeleev to conclude a law, which is called Mendeleev periodic law.

According to Mendeleev's periodic law, the physical and the chemical properties of elements
are the periodic functions of their atomic weights. It means that when elements are arranged
on the basis of increasing atomic weights, the elements having similar physical and chemical
properties are repeated after a regular interval.

Memory Plus

The meaning of periodic function is that the properties of elements go on changing when
atomic mass is increased but after a certain regular interval; they repeat the characters of
the previous element.

Mendeleev's Periodic Table

After the study of physical and chemical properties of elements, Mendeleev proposed a
periodic table called Mendeleev’s periodic table. The table or chart which is obtained after
arranging elements on the basis of increasing atomic weight is called Mendeleev's periodic
table.

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In Mendeleev's periodic table, there are seven horizontal rows called periods, and eight vertical
columns called groups. Sometimes periods are also called series and groups are also called
families. In this table, elements with similar physical and chemical properties fall under the
groups and elements with gradual change in characteristics fall under the periods. During the
time of Mendeleev, only 63 elements were known to us, so he left some gaps for undiscovered
elements. For example, scandium, gallium, germanium, etc. were discovered after the design
of Mendeleev's periodic table. To keep these elements, Mendeleev left gaps by saying eka-
boron, eka-aluminium and eka-silicon.

A part of Mendeleev's periodic table is given below:

Group I Group II Group III Group IV Group V Group VI Group VII Group VIII

Period H
1

Period Li Be B C NOF
2

Period Na Mg Al Si P S Cl
3

Period K Ca 1* Ti V Cr Mn Fe Co

4 Cu Zn 2* 3* As Se Br Ni Rh
Pd
Period Rb Sr Y Zr Nb Mo 4* Ru
lr
5 Ag Cd In Sn Sb Te l Pt

Period Cs Ba La Hf Ta W Re Os

6 Au Hg Th Pb Bi Po At

Name given by Mendeleev: 1* Eka - Aluminium,

2* Eka - Boron, 3* Eka - Silicon, 4* Eka - Manganese

Characteristics of the Mendeleev's Periodic Table

1. Mendeleev arranged elements on the basis of increasing atomic weight.

2. In this table, there are seven horizontal rows (periods) and eight vertical columns (groups).

3. In Mendeleev's periodic table, he left gaps for undiscovered elements. For example;
scandium, gallium, germanium, etc.

4. Each group of Mendeleev's periodic table was further divided into two sub- groups (sub-
group A and sub-group B).

5. Inert gases like He, Ne, Ar, Kr, Xe and Rn were not discovered at the time of Mendeleev,
so there was no separate place for them.

Merits or Advantages of Mendeleev's Periodic Table

1. Mendeleev's periodic table was the first scientific as well as systematic classification of
elements, which made their study easy, fast and systematic.

2. Many elements were not discovered during the time of Mendeleev. So he left gaps for
these elements. For example, scandium, gallium and germanium were discovered later on.

3. The atomic weight of some elements like beryllium, platinum, gold, uranium, etc. was

not confirmed before Mendeleev's periodic table came into existence. With the help of

this table, atomic masses of these elements were corrected.

CHEMISTRY 0Optional Science - 10 155

4. Mendeleev's periodic table encouraged subsequent scientists to discover new elements as
he left some gaps for them.

Demerits of Mendeleev's Periodic Table
1. In Mendeleev's periodic table, he could not arrange hydrogen properly because sometimes

it gains electrons like those of halogens (group VIIA) and sometimes it loses electrons
like those of alkali metals (group IA).
2. Mendeleev could not arrange isotopes properly as they have the same atomic numbers
but different atomic weight.

Memory Plus

According to the Mendeleev, isotopes of one element should have different places as
they have different atomic weight. For example; there are three isotopes of carbon,
such as C612, C613, and C614. On the basis of Mendeleev's periodic table, there should
be three different places for them as they have different weight.

3. In some cases of Mendeleev’s periodic table, elements with more atomic weight were
placed first than those with less atomic weight. It was completely wrong according to his
periodic law. For example: cobalt with atomic weight 58.9 was placed before nickel with
atomic weight 58.6.

4. In some places of Mendeleev's periodic table, chemically dissimilar elements were
grouped together. For example; very less reactive coinage metals (Cu, Ag, and Au) and
highly reactive alkali metals (Li, Na, K, Rb and Cs) were placed in the same group (first
group).

5. In Mendeleev's periodic table, chemically similar elements were placed in different
groups. For example; gold is similar to platinum, copper is similar to mercury, barium is
similar to lead, etc. Unfortunately they were placed separately.

6. Mendeleev could not arrange lanthanides and actinides properly in the fixed position.

7. Mendeleev's periodic table was unable to explain the atomic properties of the elements
like valency, metallic characters, ionization potential, electronegativity, reactivity, etc.

8. All groups are divided into sub-groups but the 8th group is not divided into sub-groups.

Modern Periodic Law

To overcome the defects of Mendeleev's periodic table, a group of chemists led by Henery
Moseley studied the physical and chemical properties of elements in 1913 AD with the help of
x-ray spectra. They concluded that atomic number is more fundamental property of elements
rather than their atomic weight. This conclusion was quite satisfactory for an atom because
chemical properties of the elements depend upon the number of electrons present in its atom.
With the help of the above study, Henery Moseley modified the Mendeleev's periodic law.
This modified law is called modern periodic law.

According to Modern periodic law, the physical and the chemical properties of elements are the
periodic functions of their atomic numbers. It means that when elements are arranged on the basis
of their increasing atomic number, at a regular interval, the properties of elements are repeated.

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While making the sequence of elements based on their increasing atomic number, it is found
that the elements which have similar physical and chemical properties lie in the same vertical
column. Similarly, the elements with gradual change in properties lie in the same horizontal
row one after another.

Modern Periodic Table

After complete study of the physical and chemical properties of elements and their compounds,
the team of Henery Moseley proposed the modern periodic law, which is modified form of the
Mendeleev's periodic law. With the help of this modern periodic law, Bohr arranged elements
on the basis of increasing atomic numbers. As a result, he found a table called the modern
periodic table.

The table or the chart which is obtained after arranging elements on the basis of their
increasing atomic numbers is called the modern periodic table. In the modern periodic table,
there are seven horizontal rows called periods. In these periods, elements with gradual change
in characters lie one after another. Similarly, in modern periodic table, there are 18 groups. In
these groups, elements with similar physical and chemical properties fall one below another.

Characteristics of the Modern Periodic Table

The main features of the modern periodic table are given below:

1. In modern periodic table, elements are arranged on the basis of their increasing atomic
numbers.

2. In this table, there are seven periods and 18 groups.

3. Hydrogen is kept in the first period and the first group.

4. Inert gases like He, Ne, Ar, Kr, Xe and Rn are placed in a separate group (zero group or
18th group) at the extreme right side of the table.

5. In the modern periodic table, metals are kept at the left-hand side, non-metals at the
right-hand side and metalloids are placed between metals and non-metals.

6. In this table, hydrogen is placed in 1st( IA) group and lanthanides and actinides are placed
in a separate box below the main table.

7. Reactive metals, less reactive metals, transitional metals, reactive non-metals and less-
reactive non-metals are kept separately. But, metalloids are not perfectly grouped. They
are roughly scattered in the right side of the periodic table.

8. The whole periodic table is based on atomic sub-shells and divided into four blocks. They
are s-block, p-block, d-block and f-block.

9. Most of the synthetic elements that are prepared in the laboratory are generally present
in the last periods (6th and 7th).

Memory Plus

In modern Periodic Table, the s-block elements are kept on the extreme left, p-block
elements are on the extreme right, d-block elements are placed between s-block and
p-block and f-block elements are kept separately below the main table.

CHEMISTRY 0Optional Science - 10 157

Advantages of the Modern Periodic Table

In the modern periodic table, elements are classified on the basis of increasing atomic numbers.
This system of classification helps to solve the defects of Mendeleev's periodic table in the
following ways.

1. The position of hydrogen is fixed: The accurate position of hydrogen has not yet been
solved completely. But as its atomic number is the least number (i.e. one), it is kept in IA
(group 1) group of the modern periodic table along with alkali metals.

2. The position of isotopes is fixed: Isotopes are the atoms of the same elements, which
have the same atomic numbers but different atomic weights. Therefore, isotopes of one
element are placed together in the same group. For example, there are three isotopes of
carbon (i.e. C612, C613 and C614) which have same atomic numbers. They are placed together
in IVA group of the modern periodic table.

3. The wrong position of the elements has been corrected: When elements are arranged
on the basis of their increasing atomic numbers, the wrong position of argon, potassium,
nickel, cobalt, etc. is solved automatically without changing their own places.

4. The alkali metals and coinage metals are kept separately: In the modern periodic table,
the most reactive alkali metals are kept in IA group (group 1) whereas the least reactive
coinage metals are kept in IB group (group 11).

5. Lanthanides and actinides are kept below the main table: In the modern periodic table,
fourteen elements of lanthanide series and fourteen elements of actinide series are placed
in a separate box below the main periodic table.

6. The elements are arranged according to the electronic configuration of their orbitals,
viz. s, p, d and f. This divides the periodic table into four main blocks, i.e. s, p, d and f
block.

Why hydrogen can be placed in group IA along with alkali metals
i. Both hydrogen and alkali metals have one valence electron.
ii. Both hydrogen and alkali metals can lose one electron to form electropositive radicals.
iii. Both hydrogen and alkali metals can form halides, oxides, sulphides, etc.
iv. Both hydrogen and alkali metals can react with halogens, oxygen and sulphur.

Why hydrogen can be placed in group VIIA along with halogens
i. Both hydrogen and halogens have valency one.
ii. Both hydrogen and halogens can gain one electron to form electronegative radicals.
iii. Both hydrogen and halogens need one electron to complete their stable electronic

configuration.
iv. Both hydrogen and halogens can react with metals.
v. Both hydrogen and halogens can make diatomic molecules.
vi. Both hydrogen and halogens occur in gaseous state.

Drawbacks of the modern periodic table

The modern periodic table has corrected many mistakes of the Mendeleev’s periodic table,

still it has some drawbacks. Some weaknesses of the modern periodic table are as follows:
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CHEMISTRY

1. The position of hydrogen is still controversial as it can be placed both in group IA group
(group 1) and VIIA (group 17).

2. The lanthanides and actinides do not have a clear group. But their position is still separate
from the rest of the elements.

3. In modern periodic table, helium is kept in p-block but its last electron falls into the
s-orbital.

Position of Elements in Modern Periodic Table
1. Position of hydrogen: Hydrogen has only one shell with one electron. Due to its one

atomic number, it is placed under IA group of the modern periodic table along with alkali
metals. But, hydrogen shows both the characteristics of alkali metals and halogens.

2. Position of metals, non-metals and metalloids: In the modern periodic table, metals are
present at left-hand side, non-metals are present at right-hand side and metalloids are
placed between metals and non-metals.

3. Position of lanthanides and actinides: A group of fourteen elements starting from
cerium (Ce58) to lutetium (Lu71), which have similar characters with lanthanum (La57) are
called lanthanides. Similarly, a group of fourteen elements starting from thorium (Th90) to
lawrencium (Lr103), which have similar characters of actinium (Ac89) are called actinides.
These twenty eight elements (14 lanthanides and 14 actinides) are kept separately in the
separate box below the main periodic table. These elements have dissimilar characters
with other elements of the rest of the periodic table.

4. Position of inert gases: He, Ne, Ar, Kr, Xe and Rn are called inert gases as they do not
participate in chemical reactions. These elements have zero valency. They are kept in the
zero group (group 18) at the extreme right-hand side of the modern periodic table.

