chemistry

1. Electronic Configuration
Definition: The systematic distribution of electrons in the various atomic orbitals is called its electronic
configuration. In the electronic configuration, electrons are distributed in various energy levels (various shells) of
an atom, such as K shell, L shell, M shell, N shell, etc.
The atomic number of hydrogen is 1. Therefore, a hydrogen atom
contains 1 electron, which will be placed in the s subshell of the
first shell/orbit. The electron configuration of hydrogen is 1s1, as
illustrated below.

2. • Electron Configuration of Oxygen
• The atomic number of oxygen is 8, implying that
an oxygen atom holds 8 electrons. Its electrons
are filled in the following order:
• K shell – 2 electrons
• L shell – 6 electrons
• Therefore, the electron configuration of oxygen
is 1s2 2s2 2p4, as shown in the illustration
provided below.

3. • Chlorine Electronic Configuration
• Chlorine has an atomic number of 17.
Therefore, its 17 electrons are
distributed in the following manner:
• K shell – 2 electrons
• L shell – 8 electrons
• M shell – 7 electrons
• The electron configuration of chlorine
is illustrated below. It can be written
as 1s22s22p63s23p5 or as [Ne]3s23p5

4. C Br He Tc
hydrogen
1
H
1.00794
solid liquid gas synthetic helium
2
He
4.002602
lithium
3
Li
6.941
beryllium
4
Be
9.012182
boron
5
B
10.811
carbon
6
C
12.0107
nitrogen
7
N
14.00674
oxygen
8
O
15.9994
fluorine
9
F
18.9984
neon
10
Ne
20.1797
sodium
11
Na
22.98977
magnesium
12
Mg
24.3050
aluminium
13
Al
26.981538
silicon
14
Si
28.0855
phosphorus
15
P
30.97376
sulphur
16
S
32.065
chlorine
17
Cl
35.453
argon
18
Ar
39.984
potassium
19
K
39.0983
calcium
20
Ca
40.078
scandium
21
Sc
44.95591
titanium
22
Ti
47.867
vanadium
23
V
50.9415
chromium
24
Cr
51.9961
manganese
25
Mn
54.93805
iron
26
Fe
55.845
cobalt
27
Co
58.9332
nickel
28
Ni
58.6934
copper
29
Cu
63.546
zinc
30
Zn
65.409
gallium
31
Ga
69.723
germanium
32
Ge
72.64
arsenic
33
As
74.9216
selenium
34
Se
78.96
bromine
35
Br
79.904
krypton
36
Kr
83.798
rubidium
37
Rb
85.4678
strontium
38
Sr
87.62
yttrium
39
Y
88.90585
zirconium
40
Zr
91.225
niobium
41
Nb
92.90638
molybdenum
42
Mo
95.94
technetium
43
Tc
[98]
ruthenium
44
Ru
101.07
rhodium
45
Rh
102.9055
palladium
46
Pd
106.42
silver
47
Ag
107.8682
cadmium
48
Cd
112.411
indium
49
In
114.818
tin
50
Sn
118.710
antimony
51
Sb
121.760
tellurium
52
Te
127.60
iodine
53
I
126.9045
xenon
54
Xe
131.293
caesium
55
Cs
132.90545
barium
56
Ba
137.327
lutetium
71
Lu
174.967
hafnium
72
Hf
178.49
tantalum
73
Ta
180.9479
tungsten
74
W
183.84
rhenium
75
Re
186.207
osmium
76
Os
190.23
iridium
77
Ir
192.217
platinum
78
Pt
195.078
gold
79
Au
196.96655
mercury
80
Hg
200.59
thallium
81
Tl
204.3833
lead
82
Pb
207.2
bismuth
83
Bi
208.980
polonium
84
Po
[209]
astatine
85
At
[210]
radon
86
Rn
[222]
francium
87
Fr
[223]
radium
88
Ra
[226]
lawrencium
103
Lr
[262]
rutherfordiu
m
104
Rf
[261]
dubnium
105
Db
[262]
seaborgium
106
Sg
[266]
bohrium
107
Bh
[264]
hassium
108
Hs
[269]
meitnerium
109
Mt
[268]
darmstadtium
110
Ds
[271]
roentgenium
111
Rg
[272]
Coperniciu
m112
Cn
[285]
nihonium
113
Nh
[289]
Flerovium
114
Fl
[289]
moscovium
115
Mc
[289]
Livermoriu
m116
Lv
[292]
tennessine
117
Ts
[293]
oganesson
118
Og
[294]
lanthanum
57
La
138.9055
cerium
58
Ce
140.116
praseodymium
59
Pr
140.90765
neodymium
60
Nd
144.24
promethium
61
Pm
[145]
samarium
62
Sm
150.36
europium
63
Eu
151.964
gadolinium
64
Gd
157.25
terbium
65
Tb
158.9253
dysprosium
66
Dy
162.50
holmium
67
Ho
164.930
erbium
68
Er
167.259
thulium
69
Tm
168.934
ytterbium
70
Yb
173.04
actinium
89
Ac
[227]
thorium
90
Th
232.038
protactinium
91
Pa
231.0359
uranium
92
U
238.0289
neptunium
93
Np
[237]
plutonium
94
Pu
[244]
americium
95
Am
[243]
curium
96
Cm
[247]
berkelium
97
Bk
[247]
californium
98
Cf
[251]
einsteinium
99
Es
[252]
fermium
100
Fm
[257]
mendelevium
101
Md
[258]
nobelium
102
No
[259]

5. • On the basis of electronic configuration, elements are classified into four
Blocks known as s, p, d and f- blocks.
• 1st and 2nd group elements are called s-block elements. The general electronic
configuration is ns1-2.
• 13th to 18th group elements are called p-block elements. The general electronic
configuration is ns2 np1-6.
• 3rd to 12th group elements are called d-block elements. The general electronic
configuration is (n-1)d1-10 ns1-2.
• Lanthanides and actinides elements are called f-block elements. The general
electronic configuration is (n-2)f1-14 (n-1)d0-1 ns2.

