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INORGANIC CHEMISTRY (AS)
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The Periodic table
Groups II , VII , V and VI
Transition elements
THE PERIODIC TABLE (AS)
Trends in physical and chemical
properties
Across a period from left to right
Periodic table
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1869 : proposed by Mendeleev
63 elements arranged according to
increasing atomic mass
1913 : modern periodic table proposed
by Moseley
108/109 elements arranged according
to increasing atomic number
Structure of the Table
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Elements are arranged according to their atomic no.
Period = the horizontal row
Group = the vertical column
Divided into sections:
‘s-block’ (Groups I and II) : outermost subshell ‘s’
‘p-block’ (Groups III to VIII):outermost subshell ‘p’
‘d-block’ with transition elements : partially filled d
subshell
Diagonal relationship
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First element in any group usually
shows unusual behaviour unlike the rest
of the elements in the same group
More similar to the second element in
the neighbouring group
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Example :
Group
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I
Li
II
Be
III
B
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Na
Mg
Al
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Rxn
Heat
carbonate
Typical
Group I
element
No
reaction
Li
(Group I)
Heat
nitrate
MNO2 + O2 Li2O + O2
+ NO2
Mg
(Group II)
Li2O + CO2 MgO +
CO2
MgO + O2
+ NO2
Trends in physical properties
( Physical periodicity )
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1. Atomic and ionic radii
a. Atomic radius :
i) Size of atoms = atomic (covalent) radii
= half internuclear distance between 2
atoms of the same element joined by a
single bond
Eg
Cl - Cl
Size of Cl = x/2
x
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Note :atomic radius for other elements
(1)Metals – metallic radius of the atom
in the metal
(2)Noble gases -atomic/ VDW radius is
half the average distance between
adjacent non-bonded atoms -unusually
large value
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ii) Atomic size decreases from left to right
due to increase in nuclear charge
iii) Decrease in size becomes smaller with
increasing atomic number due to increased
repulsion between electrons
Eg
Li
Be
B
C
Radius
0.123 0.089 0.080
0.077
Decrease
0.034
0.009
0.003
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Period 2
Atomic
Li
radius
Be
F
Atomic number
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Period 3 ( data from Data Booklet )
Atomic
Unusually
large value
radius
(VDW radius)
Ar
Atomic number
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b. Ionic radius
i) Positive ions :
Atom - electrons  Positive ion
(p = e)
( p > e)
In ion , nuclear charge has greater pull on
remaining electrons
Or sometimes positive ion has one shell less
than atom
Size of positive ion smaller than neutral atom
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ii) Negative ions :
Atom + electrons  Negative ion
(p = e)
( p  e)
In ion , nuclear charge has weaker pull on
the electrons
Also increased repulsion between electrons
Size of negative ion bigger than neutral atom
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c. Eg : Period 3
i) Positive ions
Na+ Mg2+
Al3+ Si4+
Ionic radius
0.095 0.065 0.050 0.041
( nm)
All ions are isoelectronic (10 e ), [Ne]
From Na to Si, nuclear charge increases
Attractive force on outer electrons increases
Decrease in ionic radii
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ii)Negative ions
P3S2ClIonic radius/nm
0.212 0.184 0.181
All ions are isoelectronic ,18 e , [Ar]
From P to Cl, nuclear charge increases
Attractive force on outer electrons increases
Decrease in ionic radii
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iii)Radii of negative ions are larger than
that of positive ions
Reasons :
Extra shell of electrons in negative ions
( 18 e vs 10 e ) (*)
Increased repulsion between larger no
of electrons
iv)Graph
Ionic
radius
Na+
Si4+
P3-
Cl-
Atomic
number
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2. Melting* and boiling points :
Depends on strength of the bonds and
on the structure in the solid state
Eg Period 3 ( Na to Ar )
Na Mg Al
Si P S Cl
Ar
Giant metallic
structure
Giant molecular
structure
Simple molecular
structure
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a. Na , Mg , Al :
Giant metallic structure
Strong metallic bonds  high m.p
M.p increases from Na to Al as metallic
bonds becomes stronger
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Reasons : Increasing no of delocalised
electrons (*)
Increasing nuclear charge
Decreasing atomic radius
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b. Si :
Giant molecular structure
Numerous strong covalent bonds
present
Highest melting point
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c. P to Ar :
Simple molecular structure
Weak VDW forces  low m.p
Strength of VDW forces  no of electrons
P4
S8
Cl2
Ar
no of electrons
most
least
VDW force
strongest
weakest
m.p / 0C
44 119
-101 -189
highest
lowest
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Note:
Structure of P4 : tetrahedral
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Sulphur, S8
Si
Melting
Point / 0 C
Mg
Al
S
Na
P
Cl
Ar
Atomic number
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3. Electrical conductivity :
In period 3 (from Na to Ar) , elements
change from metal to non metal
a. Na , Mg , Al :
Metals with metallic structures consisting of
cations in a sea of delocalised electrons
Mobile electrons  Good conductors/high
conductivity
Conductivity increases as no of delocalised
electrons increases from Na to Al
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b. Si :
Semi metal , semi conductor / moderate
conductivity
c. P to Ar :
Non metals , exists as simple discrete
molecules
All electrons paired in covalent bonds
No free electrons or ions  non conductors/
low conductivity
Electrical
conductivity
Al
Na
Si
Atomic number
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4. Ionisation energy ( first I.E ) :
a. Generally , the I.E of elements
increase across a period due to
i) increase in nuclear charge (*)
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ii)slight decrease in atomic radius
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Note : However there are anomalies :
(1)between elements in Groups II and III ,
and
(2)between elements in Groups V and VI
b. In any period , the elements in Group VIII
( inert or noble gases ) being stable have the
highest first I.E
First I.E
Ar
P
Mg
Na
Si
Cl
S
Al
Atomic number
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5. Electronegativity :
a. Definition : electronegativity of an
element is a measure of its attraction
for bonding electrons
b. Electronegativities increases from left
to right across a period due to
increase in nuclear charge
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c. Electronegativity decreases down a
group
due to increase in shielding effect
Changes in Physical Properties of
Elements in a Period
Property
Na
Mg
Al
Si
P
S
Cl
Ar
Proton #
11
12
13
14
15
16
17
18
Ionisation energy
(kJ/mol)
500
740
580
790
1010
1000
1260
1520
Covalent bond
radius (nm)
0.15
6
0.136
0.125
0.117
0.110
0.104
0.099
0.19
2
Electronegativity
1.0
1.25
1.45
1.74
.05
2.45
2.85
Structure
Giant
Metal
-lic
Giant
molec
-ular
Simpl
e
Molec
-ular
Density (gcm-3)
0.97
1.74
2.70
2.33
1.82
2.07
gas
Gas
Melting pt (oC)
98
low
890
651
high
1117
660
high
2447
1410
high
2355
44
low
280
119
low
445
-101
low
-35
-189
low
-186
high
high
high
mode
rate
low
low
low
low
Boiling pt (oC)
Electrical
conductivity