Classification of Elements and Periodicity in Properties: Overview and Evolution
C
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N
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w
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:
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N
e
w
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O
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M
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n
d
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v
’
s
p
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r
i
o
d
i
c
l
a
w
“The
properties of
the
elements
are
a
periodic
function
of
their
atomic
weights.”
Mendeleev
arranged
elements
in
horizontal rows
and
vertical
columns
of
a
table
in
order
of
their
increasing
atomic
weights
in
such
a
way
that
the
elements
with
similar
properties
occupied
the
same
vertical
column
or
group
.
M
e
n
d
e
l
e
e
v
l
e
f
t
g
a
p
s
u
n
d
e
r
s
i
l
i
c
o
n
a
n
d
a
l
u
m
i
n
i
u
m
a
n
d
t
h
o
s
e
w
e
r
e
k
n
o
w
n
a
s
e
k
a
a
l
u
m
i
n
i
u
m
a
n
d
e
k
a
s
i
l
i
c
o
n
.
L
a
t
e
r
i
t
w
a
s
d
i
s
c
o
v
e
r
e
d
t
h
a
t
e
k
a
a
l
u
m
i
n
i
u
m
h
a
d
p
r
o
p
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r
t
i
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s
s
i
m
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l
a
r
t
o
g
a
l
l
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u
m
a
n
d
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k
a
s
i
l
i
c
o
n
h
a
d
p
r
o
p
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r
t
i
e
s
s
i
m
i
l
a
r
t
o
g
e
r
m
a
n
i
u
m
.
A
c
c
o
r
d
i
n
g
t
o
m
o
d
e
r
n
p
e
r
i
o
d
i
c
l
a
w
,
“The
physical
and
chemical
properties
of
the
elements
are periodic
functions
of
their
atomic
numbers”.
Modern
form
of
periodic
table
was
discovered
by
henry
Mosley
and
hence
it
was
also
called
henry Mosley's
periodic
table.
T
h
e
P
e
r
i
o
d
i
c
L
a
w
r
e
v
e
a
l
e
d
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p
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s
a
m
o
n
g
t
h
e
9
4
n
a
t
u
r
a
l
l
y
o
c
c
u
r
r
i
n
g
e
l
e
m
e
n
t
s
(
n
e
p
t
u
n
i
u
m
a
n
d
p
l
u
t
o
n
i
u
m
l
i
k
e
a
c
t
i
n
i
u
m
a
n
d
p
r
o
t
o
a
c
t
i
n
i
u
m
a
r
e
a
l
s
o
f
o
u
n
d
i
n
p
i
t
c
h
b
l
e
n
d
e
–
a
n
o
r
e
o
f
u
r
a
n
i
u
m
)
.
T
h
e
h
o
r
i
z
o
n
t
a
l
r
o
w
s
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f
t
h
e
p
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d
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a
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w
h
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t
h
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v
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r
t
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a
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c
o
l
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m
n
s
a
r
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c
a
l
l
e
d
g
r
o
u
p
s
o
r
f
a
m
i
l
i
e
s
.
D
I
V
I
S
I
O
N
O
F
P
E
R
I
O
D
I
C
T
A
B
L
E
T
h
e
p
e
r
i
o
d
i
c
t
a
b
l
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i
s
d
i
v
i
d
e
d
i
n
t
o
4
b
l
o
c
k
s
m
a
i
n
l
y
t
h
e
s
-
b
l
o
c
k
p
-
b
l
o
c
k
d
-
b
l
o
c
k
a
n
d
f
-
b
l
o
c
k
.
S
-
b
l
o
c
k
:
-
T
h
e
e
l
e
m
e
n
t
s
o
f
G
r
o
u
p
1
(
a
l
k
a
l
i
m
e
t
a
l
s
)
a
n
d
G
r
o
u
p
2
(
a
l
k
a
l
i
n
e
e
a
r
t
h
m
e
t
a
l
s
)
w
h
i
c
h
h
a
v
e
n
s
1
a
n
d
n
s
2
o
u
t
e
r
m
o
s
t
e
l
e
c
t
r
o
n
i
c
c
o
n
f
i
g
u
r
a
t
i
o
n
b
e
l
o
n
g
t
o
t
h
e
s
-
B
l
o
c
k
E
l
e
m
e
n
t
s
.
They
are
all
reactive
metals
with
low
ionization
enthalpies.
They
lose
the
outermost
electron(s)
readily
to
form
1+
ion
in
the
case
of
alkali
metals
or
2+
ion
in
the
case
of
alkaline
earth
metals.
The
metallic
character
and
the
reactivity
increase
as
we
go
down
the
group.
Because
of
high
reactivity
they
are
never
found
pure
in
nature.
The
compounds
of
the
s-
block
elements,
with
the
exception
of
those
of
lithium
and
beryllium
are
predominantly
ionic.
P
-
b
l
o
c
k
:
T
h
e
p
-
B
l
o
c
k
E
l
e
m
e
n
t
s
c
o
m
p
r
i
s
e
t
h
o
s
e
b
e
l
o
n
g
i
n
g
t
o
G
r
o
u
p
1
3
t
o
1
8
a
n
d
t
h
e
s
e
t
o
g
e
t
h
e
r
w
i
t
h
t
h
e
s
-
B
l
o
c
k
E
l
e
m
e
n
t
s
a
r
e
c
a
l
l
e
d
t
h
e
R
e
p
r
e
s
e
n
t
a
t
i
v
e
E
l
e
m
e
n
t
s
o
r
M
a
i
n
G
r
o
u
p
E
l
e
m
e
n
t
s
.
The
outermost
electronic
configuration
varies
from
ns
2
np
1
to
ns
2
np
6
in
each
period.
At
the
end
of
each
period
is
a
noble
gas
element
with
a
closed
valence
shell
ns
2
np
6
configuration.
