Ionic Bonding and Ionic Compounds — AQA GCSE Chemistry

SPEC 4.2.1.2 4.2.1.3 IONIC

Ionic Bonding & Ionic Compounds

Electron transfer, dot-and-cross diagrams for Groups 1, 2, 6, and 7, the giant ionic lattice, and the precise language AQA demands — atoms react, ions form, chloride not chlorine.

How Ions Form

Ionic compounds form when a metal atom reacts with a non-metal atom. Outer-shell electrons are transferred from the metal to the non-metal. Both end up with the electron arrangement of a noble gas.

Metal atoms — LOSE electrons

Metals have a few electrons in the outer shell. It takes less energy to lose them than to gain enough to fill the shell.

Result: a positive ion (cation). Charge = group number.

Non-metal atoms — GAIN electrons

Non-metals have nearly full outer shells. It takes less energy to gain a few than to lose all the outer electrons.

Result: a negative ion (anion). Charge = (group number − 8).

Ion

A charged particle formed when an atom gains or loses electrons.

Ion Charge from Group Number

The charge on an ion can be predicted directly from where the element sits in the periodic table. Spec requirement: you must be able to do this for Groups 1, 2, 6, and 7.

GroupTypeOuter electronsActionIon formedExamples

1Metal1Lose 11+Na⁺, K⁺, Li⁺

2Metal2Lose 22+Mg²⁺, Ca²⁺

6Non-metal6Gain 22−O²⁻, S²⁻

7Non-metal7Gain 11−Cl⁻, F⁻, Br⁻

Pattern to memorise

Metals lose, non-metals gain, both end with a full outer shell (the electron arrangement of the nearest noble gas in Group 0). The ion is now stable.

01Worked Example 4 MARKS

Predict the charge on the ion formed by each element. Use the group number.

(a) Calcium (Group 2) (b) Bromine (Group 7)

(c) Oxygen (Group 6) (d) Lithium (Group 1)

(a)

Ca is a Group 2 metal → loses 2 outer electrons → Ca²⁺.

(b)

Br is a Group 7 non-metal → gains 1 electron → Br⁻ (the bromide ion, charge 1−).

(c)

O is a Group 6 non-metal → gains 2 electrons → O²⁻ (called the oxide ion, not "oxygen ion").

(d)

Li is a Group 1 metal → loses 1 electron → Li⁺.

Sign-error trap (P1 Jun24 Q04.4): a metal loses electrons, so its ion is positive (e.g. Zn²⁺ — never Zn²⁻). Always check the sign before writing the answer.

Electron Transfer: NaCl

Sodium chloride is the classic example. Watch the electron move — and watch the language change from atom to ion.

After transfer — IONS

Na atom loses its outer electron → Na⁺ ion with arrangement (2,8) — same as neon. Cl atom gains that electron → Cl⁻ ion with arrangement (2,8,8) — same as argon. Both ions now have full outer shells.

Dot-and-Cross Diagram Rules

The dot-and-cross diagram is a standard way to show electron transfer. Five rules — follow them every time.

#RuleWhy it matters

1Use dots for one atom's outer electrons; crosses for the other'sShows where each electron came from

2You may show only the outer shell — AQA accepts this"allow use of outer shells only" (P1 Jun21 MS)

3After transfer, draw each ion in square brackets with the charge outsideThe charge is the marking point; brackets are convention to show it's now an ion

4The metal ion's outer shell is now empty (the full shell underneath is now the outermost shell)Reflects the loss of electrons

5The non-metal ion's outer shell now has 8 electrons (full)Same arrangement as nearest noble gas

What scores marks (Jun21 MS)

For a 4-mark dot-and-cross: (1) correct ratio of atoms; (2) correct outer-electron count for each atom; (3) correct charge on metal ion (with bracket); (4) correct charge on non-metal ion (with bracket).

