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SUB: CHEM-101
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 Definition:
 Electron: An electron is a mass less subatomic particle which bears one unit of negative charge
on it. It is denoted by
0
1
e
 Proton: It can be defined as a subatomic particle which bears mass equal to 1 amu and positive
charge equal to +1 on it.
 Neutron: A subatomic particle which bears a mass equal to 1 amu and has no charge on it. It is
denoted by
1
0
n
 Atomic number: Atomic number of an element is defined as the number of protons present in
the nucleus.
Atomic No. Z = No. of protons.
= No of electron in an atom.
So, Z = P = e
 Mass number: The sum of the number of neutrons and protons present in the nucleus of an
atom is called the mass number of that atom.

1.
2.
3.
Mass No. = No. of protons + No. of neutrons
So, A = P+n
Rutherford ‘s Nuclear Atomic Model: According to this model,
Atom contains a massive (Heavy) and positively charged part at its centre. This central part of
the atom is called nucleus.
The size of the nucleus is very small as compared to that of the whole atom. Atom consists of a
lot of empty space round the nucleus.
The electrons are revolving round the nucleus in orbits with a fast speed as the planets revolved
around the sun.
 Drawbacks of the classical Rutherford ‘s Nuclear Atomic Model:
1. According to the classical electromagnetic theory, if a charged particle (i.e. electron) is
accelerating around an oppositely charged particle (i.e. nucleus), it (i.e. electron) would
continuously lose on radiate energy and its speed will decrease and hence will gradually come
closer to the nucleus by a spiral path. The ultimate result of this spiral path will be that the
electron will fall into the nucleus, thereby making the atom unstable i.e, according to
Rutherford ‘s model the atom should not exist.
2. This model could not say anything as to how and where these electrons were arranged.
3. This model could also not explain how the spectral line produced.
 Postulates of Bohr’s Theory:
1. Each orbit round the nucleus is associated with a definite amount of energy and the orbits are
therefore called energy level or main energy shells. These shells are numbered as 1, 2, 3…….
Starting from the nucleus and are designated by capital letters K, L, M……. respectively.
2. While revolving round the nucleus in a fixed orbit the electron neither loses nor gains energy.
3. Energy is emitted or absorbed by an atom when an electron jumps from one energy level to the
other.
4. The electron can move only in that orbit in which the angular momentum of the electron is
quantized, i.e. the angular momentum of the electron is a whole number multiple of
mvr =
h
i.e.
2
nh
2
Where, m = mass of the electron
v = tangential velocity of electron in its orbit
r = distance between the electron and the nucleus
n = a whole number 1, 2, 3…….
 Limitation of Bohr’s Theory:
1. No explanation of fine structure.
2. No explanation for the spectra of multi electron system.
3. He could not give any explanation for using the principle of quantization of angular
momentum and it was introduced by him arbitrary.
4. According to Bohr’s the electrons move in circular orbits round the nucleus but the modern
researches have shown that the motion of electron is not limited to a single plane, but take
place in three dimensional space.
5. No explanation for zee man and stark effect.
6. Bohr assumes that an electron in an atom is located at a definite distance from the nucleus and
is revolving round it with definite velocity, i.e. it is associated with a fixed value of
momentum. This is against the Heisenberg’s uncertainty principle according to which is
impossible to determine simultaneously with certainly the position and momentum of a particle.
Q: What do you understand by the term quantum number? How many quantum numbers has an
electron in an orbital? Name and explain them in brief.
Ans: The quantum numbers are the identification numbers for an individual electron in an atom,
since these numbers fully describe the position and energy of an electron in an atom.
Four quantum numbers which are required to completely specify the character of an electron
are,
1. Principal quantum number (n)
2. Subsidiary quantum number (l)
3. Magnetic quantum number (m)
4. Spin quantum number (s)
1. Principal Quantum Number (n): This quantum number denotes the energy level or the principal
shell to which an electron belongs. It is denoted by “n”.
It can have only integral value from 1 to  , but 1 to 7 has
so far been established. The letters K, L, M…… are also used to designate the values of n.
2. Subsidiary Quantum Number (l): This quantum number is denoted by letter “l”. It is also called
angular momentum quantum number, orbital quantum number and sometimes subsidiary
quantum number. It gives an idea of the shapes of the orbital.
“l” can have any value from 0 to (n-1) for a given value of n i.e.
l = 0, 1, 2 …… (n-1), (n-2)
The total number of different values of l is equal to n.
Different values of l = 0, 1, 2, 3 …… are symbolized by the letter’s s, p, d, f.
