What is ionic bonding in the context of Edexcel IGCSE Chemistry?

Ionic bonding is a type of chemical bond that occurs when atoms transfer electrons to achieve a full outer shell of electrons, resulting in the formation of positively and negatively charged ions. This bond is typically formed between metals and non-metals, where metals lose electrons to become cations and non-metals gain electrons to become anions. The electrostatic attraction between these oppositely charged ions holds them together in an ionic compound.

How do you determine the formula of an ionic compound?

To determine the formula of an ionic compound, first identify the ions involved and their charges. The total positive charge must balance the total negative charge. For example, in sodium chloride (NaCl), sodium (Na) forms a +1 ion and chlorine (Cl) forms a -1 ion. Since their charges balance each other, the formula is NaCl. For compounds with different charges, use the criss-cross method to balance the charges and write the simplest ratio of ions.

Why do ionic compounds have high melting and boiling points?

Ionic compounds have high melting and boiling points because of the strong electrostatic forces of attraction between the oppositely charged ions in their lattice structure. These forces require a significant amount of energy to overcome, resulting in high melting and boiling points. This is a key characteristic of ionic compounds, as covered in the Edexcel IGCSE Chemistry syllabus.

What are some common properties of ionic compounds?

Common properties of ionic compounds include high melting and boiling points, the ability to conduct electricity when molten or dissolved in water, and being generally soluble in water. These properties arise from the strong ionic bonds and the lattice structure of the compounds, which are important concepts in the Principles of Chemistry for Edexcel IGCSE.

How does ionic bonding differ from covalent bonding?

Ionic bonding differs from covalent bonding in that it involves the transfer of electrons from one atom to another, leading to the formation of ions, whereas covalent bonding involves the sharing of electron pairs between atoms. Ionic bonds typically form between metals and non-metals, while covalent bonds form between non-metals. Understanding these differences is crucial for students studying Chemistry at the Edexcel IGCSE level.

Why is "Saying 'electrons are shared' in an ionic bond" a common mistake in Ionic Bonding?

Why it happens: Confusing ionic with covalent bonding. How to avoid it: Ionic = electrons TRANSFERRED (from metal to non-metal, complete loss / gain). Covalent = electrons SHARED. The two are mutually exclusive at IGCSE level. Source: 4CH1 Examiner Reports 2022-2024

Why is "Saying 'electrons flow' to explain ionic conductivity" a common mistake in Ionic Bonding?

Why it happens: Confusing ionic with metallic conduction. How to avoid it: In ionic compounds, the CHARGE CARRIERS are IONS, not electrons. The phrase to use is 'ions free to move' (in molten or aqueous state). 'Electrons free to move' applies only to metals. Source: 4CH1 Examiner Reports 2022-2024

Why is "Forgetting square brackets / charge labels on dot-and-cross diagrams for ions" a common mistake in Ionic Bonding?

Why it happens: Treating the diagram as a drawing of atoms, not ions. How to avoid it: After electron transfer, the species are IONS, not atoms. Draw them in square brackets with the charge as a superscript: [Na]⁺, [Cl]⁻, [Mg]²⁺, [O]²⁻. The brackets and charges are mark-bearing.

Why is "Saying MgO has a higher mp than NaCl because it 'has more electrons'" a common mistake in Ionic Bonding?

Why it happens: Vague reasoning that doesn't address the question. How to avoid it: The factor is the SIZE OF THE IONIC CHARGES. Mg²⁺ + O²⁻ have larger charges (2+ and 2−) than Na⁺ + Cl⁻ (1+ and 1−), so the electrostatic forces are stronger, more energy needed to overcome them, higher mp. Source: 4CH1 Examiner Reports 2022-2024

Why is "Confusing group number with charge magnitude (e.g. claiming oxygen is +6 because it is in Group 6)" a common mistake in Ionic Bonding?

Why it happens: Misremembering the metal-cation pattern. How to avoid it: Metals (Groups 1-3) LOSE electrons → charge magnitude = group number. Non-metals (Groups 5-7) GAIN electrons → charge magnitude = 8 minus group number. Group 6 → gains 2 → forms 2− ions, not 6+.

Principles of ChemistryEdexcel IGCSE Chemistry (4CH1)

Ionic Bonding

Q: Why is "Writing 'attraction' or 'bond' instead of 'electrostatic forces of attraction'" a common mistake in Ionic Bonding?

Why it happens: Lack of precise vocabulary. How to avoid it: The mark-scheme phrase is **electrostatic forces of attraction** (between oppositely-charged ions). Use this exact phrase every time — vague 'attraction' loses the mark.

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Ionic Bonding — Pearson Edexcel International GCSE Chemistry 4CH1 Study Notes (2026 onwards, Spec Issue 4)

How metals transfer electrons to non-metals, forming oppositely-charged ions held together by strong electrostatic forces. Predicting ion charges from group number, drawing dot-and-cross diagrams, and explaining the four characteristic properties of ionic compounds (high mp / bp, hard but brittle, soluble in water, conduct only when molten or dissolved).