A
l
l
t
h
e
o
r
b
i
t
a
l
s
i
n
t
h
e
v
a
l
e
n
c
e
s
h
e
l
l
o
f
t
h
e
n
o
b
l
e
g
a
s
e
s
a
r
e
c
o
m
p
l
e
t
e
l
y
f
i
l
l
e
d
b
y
e
l
e
c
t
r
o
n
s
a
n
d
i
t
i
s
v
e
r
y
d
i
f
f
i
c
u
l
t
t
o
a
l
t
e
r
t
h
i
s
s
t
a
b
l
e
a
r
r
a
n
g
e
m
e
n
t
b
y
t
h
e
a
d
d
i
t
i
o
n
o
r
r
e
m
o
v
a
l
o
f
e
l
e
c
t
r
o
n
s
.
T
h
e
n
o
b
l
e
g
a
s
e
s
t
h
u
s
e
x
h
i
b
i
t
v
e
r
y
l
o
w
c
h
e
m
i
c
a
l
r
e
a
c
t
i
v
i
t
y
.
P
r
e
c
e
d
i
n
g
t
h
e
n
o
b
l
e
g
a
s
f
a
m
i
l
y
a
r
e
t
w
o
c
h
e
m
i
c
a
l
l
y
i
m
p
o
r
t
a
n
t
g
r
o
u
p
s
o
f
n
o
n
-
m
e
t
a
l
s
.
T
h
e
y
a
r
e
t
h
e
h
a
l
o
g
e
n
s
(
G
r
o
u
p
1
7
)
a
n
d
t
h
e
c
h
a
l
c
o
g
e
n
s
(
G
r
o
u
p
1
6
)
.
These
two
groups
of
elements
have
highly
negative
electron
gain
enthalpies
and
readily
add
one
or
two
electrons
respectively
to
attain
the
stable
noble
gas
configuration.
The
non-metallic
character
increases
as
we
move
from
left
to
right
across
a
period
and
metallic
character
increases
as
we
go
down
the
group
D
-
b
l
o
c
k
:
E
l
e
m
e
n
t
s
i
n
w
h
i
c
h
t
h
e
d
-
o
r
b
i
t
a
l
s
a
r
e
c
o
m
p
l
e
t
e
l
y
f
i
l
l
e
d
a
r
e
c
a
l
l
e
d
t
h
e
d
-
b
l
o
c
k
e
l
e
m
e
n
t
s
.
G
e
n
e
r
a
l
e
l
e
c
t
r
o
n
i
c
c
o
n
f
i
g
u
r
a
t
i
o
n
i
s
(
n
-
1
)
d
1
-
1
0
n
s
1
-
2
.
P
r
o
p
e
r
t
i
e
s
:
-
-Shows
catalytic
properties
-forms
coloured
ions
-shows
paramagnetic
properties
-forms
complexes
They
are
often
referred
to
as
transition
metals
as
they
are placed
between
metals
and
non
metals
and
also
due
to
their
charecteristic
feature.
F
-
b
l
o
c
k
:
Elements
in
which
the
f-orbital
is
completely
filled
are
called
the
f-
block
elements.
General
electronic
configuration
is
ns
2
(n-1)d
0-1
(n-2)f
1-10
.
Inner
transition
metals
are
known
as
f-block
elements.
They
have
incomplete
d-orbital
electronic
confg
in
their
normal
state
and
in
their
oxidation
state.
They
are
generally
radioactive.
T
r
e
n
d
s
i
n
p
h
y
s
i
c
a
l
p
r
o
p
e
r
t
i
e
s
W
e
w
i
l
l
b
e
s
t
u
d
y
i
n
g
,
A
t
o
m
i
c
r
a
d
i
u
s
I
o
n
i
c
r
a
d
i
u
s
I
o
n
i
s
a
t
i
o
n
e
n
t
h
a
l
p
y
E
l
e
c
t
r
o
n
g
a
i
n
e
n
t
h
a
l
p
y
E
l
e
c
t
r
o
n
e
g
a
t
i
v
i
t
y
O
x
i
d
a
t
i
o
n
s
t
a
t
e
s
C
h
e
m
i
c
a
l
r
e
a
c
t
i
v
i
t
y
A
t
o
m
i
c
r
a
d
i
u
s
The
atomic
size
generally
decreases
across
a
period
as
for
the
elements
of
the
second
period.
It
is because
within
the
period
the outer
electrons
are
in
the
same
valence
shell
and
the
effective
nuclear
increases
as
the
atomic
number
increases.
This
results
in
increased
attraction
between
nucleus
and
the
electron
Atomic
radius
decreases
from
left
to
right
as
principal
quantum
number
increases.
It
also
increases
down
a
group
as
the
number
of
shells
increase.
V
a
r
i
a
t
i
o
n
o
f
a
t
o
m
i
c
s
i
z
e
t
o
t
h
a
t
o
f
a
l
k
a
l
i
m
e
t
a
l
s
I
o
n
i
c
r
a
d
i
u
s
T
h
e
r
e
m
o
v
a
l
o
f
a
n
e
l
e
c
t
r
o
n
f
r
o
m
a
n
a
t
o
m
r
e
s
u
l
t
s
i
n
t
h
e
f
o
r
m
a
t
i
o
n
o
f
a
c
a
t
i
o
n
,
w
h
e
r
e
a
s
g
a
i
n
o
f
a
n
e
l
e
c
t
r
o
n
l
e
a
d
s
t
o
a
n
a
n
i
o
n
.
A
cation
is
always
smaller
than
its
parent
atom
while
a
anion
is
always
greater
to
its parent
atom.
This
is
due
to
electron
electron
repulsion.
Isoelectronic
species
are
those
species
which
have
Two
or
more
species
with
same
number
of
atoms,
same
number
of
valence
electrons
and
same
structure,
regardless
of
the
nature
of
elements
involved.
The
cation
with
the
greater
positive
charge
will
have
a
smaller radius
because
of
the
greater
attraction
of
the
electrons
to
the
nucleus.
Anion
with
the
greater
negative
charge
will
have
the
larger
radius.
In
this
case,
the net
repulsion
of
the
electrons
will
outweigh
the
nuclear
charge and
the
ion
will
expand
in
size.
