Electronegativity

The definition of electronegativity is “the ability of an atom to attract a shared pair of electrons in a covalent bond”.

The first thing to learn is that electronegativity increases (strengthens) as you go along the periods of the periodic table and decreases (weakens) when you go down the groups, so the top right of the periodic table contains the elements with the higher electronegativities and the bottom left corner are the elements with the lower electronegativities.

Now, if we look at the periodic table, we can state that fluorine is the most electronegative element and francium is the least.

You are probably wondering why helium isn’t the most electronegative element and the answer is that noble gases are not included as they all have a full outer shell of electrons and therefore do not form a covalent bond with another element.

Electronegativity can also be used to find out the type of bond formed between two atoms by calculating the difference in value between the electronegativites of the two atoms:

• If the difference is large then the bonding is most likely to be ionic.
• If the difference is small then the bonding is most likely to be polar covalent.
• If there is no difference then the bonding will be pure covalent.

So the two elements that show the biggest difference in electronegativity are fluorine and francium and we would expect an ionic bond to form if these two elements came together.

Another important point to know is that metal elements tend to have low electronegativity values which indicates that they only have weak attractions for electrons.

And non-metal elements tend to have high electronegativity values indicating that they have strong attractions for electrons and they will form ionic bonds.

Finally, below shows some electronegativity values in relation to the periodic table and also some worked examples showing the electronegativity difference between atoms and the type of bonding formed:

Worked examples (using the table above)

• NH₃ – electronegativity difference of 3.0 – 2.1 = 0.9 (small). Type of bonding – polar covalent.
•  CaCl₂ – 3.0 – 1.0 = 2.0 (large). Type of bonding – ionic.
• BrCl – 3.0 – 2.8 = 0.2 (small). Type of bonding – polar covalent.
• NaCl – 3.0 – 0.9 = 2.1 (large). Type of bonding – ionic.

 http://en.wikipedia.org/wiki/File:Periodic_table.svg