IB Chemistry home > Syllabus 2016 > Periodicity > Electronegativity

Syllabus ref: 3.2

The table of electronegativity values has been calculated from a knowledge of the element's ionisation energy and electron affinity.

Definition

The electronegativity of an element is a measure of its ability to attract electrons from within a covalent bond.

It is an artificial scale from 0.7 to 4.0, created by combining the ionisation energy and the electron affinity of the elements. There are several versions of electronegativity although the most commonly used is that developed by Linus Pauling.

Only elements with electronegativity values of 3.0 and above are considered electronegative.


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Descending a group

The electronegativity decreases on descending a group. This makes common sense as electronegativity is derived from the ionisation energy and electron affinity, which both decrease on descending a group.


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Crossing a period

The electronegativity increases from left to right across a period, the most electronegative elements are nitrogen, oxygen and fluorine (in increasing order)


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Comparing electronegativities

Information about the nature of chemical bonding can be predicted by looking at the electronegativities of the bonding atoms. The greater the difference in electronegativity the greater the polarisation of the bond.

Carbon has an electronegativity of 2.5 and hydrogen 2.1. These two values are relatively close to one another, so bonds formed between carbon and hydrogen are considered non-polar.

However, bonds formed between oxygen (3.5) and hydrogen (2.1) are highly polarised, (This is the basis behind intermolecular hydrogen bonding) whereas bonds formed between oxygen and elements with a lower electronegativity than hydrogen are polarised to such an extent that they are ionic.

In the same way the nature of the chlorides of the elements may be predicted.

Aluminium chloride formed from aluminium (electronegativity = 1.5) and chlorine (3.0) falls right on the borderline of ionic/covalency. In fact, aluminium chloride is ionic at low temperatures, but covalent as the temperature rises. We consider it to be a covalent compound with highly polarised bonds. (the molecule itself is non-polar due to symmetry).

These 'rules' are not empirical in that they cannot be used mathematically. There is no 'magic' difference in electronegativity at which the change between ionic and covalent occurs, if this were the case then metal nitrides and sulfides would not be ionic. It is more a case of getting a feel for the concept and using the electronegativity values as a rough guide.


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