I
o
n
i
s
a
t
i
o
n
e
n
t
h
a
l
p
y
A
q
u
a
n
t
i
t
a
t
i
v
e
m
e
a
s
u
r
e
o
f
t
h
e
t
e
n
d
e
n
c
y
o
f
a
n
e
l
e
m
e
n
t
t
o
l
o
s
e
e
l
e
c
t
r
o
n
i
s
g
i
v
e
n
b
y
i
t
s
I
o
n
i
z
a
t
i
o
n
E
n
t
h
a
l
p
y
.
I
t
r
e
p
r
e
s
e
n
t
s
t
h
e
e
n
e
r
g
y
r
e
q
u
i
r
e
d
t
o
r
e
m
o
v
e
a
n
e
l
e
c
t
r
o
n
f
r
o
m
a
n
i
s
o
l
a
t
e
d
g
a
s
e
o
u
s
a
t
o
m
(
X
)
i
n
i
t
s
g
r
o
u
n
d
s
t
a
t
e
.
I
n
o
t
h
e
r
w
o
r
d
s
,
t
h
e
f
i
r
s
t
i
o
n
i
z
a
t
i
o
n
e
n
t
h
a
l
p
y
f
o
r
a
n
e
l
e
m
e
n
t
X
i
s
t
h
e
e
n
t
h
a
l
p
y
c
h
a
n
g
e
(
∆
i
H
)
f
o
r
t
h
e
r
e
a
c
t
i
o
n
X
(
g
)
→
X
+
(
g
)
+
e
–
The
ionization
enthalpy
is
expressed
in
units
of
kJ mol–1.
We
can
define
the
second
ionization
enthalpy
as
the
energy
required
to
remove
the
second
most
loosely
bound
electron
X
+
(
g
)
→
X
2
+
(
g
)
+
e
–
Energy is
always
required
to
remove
electrons
from
an
atom
and
hence
ionization
enthalpies
are always
positive.
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.
In
the
same
way
the
third
ionization
enthalpy
will
be
higher
than
the
second
and
so
on.
The
term
“ionization
enthalpy”,
if
not
qualified,
is
taken
as
the
first
ionization
enthalpy.
Energy
is
always
required
to
remove
electrons
from
an
atom
and
hence
ionization
enthalpies
are
always
positive.
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.
In
the
same
way
the
third
ionization
enthalpy
will
be
higher
than
the
second
and
so
on.
The
term
“ionization
enthalpy”,
if
not
qualified,
is
taken
as
the
first
ionization
enthalpy.
We
will
find
maxima
at the
noble
gases
which
have
closed
electron
shells
and
very
stable
electron
configurations.
On
the other
hand,
minima
occur
at
the
alkali
metals
and
their
low
ionization
enthalpies
can
be
correlated
with
their
high
reactivity.
S
h
i
e
l
d
i
n
g
e
f
f
e
c
t
I
n
m
u
l
t
i
e
l
e
c
t
r
o
n
a
t
o
m
s
,
t
h
e
e
l
e
c
t
r
o
n
s
a
r
e
p
r
e
s
e
n
t
i
n
t
h
e
o
u
t
e
r
m
o
s
t
s
h
e
l
l
a
n
d
t
h
e
y
d
o
n
o
t
e
x
p
e
r
i
e
n
c
e
a
f
u
l
l
p
o
s
i
t
i
v
e
c
h
a
r
g
e
h
e
n
c
e
t
h
e
y
a
r
e
s
h
i
e
l
d
e
d
o
r
s
c
r
e
e
n
e
d
a
w
a
y
f
r
o
m
t
h
e
n
u
c
l
e
u
s
.
F
o
r
E
g
:
-
i
n
t
h
e
s
a
m
e
s
u
b
s
h
e
l
l
t
h
e
p
-
e
l
e
c
t
r
o
n
s
w
i
l
l
g
e
t
s
h
i
e
l
d
e
d
a
w
a
y
b
y
t
h
e
s
-
e
l
e
c
t
r
o
n
a
s
p
-
e
l
e
c
t
r
o
n
i
s
a
w
a
y
f
r
o
m
t
h
e
i
n
f
l
u
e
n
c
e
o
f
n
u
c
l
e
u
s
a
n
d
h
e
n
c
e
i
t
s
d
o
e
s
n
o
t
e
x
p
e
r
i
e
n
c
e
a
c
o
m
p
l
e
t
e
p
o
s
i
t
i
v
e
c
h
a
r
g
e
!
I
o
n
i
s
a
t
i
o
n
a
n
d
a
t
o
m
i
c
r
a
d
i
u
s
a
r
e
c
o
r
r
e
l
a
t
e
d
i
n
2
w
a
y
s
:
-
(
i
)
t
h
e
a
t
t
r
a
c
t
i
o
n
o
f
e
l
e
c
t
r
o
n
s
t
o
w
a
r
d
s
t
h
e
n
u
c
l
e
u
s
,
a
n
d
(
i
i
)
t
h
e
r
e
p
u
l
s
i
o
n
o
f
e
l
e
c
t
r
o
n
s
f
r
o
m
e
a
c
h
o
t
h
e
r
.
T
h
e
e
f
f
e
c
t
i
v
e
n
u
c
l
e
a
r
c
h
a
r
g
e
e
x
p
e
r
i
e
n
c
e
d
b
y
a
v
a
l
e
n
c
e
e
l
e
c
t
r
o
n
i
n
a
n
a
t
o
m
w
i
l
l
b
e
l
e
s
s
t
h
a
n
t
h
e
a
c
t
u
a
l
c
h
a
r
g
e
o
n
t
h
e
n
u
c
l
e
u
s
b
e
c
a
u
s
e
o
f
“
s
h
i
e
l
d
i
n
g
”
o
r
“
s
c
r
e
e
n
i
n
g
”
o
f
t
h
e
v
a
l
e
n
c
e
e
l
e
c
t
r
o
n
f
r
o
m
t
h
e
n
u
c
l
e
u
s
b
y
t
h
e
i
n
t
e
r
v
e
n
i
n
g
c
o
r
e
e
l
e
c
t
r
o
n
s
.