3. Magnetic Quantum Number (m): It is denoted by “m”. This quantum number describes the
orientation of the orbital in space.
Total possible values of “m” depend on the values of “l”. “m”
can have the following values for a given value of “l”.
m = 0,  1,  2,  3……  l giving a total number of
values of m equal (2l+1). Thus if l = 0 (s-sub shell), m = 0 (one value only), l = 1 (p-sub shell), m
= 0,  1 (total three values), l = 2 (d-sub shell), m = 0,  1,  2 (total five values), l = 3 (f-sub
shell), m = 0,  1,  2,  3 (total seven values).
4. Spin Quantum Number (s): This quantum number describes the electron while moving round
the nucleus in an orbit rotates about its own axis either in a clock wise direction or in an anti
clock wise direction.
“s” can have two values viz +
spinning of the electron in the clock wise direction) and -
1
(corresponding to the
2
1
(corresponding to the spinning in
2
anti-clock wise direction).
 Paulis Exclusion Principle: This principle states that “It is impossible that two electron in a
given atom to have all the four quantum numbers identical”.
i.e. In an atom two electrons can have maximum three
quantum numbers (n, l, m) the same value and the fourth (s) will definitely be having a
different value.
 Mendeleef’s Periodic Law:
Russian chemist Mendeleef gave a law known as Mendeleef’s periodic law which states as:
The physical and chemical properties of the elements are a periodic function of their atomic
weights, i.e. if the elements are arranged in the increasing order of their atomic weight, the
properties of the elements (i.e. similar elements) are repeated after definite regular intervals or
period.
 Defects of Mendeleef’s Periodic Table:
1.
2.
3.
4.
5.
Mendeleef’s periodic table suffers from the following defects:
Hydrogen resembles both the alkali metals and the halogens. Its position in the periodic table is
therefore anomalous.
Lanthanides and Actinides do not find their proper place in the table.
Certain elements which possess similar properties are separated in the table as for example copper
mercury while many dissimilar elements have been grouped together. For example Cu, Ag and Au
grouped along with the alkali metals.
Elements of higher weight precede those of lower atomic weight at four places as shown below:
a) Ar ( Z = 18, At wt = 40 ) precedes K ( Z = 19, At wt = 39.0 )
b) Co ( Z = 27, At wt = 59.9 ) precedes Ni ( Z = 28, At wt = 58.6 )
c) Te ( Z = 52, At wt = 127.6 ) precedes I ( Z = 53, At wt = 126.9)
d) Th ( Z = 90, At wt = 232.12 ) precedes Pa ( Z = 91, At wt = 231 )
It is not possible to accommodate large number of isotopes in the periodic table.
 Modern Periodic Law:
Mosley put forward modern periodic law which states as follows:
The physical and chemical properties of elements are a periodic function of their atomic number,
i.e. if the elements are arranged in the increasing order of their atomic numbers, the properties of
the elements (i.e. similar elements) are repeated after definite regular intervals or periods.
 Extended or Long Form of Periodic Table:
In order to remove the defects of Mendeleef’s periodic table a number of tables have been
suggested far the classification of elements. All these tables classify the elements on the basis of
modern periodic law. i.e. the elements are arranged in the increasing order of their atomic
numbers. Out of such various tables long form of periodic table is most widely used. It is also
referred to as Bohr’s table, since it apparently follows the Bohr scheme of the arrangement of
elements into four types based on their electronic configurations.
 Different Portions of the Long Form of Periodic Table:
The long form of the periodic table has the following portions:
1. The Left Portions: This portion has the elements of groups IA and IIA. These elements are
extremely electropositive in character.
2. The Right Portion: This portion consists of the elements of groups IIIA, IVA, VA, VIA, VIIA
and zero. This portion has metal, all metalloids, non-metals and noble gases.
3. The Middle Portion: This portion consists of the elements of groups IIIB, IVB, VB, VIB, VIIB,
VIII, IB and IIB. Elements present in this portion can be classified into two groups,
(i)
Transition Elements
(ii)
Inner Transition Element
 Study of Periods and Groups:
1.
2.
3.
4.
Periods: These are the horizontal rows. Long form of periodic table consists of 7 periods.
Groups: The vertical columns are called groups. These are 1b in all as shown below:
IA, IIA, IIIA, IVA, VA, VIA and VIIA groups
IB, IIB, IIIB, IVB, VB, VIB and VIIB groups
VIII group. This group has three columns
Zero group. This group has inert element.
 Merits of Long Form of Periodic Table over Mendeleef’s Periodic Table:
1.
2.
3.
4.
5.
6.
7.
