A Level - Electronegativity: An Atomic Tug Of War!

Covalent bonds aren't all purely covalent. The electron pair which resides between the nuclei of two atoms, held in place by strong electrostatic forces, doesn't always stay exactly in the middle of the two atoms. 

You can imagine that each of the two nuclei exert a force of attraction upon the bonded pair. and this force is different for different elements. The ability of an atom to draw a pair of bonded electrons within a covalent bond towards itself is known as electronegativity. 

The above diagram illustrates, in a silly way, that some atoms are much better at drawing electrons towards them than others. This means that the bonding electrons spend more of their time closer to the more electronegative atom within the bond, creating a partial charge (dipole):

Where the electrons have been drawn towards the more electronegative atom, a partial negative charge is created, represented here by the δ- symbol. In turn, where the electrons have been drawn away from one atom, this leaves it with a partial positive charge, denoted by the δ+ symbol. Collectively, these two charges make up a dipole. 

Electronegativity is influenced by three main factors:

• Atomic Radius - the smaller an atom, the closer the bonding electrons will be to the nucleus and therefore the greater the electrostatic attraction will be. 
• Nuclear Charge - the more protons in the nucleus, the greater the charge of the nucleus and the electrostatic attraction. 
• Shielding - Inner electron shells shield the nuclear charge from the bonding pair, reducing the effective nuclear charge (attraction felt by the bonding pair of electrons).

The electronegativity of elements increases as you go from left to right and from the bottom to the top of the periodic table, like so:

Note that group 0 are not accounted for with electronegativity as they rarely bond to anything. 

The electronegativity of different elements was quantified by American Chemist Linus Pauling, in the Pauling scale, as shown below:

source

So, now we return to the start of this post, where I stated that most covalent bonds are not purely covalent. Bonds actually fall onto one of three categories:

• pure covalent
• polar covalent 
• ionic

So, in summary:

• Electronegativity is the ability of an atom to draw a bonding pair of electrons from a covalent bond towards itself.
• Electronegativity increases progressively across and up the periodic table. 
• A difference in electronegativity can create a polar bond with a dipole.