A
n
o
m
a
l
i
e
s
2s
electron
of
lithium
is
shielded
away.
The
2
s
electron
in
lithium
is
shielded
from
the
nucleus
by
the
inner
core
of
1
s
electrons.
As
a
result,
the
valence
electron
experiences
a
net
positive
charge
which
is
less
than
the
actual
charge
of
+3.
In
general,
shielding
is
effective
when
the
orbitals
in
the
inner
shells
are
completely
filled.
This
situation
occurs
in
the
case
of
alkali
metals
which
have
single
outermost
ns
-electron
preceded
by
a
noble
gas
electronic
configuration.
When
we
move
from
lithium
to
fluorine
across
the
second
period,
successive
electrons
are
added
to
orbitals
in
the
same
principal
quantum
level
and
the
shielding
of
the
nuclear
charge
by
the
inner
core
of
electrons
does
not
increase
very
much
to
compensate
for
the
increased
attraction
of
the
electron
to
the
nucleus.
Thus,
across
a
period,
increasing
nuclear
charge
outweighs
the
shielding.
Consequently,
the
outermost
electrons
are
held
more
and
more
tightly
and
the
ionization
enthalpy
increases
across
a
period.
As
we
go
down
a
group,
the
outermost
electron
being
increasingly
farther
from
the
nucleus,
there
is an
increased
shielding
of
the
nuclear
charge
by
the
electrons
in
the
inner
levels.
In
this
case,
increase
in
shielding
outweighs
the
increasing
nuclear
charge
and
the
removal
of
the
outermost
electron
requires
less
energy
down
a
group.
I
t
i
s
e
a
s
i
e
r
t
o
r
e
m
o
v
e
2
p
e
l
e
c
t
r
o
n
o
f
b
o
r
o
n
t
h
a
n
2
s
e
l
e
c
t
r
o
n
o
f
b
e
r
y
l
l
i
u
m
.
T
h
e
f
i
r
s
t
i
o
n
i
s
a
t
i
o
n
e
n
t
h
a
l
p
y
o
f
b
o
r
o
n
i
s
s
l
i
g
h
t
l
y
l
e
s
s
t
h
a
n
b
e
r
y
l
l
i
u
m
t
h
o
u
g
h
t
h
e
f
o
r
m
e
r
h
a
s
a
g
r
e
a
t
e
r
n
u
c
l
e
a
r
c
h
a
r
g
e
.
W
h
e
n
w
e
c
o
n
s
i
d
e
r
t
h
e
s
a
m
e
p
r
i
n
c
i
p
a
l
q
u
a
n
t
u
m
l
e
v
e
l
,
a
n
s
-
e
l
e
c
t
r
o
n
i
s
a
t
t
r
a
c
t
e
d
t
o
t
h
e
n
u
c
l
e
u
s
m
o
r
e
t
h
a
n
a
p
-
e
l
e
c
t
r
o
n
.
In beryllium,
the
electron
removed
during
the
ionization
is an
s
-
electron
whereas
the
electron
removed
during
ionization
of
boron
is
a
p
-electron.
The
penetration
of
a
2
s-
electron
to
the
nucleus
is
more
than
that
of
a
2
p-
electron;
hence
the
2
p
electron
of
boron
is
more
shielded
from
the
nucleus
by
the
inner
core
of
electrons
than
the
2
s
electrons
of
beryllium.
Therefore,
it
is
easier
to
remove
the
2
p
-electron
from
boron
compared
to
the
removal
of
a
2
s-
electron
from
beryllium.
Thus,
boron
has
a
smaller
first
ionization
enthalpy
than
beryllium
F
i
r
s
t
i
o
n
i
s
a
t
i
o
n
e
n
t
h
a
l
p
y
o
f
o
x
y
g
e
n
i
s
l
e
s
s
t
h
a
n
n
i
t
r
o
g
e
n
.
The
smaller
first
ionization
enthalpy
of
oxygen
compared
to
nitrogen.
This
arises because
in
the
nitrogen
atom,
three
2
p-
electrons
reside
in
different
atomic
orbitals
(Hund’s
rule)
whereas
in
the
oxygen
atom,
two
of
the
four
2
p-
electrons
must
occupy
the
same
2
p-
orbital
resulting
in
an
increased
electron-electron
repulsion.
Consequently,
it
is
easier
to
remove
the
fourth
2
p
-
electron
from
oxygen
than
it
is,
to
remove
one
of
the
three
2
p
-
electrons
from
nitrogen.
G
i
v
e
r
e
a
s
o
n
s
W
h
y
a
r
e
l
a
n
t
h
a
n
i
d
e
s
a
n
d
a
c
t
i
n
i
d
e
s
p
l
a
c
e
d
s
e
p
a
r
a
t
e
l
y
a
t
t
h
e
b
o
t
t
o
m
o
f
t
h
e
p
e
r
i
o
d
i
c
t
a
b
l
e
?
(a)
F
o
r
c
o
n
v
e
n
i
e
n
c
e
a
n
d
s
y
s
t
e
m
a
t
i
c
s
t
u
d
y
o
f
e
l
e
m
e
n
t
s
h
a
v
i
n
g
s
i
m
i
l
a
r
p
r
o
p
e
r
t
i
e
s
.
(b)
T
o
m
a
i
n
t
a
i
n
t
h
e
s
t
r
u
c
t
u
r
e
o
f
p
e
r
i
o
d
i
c
t
a
b
l
e
o
f
c
l
a
s
s
i
f
i
c
a
t
i
o
n
b
y
k
e
e
p
i
n
g
e
l
e
m
e
n
t
s
w
i
t
h
s
i
m
i
l
a
r
p
r
o
p
e
r
t
i
e
s
i
n
a
s
i
m
i
l
a
r
w
a
y
.
E
l
e
m
e
n
t
s
l
i
k
e
Z
n
,
P
d
,
C
d
a
r
e
n
o
t
c
o
n
s
i
d
e
r
e
d
t
r
a
n
s
i
t
i
o
n
e
l
e
m
e
n
t
s
?
