Covalent Bonding and Dative Bonds

Spec Mapping — OCR H432 Module 2.2.2 — Bonding and structure, covering a covalent bond as a shared pair of electrons between two atoms with each pair attracted to both nuclei, the distinction between single, double and triple bonds, dot-and-cross diagrams for simple molecules and polyatomic ions, dative (coordinate) covalent bonding as a special case in which both electrons of the bonding pair originate from the same atom, and the principal exceptions to the octet rule (incomplete octets, expanded octets, odd-electron species) (refer to the official OCR H432 specification document for exact wording).

A covalent bond is the second of the three core bonding models. Where ionic bonding involves complete electron transfer between atoms of very different electronegativity, covalent bonding involves sharing — both atoms contribute electrons to a bonding pair that is attracted to both nuclei simultaneously. This pairing model, developed by Lewis in 1916 (and formalised quantum-mechanically by Heitler and London in 1927 and Pauling in the 1930s), underlies the vast majority of organic chemistry, the structure of molecular gases (O₂, N₂, CO₂), and the geometry of molecular ions (NH₄⁺, SO₄²⁻). This lesson teaches you to draw dot-and-cross diagrams systematically, to count electrons in single, double and triple bonds, to recognise dative (coordinate) bonds — a special case in which both electrons of a bonding pair come from the same atom — and to handle the principal exceptions to the simple octet rule.

Key Definitions:

• Covalent bond — a shared pair of electrons between two atoms, attracted to both nuclei.
• Single bond — one shared pair (2 electrons).
• Double bond — two shared pairs (4 electrons).
• Triple bond — three shared pairs (6 electrons).
• Dative (coordinate) covalent bond — a covalent bond in which both shared electrons come from the same atom.
• Lone pair — a non-bonding pair of valence electrons on an atom.

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What a covalent bond is

When two non-metal atoms approach, their valence electrons can interact in one of two ways. In bonding overlap, the wavefunctions of the two atomic orbitals combine constructively to form a molecular orbital with high electron density between the two nuclei — both nuclei attract the same pair of electrons, and the system is bound together. The Lewis dot picture — a "shared pair" between two atoms — is the qualitative shorthand for this quantum-mechanical bonding overlap.

Each covalent bond holds 2 electrons in a region where both nuclei feel the attraction. The atoms typically share enough electrons for each to reach a stable noble-gas configuration:

• Hydrogen needs 2 outer electrons (matching He).
• Most second-row elements (C, N, O, F) need 8 outer electrons (matching Ne) — the octet rule.
• Period-3 onwards can sometimes exceed 8 outer electrons via d-orbital involvement (expanded octets).

Covalent bonding occurs principally between two non-metals with small electronegativity difference; the electrons are shared rather than transferred outright.

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Dot-and-cross diagrams for simple molecules

Dot-and-cross diagrams use dots for one atom's outer electrons and crosses for the other's, with the bonding pair shown in the overlap region between atoms. Lone pairs sit on individual atoms.

H₂ H H · × 1 bond pair

HCl H Cl · × ·· ·· ·· 1 bond pair + 3 lone pairs on Cl

H₂O O H H · · × × ·· ·· 2 bond pairs + 2 lone pairs on O

NH₃ N H H H ·· 3 bond pairs + 1 lone pair

CH₄ C H H H H 4 bond pairs, no lone pair

CO₂ O C O ·· ·· ·· ·· 2 C=O double bonds, 4 lone pairs total

N₂ N N ·· ·· N≡N triple bond + 2 lone pairs

Dots = electrons from atom A; crosses = electrons from atom B; bond pair sits between the atoms. Lone pair = non-bonding pair on one atom only.

Six canonical examples

• H₂ — each H provides 1 electron; the shared pair gives each H an effective He configuration.
• HCl — 1 bond pair (1 electron each); Cl retains 3 lone pairs to complete its octet.
• H₂O — O provides 6 outer electrons, shares 2 (one each with each H), retains 2 lone pairs.
• NH₃ — N provides 5 outer electrons, shares 3, retains 1 lone pair. (The lone pair is responsible for NH₃'s basicity and the dative bond in NH₄⁺.)
• CH₄ — C provides 4 outer electrons, all shared as 4 C-H bond pairs; no lone pair on C.
• CO₂ — two C=O double bonds; each O contributes 4 electrons to its double bond and retains 2 lone pairs. C has no lone pairs.
• N₂ — three shared pairs (a triple bond) plus 1 lone pair on each N. Bond enthalpy 945 kJ mol⁻¹ — exceptionally high; explains why atmospheric N₂ is inert.