02Worked Example: NaCl 4 MARKS

Draw a dot-and-cross diagram to show what happens when atoms of sodium and chlorine react. Show only the outer-shell electrons.

M1

Sodium atom shown with 1 outer electron (1 dot) and chlorine atom shown with 7 outer electrons (7 crosses).

M2

1 sodium to 1 chlorine — the ratio (formula NaCl).

M3

Sodium ion drawn in square brackets with + charge; outer shell now empty.

M4

Chloride ion drawn in square brackets with − charge; outer shell now has 8 electrons (7 crosses + 1 dot from sodium).

Note: AQA "allows any combination of x, , o, e⁽⁻⁾ for electrons" — just be consistent: one symbol per atom.

When the Ratio isn't 1 : 1

The number of atoms involved depends on the charges, because the total positive charge must equal the total negative charge. The compound is always neutral.

CompoundRatioWhyFormula

NaCl1 Na : 1 Cl1+ balances 1−NaCl

MgCl₂1 Mg : 2 Cl2+ needs two 1− ionsMgCl₂

Na₂O2 Na : 1 OO²⁻ needs two 1+ ionsNa₂O

MgO1 Mg : 1 O2+ balances 2−MgO

CaCl₂1 Ca : 2 Cl2+ needs two 1− ionsCaCl₂

The balance rule

Total positive charge = total negative charge. For Mg²⁺ + Cl⁻: one 2+ charge needs two 1− charges to balance — so MgCl₂ not MgCl. The dot-and-cross diagram for MgCl₂ shows one Mg atom and two Cl atoms.

The mark for ratio (Jun21 Q04.4 MS)

For Na₂O, AQA explicitly awards a separate mark for showing "two sodium atoms to one oxygen atom". Get the ratio wrong — even with correct individual atoms drawn — and you lose this mark.

03Worked Example: MgCl₂ 4 MARKS

Magnesium reacts with chlorine to form magnesium chloride (MgCl₂). Draw a dot-and-cross diagram. Show only the outer-shell electrons.

A 1 : 2 ratio — one Mg atom transfers electrons to two Cl atoms.

M1

Mg atom shown with 2 outer electrons (2 dots); each Cl atom shown with 7 outer electrons (7 crosses).

M2

1 magnesium to 2 chlorines — the ratio (formula MgCl₂). Mg²⁺ needs two Cl⁻ ions to balance the charge.

M3

Magnesium ion drawn in square brackets with 2+ charge; outer shell now empty.

M4

Each chloride ion drawn in square brackets with − charge; outer shell now has 8 electrons (7 crosses + 1 dot transferred from Mg).

Same principle as NaCl, different ratio. Each Mg atom transfers two electrons — one to each of two Cl atoms. The same approach works for Na₂O (2 Na atoms each transfer 1 e⁻ to one O atom), MgO (1 Mg transfers 2 e⁻ to 1 O), and CaCl₂ (1 Ca transfers 2 e⁻, one to each of 2 Cl).

The Giant Ionic Lattice

Once ions form, they don't stay as isolated pairs. They pack together in a giant ionic lattice — a regular, repeating, three-dimensional arrangement of alternating positive and negative ions.

Giant ionic lattice

A regular 3D arrangement of oppositely charged ions held together by strong electrostatic forces of attraction acting in all directions.

Three things to remember

(1) Each ion is surrounded by ions of opposite charge. (2) Electrostatic forces act in all directions — not just between pairs. (3) The whole structure is held together by these forces — this is the ionic bonding.

Four Ways to Represent Ionic Compounds

The spec requires you to recognise an ionic compound from any of these four diagram types. NaCl is the spec's standard example — you need to know it in all four forms.

1. Dot-and-Cross

Shows: electron transfer, where each electron came from, charges on ions.

Missing: 3D arrangement, real positions in lattice.

2. Ball-and-Stick

Shows: 3D arrangement of ions, regular pattern, connections between ions.

