Electronegativity is a measure of how strongly atoms attract bonding electrons to themselves.

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If you know the electronegativity of each atom in a diatomic molecule, you can predict how the bond will polarize. For example, on the Pauling Scale, hydrogen's electronegativity is 2.18 and chlorine's is 3.16. This results in a polar covalent bond, with hydrogen slightly positively charged and chlorine slightly negatively charged.

Its symbol is the Greek letter chi:


The higher the electronegativity, the greater an atom's attraction for electrons.

The most electronegative element is fluorine, followed by oxygen, chlorine and nitrogen.

Atoms with high electronegativity form negative ions and, when covalently bonded to atoms of lower electronegativity, have a greater share of the electrons than the other atom. The result of this is a small negative charge on the more electronegative atom and a small positive charge on the atom it is bonded to. Such a bond is described as a polar covalent bond.

The least electronegative (most electropositive) element is cesium, followed by rubidium, potassium and barium. These elements have a very strong tendency to form positive ions.

In general, electronegativity of elements increases as you look from the top of a group to the bottom and as you look from left to right across a period.

Electronegativity is a relative scale - it is calculated rather than measured. Various scales of electronegativity have been devised - for example the Pauling Scale.

Linus Pauling assigned fluorine's electronegativity as 4, and then calculated the electronegativities of other elements relative to this number using bond energies. Fluorine is now assigned an electronegativity of 3.98, because this value gives the best internal consistency when calculations are performed over a wide range of compounds.

Another scale is the Mulliken scale, which bases electronegativity values on the following equation:

Electronegativity = 0.5 x (Ionization Potential + Electron Affinity)

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