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Multiple bonds — single, double, triple

Bond order	Shared pairs	Electrons in bond	Mode
1 (single)	1	2	1 σ bond
2 (double)	2	4	1 σ + 1 π bond
3 (triple)	3	6	1 σ + 2 π bonds

A σ (sigma) bond is end-on overlap of orbitals along the internuclear axis. A π (pi) bond is sideways overlap of p-orbitals, with electron density above and below the axis. The σ bond is always stronger than any individual π bond.

Higher bond order means shorter and stronger bonds:

Bond	Bond length (pm)	Bond enthalpy (kJ mol⁻¹)
C-C	154	347
C=C	134	612
C≡C	120	838
N-N	145	163
N=N	125	418
N≡N	110	945
O-O	148	144
O=O	121	498

Notice that a triple bond is not three times stronger than a single bond (the π bonds are weaker than the σ); but it is always shorter in length.

H-X bond trend down the halogen group

Bond	Length (pm)	Enthalpy (kJ mol⁻¹)
H-F	92	568
H-Cl	127	432
H-Br	141	366
H-I	161	298

As you descend Group 17 the halogen atom gets larger, the bonding pair sits further from the halogen nucleus, and the bond becomes longer and weaker. This trend explains why HI thermally decomposes much more readily than HF — a synoptic point linking to Module 3 (Group 7 chemistry).

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Dative (coordinate) covalent bonds

A dative covalent bond (= coordinate covalent bond) is a covalent bond in which both electrons of the shared pair come from the same atom. The donor atom must have a lone pair; the acceptor must have an empty orbital.

Once formed, a dative bond is identical in every measurable property (length, strength, polarity) to an ordinary covalent bond — you cannot tell them apart experimentally. Only the origin of the electron pair differs. We mark a dative bond in diagrams with an arrow pointing from donor (D) to acceptor (A): D → A.

Example 1 — the ammonium ion NH₄⁺

NH₃ has a lone pair on N. When NH₃ meets H⁺ (a bare proton with an empty 1s orbital), the lone pair donates into the empty orbital, forming a fourth N-H bond — a dative bond:

$\text{NH}_3 + \text{H}^+ \rightarrow \text{NH}_4^+$NH3​+H+→NH4+​

In NH₄⁺ all four N-H bonds are equivalent in length, strength, and reactivity. The +1 charge resides on the ion as a whole (formally on N for electron-counting purposes).

NH₄⁺ N H H H H dative [NH₄]⁺ — 4 equivalent N-H bonds

H₃O⁺ O H H H ·· dative [H₃O]⁺ — pyramidal, 1 lone pair, 3 bond pairs

Al₂Cl₆ dimer Cl-Al Al-Cl Cl Cl 2 dative bonds (Cl→Al) bridge two AlCl₃

Dative bond arrow points from donor (lone-pair atom) to acceptor (empty-orbital atom)

Example 2 — hydronium ion H₃O⁺

H₂O has two lone pairs on O. One lone pair donates to an incoming H⁺:

$\text{H}_2\text{O} + \text{H}^+ \rightarrow \text{H}_3\text{O}^+$H2​O+H+→H3​O+

H₃O⁺ has three O-H bonds and one remaining lone pair — pyramidal geometry, very similar to NH₃ (next lesson).

Example 3 — tetrahydroxoaluminate [Al(OH)₄]⁻

Al³⁺ has empty 3s/3p orbitals. Four OH⁻ ions each donate a lone pair to Al, forming four dative bonds:

$\text{Al}^{3+} + 4\text{OH}^- \rightarrow [\text{Al(OH)}_4]^-$Al3++4OH−→[Al(OH)4​]−

This ion forms when Al(OH)₃ is dissolved in excess NaOH — the basis of the amphoteric behaviour of Al₂O₃ you met in Lesson 1.

Example 4 — carbon monoxide CO

CO is unusually strongly bound (bond enthalpy 1077 kJ mol⁻¹, even higher than N₂). The conventional dot-and-cross has a C=O double bond plus a dative bond from an O lone pair into an empty C orbital, giving an effective C≡O triple bond. The bond length (113 pm) is consistent with a triple bond.

Example 5 — aluminium chloride dimer Al₂Cl₆