Differences between Modern and Mendeleev’s Periodic Table

Modern Periodic Table Mendeleev’s Periodic Table

1. Modern periodic table is based on 1. Mendeleev’s periodic table is based on
increasing atomic number. increasing atomic weight.

2. In this table, there are seven periods 2. In this table, there are seven periods
and eighteen groups. and eight groups.

3. In this table, hydrogen, isotopes 3. In this table, hydrogen,isotopes
lanthanides and actinides are kept lanthanides and actinides are not kept
properly. properly.

Memory Plus

Alkaline earth metals: Elements which are present in IIA group of the modern periodic
table are called alkaline earth metals. They are found on the earth’s crust and react with
water to give hydroxides. They include magnesium (Mg), calcium (Ca), etc.

Mg + 2H2O → Mg(OH)2 + H2
Ca + 2H2O → Ca(OH)2 + H2

CHEMISTRY 0Optional Science - 10 159

Importance of the Periodic Table

Periodic table is very important to make the study of elements easy, fast and systematic. It
is a useful tool for the students, chemists and scientists. The modern periodic table not only
mentions the name, atomic number and position of the elements but also their atomic weights,
electronic configuration, nature, isotopes and other characteristics. Thus, periodic table is very
useful to us. Some of the important points are listed below.

i. The complete information about the elements like the number of electrons, number of
protons, electronic configuration, atomic number, atomic weight, valency, nature, block,
etc. can be obtained from the periodic table.

ii. We can study the properties of elements with the help of the periodic table, .

iii. The periodic variations such as atomic number, valency, atomic radius, ionization
potential, electronegativity, electron affinity, etc. can be studied easily with the help of
periodic table.

iv. The characteristics of the elements can be known by knowing the exact position of the
element in the periodic table.

v. The properties of the elements that are not yet discovered can also be predicted.

vi. The periodic table makes the arrangement of elements systematic, scientific and clear.

Memory Plus

The widely accepted modern periodic table is also called long form of periodic table. It is
given by IUPAC.

Periods and Groups of Modern Periodic Table

Periods

In the modern periodic table, there are seven horizontal rows called periods. In each
period, there are fixed numbers of elements with gradual change in atomic numbers and
characteristics. As we move from left to right in the period, the size of atoms decreases but the
atomic number increases. Thus, the horizontal rows of the periodic table are called periods. In
periods, elements with gradual change in characters and atomic numbers are kept. The very
short description of the different periods is given in the table.

S.N. Period Number of Elements Types of Period
1. First 2 Very short period

2. Second 8 Short period
3. Third 8 Short period
4. Fourth 18 Long period
5. Fifth 18 Long Period
6. Sixth 32 Very long period
7. Seventh 26 Very long but incomplete period

Groups

The vertical columns of the periodic table are called groups. There are a total 18 groups in
the modern periodic table. In these groups, elements with similar characteristics fall one
below another. Thus, vertical columns of the periodic table where elements with similar
characteristics are kept are called groups.

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Except zero and VIII groups, all other groups are further divided into sub-groups A and B. In
the eighth group, there are three vertical columns but they are not represented by A, B and C.
Altogether, there are a total of eighteen vertical columns, which are also called 18 groups in
the long form of the modern periodic table.

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18
IA IIA IIIB IVB VB VIB VIIB VIII VIII VIII IB IIB IIIA IVA VA VIA VIIA 0

Fact with reason

Sodium and potassium are kept in the same group. Why?
Sodium and potassium have the same number of valence electron and valency. They also
form the similar type of radical. So they are kept in the same group.

~

Differences between Periods and Groups

Periods Groups

' 1. The vertical columns of the periodic

1. The horizontal rows of the periodic

table are called periods. I I table are called groups. -

I

2. There are a total of seven periods in the 2. There are a total of nine groups in

modern periodic table. I modern periodic table.

- 13. I-
In the group top to bottom, the
3. In the period left to right, the atomic
L___J size decreases. I atomic size increases. __J

Periodic Variation

In the modern periodic table, there are seven horizontal rows and 18 vertical columns.
Horizontal rows are called periods and vertical columns are called groups. The elements
present in a group are similar to each other in many aspects such as valence electrons, valency,
metallic nature, radical formation, etc. Similarly, the elements across a period also have some
similarities such as valence shells. Not only similarities, elements present in the same group
and same period also have many differences. In a group, elements have different atomic
size, ionization potential, electronegativity, etc. The variation in the properties of elements
in periods and groups are periodic. It means that characteristics repeat at a regular interval.
Thus, the variation in the properties of elements across a period or down the group of the
modern periodic table in a periodic manner is called periodic variation.

Memory Plus

Alkali metals: Elements which are present in IA group of the modern periodic table are
called alkali metals. Examples: lithium (Li), sodium (Na), potassium (K), etc. These metals
react with water to give hydroxides.

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

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Periodic variation of the some properties

1. Atomic size: The distance between nucleus and the outermost shell of an isolated gaseous

atom is called atomic size. As we go from top to bottom in a group, the atomic size

increases as the number of additional shell increases. Similarly, as we go from left to right

in group, the atomic size decreases as nuclear force of attraction increases.

Atomic Radius Decreases ..

Oroup

IA 2A 3A 4A SA 6A 7A 8A
(I) (2)
(13) (1 4) ( 15)(16) {17) {18)
G
Q

H He

2~
Li Be

5

fig:Variation of atomic size in a group and period of the modern periodic table

Fact with reason

Atomic size increases in a group and decreases in a period. Why?

The number of additional shell increases in a group as we go from top to bottom.
So, atomic size increases in a group. However, in a period, the numbers of shells are
fixed but the number of protons increases as we go from left to right. The additional
proton increases the nuclear force of attraction and decreases the atomic size.

2. Ionization potential (or Nuclear power)

The amount of energy required to lncr~ · · ·gY, ~
remove the loosely bonded outermost H,
electron from an isolated gaseous • H... ■ C NO F N, I --7
atom is called ionization potential.
Ionization potential is measured in • u ~' II I® \
electron volt (eV). Its value is more in
the small size and less in the big size. t~••e• iK II ' '_,I
Therefore, in the period left to right,
ionization potential increases as the -~) II @

,

____ ,,,,, I

fig:variation of ionization energy in periodic table

atomic size decreases. Similarly, the value of ionization potential decreases in the group

top to bottom, as atomic size increases.

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Fact with reason

Why does ionization potential increase in the period left to right and decrease
top to bottom in the group?

In the period left to right, the atomic size decreases and nuclear force of attraction
increases. It makes difficult to remove outermost electron from the atom. So, the
ionization potential increases in the period left to right. Similarly, in the group,
atomic size increases. It makes easy to remove the outermost electron from the
atom. So, the ionization potential decreases top to bottom in the group.

3. Electron affinity

The amount of energy released by an

atom when an electron is added to its

valence shell is called electron affinity.

Conceptually, electron affinity is opposite

to the ionization potential. When an

electron is added to the valence shell of

a neutral atom, the atom releases energy.

Therefore, the electron affinity is the

measurement of the electron pulling fig:variation of electron affinity in periodic table
ability of an atom. Higher the value of

electron affinity, higher will be the ability to pull an electron and vice-versa. Its unit is

eV. Electron affinity increases in the period left to right because atomic size decreases.

Similarly, electron affinity decreases in the group top to bottom because atomic size

increases.

Fact with reason

Why does electron affinity increase in the period left to right and decrease top to
bottom in the group?

In the period left to right, the atomic size decreases and nuclear force of attraction
increases. It makes easy to pull an electron towards the atom. So, the electron affinity
increases in the period left to right. Similarly, in the group, atomic size increases and
nuclear force of attraction decreases. It makes difficulty to pull an electron towards
the atom. So, electron affinity decreases top to bottom in the group.

4. Valence electrons Number of Valence Electrons
2 34
The total numbers of electrons present in the
outer most shell of an atom are called valence fig:variation of valence electrons in periodic table
electrons. It increases in period left to right and
remains the same in the group top to bottom.

5. Valency

The total number of electrons gained or lost or shared during the chemical combination

is called valency. In a period left to right, valency first increases up to four and then

decreases up to zero. Similarly, in the group top to bottom, it remains the same.

CHEMISTRY 0Optional Science - 10 163

For example:

Element: Li Be B C N O F Ne
Valency 12343 2 10

6. Electronegativity

The amount of energy required for an atom to attract foreign electrons towards itself is

called electronegativity. It increases as the atomic size decreases and it decreases as the

atomic size increasing. In period left to right, electronegativity increases because atomic

size decreases and nuclear force of attraction increases. Similarly, in group top to bottom

electronegativity decreases because atomic size increases and nuclear force of attraction

decreases.

l&WIIHll4¥1@111119.

fig:variation of electronegativity in periodic table

Fact with reason

Electronegativity increases in the period left to right and decreases in the group
top to bottom. Why?

In the period left to right, the atomic size decreases and nuclear force of attraction
increases. It makes easy to attract the foreign electron towards the atom. So, the
electronegativity increases in the period left to right. Similarly, in the group, atomic
size increases and nuclear force of attraction decreases. It makes difficulty to attract
the foreign electron towards the atom. So, the ionization potential decreases top to
bottom in the group.

7. Metallic characters: Metallic character increases from top to bottom in the group and
decreases from left to right in the period.

Metallic character decreases

□ bM"' -1-8 ,
T □</l
Me<allolds
--. "• ' .u~
2 □ Nonmetob 14 15 16 17 He

-;- 6 9 10

Be BC N0 F Ne
."s' . ,.u~
11 12 13 1p5 1s6 a17 II
Na Mg 3 Az
Al SI
"...' i· " .. "• .'. .. .. .. .. .. :. ., ..., .....,., ,. n N,,b ",. ,. ,,.. ,. ,.., c"'e' A:il.u2l58910 11u

19 20 21 22 z, Cr ~ 26 Z1 29 31 ~ 35 ~
V
K Ca Sc Fe Co Ni Cu Zn Ga Br

Rb Sr y Zr 47 Te I Xe

Mo Tc Ru Rh Pd Ag Cd In Sn

II 12 13

Pt Au Hg TI Pb Bi Po Al Rn
56

" "-6 Ca Ba
72 76 77

La Hf Ta Re Os Ir
.. ..., IOI 105 106 107 IOI 109 110 111 112 114 116w
s Fr Ra Ac Rf Db 5g Bh Ha Mt
-.5 ~
~ Lanthoruda C~ e~
Pr N~ d 61 S~ m~ Eu G" d TMb D" y HQo a T~ m Y~ b L~ u
Pm
Er

Adlnlda T~ h P~ a n N"p PHu A"m C~ m B~ k C~ f EWs F~ m M~ d N~ o WLr

U

fig:variation of metallic characters in periodic table

164 Optional Science - 10

CHEMISTRY

Sub-Shells (or Orbitals)

On the basis of 2n2 formula, when there are two electrons in the outermost shell of an atom,
it is called duplet (for example helium). Similarly, when there are eight electrons in the
outermost shell, it is called octet such as Ne, Ar, Kr, Xe and Rn. The electronic configuration
of all the elements cannot be explained by the 2n2 formula. Thus, the concept of sub-shells
was proposed. According to this concept, each and every main shell contains one or more
sub-shells, which are denoted by s, p, d and f. Different shells contain different numbers of
sub-shells. Thus, the shells which are present within the main shell are called sub-shell. They
are also called sub-energy level or orbitals and denoted by s, p, d and f. The shells with their
sub-shells are given in the table.