The long form of periodic table has a number of merits over the Mendeleef’s periodic table in the
following respects:
The classification of the elements is based on a more fundamental property viz atomic number.
It relates the position of an element to its electron configuration. Thus each group contains
elements with similar electronic configuration and hence similar properties.
It explains the similarities and variations in the properties of the elements in terms of their
electronic configurations and brings out clearly the trends in chemical properties across the long
period.
The inert gases having completely filled electron shells have been placed at the end of each period.
Such a location of the inert gases represents a logical completion of each period.
In this form of the periodic table, the elements of the two sub groups have been placed separately
and thus dissimilar elements do not fall together.
It provides a clear demarcation of different types of the elements like active metals, transition
metals, non-metals, metalloids, inert gases, lanthanides and actinides.
It is easier to remember, understand and reproduce.
 Defects of Long Form of Periodic Table:
Although the long form of the periodic table is superior to Mendeleef’s periodic table in many
respects, it retains some of the defects such for example:
1. The problem of the position of hydrogen still remains unsolved.
2. It also fails to accommodate the lanthanides and actinides in the main body of the table.
3. The arrangement is unable to reflect the electronic configuration of many elements.
 Aufban Principle:
This principle states as follows:
“The orbital are filled up with electrons in the increasing order of their energy i.e. The orbital of
minimum energy are filled up first with electrons and then the orbital of higher energy start to fill”.
The relative order of energy of various orbital of an atom as shown in the energy level diagram for
a multi electron atom is as
1s<2s<2p<3s<3p<4s<3d<4p<5s<4d<5p<6s<4f<5d<6P<7s<5f<6d<7p<8s………
 Hund’s Rule of Maximum Multiplicity:
This rule states as:
“The distributions of electrons in a set of degenerate orbital like P x , P y , P z set, d xy , d yz , d zx , d z ,
d x  y set etc of a given sub shell takes place in such a way as to give the maximum number of
unpaired electrons and these electron must have the same direction of their spins i.e. the orbital of a
given sub shell are first filled singly and then the pairing if electrons in each orbital begins”.
2
2
2
In order to illustrate the principal, let us consider the electronic configuration of oxygen atom,
whose atomic number is 8. Let the eight electrons in this atom are numbered as 1, 2, 3……7, 8.
Quite evidently the electrons numbered as 1 and 2 will enter 1s orbital and will have opposite
spins. Electrons 3 and 4 will occupy 2s orbital to make it completely filled. These electrons will
also have opposite spins. Now according to Hunt’s rule of maximum multiplicity, the degenerate
2P x , 2P y and 2P z orbital will first be filled singly with electrons numbered as 5, 6 and 7 and all
these three electrons will have same spins. Electron number 8, which is the last electron, will pair
with the electron 5 in 2P x orbital. Thus the complete electronic configuration of oxygen atom in its
ground state will be as given below –
Energy
Increasing
↑↓
5, 8
↑
6
↑
7
2p
2nd shell
↑↓
3, 4
2s
↑↓
1, 2
1s
1st shell
Fig: Electronic configuration of oxygen atom.
 Types of Elements on the Basis of Their Electronic Configuration:
1. Differentiating Electron Classification: This classification divides the elements into four type’s viz.
s, p, d and f block elements depending on the nature of the atomic orbital into which the last
electron (called the differentiating electron) enters.
a) s-block elements: In these elements the differentiating electron enters the s-orbital which is being
progressively filled. The elements of groups IA (H to Fr) and IIA (Be to Ra) belong to this block.
The members of this block thus lie on the extreme left of the periodic table.
b) p-block elements: The elements in which p-orbital are being progressively filled are called p-block.
The elements of the groups IIIA, IVA, VA, VIA, VIIA and zero (except He) are the members of
this block. These elements lie on the extreme right of the periodic table.
c) d-block elements: the elements in which the differentiating electron enters the (n-1) d orbital of
(n-1)th main shell is called d-block elements. These are placed in the middle of the periodic table i.e.
between s- and p-block elements.
d) f-block elements: The elements in which the extra electron enters the (n-2) f orbital of the (n-2)
the main shell are called f-block element.
s-block
elements
d-block
elements
p-block
elements
f-block
elements
2. Bohr’s Classification: In this classification proposed by Bohr the elements are grouped into four
classes depending on the number of incomplete shells of electrons in the atom.
i. Inert gases: In the atoms of these elements the s- and p-sub shells of the outermost shell are
completely filled, thus the valence shell configuration of these elements is ns2p6 with the exception
of He which has 1s2 configuration.
ii. Representative or Normal elements: In the atoms of these elements the outermost shell, i.e. n th
shell only is incompletely filled while all the underlying level are filled to their capacity.