7. Classification of Modern Periodic Table
Based on nature and characteristics, elements are classified into three main types:
• Metals – These comprises 78% of all known elements and appear on the left-hand side of the
periodic table. All the s, d and f block elements are metals. Usually, metals are solid at room
temperature with high melting and boiling points. Except for mercury because it is liquid at room
temperature. They are malleable and good conductors of heat and electricity. They are
electropositive as they form bonds by losing electrons. In general cases, oxides of metals are basic
in nature.
• Nonmetals - These elements lie on the top right-hand side of the periodic table. The number of non-
metal is very few. Non-metals are either solids or gases at room temperature except bromine which
is liquid with low boiling and melting point. They are bad conductors of heat and electricity, and they
are brittle solids. They are electronegative as they form bonds by gaining electrons. In general
cases, oxides of non-metals are acidic in nature.
• Metalloids - These have specific characteristics common to both metals and non-metals. These are
also known as semi-metals. The elements which show the properties of both metals and nonmetals
are called metalloids or semimetals. For example Boron, silicon, germanium, arsenic, antimony,
tellurium, and polonium

8. Periodic properties and their trends
The periodic properties may be defined as:
The properties of the elements are directly or indirectly related to the electronic configuration of their atoms
and show gradation (increases or decreases) in moving down a group or a longer period.
The common physical properties of the elements are melting points, boiling points, density, enthalpy of fusion and vaporization etc.
But we shall focus our attention mainly on the properties which are based on electronic configuration these are:
•Atomic and ionic radii
•Ionization enthalpy
•Electrons gain enthalpy
•Electronegativity

9. 1. Atomic and ionic Radii
Atomic Radii:
The atomic radius may be defined as the distance from the centre of the nucleus to the
outermost shell containing electrons. Depending upon the nature of bonding in the atoms these
are (i) Covalent radii (ii) van der Waals radii (iii) Metallic radii
(i) Covalent radii: One-half of the distance between the centres of the nuclei of two adjacent
similar atoms joined to each other by a single covalent bond is known as covalent radii. Eg Cl-
Cl bond distance=198 pm covalent radius of Cl= 99 pm.
(ii)van der Waals radii: Half of the internuclear distance between two similar adjacent atoms
belonging to the two neighbouring molecules of the same substance in the solid state is known
as van der Waals radii.
(iii) Metallic radii: Half the distance between the centre of the nuclei of two adjacent atoms in
the metallic crystal is known as metallic radii
As we move from left to right in a period, the atomic radius decreases due to an increase
in effective nuclear charge (Zeff). Along the group, as we move from top to bottom, atomic
radius increases due to increase in principal quantum number which causes an increase in the
number of shells and increases in shielding effect.

11. • Ionic Radii:
• The ions formed by the loss of one or more electrons from the neutral atom
are known as cation (positive ion) when the electrons added to the neutral
atom form an anion (negative ion). The effective distance from the centre of
the nucleus of the ion upto which it exerts its influence on the electron cloud is
known as the ionic radii.
• The ionic radii change in the same trend as atomic radii. It decreases along
the period from left to right and increases down the group from top to bottom.
size of cation and anion of any natural atom as: cation< neutral atom < anion

12. 2. Ionization enthalpy
The amount of energy required when an electron is removed from the outermost orbit of an isolated
gaseous atom is known as Ionisation Enthalpy (IE).
Generally left to right in period IE increases whereas on moving down the group it decreases but half-
filled orbital and fully filled orbitals are highly stable and thus have high IE.
3. Electrons gain enthalpy
The electron gain enthalpy is defined as the change in enthalpy which takes place when a gaseous
atom gains an extra electron to form a monovalent anion in the gaseous state.
Electron gain enthalpy increases across the periods while it decreases down the group.
Chlorine has the highest electron affinity than fluorine.
4. Electronegativity
Electronegativity is the tendency of an atom to attract the shared pair of electrons towards itself in a
covalent bond. Fluorine is the most electronegative element while Cesium is the least.
In the periods left to right electronegativity increases. In the groups while moving down the groups
electronegativity decreases.

13. • Periodic Trends in Chemical Properties along a Period
(i) Metallic character: Decrease across a period maximum on the extreme left (alkali
metals).
(ii) Non-metallic character: Increases along a period. (From left to right).
(iii) Basic nature of oxides: Decreases from left to right in a period.
(iv) Acidic nature of oxides: Increases from left to right in a period
• • Variation from Top to Bottom on Moving Down a Group
(i) Metallic character. Generally increases because increase in atomic size and hence
decrease in the ionization energy of the elements in a group from top to bottom.
(ii) Non-metallic character. Generally decreases down a group. As electronegativity of
elements decreases from top to bottom in a group.
(iii) Basic nature of oxides. Since metallic character or electropositivity of elements
increases in going from top to bottom in a group basic nature of oxides naturally increases.
(iv) Acidic character of oxides. Generally decreases as non-metallic character of elements
decreases in going from top to bottom in a group.
(v) Reactivity of metals. Generally increases down a group. Since tendency to lose
electron increases.
(vi) Reactivity of non-metals. Generally decreases down the group, Higher the electro-
negativity of non-metals, greater is their reactivity. Since electronegativity of non-metals in
a group decreases from top to bottom, their reactivity also decreases