Main Shell Total sub-shells Sub-shells (or Orbitals)
Main shell first (n=1) 1 s
Main shell second (n=2) 2
Main shell third (n=3) 3 s and p
Main shell fourth (n=4) 4 s, p and d
s, p, d and f

The s- sub-shell may contain maximum two electrons.
The p- sub-shell may contain maximum six electrons.
The d- sub-shell may contain maximum ten electrons.
The f- sub-shell may contain maximum fourteen electrons.

On the basis of the above data, we can co-relate the 2n2 formula with sub-shells(s, p, d and
f). The main shell, its sub-shells and the maximum numbers of electrons present there can be
summarized as follows:

Main shell Sub-shells Total electrons
K (n = 1) 1s 2
L (n = 2)
M (n = 3) 2s, 2p 2+6=8
N (n = 4) 3s, 3p, 3d 2 + 6 + 10 = 18
4s, 4p, 4d, 4f 2 + 6 + 10 + 14 = 32

Aufbau Principle

The Aufbau is a German word. It means “building up or construction”. This concept was
formulated by the scientists Niels Bohr and Wolfgang Pauli. According to the concept of the
sub-shell, each and every main shell contains one or more sub-shells, which are represented
by s, p, d and f. There is difference in the energy level of these sub-shells. During electronic
configuration, entering of electrons always occur from the lower energy level to the higher
energy level. Thus, according to Aufbau principle, “filling of electrons in the sub-shells always
occurs from the lower energy level to the higher energy level.”

In absence of magnetic field, the energy of an orbital depends upon the value of principal
quantum number (n) and azimuthal quantum number (l). Hence, the sequence of filling of
orbitals can be obtained from the (n+I) rule. It can be summarized as:

CHEMISTRY 0Optional Science - 10 165

i. The orbitals with lower value of (n+l) have lower energy and are filled first.

ii. When two orbitals have the same value of (n+l), the orbital with lower value of "n" has
lower energy and is filled first.

Memory Plus

For the sub-shells s, p, d and f, the value of azimuthal quantum number (l) is represented
as 0, 1, 2 and 3 respectively. For example: in 3d, the value of n = 3 and the value of "l" is 2.

The (n+l) value of the some orbitals is given in the table

Orbitals Principal quantum number (n) Azimuthal Sum of n + l
1s 1 quantum number (l ) 1+0=1

0

2s 2 0 2+0=2
2p 2 1 2+1=3

3s 3 0 3+0=3
3p 3 1 3+1=4

3d 3 2 3+2=5

4s 4 0 4+0=4

From the above table, we can conclude the following points. l=O 1=1 1=2 1=3
n=1
1. The first electron always goes to the 1s orbital as it has the n=2
lowest (n+l) value. n=3
n=4
2. In between 2s and 2p, the (n+l) value of 2s is less than 2p. So,
the 2s orbital is filled first. n=S
n=6
3. In between 2p and 3s, both have the same value of (n+l). In this n=7
case electron enters into 2p orbital which has less value of "n".
n=8
4. In between 3s and 3p, the (n+l) value of 3s is less than 3p. So, fig:Order of filling electrons in
the 3s orbital is filled first.
orbitals
5. In between 3d and 4s, the (n+l) value of 4s is less than 3d. So,
the 4s orbital is filled first.

On the basis of above discussion, the sequence of filling of electrons
in different sub-shells is as follows.

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

To find out the sequence of energy level of different sub-shells, an alternative arrow diagram
method is shown in the given figure above.

Limitations to Aufbau principle

The Aufbau principle is applicable in most of the cases. However, some elements do not follow
this rule. For example: copper and chromium do not follow Aufbau principle.

166 Optional Science - 10

CHEMISTRY

According to Aufbau Principle: Chromium (24) = 1s2, 2s2 2p6, 3s2 3p6, 4s2, 3d4
Actual electronic configuration: Chromium (24) = 1s2, 2s2 2p6, 3s2 3p6, 4s1, 3d5
According to Aufbau Principle: Copper (29) = 1s2, 2s2 2p6, 3s2 3p6, 4s2, 3d9
Actual electronic configuration: Copper (29) = 1s2, 2s2 2p6, 3s2 3p6, 4s1, 3d10

Electronic Configuration

The electronic configuration in different shells is done according to 2n2 formula. According
to this formula, in the first shell, we can fill maximum 2 electrons, in second shell, we can fill
8 electrons and in the third shell 18 electrons. This method is easy and clear for the first 18
elements. But, beyond atomic number 18, it does not express the true electronic configuration.
For example: on the basis of the 2n2 formula, the electronic configuration of iron (atomic
number 26) is 2,8,8,8. But, it is not the case in iron. The actual electronic configuration of iron
is 2,8,14,2. It can be easily understand by Aufbau principle.

The systematic distribution of electrons in different shells and sub-shells is called electronic
configuration. On the basis of Aufbau principle, the electronic configuration of the some
elements in different shells and the sub-shells is given in the table below.

At.

Elements KL M N OP

No.
No. 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 6f

H 11

He 2 2

Li 3 2 1

Be 4 2 2

B 5 221

C 6 222

N 7 223

O 8 224

F 9 225

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

Ba 56 2 2 6 2 6 10 2 6 10 2 6 2

U 92 2 2 6 2 6 10 2 6 10 14 2 6 10 3 2 6 1 2

CHEMISTRY 0Optional Science - 10 167

S.N. Element Atomic number Sub-shell electronic configuration
1 Hydrogen 1 1s1

2 Helium 2 1s2
3 Lithium 3 1s2, 2s1
4 Beryllium 4 1s2, 2s2
5 Boron 5 1s2, 2s2 2p1
6 Carbon 6 1s2, 2s2 2p2
7 Nitrogen 7 1s2, 2s2 2p3
8 Oxygen 8 1s2, 2s2 2p4
9 Fluorine 9 1s2, 2s2 2p5
10 Neon 10 1s2, 2s2 2p6
11 Sodium 11 1s2, 2s2 2p6, 3s1
12 Magnesium 12 1s2, 2s2 2p6, 3s2
13 Aluminium 13 1s2, 2s2 2p6, 3s2 3p1
14 Silicon 14 1s2, 2s2 2p6, 3s2 3p2
15 Phosphorus 15 1s2, 2s2 2p6, 3s2 3p3
16 Sulphur 16 1s2, 2s2 2p6, 3s2 3p4
17 Chlorine 17 1s2, 2s2 2p6, 3s2 3p5
18 Argon 18 1s2, 2s2 2p6, 3s2 3p6

19 Potassium 19 1s2, 2s2 2p6, 3s2 3p6, 4s1
20 Calcium 20 1s2, 2s2 2p6, 3s2 3p6, 4s2

Classification of Elements based on Electronic Configuration (s, p, d and f blocks)

According to the sub-shells electronic configuration, the periodic table is divided into four
parts. They are called blocks. The four different blocks are given below.

1. s- block 2. p- block 3. d- block 4. f- block

s - Block ………….. p - Block …………

Period IA 0

n=1 IIA IIIA IVA VA VIA VIIA

n=2 ……..………………… d - Block …….…………………
n=3
IIIB IVB VB VIB VIIB VIII IB IIB

n=4

n=5

n=6

n=7

0168 Optional Science - 10

CHEMISTRY

1=---=~[JJTT--+-+-*seriesofLanthanide TT~~
Elements (58 to 71)
**Series of Actinide Elements

(90 to 103)

…………………………………….. f – Block elements ……………………………………………

1. s-block

The block of alkali metals (IA) and alkaline earth metals (IIA) which is present at the left-
hand side of the periodic table is called s-block. The elements which have the last electron in
"s"sub-shell are called s-block elements. Examples: Li, Na, K, Rb, Cs, Be, Mg, Ca, etc.

Fact with reason

Sodium is a s-block element. Why?
The electronic configuration of sodium (1s2, 2s2 2p6, 3s1) shows that its last electron
enters into "s" sub-shell. So, sodium is a s-block element.

2. p-block
The block of metals, non-metals, metalloids and inert gases which are present in IIIA, IVA,
VA, VIA, VIIA and zero groups is called p-block. It is present at the right-hand side of
the modern periodic table. The elements which have the last electron in "p" sub-shell are
called p-block elements. For example: C, N, O, F, Cl, etc.

Fact with reason

Chlorine is a p-block element. Why?
The electronic configuration of chlorine (1s2, 2s2 2p6, 3s2 3p5) shows that its last electron
enters into "p" sub-shell. So, chlorine is a p-block element.

3. d-block

The block of transitional metals which is present in between s-block and p-block of the
modern periodic table is called d-block. This block includes elements of IB to VIIB and
VIIIth groups. As the d-block is present between s-block and p-block, it is also known as
the transitional block. The elements which have the last electron in d-sub shell are called
d-block elements. Examples: Ag, Au, Fe, Cu, etc.

Fact with reason

Iron is a d-block element. Why?

The electronic configuration of iron (1s2, 2s2 2p6, 3s2 3p6, 4s2, 3d6) shows that its last
electron enters into "d" sub-shell. So, iron is a d-block element.

4. f-block

The block of lanthanides and actinides which is present below the main periodic table
is called f-block. The elements which have last electron in "f" sub-shell are kept in this
block. It is also called the block of inner transitional elements as it is taken out from the
transitional block. In this block there are only 28 elements. Examples: Ce, Th, U, etc.

Optional Science - 10 169

CHEMISTRY

Valence shell

The outermost shell of an atom is called valence shell. For example, in sodium the third shell
is called valence shell and in potassium the fourth shell is called valence shell. Loss or gain or
share of electrons takes place from the valence shell.

Valence electrons

The total number of electrons present in the valence shell of an atom are called valence electrons.
For example, in sodium there is one valence electron, in oxygen, there are six valence electrons
and in chlorine there are seven valence electrons.

Valency

Except six inert gases (He, Ne, Ar, Kr, Xe and Rn), other elements combine with each other
to make compounds. During the chemical combination, these elements either give or take
or share the certain number of electrons from their outermost shell. After they give, take or
share of these electrons, elements attain stable electronic configuration. Thus, the valency is
the number of electrons gained, lost or shared by an element during the chemical reaction.
For example: the valency of sodium is one as it loses one electron from its outermost shell.
Similarly, the valency of oxygen is two as it gains two electrons in its outermost shell.

In modern periodic table, the valency remains the same in the group top to bottom. Similarly,
in the period left to right, valency first increases up to four and decrease up to zero.

Element Li Be B CN O F Ne
Valency 123 43 2 1 0

Some elements, their electronic configuration and valency is given in the table.

S.N. Element Electronic configuration Valency

1 Hydrogen 1 1 (can give or take 1)

2 Helium 2 0 (already duplet)
3 Boron 2, 3 3 (can give 3 electrons)
4 Magnesium 2, 8, 2 2 (can give 2 electrons)
5 Argon 2, 8, 8 0 (already octet)

6 Chlorine 2, 8, 7 1 (can take 1 electron)
7 Sulphur 2, 8, 6 2 (can take 2 electrons)
8 Potassium 2, 8, 8, 1 1 (can give 1 electron)
9 Sodium 2, 8, 1 1 (can give 1 electron)

10 Aluminium 2, 8, 3 3 (can give 3 electrons)
11 Calcium 2, 8, 8, 2 2 (can give 2 electrons)
12 Fluorine 2, 7 1 (can take 1 electron)
13 Neon 2, 8 0 (already octet)

Valency of the radical

Atoms or group of atoms which have common charge on them are called radicals. There

are two types of radicals. They are electropositive radicals and electronegative radicals. The

valency of the radicals is the amount of charge present on them. Some electropositive radicals

and electronegative radicals with their valencies are given below.

