1The idea in one paragraph

A metal atom that loses electrons becomes a positive ion, a cation. A non-metal atom that gains them becomes a negative ion, an anion. Opposite charges attract, and that attraction, acting in every direction at once, is the ionic bond. Nothing stops at two ions: each ion attracts all the oppositely charged ions around it, so the ions pack into a huge, regular, three-dimensional lattice. Almost every property of an ionic compound follows from that picture. It melts only at high temperature because millions of strong attractions must be overcome. It conducts electricity only when melted or dissolved, because only then can its ions move. And it often dissolves in water, because water molecules can attract the ions strongly enough to pull them out of the lattice.

2Making ions: what the electron configuration tells you

An atom is neutral because its protons and electrons are equal in number. Take electrons away and the protons win, so the particle carries a positive charge. Add electrons and it carries a negative charge. The charge is simply the imbalance.

Which way an atom goes depends on its outer shell. Sodium is 2,8,1. Losing one electron leaves 2,8, the arrangement of neon, and that takes far less energy than gaining seven. Chlorine is 2,8,7. Gaining one electron gives 2,8,8, the arrangement of argon. Figure 1 shows the transfer, with sodium's electrons drawn as crosses so you can follow the one that moves.

The sodium ion has 11 protons and 10 electrons, so its charge is 1+. The chloride ion has 17 protons and 18 electrons, so its charge is 1−. Notice the names: the cation keeps the element's name, sodium ion, while the anion takes the ending -ide, chloride ion.

The rule this gives you is to count how many electrons the atom must lose or gain to reach the nearest noble gas configuration. For the main-group elements that count is fixed by the group, which is why Figure 2 works.

Work one from the configuration alone. An atom is 1s² 2s² 2p⁶ 3s² 3p⁶ 4s². It has two electrons beyond the argon core, so it loses both and forms a 2+ ion with configuration 1s² 2s² 2p⁶ 3s² 3p⁶. The element is calcium and the ion is Ca²⁺. Another atom is 1s² 2s² 2p³. It is three electrons short of neon, so it gains three and forms a 3− ion, 1s² 2s² 2p⁶: the nitride ion, N³⁻.

Group 14 does not appear in Figure 2 for a reason. Carbon and silicon would need to lose or gain four electrons, which costs too much energy, so they share electrons instead. That is S2.2.

Transition elements are the exception. Iron is [Ar] 3d⁶ 4s². The 4s and 3d sub-levels are so close in energy that iron can lose two electrons or three, and both ions are common. Figure 3 shows the rule that decides which electrons go.

The 4s electrons are lost first, even though 4s filled first. So Fe²⁺ is [Ar] 3d⁶, not [Ar] 3d⁴ 4s², and Fe³⁺ is [Ar] 3d⁵. Copper, [Ar] 3d¹⁰ 4s¹, forms Cu⁺ ([Ar] 3d¹⁰) and Cu²⁺ ([Ar] 3d⁹) in the same way. Because the charge cannot be read from the group, it goes in the name as a Roman numeral: iron(II) is Fe²⁺, iron(III) is Fe³⁺, copper(II) is Cu²⁺.

3The ionic bond

An ionic bond is the electrostatic attraction between oppositely charged ions.

Two points in that definition earn marks. The word electrostatic says what the force is: the attraction between opposite charges, the same force that holds an electron to a nucleus. The words oppositely charged ions say between what. An answer that says "an ionic bond is when electrons are transferred" describes how the ions were made, not what the bond is, and scores nothing for the definition.

The attraction is non-directional. A Na⁺ ion does not bond to one particular Cl⁻ ion; it attracts every Cl⁻ around it, and more weakly the further away they are. That is why an ionic compound is never a small unit but always a lattice.

Ionic bonding only happens between different elements. Ions of opposite charge are needed, so one atom must give electrons and another must take them. Two chlorine atoms both want to gain an electron, and neither will give one up. Covalent bonding, where electrons are shared, does not have this limit, which is why Cl₂ exists and "chlorine chloride" does not. In practice ionic bonding is expected between a metal on the left of the table and a non-metal on the right, where the difference in electronegativity is large. S2.4 turns that into a scale.

The formation of an ionic compound from its elements is also a redox reaction: sodium loses electrons and is oxidised, chlorine gains them and is reduced. You meet the language of that in Reactivity 3.2.

4Formulas and names

An ionic compound has no overall charge, so the positive and negative charges in its formula must cancel. That is the only rule you need.

Binary compounds contain two elements. Name them cation first, then the anion with the ending -ide: sodium chloride, magnesium oxide, calcium nitride, lithium sulfide. To write the formula, find the smallest whole numbers of each ion that make the charges cancel.

The reverse direction is just as mechanical. CuO contains O²⁻, so the copper must be 2+: copper(II) oxide. Fe₂O₃ carries three O²⁻, a total of −6, shared between two iron ions, so each is 3+: iron(III) oxide. The Roman numeral is always the charge on one metal ion, never the number of atoms.

Polyatomic ions are groups of covalently bonded atoms that carry an overall charge. They move through a formula as one unit. The guide requires these seven, by name and formula:

Ammonium is the only positive one, and it behaves like a metal ion in a formula. When you need more than one polyatomic ion, put it in brackets with the number outside.

