Introduction
In the last lesson we met the covalent bond as a shared pair of electrons holding two nuclei together. Today we go further: why are some covalent bonds polar and others not, why do double and triple bonds behave differently from single bonds, and what happens when one Lewis formula just isn't enough to describe a real molecule?
Polar vs non-polar covalent bonds
Not all shared electron pairs are shared equally. When two identical atoms bond (like H–H or Cl–Cl), the electrons are shared evenly — this is a non-polar covalent bond. When atoms of different electronegativity bond, the electron pair sits closer to the more electronegative atom, creating a polar covalent bond with a slightly negative end (δ−) and slightly positive end (δ+), as seen in I–Cl (Talbot, p.143).
We can predict bond type using electronegativity difference (Δχ), based on the Pauling scale (Talbot, p.134):
- Δχ < 0.5 → essentially non-polar covalent (e.g. C–H)
- 0.5 < Δχ < 1.7 → polar covalent
- Δχ > 1.7 → likely ionic (e.g. FrF)
Remember: ionic and covalent bonding are two ends of a spectrum, not strict categories — polar bonds sit in between, showing "covalent bonds with ionic character."
Worked example 1: Classify the H–F bond. Electronegativities: H = 2.2, F = 4.0. Δχ = 1.8 → this is close to the ionic border but conventionally treated as a very polar covalent bond.
Worked example 2: Classify the C–Cl bond. Electronegativities: C = 2.6, Cl = 3.2. Δχ = 0.6 → polar covalent.
Single, double, and triple bonds
Sometimes one shared pair isn't enough to give both atoms a stable electron arrangement, so atoms share two or three pairs. Using a simple electrostatic model (Talbot, p.143):
- O₂ contains a double bond (two shared pairs)
- N₂ contains a triple bond (three shared pairs)
More shared pairs mean stronger electrostatic attraction between the nuclei and the electrons, so:
- Bond enthalpy increases: single < double < triple
- Bond length decreases: single > double > triple
Worked example 3: Compare C–C (348 kJ mol⁻¹, 0.154 nm) with C=C (612 kJ mol⁻¹, shorter length) (Talbot, p.151). The double bond is both shorter and stronger, because more electron pairs pull the nuclei closer together with greater force.
Resonance: when one structure isn't enough
For some molecules and ions, a single Lewis formula fails to represent reality. Ozone, O₃, can be drawn with the double bond on either side of the central oxygen atom — two equally valid Lewis structures called resonance structures, linked with a double-headed arrow (not to be confused with an equilibrium sign!) (Talbot, p.186).
Neither structure alone is correct — experimentally, both O–O bonds in ozone have the same length, intermediate between a single and double bond (0.121 nm vs 0.132 nm predicted extremes). The real molecule is best described as an average, or hybrid, of the resonance structures.
Worked example 4: Why can't ozone simply be drawn as one fixed structure with one single and one double bond? Because X-ray data shows both O–O bonds are identical in length — resonance explains this by delocalising the double bond character across both bonds.
Key takeaways
- Bond polarity depends on electronegativity difference: <0.5 non-polar, 0.5–1.7 polar, >1.7 ionic.
- Polar covalent bonds have partial charges (δ+/δ−) due to unequal electron sharing.
- More shared electron pairs = shorter, stronger bonds (single < double < triple).
- Resonance structures are needed when a single Lewis formula can't explain equal bond lengths/energies (e.g. ozone).
- The true structure of a resonance molecule is a hybrid, not any one single Lewis formula.