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A covalent bond is, at A-Level, defined as a shared pair of electrons attracted simultaneously to the nuclei of two atoms. This simple electrostatic picture — formalised by G. N. Lewis in 1916 — is the foundation of almost all molecular chemistry. In this lesson we go beyond the bare definition: we treat single, double and triple bonds; we introduce sigma (σ) and pi (π) character qualitatively to rationalise why C=C is not exactly twice C–C in enthalpy; we work through dot-cross diagrams for the canonical AQA examples (H₂, Cl₂, HCl, O₂, N₂, CO₂, CH₄, NH₃, H₂O); and we extend to dative (coordinate) covalent bonds where both electrons in the shared pair originate from the same atom — the key concept needed to handle NH₄⁺, H₃O⁺, the Al₂Cl₆ dimer and, looking ahead, transition-metal complexes. We close with bond-length and bond-enthalpy trends, and with the major exceptions to the octet rule (BF₃ as electron deficient; SF₆ via "expanded octet" — flagged here as a useful A-Level shorthand that undergraduate treatment refines).
Spec mapping (AQA 7405): This lesson anchors §3.1.3.2 (covalent and dative covalent bonding) within the broader §3.1.3 "Bonding" topic. It is the conceptual partner of §3.1.3.1 (ionic bonding — lesson 0 of this course) and feeds directly into §3.1.3.6 (shapes of molecules and ions — lesson 2), §3.1.3.5 (electronegativity and bond polarity — lesson 3), §3.1.3.7 (forces between molecules — lesson 4) and the whole of §3.3 (organic chemistry), which presupposes fluency with σ/π skeletons and electron-pair counting. Refer to the official AQA specification document for the exact wording of each section.
Assessment objectives: Definitions of covalent bond, dative covalent bond and lone pair are AO1 recall items, frequently asked as one- or two-mark openers. Drawing dot-cross diagrams for CO₂, NH₃, H₂O, NH₄⁺ and AlCl₃ dimers is AO2 — application of the electron-pair-counting procedure to specified molecules. Comparing bond enthalpies and bond lengths (e.g. C–C vs C=C vs C≡C, or rationalising why C=C is less than 2 × C–C) and explaining the dimerisation of AlCl₃ in the gas phase test AO3 — analysis and synthesis. Expect Paper 2 (Section A, inorganic and physical) to combine all three on a single multi-part question.
Key Definition: A covalent bond is a shared pair of electrons between two atoms, where the bonding pair of electrons is attracted simultaneously to the nuclei of both atoms.
Covalent bonding typically occurs between non-metal atoms. Each atom contributes one electron to the shared pair (except in dative bonds — see below). The electrostatic attraction between the negatively charged shared pair and the two positively charged nuclei provides the bonding force; the inter-nuclear repulsion between the two nuclei is overcome at the equilibrium internuclear distance, where attractive and repulsive forces balance to give a potential-energy minimum. The depth of that minimum is the bond enthalpy; the position of the minimum on the internuclear-distance axis is the bond length. Both quantities are routinely tabulated for A-Level use.
Because the bonding electrons are localised between the two nuclei, simple covalent molecules tend to have low melting and boiling points (no extended electrostatic lattice — see lesson 0 for the contrast with ionic bonding), do not conduct electricity in any state (no free charge carriers), and dissolve preferentially in non-polar solvents. These macroscopic properties are downstream of the localised-pair model.
A single bond involves one shared pair of electrons.
Each hydrogen atom has 1 electron. They share one pair to reach the helium duplet:
Each Cl has 7 outer electrons; they share one pair to reach an octet:
Each Cl ends with 1 bonding pair and 3 lone pairs.
Oxygen has 6 outer electrons and needs 2 more. It shares one pair with each of two hydrogen atoms:
Oxygen has 2 bonding pairs and 2 lone pairs (4 electron pairs total).
Nitrogen has 5 outer electrons and shares one pair with each of three hydrogen atoms:
Nitrogen has 3 bonding pairs and 1 lone pair.
Carbon has 4 outer electrons; each H provides 1. Four shared pairs give carbon a full octet:
Carbon has 4 bonding pairs and 0 lone pairs → tetrahedral, all H–C–H angles 109.5°.
Oxygen (O₂): Each oxygen needs 2 more electrons. They share two pairs:
Each oxygen has 2 bonding pairs (in one double bond) and 2 lone pairs. (At A-Level we treat O₂ as having a double bond; the experimentally observed paramagnetism of O₂ — which the simple Lewis picture cannot explain — is taken up in the Going Further section.)
Carbon dioxide (CO₂): Carbon shares two pairs with each oxygen:
Carbon has 4 bonding pairs (in two double bonds) and no lone pairs.
Ethene (C₂H₄):
The C═C double bond consists of one sigma (σ) bond and one pi (π) bond.
Nitrogen (N₂): Each nitrogen needs 3 more electrons. They share three pairs:
Each nitrogen has 3 bonding pairs (in one triple bond) and 1 lone pair. The N≡N triple bond is very strong (mean bond enthalpy = 944 kJ mol⁻¹), which is why N₂ is so kinetically inert and why fixing atmospheric nitrogen industrially (Haber process) requires aggressive conditions.
Hydrogen cyanide (HCN):
Carbon forms a single bond with hydrogen and a triple bond with nitrogen.
Counting up:
This bookkeeping immediately rationalises a key A-Level observation: the enthalpy of a C=C bond (612 kJ mol⁻¹) is less than 2 × the enthalpy of a C–C bond (2 × 347 = 694 kJ mol⁻¹). The σ component is roughly the same in both, but the additional π in C=C is intrinsically weaker than the σ, so the doubling falls short. The same logic explains why C≡C (839) is well below 3 × C–C (1041).
Key Definition: A dative covalent bond (also called a coordinate bond) is a covalent bond in which both electrons in the shared pair come from the same atom — the donor. The other atom — the acceptor — must have a vacant orbital able to receive the pair.
The atom that donates both electrons must possess a lone pair. Once formed, a dative bond is identical in length, strength and properties to an ordinary covalent bond — the historical distinction lives only in the bookkeeping of where the electrons originated.
Ammonia has a lone pair on nitrogen. When it reacts with H⁺ (which has an empty 1s orbital), nitrogen donates its lone pair:
The arrow → represents the dative bond pointing from donor (N) to acceptor (H⁺). In the finished NH₄⁺ ion, all four N–H bonds are crystallographically and spectroscopically identical — you cannot pick out the dative one. The shape is tetrahedral, H–N–H = 109.5°.
Water has two lone pairs on oxygen. One of them is donated to H⁺:
All three O–H bonds in H₃O⁺ are identical; the shape is pyramidal with bond angle ≈ 107° (three bonding pairs, one lone pair on O).
Carbon has 4 outer electrons, oxygen has 6. A naive double bond gives C only 6 electrons. To attain an octet on both atoms, the picture is: two shared pairs (one σ, one π) drawn from one electron each, plus a third dative pair donated from oxygen to carbon:
Each atom ends with three bonding pairs (in a triple bond) and one lone pair. CO is isoelectronic with N₂ and — like N₂ — has a very high bond enthalpy (~1077 kJ mol⁻¹).
In the gas phase, monomeric AlCl₃ is electron deficient — Al has only 6 electrons in its valence shell after sharing one pair with each of three chlorines. Two AlCl₃ units therefore dimerise: a chlorine on one monomer donates a lone pair to the empty orbital on the aluminium of the other, and vice versa, generating two bridging dative bonds:
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