Periodic classification of elements -One Shot.pdf

Content-
● Why do we need to classify elements?
● Genesis of periodic elements
● (triads, law of octaves, Mendeleev)
● Modern periodic law and the present
form of the periodic table
● Nomenclature of elements with
atomic numbers > 100
● Electronic configuration of elements
and the periodic table
● Electronic configurations and types of
elements- s,p, d and f blocks
● The S block elements
● The p block elements
● The d block elements
● The f block elements
● Metals, non metals, and metalloids
● Periodic trends in properties of
elements
● Trends in physical properties
● Ionization enthalpy
● Electron gain enthalpy
● Electronegativity
● periodic trends in chemical properties
● Oxidation states
● Anomalous properties of second
period elements
● Periodic trends and chemical
reactivity

Need for classification of elements
● 1800 → 31 elements
● 1865 → 63 elements
● 2016 → 118 elements
● It is difficult to individually study the chemistry of each element and its
compounds.
● As per IUPAC there are currently; 2 million inorganic compounds & 18
million Organic compounds.
● An arrangement in which elements having similar properties were
classified in tabular form which lead to the formation of the periodic table.
● Formulation of Periodic table provided a way for the study of chemistry in
a much simpler way.

Dobereiner’s Triads
Grouped the elements with similar chemical properties into clusters of three called 'Triads'.
When 3 elements having similar properties are arranged in an increasing order of their
atomic mass, the atomic mass of middle element is mean of atomic mass of the other two
elements

Newland’s Law of Octaves
When elements are arranged in the increasing order of their atomic
mass, the properties of every eighth element resembles that of the first.

Demerits of Newland’s Law of Octaves
● Worked only for the lighter elements (only till Ca).
● Discovery of noble gases failed this law.
● Not valid for heavier elements as O and Fe were placed in the
same group despite having vastly different properties.

Lothar meyer’s curve
● Alkali metals are at the peak of the curve.
● Alkaline earth metals are at the descending part of the curve.
● Halogens are at the ascending part of the curve.
● Most of the transition metals are lying in the broad minimas.

Lothar meyer’s curve
Many other physical properties of the elements also change in a
periodic manner like melting point, boiling point, density,
thermal expansion coefficient, etc.
Based on these observations, he proposed that-
“The physical properties of elements are periodic functions of
their atomic weights”
This formed the basis of Mendeleev’s periodic table

Mendeleev's Periodic Table
Physical and chemical properties of an element are periodic
functions of their atomic masses.

Merits of Mendeleev's Periodic Table
• Classification of elements was based on the properties of elements.

● If the elements are arranged in the order of their increasing atomic
weights, after a regular interval elements with similar properties are
repeated. The table is divided into nine vertical columns and seven
horizontal rows.
● He placed elements having similar nature in the same vertical column or
group with respect to atomic weight.
● At some places, the order of atomic weight was changed in order to
justify the physical and chemical nature of elements.
● Some places were left vacant for new elements which were not
discovered at that time.

● Left gaps in his periodic table, for elements yet to be discovered.

● Determination & Correction of doubtful Atomic weights
Beryllium
Atomic
Weight
Valency
Earlier 13.5 3
Corrected 9.01 2

Demerits of Mendeleev's Periodic Table
• Position of Hydrogen element was not justified.
• No separate positions were given to isotopes.
• Order of increasing atomic weights is not strictly followed in the
arrangement of elements in the periodic table.
For e.g. – Co (At. wt. 58.9) is placed before I (127) and Ar (39.9)
before K (39).
• There was no explanation for the periodicity of elements.
• No place for Noble or Inert gases.

Henry Moseley observed regularities in the characteristic X-ray spectra of the
elements. A plot of ( where ν is frequency of X-rays emitted) against atomic
number (Z ) gave a straight line and not the plot of vs atomic mass.
Modern periodic Table

Modern periodic Table
Henry Moseley stated that
“the properties of elements are the periodic function of
their atomic numbers.”

Key points about Modern periodic Table
❏ The classification of elements is based on the atomic number, which is a
more fundamental property.
❏ The reason for placing isotopes at one place is justified as the
classification is on the basis of atomic number.
❏ It explains the periodicity of the properties of the elements and relates
them to their electronic configurations.
❏ The lanthanides and actinides are placed separately at the bottom of the
periodic table.

