Chemistry
1. Explain the principles and rules governing the electronic configuration of atoms.
The arrangement of electrons in various atomic orbitals, known as electronic
configuration of atom is done by following Aufbau principle, Pauli exclusion principle
and Hund‟s rule. They are discussed below:
Aufbau Principle: “ Electrons are filled in the increasing order of energy level”.
According to this principle the electrons occupy the orbitals with lowest energy.
This is decided by the sum of the principle quantum number and azimuthal quantum
number. This is called (n + l) rule.
Rule 1: The electrons first occupy that orbital for which (n + l) value is the lowest.
Rule 2: When (n + l) values for two orbitals are equal, then the electrons first
occupy the orbital with lower value of n.
Illustration of (n + l) rule of Aufbau principle:
* For 1s orbital n + l = 1 + 0 and for 2s orbital n + l = 2 + 0 = 2. Therefore, according
to rule 1, first the electrons occupy 1s orbital, then 2s orbital.
* For 2p orbital n + l = 2 + 1 = 3, and for 3s orbital n + l = 3 + 0 = 3, the values of
n + l are equal. Now according to rule 2, first the electrons will occupy 2p orbitals then 3s
orbital.
Following the (n + l) of Aufbau principle, the orbitals in increasing order of energy
are arranged as:
1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p.
Pauli’s Exclusion Principle: “In an atom no two electrons can have the same set
of four quantum numbers”:
Illustration of Pauli’s exclusion principle:
1. In an atom if one electron is assigned a set of four quantum number n = 1, l = 0,
m = 0, s = + ½, then other electrons cannot be assigned the same set of quantum
numbers.
2. If three quantum numbers for two electrons are the same, then these electrons must
have different fourth quantum number.
n L m s
+½ Different
First electron 1 0 0
–½
Second electron 1 0 0
Hund’s Rule: “Among the orbitals of same energy, electrons do not start pairing,
until all these orbitals are singly occupied”.
Importance of Hund’s rule: This rule provides guidelines of filling electrons in
the degenerate orbitals (orbitals having equal energies) of an atom. According to this rule:
1. If the number of electrons is equal to (or less than) the number of degenerate
, orbitals, then orbitals are singly occupied.
2. Pairing of electron spins takes place only when each one of the degenerate orbitals
is singly occupied. This is possible only when the number of electrons to be filled
is grater than the number of degenerate orbitals.
2. Illustrate with examples the use of Aufbau principle, Pauling’s exclusion principle
and Hunds rule in writting the electronic configuration of atom.
In Hydrogen atom, these is only one electron which occupies 1s orbital and the electronic
state is represented by
1s
Hydrogen 1s1
In Helium atom, the second electron is also in the 1s state and its spin is paired,
with the first electron so that this orbital is complete.
1s
Helium 1s2
The third electron is Lithium would occupy 2s orbital which has the minimum
energy in this shell. In the atom beryllium, the fourth electron completes the 2s orbital and
thus with boron, the fifth electron must enter 2px orbital.
1s 2s 2px 2py 2pz
Lithium 1s2 2s1
1s 2s 2px 2py 2pz
Beryllium 1s2 2s2
1s 2s 2px 2py 2pz
Boron 1s2 2s2 2p1
There are three 2p orbitals. (Six electrons can be accommodated in 2p sub shell).
The sixth electron of carbon occupies the 2py orbital and similarly, in nitrogen, the seventh
electron occupies the 2pz orbital.
1s 2s 2px 2py 2pz
Carbon 1s2 2s2 2p2
1s 2s 2px 2py 2pz
Nitrogen 1s2 2s2 2p3
, Once these three orbitals contain an electron each, the introduction of further
electrons, as in oxygen, fluorine and neon brings out completion. With neon the shell
(energy level) n = 2 is complete.
Continuation of Qns. No. 2
1s 2s 2px 2py 2pz
Oxygen 1s2 2s2 2p4
1s 2s 2px 2py 2pz
Fluorine 1s2 2s2 2p5
1s 2s 2px 2py 2pz
Neon 1s2 2s2 2p6 = [Ne]
3. What is oxidation number? Explain the rules which are used to deter mine the
oxidation number.
Oxidation number is defined as the apparent charge assigned to an atom of the
element in a compound according to set rules.
The rules are:
(i) In the elementary state, oxidation number of an atom is zero
e.g., O2, H2, S, Na etc.
(ii) In a monoatomic ion, charge on the ion itself represents oxidation number.
e.g., Oxidation numbers of Na+ , Ba2+ , Al3+, Cl– , S2–, are + 1, + 2, + 3, –1, –2
among cases repectively.
(iii) Fluorine, the most electronegative element has an oxidation number – 1, in all its
compounds.
(iv) Oxidation number of hydrogen is always + 1, in its compounds except metal
hydrides (e.g., LiH NaH) in which it is – 1.
(v) Oxidation number of oxygen is – 2 in all its compounds except in peroxides (e.g.,
H2O2, Na2 O2) and in OF2 where it is – 1 and + 2 respectively.
(vi) Oxidation number of alkali metals (Li, Na, K, Rb, Cs and Fr) is always + 1 in their
compounds.
(vii) Oxidation number of alkaline earth metals (Be, Mg, Ca, Ba, Sr, Ra) is always + 2 in
their compounds.