0170 Optional Science - 10

CHEMISTRY

S.N. IElectropositive radicals I Valency IElectronegative radicals Valency l
2
1 l Ammonium (NH4+) 11 Sulphate (SO4--) 1 -
2
2 Sodium (Na+) 1 Nitrate (NO3-) 1 -
3
~ Calcium (Ca++) 2 Carbonate (CO3--) 1 I
2
3 I Aluminium (Al+++) I3 Hydroxide (OH-) 1
1
4 Potassium (K+ ) 1 Phosphate (PO4---) 1
2 Chloride (Cl-)
5 Magnesium (Mg++)
I1 Oxide (O--)
6 IHydrogen (H+) I1 Cyanide (CN-)
ISilver (Ag+ )
7 IZinc (Zn++) 12 Bromide (Br-)

- 1Beryllium (Be++) l2 Fluoride (F-)

8 I

~

9

-

L10

Variable Valency

We know that iron makes two types of compounds with chlorine. They are FeCl2 and FeCl3.
Similarly copper makes CuCl and CuCl₂. Did you notice two types of valency in FeCl2 and
FeCl3? In FeCl2, the valency of iron is 2 and in FeCl3, the valency of iron is 3. It means iron
shows two valency 2 and 3. Similarly, copper shows two valency 1 and 2. Such type of valency

is called variable valency. Thus, occurrence of more than one valency of an element is called

variable valency. The existence of variable valency of the elements means that they can give,

take or share different number of electrons with different atoms or different radicals. The

transitional elements (d-block elements) in the modern periodic table show variable valency

because they have two incomplete outer shells.

The elements and their variable valency are given in the table.

S.N. Elements Radicals Valency
1 Iron Ferrous (Fe2+) 2
2 Copper
3 Tin Ferric (Fe3+) 3
4 Lead Cuprous (Cu+) 1
5 Mercury
6 Gold Cupric (Cu2+) 2
7 Antimony Stannous (Sn2+) 2
4
Stannic (Sn4+) 2
Plumbous (Pb2+) 4
1
Plumbic (Pb4+)
Mercurous (Hg+) 2
1
Mercuric (Hg2+)
Aurous (Au+) 3
3
Auric (Au3+) 5
Antimonous (Sb3+)

Antimonic (Sb5+)

Fact with reason

Why do d-block elements show variable valency?

The d-block elements have two incomplete outer shells. Electrons present in these two

shells participate in chemical bonding. So, d-block elements show variable valency.

CHEMISTRY 0Optional Science - 10 171

Answer writing skill

1. State Aufbau principle and write down its sequence.
Ans: According to Aufbau principle, “filling of electrons always occurs from the lower

energy level to the higher energy level.” The sequence is given below:

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

2. What are d-block elements? Why are they called transitional elements?

Ans: Those elements whose last electron enters into d-orbital are called d-block elements.
They are called transitional elements because they are present in between s-block and
p-block.

3. What is ionization potential? How does it change across the period and down the
group?

Ans: The amount of energy required to remove the loosely bonded outermost electron
from an isolated gaseous atom is called ionization potential. In the period left to right,
ionization potential increases as the atomic size decreases. Similarly, the value of
ionization potential decreases in the group top to bottom, as atomic size increases.

4. Write down a main difference between Mendeleev's Periodic table and Modern
Periodic table?

Ans: Mendelveev's periodic table is based on increasing atomic weight but Modern periodic
table is based on increasing atomic number.

5. Which has bigger atomic size between sodium and potassium? Why?

Ans: Potassium has bigger atomic size between sodium and potassium. This is because
potassium has four atomic shells and sodium has only three atomic shells. Due to more
number of shells, potassium has bigger atomic size.

6. Why are fluorine and chlorine kept in p-block?

Ans: The electronic configuration of fluorine is 1s2, 2s2 2p5 and the electronic configuration of
chlorine is 1s2, 2s2 2p6, 3s2 3p5. This electronic configuration shows that the last electron
of the fluorine and chlorine lie in the "p" sub-shell. So, they lie in p-block.

7. Write any two differences between s-block and p-block elements.
Ans:

S.N. s-block elements S.N. p-block elements

1 s-block elements have the last 1 p-block elements have the last electron in

electron in "s" sub-shell. "p" sub-shell.

2 All s-block elements are metal. 2 p-block elements are metals, non-metals,
metalloids and inert gases.

0172 Optional Science - 10

CHEMISTRY

8. Write down any three characteristics of the Modern Periodic Table.

Ans: The main characteristics of the modern periodic table are given below:

i. In modern periodic table, elements are arranged on the basis of their increasing
atomic numbers.

ii. In this table, there are seven periods and 18 groups.

iii. Inert gases like He, Ne, Ar, Kr, Xe and Rn are placed in a separate group (zero
group or 18th group) at the extreme right side of the table.

9. Write down any thee advantages of the periodic table.

Ans: The advantages of the periodic table are given below.

i. The detail information about the elements like the number of electrons, number of
protons, electronic configuration, atomic number, atomic weight, valency, nature,
block, etc. can be obtained from the periodic table.

ii. With the help of periodic table, the properties of elements can be studied easily.

iii. The periodic variations such as atomic number, valency, atomic radius, ionization
potential, electronegativity, electron affinity, etc. can be studied easily with the
help of periodic table.

10. How does modern periodic table remove the defects of the Mendeleev's Periodic table?

Ans: In the modern periodic table, elements are classified on the basis of increasing atomic
numbers. This system of classification helped to solve the defects of Mendeleev's
periodic table in the following ways.

1. The position of hydrogen is fixed: The accurate position of hydrogen has not yet
been solved completely. But as its atomic number is the least number (i.e. one), it
is kept in IA (group 1) group of the modern periodic table along with alkali metals.

2. The position of isotopes is fixed: Isotopes are the atoms of the same elements,
which have the same atomic numbers but different atomic weights. Therefore,
isotopes of one element are placed together in the same group. For example,
there are three isotopes of carbon (i.e. C612, C613 and C614) which have same atomic
numbers. They are placed together in IVA group of the modern periodic table.

3. The wrong position of the elements has been corrected: When elements are
arranged on the basis of their increasing atomic numbers, the wrong position of
argon, potassium, nickel, cobalt, etc. is solved automatically without changing
their own places.

4. The alkali metals and coinage metals are kept separately: In the modern periodic
table, the most reactive alkali metals are kept in IA group (group 1) whereas the
least reactive coinage metals are kept in IB group (group11).

5. Lanthanides and actinides are kept below the main table: In the modern periodic
table, fourteen elements of lanthanide series and fourteen elements of actinide
series are placed in a separate box below the main periodic table.

Optional Science - 10 173

CHEMISTRY

Exercise

1Step

1. Define the following terms:

i. Mendeleev's periodic table ii. Modern periodic table iii. Aufbau principle

iv. Electronic configuration v. Valency vi. Variable valency

vii. Ionization Potentaial viii. Electronegativity ix. Electron affinity

x. Alkali metals xi. Alkaline earth metals xii. s-block elements

xiii. p-block elements xiv. d- block elements xv. f-block elements

2. Write electronic configuration of given elements on the basis of sub-shells.

i. Na1123 ii. Cl1735 iii. K 39 iv. Ca2040
19

2Step

1. Write two differences between:
i. Modern periodic table and Mendeleev’s periodic table
ii. Periods and groups

iii. s- block elements and p- block elements

2. Give reason.
i. Elements of IA group of the modern periodic table are called alkali metals.
ii. Elements of IIA group of the modern periodic table are called alkaline earth metals.
iii. Fluorine is kept in p-block of the modern periodic table.
iv. Potassium is kept in the s-block of the periodic table.
v. Electron affinity increases in the period left to right and decreases top to bottom in
the group.

174 Optional Science - 10

CHEMISTRY

vi. Electronegativity increases in the period left to right and decreases in the group top to
bottom.

vii. The electronic configuration of chromium and copper does not follow the Aufbau
Principle.

viii. As we move from top to bottom, atomic size increases.
ix. As we move from top to bottom, the ionization potential decreases.

3Step

1. What is Mendeleev's Periodic table? Write down its characteristics.
2. What is Modern Periodic table? Write down its characteristics.
3. Explain the demerits of the Mendeleev's periodic table in the points.
4. Enlist the advantages of the Mendeleev's periodic table.
5. What are the advantages of the periodic table? Write down in points.
6. What are the characteristics that is why hydrogen can be placed in the IAgroup of the

modern periodic table?
7. Discuss the characteristics that is why hydrogen can be placed in the VIIA group of the

modern periodic table.
8. Discuss the drawbacks of the Mendeleev's Periodic table.
9. What is valency and variable valency? Generally d-block elements show variable

valency, why?
10. Write down two examples of elements and their electronic configuration where Aufbau

principle is not followed.

4Step

1. Describe the s, p, d and f blocks of the modern periodic table.
2. Explain the periodic variation of ionization potential, electronegativity, electron affinity

and atomic radius.

CHEMISTRY 0Optional Science - 10 175

Multiple choice questions:

1. Mendeleev's periodic table is based on increasing

a. Atomic number b. Atomic weight

c. Electron numbers d. Protons number

2. Mendeleev's periodic table has

a. 7 periods and 18 groups b. 7 periods and 8 groups

c. 8 periods and 7 groups d. 18 periods and 7 groups

3. The maximum number of electrons in s and p orbitals are

a. 2 and 5 b. 2 and 9

c. 3 and 6 d. 2 and 6

4. The valency of SO4 in H2SO4 is b. 2
a. 1

c. 3 d. 4

5. The symbol of ferrous and ferric are:

a. Fe3+ and Fe2+ b. Fe2+ and Fe

c. Fe2+ and Fe3+ d. Fe and Fe+

6. The unit of ionization potential is

a. Potential volt b. Electron volt

c. Pascal d. Joule

7. Which of the following shows variable valency?

a. Sodium b. Calcium

c. Iron d. Potassium

8. Ionization potential increases in

a. Group b. Period

c. Block d. None of the above

9. What is the value of azimuthal quantum number for d sub-shell?

a. 1 b. 2

c. 3 d. 4

10. What is the correct electronic configuration of chromium?

a. 1s2, 2s2 2p6, 3s2 3p6, 4s2, 3d4 b. 1s2, 2s2 2p6, 3s2 3p6, 4s1, 3d5

c. 1s2, 2s2 2p6, 3s2 3p6, 4s2, 3d9 d. 1s2, 2s2 2p6, 3s2 3p6, 4s1, 3d10

0176 Optional Science - 10

CHEMISTRY

UNIT CHEMICAL
BONDING AND
9
CHEMICAL
ARITHMETIC

About the inspiring personality

Avogadro born in 1776 and died in 1856. He is best known for his law that "equal volumes
of different gases contain an equal number of molecules, provided they are at the same
temperature and pressure". He worked as a professor in mathematical physics in Turin
University.

Syllabus issued by CDC Learning objectives:

Theory 6 At the end of this unit, students will be able to:

Practical 2 • Explain sigma bond, pi bond and coordinate
bond along with the chemical compounds.
• Types of bonding (Ionic, covalent • Explain Avogadro’s number and chemical
arithmetic.
and coordinate)

• Sigma and pi bonds

• Avogadro's law • Solve the simple numerical problems based on
• Chemical calculation Avogadro’s number and chemical arithmetic.

• Structure of electrovalent compounds
(NaCl, MgCl2 and CaO)

• Structure of covalent compounds
(H2, O2, N2, H2O, NH3 and CH4)

• Structure of coordinate compounds
(O3 and SO3)

Key terms and terminologies of the unit I

1. Bond: The force of attraction which binds two or more elements together to make a
stable chemical compound is called bond.

2. Octet: Octet state is the electronic configuration in which there are eight electrons in the
valence (outermost) shell of an atom.