What you’ll learn

Mapped to the Cambridge IGCSE 4CH1 syllabus (2026 onwards).

1.46 — Describe the formation of ions by electron transfer for elements in Groups 1, 2, 3, 5, 6 and 7.
1.47 — Recall the charges of ions formed by elements in Groups 1, 2, 3, 5, 6 and 7 of the periodic table.
1.48 — Use the formulae and charges of common ions to write the formulae of ionic compounds.
1.49 — Draw dot-and-cross diagrams to show the transfer of electrons during ionic bond formation, using NaCl, MgO, CaCl₂, Na₂O and Al₂O₃ as examples.
1.50 — Understand that an ionic bond is a strong electrostatic force of attraction between oppositely-charged ions.
1.53 — Recall that ionic compounds form a giant lattice structure containing oppositely-charged ions.
1.54 — Explain the high melting and boiling points, hardness, and solubility patterns of ionic compounds in terms of the giant ionic lattice and electrostatic forces.
1.55 — Explain the electrical conductivity of ionic compounds in their solid, molten and aqueous states.

Forming ions by electron transfer (spec 1.46, 1.47)

Metals lose; non-metals gain; both reach noble-gas configuration.

Atoms form ions to achieve a stable noble-gas electron configuration (full outer shell). Whether the atom loses or gains electrons depends on whether it is closer to the previous or next noble gas in the periodic table.

Metals (Groups 1-3): LOSE electrons → form POSITIVE ions (cations).

Group	Example	Outer e⁻	Loses	Ion
1	Na (2,8,1)	1	1 e⁻	Na⁺ (2,8)
1	K (2,8,8,1)	1	1 e⁻	K⁺
2	Mg (2,8,2)	2	2 e⁻	Mg²⁺
2	Ca (2,8,8,2)	2	2 e⁻	Ca²⁺
3	Al (2,8,3)	3	3 e⁻	Al³⁺

For example: $\text{Na} \rightarrow \text{Na}^+ + e^-$ — the single outer-shell electron is lost, leaving the previous shell (2, 8) — neon configuration.

Non-metals (Groups 5-7): GAIN electrons → form NEGATIVE ions (anions).

Group	Example	Outer e⁻	Gains	Ion
5	N (2,5)	5	3 e⁻	N³⁻ (2,8)
6	O (2,6)	6	2 e⁻	O²⁻ (2,8)
6	S (2,8,6)	6	2 e⁻	S²⁻ (2,8,8)
7	F (2,7)	7	1 e⁻	F⁻ (2,8)
7	Cl (2,8,7)	7	1 e⁻	Cl⁻ (2,8,8)

For example: $\text{Cl} + e^- \rightarrow \text{Cl}^-$ — the chlorine atom gains one electron, completing its outer shell to 8 electrons (argon configuration).

Quick rule. For metals, the charge magnitude equals the group number. For non-metals, the charge magnitude equals (8 minus the group number), with a negative sign. Group 4 elements usually share electrons (covalent bonding) rather than form ions, and Group 0 noble gases form no ions at all.

Metals LOSE electrons → cations.
Non-metals GAIN electrons → anions.
Charge magnitude = group number (for metals) or 8 − group number (for non-metals).
Both species achieve a noble-gas configuration.

See the full worked example for ionic bonding →

Writing ionic formulae from charges (spec 1.48)

Balance + and − charges.

An ionic compound is always electrically neutral — total positive charge = total negative charge. Use this rule to deduce the chemical formula from the ion charges.

Procedure.

Write the cation and anion with their charges.
Find the lowest common multiple (LCM) of the charge magnitudes.
The number of each ion needed = LCM ÷ that ion's charge magnitude.
Write the formula with the cation first and the anion second.

Worked examples.

Sodium chloride. Na⁺ + Cl⁻. LCM = 1. So 1 Na⁺ + 1 Cl⁻ → NaCl.
Magnesium oxide. Mg²⁺ + O²⁻. LCM = 2. So 1 Mg²⁺ + 1 O²⁻ → MgO.
Calcium chloride. Ca²⁺ + Cl⁻. LCM = 2. So 1 Ca²⁺ + 2 Cl⁻ → CaCl₂.
Sodium oxide. Na⁺ + O²⁻. LCM = 2. So 2 Na⁺ + 1 O²⁻ → Na₂O.
Aluminium chloride. Al³⁺ + Cl⁻. LCM = 3. So 1 Al³⁺ + 3 Cl⁻ → AlCl₃.
Aluminium oxide. Al³⁺ + O²⁻. LCM = 6. So 2 Al³⁺ + 3 O²⁻ → Al₂O₃.
Magnesium nitride. Mg²⁺ + N³⁻. LCM = 6. So 3 Mg²⁺ + 2 N³⁻ → Mg₃N₂.

Polyatomic ions require brackets when more than one is present:

Calcium hydroxide: Ca²⁺ + 2 OH⁻ → Ca(OH)₂.
Ammonium sulfate: 2 NH₄⁺ + SO₄²⁻ → (NH₄)₂SO₄.
Aluminium nitrate: Al³⁺ + 3 NO₃⁻ → Al(NO₃)₃.
Calcium carbonate: Ca²⁺ + CO₃²⁻ → CaCO₃ (no brackets, only one CO₃²⁻).