(
a
)
A
s
t
h
e
y
h
a
v
e
i
n
c
o
m
p
l
e
t
e
l
y
f
i
l
l
e
d
d
-
o
r
b
i
t
a
l
s
C
h
e
m
i
s
t
r
y
o
f
A
c
t
i
n
o
i
d
s
i
s
c
o
m
p
l
i
c
a
t
e
d
w
h
y
?
A
s
t
h
e
y
s
h
o
w
l
a
r
g
e
n
o
o
f
o
x
i
d
a
t
i
o
n
s
t
a
t
e
s
E
l
e
c
t
r
o
n
g
a
i
n
e
n
t
h
a
l
p
y
w
h
e
n
a
n
e
l
e
c
t
r
o
n
i
s
a
d
d
e
d
t
o
a
n
e
u
t
r
a
l
g
a
s
e
o
u
s
a
t
o
m
(
X
)
t
o
c
o
n
v
e
r
t
i
t
i
n
t
o
a
n
e
g
a
t
i
v
e
i
o
n
,
t
h
e
e
n
t
h
a
l
p
y
c
h
a
n
g
e
a
c
c
o
m
p
a
n
y
i
n
g
t
h
e
p
r
o
c
e
s
s
i
s
d
e
f
i
n
e
d
a
s
t
h
e
E
l
e
c
t
r
o
n
G
a
i
n
E
n
t
h
a
l
p
y
(
∆
e
g
H
)
.
Electron
gain
enthalpy
provides
a
measure
of
the
ease
with
which
an
atom
adds
an
electron
to
form
anion
X
(
g
)
+
e
–
→
X
–
(
g
)
Depending
on
the
element,
the
process
of
adding
an
electron
to
the
atom
can
be
either
endothermic
or
exothermic
Added
to
the
atom
and
the
electron
gain
enthalpy
is
negative.
For
example,
group
17
elements
(the
halogens)
have
very
high
negative
electron
gain
enthalpies
because
they
can
attain
stable
noble
gas
electronic
configurations
by
picking
up
an
electron
On
the
other
hand,
noble
gases
have
large
positive
electron
gain
enthalpies
because
the
electron
has
to
enter
the
next
higher
principal
quantum
level
leading
to
a
very
unstable
electronic
configuration.
It
may
be
noted
that
electron
gain
enthalpies
have
large
negative
values
toward
the
upper
right
of
the
periodic
table
preceding
the
noble
gases.
As
a
general
rule,
electron
gain
enthalpy
becomes
more
negative
with
increase
in
the
atomic
number
across
a
period.
The
effective
nuclear
charge
increases
from
left
to
right
across
a
period
and
consequently
it
will
be
easier
to
add
an
electron
to
a
smaller
atom
since
the
added
electron on
an
average
would
be
closer
to
the
positively
charged
nucleus.
We
should
also
expect
electron
gain
enthalpy
to
become
less
negative
as
we
go
down
a
group
because
the
size
of
the
atom
increases
and
the
added
electron
would
be
farther
from
the
nucleus
e
l
e
c
t
r
o
n
e
g
a
t
i
v
i
t
y
A
q
u
a
l
i
t
a
t
i
v
e
m
e
a
s
u
r
e
o
f
t
h
e
a
b
i
l
i
t
y
o
f
a
n
a
t
o
m
i
n
a
c
h
e
m
i
c
a
l
c
o
m
p
o
u
n
d
t
o
a
t
t
r
a
c
t
s
h
a
r
e
d
e
l
e
c
t
r
o
n
s
t
o
i
t
s
e
l
f
i
s
c
a
l
l
e
d
e
l
e
c
t
r
o
n
e
g
a
t
i
v
i
t
y
.
F
l
o
u
r
i
n
e
h
a
s
t
h
e
h
i
g
h
e
s
t
e
l
e
c
t
r
o
n
e
g
a
t
i
v
i
t
y
v
a
l
u
e
o
f
4
.
I
t
i
s
t
h
e
m
o
s
t
e
l
e
c
t
r
o
n
e
g
a
t
i
v
e
e
l
e
m
e
n
t
.
Non-metallic
elements
have
strong
tendency
to
gain
electrons.
Therefore,
electronegativity
is
directly
related
to
that
non-metallic
properties
of
elements.
It
can
be
further
extended
to
say
that
the
electronegativity
is
inversely
related
to
the
metallic
properties
of
elements.
Thus,
the
increase
in
electronegativities
across
a
period
is
accompanied
by
an
increase
in
non-metallic
properties
(or
decrease
in metallic
properties)
of
elements.
Similarly,
the
decrease
in
electronegativity
down
a
group
is
accompanied
by
a
decrease
in
non-metallic
properties
(or
increase
in
metallic
properties)
of
elements.
O
x
i
d
a
t
i
o
n
s
t
a
t
e
T
h
e
v
a
l
e
n
c
e
e
q
u
a
l
s
t
o
t
h
e
n
u
m
b
e
r
o
f
e
l
e
c
t
r
o
n
s
p
r
e
s
e
n
t
i
n
t
h
e
o
u
t
e
r
m
o
s
t
s
h
e
l
l
.
T
h
e
t
e
r
m
o
x
i
d
a
t
i
o
n
s
t
a
t
e
i
n
t
e
r
c
h
a
n
g
e
a
b
l
y
m
e
a
n
s
v
a
l
e
n
c
e
.
F
l
u
o
r
i
n
e
h
a
s
t
h
e
o
x
i
d
a
t
i
o
n
s
t
a
t
e
o
f
-
1
b
u
t
i
n
c
o
m
p
o
u
n
d
s
l
i
k
e
O
F
2
i
t
e
x
h
i
b
i
t
s
+
2
o
x
i
d
a
t
i
o
n
s
t
a
t
e
.
O
x
i
d
a
t
i
o
n
s
t
a
t
e
d
e
t
e
r
m
i
n
e
s
t
h
e
o
x
i
d
a
t
i
o
n
n
u
m
b
e
r
o
f
t
h
e
e
l
e
m
e
n
t
.
