Electronegativity

Electronegativity is how strongly an atom pulls bonded electrons toward itself. It is the single most useful concept for predicting whether a bond will be covalent or ionic, where a molecule’s partial charges will sit, and how reactive atoms are likely to be in chemical reactions. The Pauling scale, introduced by Linus Pauling in 1932, runs from 0.7 (cesium) to 3.98 (fluorine). Once you know an atom’s electronegativity, you can predict a remarkable amount about its chemistry without doing any further work.

The Pauling Scale

Linus Pauling defined electronegativity in 1932 by comparing bond energies. He noticed that an A-B bond between atoms of different electronegativity is stronger than the average of A-A and B-B bonds — the extra strength is the ionic character of the bond. From the bond-energy difference he derived a dimensionless scale anchored so that hydrogen has electronegativity 2.20.

Selected values on the modern Pauling scale:

Element	Electronegativity	Element	Electronegativity
Fluorine (F)	3.98 (highest)	Hydrogen (H)	2.20
Oxygen (O)	3.44	Carbon (C)	2.55
Chlorine (Cl)	3.16	Sodium (Na)	0.93
Nitrogen (N)	3.04	Potassium (K)	0.82
Bromine (Br)	2.96	Cesium (Cs)	0.79 (lowest)

Other electronegativity scales exist — Mulliken (averaging ionization energy and electron affinity), Allred-Rochow (based on effective nuclear charge), Allen (spectroscopic). They differ in details but agree on the basic trends. Pauling is the most widely used.

The Two Trends on the Periodic Table

Electronegativity follows two clean trends across the periodic table:

• Increases across a period (left to right). Atoms in the same row have the same number of electron shells, but the nuclear charge increases as you go right. More protons pulling on the same number of shells means a stronger pull on bonded electrons. Sodium (0.93) → magnesium (1.31) → … → chlorine (3.16). The jump from left to right is large.
• Decreases down a group (top to bottom). Atoms in the same column have similar valence-electron arrangements, but more electron shells are added as you go down. The bonded electrons sit farther from the nucleus and are also shielded by inner-shell electrons. So the pull weakens. Fluorine (3.98) → chlorine (3.16) → bromine (2.96) → iodine (2.66).

Combining the two trends, the most electronegative elements sit in the top right (excluding the noble gases, which usually do not form bonds): fluorine, oxygen, chlorine, nitrogen. The least electronegative sit at the bottom left: cesium, francium, rubidium.

Predicting Bond Type from Electronegativity Difference

The most useful application of electronegativity is predicting whether a bond between two atoms will be nonpolar covalent, polar covalent, or ionic. The rule of thumb based on electronegativity difference (Δχ):

Δχ range	Bond type	Example
0 to 0.4	Nonpolar covalent	H-H (Δχ = 0), C-H (Δχ = 0.35)
0.5 to 1.7	Polar covalent	O-H (Δχ = 1.24), C-O (Δχ = 0.89)
1.8 and above	Ionic	Na-Cl (Δχ = 2.23), K-F (Δχ = 3.16)

The 1.7 cutoff is a convention, not a hard rule. In reality, bond character lies on a continuum from pure covalent to pure ionic, and many bonds (like HF, Δχ = 1.78) sit awkwardly at the boundary. The general lesson stands: bigger electronegativity difference means more ionic character.

Polar Molecules and Dipoles

In polar covalent bonds, the more electronegative atom carries a partial negative charge (\\(\\delta^-\\)), and the less electronegative atom carries a partial positive charge (\\(\\delta^+\\)). The result is a bond dipole. The dipole’s magnitude is roughly proportional to the electronegativity difference.

Whether a whole molecule has a net dipole depends on geometry. Water (H-O-H) has two polar O-H bonds, and the molecule is bent rather than linear, so the dipoles add to a net molecular dipole. Carbon dioxide (O=C=O) also has polar C-O bonds, but the molecule is linear and the two dipoles point in opposite directions — they cancel, so CO₂ has no net dipole. Geometry matters.

Why Electronegativity Matters in Chemistry

• Predicting reactivity. Electronegative atoms attract electrons, which is why the most reactive nonmetals (F, O, Cl) are the most electronegative.
• Understanding acid strength. The acidity of HX (where X is a halogen) is influenced by the H-X bond strength and the stability of X⁻ — both relate to electronegativity.
• Drawing Lewis structures. Formal charges and resonance structures depend on knowing which atom is more electronegative.
• Predicting intermolecular forces. Hydrogen bonding requires a hydrogen bonded to N, O, or F — all highly electronegative.
• Organic chemistry mechanisms. Nucleophiles attack electrophilic (electron-poor) sites; electrophiles attack nucleophilic (electron-rich) sites. Both concepts trace back to electronegativity gradients within molecules.

Related study notes: Periodic Table, Chemical Bonding, Avogadro’s Number, Mole Concept.

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