Ionic Lattices & Compound Ions Revision Notes | A Level AQA Chemistry

1. What Holds an Ionic Compound Together?

Picture a grain of table salt — $\text{NaCl}$ — sitting on your fingertip. It is hard, brittle, and has a melting point of 801 °C. None of that happens by accident. The reason lies in what is happening at the atomic scale: billions of charged particles locked into a precise, repeating arrangement, each one held in place by powerful electrical forces pulling in every direction at once.

That arrangement is called an ionic lattice, and the force holding it together is electrostatic attraction between oppositely charged ions. This is the AQA definition of ionic bonding, and every word in it matters for the exam.

When a metal atom loses one or more electrons, it becomes a positively charged cation. The non-metal atom that gains those electrons becomes a negatively charged anion. Opposite charges attract, and that attraction is the ionic bond. The key point is that this attraction does not stop at one pair — it extends in all directions throughout the solid.

2. The Ionic Lattice — Structure and Significance

Now that you know what ionic bonding is, the next question is: how do those ions arrange themselves? The answer is the ionic lattice — a giant, regular, three-dimensional structure where every ion is surrounded by ions of the opposite charge.

In sodium chloride, each $\text{Na}^+$ ion is surrounded by six $\text{Cl}^-$ ions, and each $\text{Cl}^-$ is surrounded by six $\text{Na}^+$ ions. This 6:6 coordination repeats throughout the entire crystal. The result is a structure with no discrete molecules — the whole crystal is one giant ionic lattice.

The strength of the electrostatic attraction depends on two things: the size of the ionic charges and the distance between the ions. Higher charges and smaller ionic radii both increase the strength of the attraction. This is why magnesium oxide ($\text{Mg}^{2+}$ and $\text{O}^{2-}$) has a much higher melting point (2852 °C) than sodium chloride — the 2+ and 2− charges create a far stronger lattice than 1+ and 1−.

This connects directly to the spec requirement: ionic bonding involves electrostatic attraction between oppositely charged ions in a lattice. The lattice is not a side detail — it is the reason ionic compounds have their characteristic properties of high melting points, hardness, and brittleness.

3. Properties Explained by the Lattice

The ionic lattice explains several physical properties that examiners regularly ask about. Each property links directly back to the strength and arrangement of the electrostatic attractions.

4. Compound Ions — Charged Groups That Act as One

So far, the ions we have discussed — $\text{Na}^+$, $\text{Cl}^-$, $\text{Mg}^{2+}$, $\text{O}^{2-}$ — are all single atoms that have gained or lost electrons. At A Level, you also need to know about compound ions: groups of atoms that are covalently bonded together but carry an overall charge, allowing them to participate in ionic compounds.

Think of a compound ion as a team. The atoms within it are held together by covalent bonds, but the whole team carries a charge and behaves as a single unit when forming an ionic lattice. The five compound ions you must know for AQA are listed below — their formulae and charges are non-negotiable exam knowledge.

Notice that ammonium ($\text{NH}_4^+$) is the only cation in this group. All the others are anions. This matters when you are writing formulae for ionic compounds: ammonium compounds pair a positive compound ion with a negative ion, whereas sulfate, carbonate, hydroxide, and nitrate compounds pair a positive metal ion with a negative compound ion.

5. Writing Formulae for Ionic Compounds Containing Compound Ions

Knowing the compound ion formulae is only half the task — you also need to use them correctly when writing the formula of an ionic compound. The rule is the same as for simple ions: the total positive charge must equal the total negative charge, so the compound is electrically neutral overall.

When you need more than one compound ion in a formula, you must place the entire ion in brackets before adding the subscript. This is a point where many students lose marks — writing $\text{Ca(OH)}_2$ correctly requires the brackets to show that both $\text{OH}^-$ units are complete ions, not just two separate O and H atoms.

Worked example: Writing the Formula of Calcium Hydroxide

Write the correct formula for calcium hydroxide, given that calcium forms $\text{Ca}^{2+}$ and hydroxide is $\text{OH}^-$.

1. $\text{Ca}^{2+}$ (charge: 2+) and $\text{OH}^-$ (charge: 1−) — Always start by writing out both ions with their charges clearly. This prevents sign errors later.
2. One $\text{Ca}^{2+}$ provides 2+ charge. Two $\text{OH}^-$ provide 2− charge. Ratio is 1: 2. — The total positive charge (2+) must equal the total negative charge (2−). You need two hydroxide ions to balance one calcium ion.
3. $\text{Ca(OH)}_2$ — The brackets around OH are essential — they show that the subscript 2 applies to the entire hydroxide ion (both the O and the H), not just the H atom.
4. 1 × (2+) + 2 × (1−) = 2+ − 2 = 0 — The overall charge is zero, confirming the formula is correct. This verification step takes seconds and prevents errors.

Answer: $\text{Ca(OH)}_2$

Worked example: Writing the Formula of Ammonium Sulfate

Write the correct formula for ammonium sulfate, given that ammonium is $\text{NH}_4^+$ and sulfate is $\text{SO}_4^{2-}$.

1. $\text{NH}_4^+$ (charge: 1+) and $\text{SO}_4^{2-}$ (charge: 2−) — Note that both ions here are compound ions — this is unusual but entirely valid. The same neutrality rule applies.
2. Two $\text{NH}_4^+$ provide 2+ charge. One $\text{SO}_4^{2-}$ provides 2− charge. Ratio is 2: 1. — You need two ammonium ions to balance the 2− charge of one sulfate ion.
3. $(\text{NH}_4)_2\text{SO}_4$ — Brackets around $\text{NH}_4$ are required because the subscript 2 must apply to the whole ammonium ion. The sulfate ion does not need brackets here because there is only one of them.
4. 2 × (1+) + 1 × (2−) = 2+ − 2 = 0 — Overall charge is zero. The formula $(\text{NH}_4)_2\text{SO}_4$ is correct.

Answer: $(\text{NH}_4)_2\text{SO}_4$

6. Exam Technique — Writing Ionic Formulae and Explaining Ionic Bonding

AQA examiners test ionic bonding and compound ions in predictable ways. Knowing the traps in advance is the difference between full marks and dropping marks on questions you actually understand.

One more examiner insight: when a question asks you to 'explain' rather than 'describe', you must include a reason. 'The ions are attracted to each other' scores fewer marks than 'the oppositely charged ions experience strong electrostatic attraction, which requires a large amount of energy to overcome.' The second version gives the observation and the scientific principle — that is what 'explain' demands.

7. Consolidation — Bringing It All Together

You have now covered both specification requirements for this topic. Ionic bonding is the electrostatic attraction between oppositely charged ions in a lattice — and that lattice is a giant, regular, three-dimensional arrangement where each ion is surrounded by ions of the opposite charge. The strength of those attractions determines the physical properties of the compound.

The five compound ions — sulfate ($\text{SO}_4^{2-}$), hydroxide ($\text{OH}^-$), nitrate ($\text{NO}_3^-$), carbonate ($\text{CO}_3^{2-}$), and ammonium ($\text{NH}_4^+$) — behave as single units in ionic compounds. When writing formulae, the rule is always the same: total positive charge must equal total negative charge, and brackets are required whenever more than one compound ion is needed.

These two specification points — the nature of ionic bonding in a lattice, and the formulae of compound ions — underpin a significant portion of AQA inorganic and physical chemistry. Every time you encounter an ionic compound in later topics, the same principles of charge balance and electrostatic attraction apply.