Covalent bonding and coordinate (dative covalent) bonding

· Covalent bonding = electrostatic attraction between the nuclei of two atoms and a shared pair of electrons.
· A single covalent bond contains 1 shared pair of electrons.
· A double bond contains 2 shared pairs of electrons.
· A triple bond contains 3 shared pairs of electrons.
· Covalent bonds usually form between non-metal atoms.
· In exam answers, always mention attraction between nuclei and shared pair of electrons, not just “sharing electrons”.

This diagram shows a sigma bond formed by direct overlap along the internuclear axis. It supports the definition of covalent bonding as attraction between nuclei and a shared electron pair. Source

Covalent molecules you must know

· H₂: single covalent bond, H–H; each H shares 1 electron to achieve a full first shell.
· Cl₂: single covalent bond, Cl–Cl; each Cl completes an outer shell of 8 electrons.
· O₂: double covalent bond, O=O; 2 shared pairs.
· N₂: triple covalent bond, N≡N; 3 shared pairs, very strong bond.
· HCl: single covalent bond, H–Cl; bond is polar due to higher electronegativity of Cl.
· CO₂: O=C=O; carbon forms two double bonds.
· NH₃: three N–H single bonds and one lone pair on nitrogen.
· CH₄: four C–H single bonds.
· C₂H₆: all single bonds; contains a C–C single bond and C–H bonds.
· C₂H₄: contains a C=C double bond; one σ bond and one π bond between the carbon atoms.

These diagrams show how outer-shell electrons are arranged in covalent molecules. They are useful for practising dot-and-cross style reasoning before drawing exam diagrams. Source

Expanded octets in Period 3

· Elements in Period 3 can expand their octet because they can form compounds with more than 8 electrons around the central atom.
· Required examples: SO₂, PCl₅, SF₆.
· PCl₅: phosphorus forms 5 covalent bonds.
· SF₆: sulfur forms 6 covalent bonds.
· Do not apply the octet rule too rigidly to Period 3 central atoms in these examples.
· In dot-and-cross diagrams, show the correct number of bonding pairs around the central atom, even if this exceeds 8 electrons.

Coordinate / dative covalent bonding

· Coordinate bond / dative covalent bond = a covalent bond in which both electrons in the shared pair come from one atom.
· Once formed, a coordinate bond behaves like an ordinary covalent bond.
· The electron-pair donor must have a lone pair.
· The electron-pair acceptor must be electron-deficient or have an available orbital.
· Use an arrow from the atom donating the lone pair to the atom accepting it.
· Example: NH₃ + HCl → NH₄Cl; ammonia donates its lone pair to H⁺, forming NH₄⁺.
· In NH₄⁺, one N–H bond is originally coordinate, but all four N–H bonds are equivalent after formation.
· Example: Al₂Cl₆ has coordinate bonding from a lone pair on Cl to an electron-deficient Al atom.

This image helps visualise the ammonium ion formed when ammonia donates a lone pair to H⁺. It is directly relevant to coordinate bonding in NH₄⁺. Source

This diagram shows the dimeric structure of aluminium chloride, Al₂Cl₆. The bridging chlorine atoms use lone pairs to form coordinate bonds to electron-deficient aluminium atoms. Source

Sigma and pi bonds

· σ bond = formed by direct overlap of orbitals between bonding atoms.
· π bond = formed by sideways overlap of adjacent p orbitals above and below the σ bond.
· A single bond is always 1 σ bond.
· A double bond = 1 σ bond + 1 π bond.
· A triple bond = 1 σ bond + 2 π bonds.
· Required examples: H₂, C₂H₆, C₂H₄, HCN, N₂.
· H₂: one σ bond from overlap of two 1s orbitals.
· C₂H₆: C–C and C–H bonds are all σ bonds.
· C₂H₄: C=C contains one σ bond and one π bond.
· HCN: C≡N contains one σ bond and two π bonds.
· N₂: N≡N contains one σ bond and two π bonds.
· π bonds restrict rotation, so molecules with double bonds often have fixed planar regions.

Hybridisation: sp, sp² and sp³

· Hybridisation = mixing of atomic orbitals to form new hybrid orbitals used in bonding.
· sp³: one s orbital + three p orbitals → four sp³ orbitals.
· sp³ examples: CH₄, C₂H₆, central N in NH₃.
· sp²: one s orbital + two p orbitals → three sp² orbitals, with one unhybridised p orbital left for π bonding.
· sp² example: carbon atoms in C₂H₄.
· sp: one s orbital + one p orbital → two sp orbitals, with two unhybridised p orbitals left for π bonding.
· sp examples: carbon in HCN and atoms involved in triple bonds such as N₂.
· Link hybridisation to bonding: sp³ = single-bond framework, sp² = double-bond framework, sp = triple-bond framework.

This diagram shows how four equivalent sp³ hybrid orbitals form a tetrahedral bonding arrangement. It is useful for understanding bonding in CH₄ and C₂H₆. Source

Bond energy, bond length and reactivity

· Bond energy = energy required to break one mole of a particular covalent bond in the gaseous state.
· Higher bond energy means a bond is generally stronger and harder to break.
· Bond length = internuclear distance between two covalently bonded atoms.
· Shorter bonds are usually stronger.
· Multiple bonds are generally shorter and stronger than single bonds between the same atoms.
· Reactivity comparisons often depend on how easily a bond breaks: lower bond energy = more reactive bond.
· Exam comparison example: N₂ is relatively unreactive because the N≡N triple bond has very high bond energy.