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3. Octet rule: The tendency of an element to make eight electrons in the outermost shell
as a result of gain or lose or share of electrons is called octet rule.

4. Duplet: Duplet is the state of electronic configuration in which an atom has only one
shell (K-shell) with 2 electrons in it.

5. Electrovalent bond: The chemical bond which is formed in between two opposite
charges as a result of transfer of electrons is called electrovalent bond.

6. Covalent bond: The chemical bond which is formed by mutual sharing of electrons
in between two or more non-metal atoms is called covalent bond.

7. Single covalent bond: In single covalent bond one pair of electrons is shared between
two non-metal atoms.

8. Double covalent bond: In double covalent bond, two pairs of electrons are shared
between two non-metal atoms.

9. Triple covalent bond: In triple covalent bond, three pairs of electrons are shared
between two non-metal atoms.

10. Electrovalent compounds: The chemical compounds which are formed as a result of
electrovalent bonding are called electrovalent compounds or ionic compounds.

11. Covalent compounds: The chemical compounds which are formed as a result of
covalent bonding are called covalent compounds.

12. Coordinate bond: The type of chemical bond in which one of the combining atoms
contributes both of the shared electrons is called coordinate bond.

13. Coordinate compounds: Those chemical compounds which are formed as a result of
coordinate bonding are called coordinate compounds.

14. Sigma and pi bonds: The covalent bond which are formed by the overlapping of the
atomic orbitals is called sigma and pi bonds.

15. Sigma-bond: The type of chemical bond which is formed due to head to head
overlapping of the atomic orbitals is called sigma bond.

16. Pi- bond: The type of chemical bond which is formed due to side to side overlapping
of the atomic orbitals is called pi-bond.

17. Avogagro's Principle: According to Avogagro's Principle "Under the constant
temperature and pressure, equal volumes of all gases contain equal number of
molecules".

Introduction

Scientists discovered altogether 118 elements till today. Most of these elements are unstable.
Except six inert gases all other elements react with each other to make stable compounds. When
two or more combining elements come near together, a force of attraction occurs between the
atoms. This force of attraction is responsible to bind these atoms. This attractive force is called
chemical bond. Therefore, the force of attraction which binds two or more elements together
to make a stable chemical compound is called bond.

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Elements are unstable because they have incomplete outermost shell. For example, sodium
is unstable because it has one electron in its outermost shell. Similarly, chlorine is unstable
because it has seven electrons in its outermost shell. Thus, those elements which have eight
electrons in their outermost shell (except helium) are stable in nature. In order to make
chemical bonding, it is necessary for an atom to attain stable electronic configuration like inert
gages. This stable electronic configuration is called duplet state and octet state.

Fact with reason

Except inert gases other elements are chemically unstable. Why?

Except inert gases other elements do not have complete outermost shell. This is why
they want to combine with other elements to gain stability. Thus, except inert gases other
elements are chemically unstable.

Octet and Octet Rule

Out of 118 elements, five elements, viz. Neon (Ne), Argon (Ar), Krypton (Kr), Xenon (Xe) and
Radon (Rn) have eight electrons in their outermost shell. The electronic configuration of these
elements is called octet state. Thus, octet state is the electronic configuration in which there
are 8 electrons in the valence (outermost) shell of an atom. It is observed that if an atom has
8 electrons in its outermost shell, it becomes chemically stable. These stable elements neither
give or take nor share electrons. Except these five elements, other elements are chemically
unstable. These elements give or take or share electrons to make complete octet. This tendency
of an element is called octet rule. Therefore, the tendency of an element to make eight electrons
in the outer most shell as a result of gain or lose or share of electrons is called octet rule. For
example, sodium loses one electron and chlorine gains one electron to obtain octet.

Atomic structure of neon (Ne), Argon (Ar), Krypton (Kr), Xenon (Xe) and Radon (Rn)

II Inert gas II Atomic Electronic configuration 7
number
S.N. Symbol K L IM N O -
10 P-
2 Neon Ne 18 2 8I
3 Argon Ar 36 -
4 Krypton Kr 54 2 8 I8 I -
5 Xenon Xe 86 -
6 Radon Rn 2 8 I 18 I 8 -
8 _J
I I 2 8 I 18 I 18 8

2 I 8 I 18 I 32 I 18

Duplet and Duplet Rule

Helium is a single element which has only one shell (K-shell) in its atomic structure. In this
K-shell, there are two electrons. It has been observed that, chemically helium is stable. So,
the state of electronic configuration of helium is called duplet. Duplet is the state of electronic
configuration in which an atom has only one shell (K-shell) with 2 electrons in it. There are
few elements like hydrogen, lithium, beryllium and boron which can form duplet by giving,
taking or sharing of electrons. It is called duplet rule. Therefore, the tendency of an element
to make two electrons in the outermost K-shell as a result of gain or lose or share of electrons
is called duplet rule.

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Fact with reason

Why are inert gases chemically stable?

Inert gases like helium, neon, organ, krypton, xenon and radon are chemically stable
because they have eight electrons in their outermost shell except helium. These elements
do not gain, take and share electrons. So, inert gases are chemically stable.

Chemical bonding

As we have discussed that, except inert gases other elements are chemically unstable and
combine together to make stable chemical compounds. In the chemical compounds, atoms
of the elements are combined together by a force of attraction called chemical bond. Due to
chemical bond, it is possible to make stable chemical compounds. According to the nature,
different kinds of chemical bonds are present in the compounds. In this unit, we will discuss
three different types of chemical bonds. They are given below.

i. Ionic or Electrovalent Bond ii. Covalent Bond iii. Coordinate Bond

Ionic or electrovalent bond

Metals have one, two or three electrons in their valence shell. Similarly, non-metals have five
or six or seven electrons in their valence shell. According to octet and duplet rules metals lose
electrons and non-metals gain electrons to complete octet or duplet. As a result metals become
positively charged and non-metals become negatively charged. In between these opposite
charges there occurs an electrostatic force of attraction called electrovalent bond. Thus, the
chemical bond which is formed in between two opposite charges as a result of transfer of
electrons is called electrovalent bond.

Na2, 8, 1 →—e— Na+2, 8 NaCl

Cl2, 8, 7 →+ e— Cl–2, 8, 8

Fig: Process of formation of electrovalent bonding

Memory Plus

The chemical bond which is formed by giving and taking of electrons between metals and
non-metals is called electrovalent bond or ionic bond.

Fact with reason

Electrovalent bond is also called ionic bond. Why?

In electrovalent bond transfer of electrons takes place from metals to non-metals. As a result,
metal gains positive charge and non-metal gains negative charge. In between these two
opposite charges electrostatic force of attraction takes place. Due to this force of attraction
opposite ions combine together to give compound. If this compound is kept in water, it also
gives corresponding opposite ions. Thus, electrovalent bond is also called ionic bond.

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Electrovalent compounds or ionic compound
There are large numbers of chemical compounds like, NaCl, KCl, MgCl2, NH4Cl, CaCl2,
CaCO3, etc. which are formed as a result of electrovalent bonding are called electrovalent
compounds. They are also called ionic compounds as they give opposite ions (cations and
anions) after the process of ionization. Thus, the chemical compounds which are formed as a
result of electrovalent bonding are called electrovalent compounds or ionic compound.

Characteristics of electrovalent compounds

1. Electrovalent compounds are formed as a result of electrovalent bonding.

2. They have high melting and boiling point.

2. They are strong chemical compounds as there is strong electrostatic force of attraction.

3. They are soluble in water but insoluble in organic solvents like ether, benzene, carbon
tetra chloride, etc.

4. They do not have the fixed geometrical shape.

5. In fused or solution state electrovalent compounds conduct electricity.

Memory Plus

Electronegative ions are also called acid radicals or anions. Similarly electropositive ions
are also called basic radicals or cations.

Covalent bond

Electrovalent bond is limited only in between two opposite charges (generally in between
metals and non-metals). But there are many chemical compounds which are formed by
combining two or more non-metals. For example: CH4, CO2, NH3, O2, N2, Cl2, etc. These
compounds are formed by mutual sharing of one or more electrons between the combining
atoms. The bond present in these compounds is called covalent bond. Thus, the chemical bond
which is formed by mutual sharing of electrons in between two or more non-metal atoms is
called covalent bond.

In a covalent bond, the pairing electrons come from both atoms in equal number. It means that

each sharing atom contributes equal number of electrons to the bond. For example, in hydrogen

there is one electron in its valence shell and chlorine has seven electrons in its valence shell.

Hydrogen needs one electron to complete duplet and chlorine needs one electron to make

octet. While making hydrochloric acid (HCl), hydrogen makes duplet and chlorine makes

octet. Therefore, both elements provide one electron each to make single covalent bond. The

pair of electrons that are shared by hydrogen and

chlorine are called share-pair electron. Similarly, the

pair of electrons which do not use in bonding are H ~ lone pair of

called lone pairs of electrons. In the compound HCl, / electrons

hydrogen does not have lone pair of electrons but shared pair ofelectrons in which each atom
chlorine has three lone pair of electrons.
contributes one electron
fig:formation of chemical bond in HCl

Types of covalent bonds

While making covalent bond, non-metal atoms share one or more pair of electrons. On the

basis of number of share-pair of electrons, there are three types of covalent bonds. They are:

CHEMISTRY 0Optional Science - 10 181

i) Single covalent bond

In single covalent bond one pair of electrons H-H 0 H
H-CI -H
is shared between two non-metal atoms. For /\
HH I
example, the bond between two hydrogen
H
atoms in H2, the bond between hydrogen
and oxygen atoms in H2O, the bond between II, molecule 11, 0 m olecule CII4 m olecule
carbon and hydrogen atoms in CH4, etc. The
single covalent bond is represented by using examples of single covalent bond in different molecules

single line (–).

ii) Double covalent bond

In double covalent bond, two pairs of elec- CO2molecule Oxygenmolecule

trons are shared between two non-metal examples of double covalent bond in different molecules

atoms. For example: The covalent bond be-

tween two oxygen atoms in O2, the bond be-
tween carbon and oxygen atoms in CO2, the
covalent bond between two carbon atoms in

ethene (CH2= CH2), etc. The double covalent
bond is represented by using double lines (=)

between two atoms.

iii) Triple covalent bond

In triple covalent bond, three pairs of elec- N=N H-C=N
trons are shared between two non-metal at-
oms. For example: The bond between nitro- Nitrogen molecule HCN molecule (Hydrogen cyanide)

examples of triple covalent bond in different molecules

gen atoms in N2 (N≡N), the bond between carbon and nitrogen atoms in HCN (H-C≡N),
the bond between two carbon atoms of ethyne C2H2(CH≡CH),etc. The triple covalent
bond is represented by using triple lines (≡).

Covalent compounds

The chemical compounds which are formed as a result of covalent bonding are called covalent
compounds. For example: CH4, CO2, H2O, CCl4, NH3, O2, N2, Cl2, H2, etc. All the covalent
compounds are formed between the non-metal atoms. These compounds have low melting
and boiling point.

Characteristics of the covalent compounds

1. Covalent compounds are formed by covalent bonding.
2. Covalent compounds have low melting and boiling point as compared to the electrovalent

compounds.
2. They are weaker than electrovalent compounds.
3. They are insoluble in water but soluble in organic solvents.
4. They have fixed geometrical shape.
5. They do not conduct electricity in aqueous solution.

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Differences between electrovalent and covalent bonds

S.N. Electrovalent bond S.N. Covalent bond

1. Electrovalent bond is formed by the 1. Covalent bond is formed by the sharing

transfer of electrons from metal to of electrons between two non-metal

non-metal. atoms.

2. It is a stronger bond. 2. It is a weaker bond.

3. It is formed between metal and 3. It is formed between non-metal atoms.

non-metal atoms.