Features of modern periodic table
➢ Vertical columns are known as groups
➢ Horizontal rows are known as periods.
➢ Total number of periods = 7
➢ Total number of groups = 18
➢ Elements are placed in periods based on the number of shells in
their atoms.

➢ The left-most group is known as alkali metals group
➢ Second group is also known as alkaline earth metals group
➢ Groups 3 to 12 are transition elements.
➢ Group 17 is known as halogen family.
➢ Group 18 is known as noble gas group, and also called as zero
group.

Period (n) Orbit filled up Number of
elements in the
period
Atomic number Remark
From To
First (1) 1s 2 H(1) He(2) Very short period
Second (2) 2s2p 2+6=8 Li(3) Ne(10) Short period
Third (3) 3s3p 2+6=8 Na(11) Ar(18) Short period
Fourth (4) 4s3d4p 2+10+6+18 K(19) Kr(36) Long period
Fifth (5) 5s4d5p 2+10+6=18 Rb(37) Xe(54) Long period
Sixth (6) 6s4f5d6p 2+14+10+6=32 Cs(55) Rn(86) Very long period
Seventh (7) 7s5f6d7p 2+14+10+6=32 Fr(87) Og(118) Very long period

Nomenclature of elements with atomic numbers
>100
● The digits of atomic number of an element are expressed serially using numerical roots.
● Then successive roots are put together and the name is ended with ium.
Digit Word root Digit Word root
0 nil 5 pent
1 un 6 hex
2 bi 7 sept
3 tri 8 oct
4 quad 9 enn

The atomic number of unnilunium is—-------------
[JEE2020]

Screening effect /Shielding effect
➢ In a multi electronic system, the electrons are attracted
towards the Nucleus as well as repelled by each other. Thus the
force attraction between the Nucleus and the valence shell
electrons is somewhat decreased.

Effective nuclear Charge
Z* = Z – σ
σ = Screening constant / Shielding constant
Z* = effective nuclear charge
Effective Nuclear Charge :-
The force of attraction on valence shell electrons

Slater’s Rule
If electron resides in ns or np subshell
❏ All electrons present in shells higher than concerned
electron contribute zero to σ.
❏ All electrons present in same shell contribute 0.35 toσ .
❏ All electrons present in (n-1) shell contribute 0.85 to σ .
❏ All electrons present in deeper shells contribute 1 toσ .

Calculate σ and Zeff for 4s electron in Mn(Z =25) ?

σ for 4s1 electron in Mn atom = 0.35×[No. Of the remaining electrons in 4s orbital] +
0.85 [No. of electrons in 3s, 3p and 3d orbitals] +
1.0×[No. of electrons in the inner shells]
= 0.35 x 1 +0.85×13+1.0×10 =21.40.
∴ Zeff experienced by 4s1 electron = 25 -21.40 = 3.60
Solution:

Factors affecting shielding effect
❏ No. of inner shell electrons
More the no. of inner shell electrons, more will be the value of σ.
❏ Shape of orbital
Order of screening effect
(max screening) ns > np > nd > nf (min screening)

Atomic radii
The average distance of valence shell e– from nucleus is called atomic
radius. It is very difficult to measure the atomic radius because –
❏ The isolation of single atom is very difficult.
❏ There is no well defined boundary for the atom. (The probability of
finding the e– is 0 only at infinity).
❏ So, the more accurate definition of atomic radius is –
Half the inter-nuclear distance(d) between two atoms in a diatomic
molecule is known as atomic radius.

Covalent radius Metallic radius
Vanderwaal
radius
Atomic Radius

Covalent radius
It is defined as half of the inter-nuclear distance between two nuclei of
atoms bonded by a single covalent bond.
For homonuclear molecule
r =
dA–A
2

Covalent radius
For hetronuclear molecule

Given that covalent radius of F(EN=4) and Cl(EN=3) are 71 pm and 99 pm, find the
bond length of F-Cl molecule.
A
B
C
D
170 pm
71 pm
165 pm
161 pm

Solution:
In picometer units,
Bond length = RA + RB - 9 |XA - XB| = 71 + 99 - 9(4-3) = 161 pm

Metallic radius
It is defined as half of the internuclear distance between the nuclei of two
adjacent atoms in a metallic crystal.
r =
dM–M
2

Van Der Waals radius
One half of the distance between the nuclei of two adjacent
atoms belonging to two neighbouring molecules of a compound
in the solid state is called Van der Waals radius.