The element
may be divided into two types,
a) s-block elements: The outermost shell configurations of these elements vary from ns1 to ns2.
b) p-block elements: The atoms of these elements have their outermost shell configuration represented
by ns2p1 to ns2p5.
iii. Transition elements: In these the outermost two shells are incomplete. Atoms of these elements
have (n-1) d 1-10 ns 0, 1, 2 general electronic configurations.
iv. Inner Transition Elements: The atom of these elements has outer three shells incomplete. These
three incomplete shells are,
a) Outer most ns-orbital
b) (n-1) d orbital and
c) (n-2) f orbital.
The atoms of these elements have their general electronic configuration as
(n-2) f 1-14, (n-1) d o or 1 or 2 ns2.
 Hund’s Rule of Maximum Multiplicity:
This rule states as:
“The distributions of electrons in a set of degenerate orbital like P x , P y , P z set, d xy , d yz , d zx , d z ,
d x  y set etc of a given sub shell takes place in such a way as to give the maximum number of
unpaired electrons and these electron must have the same direction of their spins i.e. the orbital of a
given sub shell are first filled singly and then the pairing if electrons in each orbital begins”.
2
2
2
In order to illustrate the principal, let us consider the electronic configuration of oxygen atom,
whose atomic number is 8. Let the eight electrons in this atom are numbered as 1, 2, 3……7, 8.
Quite evidently the electrons numbered as 1 and 2 will enter 1s orbital and will have opposite
spins. Electrons 3 and 4 will occupy 2s orbital to make it completely filled. These electrons will
also have opposite spins. Now according to Hunt’s rule of maximum multiplicity, the degenerate
2P x , 2P y and 2P z orbital will first be filled singly with electrons numbered as 5, 6 and 7 and all
these three electrons will have same spins. Electron number 8, which is the last electron, will pair
with the electron 5 in 2P x orbital. Thus the complete electronic configuration of oxygen atom in its
ground state will be as given below –
Energy
Increasing
↑↓
5, 8
↑
6
↑
7
2p
2nd shell
↑↓
3, 4
2s
↑↓
1, 2
1s
1st shell
Fig: Electronic configuration of oxygen atom.
 Types of Elements on the Basis of Their Electronic Configuration:
1. Differentiating Electron Classification: This classification divides the elements into four type’s viz.
s, p, d and f block elements depending on the nature of the atomic orbital into which the last
electron (called the differentiating electron) enters.
a) s-block elements: In these elements the differentiating electron enters the s-orbital which is being
progressively filled. The elements of groups IA (H to Fr) and IIA (Be to Ra) belong to this block.
The members of this block thus lie on the extreme left of the periodic table.
b) p-block elements: The elements in which p-orbital are being progressively filled are called p-block.
The elements of the groups IIIA, IVA, VA, VIA, VIIA and zero (except He) are the members of
this block. These elements lie on the extreme right of the periodic table.
c) d-block elements: the elements in which the differentiating electron enters the (n-1) d orbital of
(n-1)th main shell is called d-block elements. These are placed in the middle of the periodic table i.e.
between s- and p-block elements.
d) f-block elements: The elements in which the extra electron enters the (n-2) f orbital of the (n-2)
the main shell are called f-block element.
s-block
elements
d-block
elements
p-block
elements
f-block
elements
2. Bohr’s Classification: In this classification proposed by Bohr the elements are grouped into four
classes depending on the number of incomplete shells of electrons in the atom.
iii. Inert gases: In the atoms of these elements the s- and p-sub shells of the outermost shell are
completely filled, thus the valence shell configuration of these elements is ns2p6 with the exception
of He which has 1s2 configuration.
iv. Representative or Normal elements: In the atoms of these elements the outermost shell, i.e. n th
shell only is incompletely filled while all the underlying level are filled to their capacity.
The element
may be divided into two types,
c) s-block elements: The outermost shell configurations of these elements vary from ns1 to ns2.
d) p-block elements: The atoms of these elements have their outermost shell configuration represented
by ns2p1 to ns2p5.
v. Transition elements: In these the outermost two shells are incomplete. Atoms of these elements
have (n-1) d 1-10 ns 0, 1, 2 general electronic configurations.
vi. Inner Transition Elements: The atom of these elements has outer three shells incomplete. These
three incomplete shells are,
a) Outer most ns-orbital
b) (n-1) d orbital and
c) (n-2) f orbital.
The atoms of these elements have their general electronic configuration as
(n-2) f 1-14, (n-1) d o or 1 or 2 ns2.
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