Today we're starting a two-part look at covalent bonding — what happens when atoms share electrons instead of transferring them completely, as they do in ionic bonding.
What is a covalent bond?
A covalent bond forms when two atoms (usually non-metals) share a pair of electrons so that both nuclei are attracted to the same shared pair. This mutual attraction holds the atoms together (Talbot, p.143). A simple electrostatic picture shows two positive nuclei both pulling on the negative electron pair sitting between them — like a tug-of-war where both sides win by sharing the "rope."
In some molecules, more than one pair of electrons is shared:
- Single bond – one shared pair (e.g. H–H)
- Double bond – two shared pairs (e.g. O=O in O₂)
- Triple bond – three shared pairs (e.g. N≡N in N₂)
(Talbot, p.143)
Bond strength and length
The more electron pairs shared between two nuclei, the stronger and shorter the bond. This is because more shared electrons mean more electrostatic attraction pulling the nuclei together (Talbot, p.151).
Worked Example 1: Carbon–carbon bonds show this pattern clearly (Talbot, p.151):
Bond type Bond enthalpy (kJ mol⁻¹) Bond length (nm)
C–C (single) 348 0.154
C=C (double) 612 shorter
So going from single → double → triple bonds: enthalpy increases, bond length decreases.
Polar vs non-polar covalent bonds
Not all covalent bonds share electrons equally. If the two atoms have different electronegativities, one atom pulls the shared pair closer, creating a polar covalent bond with a slightly negative end (δ−) and slightly positive end (δ+). A classic example is I–Cl, where chlorine (more electronegative) pulls the electron density toward itself (Talbot, p.143).
We can predict bond type using electronegativity difference (ΔEN), based on the Pauling scale (Talbot, p.134):
- ΔEN < 0.5 → non-polar covalent (e.g. C–H)
- 0.5 < ΔEN < 1.7 → polar covalent
- ΔEN > 1.7 → ionic
Remember: ionic and covalent bonding are two ends of a spectrum — polar covalent bonds sit in between, since they involve partial electron transfer rather than complete transfer or perfectly equal sharing.
Worked Example 2: Predict the bonding type in Cl₂, HI, and LiF using electronegativity values (Talbot, p.135).
- Cl₂: ΔEN = 3.2 − 3.2 = 0 → non-polar covalent
- HI: ΔEN = 2.7 − 2.2 = 0.5 → borderline, essentially non-polar/weakly polar covalent
- LiF: ΔEN = 4.0 − 1.0 = 3.0 → ionic
Worked Example 3: Is the bond in H₂O polar or non-polar?
- O has EN ≈ 3.4, H has EN ≈ 2.2
- ΔEN = 3.4 − 2.2 = 1.2 → falls between 0.5 and 1.7 → polar covalent bond
This explains why water molecules have δ+ hydrogens and a δ− oxygen — a fact that becomes hugely important later when we study intermolecular forces!
Key takeaways
- A covalent bond arises from two nuclei being attracted to a shared pair (or pairs) of electrons.
- Bonds can be single, double, or triple — more shared pairs means stronger and shorter bonds.
- Electronegativity difference (ΔEN) tells us whether a bond is non-polar covalent (<0.5), polar covalent (0.5–1.7), or ionic (>1.7).
- Polar covalent bonds have unequal sharing, creating δ+ and δ− ends on the molecule.
- Ionic and covalent bonding are extremes of the same spectrum, not completely separate categories.