Periodic properties

1
Late Ku. Durga K. Banmeru Science College,
LONAR DIST. BULDANA (Maharashtra), India.
Section –I
Unit-1
A) Periodic properties
B. Sc. Ist year Sem-Ist
Subject:- Chemistry

2
Dr. Suryakant B. Borul
(M.Sc., M.Phil., Ph.D.)
Head Of Department
Department of Chemistry
Late Ku. Durga K. Banmeru Science College,
Lonar
Teacher Profile

 Periodic Classification of the elements
 Modern periodic table i.e. long form of periodic table is based on “modern
periodic law”, which states that the properties of “the element are a periodic
function of their atomic numbers.”
 If the elements are arranged in the increasing order of their atomic numbers,
after certain intervals, there will be repetition in the properties of the elements.
 The long form of periodic table is derived from original Mendeleeff’s periodic
table with modifications.
 We know that, as we move from left to right, electrons are successively added at
every element with increase in atomic number.
A) “Periodic properties”

 s, p, d and f-Block Elements
In long form of periodic elements divided into four types on basis of nature
of atomic orbital in which last electron enters. These types are as s, p, d and f
block elements.
For example – Sodium element
Q- What is the block position of Na in periodic table?
Ans- Atomic number of Sodium is 11, therefore electronic configuration of Sodium
as-
11Na- 1s2 , 2s2 sp6, 3s1
the last electron of Na enter in s-orbital, thus the sodium is s-block element.

a) s-Block elements- “The element their last electron enters in ‘s’
atomic orbital such elements are called as s-block elements.”
The elements in these group as –IA and IIA
Groups Elements
IA H, Li, Na, K, Rb, Cs AND Fr
IIA Be, Mg, Ca, Sr, Ba, AND Ra
The general valence shell electronic configuration of these elements varies from
ns1 to ns2 where n= number of valence shell of an element. (i.e. for IA 1s1 to 1s2
and IIA 2s1 to 2s2
The elements of this block lie on the extreme left to periodic table.

b) p-Block elements- “The element their last electron enters in ‘p’ atomic
orbital such elements are called as p-block elements.”
Group IIIA IVA VA VIA VIIA Zero group
13 14 15 16 17 18
B C N O F Ne
Al Si P S Cl Ar
Ga Ge As Se Br Kr
In Sn Sb Te I Xe
Ti Pb Bi Po At Rn
Nh Fl Mc Lv Ts Og
 The general valence shell electronic configuration of these elements varies from ns2 np1 to
ns2 np6 where n= number of valence shell of an element.
The elements in these group as –IIIA and VIIA groups

c) d-Block elements-
“The element their last electron enters in ‘d’atomic orbital such elements are
called as d-block elements.”
This block elements are present in the middle of the periodic table in between s
and p-block elements.
The valence shell electronic configuration of these elements can be given as (n-
1)d1-10, ns 0, 1, 2
These elements also called transition elements.
These elements are divided into four series corresponding to the filling of 2d,
4d, 5d, and 6d orbital’s of (n-1)th main shell.

3d- Series: This is a first transition series and consists of ten elements. (21Sc, 22Ti,
23V, 24Cr, 25Mn, 26Fe, 27Co, 28Ni, 29Cu, 30Zn ). The last electron of these elements
are enter in 3d orbital, hence known as 3d series elements.
4d- Series: This is a second transition series and consists of ten elements. (39Y, to
48Cd). The last electron of these elements are enter in 4d orbital, hence known as
4d series elements.
5d- Series: This is a third transition series and consists of ten elements. (57La, 72Hf
to 80Hg). The last electron of these elements are enter in 5d orbital, hence known
as 5d series elements.
6d- Series: This is a fourth transition series and consists of ten elements.
(89Ac,104Rf to112Cn). The last electron of these elements are enter in 6d orbital,
hence known as 4d series elements.

d) f-Block elements-
“The element their last or additional electron enters in ‘f’ atomic orbital
such elements are called as f-block elements.”
This block elements are present at the bottom of the periodic table.
The general valence shell electronic configuration of these elements can be given
as (n-2) f0-14 (n-1)d0, 1, 2, ns2
These elements also called inner transition elements.
These elements are divided into two series corresponding to the filling of 4f,
5f, orbital’s of (n-2)th main shell.

