Ionic bonding and lattice strength

Ionic bonding is the strong electrostatic attraction between oppositely charged ions, arranged in a giant ionic lattice. Metals lose electrons to form cations; non-metals gain them to form anions.

• Attraction is stronger when the ionic charge is larger and the ionic radius is smaller. That is why MgO (2+ / 2−, small ions) melts at 2852 °C but NaCl (1+ / 1−) melts at 801 °C.
• Ionic radii increase down a group (extra shells). Across an isoelectronic set (N³⁻, O²⁻, F⁻, Na⁺, Mg²⁺) the radius falls as nuclear charge rises with the same number of electrons.
• Ionic solids do not conduct (ions are fixed) but conduct when molten or aqueous — the ions are free to migrate: cations to the negative electrode, anions to the positive.

Evidence: electron-density maps show real ionic compounds are not perfectly ionic — bonding is a continuum from pure ionic to pure covalent, and a small, highly charged cation can polarise a large anion, giving partial covalent character.

Covalent and dative covalent bonds

A covalent bond is the strong electrostatic attraction between two nuclei and the shared pair of electrons between them. In a dative (coordinate) bond both electrons come from the same atom — shown by an arrow, e.g. in NH₄⁺ and in Al₂Cl₆.

• Shorter bonds are stronger: C≡C (0.120 nm, 838 kJ mol⁻¹) > C=C (0.134 nm, 612) > C–C (0.154 nm, 347).
• Once formed, the four N–H bonds in NH₄⁺ are identical — you cannot tell which one was dative.
• Some molecules expand the octet (PCl₅, SF₆) because the third shell has d orbitals available; others are electron deficient (BF₃, 6 electrons).

Dative bonds

Electron-pair repulsion (VSEPR)

Electron pairs in the outer shell of the central atom repel each other and take up positions as far apart as possible. Repulsion strength: lone pair–lone pair > lone pair–bonding pair > bonding pair–bonding pair. Each lone pair squeezes the bond angle by roughly 2.5°.

• 2 pairs → linear, 180° (CO₂, BeCl₂)
• 3 pairs → trigonal planar, 120° (BF₃)
• 4 pairs → tetrahedral, 109.5° (CH₄); 3 bonding + 1 lone → pyramidal, 107° (NH₃); 2 bonding + 2 lone → bent, 104.5° (H₂O)
• 5 pairs → trigonal bipyramidal, 120° and 90° (PCl₅)
• 6 pairs → octahedral, 90° (SF₆); 4 bonding + 2 lone → square planar, 90° (XeF₄)

Method: count the electrons in the central atom's outer shell, add one per bonded atom (adjust for charge), divide by 2 = number of pairs. Then subtract the bonding pairs to get the lone pairs.

Your turn — bond angle in ammonia

Your turn — bond angle in SF₆

Your turn — counting lone pairs

Predicting a shape

Electronegativity, polar bonds and polar molecules

Electronegativity is the ability of an atom to attract the bonding electrons in a covalent bond. It increases across a period and decreases down a group — F (4.0) > O (3.5) > N = Cl (3.0) > C (2.5) > H (2.1) on the Pauling scale.

• A difference in electronegativity gives a polar bond with a permanent dipole (δ+ and δ−).
• A molecule with polar bonds is only polar overall if the dipoles do not cancel. CO₂ is linear and symmetrical → the two dipoles cancel → non-polar. H₂O is bent → they do not cancel → polar. CCl₄ is tetrahedral and symmetrical → non-polar.
• A very large difference (roughly > 1.7) gives ionic bonding — bonding is a continuum.

Solvents: like dissolves like. Ionic and polar solutes dissolve in polar solvents (water solvates ions with ion–dipole attraction); non-polar solutes dissolve in non-polar solvents. If the solvation energy cannot repay the lattice energy, the salt stays insoluble.

Your turn — electronegativity difference

Polar bonds, non-polar molecule

The three intermolecular forces

Intermolecular forces are much weaker than covalent bonds (a few kJ mol⁻¹ versus hundreds).

• London forces (induced dipole–induced dipole) act between all molecules. Random electron movement gives an instantaneous dipole, which induces a dipole in a neighbour. They get stronger with more electrons and with a larger contact surface area — hence the rise in boiling temperature down Group 7 and along the alkanes (and branching lowers b.p. because contact is reduced).
• Permanent dipole–dipole forces act in addition between polar molecules (e.g. HCl).
• Hydrogen bonds — the strongest — form when H is bonded to N, O or F and is attracted to a lone pair on the N, O or F of another molecule.

Common error: boiling a molecular liquid does not break covalent bonds — only the intermolecular forces are overcome.

Hydrogen bonding and the anomalies of water

A hydrogen bond needs H attached to N, O or F (very electronegative, so H is strongly δ+) and a lone pair on the N, O or F of the next molecule. The bond is directional — 180° through the H atom.

• High boiling temperature: H₂O, NH₃ and HF all boil far above the trend of their groups because H bonds must be overcome.
• Ice is less dense than water: each molecule forms up to four hydrogen bonds in an open tetrahedral lattice, holding the molecules further apart than in the liquid — so ice floats.
• Water has a high surface tension and viscosity for its size, for the same reason.

Why does ice float?

Boiling temperature down Group 7

Giant structures and the carbon allotropes

Predict properties from structure + bonding:

• Giant ionic (NaCl, MgO): high m.p., brittle, conducts only when molten/aqueous, often soluble in water.
• Giant covalent (diamond, SiO₂): very high m.p., hard, insoluble, non-conducting — except graphite.
• Graphite: layers of hexagonal rings, each C bonded to 3 others; the fourth electron is delocalised → conducts along the layers; weak London forces between layers → soft, a lubricant. Graphene is a single layer. Diamond: each C bonded to 4 in a tetrahedral lattice → hardest, does not conduct.
• Simple molecular (I₂, CO₂): low m.p., soft, non-conducting — the lattice is held by London forces, not covalent bonds.
• Metallic: positive ions in a sea of delocalised electrons — conducts (electrons move), malleable (layers slide), strong bonding for Mg²⁺ > Na⁺.

Which force must be overcome on melting?

Tap a substance, then tap the strongest force that has to be broken or overcome when it melts.

Diamond vs graphite

Match the molecule to its shape and bond angle

Tap an item on the left, then its partner on the right.

The big ideas to know

You've now covered Topic 2: Bonding and Structure of the Edexcel A-level Chemistry (9CH0) specification. Press Finish to see your score.