Differences between electrovalent and covalent compounds

S.N. Electrovalent compounds S.N. Covalent compounds

1. Electrovalent compounds are formed 1. Covalent compounds are formed by

by electrovalent bonding. covalent bonding.

2. Generally, they are harder, stronger, 2. Generally, they are gases, liquids or soft

brittle and crystalline solid. solids.

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

point. point.

4. Generally, they are soluble in water 4. Generally, they are soluble in organic

but insoluble in organic solvent. solvent but insoluble in water.

5. They are good conductor of electricity 5. They are bad conductor of electricity.

in molten and solution state.

Coordinate Bond

The type of chemical bond in which one of the combining atoms contributes both of the shared
electrons is called coordinate bond. A coordinate bond is similar to a covalent bond but in a
coordinate bond, only one atom contributes for the bonding. We represent a coordinate bond
by using an arrow (→) instead of a dash (–). In a coordinate bond, the atom which donates a
pair of electrons is called donor atom and the atom which receives the electron pair is called
recipient atom. The arrow is always faces from donor to the recipient atom.

Example 1: Formation of ammonium ion (NH4+)

In ammonia, the nitrogen atom contains a lone pair of electrons after completing its octet. A

hydrogen ion needs two electrons to complete its duplet. So, hydrogen ion accepts the lone

pair of electrons from nitrogen atom. Thus, a

coordinate bond is formed between ammonia

and hydrogen ion to make ammonium

ion(NH4+).

NH3 + H+ NH4+ formation of ammonium ion

Example 2: Formation of hydronium ion (H3O+)

In water molecule, there are two lone pair of electrons in oxygen atom. Similarly, a hydrogen

ion does not have electrons. So, it needs two electrons to complete duplet. Now, hydrogen

ion accepts a lone pair of electrons from the oxygen atom. Thus, a coordinate bond is formed

between water and hydrogen ion.

H2O + H+[H3O]+

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Memory Plus

Coordinate bonding is a combination of covalent and electrovalent bonding.

Sigma and Pi bonds

In the atomic structure, different orbitals have different shapes. The s-orbital is spherical

shaped, p-orbital is dumb bell shaped, d-orbital is double dumb bell shaped and the shape of

f-orbital is complex. When atoms combine with each other, these orbitals overlap in different

ways. After their overlapping, covalent bond is formed. The way of overlapping is different

in different compounds. The type

of bond formed between atoms

depends upon how their orbitals y
overlap. The overlapping of orbitals

can occur between s and s orbitals,

p and p orbitals, s and p orbitals

and so on. Thus, the covalent bond p d
formed by the overlapping of the
atomic orbitals is called sigma and fig:shapes of atomic orbitals

pi bonds.

1. Sigma Bond (σ) (a) ~ •~ ~ ffi
s s
The type of chemical bond which is formed (a-) Sigm a bond
due to head to head overlapping of the atomic ~(bJ ~
orbitals is called sigma bond. The sigma bond c.+a. X X
is represented by a Greek letter sigma (σ). It s z • z - ()3jo z
is stronger than a pi bond. There are three
possibilities of forming sigma bond. They are: P Sigma (a-) bond

a. Overlapping of two s-orbitals o ioc,i ~El z • z - ojEDjo z
b. Overlapping of two p-orbitals
c. Overlapping of one s and one p orbitals p ,, Pz Sigm a (a") bond

fig: formation of a sigma bond due to (a) s-s overlap
(b) s-p overlap (c) p-p overlap

2. Pi Bond (p-bond)

The type of chemical bond which is formed due to side to side

overlapping of the atomic orbitals is called pi bond. It is represented

by a Greek letter Pi (p). Pi bond is a weaker covalent bond compared to

the sigma bond. There is only one way by which a pi bond is formed fig:sidewise overlap of
i.e. side to side overlapping of the atomic orbitals. p-orbitals forπ bond

It has been observed that a compound having single covalent bond between two carbon
atoms (C-C) or between carbon and hydrogen atoms (C-H) is always sigma bond.
Similarly, in any double covalent bond, there is one sigma bond and one bond. In the
same way, in a triple covalent bond, there is one sigma bond and two pi bonds.

In organic chemistry, one carbon atom can combine with another carbon atom making single,
double or triple covalent bonds. In these covalent bonds, the total numbers of sigma and pi
bonds are given in the table:

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S.N. Type of Bond Total number of Total number of pi bond
sigma bond
1 C-C (sigma bond) 1 0
2 C=C (double bond) 1
3 C≡C (triple bond) 1 2
1

Examples of ionic compounds

Those chemical compounds which are formed as a result of electrovalent bonding are called
electrovalent compounds or ionic compounds. The formation and structure of different ionic
compounds like NaCl, MgCl2, CaO, etc. are described below.

a. Formation of sodium chloride (NaCl)

Sodium chloride (NaCl) is an example of

electrovalent compound. It is formed by sodium contributes
combining sodium metal and chlorine non-metal.
Sodium has electronic configuration 2, 8, 1 and electro11, leaving It
chlorine has electronic configuration 2, 8, 7.
According to octet rule, the above electronic v~lh aclosed sholl
configuration shows that, sodium has one extra .. .Na•
+ •• - + + ,. ....
Na
• Cl : Cl ,
. /\

forming ionic .----' ~---..
bond chlorine gains

electron, leaving II

electron and chlorine requires one more electron with aclosed shell

to attain stable electronic configuration. Thus, in fig:formation of NaCl

order to make sodium chloride, sodium loses one electron and makes sodium ion (Na+).

Similarly, chlorine gains one electron and makes chloride ion (Cl-). In between these opposite

ions there occurs a force of attraction called electrostatic force of attraction or electrovalent

bond. As a result, sodium and chloride combine together to give sodium chloride (NaCl)

molecule.

Na (2,8,1) + Cl (2,8,7)  Na+ (2,8) Cl- (2,8,8)

+

~ ..
~ II

Na Cl Na+ Ct"
Sodium atom Chlorine atom Chloride ion
Sodium ion
(a cation) (an anion)

Sodium chloride (NaCl)

Fact with reason

Sodium chloride is an electrovalent compound. Why?
Sodium chloride is formed by the transfer of electron from sodium to chlorine.
Transfer of electrons from metal to non-metal makes electrovalent bonding. So,
sodium chloride is an electrovalent compound.

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b. Formation of magnesium chloride (MgCl2)

Magnesium chloride is an example of electrovalent compound. It is formed by combining
magnesium metal and chlorine non-metal. Magnesium has electronic configuration 2,
8, 2 and chlorine has electronic configuration 2, 8, 7. According to octet rule, the above
electronic configuration shows that, magnesium has two extra electrons and chlorine
requires one more electron to attain stable octet state. Thus, in order to make magnesium
chloride, magnesium loses two electrons and makes magnesium ion (Mg++). Similarly,
two chlorine atoms gain one electron each to make chloride ions (2Cl-). In between one
magnesium ion and two chloride ions there occurs a force of attraction called electrostatic
force of attraction or electrovalent bond. As a result, magnesium chloride (MgCl2)
molecule is formed. Here, magnesium chloride is formed by the transfer of electrons from
magnesium to chlorine (electrovalent bonding), it is an electrovalent compound.

Mg (2,8,2) + 2Cl (2,8,7)  Mg++ (2,8) Cl- (2,8,8) Cl- (2,8,8)

Each magnesium atom loses

(W}~r,/~- '°two electrons to form a

magnesium ion. ~

I
Each chlorine atom gains an xx
booomo ,cbloriddoo.
X Cl !
@»:~! •[@)J[@r[@)l ~
Mg: + xx
'--A@); I I
~ xx - - -

Cl !

xx

Electrostatic attraction arises between
magnesium and chloride ions.
Magnesium chloride is formed .

fig:Formation of magnesium chloride (MgCl2)

Fact with reason

Magnesium chloride is an electrovalent compound. Why?
Magnesium chloride is formed by the transfer of electron from magnesium to chlorine.
The transfer of electrons from metal to non-metal makes electrovalent bonding. So,
magnesium chloride is an electrovalent compound.

c. Formation of calcium oxide (CaO)

Calcium oxide is an example of electrovalent compound. It is formed by combining
calcium metal and oxygen non-metal. Calcium has electronic configuration 2, 8, 8, 2 and
oxygen has electronic configuration 2, 6. According to octet rule, the above electronic
configuration shows that, calcium has two extra electrons and oxygen requires two more
electrons to attain stable octet state. Thus, in order to make calcium oxide, calcium loses
two electrons and makes calcium ion (Ca++) resembling to electronic configuration of
argon. Similarly, oxygen atom gains two electrons to make oxide ion (O--) resembling
to electronic configuration of neon. In between one calcium ion and one oxide ion there
occurs a force of attraction called electrostatic force of attraction or electrovalent bond. As
a result,calcium oxide (CaO) molecule is formed. Here, calcium oxide is formed by the
transfer of electrons from calcium to oxygen (electrovalent bonding), it is an electrovalent
compound.

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({i/~ @( : :§:o:Ca ,:.,---+--..__, ..
[ ] 2-
[~ p=8Oxygen 2,6 - - > Caz+ (or)

l~ .rp~20Calcium2,8,8,2 ' ~ J"' Calcium atom Oxygen atom

~ J """ 2, 8, 8, 2 2, 6 Calcium u,n Oxide ion

Calcium lon(2. 8, 8) 2, 8, 8 2, 8

Calcium oxide

Oxide lon(2, 8)
fig:formation of calcium oxide (CaO)

Fact with reason

Calcium oxide is an electrovalent compound. Why?
Calcium oxide is formed by the transfer of electrons from calcium to oxygen. The
transfer of electrons from metal to non-metal makes electrovalent bonding. So, calcium
oxide is an electrovalent compound.

Examples of covalent molecules
Those molecules which are formed as a result of covalent bonding are called covalent
molecules. For example: H2, O2, N2, Cl2, H2O, CH4 , NH3, etc. The formation and structure of
some covalent molecules are described below:

a. Formation of hydrogen molecule (H2)

Hydrogen has one electron in its k-shell. According to
duplet rule, it needs one more electron to attain stable
electronic configuration similar to helium atom. Therefore,
two hydrogen atoms share one electron each to make fig:formation of hydrogen molecule (H2)
hydrogen molecule. Here, hydrogen molecule is formed as a result of sharing of electrons
(covalent bonding), it is a covalent molecule.

Fact with reason

Hydrogen (H2) is a covalent molecule. Why?
Hydrogen molecule is formed as a result of sharing of electrons between two
hydrogen atoms. The sharing of electrons makes covalent bonding. So, hydrogen (H2)
is a covalent molecule as it has covalent bonding.

b. Formation of nitrogen molecule (N2)

Nitrogen is a non-metal. Its electronic configuration is 2, No=2.5 · O·. N.. - ~
5. The electronic configuration shows that nitrogen has
five electrons in its valence shell. According to octet rule, ~
it needs three more electrons to attain stable electronic
N+N - N2

fig:formation of nitrogen molecule (N2)

configuration similar to electronic configuration of neon. Therefore, two nitrogen atoms

share three electrons each to make nitrogen molecule. Here, nitrogen molecule is formed

by the sharing of electrons (covalent bonding), it is a covalent molecule.

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Fact with reason

Nitrogen (N2) is a covalent molecule. Why?
Nitrogen molecule is formed as a result of sharing of electrons between two
nitrogen atoms. The sharing of electrons makes covalent bonding. So, nitrogen (N2)
is a covalent molecule as it has covalent bonding.

c. Formation of oxygen molecule (O2)

Oxygen is a non-metal. Its electronic configuration
is 2, 6. This electronic configuration shows that

@ @ - @ @oxygen has six electrons in its valence shell.