4f- Series ( Lanthanides or Lanthanones) : This is a first inner transition series
and consists of fourteen elements. (58Ce, to 71Lu ). The last electron of these
elements are enter in 4f orbitals.
5f- Series ( Actinides or Actinones) : This is a second inner transition series and
consists of fourteen elements. (90Th, to 103Lr ). The last electron of these elements
are enter in 5f orbital’s.

Periodic Properties
 Covalent Radius
Defn- half the inter-nuclear distance between two identical atoms which
are linked by a covalent bond is called as Covalent radius.
Covalent
radius = d/2
Covalent radius.

Van der waal Radius or Collision radius

Van der waal Radius or Collision radius
Q. Explain the term Collision radius.
Ans- Definition- “The half of the distance between the nuclei of two non-
bonded neighboring atoms of two adjacent molecules is called as collision
radius.”
It is also called as Van der Waals radius.
For Example- In solid crystalline form of B.H.C. (C6H6Cl6), the molecules
are so arranged that the shortest distance between chloride nuclei in different
molecules is 3.60 .
Thus, in this molecules the van der Waals radius of chlorine is 3.60/2 = 1.80 .

Atomic and Ionic Radii
Q. Explain the term Atomic and ionic radius.
Ans- Atomic Radius- “The distance between the nucleus and outermost
shell of electrons of an atom ion is called as atomic radius.”
Ionic Radius- “The distance between the nucleus and outermost shell of
electrons of an ion is called as ionic radius.”

Periodic Variations of Atomic and ionic radius.
 In Period-
If we move from left to right in a period the values of atomic and
ionic radii decrease with increase in atomic number.
For example- Consider 2nd period it is observed that the atomic radii goes
decreasing from Li to F with increase in atomic number.
2nd period 3Li 4Be 5B 6C 7N 8O 9F 10Ne
Electronic
configuration
2, 1 2,2 2,3 2,4 2,5 2,6 2,7 2,8
Atomic radii in
pm
123 90 82 77 75 73 72 112

Metallic Radius-
Q. Explain the term Metallic radius.
Ans- Metallic Radius- “The half distance between the nuclei of two
adjacent metal atoms in the metallic close packed crystal lattice in which the
metal has coordination number of twelve is called as metallic radius.”
Metallic radii are generally smaller than Van der Waals radii because the
bonding in metallic crystal lattice is much stronger than Van der Waals forces.
Metallic radii are 10% to 15% larger than covalent radii.

 Ionization Potential Or
Ionization Energy-
Defn- “The amount of energy required to remove the outermost electron
from an isolated gaseous atom of an element in its ground state to form
cation is called as ionization energy or potential.”
It is represented as IE or IP.
This is an energy required process (endothermic process).

“The amount of energy required to remove the outermost or
valence electron from an isolated gaseous atom of an element in its
ground state to form cation is called as ionization energy or potential.”

The above process is specifically called as first ionization energy (I1) and
corresponds to removal of first electron.
Similarly 2nd, 3rd, 4th etc.(I2, I3, I4,….etc.) ionization energies corresponding
to the removal of 2nd, 3rd, 4th …etc. electron to form M2+, M3+, M4+ cation
respectively.

Periodic Variations of Ionization Energy.
 In a Period-
 If we move from left to right in a period the values of ionization energy
goes on increasing with increase in atomic number.
 We know that atomic size goes on decreasing from left to right in a period.
 i.e. distance between nucleus and outermost electron is becoming lesser
and lesser in a period.
 Now, as atomic number increases, the positive charge or nuclear charge
also increases i.e. electron is attached more strongly by the nucleus.
 Thus removal of electron required more amount of energy.

 Therefore, ionization energy goes on increasing fro left to right in a period.

