Periodicity
In chemistry periodicity of elements means the recurrence of similar properties of the elements
after certain regular intervals when they are arranged in the order of increasing atomic numbers.
Cause of Periodicity of Elements
The cause of periodicity of the properties of elements is the repetition of similar electronic
configuration of their atoms in the outermost energy shell (or valence shell) after certain regular
intervals. To explain the cause of periodicity of elements we are taking two examples.
The chemical properties of all alkali metals resemble each other. This is due to the fact that
chemical properties of elements depend primarily on the arrangement of electrons in the outermost
shell (valence shell). The electronic configuration of the atoms of alkali metal group shows that all
the alkali metals have one electron in their outermost s-orbital. So, their electronic configuration
may be represented as [noble gas] ns1.
Electronic configuration:
Arrangement of electrons in various orbits and orbitals is called electronic configuration. It is
determined by the principles of quantum mechanics. Electrons can be distributed in orbits(shells)
and orbitals(subshells) of an atom whereas this distribution follows a pattern which is governed by
some rules. So to understand the concept of electronic configuration, it is important to know about
the basic concept of orbit, orbital and rules for the filling of electrons in orbitals.
6.3.1 Orbit(shells):
Orbit can be defined as the hypothetical pathway of an electron present around the nucleus. Shells
can be named as K, L, M, N, O, P and Q; or 1, 2, 3, 4, 5, 6 and 7 going from the innermost to the
outermost. The shell that is closest to the nucleus is named as “K shell” (also called “1 shell”),
while this shell is followed by the “L shell” (also called “2 shell”), then comes the “M shells”
(also called “3 shell”) and so on farther from the nucleus. A fixed electron number can be contained
by a shell. such as
K shell: only 2 electrons can be hold by the K shell.
L shell: L shell can have up to eight (2+6) electrons.
M shells: 18 electrons (2+6+10) can be occupied by the M shell.
And so on.
The filling of electrons in shells follow the general formula
2n2
where n= 1 for K shell, 2 for L and so on
for example Na has 11 electrons. The distribution of electrons in different shells will be as
follows
K shell = 2 electrons, L shell = 8 electrons, M shell = 1 electrons
,6.3.2 Orbitals(subshells):
Each orbit contains some subshells or orbitals. Atomic orbital is a physical space or region where
probability of finding the electron in an atom is most or where the electron can be calculated to
present. Each orbital that is characterized by the set of values of the three quantum numbers such
as n, l and m can have only two electron each with its own spin quantum number. The orbitals are
named as s orbital, p orbital, d orbital and f orbital. Each shell is composed of one or more
subshells.
s subshell: s subshell is a spherical orbital. Only two electrons can accommodate in it. Each of the
energy level has at least one s subshell.
P orbital: This subshell has dumbbell-shaped three orbitals. All are present at right angle to one
another. Each of the orbital can hold only two electrons. In the energy states of two or higher p
orbital can be present.
S px py pz
Shapes of s and porbitals
d orbital: This subshell contain five orbitals and they are arranged in x, y and z axis. Each of the
orbitals can hold only two electrons. So, when a d subshell is fully completed it will have 10
electrons. In the energy level of three or higher this orbital is present.
Shapes of d orbitals
, f subshell: f subshell contains seven orbitals with symmetrical distribution in x, y and z axis. Each
of the orbital can hold only two electrons. If a f subshell is fully loaded then it will have 14
electrons. In the energy states of four or higher this subshell can be found.
Rules for filling of orbitals
Filling of electrons in different orbitals follow some general rules, which are following
1. Pauli Exclusion principle
2. Auf Bau principle
3. Hund’s rule
4. n+l rule
Pauli Exclusion principle:
This principle states that for electrons in an atom, in a poly electron system it is impossible for two
electrons to have same values of four quantum numbers, i.e. principle quantum number (n),
azimuthal quantum number (l), magnetic quantum number (mℓ) and spin quantum number (ms).
Auf Bau principle:
Also known as building up principle states that atomic orbitals of lower energy levels will be filled
first. For example 1s subshell is filled before the filling of 2s subshell.
Hund’s rule:
This rule states that in sublevel every orbital is singly filled before any orbital is doubly filled. And
also, the electrons have same spin in singly occupied orbital.
n+l rule:
n+l rule is used to know which of the orbital will be filled first with the electrons. For example,
the orbitals which have same values of n+l, the orbital with lower value of n (principle quantum
number) will be first filled. i.e. 3d will be filled before the filling of 4p.
Rules for Assigning Electron Orbitals:
Orbitals occupation:
Orbitals are filled by electrons in a way to minimize the energy of atom. In an atom, electrons fill
the orbitals in principle energy levels in order of increasing energy (electrons getting far from
nucleus). The order look like
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, and 7p
To remember this pattern, probably the easiest, is to refer to the periodic table and remember where
each orbital block falls to logically deduce this pattern. Another way is to make a table like the one
below and use vertical lines to determine which subshells correspond with each other.
