AS Chemistry: Ionic, Covalent and Metallic Bonding — Revision Notes

Condensed for the final weeks. For the full explanation, use the Electronegativity, Ionic and Metallic Bonds study guide.

The three bond types

Bond	Formed between	Held together by
Ionic	Metal + non-metal	Electrostatic attraction between oppositely charged ions
Covalent	Non-metal + non-metal	Attraction between the shared pair and both nuclei
Metallic	Metal atoms	Attraction between positive ions and delocalised electrons

Every definition must mention electrostatic attraction. “Atoms share electrons” alone does not describe the force.

A dative (coordinate) covalent bond is one where both electrons of the shared pair come from the same atom — NH₄⁺, H₃O⁺, and every metal–ligand bond. Once formed it is identical in every respect to a normal covalent bond.

Worked electron-transfer examples, the three ionic compounds most often examined directly: NaCl — sodium (2,8,1) transfers one electron to chlorine (2,8,7), giving Na⁺ (2,8) and Cl⁻ (2,8,8). MgO — magnesium (2,8,2) transfers two electrons to oxygen (2,6), giving Mg²⁺ (2,8) and O²⁻ (2,8). CaF₂ — calcium (2,8,8,2) transfers one electron to each of two fluorines (2,7), giving Ca²⁺ (2,8,8) and two F⁻ (2,8). In every case the resulting ions pack into a giant ionic lattice, held together by electrostatic attraction extending in all directions, not confined to one ion pair.

Electronegativity

The ability of an atom to attract the bonding electrons in a covalent bond.

Increases across a period (greater nuclear charge, same shielding, smaller radius) and decreases down a group (more shells, more shielding, greater distance). Fluorine is the most electronegative element.

The electronegativity difference determines bond character — but the boundary is a continuum, not a switch:

difference ~0        pure covalent
small difference     polar covalent
large difference     ionic

Polarisation — the bridge between ionic and covalent

A small, highly charged cation distorts the electron cloud of a large anion, pulling electron density into the space between the nuclei. This gives an ionic compound covalent character.

Polarising power of a cation increases with higher charge and smaller radius. Polarisability of an anion increases with larger radius and higher charge.

This is why AlCl₃ behaves covalently — Al³⁺ is small and triply charged, so it polarises chloride heavily — while NaCl does not. It also explains why lithium compounds are the most covalent in Group 1 and why carbonates of small cations decompose at lower temperatures.

Structure and properties

Structure	Example	m.p.	Conducts	Solubility
Giant ionic	NaCl	High	Molten/aqueous only	Soluble in water
Giant covalent	Diamond, SiO₂	Very high	No	Insoluble
Giant covalent (layered)	Graphite	Very high	Yes	Insoluble
Simple molecular	I₂, CO₂	Low	No	Non-polar solvents
Giant metallic	Cu, Fe	High	Yes, solid and molten	Insoluble

Ionic solids do not conduct because the ions are fixed in the lattice; molten or dissolved, they are free to move. Metals conduct in both states because the delocalised electrons are mobile regardless.

Graphite is the standard question. Each carbon forms three covalent bonds, leaving one delocalised electron per atom free to move along the layers — hence conduction. The layers are held together only by weak induced-dipole forces, so they slide over one another, making graphite soft and a good lubricant. Diamond has four bonds per carbon, no free electrons, and a rigid tetrahedral network — hard, and an insulator.

Simple molecular substances melt at low temperatures because the weak intermolecular forces break, not the strong covalent bonds. Saying “the covalent bonds break” is a common error worth watching for — the covalent bonds within each molecule stay intact.

Metallic bond strength

Increases with:

• More delocalised electrons per atom — Mg (2 per atom) melts far higher than Na (1).
• Smaller ionic radius — stronger attraction to the electron sea.
• Higher ionic charge.

Metallic bonding follows the same electrostatic logic: metal atoms release outer-shell electrons into a shared “sea” free to move through the whole structure, leaving a regular lattice of positive ions. This is not a bond between any two specific atoms but a force extending through the entire structure — which is exactly why metals conduct electricity in both solid and molten states (the delocalised electrons carry charge freely) and are malleable (layers of ions slide past each other while the electron sea simply redistributes, so no specific bond has to break).

Exam traps

• Defining bonds without “electrostatic attraction”.
• Saying covalent bonds break when a simple molecular solid melts.
• Saying ionic solids conduct.
• Confusing polarising power (cation) with polarisability (anion).
• Forgetting graphite’s fourth electron is delocalised, not unbonded.
• Treating ionic and covalent as a strict binary rather than a continuum.

Self-test

1. Define electronegativity and give its two trends.
2. What is a dative covalent bond, and how does it differ once formed?
3. Which cation properties increase polarising power?
4. Why does graphite conduct but diamond not?
5. Why do simple molecular substances have low melting points?

Answers: 1. The ability of an atom to attract the bonding electrons in a covalent bond; it increases across a period and decreases down a group. 2. One in which both electrons of the shared pair come from the same atom; once formed it is indistinguishable from an ordinary covalent bond. 3. High charge and small ionic radius. 4. Each carbon in graphite forms only three covalent bonds, leaving one delocalised electron per atom free to move along the layers; in diamond all four outer electrons are used in covalent bonds, so none are free. 5. Only the weak intermolecular forces between molecules need to be overcome; the strong covalent bonds within the molecules remain intact.