Total + charge = total − charge.
Use LCM method to balance charges.
Brackets around polyatomic ions when more than one is needed.

Dot-and-cross diagrams (spec 1.49)

Show only OUTER-SHELL electrons; ions in [brackets] with charge.

A dot-and-cross diagram shows the transfer of outer-shell electrons during ionic bond formation. Conventions:

Draw only the outer-shell electrons (inner shells are not relevant for bonding).
Use dots (•) for one element's electrons and crosses (×) for the other's. The choice is arbitrary but must be consistent.
After electron transfer, show the resulting ions in square brackets with the charge as a superscript: [Na]⁺, [Cl]⁻, [Mg]²⁺, [O]²⁻.
The transferred electrons appear inside the bracket of the anion (the species that GAINED them) — you can see they have been moved across.

Example 1 — sodium chloride (NaCl). Na atom (×) has 1 outer electron; Cl atom (•) has 7. The single × transfers to Cl. Result:

[Na]⁺ — empty outer shell shown (the M shell is gone, the 2,8 configuration is now the outer).
[Cl • • • • • • • ×]⁻ — 8 electrons in the outer shell (7 dots + 1 cross).

Example 2 — magnesium oxide (MgO). Mg (×) has 2 outer electrons; O (•) has 6. Both × electrons transfer to O. Result:

[Mg]²⁺
[O • • • • • • ××]²⁻

Example 3 — calcium chloride (CaCl₂). Ca (×) loses 2 electrons; each Cl (•) gains 1. Result:

[Ca]²⁺ in the middle.
TWO [Cl • • • • • • • ×]⁻ ions, one on each side.

Example 4 — sodium oxide (Na₂O). Each Na has 1 outer electron; O needs 2. So TWO Na atoms (each × shown) donate 1 electron each to ONE O atom. Result:

TWO [Na]⁺ ions.
ONE [O • • • • • • ××]²⁻ ion.

Example 5 — aluminium oxide (Al₂O₃). Each Al loses 3 electrons; each O gains 2. LCM = 6: 2 Al and 3 O are needed. The 6 transferred electrons fill the three oxide ions. Result:

2 × [Al]³⁺.
3 × [O]²⁻ (each with 8 outer electrons).

Mark-scheme tip. The square brackets and charges are MARK-BEARING. Don't draw "atoms with arrows" instead of "ions in brackets" — Edexcel mark schemes specifically require the brackets and labels.

Show only outer-shell electrons.
Use dots for one element, crosses for the other.
Ions go in square brackets with the charge labelled.
Inner shells are not drawn (not relevant for bonding).

See the full worked example for ionic bonding →

The giant ionic lattice (spec 1.50, 1.53)

3D regular arrangement held by strong electrostatic forces.

An ionic compound does NOT exist as isolated "molecules" of one cation + one anion (despite what a formula like NaCl might suggest). Instead, the ions arrange in a giant ionic lattice — a regular 3D pattern with billions of alternating + and − ions held together by electrostatic forces in every direction.

Sodium chloride lattice. Each Na⁺ ion is surrounded by 6 Cl⁻ neighbours (one above, one below, four in the same plane). Each Cl⁻ is surrounded by 6 Na⁺ neighbours in the same way. This 6:6 arrangement extends indefinitely in all three dimensions.

Why the lattice is so stable. Each ion experiences ELECTROSTATIC ATTRACTION not just to its 6 nearest neighbours but, more weakly, to every other oppositely-charged ion in the crystal. The net effect is an extremely strong "glue" holding the entire crystal together — far stronger than any individual ion-ion attraction would suggest.

Energy to overcome the lattice. To melt an ionic compound, you must supply enough thermal energy to overcome all these electrostatic forces simultaneously. This requires a lot of energy → high melting and boiling points.

Factors affecting lattice strength.

Charge magnitude. Larger charges → stronger forces (force ∝ product of charges). Mg²⁺O²⁻ is held by forces ~4× stronger than Na⁺Cl⁻ (2×2 vs 1×1).
Ionic radius / distance. Smaller ions can pack closer → stronger forces (force ∝ 1/distance²). Mg²⁺ < Na⁺ and O²⁻ < Cl⁻, so MgO ions are closer than NaCl ions.

Result: MgO mp (2852 °C) ≫ NaCl mp (801 °C) — both factors push in the same direction.

Giant ionic lattice = 3D regular arrangement.
Each ion surrounded by oppositely-charged neighbours in every direction.
Strong electrostatic forces of attraction → high mp / bp.
Larger charges + smaller ions = stronger lattice = higher mp.

See the full worked example for ionic bonding →

Properties of ionic compounds (spec 1.54, 1.55)

High mp / bp, hard but brittle, soluble in water, conducts when molten or aqueous.