Exceptions-
 Now, the exceptions to this general trends in 2nd period Be, N, Ne and in 3rd
period Mg, P and Ar.
 The high values of these elements can be explained on the ground that
either completely filled orbital’s or half filled orbital’s are extra stable and
therefore required more energy to remove an electron from these atoms.
 For example- electronic configuration as Be-1s2,2s2 completely filled 2s
sub shell and Ne-1s2,2s22p6 completely filled.
 In N-1s2,2s22p3 exactly half filled p sub shell . Removal of electron from
such atoms difficult and required more amount of energy.

 In a Group-
 In a group if we move from top to bottom ionization energy goes on
decreasing with increase in atomic number.
 This decrease in the values of ionization energy can be explained by
taking into consideration the following three factors mainly.
1. Increase in atomic size-
2. Screening effect (Shielding effect) -
3. Increase in atomic number-

1. Increase in atomic size- the atomic size increase moving from top to
bottom in a group, due to this lesser attraction force in between nucleus
and outermost electrons.
Thus removal of outermost electron more easy thereby required
lesser amount of energy and hence decrease in ionization energy.
2. Screening effect (Shielding effect) of the inner electron on the
outermost electron-
“The repulsion experienced by the valence electron due to presence of
inner shell electrons is called screening effect”.

• More the number of inner shells, stronger is the screening effect.
• Since in a group, every time there is the addition of one new shell, this
effect becomes more & more stronger and contributes so to the increased
repulsive force between nucleus and the outermost electron and hence
lesser attraction between them.

3. Increase in atomic number-
• We move downwards in a group positive charge (nuclear charge) goes on
increasing regularly due to increase in atomic number.
• Due to this, the outermost electron is attracted more and more strongly
towards the nucleus. Therefore, the effect of increase in atomic number
should cause increase in the values of ionization energy.
First two effects nullify the effect of increase in atomic number.
Hence , though nuclear charge increases while move downwards in a
group, it is not that much enough to counter balance the effect of increase
in the atomic size as well as increase in the screening effect.

First two effects nullify the effect of increase in atomic number.
Hence, though nuclear charge increases while move downwards in a
group, it is not that much enough to counter balance the effect of
increase in the atomic size as well as increase in the screening effect.

“Effect of IE on different Properties of
Element”
The Magnitude of ionization energy affects different properties of
elements like
 metallic and nonmetallic character,
 relative reactivity,
 oxidizing and reducing power and
 acidic and basic nature of oxides and hydrides etc.

1. Metallic and nonmetallic-
Metallic nature is nothing but ability of metal to lose one or more
electrons to forms cation. This ability depends upon the ionization
energy. That is with increase in ionization energy decrease the metallic
character of an element.
When we move from left to right in a period, ionization energy
increases which in turn decrease metallic character.
Thus alkali and alkaline earth metals are metallic in nature whereas
halogens are typical non metals.

2. Relative reactivity-
 The element with higher value of ionization potential are less reactive.
(for example- inert gases).
 While those have lower value are highly reactive. For example alkali and
alkaline earth metals.
 Ionization energy of elements in group decrease from top to bottom,
therefore, reactivity of elements increase.

3. Oxidizing and reducing power of elements-
 An element which reduces other elements by providing electrons and
itself undergoes oxidation such elements are called as reducing agent.
 The ability of an element reduces to other is called as reducing power of
an element.
 It is depends upon the ionization energy, lower the ionization energy
greater the ease of its oxidation and stronger the reducing agent.
 Therefore alkali metals have lower IP values are stronger reducing
agents.

4. Acidic and Basic nature of oxides and hydroxides-
 An element which loses an electron easily and gets converted into
positively charged ion (cation) is said to show basic character.
 Lower the values of IP easy removal of an electron.
 Therefore lower IP values elements are greater will be its basic
character.
 Thus metal oxides and hydroxides are basic in nature.
 Basic properties of elements increase with increase in atomic number in
a group.

