Predicting Ion Charges & Formula Construction Revision Notes | A Level AQA Chemistry

1. Why Ion Charges Matter

Every ionic formula you write — from sodium chloride to iron(III) oxide — depends on one thing: knowing the charge on each ion. Get the charge wrong and the formula is wrong, and that costs marks across bonding, equations, and calculations.

The good news is that the Periodic Table encodes most of the information you need. Once you see the pattern, predicting ion charges becomes a rapid, reliable skill rather than a memory exercise.

2. Predicting Ion Charges from the Periodic Table

When a main-group element forms an ion, it loses or gains electrons to achieve a full outer shell — the same electron configuration as the nearest noble gas. The number of electrons lost or gained is determined directly by the element's group number.

For metals in Groups 1, 2, and 3, the outer-shell electrons are lost, producing cations with charges of $+1$, $+2$, and $+3$ respectively. For non-metals in Groups 5, 6, and 7, electrons are gained to complete the outer shell, producing anions with charges of $-3$, $-2$, and $-1$ respectively.

Notice that main-group elements in the same group typically form ions with the same charge. Sodium and potassium are both in Group 1, so both form $+1$ ions. Oxygen and sulphur are both in Group 6, so both form $-2$ ions. This pattern applies reliably to simple main-group ions in Groups 1, 2, 3, 5, 6, and 7; heavier p-block elements and transition metals can show variable charges and are treated separately.

Transition metals are a different case. They can form ions with more than one possible charge — iron forms both $\text{Fe}^{2+}$ and $\text{Fe}^{3+}$, for example. The charge cannot be predicted from group number alone; it must be stated explicitly using Roman numerals, as in iron(II) and iron(III).

There is also a set of common polyatomic ions — ions made of more than one atom — whose charges you need to know. These cannot be derived from group number, so they require direct recall.

Ammonium ($\text{NH}_4^+$) is the only common polyatomic cation at A Level. Every other polyatomic ion in the table above is an anion. Knowing these charges is a prerequisite for constructing formulae involving compounds such as calcium carbonate or aluminium sulphate.

3. Constructing Ionic Formulae

Now that you can predict ion charges, the next step is combining ions into a correct ionic formula. The governing rule is charge balance: the total positive charge must exactly cancel the total negative charge, giving an overall charge of zero.

The most reliable method is the cross-multiplication method, sometimes called the 'swap and simplify' approach. You swap the numerical values of the charges to become the subscripts, then simplify to the lowest whole-number ratio.

Worked example: Constructing the Formula for Magnesium Chloride

Construct the correct ionic formula for magnesium chloride.

1. $\text{Mg}^{2+}$ and $\text{Cl}^-$ — Magnesium is in Group 2, so it loses 2 electrons to form a $2+$ ion. Chlorine is in Group 7, so it gains 1 electron to form a $1-$ ion.
2. Charge on $\text{Mg}^{2+}$ is 2; charge on $\text{Cl}^-$ is 1. Swap: 2 chloride ions are needed per magnesium ion. — The numerical value of the magnesium charge (2) becomes the subscript for chloride. The numerical value of the chloride charge (1) becomes the subscript for magnesium.
3. $\text{Mg}_1\text{Cl}_2$ — A subscript of 1 is not written explicitly in chemical formulae.
4. $\text{MgCl}_2$ — The ratio 1:2 is already in its simplest form. No further simplification is needed.
5. $(+2) + 2 \times (-1) = +2 - 2 = 0$ — Total charge is zero, confirming the formula is correct.

Answer: $\text{MgCl}_2$

The verification step in Step 5 is worth doing in every question. It takes five seconds and catches errors before they cost marks. If the total charge is not zero, the formula is wrong.

Worked example: Constructing the Formula for Aluminium Oxide

Construct the correct ionic formula for aluminium oxide.

1. $\text{Al}^{3+}$ and $\text{O}^{2-}$ — Aluminium is in Group 3, forming a $3+$ ion. Oxygen is in Group 6, forming a $2-$ ion.
2. Swap the charge values: aluminium gets subscript 2, oxygen gets subscript 3. — The numerical value of the oxide charge (2) becomes the subscript for aluminium. The numerical value of the aluminium charge (3) becomes the subscript for oxygen.
3. $\text{Al}_2\text{O}_3$ — Check whether this ratio can be simplified. The highest common factor of 2 and 3 is 1, so no simplification is possible.
4. $2 \times (+3) + 3 \times (-2) = +6 + (-6) = 0$ — Total charge is zero. The formula $\text{Al}_2\text{O}_3$ is correct.