C
h
e
m
i
c
a
l
r
e
a
c
t
i
v
i
t
y
D
i
a
g
o
n
a
l
r
e
l
a
t
i
o
n
s
h
i
p
r
e
f
e
r
s
t
o
t
h
e
r
e
l
a
t
i
o
n
s
h
i
p
b
e
t
w
e
e
n
c
e
r
t
a
i
n
p
a
i
r
s
o
f
d
i
a
g
o
n
a
l
l
y
a
d
j
a
c
e
n
t
e
l
e
m
e
n
t
s
i
n
t
h
e
s
e
c
o
n
d
a
n
d
i
n
t
h
e
3
r
d
p
e
r
i
o
d
r
e
s
p
e
c
t
i
v
e
l
y
.
This
characteristic
property
is
shown
by
s-
and
p-blocks
respectively.
he anomalous
behaviour
is
attributed
to
their
small
size,
large
charge/
radius
ratio
and
high
electronegativity
of
the
elements.
In
addition,
the
first
member
of
group
has
only
four
valence
orbitals
(2
s
and
2
p
)
available
for
bonding,
whereas
the
second
member
of
the
groups
have
nine
valence
orbitals
(3
s
,
3
p
,
3
d
).
A
s
a
c
o
n
s
e
q
u
e
n
c
e
o
f
t
h
i
s
,
t
h
e
m
a
x
i
m
u
m
c
o
v
a
l
e
n
c
y
o
f
t
h
e
f
i
r
s
t
m
e
m
b
e
r
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f
e
a
c
h
g
r
o
u
p
i
s
4
(
e
.
g
.
,
b
o
r
o
n
c
a
n
o
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l
y
f
o
r
m
[
B
F
4
]
−
,
w
h
e
r
e
a
s
t
h
e
o
t
h
e
r
m
e
m
b
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r
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o
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g
r
o
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a
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p
a
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m
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g
.
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a
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s
[
A
l
F
6
]
3
-
.
F
u
r
t
h
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r
m
o
r
e
,
t
h
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f
i
r
s
t
m
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m
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p
-
b
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o
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m
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n
t
s
d
i
s
p
l
a
y
s
g
r
e
a
t
e
r
a
b
i
l
i
t
y
t
o
f
o
r
m
p
π
–
p
π
m
u
l
t
i
p
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e
b
o
n
d
s
t
o
i
t
s
e
l
f
(
e
.
g
.
C
=
C
,
C
≡
C
,
N
=
N
,
N
≡
Ν
)
a
n
d
t
o
o
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r
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(
e
.
g
.
,
C
=
O
,
C
=
N
,
C
≡
N
,
N
=
O
)
c
o
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t
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.
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i
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i
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d
u
e
t
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a
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i
t
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t
r
o
n
s
i
n
v
a
r
i
o
u
s
o
r
b
i
t
a
l
s
.
E
l
e
m
e
n
t
s
w
h
i
c
h
c
o
m
b
i
n
e
o
n
t
h
e
t
w
o
e
x
t
r
e
m
e
s
o
f
t
h
e
p
e
r
i
o
d
i
c
t
a
b
l
e
f
o
r
m
o
x
i
d
e
s
l
i
k
e
p
e
r
o
x
i
d
e
s
a
n
d
s
u
p
e
r
o
x
i
d
e
s
.
S
o
m
e
a
r
e
a
c
i
d
i
c
b
a
s
i
c
a
n
d
n
e
u
t
r
a
l
o
r
b
e
t
t
e
r
k
n
o
w
n
a
s
a
m
p
h
o
t
e
r
i
c
o
x
i
d
e
s
M
e
t
a
l
s
f
r
o
m
b
a
s
i
c
o
x
i
d
e
s
,
n
o
n
m
e
t
a
l
f
o
r
m
a
c
i
d
i
c
o
x
i
d
e
s
.
Among
transition
metals
(3
d
series),
the
change
in
atomic
radii
is
much smaller
as
compared
to
those
of
representative
elements
across
the
period.
The
change
in
atomic
radii
is
still
smaller
among
inner-transition
metals
(4
f
series).
The
ionization
enthalpies
are
intermediate
between
those
of
s-
and
p-
blocks.
As
a
consequence,
they
are
less
electropositive
than
group
1
and
2
metals.
T
h
e
e
n
d
.
The journey of understanding the classification of elements and periodicity in properties begins with early laws like the Law of Triads and Newland's Law of Octaves. Mendeleev's Periodic Law revolutionized the organization of elements, leading to the modern periodic table. Discoveries of eka-aluminium and eka-silicon further enriched our understanding of elemental properties. The Periodic Law highlights analogies among elements and the division of the periodic table into s-block, p-block, d-block, and f-block. Explore the periodic table's structure and the variations in element properties across periods and groups.
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Presentation Transcript
Classification of elements and periodicity in properties Chapter3
Lets recall andrevise Law of triads:- According to law of triads elements were placed in a group of 3 andthey were called triads. These were placed in the increasing order of atomic masses Newland s law of octaves:- the elements were placed in increasing order of their atomic weights and it was noted that every eighth element had properties similar to the first element . The relationship was just like every eighth note that resembles the first in octaves of music. Newlands s Law of Octaves seemed to be true to elements only upto calcium.
Mendeleevs periodiclaw The properties of theelements are a periodic function of their atomic weights. Mendeleev arranged elements in horizontal rows and vertical columns of a table in order of their increasingatomic weights in such a way that the elements with similar properties occupied the same vertical column or group .
Mendeleevleft gapsunder silicon and aluminium and those were known as eka aluminium and eka silicon. Later it was discovered that eka aluminium had properties similarto gallium and eka silicon had properties similar togermanium. According to modern periodiclaw, The physical and chemical propertiesof the elements are periodic functions of their atomicnumbers . Modern form of periodic table was discovered by henry Mosley andhence it was also called henry Mosley's periodic table.