Electron Affinity/Affinity Energy-
Defn- “The amount of energy released when an electron is added to an
isolated neutral gaseous atom in its ground state to produce an anion
affinity or affinity energy.”
It is represented as E or EA.
Its unit electron Volts (eV), Kilocalories (Kcal) per gram or Kilojule per
mole (KJ. Mole1-)
This is an exothermic process because energy released in it.

 Periodic Variations-
 Generally electron affinity decreases with increase in atomic radius and
increases with decrease in screening effect by the inner electrons.
 Apart from these factors, nature of orbital also affects the E values.
 If other factors are common, the electron affinity is largest for an
electron entering as s-orbital and decreases in the order, s>p>d<f
orbital.

In a Period-
If we moving from left to right in a period, electron affinity generally
goes on increasing with increase in atomic number.

In a Group-
If we moving from top to bottom in a group, electron affinity goes on
decreasing with increase in atomic number.

Q. Distinguish between Electronegativity and Electron affinity.

Electronegativity-
Defn- “The tendency or ability of an element to attract the electron pair or
shared pair towards itself is called as electronegativity of an element.”
Electronegativity of an atom is represented as cA.
The term "electronegativity" was introduced by Jöns Jacob
Berzelius in 1811.
Pauling first time proposed this term. He assigned an imaginary
value of 4 for fluorine, the most electronegative element among all
elements.

For example- HCl molecule which is formed by combination of two
dissimilar atoms, hydrogen and chlorine. The electron pair shared in between
H & Cl, this electron pair nearer to the Cl i.e. does not lie in center. The reason
for this unequal sharing of electron pair is that the Cl has a greater tendency to
attract electron pair than H-atom.

In a Period-
If we moving from left to right in a period, electronegativity increases
with increase in atomic number.
Periodic Variations

Explanation-
 Electronegativity increases from left to right is due to that the radius
decreases gradually.
 But since there is increase in the nuclear charge, the added electron is
held more strongly and this process goes on increasing from left to right
with increase in atomic number.
 Thus resulting in the increase in electronegativity values

In a Group-
If we move from top to bottom in a group, electronegativity decrease with
increase in atomic number.

Explanation-
 Electronegativity decrease when we proceed downwards in group.
 because increase in atomic radius due to which added electron experiences
lesser force of attraction towards nucleus.
Another factor, screening effect, which along the group due to increase in
the number on inner shells, is also responsible for decreased attraction
between nucleus and the added electron.
Therefore, small atoms attracts electrons more strongly than larger ones and
are thus these are more electronegative.

Determination of Electronegativity values
There are two different scales to determine the electronegativity
values of elements-
1) Pauling’s electronegativity scale or Pauling’s Bond
Energy Scale-
 This method takes into consideration bond energies i.e. the energy required
to break a bond to form neutral atoms.
Consider a bond A-B between two dissimilar atoms A and B of a
molecule AB. Let the bond energies of A-A, B-B and A-B bonds be denoted
as EA-A, EB-B and EA-B respectively.
The bond formed between two atoms A and B is generally
intermediate between pure covalent A-B and pure ionic A+B- .

Because of the partial ionic character the bond is strengthened or in
other words, the bond energy is increased. The bond in fact is stabilized by
resonance.
 In the determination of electronegativity scale, the following measurements
are made.
1. Actual or Experimental bond energy = H
2. Bond energy when the bond is truly covalent = Q
3.Resonance energy () due to ionic character of covalent bond called
Ionic Resonance energy () = H- Q
Since resonance energy is a measure of partial ionic character of a
covalent bond and the difference in electronegativity between the bonded
atoms is also related to the ionic character of the bond,

so  is also related to the difference in the electronegativity of the
bonded atoms. If XA and XB are the electronegativity of atoms A and B and
XA > XB , then
The factor 0.208 comes from the conversion of experimental value,
measured in kcal/mole into eV energy. Pauling fixed arbitrarily 4.0 as
electronegativity value for fluorine. From this, other electronegativity values
were calculated using the above equation (2).
………..(1)
………..(2)