According to octet rule, it needs two more electrons 0 +0 O= O

to attain stable electronic configuration similar to fig:formation of oxygen molecule (O2)

electronic configuration of neon. Therefore, two oxygen atoms share two electrons each

to make oxygen molecule. Here, oxygen molecule is formed by the sharing of electrons

(covalent bonding), it is a covalent molecule.

Fact with reason

Oxygen (O2) is a covalent molecule. Why?
Oxygen molecule is formed as a result of sharing of electrons between two oxygen
atoms. The sharing of electrons makes covalent bonding. So, oxygen (O2) is a covalent
molecule as it has covalent bonding.

d. Formation of water molecule (H2O)

In water, there is one oxygen @(!) - (-0f'@o)
atom and two hydrogen atoms.
The electronic configuration of w
oxygen is 2, 6. In this electronic

configuration, there are six ~ H- 0
electrons in the valence shell
of oxygen. According to octet 2H + 0 I

H

fig:formation of water molecule (H2O) fig:3D structure of water molecule

rule it needs two electrons to attain stable electronic configuration similar to electronic

configuration of neon. Similarly, in hydrogen, there is one electron in its valence shell (or

K-shell). According to duplet rule, hydrogen needs one electron to attain stable electronic

configuration similar to helium. Therefore, one oxygen atom shares two electrons with

two hydrogen atoms and forms water molecule. Here, water molecule is formed as a

result of sharing of electrons (covalent bonding), it is a covalent molecule.

Fact with reason

Water (H2O) is a covalent molecule. Why?
Water molecule is formed as a result of sharing of electrons between one oxygen atom
and two hydrogen atoms. The sharing of electrons makes covalent bonding. So, water
(H2O) is a covalent molecule as it has covalent bonding.

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e. Formation of methane molecule (CH4)

H

I
or 1-1-C- H

k

fig:formation of methane molecule (CH₄) fig:3D structure of methane molecule

In methane there is one carbon atom and four hydrogen atoms. The electronic configuration
of carbon is 2, 4. In this electronic configuration, there are four electrons in the valence
shell of carbon. According to octet rule it needs four electrons to attain stable electronic
configuration similar to electronic configuration of neon. Similarly in hydrogen, there is
one electron in its valence shell (or k-shell). According to duplet rule, hydrogen needs one
electron to attain stable electronic configuration similar to helium. Therefore, one carbon
atom shares four electrons with four hydrogen atoms and forms methane molecule. Here,
methane molecule is formed as a result of sharing of electrons (covalent bonding), it is a
covalent molecule.

Fact with reason

Methane (CH4) is a covalent molecule. Why?
Methane molecule is formed as a result of sharing of electrons between one carbon
atom and four hydrogen atoms. The sharing of electrons makes covalent bonding. So,
methane (CH4) is a covalent molecule as it has covalent bonding.

f. Formation of ammonia molecule (NH3)
In ammonia, there is one nitrogen atom and three hydrogen atoms. The electronic
configuration of nitrogen is 2, 5. In this electronic configuration, there are five electrons
in the valence shell of nitrogen. According to octet rule it needs three electrons to attain
stable electronic configuration similar to electronic configuration of neon. Similarly, in
hydrogen, there is one electron in its valence shell (or K-shell). According to duplet rule,
hydrogen needs one electron to attain stable electronic configuration similar to helium.
Therefore, one nitrogen atom shares three electrons with three hydrogen atoms and forms
ammonia molecule. Here, ammonia molecule is formed as a result of sharing of electrons
(covalent bonding), it is a covalent molecule.

•• Ammonia or H

•N• + 3[xH] ➔ ~ H-NI -H

II

t

lone pair

fig:formation of ammonia molecule (NH₃) fig:3D structure of ammonia molecule

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Fact with reason

Ammonia (NH3) is a covalent molecule. Why?
Ammonia molecule is formed as a result of sharing of electrons between one nitrogen
atom and three hydrogen atoms. The sharing of electrons makes covalent bonding.
So, ammonia (NH3) is a covalent molecule as it has covalent bonding.

Examples of coordinate compounds
Those chemical compounds which are formed as a result of coordinate bonding are called
coordinate compounds. For example: ozone molecule (O3), sulphur trioxide molecule (SO3),
carbon monoxide molecule (CO), etc.

a. Structure of Ozone molecule (O3)

Ozone is a molecule which contains three oxygen atoms. It is also called trioxygen

molecule. The electronic configuration of oxygen atom is 2,6. According to the electronic

configuration, there are six electrons in the valance shell of an oxygen atom. On the basis

of octet rule, each oxygen atom needs two electrons to complete its octet similar to the

electronic configuration of neon. Now, the two oxygen atoms combine together to make

oxygen molecule (O2).

In an oxygen molecule, each oxygen atom has two lone pair 2+2=4 bonded electrons
of electrons. When one more oxygen atom comes to join with
the oxygen molecule, one of the oxygen atoms from the •• /o•.. . _ two lone pairs of
oxygen molecule donates one lone pair of electrons to the
O• • - ••/ electrons

third oxygen atom to complete octet of the third atom. So, in ozone molecule, there are

two types of bonds. They are covalent bond and coordinate bond.

. ~o's .0.. ..o.. • •Resonance ....o../Co~0..·.

fig:Ozone resonance structures fig:3D structure of ozone molecule

The actual structure of ozone is angular with O-O-O bond angle of 116.80. But, for our
convenient, we consider the structure of ozone linear planar. Actually, the double bond
of the ozone molecule can be found between any two oxygen atoms with the central
atom. Therefore, the structure of an ozone molecule is a resonance structure in which two
possible structures are present.

Fact with reason

Ozone is a coordinate covalent compound. Why?

Ozone molecule has three oxygen atoms. Among these atoms, there is one covalent
double bond between any two oxygen atoms and one coordinate bond between one
of the covalent bonding oxygen atoms and the third oxygen atom. So, ozone is a
coordinate covalent compound.

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Memory Plus

Abundant amount of ozone molecules are present in the stratosphere making ozone layer.

b. Structure of sulphur trioxide molecule (SO3)
Sulphur trioxide is a molecule which contains one sulphur and three oxygen atoms. The
electronic configuration of sulphur atom is 2, 8, 6 and oxygen atom is 2, 6. According to the
above electronic configuration, there are six electrons in the valance shell of sulphur and
oxygen atoms. On the basis of octet rule, sulphur and oxygen atoms need two electrons
each to complete their octet. Now, one sulphur atom and one oxygen atom combine
together to make double covalent bond. In this structure, sulphur and oxygen both have
two lone pair of electrons in each atom. Later on, sulphur atom donates its both lone pair
of electrons to the two oxygen atoms making two coordinate bonding between two S-O
atoms. As a result, sulphur trioxide molecule is formed. In sulphur trioxide molecule,
there are two types of bonds. They are covalent bond and coordinate bond.

00 0
11 < , I < ) II
/ , ~\ / s\
00 OO O0

fig:sulphur trioxide resonance structures fig:3D structure of sulphur trioxide molecule

The structure of sulphur trioxide is trigonal planar. The bond angle between O-S-O in

SO3 is 120 degrees. Since, the sulphur can form double bond between sulphur and oxygen
atoms, the SO3 have three possible structures called resonance structures.

Fact with reason

Sulphur trioxide is a coordinate covalent compound. Why?

In sulphur trioxide molecule there is one sulphur atom and three oxygen atoms. Among
these atoms, there is one covalent double bond between sulphur atom and one of the
oxygen atoms. Similarly, two coordinate bonds are present between one sulphur and two
oxygen atoms. So, sulphur trioxide is a coordinate covalent compound.

Avogadro’s Law

We have discussed different gas laws like Boyle's law, Charle's law, Graham's law, etc. in
class nine. They established the relation of temperature, pressure and volume of the gases.
The Boyle’s law gives the relationship between pressure and volume of a gas, Charles law
gives the relationship between temperature and volume of a gas and Graham law gives
the relationship between rate of diffusion and density. Similarly, Avogadro’s Law gives the
relationship between volume of a gas with its amount (moles). Thus, Avogadro’s law states
that, “under constant temperature and pressure, equal volumes of all gases contain equal
number of molecules.”

Memory Plus

All ideal gases follow Avogadro's Principle in all conditions of temperature and pressure
but real gases follow this principle only at high temperature and low pressure.

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Memory Plus

• Under constant temperature and pressure, volume of a gas (V) is directly proportional
to the number of molecules (n). It means that V α n

• Under constant temperature and pressure, molar mass of all gases contain equal
number of molecules i.e. 6.022×1023 and equal volume i.e. 22.4 litres.

Mathematical derivation of Avogadro’s Law V1 V2
n1
Take a balloon and fill a certain amount of gas (say n1 mol) into the
balloon. Suppose the volume of gas in the balloon is V1. Now, what n2
happens to the volume of air inside the balloon if more air is blown
into it? Yes, the volume of air increases inside the balloon. Similarly, V2 V1
if we release the neck of the balloon gently to let air release out of the n1
balloon, the volume of air decreases. Thus, if we blow more air into the
balloon, there will be more volume of the gas and vice versa. It shows n2
that, the volume of a gas (V) is directly proportional to the amount of
gas (n).

i.e. V α n

or, V = nk where 'k' is a proportionality constant.

i.e. V = k
n

Let, us consider a sample of gas has volume V1 and amount of gas n1. If the amount of gas is
increased to make n2, then the volume of the gas also becomes V2. Then, we know that,

V1 = k ---------- (i)
n1

V2 = k ---------- (ii)
n2

Equating equations (i) and (ii),

V1 = V2
n1 n2

Explanation of Avogadro’s Law

Take equal volume of three gases like hydrogen (H2),
nitrogen (N2) and oxygen (O2) at equal temperature
and pressure. If we take molar mass of these gases,

then hydrogen contains 2g, nitrogen contains 28 g

and oxygen contains 32 g. At STP the molar mass 1mol 1mol 1mol
of these gases contain equal volume i.e. 22.4 litres.

The molar mass of these gases also contain equal

number of molecules i.e. 6.022 × 1023.

Using the above relationship between volume and 22.4 litres
amount of substance, we have
0

6,022 x 1023 molecules

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V1 = V2 = V3
n1 n2 n3

i.e., V1 = V2 = V3
mol mol 1 mol
1 1

or, V1 = V2 = V3 = 22.4 litre

From the above explanation, we can conclude that:

i. At STP 1 mole of Hydrogen gas = 2 g = 22.4 litres= 6.022 × 1023molecules.

ii. At STP 1 mole of Oxygen gas = 32 g = 22.4 litres = 6.022 ×1023molecules.

iii. At STP 1 mole of Nitrogen gas = 28 g = 22.4 litres = 6.022 × 1023molecules.

iv. At STP 1 mole of Carbon dioxide gas = 44 g = 22.4 litres = 6.022 × 1023molecules and so on.

Memory Plus

The volume occupied by one mole of an ideal gas is called molar volume. It is 22.4 litres.

Example: 1

A sample of 1.3 mole of gas has volume 10 litres. Calculate the amount of gas present in 20
litres if pressure and volume are kept constant.

Solution: Given,

Initial volume of the gas (V1)= 10 litres
Initial amount of gas (n1) = 1.3 mole
Final volume of the gas (V2) = 20 litres
Final amount of the gas (n2) = ?
According to Avogadro’s Law

V1 = V2
n1 n2

Or, V1n2 = V2n1

Or, 10 × n2 = 20 × 1.3

Or, n2 = 20× 1.3 mole
10

Or, n2 = 2.6 mole

 The final volume of the gas is 2.6 moles.

