Variation of periodic properties across the periodic table

1. Periodicity
&
their properties

2. PERIODICITY
Repetition of properties of elements after a certain interval when the elements all
arranged in increasing order of atomic number.
PERIODIC PROPERTIES OF ELEMENTS
ATOMIC RADIUS
It is distance between outermost electron and nucleus.
Atomic radius depends on the type of chemical bond between atoms in a
molecule. These are :
1. Covalent radius 2. Ionic radius 3. Metallic radius 4. Vander waal’s radius

3. PERIODIC TRENDS
ATOMIC SIZE/ATOMIC RADIUS
atomic size/radius increases from top to bottom (due to increasing
number of shells thus decreasing effective nuclear charge, Zeff)
atomic size/radius decreases with
increasing atomic number, number of
shells remain same but number of
electrons & protons increases, thus
increasing effective nuclear charge, Zeff)
Along the period
Top
to
bottom

4. Factors affecting atomic size
i) Effective nuclear charge
Atomic radius ∝
(ii) Number of shells
Atomic radius ∝ No. of shells
(iii) screening effect
Atomic radius ∝ Screening effect
(iv) Magnitude of +ve charge
Atomic radius ∝
(v) Magnitude of -ve charge
Atomic radius ∝ Magnitude of -ve charge
(vi) Bond order
Atomic radius ∝
Effective nuclear charge (Zeff)
1
1
Magnitude of +ve charge
1
Bond order

5. IE increases with atomic number increase due
to the decrease in size & increase in effective
nuclear charge (Zeff) along the period
IE decrease with increase in size & number of shells increases
So effective nuclear charge (Zeff ) decreases along the group
IONISATION ENERGY
Ionisation energy (IE) is defined as the amount of energy required to remove the most
loosely bound electron from an isolated gaseous atom to form a cation.
Mg
+ Mg
++ + e-IE2
Mg
++ Mg
+++ + e-IE3
Mg Mg
+ + e-IE1
(IE)1 < (IE)2 < (IE)3
Along the period
Top
to
bottom
Perioddic trends

6. Factors Influencing Ionisation energy
(i) Ionisation energy ∝
(ii) Ionisation energy ∝ Effective nuclear charge (Zeff)
(iii) Ionisation energy ∝
(iv) Electronic Configuration:
If an atom has exactly half-filled or completely filled orbitals, then such an
arrangement has extra stability.
1
Atomic size
1
Screening effect

7. ELECTRON AFFINITY
Electron affinity is conventionally defined as the energy released when an electron is
added to the valence shell of an isolated gaseous atom.
Mg + e- Mg
- Exothermic reactionEA1
Mg
- + e- Mg
- - Endothermic reactionEA2
EA increases
atomic number increase
decrease in size
effective nuclear charge (Zeff) increases
EA decrease
size increase
number of shells increases
effective nuclear charge (Zeff ) decreases
Along the period
Top
to
bottom

8. (i) Electron affinity ∝
(ii) Electron affinity ∝ Effective nuclear charge (zeffs)
(iii) Electron affinity ∝
(iv) Stability of half filled and completely filled orbitals of a subshell is comparatively
more and the addition of an extra electron to such an system is difficult and hence the
electron affinity value decreases.
1
Atomic size
1
Screening effect
Factors Influencing Electron affinity

9. Variation of electronegativity in a group Variation of electronegativity in a period
On moving down the groups, Z increases
but Zeff almost remains constant, number
of shells (n) increases, rn (atomic radius)
increases. Therefore, electronegativity
decreases moving down the groups.
While moving across a period left to
right, Z, Zeff increases & rn decreases.
Therefore, electronegativity increases
along a period.
Electronegativity:
Electronegativity is a measure of the tendency of an element to attract shared
electrons towards itself in a covalently bonded molecules.

10. Method to measure the value of electronegativity
(a) Pauling’s scale :
Linus Pauling developed a method for calculating relative electronegativities of most
elements. According to Pauling
(b) Mulliken’s scale :
Electronegativity can be regarded as the average of the ionisation energy (IE) and
the electron affinity (EA) of an atom (both expressed in electron volts).

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