Electronegativity

Electronegativity is a fundamental chemical property describing the tendency of an atom to attract shared electrons or electron density when participating in a chemical bond. Represented by the symbol χ, it provides a quantitative means of interpreting bond behaviour, particularly the distribution of electron density and the resulting polarity of covalent or partially ionic bonds. As a comparative measure, electronegativity helps predict the character of chemical bonding across the full spectrum from purely covalent to predominantly ionic interactions. Its conceptual opposite is electropositivity, describing an atom’s propensity to donate electrons.
Electronegativity is influenced by two principal factors: the effective nuclear charge, which draws electrons closer to the nucleus, and the distance of valence electrons from the nucleus, shaped by electron-shell structure and shielding by inner electrons. Atoms with high effective nuclear charge and minimal shielding typically exhibit high electronegativity. Conversely, large atomic radii and significant shielding reduce the nuclear pull experienced by valence electrons, leading to lower electronegativity values.

Historical Development

The term electronegativity was introduced in 1811 by the Swedish chemist Jöns Jacob Berzelius, although earlier chemists such as Amedeo Avogadro had explored related ideas. A fully quantified scale, however, did not emerge until 1932 when Linus Pauling developed the first widely adopted electronegativity scale. Pauling’s system, grounded in bond-energy considerations within valence bond theory, provided a coherent numerical framework that correlated well with observed chemical properties.
Because electronegativity cannot be measured directly, various methods have been devised to infer it from atomic or molecular data. Despite numerical differences between scales, all approaches reveal consistent periodic trends: electronegativity generally increases across a period and decreases down a group. In its most familiar form—the Pauling scale—fluorine is the most electronegative element (χ ≈ 3.98) and caesium among the least (χ ≈ 0.79). These values represent relative tendencies and assume approximate transferability across different chemical environments.
Electronegativity correlates strongly with ionisation energy and, for many elements, with electron affinity. Although variations arise according to bonding context, these correlations support the broader interpretation of electronegativity as a measure of an atom’s electron-attracting power.

Pauling’s Method

Pauling’s scale is founded on observations that heteronuclear bonds (A–B) often possess greater bond dissociation energies than would be predicted solely from the average of homonuclear bonds (A–A and B–B). This excess stabilisation is attributed to partial ionic character arising from unequal electronegativity. By analysing the difference between the actual dissociation energy and the expected geometric or arithmetic mean of related homonuclear bonds, Pauling derived an expression for the electronegativity difference:
∣χA−χB∣=EAB−EAA+EBB2|\chi_A – \chi_B| = \sqrt{E_{AB} – \frac{E_{AA} + E_{BB}}{2}}∣χA​−χB​∣=EAB​−2EAA​+EBB​​​
where E values are dissociation energies (commonly expressed in electronvolts). Because the scale depends on differences rather than absolute values, Pauling assigned hydrogen an initial reference value and calculated other values relative to it. Later refinements by A. L. Allred in 1961 incorporated expanded thermodynamic data, producing the set of Pauling electronegativity values still widely used.
The method provides a semiempirical yet theoretically grounded interpretation of bond polarity. It works most effectively for single covalent bonds and becomes less accurate in environments where bond multiplicity or conjugation significantly modifies energy contributions.

Mulliken Electronegativity

Robert S. Mulliken proposed an alternative definition based on atomic energy characteristics rather than bond energies. He argued that the electronegativity of an atom can be approximated by the arithmetic mean of its first ionisation energy (Ei) and electron affinity (Eea):
χ=Ei+Eea2\chi = \frac{E_i + E_{ea}}{2}χ=2Ei​+Eea​​
This formulation, sometimes referred to as absolute electronegativity, carries physical units such as electronvolts or kilojoules per mole. Conversion factors are used to align Mulliken values with the Pauling scale for comparative convenience. Because it depends on electron affinity, the Mulliken method is limited to elements for which such data are available or reliably estimated.
Conceptually, Mulliken electronegativity is closely related to the chemical potential, with higher values indicating stronger tendencies to attract electrons. The approach aligns well with modern frontier molecular orbital theory and provides a useful alternative in theoretical and computational chemistry.

Periodic Trends and Transferability

Electronegativity displays characteristic periodic behaviour. Across a period, increasing nuclear charge draws valence electrons closer, raising electronegativity. Down a group, expanding atomic radius and increased shielding reduce the effective nuclear pull, decreasing electronegativity. These trends underpin many patterns in chemical reactivity, such as the progressive increase in oxidising power toward the top right of the periodic table.
Although electronegativity values can vary with molecular environment—especially in cases of unusual oxidation states or coordination—chemists often treat the property as broadly transferable. This assumption permits the use of electronegativity values in predicting molecular polarity, assessing bond type and estimating approximate bond energies. Such predictions offer qualitative insight into molecular behaviour despite expected uncertainties of around 10 per cent in energy estimates.

Significance in Chemical Bonding

Electronegativity is widely used to interpret:

• Bond polarity, determined by differences in electronegativity between bonded atoms.
• Bond type, with large differences favouring ionic character and small differences favouring covalent character.
• Reactivity patterns, such as nucleophilic and electrophilic tendencies.
• Molecular dipole moments, which depend on both charge separation and molecular geometry.

Because bond polarity influences intermolecular forces, electronegativity also helps explain physical properties such as solubility, boiling points and crystal structures.