Four characteristic properties of ionic compounds, each explained by the giant ionic lattice + electrostatic forces story:

1. High melting and boiling points. The lattice contains strong electrostatic forces of attraction between oppositely-charged ions, acting in every direction throughout the crystal. A large amount of energy is needed to overcome these forces and break the lattice apart. Hence ionic compounds are typically SOLID at room temperature with high mp / bp (e.g. NaCl mp 801 °C).

2. Hard but brittle. The lattice is held tightly by electrostatic attraction, so the solid resists scratching and indentation — it is hard. However, when a force is applied that causes layers of the lattice to slide past one another, ions of LIKE CHARGE may end up adjacent. The resulting like-charge repulsion causes the crystal to split / shatter along that plane — the solid is brittle. (Compare with metals, which are also hard but DUCTILE because metallic bonds don't depend on a fixed arrangement.)

3. Soluble in water; usually insoluble in organic solvents. Water is a POLAR molecule — its oxygen is slightly negative and its hydrogens slightly positive. When NaCl is placed in water:

Slightly-positive H atoms attract Cl⁻ ions.
Slightly-negative O atoms attract Na⁺ ions.
These attractions pull ions out of the lattice, surrounding each ion with a "shell" of water molecules → the lattice dissolves.

Non-polar solvents (petrol, oil) lack a strong dipole and cannot break the lattice — so ionic compounds are insoluble in them.

4. Electrical conductivity depends on state.

State	Conducts?	Reason
Solid	NO	Ions are fixed in lattice positions and cannot move to carry charge.
Molten	YES	Heat overcomes electrostatic forces; ions become free to move.
Aqueous (dissolved)	YES	Lattice broken by water; ions free to move through the solution.

In ionic compounds, the charge carriers are IONS (not electrons — that's metals). Always use the phrase "ions free to move" in your answer; the mark-scheme phrase "ions free to carry charge" is also credited.

Common Edexcel exam style. A table or figure of conductivity data; you must (a) identify the substance as ionic, (b) explain WHY the solid does not conduct and the molten / aqueous form does. This is worth 4 marks and is very high-frequency.

High mp / bp: strong electrostatic forces in the lattice.
Hard but brittle: shift of layers → like-charge repulsion → crystal shatters.
Soluble in water (polar) but not petrol (non-polar).
Conducts when molten or aqueous (ions free to move); not as solid.

See the full worked example for ionic bonding →

How it’s examined

Ionic bonding is a guaranteed Paper 1 topic on every 4CH1/1C. Most-common items: 5-6 mark extended response on the formation of an ionic bond (using a worked example such as NaCl, MgO or CaCl₂); 4-5 mark dot-and-cross diagram (or description); 4-mark conductivity-in-different-states explanation; 4-5 mark melting-point comparison (MgO vs NaCl or similar); 3-mark formula-from-charges derivation. Paper 2 (4CH1/2C) sometimes asks for the polar-water solvation mechanism or the like-charge-repulsion mechanism for brittleness. Worth 10-15 marks across both papers in a typical session.

Sources: Pearson Edexcel International GCSE Chemistry 4CH1 specification — Issue 4, September 2024 (valid for 2026 onwards); Pearson Edexcel 4CH1 Examiner Reports 2022-2024; 4CH1/1C May/Jun 2024 question paper and mark scheme; 4CH1/1C January 2024 question paper and mark scheme; 4CH1/1CR January 2024 question paper and mark scheme; 4CH1/2C May/Jun 2024 question paper and mark scheme. Last reviewed 2026-05-12.

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Step-by-step worked examples — Ionic Bonding

Step-by-step solutions to past-paper-style questions on ionic bonding, written exactly the way a tutor would explain them at the board.