Missing: ions aren't really "balls"; "sticks" don't really exist (no discrete bonds — the attraction acts in all directions); the model shows gaps, but in reality ions touch.

3. 2D diagram

Shows: alternating + and − pattern, ratio of ions, simple visual.

Missing: 3D structure entirely — gives no sense of forces in all directions.

4. 3D space-filling

Shows: realistic relative sizes, that ions touch, packing in 3D.

Missing: cannot see ions inside the structure, no electron information.

Spec requirement (4.2.1.3)

You must be able to describe the limitations of each representation. Each diagram type sacrifices some information to show others — no single diagram captures everything about an ionic compound.

Empirical Formula from a Diagram

Given a model or diagram of an ionic compound, you may be asked to work out its empirical formula — the simplest whole-number ratio of ions.

Three-step method

Count the positive ions shown in the diagram.

Count the negative ions shown.

Simplify to the smallest whole-number ratio — that's your formula.

Worked example: A diagram shows 8 calcium ions (Ca²⁺) and 16 chloride ions (Cl⁻). What is the empirical formula?

1
Count: 8 Ca²⁺ ions.
2
Count: 16 Cl⁻ ions.
3
Ratio: 8 : 16 = 1 : 2. Empirical formula: CaCl₂.

Sense check: Ca²⁺ + 2 × Cl⁻ = (2+) + (2−) = 0. Charges balance. ✓

Why "empirical"?

An ionic compound is a giant lattice with no fixed number of ions — the formula gives the ratio, not a count of atoms in a molecule. NaCl means "1 Na for every 1 Cl" — not "one molecule of NaCl".

Examiner Alerts: Language Traps

Three specific language traps from real examiner reports on ionic bonding. Each one loses easy marks if you slip.

Trap 1 — chlorine vs chloride (P1 Jun24 Q05.1)

Chlorine = the element (Cl, the green gas). Chloride = the negative ion (Cl⁻). Same with oxygen/oxide, sulfur/sulfide, nitrogen/nitride. Writing "chlorine ions" instead of "chloride ions" loses the mark every time.

Trap 2 — "atoms" missing in description (P1 Jun24 Q05.1)

Examiner report: "Many failed to include the word 'atoms' when referring to calcium and chlorine." Write: "calcium atoms lose electrons" — not just "calcium loses electrons". The atom/ion distinction is part of the description.

Trap 3 — atoms vs ions for full outer shells (P1 Jun24 Q05.1)

Examiner report: "It is the resulting IONS which have complete outer shells, not atoms." Atoms start with incomplete shells — that's why they react. After transfer, refer to ions, not atoms.

Weak

"Calcium reacts with chlorine. The chlorine gains electrons. Both atoms have full outer shells."

Missing "atoms" early on, says "chlorine" not "chloride", says "atoms" have full shells.

Strong

"Calcium atoms transfer two electrons to chlorine atoms. The Ca²⁺ ions and Cl⁻ chloride ions formed each have a full outer shell."

Atoms named, electron count given, ions correctly named with charges.

Common Errors

Students often writeCorrect versionWhy it matters

"Chlorine ions" / "oxygen ions"" Chloride ions" (Cl⁻) / " oxide ions" (O²⁻)The element ends in "-ine"; the ion ends in "-ide"

"Atoms have full outer shells" (after transfer)" Ions have full outer shells"Post-transfer they are ions, not atoms

"Molecule of NaCl""Formula unit of NaCl" / "the giant ionic lattice of NaCl"Ionic compounds are lattices, not molecules

Missing brackets / charge on dot-cross ionsAlways draw [ion]^charge — charge is the marking point; brackets are convention2 marks for Na₂O are charge-related (Jun21 MS)

"Ca²⁻" or similar sign errorMetal loses → positive (Ca²⁺); non-metal gains → negative (O²⁻)Sign confusion (P1 Jun24 Q04.4)

"Sodium and chlorine react…" (no "atoms")"Sodium atoms and chlorine atoms react…"Examiner specifically flagged missing "atoms" (Jun24 Q05.1)

Practice Question 1

Describe what happens when atoms of sodium and oxygen react to produce sodium oxide. Refer to electrons in your answer.4 MARKS

(Style based on AQA 8462/1H Jun 2021 Q04.4)

•

P Sodium oxide is an ionic compound formed when sodium atoms react with oxygen atoms.

