Periodic Properties
By-
A.P.S. BHADOURIYA
M.Sc. (L.U.), B.Ed.
NET (CSIR-UGC) Qualified
Periodic Properties
Periodic Trends in Physical Properties
Shielding effect & Effective Nuclear Charge
Atomic Radius
Ionic Radius
Ionization Enthalpy
Electron Gain Enthalpy
Electronegativity
Periodic Properties
Periodic Trends in Chemical Properties
 Periodicity of Valence or Oxidation States
 Anomalous Properties of Second Period
Elements
 Chemical Reactivity
Shielding effect & Effective Nuclear Charge
The decrease in nuclear charge ( nuclear
force of attraction) on outermost shell
electrons due to repulsion caused by inner
shell electron is known as shielding effect of
inner shell or intervening electrons on outer
shell electron.
Shielding effect & Effective Nuclear Charge
Due to shielding effect the nuclear charge is lowered
on outermost shell electrons, the net nuclear
charge acting on outermost shell electrons is known
as Effective Nuclear Charge. It is denoted by Z* or
Zeff.


Z* or Zeff. = Z - σ
where Z = nuclear charge( = atomic No.) & σ =
shielding constant or screening constant , it is a
measure of shielding effect
Shielding effect & Effective Nuclear Charge
Determination of ENC (Z*)
If the electron resides in s or p orbital
1. Electrons in principal shell higher than the e- in
question contribute 0 to σ .
2. Each electron in the same principal shell contribute
0.35 to σ (0.30 if it is 1S shell).
3. Electrons in (n-1) shell each contribute 0.85 to σ .
4. Eelectrons in deeper shell each contribute 1.00 to σ
Shielding effect & Effective Nuclear Charge
Determination of ENC (Z*)
If the electron resides in d or f orbital
1. All e-s in higher principal shell contribute 0 to
2. Each e- in same shell contribute 0.35 to σ
3. All inner shells in (n-1) and lower contribute
1.00 to σ
σ
Shielding effect & Effective Nuclear Charge
Determination of ENC (Z*)
e.g. Calculate the Z* for the 2p electron Fluorine
(Z = 9) 1s2, 2s 2p5.
Soln. Screening constant for one of the outer electron
 6 (six) (two 2s e- and four 2p e-) = 6 X 0.35 = 2.10
 2 (two)1s e- = 2 X 0.85 = 1.70
 σ = 1.70+2.10 = 3.80
 Z* = 9 - 3.80 = 5.20
Shielding effect & Effective Nuclear Charge
Trend of ENC in Periodic Table
In a Period - Effective nuclear charge Z*
increases increases rapidly along a
period(0.65 per next group)
e.g.
Li
1.3
Be
B
1.95 2.6
C
N
O
F
Ne
3.3
3.9
4.6
5.2
5.9
Shielding effect & Effective Nuclear Charge
Trend of ENC in Periodic Table
 In a Group - Effective nuclear charge Z* increases
slowly along a group.
e.g.
Gr-1
H
Li
Na
K
Rb
Cs
Z*
1.0
1.3
2.2
2.2
2.2
2.2
PERIODIC TREND OF ATOMIC RADIUS
In A Period
atomic radius decreases with increase in atomic number
(in a period left to right)
BECAUSE in a period left to right-
 1. n (number of shells) remain constant.
 2. Z increases (by one unit)
 3. Z* increases (by 0.65 unit)
 4. Electrons are pulled close to the nucleus by the increased
Z*
 In a groupAtomic radius increases moving down the group

Because, along a group top to bottom
1. n increases
2. Z increases
3. No dramatic increase in Z* - almost remains
constant
IONIC RADII

All anions are larger than their parent atoms.
because the addition of one or more electrons would result
in increased repulsion among the electrons and a decrease
in ENC.