Comparison of Atomic radius
❏ For a species:
Van der Waals Radius > Metallic Radius > Covalent Radius
❏ Van der waals radius is largest .
❏ Covalent radius is smallest .

Number of shells :
Atomic radius increases, with increase in no. of shells
No. of shells ∝ Atomic radius
Factor affecting Atomic radius

Effective Nuclear charge :
❏ With addition of protons
→ Nuclear charge increases
❏ With increase in nuclear charge
→ Atomic size decreases
Nuclear Charge ∝
1
Atomic Radius

Shielding Effect :
Inner shell electrons are “shielding” the outer shell electron
The attractive force of the nucleus does not
reach the outermost electron. As a result of the
shielding effect, the atomic radii increases
Shielding Effect ∝ Atomic radius

Multiplicity of bond
Covalent radii decrease as the multiplicity of bond increases
Bond CーC C=C C☰C
dc-c(Å) 1.54 1.34 1.20
rc(Å) 0.77 0.67 0.60

General trend in periodic Table
for atomic Radius

Across a Period
Atom (Period II) Li Be B C N O F
Atomic radius 152 111 88 77 74 66 64
Atomic size decrease
Zeff increase

Down the Group
Atom
(Group
I)
Atomic radius
(pm)
Atom
(Group 17)
Atomic
radius
(pm)
Li 152 F 72
Na 186 Cl 99
K 231 Br 114
Rb 244 I 133
Cs 262 At 140
No of shell
increase
Atomic
radius
increase

In group 13 the trend of atomic radius is
B < Al ≈ Ga < In < Tl
This is due to poor Shielding of 3d electrons in
Ga as a result effective nuclear charge
increases hence size of Ga decreases.
Special cases :

Element Sc Ti V Cr Mn Fe Co Ni Cu Zn
Atomic radius (A) 1.44 1.32 1.22 1.18 1.17 1.17 1.16 1.15 1.17 1.25
Nearly same
Due to increase in nuclear
charge
Screening effect > nuclear charge
Screening effect = nuclear charge
Increases
Decreases
● In the transition series (e.g. in first transition series), the covalent
radii of the elements decrease from left to right across a row until
near the end when the size increases slightly
Special cases :

Which has maximum atomic radius -
A
B
C
D
Al
Si
P
Mg

Solution:
Mg, as we move across the
period atomic radius
decreases.

● Ions are formed by either gain of electrons by an atom or loss of
electrons by an atom
● On the basis of this Ions can be classified in two types
Ionic radius

Radius of an anion is invariably bigger than that of the corresponding atom.
Reason:
❏ The effective nuclear charge decrease in the formation of anion. Thus
the electrostatic force of attraction between the nucleus and the
outer electrons decreases as the size of the anion increases.
❏ Interelectronic repulsion increases.
Anionic radius

Radius of cation is smaller than that of corresponding atom.
Reason:
During the formation of cation ,effective nuclear charge increase as a
result atomic size decreases.
Cationic radius

Ionic radii of
A
B
C
D
Ti4+ < Mn7+
35Cl- < 37Cl-
K+ > Cl-
P3+ > P5+

The species containing the same number of electrons but
differ in the magnitude of their nuclear charges.
Isoelectronic Species

The correct order of the ionic radii of O2-, N3-, F-, Mg2+, Na+ and Al3+ is:
A
B
C
D
N3- < O2- < F- < Na+ < Mg2+ < Al3+
Al3+ < Na+ < Mg2+ < O2- < F- < N3-
Al3+ < Mg2+ < Na+ < F- < O2- < N3-
N3- < F- < O2- < Mg2+ < Na+ < Al3+

Solution:
All are isoelectronic species, so more is the Zeff less will be the ionic size.
∴ Correct order of ionic radii is
Al3+ < Mg2+ < Na+ < F- < O2- < N3-

Chloride ion and potassium ion are isoelectronic. Then
A
B
C
D
Potassium ion is relatively
bigger.
Depends on the other
cation and anion
Their size are same
Chloride ion is bigger
than potassium ion.
HOMEWORK

Ionisation energy (IE) - Amount of energy required to
remove the most loosely bound electron from an isolated
gaseous atom to form a cation.
Ionisation Energy

IE1
IE2
IE3
IE1 < IE2 < IE3
Ionisation Energy

Atomic size :
Greater the size
of atom →
Easier is the removal of the
valence electron, hence energy
required is less.
Factor affecting Ionisation Energy

Nuclear charge :
Greater the
charge →
Greater is attraction between nucleus
and electrons present in outermost
shell, more will be difficult to remove
electron.