1Forming an ionic bond between Na and Cl (5 marks, Paper 1)
Core• Adapted from 4CH1/1C Jan 2024 Q5• ionic bonding, dot-cross, spec-1.46
Question
Sodium (Z = 11, electron configuration 2,8,1) reacts with chlorine (Z = 17, electron configuration 2,8,7) to form sodium chloride. Describe how an ionic bond forms between a sodium atom and a chlorine atom. (5 marks)
Step-by-step solution
1. Step 1
  Identify what each atom needs to do (1 M). Sodium has 1 electron in its outer shell; it can achieve the stable, full noble-gas configuration of neon (2, 8) by LOSING its outer electron. Chlorine has 7 outer electrons; it can achieve the stable, full noble-gas configuration of argon (2, 8, 8) by GAINING 1 electron.
2. Step 2
  Electron transfer (1 A). The outer-shell electron of sodium is TRANSFERRED to the outer shell of chlorine. The transfer happens because both atoms can attain a more stable noble-gas configuration.
3. Step 3
  Identify the ions formed (1 A). Sodium becomes a positively charged ion: $\text{Na} \rightarrow \text{Na}^+ + e^-$ . Chlorine becomes a negatively charged ion: $\text{Cl} + e^- \rightarrow \text{Cl}^-$ .
4. Step 4
  Bonding (1 B). The oppositely-charged ions (Na⁺ and Cl⁻) are held together by strong electrostatic forces of attraction — this is the ionic bond.
5. Step 5
  Lattice (1 B). In solid NaCl, many Na⁺ and Cl⁻ ions arrange in a giant ionic lattice — a regular 3D structure where each Na⁺ is surrounded by 6 Cl⁻ and vice versa, held together by strong electrostatic forces in all directions.
Answer
Sodium LOSES its 1 outer electron (achieves Ne configuration, becomes Na⁺); chlorine GAINS 1 electron (achieves Ar configuration, becomes Cl⁻). The opposite charges attract by strong electrostatic forces — the ionic bond. Many Na⁺ + Cl⁻ pairs form a giant ionic lattice.
Examiner tip
Edexcel 5-mark mark scheme: 1 for 'noble gas configuration' identified; 1 for 'transfer of electron' (NOT 'sharing'); 1 for both ions named with correct charges; 1 for 'electrostatic forces of attraction'; 1 for 'lattice / regular 3D arrangement'. The phrase electrostatic forces of attraction is the critical mark-bearing keyword — vague phrases like 'attraction' or 'bond' alone won't score.
2Comparing melting points (5 marks, Paper 1)
Core• Adapted from 4CH1/1C May/Jun 2024 Q6• ionic properties, spec-1.54
Question
Magnesium oxide (MgO) has a much higher melting point (2852 °C) than sodium chloride (NaCl, 801 °C). Explain this difference in terms of the ions present and the forces between them. (5 marks)
Step-by-step solution
1. Step 1
  Both are ionic with the same lattice type (1 M). Both MgO and NaCl form giant ionic lattices with strong electrostatic forces between oppositely-charged ions.
2. Step 2
  Identify the ions (1 M). MgO contains Mg²⁺ and O²⁻ ions (charges 2+ and 2−). NaCl contains Na⁺ and Cl⁻ ions (charges 1+ and 1−).
3. Step 3
  Larger charges → stronger forces (1 A). The electrostatic force between two ions is proportional to the product of their charges. For MgO: (2)(2) = 4. For NaCl: (1)(1) = 1. The electrostatic forces in MgO are roughly four times stronger.
4. Step 4
  Smaller ionic radii in MgO (1 B). Mg²⁺ is smaller than Na⁺ (more positive charge contracts the electron cloud); O²⁻ is smaller than Cl⁻ (different period). Smaller distance → even stronger electrostatic attraction.
5. Step 5
  Link to melting point (1 B). Therefore MgO requires MUCH more energy to overcome these stronger electrostatic forces of attraction, so it has a much higher melting point than NaCl.
Answer
Both ionic with strong electrostatic forces in a giant lattice. MgO has 2+ and 2− ions (larger charges than Na⁺ / Cl⁻), and the ions are smaller, so electrostatic forces are MUCH stronger. More energy needed to overcome these forces → higher mp.
Examiner tip
Mark scheme awards both factors (charge AND ionic radius) but the larger-charges argument is the dominant one and credited first. Examiner reports note candidates often state 'MgO has more electrons' — this is irrelevant and scores 0. The factor is electrostatic FORCE strength, which depends on ionic charge and distance.
3Conductivity of ionic compounds (4 marks, Paper 1)
Core• Adapted from 4CH1/1CR Jan 2024 Q7• conductivity, spec-1.55
Question
Sodium chloride does not conduct electricity when solid but does conduct electricity when molten or when dissolved in water. Explain these observations. (4 marks)
Step-by-step solution
1. Step 1
  State the requirement for conduction (1 M). Electrical conductivity requires charge carriers that are free to move. In ionic compounds the charge carriers are the ions (not electrons).
2. Step 2
  Solid NaCl — no conduction (1 A). In a solid ionic lattice, the Na⁺ and Cl⁻ ions are fixed in positions by strong electrostatic forces. The ions cannot move and therefore CANNOT carry charge → no conduction.
3. Step 3
  Molten NaCl — conducts (1 A). When NaCl is heated above its melting point (801 °C), the electrostatic forces between ions are overcome and the ions are free to move. Mobile ions carry charge → conduction.
4. Step 4
  Aqueous NaCl — conducts (1 B). When NaCl dissolves in water, the lattice is broken apart and the Na⁺ and Cl⁻ ions are separated and free to move through the solution. Mobile ions carry charge → conduction.
Answer
Solid: ions fixed in the lattice, cannot move, no conduction. Molten or dissolved: ions free to move and carry charge, so conducts.
Examiner tip
Edexcel mark schemes use the exact phrase 'free to move' and 'carry charge' / 'carry current'. Candidates who write 'electrons flow' for an ionic substance lose the marks — in ionic compounds the charge carriers are ions, not free electrons (only metals have delocalised electrons). This distinction is tested most papers.
4Writing the formula of aluminium oxide from ion charges (4 marks, Paper 1)
Core• Adapted from 4CH1/1C May/Jun 2024 Q1• ionic formulae, spec-1.50
Question
Aluminium is in Group 3 and forms Al³⁺ ions. Oxygen is in Group 6 and forms O²⁻ ions. (a) Explain why aluminium forms a 3+ ion. (b) Deduce the formula of aluminium oxide, showing your reasoning in terms of the charges. (4 marks)
Step-by-step solution
1. Step 1
  (a) Al electron configuration is 2,8,3 (1 M). Aluminium has 3 electrons in its outer shell.
2. Step 2
  (a) Losing 3 electrons → Ne configuration (1 A). It loses these 3 outer electrons to achieve the stable noble-gas configuration of neon (2, 8). Loss of 3 electrons gives a charge of 3+.
3. Step 3
  (b) Balance charges (1 M). To make a neutral compound, total positive charge = total negative charge. The lowest common multiple of 3 and 2 is 6: need 2 Al³⁺ ions (total +6) and 3 O²⁻ ions (total −6).
4. Step 4
  (b) Write the formula (1 A). Formula: Al₂O₃.
Answer
(a) Al has 2,8,3 → loses 3 outer electrons to reach 2,8 (Ne) → forms Al³⁺. (b) 2 × Al³⁺ (+6) + 3 × O²⁻ (−6) → neutral. Formula = Al₂O₃.
Examiner tip
Standard 4CH1 formula-derivation question. Mark scheme awards: 1 for outer-electron count; 1 for 'noble-gas configuration'; 1 for charge-balancing reasoning; 1 for the correct formula. ECF applies to part (b) if part (a) is wrong.