1. Explain the principles and rules governing the electronic configuration of atoms.
The arrangement of electrons in various atomic orbitals, known as electronic
configuration of atom is done by following Aufbau principle, Pauli exclusion principle
and Hund‟s rule. They are discussed below:
Aufbau Principle: “ Electrons are filled in the increasing order of energy level”.
According to this principle the electrons occupy the orbitals with lowest energy.
This is decided by the sum of the principle quantum number and azimuthal quantum
number. This is called (n + l) rule.
Rule 1: The electrons first occupy that orbital for which (n + l) value is the lowest.
Rule 2: When (n + l) values for two orbitals are equal, then the electrons first
occupy the orbital with lower value of n.
Illustration of (n + l) rule of Aufbau principle:
* For 1s orbital n + l = 1 + 0 and for 2s orbital n + l = 2 + 0 = 2. Therefore, according
to rule 1, first the electrons occupy 1s orbital, then 2s orbital.
* For 2p orbital n + l = 2 + 1 = 3, and for 3s orbital n + l = 3 + 0 = 3, the values of
n + l are equal. Now according to rule 2, first the electrons will occupy 2p orbitals then 3s
orbital.
Following the (n + l) of Aufbau principle, the orbitals in increasing order of energy
are arranged as:
1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p.
Pauli’s Exclusion Principle: “In an atom no two electrons can have the same set
of four quantum numbers”:
Illustration of Pauli’s exclusion principle:
1. In an atom if one electron is assigned a set of four quantum number n = 1, l = 0,
m = 0, s = + ½, then other electrons cannot be assigned the same set of quantum
numbers.
2. If three quantum numbers for two electrons are the same, then these electrons must
have different fourth quantum number.
n L m s
+½ Different
First electron 1 0 0
–½
Second electron 1 0 0
Hund’s Rule: “Among the orbitals of same energy, electrons do not start pairing,
until all these orbitals are singly occupied”.
Importance of Hund’s rule: This rule provides guidelines of filling electrons in
the degenerate orbitals (orbitals having equal energies) of an atom. According to this rule:
1. If the number of electrons is equal to (or less than) the number of degenerate
, orbitals, then orbitals are singly occupied.
2. Pairing of electron spins takes place only when each one of the degenerate orbitals
is singly occupied. This is possible only when the number of electrons to be filled
is grater than the number of degenerate orbitals.
2. Illustrate with examples the use of Aufbau principle, Pauling’s exclusion principle
and Hunds rule in writting the electronic configuration of atom.
In Hydrogen atom, these is only one electron which occupies 1s orbital and the electronic
state is represented by
1s
Hydrogen 1s1
In Helium atom, the second electron is also in the 1s state and its spin is paired,
with the first electron so that this orbital is complete.
1s
Helium 1s2
The third electron is Lithium would occupy 2s orbital which has the minimum
energy in this shell. In the atom beryllium, the fourth electron completes the 2s orbital and
thus with boron, the fifth electron must enter 2px orbital.
1s 2s 2px 2py 2pz
Lithium 1s2 2s1
1s 2s 2px 2py 2pz
Beryllium 1s2 2s2
1s 2s 2px 2py 2pz
Boron 1s2 2s2 2p1
There are three 2p orbitals. (Six electrons can be accommodated in 2p sub shell).
The sixth electron of carbon occupies the 2py orbital and similarly, in nitrogen, the seventh
electron occupies the 2pz orbital.
1s 2s 2px 2py 2pz
Carbon 1s2 2s2 2p2
1s 2s 2px 2py 2pz
Nitrogen 1s2 2s2 2p3
, Once these three orbitals contain an electron each, the introduction of further
electrons, as in oxygen, fluorine and neon brings out completion. With neon the shell
(energy level) n = 2 is complete.
Continuation of Qns. No. 2
1s 2s 2px 2py 2pz
Oxygen 1s2 2s2 2p4
1s 2s 2px 2py 2pz
Fluorine 1s2 2s2 2p5
1s 2s 2px 2py 2pz
Neon 1s2 2s2 2p6 = [Ne]
3. What is oxidation number? Explain the rules which are used to deter mine the
oxidation number.
Oxidation number is defined as the apparent charge assigned to an atom of the
element in a compound according to set rules.
The rules are:
(i) In the elementary state, oxidation number of an atom is zero
e.g., O2, H2, S, Na etc.
(ii) In a monoatomic ion, charge on the ion itself represents oxidation number.
e.g., Oxidation numbers of Na+ , Ba2+ , Al3+, Cl– , S2–, are + 1, + 2, + 3, –1, –2
among cases repectively.
(iii) Fluorine, the most electronegative element has an oxidation number – 1, in all its
compounds.
(iv) Oxidation number of hydrogen is always + 1, in its compounds except metal
hydrides (e.g., LiH NaH) in which it is – 1.
(v) Oxidation number of oxygen is – 2 in all its compounds except in peroxides (e.g.,
H2O2, Na2 O2) and in OF2 where it is – 1 and + 2 respectively.
(vi) Oxidation number of alkali metals (Li, Na, K, Rb, Cs and Fr) is always + 1 in their
compounds.
(vii) Oxidation number of alkaline earth metals (Be, Mg, Ca, Ba, Sr, Ra) is always + 2 in
their compounds.
