An element's electronegativity measures its atom's tendency to attract the shared pair of electrons towards itself in a covalent bond.
  • Consider the \(HCl\) molecule. The hydrogen and chlorine atoms give one electron each to form a covalent bond.
  • A chlorine atom has a higher electronegativity, and hence it pulls the shared electrons towards itself more strongly than hydrogen. Thus, the bonding electrons are left with chlorine, forming \(H^+\) and \(Cl^–\) ions when the bond breaks.
Relative electronegativity of \(H\) and \(Cl\).

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  • Electronegativity is determined by various experimental data, including bond energy, ionisation potential, electron affinity, etc.
  • The Pauling scale is the most used scale for determining electronegativity, which predicts the type of bonding (ionic or covalent) between atoms in a molecule.
Some of the elements' electronegativity is listed below.
F = \(4.0\), Cl = \(3.0\), Br = \(2.8\), I = \(2.5\), H = \(2.1\), Na = \(1\)
  • If the electronegativity difference between two elements is \(1.7\), then the bond has \(50\)% ionic character and \(50\)% covalent character.
  • The bond is considered more covalent if the difference is less than \(1.7\).
  • If the difference is more than \(1.7\), the bond is considered more ionic.
  • From left to right in the periodic table, the electronegativity increases due to an increase in nuclear charge, which in turn attracts electrons more strongly.
  • Conversely, on moving down a group, the electronegativity of the elements decreases because of the increased number of valence shells.
Periodic Property
In Periods
In Groups
Atomic radiusDecreasesIncreases
Ionic radiusDecreasesIncreases
Ionisation energyIncreasesDecreases
Electron affinityIncreasesDecreases