The cations are smaller than their parent atoms
because it has fewer electrons while its nuclear charge
remains the same & hence ENC is greater in cation than its
parent atom
ISOELECTRONIC SPECIES
 Atoms and ions which contain the same number of
electrons, are called as isoelectronic species.
For example, F–, Na+ and Mg2+ have the same number of
electrons(=10).
 The size of isoelectronic species decreases with increase in
nuclear charge. e.g.-
o2->F- >Ne>Na+>Mg2+>Al3+
---------SIZE DECREASING------
Atomic Radius
NOTE:
Metallic radii in the third row d-block are similar to
the second row d-block, but not larger as one
would expect given their larger number of
electrons.
This is due to Lanthanide Contraction as f-orbitals
have poor shielding properties.
Ionisation Energy (IE) or
Ionisation Enthalpy (ΔiH )
 Ionization: removing an electron from an atom or ion
 Ionization energy: energy required to remove an electron
from an isolated, gaseous atom or ion is called as Ionization
energy or ionisation enthalpy.
 If the atom is neutral the above defined ionisation energy
is called as first ionisation enthalpy.
 Energy required to remove an electron from an isolated,
monovalent cation is called as second Ionization energy.
 The ionization enthalpy is expressed in units of kJ /mol
Ionisation Energy (IE) or
Ionisation Enthalpy (ΔiH )
X(g) + energy → X+(g) + e–.
1st ionisation enthalpy
X+(g) + energy → X++(g) + e–.
2nd ionisation enthalpy
Ionisation Energy (IE) or
Ionisation Enthalpy (ΔiH )
The second ionization enthalpy will be higher than
the first ionization enthalpy because it is more
difficult to remove an electron from a positively
charged ion than from a neutral atom because a
cation has greater ENC than a neutral atom.
In the same way the third ionization enthalpy will
be higher than the second and so on.
Factors affecting Ionisation Enthalpy (ΔiH )
(a) Size of the atom - IE decreases as the size of the
atom increases
(b) Nuclear Charge - IE increases with increase in
nuclear charge
(c) The type of electron - Shielding effect, Penetration
effect
(e)Electronic configuration: e.g. noble gases passes
very high value of IE due to stable octet
configuration
Periodic Trend of Ionisation Enthalpy (ΔiH )
On moving down a group
1. nuclear charge increases
2. Z* due to screening is almost constant
3. number of shells increases, hence atomic size
increases.
4. there is a increase in the number of inner electrons
which shield the valence electrons from the nucleus
Thus IE decreases down the group
Periodic Trend of Ionisation Enthalpy
On moving across a period(L--->R)
1. the atomic size decreases
2. Effective nuclear charge increases
Thus IE increases along a period
However there are some exceptions also e.g.

IE of Be is higher than that of B.

IE of N is higher than that of O.
Periodic Trend of Ionisation Enthalpy (ΔiH )
Explain why-
(a). IE of Be is higher than that of B.
Ans. - In beryllium(1s2,2s2 ), the electron removed during the
ionization is an s-electron whereas the electron removed
during ionization of boron(1s2,2s2,2p1) is a p-electron. The
penetration of a 2s-electron to the nucleus is more than
that of a 2p-electron; hence the 2p electron of boron is more
shielded from the nucleus by the inner core of electrons
than the 2s electrons of beryllium.
Therefore, it is easier to remove the 2p-electron from boron
compared to the removal of a 2s- electron from beryllium.
Thus, boron has a smaller first ionization.
Periodic Trend of Ionisation Enthalpy (ΔiH )
(b) Why IE of N is higher than that of O.
Ans. The first ionization enthalpy of oxygen compared to
nitrogen is smaller. This arises because in the nitrogen
atom(1s2,2s2,2p3) three 2p-electrons reside in different
atomic orbitals (Hund’s rule) whereas in the oxygen atom
(1s2,2s2,2p4), two of the four 2p-electrons must occupy the
same 2p-orbital resulting in an increased electron-electron
repulsion. Consequently, it is easier to remove the fourth
2p-electron from oxygen than it is, to remove one of the
three 2p-electrons from nitrogen.
Electron Gain Enthalpy (ΔegH)
 When an electron is added to a neutral gaseous atom (X) to
convert it into a negative ion, the enthalpy change
accompanying the process is defined as the Electron
GainEnthalpy (ΔegH) or Electron Affinity.