•

ETwo sodium atoms each lose 1 electron — both transferred to one oxygen atom, which gains 2 electrons.

•

E Two Na⁺ ions form (each with a + charge) and one O²⁻ oxide ion forms (with a 2− charge). The ions are held together by strong electrostatic forces of attraction — this is the ionic bond. Formula: Na₂O.

Mark scheme breakdown (describe-question pattern): 1 = atoms named (Na atoms AND O atoms) 1 = electrons transferred (correct numbers, e.g. 2 from Na atoms total / 2 to one O atom) 1 = ions formed with correct charges (Na⁺ AND O²⁻) 1 = electrostatic attraction between oppositely charged ions. Common slip: writing "oxygen ions" instead of "oxide ions" — uses the element name where the ion name is needed.

Practice Question 2

Potassium reacts with fluorine to form potassium fluoride (KF). Draw a dot-and-cross diagram to show what happens. Show only the outer-shell electrons.4 MARKS

1

K is in Group 1 → 1 outer electron (1 dot). F is in Group 7 → 7 outer electrons (7 crosses).

2

Ratio: 1 K : 1 F (since 1+ balances 1−).

3

After transfer: [K]⁺ — outer shell empty — with + charge in square brackets. [F with 7 crosses + 1 dot]⁻ — full outer shell of 8 electrons — with − charge in square brackets.

Mark scheme alignment: 1 mark = K and F atoms with correct outer-electron counts 1 mark = 1:1 ratio 1 mark = K⁺ with bracket and charge 1 mark = F⁻ (fluoride ion) with bracket and charge. Watch: the F⁻ ion is called the fluoride ion — not "fluorine ion".

Practice Question 3

A diagram of an ionic compound shows 4 Mg²⁺ ions and 8 Cl⁻ ions.2 MARKS

(a) Work out the empirical formula. (b) Show that the charges balance.

(a)

Ratio: 4 Mg²⁺ : 8 Cl⁻ = 1 : 2. Empirical formula: MgCl₂.

(b)

Charge check: 1 × (2+) + 2 × (1−) = +2 − 2 = 0. The compound is neutral. ✓

Reminder: the empirical formula is the simplest whole-number ratio — ionic compounds are giant lattices, not discrete molecules. There's no "molecule of MgCl₂" — just the repeating ratio across the lattice.

Summary

Ion formation — the rule

Group 1 metal → 1+ ion (lose 1 e⁻)

Group 2 metal → 2+ ion (lose 2 e⁻)

Group 6 non-metal → 2− ion (gain 2 e⁻)

Group 7 non-metal → 1− ion (gain 1 e⁻)

Both metal and non-metal end with the electron arrangement of a noble gas

Dot-and-cross technique

Outer electrons only is allowed

Dots for one atom, crosses for the other

Brackets & charge on every ion

Ratio matters: balance positive and negative charges

Giant ionic lattice

Regular 3D arrangement of oppositely charged ions, held by strong electrostatic forces of attraction in all directions. The formula gives the simplest ratio of ions — no molecules exist in the lattice.

Examiner alert — the language

Use "atoms" before transfer and "ions" after. The non-metal ion is named "chloride" not "chlorine" (likewise oxide, sulfide, fluoride). Don't say "molecule of NaCl".

Ready for 2.03

Next up: Covalent Bonding — shared pairs, dot-and-cross diagrams for H₂, HCl, H₂O, NH₃, CH₄, CO₂, and the three errors AQA catches every year.