2. Mullikan’s Scale-
According to Mullikan electronegativity is some type of mean of the
difference of ionization potential (IP) and electron affinity (EA) of an atom.
To approximate his values to those on Pauling’s electronegativity scale, he
defined the electronegativity of an atom A by the relation.
Determination of Electronegativity values
There are two different scales to determine the electronegativity
values of elements-
1) Pauling’s electronegativity scale or Pauling’s Bond
Energy Scale-

For example 1) Calculate the effective nuclear charge felt by the 3p electron
of Silicon.
Solution- Si atomic number is (z)=14
The electronic of Si= 1s2, 2s2, 2p6, 3s2, 3p2
we know, Zeff = Z - 
According to slater’s rules,
For Si,  = [0.35(no. of electron in the n group)+0.85(no. of electron in the (n-1)
group) + 1.00 (no. of electron for closer than (n-1) group)]
= [ (0.35x 3) + (0.85 x 8) + (1 x 2)]
therefore  = 9.85
Now, Zeff = Z -  = 14 - 9.85
= 4.15

2) Calculate the effective nuclear charge on 4s electron in Potassium. (K-
At. No. 19)
Solution- K atomic number is (z)=19
The electronic of K= 1s2, 2s2, 2p6, 3s2, 3p6 , 4s1
We know, Zeff = Z - 
For K,  = [0.35(no. of electron in the n group)+0.85(no. of electron in the (n-1)
= [ (0.35x 0) + (0.85 x 8) + (1 x 10)]= 6.8+10.00
therefore  = 16.8
Now, Zeff = Z -  = 19 - 16.8
= 2.20

3) Calculate the effective nuclear charge on 3s electron in Sodium. (K-At.
No. 11)
Solution- Na atomic number is (z)=11
The electronic of Na = 1s2, 2s2, 2p6, 3s1
We know, Zeff = Z - 
For Na,  = [0.35(no. of electron in the n group)+0.85(no. of electron in the (n-1)
= [ (0.35x 0) + (0.85 x 8) + (1 x 2)] = 6.8+2.00
therefore  = 8.8
Now, Zeff = Z -  = 11 - 8.8
= 2.2

Effect of Electronegativity on the properties of elements
The Magnitude of electronegativity affects different properties of
elements like
 Metallic and nonmetallic character,
 Relative Reactivity,
 Oxidizing and Reducing Property

1. Metallic and nonmetallic-
 Metallic nature is nothing but ability of metal to lose one or more
electrons to forms cation.
 Electronegativity value increase with the metallic character decrease of
an element.
 Hence metals have low electronegativity values while nonmetals have
high eletronegativity values because they have greater tendency to gain
electrons.

2. Relative Reactivity-
Relative reactivity of a nonmetallic element in general increases with
increase in the value of electronegativity. E.g. among halogens , fluorine the
most electronegative element (4) is also the most reactive among them.
Relative reactivity order of halogens can be given as
F (4.0) > Cl (3.0) > Br (2.8) > I (2.1).

3. Oxidizing and Reducing Property
More the electronegativity value, higher the tendency to gain
electrons and consequently greater the oxidizing power of an element, hence
elements having higher values of electronegativity acts as stronger oxidizing
agents. E. g. nonmetals are powerful oxidizing agents. On the other hand, the
metals have low values of electronegativity. Metals therefore acts as powerful
reducing agents. Decrease in the electronegativity value increases reducing
power of elements.

Electronegativity and Partial ionic character of a
covalent bond
In general, larger electronegativity difference between the atoms A
and B in a molecule AB, more will be its ionic nature. In other words,
“greater the value of XA- XB, greater the % of ionic character in a bond”. This
relationship between % ionic character and electronegativity difference can
be shown graphically as-

From the graph it follows that,
1. XA–XB = 1.7, then the bond A-B is 50% ionic and 50% covalent.
2. XA–XB > 1.7, then the bond A-B is more than 50% ionic and such
compounds are considered to be ionic.
3. XA–XB < 1.7, then the bond A-B is more than 50% covalent and less
than 50% ionic. Such compounds are called as covalent compounds.

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