Model Answers — Ionic Bonding

High-scoring sample answers for ionic bonding on the Pearson Edexcel IGCSE 4CH1 paper, with examiner-style notes mapping each response to the mark scheme and assessment objectives.

Question
4CH1/1C-style — ionic properties extended response6 marks
Sodium chloride is a typical ionic compound. (a) State THREE physical properties of sodium chloride. (b) For each property in (a), explain why it occurs in terms of the structure of sodium chloride and the forces involved. (6 marks)
Model answer
(a) Three physical properties.

High melting point (801 °C) and high boiling point (1413 °C).

Hard but brittle — does not deform under force.

Soluble in water but insoluble in most organic solvents (e.g. petrol).

Does not conduct electricity as a solid, but conducts when molten or dissolved in water.

(b) Explanations.

High mp / bp. Sodium chloride has a giant ionic lattice — a regular 3D arrangement of Na⁺ and Cl⁻ ions held together by strong electrostatic forces of attraction. To melt the solid, a large amount of energy is needed to overcome these strong electrostatic forces acting in every direction throughout the lattice. Hence the high melting and boiling points.

Hard but brittle. The lattice is held tightly by electrostatic attraction, making the solid hard. However, when a force is applied, the layers of ions can shift slightly so that ions of the same charge align next to each other. The resulting repulsion between like-charged ions causes the crystal to split / shatter along the plane of force. So it is hard but brittle.

Conductivity. In a solid, the ions are fixed in their lattice positions by electrostatic forces — they cannot move, so the solid does NOT conduct. When molten, or dissolved in water, the lattice breaks down and the ions become free to move. Mobile ions carry electric charge through the liquid → it conducts electricity.
Why this scores
Edexcel 6-mark mark scheme: 1 for each of three properties stated correctly; 1 for each explanation linking to structure / forces. Key keywords: 'giant ionic lattice', 'strong electrostatic forces of attraction', 'ions free to move'. Examiner reports flag the conductivity explanation as the most frequently incomplete — many candidates explain why the molten state conducts but forget to explain why the solid does NOT.
Question
4CH1/2C-style — dot-and-cross diagram for Na₂O5 marks
Draw and describe in words a dot-and-cross diagram showing the formation of sodium oxide (Na₂O) from sodium atoms (electron configuration 2,8,1) and oxygen atoms (2,6). State the formula and explain the electron transfer. (5 marks)
Model answer
Electron transfer. Each sodium atom (2, 8, 1) has one outer electron and loses it to achieve the stable Ne configuration (2, 8). An oxygen atom (2, 6) has six outer electrons and needs to gain TWO electrons to achieve the stable Ne configuration (2, 8).

So we need TWO sodium atoms to provide the two electrons required by ONE oxygen atom. The two outer-shell electrons (one from each Na) are transferred to the oxygen atom.

Description of the dot-and-cross diagram.

Each sodium atom is drawn first as Na with its single outer-shell electron shown as a "cross" (×).

The oxygen atom is drawn with its 6 outer-shell electrons shown as "dots" (•).

After electron transfer the diagram shows TWO sodium ions, each as [Na]⁺ (with an EMPTY outer shell — the previous M shell is now lost), and ONE oxide ion as [O •• •• ••×× ]²⁻ in square brackets with charge 2− shown.

The oxide ion has 8 outer-shell electrons: the original 6 dots PLUS the two crosses transferred from the two sodium atoms.

Ions formed. 2 × Na⁺ and 1 × O²⁻. Total charges: 2(+1) + (−2) = 0 (neutral compound).

