✓ Revise Right course
Swap the numerical values of the ion charges to become subscripts, then simplify to the lowest whole-number ratio. Free AQA notes and 14 questions answered.
Predicting Ion Charges & Formula Construction is part of Bonding in A Level AQA Chemistry.
6 min
Reading time
68
Lessons in this topic
A Level AQA
Pathway
Reviewed by Revise Right
Content review
Swap the numerical values of the ion charges to become subscripts, then simplify to the lowest whole-number ratio. Free AQA notes and 14 questions answered.
AQA A-Level Chemistry mock draft, 26–31 Oct. Draft coverage: Chemistry 7405/1 · Paper 1. Reviewed for an original-practice sitting; check student availability before registering. Marking and results will follow later, after separate approval. Eligible new accounts receive a five-day Pro trial when signup email verification completes. Previous trials keep their original expiry. After expiry, new papers need Pro unless you have an explicitly confirmed free sitting offer for this event. That offer adds no Pro time or AI learning. Saved attempts and released results remain readable. No card or automatic charge.
Use this lesson to stay inside Bonding while connecting the detail back to the wider Physical chemistry topic.
Read the core explanations from Predicting Ion Charges & Formula Construction before testing yourself from memory.
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.
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.
| Group | Outer electrons | Electrons lost/gained | Ion charge | Example |
|---|---|---|---|---|
| 1 | 1 | Loses 1 | $+1$ | $\text{Na}^+$, $\text{K}^+$ |
| 2 | 2 | Loses 2 | $+2$ | $\text{Mg}^{2+}$, $\text{Ca}^{2+}$ |
| 3 | 3 | Loses 3 | $+3$ | $\text{Al}^{3+}$ |
| 5 | 5 | Gains 3 | $-3$ | $\text{N}^{3-}$, $\text{P}^{3-}$ |
| 6 | 6 | Gains 2 | $-2$ | $\text{O}^{2-}$, $\text{S}^{2-}$ |
| 7 | 7 | Gains 1 | $-1$ | $\text{F}^-$, $\text{Cl}^-$ |
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.
| Polyatomic ion | Formula | Charge |
|---|---|---|
| Hydroxide | $\text{OH}^-$ | $-1$ |
| Nitrate | $\text{NO}_3^-$ | $-1$ |
| Carbonate | $\text{CO}_3^{2-}$ | $-2$ |
| Sulphate | $\text{SO}_4^{2-}$ | $-2$ |
| Phosphate | $\text{PO}_4^{3-}$ | $-3$ |
| Ammonium | $\text{NH}_4^+$ | $+1$ |
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.
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.
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.
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.
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.
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.
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.
Answer: $\text{Fe}_2(\text{SO}_4)_3$
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.
Published by Revise Right. Use the official specification reference to check the requirements for your course. Content updated: 2026-09-05.
Move to the related lessons or back to the topic page when you need the wider sequence before exam-style practice.
The formula of aluminium oxide is Al₂O₃. Aluminium forms Al³⁺ and oxygen forms O²⁻, so the swap method gives 2 aluminium and 3 oxygen; the charges balance because 2 × (+3) + 3 × (−2) = 0.
A cation is a positively charged ion formed when an atom loses one or more electrons, with the charge equal to the number of electrons lost. Cations are formed by metals and hydrogen, for example Na → Na⁺ + e⁻.
Use the cross-multiplication or 'swap and simplify' method: swap the numerical values of the ion charges to become subscripts, then simplify to the lowest whole-number ratio. For aluminium oxide, Al³⁺ and O²⁻ give Al₂O₃, and the total charge 2 × (+3) + 3 × (-2) = 0 confirms it.
The formula of iron(III) sulphate is Fe₂(SO₄)₃. The Roman numeral gives iron a charge of +3 and sulphate is SO₄²⁻, so two iron(III) ions (+6) balance three sulphate ions (−6), with brackets around SO₄.
Group 5 non-metals gain 3 electrons to complete their outer shell, forming ions with a charge of −3, such as the nitride ion N³⁻ and P³⁻. Group 6 forms −2 ions and Group 7 forms −1 ions.
For main-group elements, the group number gives the charge. Metals in Groups 1, 2 and 3 lose their outer electrons to form +1, +2 and +3 cations, while non-metals in Groups 5, 6 and 7 gain electrons to form -3, -2 and -1 anions, to achieve the electron configuration of the nearest noble gas.
Main-group elements in the same group have the same number of outer electrons, so they lose or gain the same number of electrons to achieve a full outer shell. Sodium and potassium are both in Group 1 and form +1 ions, while oxygen and sulphur are both in Group 6 and form -2 ions.
Transition metals can form ions with more than one possible charge, so the charge cannot be predicted from group number alone. Iron forms both Fe²⁺ and Fe³⁺, so the charge is stated explicitly with Roman numerals, as in iron(II) and iron(III); iron(III) sulphate is therefore Fe₂(SO₄)₃.
Brackets are used in an ionic formula when more than one polyatomic ion is needed: the whole ion goes in brackets with the subscript outside. Calcium nitrate is Ca(NO₃)₂ because two nitrate ions balance one Ca²⁺ ion. A formula that omits these brackets, or is not in its simplest ratio, is wrong.
The swap and simplify method, or cross-multiplication method, swaps the numerical values of the ion charges to become the subscripts, then simplifies to the lowest whole-number ratio. For Mg²⁺ and Cl⁻ this gives MgCl₂, and the total charge is checked to be zero.
Calcium oxide is CaO because the swap method applied to Ca²⁺ and O²⁻ gives Ca₂O₂, and the ratio 2:2 simplifies to 1:1. Ionic formulae are always written in the simplest whole-number ratio of ions.
Ammonium, NH₄⁺, is the only common polyatomic cation at A Level. Every other common polyatomic ion, such as hydroxide OH⁻, nitrate NO₃⁻, carbonate CO₃²⁻, sulphate SO₄²⁻ and phosphate PO₄³⁻, is an anion.
The phosphate ion, PO₄³⁻, has a charge of −3. Like other polyatomic ions, its charge cannot be derived from group number, so it requires direct recall.
An ionic formula represents the simplest ratio of ions, the empirical formula of the compound, not a single molecule. Ionic compounds do not exist as discrete molecules, so the formula describes the repeating unit of the lattice.