The Periodic Law revealed important analogies among the 94 naturally occurring elements (neptunium and plutonium like actinium andprotoactinium arealsofound in pitch blende anore of uranium). The horizontalrowsof the periodic table are calledperiods while the vertical columns are called groups orfamilies.
DIVISION OF PERIODICTABLE The periodic table is divided into 4 blocks mainly the s-block p-block d- block andf-block. S-block:- Theelements of Group1(alkali metals) andGroup2(alkaline earth metals) which have ns1and ns2outermost electronic configuration belong to the s- BlockElements. They are all reactive metals with low ionization enthalpies. They lose the outermost electron(s)readilyto form 1+ion inthe caseof alkalimetalsor 2+ ionin the case of alkalineearth metals. Themetalliccharacterandthe reactivity increaseaswe go down the group. Becauseof highreactivity they areneverfound pureinnature. Thecompoundsof the s-blockelements,with the exception of those of lithium andberyllium arepredominantly ionic.
P-block: Thep-BlockElementscomprise those belonging to Group13to 18andthese together with the s-BlockElementsare called the Representative Elements or Main GroupElements. Theoutermost electronicconfigurationvariesfrom ns2np1to ns2np6ineach period. At the end of each period is a noble gas element with a closed valence shell ns2np6configuration. Allthe orbitalsin the valenceshellof the noblegasesarecompletely filled byelectronsandit isverydifficult to alterthis stablearrangement by the addition or removalof electrons. The noble gases thus exhibit very low chemical reactivity. Preceding the noble gasfamily aretwo chemicallyimportant groupsof non-metals.They arethe halogens(Group 17)andthe chalcogens(Group 16). Thesetwo groupsof elements havehighlynegativeelectrongainenthalpies andreadilyaddone or two electronsrespectivelyto attain the stablenoble gasconfiguration. Thenon-metalliccharacterincreasesaswe move from left to right acrossa period andmetalliccharacterincreasesaswe go down the group
D-block: Elements in which the d-orbitals are completely filled are calledthe d-blockelements. General electronic configuration is (n-1)d1-10ns1-2 . Properties:- -Shows catalytic properties -forms colouredions -shows paramagneticproperties -formscomplexes They are often referred to as transition metals as they are placed between metals andnon metals andalsodueto their charecteristic feature.
F-block: Elements in which the f-orbital iscompletely filled arecalled the f- block elements. General electronic configuration is ns2(n-1)d0-1(n-2)f1-10. Inner transition metalsareknown asf-block elements. Theyhaveincomplete d-orbital electronic confg in their normal state and in theiroxidation state. They are generallyradioactive.
Trends in physicalproperties We will bestudying, Atomicradius Ionicradius Ionisation enthalpy Electron gainenthalpy Electronegativity Oxidationstates Chemicalreactivity
Atomicradius The atomic size generallydecreases across a period as for the elements of the second period. It is because within the period the outer electrons are in the same valence shell and the effective nuclear increases as the atomic number increases. This results in increasedattraction between nucleus and theelectron Atomic radiusdecreasesfrom left to right as principal quantum number increases. It alsoincreasesdown a group as the numberof shellsincrease.
Variation of atomicsize to that of alkalimetals
Ionicradius The removal of an electron from an atom results in the formation of a cation,whereas gainof anelectronleadsto ananion. Acationisalwayssmallerthan itsparent atom while a anion is always greater to its parent atom. Thisisdueto electronelectronrepulsion. Isoelectronic species are thosespecies which haveTwoor morespecieswith samenumber of atoms, samenumber of valenceelectrons andsamestructure, regardlessof the nature of elementsinvolved. The cation with the greater positive charge will have a smaller radius because of the greater attraction of the electrons to the nucleus. Anion with the greater negative chargewill havethe largerradius.In this case, the net repulsion of the electrons will outweigh the nuclear charge and the ion will expand insize.
Ionisationenthalpy A quantitative measure of the tendency of an element to lose electron is given by its Ionization Enthalpy. It represents the energy required to remove anelectron from anisolated gaseous atom (X) in its ground state. In other words, the first ionization enthalpy for an element X is the enthalpy change ( i H) for the reaction X(g) X+(g)+ e Theionizationenthalpy isexpressed in unitsof kJ mol 1. We can define the second ionization enthalpy asthe energyrequired to remove the second most loosely boundelectron X+(g) X2+ (g)+ e Energyis always required to remove electrons from an atom and hence ionization enthalpies are alwayspositive. 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 aneutral atom. In the same way the third ionization enthalpy will behigher thanthe secondandsoon.The term ionization enthalpy , if not qualified, is taken as the firstionization enthalpy.
Energy is always required toremove electrons from an atom and hence ionization enthalpies are always positive. The second ionization enthalpy will be higher than the first ionization enthalpybecauseit ismore difficult to remove an electron from a positively chargedion than from aneutralatom. In the sameway the third ionization enthalpy will be higher than the second and so on. Theterm ionization enthalpy , if not qualified, is taken as the first ionization enthalpy. We will find maxima at the noble gases which have closed electron shells and very stable electron configurations. On the other hand, minima occurat the alkalimetalsand their low ionization enthalpiescanbe correlated with their highreactivity.
Shieldingeffect In multi electron atoms , the electrons arepresent in the outermostshell and they do not experience a full positive charge hence they are shielded or screened away from thenucleus. For Eg:-in the same subshell the p-electrons will get shielded away bythes-electron as p-electron is away from the influence of nucleus and hence it s does not experience a complete positive charge!
Ionisation and atomic radius are correlated in2 ways:- (i) the attraction of electrons towards the nucleus, and (ii) the repulsion of electrons from each other. The effective nuclear chargeexperienced by a valence electron in an atom will be less than the actual charge on the nucleus because of shielding or screening of the valence electron from the nucleus by the intervening coreelectrons.