Formula. Na₂O (sodium oxide).
Why this scores
Edexcel 5-mark mark scheme: 1 for identifying outer-shell electron counts (Na = 1, O = 6); 1 for stating 2 Na atoms needed; 1 for noble-gas configuration (Ne) achieved; 1 for square brackets + charges shown correctly; 1 for the formula Na₂O. The square brackets and charge labels on the ions are MARK-BEARING — candidates lose easy marks by omitting them.
Question
4CH1/2C-style — ion charges from periodic table6 marks
Explain, using examples, how the position of an element in the periodic table determines the charge on the ion it forms. Refer to Groups 1, 2, 3, 5, 6 and 7 in your answer. (6 marks)
Model answer
Elements form ions to achieve a stable noble-gas electron configuration (full outer shell). The group number of the element determines how many electrons are in the outer shell — and therefore how many must be lost or gained.

Groups 1, 2, 3 (metals — LOSE electrons → POSITIVE ions).

Group 1 (1 outer electron, e.g. Na, K) → LOSE 1 electron → +1 ions (Na⁺, K⁺).

Group 2 (2 outer electrons, e.g. Mg, Ca) → LOSE 2 electrons → +2 ions (Mg²⁺, Ca²⁺).

Group 3 (3 outer electrons, e.g. Al) → LOSE 3 electrons → +3 ions (Al³⁺).

Metals lose electrons because losing a few outer electrons is energetically easier than gaining 5, 6 or 7 electrons.

Groups 5, 6, 7 (non-metals — GAIN electrons → NEGATIVE ions).

Group 5 (5 outer electrons, e.g. N, P) → GAIN 3 electrons → −3 ions (N³⁻, P³⁻).

Group 6 (6 outer electrons, e.g. O, S) → GAIN 2 electrons → −2 ions (O²⁻, S²⁻).

Group 7 (7 outer electrons, e.g. F, Cl, Br, I — the halogens) → GAIN 1 electron → −1 ions (F⁻, Cl⁻, Br⁻, I⁻).

Non-metals gain electrons because gaining 1-3 is energetically easier than losing 5, 6 or 7 to reach the previous noble gas.

Group 4 and Group 0. Group 4 elements (C, Si) usually share electrons by COVALENT bonding rather than form ions. Group 0 (noble gases) already have full outer shells — they do not form ions under normal conditions.

Rule of thumb. Charge magnitude = number of outer electrons (Groups 1-3) OR 8 minus number of outer electrons (Groups 5-7).
Why this scores
Edexcel 6-mark mark scheme: 1 mark each for correctly stating the ion charge for Groups 1, 2, 3, 5, 6 and 7 with the noble-gas-configuration justification (1 mark each, six groups). The 'rule of thumb' summary is a memory aid — examiners credit the explicit reasoning more than the rule itself.
Question
4CH1/2C-style — extended explanation4 marks
Sodium chloride dissolves easily in water but does not dissolve in petrol (a non-polar organic solvent). Explain this observation in terms of intermolecular forces and ions. (4 marks)
Model answer
Water is a polar molecule. Water molecules (H₂O) have a permanent dipole — the oxygen is slightly negative ( $\delta^-$ ) and the hydrogens are slightly positive ( $\delta^+$ ).

Solvating ions. When solid NaCl is placed in water:

The slightly positive hydrogens of water molecules are attracted to the negative Cl⁻ ions.

The slightly negative oxygens of water molecules are attracted to the positive Na⁺ ions.

These attractions are strong enough to pull individual ions away from the lattice and surround each ion with a "shell" of water molecules. The lattice gradually breaks apart and dissolves.

Why petrol does not work. Petrol (a mixture of hydrocarbons) is a non-polar solvent. Its molecules have no permanent dipole and so cannot attract / stabilise the ions strongly enough to break the ionic lattice. The strong electrostatic forces in the NaCl lattice remain stronger than any interaction with the non-polar solvent, so the lattice does not break apart — NaCl is insoluble.

Generalisation. "Like dissolves like": polar / ionic substances dissolve in polar solvents (like water); non-polar substances dissolve in non-polar solvents (like petrol).
Why this scores
Edexcel 4-mark mark scheme: 1 for 'water is polar / has permanent dipole'; 1 for description of water molecules attracting individual ions; 1 for 'lattice breaks apart' or 'ions surrounded by water'; 1 for 'petrol is non-polar, cannot break lattice'. The keyword 'polar' is mark-bearing — candidates who say only 'water dissolves NaCl' without explaining why score poorly.

Key Definitions and Keywords — Ionic Bonding

Definitions to memorise and the exact keywords mark schemes credit for ionic bonding answers — sharpened from recent examiner reports for the 2026 Pearson Edexcel IGCSE 4CH1 sitting.