Answer: $\text{Al}_2\text{O}_3$

Notice that when the cross-multiplication gives a ratio that can be simplified, you must simplify it. For example, if you were combining $\text{Ca}^{2+}$ and $\text{O}^{2-}$, the raw cross gives $\text{Ca}_2\text{O}_2$, but this simplifies to $\text{CaO}$ because the ratio 2:2 reduces to 1:1.

4. Formulae Involving Polyatomic Ions

The charge balance principle works identically when polyatomic ions are involved. The only additional rule is the bracket convention: if you need more than one polyatomic ion, the whole ion goes in brackets before the subscript is added.

Worked example: Constructing the Formula for Calcium Nitrate

Construct the correct ionic formula for calcium nitrate.

1. $\text{Ca}^{2+}$ and $\text{NO}_3^-$ — Calcium is in Group 2, forming a $2+$ ion. Nitrate is a polyatomic anion with a $1-$ charge — this must be recalled from the polyatomic ion table.
2. Charge on $\text{Ca}^{2+}$ is 2; charge on $\text{NO}_3^-$ is 1. Swap: calcium gets subscript 1, nitrate gets subscript 2. — Two nitrate ions are needed to balance one calcium ion.
3. $\text{Ca(NO}_3)_2$ — Because there are 2 nitrate groups, the entire $\text{NO}_3$ unit is enclosed in brackets and the subscript 2 is placed outside. The subscript 1 on calcium is not written.
4. $(+2) + 2 \times (-1) = +2 - 2 = 0$ — Total charge is zero. The formula is correct.

Answer: $\text{Ca(NO}_3)_2$

Earlier, you saw that the group number directly predicts the charge for simple ions. That same logic now feeds directly into polyatomic formula construction — the calcium charge of $+2$ came straight from Group 2. The two skills are always used together.

5. Worked Exam-Style Practice

Exam questions on ionic formulae often embed the formula construction inside a broader question — for example, asking you to write an equation for a reaction, or to identify the compound formed. The method is always the same: identify the ions, determine their charges, balance, and verify.

Worked example: Constructing the Formula for Iron(III) Sulphate

Construct the correct ionic formula for iron(III) sulphate.

1. $\text{Fe}^{3+}$ and $\text{SO}_4^{2-}$ — Iron(III) means the iron ion has a charge of $+3$ — the Roman numeral gives the charge directly. Sulphate is a polyatomic anion with a $2-$ charge.
2. Charge on $\text{Fe}^{3+}$ is 3; charge on $\text{SO}_4^{2-}$ is 2. Swap: iron gets subscript 2, sulphate gets subscript 3. — Two iron(III) ions carry a total charge of $+6$; three sulphate ions carry a total charge of $-6$.
3. $\text{Fe}_2(\text{SO}_4)_3$ — Three sulphate groups require brackets around $\text{SO}_4$ with subscript 3 outside.
4. $2 \times (+3) + 3 \times (-2) = +6 + (-6) = 0$ — Total charge is zero. The formula $\text{Fe}_2(\text{SO}_4)_3$ is correct.

Answer: $\text{Fe}_2(\text{SO}_4)_3$

6. Exam Technique and Common Pitfalls

AQA typically awards a single mark for the correct formula in formula-construction questions. Showing ion charges as working is good practice and may help you spot errors before committing to a final answer, but formula-construction questions do not usually carry separate method marks. The value of showing working lies in self-checking, not in mark recovery.

When a question asks you to 'construct' or 'write' a formula, the examiner expects the simplest whole-number ratio with correct use of brackets. Formulae that are not in their simplest ratio, or that omit brackets around polyatomic ions, will not receive the accuracy mark.

One final distinction worth noting: the formula you construct represents the empirical formula of the ionic compound — the simplest ratio of ions. Ionic compounds do not exist as discrete molecules, so the formula does not represent a single molecule but rather the repeating unit of the lattice. AQA questions may ask you to explain this distinction in the context of ionic bonding.

Taken together, the two specification requirements — predicting ion charges from the Periodic Table and constructing ionic formulae — form a single, connected skill: the group number gives you the charge, and the charge balance rule gives you the formula. Practise both steps together on every example, and the process becomes automatic.