Without the brackets, CaOH₂ would mean one oxygen and two hydrogens, which is a different thing. Never change the subscripts inside a polyatomic ion to make the charges fit: sulfate is always SO₄²⁻.

5The lattice, and why its formula is empirical

Figure 4 shows a small part of the sodium chloride lattice. The ions alternate in all three directions. Each Na⁺ is surrounded by six Cl⁻ ions, and each Cl⁻ by six Na⁺ ions.

There is no molecule of sodium chloride anywhere in that picture. No Na⁺ belongs to one particular Cl⁻. A single grain of salt is one lattice containing around 10¹⁸ ions or more, and the formula NaCl tells you only the ratio of the ions in it: one to one. A formula that gives the simplest whole-number ratio of the particles is an empirical formula, and every ionic formula is one. That is why MgCl₂ means "twice as many chloride ions as magnesium ions" and why chemists call NaCl a formula unit rather than a molecule.

Not every ionic lattice has the six-around-one arrangement of sodium chloride. The pattern depends on the ratio of the ions and on their relative sizes. You are not asked to know the other arrangements, only that the structure is a three-dimensional lattice of alternating charges.

6How strong is an ionic lattice? Lattice enthalpy

Lattice enthalpy measures the strength of the ionic bonding in a compound. It is the enthalpy change when one mole of a solid ionic compound is separated into its gaseous ions, far enough apart that they no longer attract. Pulling opposite charges apart always takes energy, so lattice enthalpy is always positive: the bigger the number, the stronger the lattice. Values are listed in the data booklet. Figure 5 compares five.

Two things set the size of the attraction between two ions, and both come straight from the physics of charges.

The charge on the ions. The force between two charges is proportional to the product of the charges. A 2+ ion and a 2− ion attract with a product of 4, against 1 for a 1+ and a 1− ion. That is why MgO and CaO sit roughly four times higher than the 1+/1− compounds in Figure 5.

The radius of the ions. The force falls as the distance between the centres of the ions grows. Small ions sit closer together, so they attract more strongly. Going from KCl to NaCl swaps a larger cation for a smaller one; going from NaCl to NaF swaps a larger anion for a smaller one. Each step raises the lattice enthalpy.

Higher charges and smaller ions give a larger lattice enthalpy, stronger ionic bonding and a higher melting point.

When you compare two compounds, check charge first, because it has the bigger effect. Only if the charges are the same does size decide. MgO against NaCl is a charge argument: both are made of fairly small ions, but MgO's ions carry double charges. KCl against NaCl is a size argument, because the charges are the same.

Ion size itself follows from S1.3. A cation is smaller than its atom, because it has lost its outer shell and the remaining electrons are pulled in by the same number of protons. An anion is larger than its atom, because the extra electrons repel each other and the nucleus has no extra protons to hold them. Ionic radii are given in the data booklet.

7Properties: volatility, conductivity, solubility

The guide asks you to explain three properties. Each explanation starts from the lattice.

Volatility. A volatile substance evaporates easily. Ionic compounds have very low volatility: they are solids at room temperature with high melting and boiling points. Sodium chloride melts at about 801 °C and magnesium oxide at about 2800 °C. To melt the solid the ions must gain enough energy to move past each other, which means overcoming a very large number of strong electrostatic attractions throughout the lattice. The order of melting points usually follows the order of lattice enthalpies, which is why MgO is so much higher than NaCl.

Electrical conductivity. An electric current is a flow of charged particles. Figure 6 shows why an ionic compound conducts in one state and not the other.

In the solid the ions are held in fixed positions by the lattice. The charges are there, but they cannot move, so there is no current. When the compound is melted, or dissolved in water, the lattice breaks down and the ions are free to move: cations drift towards the negative electrode, anions towards the positive one, and a current flows. The charge carriers are the ions. There are no free electrons in an ionic compound in any state, and saying there are is the most expensive mistake in this subtopic.

Solubility. Many ionic compounds dissolve in water, and almost none dissolve in non-polar solvents such as hexane. The reason is in Figure 7.

A water molecule is polar: its oxygen end carries a small negative charge, written δ−, and its hydrogen ends carry small positive charges, δ+ (you meet the reason in S2.2). Water molecules therefore turn their δ− oxygen towards a cation and a δ+ hydrogen towards an anion. Those attractions between ions and polar molecules are strong enough to compensate for the lattice attractions the ions give up, so ions at the surface of the crystal are pulled into solution, each wrapped in water molecules. This is hydration. Hexane molecules are non-polar, so they cannot attract ions strongly, and the lattice stays intact.

"Many" is not "all". Calcium carbonate, silver chloride and barium sulfate barely dissolve at all, because for them the lattice wins. Do not write "ionic compounds are soluble in water" as a rule; write "many ionic compounds are soluble in water and insoluble in non-polar solvents".

Brittleness is not on the guide's list, but it follows from the same picture and it is a useful check that you understand the lattice. Hit an ionic crystal hard enough to shift one layer of ions by half a step, and ions of the same charge end up next to each other. They repel, and the crystal splits cleanly along that plane.

Put the three explained properties together and you have a test for ionic bonding. A substance that is solid with a high melting point, does not conduct as a solid, conducts when molten or in solution, and dissolves in water but not in hexane, is almost certainly ionic. Paper 1B likes to hand you exactly this kind of table and ask which substance is which.