Introduction
Why do atoms bother bonding at all? The short answer: bonding lowers the overall energy of a system, giving atoms a more stable electron arrangement (often resembling a noble gas). But how atoms bond depends heavily on what kinds of atoms are involved. In this lesson, we'll link electronegativity — a concept you've already met — to the three main types of chemical bonding: ionic, covalent, and metallic.
The bonding spectrum
Instead of thinking of ionic and covalent bonding as two totally separate boxes, it's more accurate to picture a continuous spectrum, with pure covalent bonding at one end and pure ionic bonding at the other. Polar covalent bonds sit in between — they are covalent bonds that have gained some "ionic character" because electrons are shared unequally (Talbot, p. 134).
Using electronegativity differences
We can estimate where a bond sits on this spectrum by calculating the difference in electronegativity (ΔEN) between the two bonded atoms:
- ΔEN < 0.5 → essentially non-polar covalent (e.g. the C–H bond)
- 0.5 ≤ ΔEN ≤ 1.7 → polar covalent
- ΔEN > 1.7 → likely ionic (e.g. francium fluoride, FrF)
(Talbot, p. 134)
Remember: these numbers are guidelines, not strict laws of nature — real bonding is more subtle, and chemists still study unusual cases (see the box below).
Worked examples
Example 1: Cl₂ EN(Cl) = 3.2 for both atoms ΔEN = 3.2 − 3.2 = 0 → Non-polar covalent bond, Cl–Cl (Talbot, p. 135)
Example 2: HI EN(H) = 2.2, EN(I) = 2.7 ΔEN = 2.7 − 2.2 = 0.5 → Borderline, but classified as polar covalent (Talbot, p. 135)
Example 3: LiF EN(Li) = 1.0, EN(F) = 4.0 (approx