3.4 Covalent bonding and coordinate (dative covalent) bonding

A covalent bond is the electrostatic attraction between the positive nuclei of two atoms and a shared pair of outer-shell electrons. The atoms share rather than transfer electrons, allowing the specified non-metal atoms to reach a more stable outer-electron arrangement.

One shared pair gives a single covalent bond; two shared pairs give a double bond; three shared pairs give a triple bond. Use the outer-shell electrons to account for the shared pairs in the specified molecules, including H₂, O₂, N₂, Cl₂, HCl, CO₂, NH₃, CH₄, C₂H₆ and C₂H₄.

In a dot-and-cross representation, show only the relevant outer-shell electrons, use dots and crosses to distinguish the two atoms, and place shared electrons as pairs between the nuclei. Keep lone pairs and shared pairs visible when they are needed to justify the structure.

For the supported Period 3 cases SO₂, PCl₅ and SF₆, the central atom may have an expanded octet. A coordinate bond is still a shared pair, but both electrons originate from one donor atom: use NH₃ + H⁺ → NH₄⁺ and Al₂Cl₆ only at this source-supported level. Orbital overlap, hybridisation, bond energy and bond length belong to neighbouring objectives.

A covalent bond forms when atomic orbitals containing unpaired electrons overlap to form a molecular orbital containing the bonding pair. Greater effective overlap gives a stronger bond. The overlap model describes the shared electron pair without changing the covalent-bond definition.

A sigma (σ) bond forms by direct end-on overlap along the line joining the two nuclei, with electron density centred between them. A pi (π) bond forms by sideways overlap of adjacent p orbitals, placing electron density above and below the plane of the σ bond.

A double bond contains one σ and one π bond; a triple bond contains one σ and two mutually perpendicular π bonds. Use the supported cases: C₂H₄ is sp² at each carbon with one C–C σ and one π bond; C₂H₂/HCN/N₂ use sp hybridisation with one σ and two π components in the triple-bond region.

Hybridisation mixes orbitals from the same shell: sp³ gives four hybrid orbitals, sp² gives three, and sp gives two. Apply the source-supported mapping to tetrahedral, trigonal-planar and linear arrangements, while keeping detailed molecular-shape treatment for the neighbouring topic.

Bond energy is the energy required to break one mole of a particular covalent bond in the gaseous state, measured in kJ mol⁻¹. Bond length is the internuclear distance between two covalently bonded atoms.

Greater attraction between bonding electrons and nuclei pulls atoms closer, giving a shorter bond length and a stronger bond. Greater bond energy therefore indicates a stronger covalent bond.

Bond multiplicity changes electron density between nuclei: triple bonds are generally shortest and strongest, double bonds intermediate, and single bonds longest and weakest. Use supplied values and units when comparing bonds.

Use bond polarity, bond strength and bond type (σ/π) as the source-supported factors when comparing covalent-molecule reactivity. A longer, weaker bond may require less energy to break, but do not treat one factor as a universal complete reactivity prediction.