Electronegativity
Definition
of Electronegativity
Ability of an atom to attract the bonding electrons in a covalent bond.
Electronegativity trends in the Periodic table
Whilst you will not be required to memorise the electronegativity of atoms you will have to be able to describe patterns in electronegativity in the periodic table.
· Across the periodic table electrons are added to the same shell, therefore shielding stays the same but the number of protons increases, therefore the atom has a greater ability to draw the bonding electrons in a covalent bond thus the electronegativity increases.
·
Down the periodic table shielding of inner electrons
increases, therefore attraction of the outer electrons decreases. Therefore
the electronegativity decreases.
| Li 1.0 |
Be 1.5 |
B 2.0 |
C 2.5 |
N 3.0 |
O 3.5 |
F 4.0 |
| Na 0.9 |
Mg 1.2 |
Al 1.5 |
Si 1.8 |
P 2.1 |
S 2.5 |
Cl 3.0 |
Elements
with a higher electronegativity are more able to draw electrons towards them.
Same
electronegativity
The
electrons are shared equally and the bond is referred to as non-polar.
Example:
I2 Iodine I –
I
Different electronegativity
The
electrons are NOT shared equally, this means one atom has a greater ‘concentration’
of electrons, thus giving it a slight negative charge. The atoms with a lower
electronegativity will therefore have a lower ‘concentration’ of electrons and
gains a slight positive charge. This is referred to as polar and results
in a permanent dipole (uneven charge distribution).
Note:
d is used to represent a
‘slight’ charge
Example:
Mg – O
d+ d-
Mg
has an electronegativity of 1.2, however Oxygen has an electronegativity of
3.5 and therefore the bond is polar and results in a permanent dipole.
CAUTION
Whilst
each bond separately may be polar the overall polarity may not!
O = C = O
The element has polar bonds, however it is linear with a bond angle of
180°. It is therefore symmetrical so the charges
cancel and overall is non-polar.
When looking
at whether a molecule is polar always consider its shape.