Screening Effect :
With increase in Screening Effect, nuclear charge on electrons
in the outermost shell reduces. Since effective nuclear charge
decreases, Ionization Enthalpy decreases.

Electronic configuration :

● Moving from left to right in a period I.E. ↑ as (Zeff ↑)
● Moving from top to bottom in a group I.E. ↓ as (Size↑)
● Special Cases
● For transition elements
Generally in period IE increases, but the increase is not so
regular (Sc, Ti, V, Cr) differ only slightly from each other and
Fe, Co, Ni, Cu values are fairly close to each other from Cu - Zn
again increasing
Ionisation Energy

Electron Gain enthalpy
Electron affinity : Energy released when an e- is
added to the valence shell of an isolated
gaseous atom.

Effective Nuclear Charge(Z ⃰ )
Higher the effective nuclear charge, higher is the electron
affinity
EA ∝ Z*
Factor affecting electron Gain
enthalpy

Atomic radius
Higher the atomic radius, lesser is the effective
nuclear charge and lesser is the electron affinity
∝
1
Atomic radius
EA
Factor affecting electron Gain enthalpy

Screening effect
Electron gain enthalpy value of the elements decreases with
the increasing shielding or screening effect. The shielding
effect between the outer electrons and the nucleus increases
as the number of electrons increases in the inner shells.

Half filled and fully filled electronic configuration :
Atoms having half filled and fully filled configuration is
most stable and they do not show tendency to accept
the extra electron.

Across a period
● Across a period, with increase in atomic number, ΔegH becomes more ‘-ve’
because across a period Zeff increases and consequently it will be easier to
add an e- to a small atom.

Down the group
● Within a group from top to bottom, ΔegH becomes less negative because
the size of the atom increases and the added e- would be at larger
distance from the nucleus.
Group 16 ΔegH Group 17 ΔegH Group 18 ΔegH
He +48
O -141 F -328 Ne +116
S -200 Cl -349 Ar +96
Se -195 Br -325 Kr +96
Te -190 I -295 Xe +77
Po -174 At -270 Rn +68
Atomic
radii
increase
ΔegH Decrease

● Ability of an atom in a chemical compound to attract
shared electrons to itself
● Defined in bonded state
Electronegativity

Mulliken Scale
Mulliken Scale: Mulliken gave the electronegativity as the
average value of ionisation potential and electron gain
enthalpy of an atom.

Pauling Scale
● According to Pauling, electronegativity difference (XA – XB)
in between two atoms A & B is given by :

Factor affecting electronegativity
(A) Atomic size Electronegativity ∝
(B) Effective nuclear charge (Zeff)
Electronegativity ∝ Zeff
(C) Hybridisation state of an atom
Electronegativity ∝ % s character in
hybridised atom
sp >sp2 >sp3
s character 50% 33% 25%
Electronegativity 3.25 2.75 2.5
(D) Oxidation state
The electronegativity value increases as
the oxidation state (i.e. the number of
positive charge) of the atom increases.

Electronegativity increases across the period as atomic size decreases.
Trend of electronegativity

Electronegativity decreases down the group as atomic size increases.
Trend of electronegativity

Metallic and non metallic nature :
a) On moving from left to right in a periods, the electronegativity of the
elements increases. So the metallic character decreases.
b) On moving down a group, the electronegativity of the elements
decreases. So the metallic character increases.
Application of electronegativity

Bond strength:
If the electronegativity difference of covalently bonded atoms Δ(x)
increases, the bond energy of the covalent bond also increases.
For Example :
the order of the H-X bond strength is –
H – F > H – Cl > H – Br > H – I

Partial ionic character in covalent Bond:

Acidic and basic nature of oxides

Periodic classification of elements -One Shot.pdf

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