Electron gain enthalpy provides a measure of the ease
with which an atom adds an electron to form anion as
represented by equation –
X(g) + e --- X- (g)+energy
(electrongainenthalpy)
Electron Gain Enthalpy (ΔegH)
 Depending on the element, the process of
adding an electron to the atom can be either
endothermic or exothermic.
 For many elements energy is released when an
electron is added to the atom and the electron gain
enthalpy is negative.
Factors Affecting E G E (ΔegH)
 ENC- With increase in ENC, the force of attraction exerted by
the nucleus on the electrons increases. Consequently, the
atom has a greater tendency to attract additional electron
i.e., its EGE increases i.e. become more negative.
 ATOMIC SIZEWith decrease in size ENC increases & hence EGE
increases.
 ELECTRONIC CONFIGURATIONThe value of EGE depends effectively upon electronic
configuration of elements, elements with stable electronic
configuration posses lower (less -ve) value of EGE, e.g.-
Factors Affecting E G E (ΔegH)
A. Noble gases have practically zero or +ve EGEs. This
is because they have no tendency to gain an
additional electron as they already have the stable
ns2np6 configuration
B. Halogens have high electron affinities. This is due
to their strong tendency to gain an additional
electron to change into the stable ns2np6
configuration.
PERIODIC TREND OF EGE (ΔegH)
IN A PERIODThe EGE increases i.e. become more negative as we move
across a period because the atomic size decreases and hence
the force of attraction exerted by the nucleus on the
electrons increases. Consequently, the atom has a greater
tendency to attract additional electron i.e., its electron
affinity increases
IN A GROUPThe EGE decreases (-)vely because the atomic size increases
and therefore, the effective nuclear attraction decreases and
thus electron affinity decreases
Electron Gain Enthalpy (ΔegH)
 Explain why –
(a). electron gain enthalpy of O is less than that of the S.
(b). electron gain enthalpy of F is less than that of the Cl.
 Ans:- The electron gain enthalpy of O or F is less than that
of the succeeding element. This is because when an electron
is added to O or F, the added electron goes to the smaller n
= 2 quantum level and suffers significant repulsion from the
other electrons present in this level. For the n = 3 quantum
level (S or Cl ), the added electron occupies a larger region
of space and the electron-electron repulsion is much less.
Electronegativity
 The tendency of an element in a molecule
to attract the shared pair of electrons
towards itself is known as electronegativity.
It is measured on Pauling scale in which F
(most EN element)is attributed to a value of
4.
Periodic trend of EN
In a Group- on moving down the group,
 Z increases but Z* almost remains constant
 number of shells (n) increases
 atomic radius increases
 force of attraction between added electron and
nucleus decreases
 Therefore EN decreases moving down the group
Periodic trend of EN
In a Period- On moving across a period left to right
 Z and Z* increases
 number of shells remains constant
 atomic radius decreases
 force of attraction between shared electron and
nucleus increases
Hence EN increases along a period
Periodic Trends in Chemical Properties
Periodicity of Valence or Oxidation States
 Anomalous Properties of Second Period
Elements
 Chemical Reactivity
Periodicity of Valence or Oxidation States
The valence of representative elements is usually (though
not necessarily) equal to the number of electrons in the
outer most orbitals and / or equal to eight minus the
number of outermost Electrons(w.r.t. H)
Some periodic trends observed in the valence of elements
(hydrides and oxides) are shown in Table
Group
1
2
13
14
Number of
valence
electron
1
2
3
3
Valence
1
2
4
3
15
16
17
5
6
7
3,5
2,6
1,7
Periodicity of Valence or Oxidation States
The oxidation state of an element in a particular
compound can be defined as the charge acquired by
its atom on the basis of electronegative
consideration from other atoms in the molecule.