Anomalies 2s electron oflithium is shielded away. The 2selectron in lithium isshielded from the nucleus by the inner core of 1selectrons.Asaresult, the valenceelectron experiences a net positive charge which is less than the actual charge of +3. In general, shielding is effective when the orbitals in the inner shells are completely filled. This situation occurs in the case of alkali metals which have single outermost ns-electron preceded by a noble gas electronicconfiguration.
Whenwe movefrom lithium to fluorine acrossthe secondperiod, successive electrons are added to orbitals in the same principal quantum level andthe shielding of the nuclear charge by the inner core of electrons does not increase very much to compensate for the increased attraction of the electron to the nucleus. Thus, across a period, increasing nuclear charge outweighs the shielding. Consequently, the outermost electronsareheld more and more tightly and the ionization enthalpy increases across a period. As we go down a group, the outermost electron being increasingly farther from the nucleus, there is an increased shielding of the nuclear charge by the electronsin the inner levels. In this case,increasein shielding outweighs the increasingnuclear charge and the removal of the outermost electron requires less energy down agroup.
It is easier to remove 2p electron of boron than 2s electronof beryllium. The first ionisation enthalpy of boron is slightly less thanberyllium though the former hasagreater nuclear charge. Whenwe consider the same principal quantum level, an s-electron is attracted to the nucleus more than ap-electron. In beryllium, the electron removed during the ionization is an s- electron whereas the electron removed during ionization of boron is ap-electron. Thepenetration of a2s-electronto the nucleus ismore than that of a 2p-electron; hencethe 2pelectron of boron ismore 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 comparedto the removalof a2s-electron from beryllium.Thus, boron hasasmallerfirst ionization enthalpy than beryllium
First ionisation enthalpy of oxygen is less thannitrogen. The smaller first ionization enthalpy of oxygen compared to nitrogen. This arises because in the nitrogen atom, three 2p- electronsresidein different atomic orbitals (Hund s rule) whereas in the oxygen atom,two of the four 2p-electrons must occupythe 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 2 p- electrons fromnitrogen.
Givereasons Why are lanthanides and actinides placed separately at the bottom of the periodic table ? (a) For convenience and systematic study of elements having similar properties. (b) To maintain the structure of periodic table of classification by keeping elements with similar properties in a similar way. Elements like Zn , Pd, Cd are not considered transition elements ? (a) As they have incompletely filled d-orbitals Chemistry of Actinoids is complicated why ? As they show large no of oxidation states
Electron gainenthalpy when an electron is added to a neutral gaseous atom (X) to convert it into a negative ion, the enthalpy change accompanying the process isdefined asthe Electron Gain Enthalpy( egH). Electron gain enthalpy provides ameasure of the easewith which anatom adds an electron to formanion X(g)+e X (g) Depending on the element,the process of adding anelectron to the atom canbe either endothermic orexothermic Added to the atom and the electron gain enthalpy isnegative. Forexample, group 17 elements (the halogens) have very high negative electron gain enthalpies becausethey canattain stable noble gaselectronic configurations bypicking up an electron On the other hand, noble gases have large positive electron gain enthalpies becausethe electron hasto enter the next higher principal quantum level leading to a very unstable electronic configuration. It may be noted that electron gain enthalpies havelarge negative valuestoward the upper right of the periodic table preceding the noble gases.
As a general rule, electron gain enthalpy becomes morenegative with increase in the atomic number across a period. The effective nuclear charge increasesfrom left to right across a period and consequently it will be easier to add an electron to a smaller atom since the added electron on an average would be closer to the positively charged nucleus. We should also expect electron gain enthalpy tobecome less negative as we go down a group because the size of the atom increases and the added electron would be farther from the nucleus
electronegativity A qualitative measureof the ability of anatom in a chemical compound to attract shared electrons to itself is called electronegativity. Flourine has the highest electronegativityvalue of 4. It is the most electronegative element.
Non-metallic elements have strong tendency to gain electrons. Therefore, electronegativity isdirectly related to that non-metallic properties ofelements. It can be further extended to say that the electronegativity is inversely related to the metallic properties of elements.Thus,the increase in electronegativities across a period is accompanied by an increase in non-metallic properties (or decrease in metallic properties) ofelements. Similarly, the decrease in electronegativity down a group is accompanied by adecreasein non-metallic properties (or increase in metallic properties) of elements.
Oxidationstate The valence equals to the number of electrons present inthe outermostshell. The term oxidation state interchangeably meansvalence. Fluorine has the oxidation state of -1 but in compounds like OF2 it exhibits +2 oxidationstate. Oxidation state determines the oxidation number of theelement.
Chemicalreactivity Diagonal relationship refers to the relationship between certain pairs of diagonally adjacent elements in the second and in the 3rd period respectively. Thischaracteristicproperty isshown by s-andp-blocks respectively.
he anomalous behaviour is attributed to their small size, large charge/radiusratio andhigh electronegativity of the elements. In addition, the first memberof group hasonly four valenceorbitals (2sand 2p) available for bonding, whereas the secondmemberof the groups havenine valenceorbitals (3s,3p,3d). As a consequence of this, the maximum covalency of the first member of each group is 4 (e.g., boron can only form [BF4 ] , whereas the other members of the groups can expand their valenceshell to accommodatemore than four pairsof electrons e.g., aluminium forms [AlF6 ]3-. Furthermore, the first member of p-block elements displays greater ability to form p p multiple bonds to itself (e.g. C=C,C C,N=N,N ) andto other secondperiod elements (e.g.,C =O,C=N,C N,N=O)compared to subsequent membersof the samegroup.
Metallic character decreases from right to leftwhereas non metallic character increases from left toright. This is due to addition of electrons in variousorbitals. Elements which combine on the two extremes of theperiodic table form oxides like peroxides and superoxides. Some are acidic basic and neutral or bette rknown as amphoteric oxides Metals from basic oxides , non metal form acidicoxides. Among transition metals (3d series), the change in atomic radii is much smaller as compared to those of representative elements acrossthe period. Thechange in atomic radii isstill smalleramong inner-transition metals (4f series). The ionization enthalpies are intermediate between those of s-andp-blocks.Asaconsequence, they arelesselectropositive than group 1and2metals.
