Revise The ionic model for IB Chemistry (S2.1)

Ionic bonding occurs between elements with a large difference in electronegativity, almost always a metal and a non-metal. The metal has a low ionisation energy, so it loses its valence electrons readily; the non-metal has a high electron affinity, so it accepts them readily. Electrons transfer from metal to non-metal until each atom reaches a stable, full outer shell — the electron configuration of the nearest noble gas. Sodium, [Ne]3s¹, loses its single 3s electron to become Na⁺ (2,8), isoelectronic with neon; chlorine, [Ne]3s²3p⁵, gains one electron to become Cl⁻ (2,8,8), isoelectronic with argon.

Because the charge on a monatomic ion is set by how many electrons it must lose or gain to fill its outer shell, you can read the likely charge straight off the periodic table. Group 1 forms 1+, Group 2 forms 2+, Group 13 forms 3+; Group 15 forms 3−, Group 16 forms 2−, Group 17 forms 1−. To build a formula, remember that the compound must be electrically neutral: the total positive charge from the cations must exactly cancel the total negative charge from the anions. The quickest route is to find the lowest whole-number ratio of ions that makes the charges balance.

Once formed, ions do not simply pair off. The electrostatic attraction of an ion is non-directional — a positive ion attracts every nearby negative ion equally in all directions, and vice versa. Ions therefore pack together in the arrangement that maximises attraction between opposite charges and minimises repulsion between like charges, producing a regular, repeating three-dimensional array called a giant ionic lattice. In sodium chloride each Na⁺ is surrounded by six Cl⁻ and each Cl⁻ by six Na⁺ (a coordination number of 6), and this pattern repeats over the entire crystal. There are no molecules: a formula such as NaCl gives only the simplest ratio of ions, called a formula unit.

Every characteristic property of an ionic compound follows from one idea: a giant lattice held by strong, non-directional electrostatic attraction, in which the ions are fixed in the solid but can move once the lattice is broken.

Not all ionic solids melt at the same temperature. The strength of the electrostatic attraction between two ions — and hence the energy needed to break the lattice — increases when the ionic charges are larger and when the ionic radii are smaller, because both bring more charge closer together. These are precisely the two factors that make a lattice enthalpy larger in magnitude. So magnesium oxide (Mg²⁺, O²⁻ — both doubly charged and small) has a much stronger lattice and a far higher melting point than sodium chloride (Na⁺, Cl⁻ — singly charged and larger). Down a group, as ionic radius grows, melting points of comparable compounds tend to fall (LiF > NaCl > KBr).

Structured 'explain' questions in Paper 2 are marked analytically: the marks are a checklist of distinct valid points, not a holistic markband. Each separate correct idea scores 1 mark, and the engine credits any wording that carries the same chemistry ('equivalent'). A mark tagged M is a reasoning point (the correct link, e.g. 'because the ions are fixed'); a mark tagged A is a correct conclusion or statement of fact. Error-carried-forward (ECF) means a correct consequence drawn from an earlier wrong statement can still score. Study how each of the four marks below is tied to one specific idea.

The ionic model is one of three bonding models in this unit. Next you will meet the covalent model (shared electrons and discrete molecules) and the metallic model (a lattice of cations in a sea of delocalised electrons), and then learn to place real bonds on a continuum of bonding using electronegativity difference. The same habit carries through: describe the structure, identify the forces, and let the forces explain the properties. Master that for ions here and the covalent and metallic models become variations on a method you already own.