3.4 Covalent bonding and coordinate (dative covalent) bonding
- Syllabus
- 9701–2028–2029
- Topic
- 3.4
- Level
- AS
Learning objectives
3.4.1Covalent bonding examples• Define covalent bonding as attraction between two nuclei and a shared electron pair.- Describe covalent bonding in H2, O2, N2, Cl2, HCl, CO2, NH3, CH4, C2H6 and C2H4- Period 3 expanded octets: SO2, PCl5 and SF6- Coordinate bonding: NH3 + HCl -> NH4+ and Al2Cl63.4.2Sigma and pi bonds• Describe covalent bonds by orbital overlap.- Sigma bonds: direct overlap between bonding atoms- Pi bonds: sideways overlap of adjacent p orbitals above and below the sigma bond- Apply to H2, C2H6, C2H4, HCN and N2; use sp, sp2 and sp3 hybridisation3.4.3Terms• Gaseous state- (a) define the terms:- bond energy as the energy required to break one mole of a particular covalent bond in the- bond length as the internuclear distance of two covalently bonded atoms- (b) use bond energy values and the concept of bond length to compare the reactivity of covalent molecules
Covalent bonds attract two nuclei to shared electron pairs
A covalent bond is the electrostatic attraction between the nuclei of two atoms and a shared pair of electrons. One, two or three shared pairs form a single, double or triple bond respectively.
| Shared-pair pattern | Required examples |
|---|---|
| one single bond | H₂, Cl₂, HCl |
| double or triple bond | O₂ has a double bond; N₂ has a triple bond |
| several bonds around a central atom | CO₂ has two C=O bonds; NH₃ has three N–H bonds; CH₄ has four C–H bonds |
| carbon–carbon bonding | C₂H₆ has a C–C single bond; C₂H₄ has a C=C double bond |
An octet is not an absolute limit for the central Period 3 atom in the specified cases: sulfur has an expanded octet in SO₂ and SF₆, while phosphorus has an expanded octet in PCl₅. This exception does not change the covalent-bond definition.
A coordinate (dative covalent) bond is a covalent bond in which both electrons in the shared pair come from the same donor atom. The donor must have a lone pair and the acceptor must have an available orbital.
NH3(g)+HCl(g)⟶NH4Cl(s)
In this reaction, nitrogen donates its lone pair to hydrogen to form the fourth N–H bond in NH₄⁺; Cl⁻ is the counter-ion. In 2AlCl₃ ⇌ Al₂Cl₆, lone pairs on bridging chlorine atoms are donated to electron-deficient aluminium atoms, forming two coordinate bonds in the dimer.
The arrow used to show a coordinate bond points from the lone-pair donor to the acceptor, but detailed dot-and-cross drawing belongs to section 3.7. Once formed, the shared pair is attracted by both nuclei; its origin is what makes the bond coordinate.
Sigma and pi bonds come from different orbital overlaps
A σ bond forms by direct overlap of orbitals along the internuclear axis. A π bond forms by sideways overlap of adjacent unhybridised p orbitals, giving electron density above and below the σ-bond axis.
Every single bond is one σ bond. A double bond contains one σ and one π bond; a triple bond contains one σ and two π bonds formed from two perpendicular pairs of p orbitals.
s+3p→4 sp3s+2p→3 sp2s+p→2 sp
Hybridisation mixes orbitals on the same atom. sp³ leaves no unhybridised p orbital, sp² leaves one for one π bond, and sp leaves two for two π bonds.
| Molecule | Relevant hybridisation | σ and π description |
|---|---|---|
| H₂ | H 1s orbitals | direct 1s–1s overlap gives one σ bond |
| C₂H₆ | each C is sp³ | all C–H and C–C bonds are σ |
| C₂H₄ | each C is sp² | C=C contains one σ and one π bond |
| HCN | C and N are sp | H–C is σ; C≡N contains one σ and two π bonds |
| N₂ | each N is sp | N≡N contains one σ and two π bonds |
Do not count a double bond as two σ bonds or hybridise the p orbital needed for a π bond. Molecular shapes and bond angles are developed separately in section 3.5.
Bond energy and bond length help compare covalent reactivity
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.
A larger bond energy means a stronger bond and more energy is required to break it. For the same pair or a comparable series of atoms, a shorter bond generally reflects stronger attraction between nuclei and bonding electrons and is associated with a stronger bond.
| Molecule | H–X bond length / pm | bond energy / kJ mol⁻¹ |
|---|---|---|
| HCl | 127 | 431 |
| HBr | 141 | 366 |
| HI | 161 | 299 |
From HCl to HI, the halogen atom becomes larger, the H–X bond becomes longer and attraction across the bond weakens. Less energy is therefore needed to break H–I, so HI is the most reactive of these hydrogen halides in a comparison where breaking the H–X bond controls the reaction.
Bond energy and length compare the ease of breaking a specified bond; they are not universal predictors of a whole molecule's reactivity. Always identify which bond must break and use the supplied values and conditions.