3.4 Covalent and Coordinate Bonding
Candidates should be able to:
- Define covalent bonding, and describe covalent bonding in H₂, O₂, N₂, Cl₂, HCl, CO₂, NH₃, CH₄, C₂H₆ and C₂H₄; understand that Period 3 elements can expand their octet (e.g. SO₂, PCl₅, SF₆); describe coordinate (dative covalent) bonding, including NH₃ + HCl → NH₄⁺ and the Al₂Cl₆ molecule.
- Describe covalent bonds in terms of orbital overlap giving σ and π bonds, including how these form in H₂, C₂H₆, C₂H₄, HCN and N₂; use the concept of hybridisation to describe sp, sp² and sp³ orbitals.
- Define bond energy and bond length, and use these to compare the reactivity of covalent molecules.
1 Covalent Bonding — The Basics ▾
When two non-metal atoms combine, they share one or more pairs of electrons. A shared pair of electrons is called a single covalent bond, or a bond pair, and is represented by a single line between the atoms, e.g. Cl—Cl.
Not all outer-shell electrons are used in bonding. The pairs not involved in a bond are called lone pairs.
Covalent Bonding Across Common Molecules
The same principle — sharing electron pairs to reach a full outer shell — explains the bonding in all of the molecules below.
| Molecule | Structure | Bonding Summary |
|---|---|---|
| Hydrogen, H₂ | H—H | 1 single bond, no lone pairs |
| Oxygen, O₂ | O=O | 1 double bond, 2 lone pairs per O |
| Nitrogen, N₂ | N≡N | 1 triple bond, 1 lone pair per N |
| Chlorine, Cl₂ | Cl—Cl | 1 single bond, 3 lone pairs per Cl |
| Hydrogen chloride, HCl | H—Cl | 1 single bond, 3 lone pairs on Cl |
| Carbon dioxide, CO₂ | O=C=O | 2 double bonds, linear, 2 lone pairs per O |
| Ammonia, NH₃ | 3 × N—H | 3 single bonds + 1 lone pair on N, pyramidal |
| Methane, CH₄ | 4 × C—H | 4 single bonds, no lone pairs, tetrahedral |
| Ethane, C₂H₆ | H₃C—CH₃ | 1 C—C + 6 C—H single bonds |
| Ethene, C₂H₄ | H₂C=CH₂ | 1 C=C double bond + 4 C—H single bonds |
2 Coordinate (Dative Covalent) Bonding ▾
In most covalent bonds, each atom supplies one electron to the shared pair. But it’s possible for just one atom to provide both bonding electrons.
For example, when gaseous ammonia and gaseous boron trifluoride react, a white solid forms with the formula NH₃BF₃. The nitrogen atom in ammonia has a lone pair; the boron atom in BF₃ has an empty 2p orbital.
The same idea explains what happens when ammonia gas meets hydrogen chloride gas: the lone pair on nitrogen is donated into the empty (antibonding) orbital that appears on hydrogen as the H—Cl bond breaks heterolytically, forming the ammonium ion.
3 Hybridisation ▾
Covalent bonds form when atomic orbitals overlap. The mixing of atomic orbitals to create a new set of equivalent orbitals is called hybridisation.
sp Hybrid Orbitals
When one s orbital and one p orbital are hybridised, the result is two sp orbitals, arranged linearly (180° apart).
sp² Hybrid Orbitals
When one s orbital and two p orbitals are hybridised, the result is three sp² hybrid orbitals, arranged in a trigonal planar shape (120° apart).
sp³ Hybrid Orbitals
When one s orbital and three p orbitals are hybridised, the result is four sp³ hybrid orbitals. Each hybrid orbital has both s character and p character, and they point towards the corners of a tetrahedron, 109°28′ apart.
4 Orbital Overlap — σ and π Bonds ▾
A single bond forms when a pair of electrons is shared between two atoms. This electron pair occupies a molecular orbital, formed by the overlap of an atomic orbital on each atom (which may be s, p, or hybrid orbitals).
The electron density of a σ bond is symmetrical about the line joining the two nuclei — it looks the same from every angle around that axis.
σ and π Bonding in Specific Molecules
Every single bond is a σ bond. A double bond is always 1σ + 1π; a triple bond is always 1σ + 2π. This gives the following picture for some key examples:
| Molecule | Central Atom Hybridisation | Sigma Bonds | Pi Bonds |
|---|---|---|---|
| H₂ | 1s (no hybridisation) | 1 (H—H) | 0 |
| C₂H₆ (ethane) | sp³ at each C | 7 (1 C—C + 6 C—H) | 0 |
| C₂H₄ (ethene) | sp² at each C | 5 (1 C—C + 4 C—H) | 1 (in the C=C) |
| HCN | sp at C | 2 (1 H—C + 1 in C≡N) | 2 (in the C≡N) |
| N₂ | sp at each N | 1 (in the N≡N) | 2 (in the N≡N) |
5 Bond Energy & Bond Length ▾
Bond length is the internuclear distance between two covalently bonded atoms.
Double and triple covalent bonds are successively shorter than single covalent bonds between the same two elements, and are generally associated with higher bond energies (more shared electron pairs pull the nuclei closer together and hold them more strongly).
- Describe a covalent bond as the attraction between two nuclei and a shared pair of electrons.
- Remember a dative bond is identical to a normal covalent bond once it has formed.
- State that σ bonds are end-on/direct overlap, and π bonds are sideways overlap.
- Count a double bond as 1σ + 1π, and a triple bond as 1σ + 2π.
- Match hybridisation to shape: sp = linear, sp² = trigonal planar, sp³ = tetrahedral.
- Don’t say a double bond is “two σ bonds” — it’s one σ and one π bond, which are not identical.
- Don’t forget the donor atom supplies both electrons in a dative bond, not one each.
- Don’t describe a π bond as symmetrical about the internuclear axis — only σ bonds are.
- Don’t mix up bond energy (breaks the bond) with bond length (distance between nuclei) — higher energy usually means shorter length, not the same thing.
- Covalent bond = electrostatic attraction between two nuclei and a shared pair of electrons.
- Dative (coordinate) bond = both shared electrons come from one atom (the donor) into an empty orbital on the other (the acceptor).
- Hybridisation mixes s and p orbitals: sp (linear), sp² (trigonal planar), sp³ (tetrahedral).
- σ bonds = end-on overlap (every single bond); π bonds = sideways overlap (found in double/triple bonds).
- Bond energy = energy to break 1 mole of a bond; bond length = internuclear distance. Multiple bonds are shorter and stronger than single bonds.