Example: 2

How many molecules are present in 10 litres of carbon dioxide at STP?
Solution: Given,

Volume of carbon dioxide at STP (V)= 10 litres
Number of molecules = ?
According to Avogadro’s Law

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22.4 litres of carbon dioxide at STP = 6.022 × 1023molecules

or, 1 litres of carbon dioxide at STP = 6.022× 1023 molecules
22.4

or, 10 litres of carbon dioxide at STP = 6.022× 1023 × 10 molecules
22.4

 10 litres of carbon dioxide at STP = 2.69 × 1023molecules

Example: 3

Calculate the volume of 5 mole of oxygen gas at STP.
Solution: Given,

Amount of oxygen gas (n) = 5 mole
Volume of oxygen (V) = ?
According to Avogadro’s Law,
1 mole of oxygen gas occupies = 22.4 litres volume
5 moles of oxygen gas occupies = 5× 22.4 litres = 112 litres

 The volume of 5 mol of oxygen gas is 112 L.

Example: 4

How much gram of oxygen gas is present in 90 litres of oxygen at STP?

Solution: Given,

Volume of oxygen gas (V) = 90 litres

Amount of oxygen = ?

According to Avogadro’s Law,

1 mole of oxygen gas= 22.4 litres = 32 grams of oxygen

Form the above relation,

or, 22.4 litres = 32 grams of oxygen

or, 1 litre = 32 grams of oxygen
22.4

or, 90 litres = 32 × 90 grams of oxygen
22.4

= 128.57grams of oxygen

 The 90 litres of oxygen contains 128.57 grams of oxygen

Answer writing skill

1. What is chemical bonding? Write down the three types of chemical bonds.

Ans: The force of attraction which binds two or more elements together to make a stable
chemical compound is called chemical bonding. Mainly there are three types of chemical
bonding. They are:

i. Ionic or Electrovalent Bond ii. Covalent Bond iii. Coordinate Bond

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2. Define electrovalent bond, covalent bond and coordinate bond.

Ans: The chemical bond which is formed in between metal and non-metal as a result of
transfer of electrons is called electrovalent bond. The chemical bond which is formed
by mutual sharing of electrons in between two or more non-metal atoms is called
covalent bond. Similarly, the type of chemical bond in which one of the combining
atoms contributes both of the shared electrons is called coordinate bond.

3. What are electrovalent compounds, covalent compounds and coordinate compounds?
Write down any two examples of each.

Ans: The chemical compounds which are formed as a result of electrovalent bonding are
called electrovalent compounds or ionic compound. Examples: NaCl, MgCl2

The chemical compounds which are formed as a result of covalent bonding are called
covalent compounds. Examples: H2O, NH3

Those chemical compounds which are formed as a result of coordinate bonding are
called coordinate compounds. Examples: O3, SO3

4. State Avogadro’s Law. Write down the relation among molar mass, number of
molecules and volume of the gas.

Ans: According to Avogadro’s law, under constant temperature and pressure, equal volumes
of all gases contain equal number of molecules. In a molar mass of a gas, the number of
molecules is 6.022 1023and volume is 22.4 litres.

5. SO3 and O3 are called coordinate covalent compounds. Why?

Ans: The molecules SO3 and O3 contain both covalent and coordinate bonds. So, they are
called coordinate covalent compounds.

6. Write any two differences between sigma bond and pi bond.

S.N. Sigma bond S.N. Pi bond

1. Sigma bond is formed by end to end 1. Pi bond is formed by side to side

overlapping of the atomic orbitals. overlapping of the atomic orbitals.

2. It is formed between s and s orbitals, 2. It is formed between p and p orbitals.
p and p orbitals and s and p orbitals.

7. Write any two differences between NaCl and CH4.

S.N. NaCl S.N. CH4

I I I I1. NaCl is an electrovalent compound. 1. CH4 is a covalent compound.

2. It has high melting and boiling point. 2. It has low melting and boiling point.

8. Write down any three characteristics of each electrovalent and covalent compounds.

(a) Characteristics of electrovalent compounds

i. Electrovalent compounds are formed as a result of electrovalent bonding.
ii. They have high melting and boiling point.
iii. They are strong chemical compounds as there is strong electrostatic force of

attraction.

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(b) Characteristics of covalent compounds

i. Covalent compounds are formed by covalent bonding.
ii. They have low melting and boiling point as compared to the electrovalent

compounds.
iii. They are weaker than electrovalent compounds.

9. In a balloon, 2.5 litres of helium gas is filled at STP. Calculate its mass?

Solution: Given,

Volume of helium = 2.5 litres

Mass of helium = ?

We know that,

22.4 litres of helium = 2 g helium

1 litre of helium = 2 grams of helium
22.4

2.5 litres of helium = 2 × 2.5 grams of helium
22.4

= 0.22 grams of helium

 The 2.5 litres of helium contains 0.22 grams of helium

10. Describe the formation and structure of sodium chloride (NaCl) molecule.

Ans: Sodium chloride (NaCl) is an example of electrovalent compound. It is formed by
combining sodium metal and chlorine non-metal. Sodium has electronic configuration
2, 8, 1 and chlorine has electronic configuration 2, 8, 7. According to octet rule, the
above electronic configuration shows that, sodium has one extra electron and chlorine
requires one more electron to attain stable electronic configuration. Thus, in order
to make sodium chloride, sodium loses one electron and makes sodium ion (Na+).
Similarly, chlorine gains one electron and makes chloride ion (Cl-). In between these
opposite ions there occurs a force of attraction called electrostatic force of attraction or
electrovalent bond. As a result, sodium and chloride combine together to give sodium
chloride (NaCl) molecule.

Na2, 8, 1 →—e— Na+2, 8

NaCl

- - - _ _ _ _ /Cl2, 8, 7 →+ e— Cl–2, 8, 8

~ ~7 ~ C® G
Cl Na+ c1-
Na Chlorine atom
Sodium atom Sodium Ion Chloride Ion
Sodium chlor ide (NaCl)
Na· =Cl=
Na+ =Cl~
1s22s22p63s; ...1_s_2_2_s_2_2p_e_3_523':i
' 1s22s22p6 1s22s22p6 3s23p6

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Exercise

1Step

1. Define the following terms.

i. Bond ii. Octet iii. Octet rule

iv. Duplet v. Electrovalent bond vi. Covalent bond

vii. Coordinate bond viii. Avogagro's Principle ix. Duplet rule

2. Write any four examples of the chemical compounds which have electrovalent bond.

3. Mention any four examples of the chemical compounds which have covalent bond.

4. Write down any two examples of the chemical compounds which have coordinate bond.

5. Separate the following into electrovalent, covalent and coordinate compounds.

i. NaCl ii. MgCl2 iii. CaO
iv. NH3 v. CH4 vi. O3
vii. SO3 viii. H2 ix. N2

2Step

1. Differentiate between:
i. Electrovalent bond and covalent bond
ii. Electrovalent compounds and covalent compounds
iii. Sigma bond and pi bond
iv. Structure of NaCl and CH4

2. Give reason.
i. Except inert gases other elements are chemically unstable.
ii. Inert gases are chemically stable.
iii. Electrovalent bond is also called ionic bond.
iv. Sodium chloride is an electrovalent compound.
v. Magnesium chloride is an electrovalent compound.
vi. Calcium oxide is an electrovalent compound.

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vii. Water (H2O) is a covalent molecule.
viii. Methane (CH4) is a covalent molecule.
ix. Ammonia (NH3) is a covalent molecule.
x. Sulphur trioxide is a coordinate covalent compound.
3. What are electrovalent compounds? Write any four examples.
4. Define covalent compounds and write down any two examples.
5. What are coordinate compounds? Write any two examples.

3Step

1. What are covalent compounds? Write down any four characteristics of the covalent
compounds.

2. Define electrovalent compounds and write down any four characteristics of the
lectrovalent compounds.

3. What are Sigma and Pi bonds? How are they formed?
4. Explain in short about the formation of hydrogen molecule (H2)
5. How is oxygen molecule (O2) formed? Describe in short.
6. Describe in short about the formation of nitrogen molecule (N2).
7. A sample of 2.6 mole of gas has volume 20 litres. Calculate the amount of gas present in

40 litres if pressure and volume are kept constant.
8. How many molecules are present in 5 litres of carbon dioxide at STP?
9. Calculate the volume of 10 mole of oxygen gas at STP.
10. How much gram of oxygen gas is present in 30 litres of oxygen at STP?
11. In a balloon, 4.0 litres of helium gas is filled at STP. Calculate its mass?

4Step

1. Describe the formation of the following molecules with the help of diagrams.

i. NaCl ii. MgCl2 iii. CaO

iv. NH3 v. CH4 vi. O3

2. vii. SO3 V1 = V2
State Avogagro's Principle and derive the relation
n1 n2

3. Describe an activity to prove that volume of a gas is directly proportional of its mass.

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Multiple Choice Questions

1. Electrovalent bond is formed between

a. Metals and metals b. Metals and non-metals
c. Non-metal and non-metal d. None of the above

2. Covalent bond is formed between

a. Metals and metals b. Metals and non-metals
c. Non-metal and non-metal d. None of the above

3. What type of bond is present between H and H in H2 ?

a. Covalent bond b. Ionic bond

c. Coordinate covalent bond d. Both a and b

4. What type of bond is present between Ca and O in CaO?

a. Covalent bond b. Ionic bond

c. Coordinate bond d. None

5. What type of chemical compound is MgCl2 ?

a. Covalent compound b. Ionic compound

c. Coordinate compound d. None of the above

6. Sigma bond is formed by the overlapping between:

a. s and s orbitals b. s and p orbitals

c. p and p orbitals d. All of the above

7. Pi bond is formed by the overlapping between

a. s and s orbitals b. s and p orbitals

c. p and p orbitals d. All of the above

8. How many Pi bonds are present in HC ≡ CH?

a. 1 b. 2

c. 3 d. 4

9. How many molecules are there in 2 g of H2 gas?

a. 6.022 ×1027molecules b. 6.022 ×1023 molecules

c. 6.022 ×1025molecules d. 6.022 ×1024 molecules

10 How much volume is present in 32g of O2 gas?

a. 22.2 litres b. 22.3 litres

c. 22.4 litres d. 22.5 litres

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UNIT ELECTROCHEMISTRY

10

About the inspiring personality

Svante Arrhenius was born on February 19, 1859, in Vik, Sweden and died on October 2,
1927. Arrhenius propounded many theories in chemistry, astronomy, earth’s ecology, etc.
He described Arrhenius theory of ionization.

...........................................................................................................................................................................................................................:

Syllabus issued by CDC Learning objectives:

• Ionic product of water After completing the study of this unit, students

• Introduction to pH and pOH along will be able to:

with their numerical problems. • Explain the ionic product of water.

• pH meter and pH range • Describe pH and pOH along with the

• Application of the neutralization numerical problems.

reaction (soil test, treatment of • Describe pH meter and pH range.

acidity, treatment of insect bite) • Explain the application of the neutralization

reactions in soil test, use of antacid, biting of

insects, etc.

Key terms and terminologies of the unit

1. Electrochemistry: Electrochemistry is a branch of physical chemistry which deals
with relationship between chemical energy and electrical energy and how one is
converted into another.

2. Ionic product of water (Kw): The product of molar concentration of hydrogen ion
(H+) and hydroxide ion (OH–) produced by self-ionization of water at a particular

temperature is called ionic product of water.

3. pH: The negative logarithm of molar concentration of hydrogen ions is called pH.

4. pH scale: The scale of hydrogen ions concentration which is used to express acidic
and basic strength of an aqueous solution is called pH scale.

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