Ion
Examiner keyword
An atom (or group of atoms) that has lost or gained one or more electrons, giving it a net positive (cation) or negative (anion) charge.
Cation (positive ion)
Examiner keyword
A positively charged ion. Formed when a metal atom LOSES one or more electrons. Examples: Na⁺, Mg²⁺, Al³⁺, NH₄⁺.
Anion (negative ion)
Examiner keyword
A negatively charged ion. Formed when a non-metal atom GAINS one or more electrons. Examples: Cl⁻, O²⁻, OH⁻, NO₃⁻.
Ionic bond
Examiner keyword
The strong electrostatic force of attraction between oppositely-charged ions in an ionic compound. Formed by transfer of electrons from a metal to a non-metal.
Electrostatic force of attraction
Examiner keyword
The force between two charges of opposite sign. The strength of the force increases with larger charges and decreases with greater distance between the centres of the ions.
Giant ionic lattice
Examiner keyword
A regular 3D arrangement of alternating positive and negative ions held together by strong electrostatic forces in every direction. Each ion has multiple neighbours of opposite charge.
Noble gas configuration
Examiner keyword
An electronic configuration where the outer shell is full (2 for He; 8 for Ne, Ar, Kr...). This is a particularly stable arrangement. Ions form by gaining or losing electrons to achieve a noble gas configuration.
Dot-and-cross diagram
Examiner keyword
A diagram showing electron transfer or sharing during bond formation. Outer-shell electrons of one atom are shown as dots (•), the other as crosses (×). Square brackets and charge labels [X]ⁿ⁺ or [X]ⁿ⁻ are required for ions.

Common Mistakes and Misconceptions — Ionic Bonding

The traps other students keep falling into on ionic bonding questions — taken from recent Pearson Edexcel IGCSE 4CH1 examiner reports and mark schemes — and how to avoid them.

✕Saying 'electrons are shared' in an ionic bond
4CH1 Examiner Reports 2022-2024
Why it happens
Confusing ionic with covalent bonding.
How to avoid it
Ionic = electrons TRANSFERRED (from metal to non-metal, complete loss / gain). Covalent = electrons SHARED. The two are mutually exclusive at IGCSE level.
✕Saying 'electrons flow' to explain ionic conductivity
4CH1 Examiner Reports 2022-2024
Why it happens
Confusing ionic with metallic conduction.
How to avoid it
In ionic compounds, the CHARGE CARRIERS are IONS, not electrons. The phrase to use is 'ions free to move' (in molten or aqueous state). 'Electrons free to move' applies only to metals.
✕Forgetting square brackets / charge labels on dot-and-cross diagrams for ions
Why it happens
Treating the diagram as a drawing of atoms, not ions.
How to avoid it
After electron transfer, the species are IONS, not atoms. Draw them in square brackets with the charge as a superscript: [Na]⁺, [Cl]⁻, [Mg]²⁺, [O]²⁻. The brackets and charges are mark-bearing.
✕Saying MgO has a higher mp than NaCl because it 'has more electrons'
4CH1 Examiner Reports 2022-2024
Why it happens
Vague reasoning that doesn't address the question.
How to avoid it
The factor is the SIZE OF THE IONIC CHARGES. Mg²⁺ + O²⁻ have larger charges (2+ and 2−) than Na⁺ + Cl⁻ (1+ and 1−), so the electrostatic forces are stronger, more energy needed to overcome them, higher mp.
✕Writing 'attraction' or 'bond' instead of 'electrostatic forces of attraction'
Why it happens
Lack of precise vocabulary.
How to avoid it
The mark-scheme phrase is electrostatic forces of attraction (between oppositely-charged ions). Use this exact phrase every time — vague 'attraction' loses the mark.
✕Confusing group number with charge magnitude (e.g. claiming oxygen is +6 because it is in Group 6)
Why it happens
Misremembering the metal-cation pattern.
How to avoid it
Metals (Groups 1-3) LOSE electrons → charge magnitude = group number. Non-metals (Groups 5-7) GAIN electrons → charge magnitude = 8 minus group number. Group 6 → gains 2 → forms 2− ions, not 6+.

Practice questions

Exam-style questions with step-by-step worked solutions. Try one before checking the method.

Past paper style quiz

Get a report showing which sub-topics you've nailed and which ones still need work.

Video lesson

Short walkthrough of the concepts students most often get stuck on.

Ionic Bonding — frequently asked questions

The things students keep getting wrong in this sub-topic, answered.

Ionic Bonding

Short Study Notes in the form of Slides

Detailed Study Notes

What you’ll learn

How it’s examined

1Forming an ionic bond between Na and Cl (5 marks, Paper 1)

2Comparing melting points (5 marks, Paper 1)

3Conductivity of ionic compounds (4 marks, Paper 1)

4Writing the formula of aluminium oxide from ion charges (4 marks, Paper 1)

Ion

Cation (positive ion)

Anion (negative ion)

Ionic bond

Electrostatic force of attraction

Giant ionic lattice

Noble gas configuration

Dot-and-cross diagram

✕Saying 'electrons are shared' in an ionic bond

✕Saying 'electrons flow' to explain ionic conductivity

✕Forgetting square brackets / charge labels on dot-and-cross diagrams for ions

✕Saying MgO has a higher mp than NaCl because it 'has more electrons'

✕Writing 'attraction' or 'bond' instead of 'electrostatic forces of attraction'

✕Confusing group number with charge magnitude (e.g. claiming oxygen is +6 because it is in Group 6)

Practice questions

Past paper style quiz

Assess your understanding

Video lesson

Ionic Bonding — frequently asked questions