In chemistry periodicity of elements means the recurrence of similar properties of the elements
after certain regular intervals when they are arranged in the order of increasing atomic numbers.
Cause of Periodicity of Elements
The cause of periodicity of the properties of elements is the repetition of similar electronic
configuration of their atoms in the outermost energy shell (or valence shell) after certain regular
intervals. To explain the cause of periodicity of elements we are taking two examples.
The chemical properties of all alkali metals resemble each other. This is due to the fact that
chemical properties of elements depend primarily on the arrangement of electrons in the outermost
shell (valence shell). The electronic configuration of the atoms of alkali metal group shows that all
the alkali metals have one electron in their outermost s-orbital. So, their electronic configuration
may be represented as [noble gas] ns1.
Electronic configuration:
Arrangement of electrons in various orbits and orbitals is called electronic configuration. It is
determined by the principles of quantum mechanics. Electrons can be distributed in orbits(shells)
and orbitals(subshells) of an atom whereas this distribution follows a pattern which is governed by
some rules. So to understand the concept of electronic configuration, it is important to know about
the basic concept of orbit, orbital and rules for the filling of electrons in orbitals.
6.3.1 Orbit(shells):
Orbit can be defined as the hypothetical pathway of an electron present around the nucleus. Shells
can be named as K, L, M, N, O, P and Q; or 1, 2, 3, 4, 5, 6 and 7 going from the innermost to the
outermost. The shell that is closest to the nucleus is named as “K shell” (also called “1 shell”),
while this shell is followed by the “L shell” (also called “2 shell”), then comes the “M shells”
(also called “3 shell”) and so on farther from the nucleus. A fixed electron number can be contained
by a shell. such as
K shell: only 2 electrons can be hold by the K shell.
L shell: L shell can have up to eight (2+6) electrons.
M shells: 18 electrons (2+6+10) can be occupied by the M shell.
And so on.
The filling of electrons in shells follow the general formula
2n2
where n= 1 for K shell, 2 for L and so on
for example Na has 11 electrons. The distribution of electrons in different shells will be as
follows
K shell = 2 electrons, L shell = 8 electrons, M shell = 1 electrons
,6.3.2 Orbitals(subshells):
Each orbit contains some subshells or orbitals. Atomic orbital is a physical space or region where
probability of finding the electron in an atom is most or where the electron can be calculated to
present. Each orbital that is characterized by the set of values of the three quantum numbers such
as n, l and m can have only two electron each with its own spin quantum number. The orbitals are
named as s orbital, p orbital, d orbital and f orbital. Each shell is composed of one or more
subshells.
s subshell: s subshell is a spherical orbital. Only two electrons can accommodate in it. Each of the
energy level has at least one s subshell.
P orbital: This subshell has dumbbell-shaped three orbitals. All are present at right angle to one
another. Each of the orbital can hold only two electrons. In the energy states of two or higher p
orbital can be present.
S px py pz
Shapes of s and porbitals
d orbital: This subshell contain five orbitals and they are arranged in x, y and z axis. Each of the
orbitals can hold only two electrons. So, when a d subshell is fully completed it will have 10
electrons. In the energy level of three or higher this orbital is present.
Shapes of d orbitals
, f subshell: f subshell contains seven orbitals with symmetrical distribution in x, y and z axis. Each
of the orbital can hold only two electrons. If a f subshell is fully loaded then it will have 14
electrons. In the energy states of four or higher this subshell can be found.
Rules for filling of orbitals
Filling of electrons in different orbitals follow some general rules, which are following
1. Pauli Exclusion principle
2. Auf Bau principle
3. Hund’s rule
4. n+l rule
Pauli Exclusion principle:
This principle states that for electrons in an atom, in a poly electron system it is impossible for two
electrons to have same values of four quantum numbers, i.e. principle quantum number (n),
azimuthal quantum number (l), magnetic quantum number (mℓ) and spin quantum number (ms).
Auf Bau principle:
Also known as building up principle states that atomic orbitals of lower energy levels will be filled
first. For example 1s subshell is filled before the filling of 2s subshell.
Hund’s rule:
This rule states that in sublevel every orbital is singly filled before any orbital is doubly filled. And
also, the electrons have same spin in singly occupied orbital.
n+l rule:
n+l rule is used to know which of the orbital will be filled first with the electrons. For example,
the orbitals which have same values of n+l, the orbital with lower value of n (principle quantum
number) will be first filled. i.e. 3d will be filled before the filling of 4p.
Rules for Assigning Electron Orbitals:
Orbitals occupation:
Orbitals are filled by electrons in a way to minimize the energy of atom. In an atom, electrons fill
the orbitals in principle energy levels in order of increasing energy (electrons getting far from
nucleus). The order look like
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, and 7p
To remember this pattern, probably the easiest, is to refer to the periodic table and remember where
each orbital block falls to logically deduce this pattern. Another way is to make a table like the one
below and use vertical lines to determine which subshells correspond with each other.
