Ionic compounds — AQA GCSE Combined Science
Test yourself on Ionic compounds with AQA GCSE practice questions.
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Ionic compounds explained
Ionic bonding occurs when a metal atom transfers electrons to a non-metal atom, forming positive and negative ions.
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These oppositely charged ions pack into a regular three-dimensional giant lattice, not separate molecules. The attraction between a positive ion and a negative ion is electrostatic, and because each ion is surrounded by ions of the opposite charge, the forces act in all directions throughout the lattice. For example, in sodium chloride each Na⁺ ion is attracted to six Cl⁻ ions and each Cl⁻ ion to six Na⁺ ions. This arrangement is called ionic bonding. The many strong attractions explain why ionic compounds have high melting and boiling points and are usually soluble in water.
The structure of sodium chloride can be represented in the following forms:
Sodium chloride is a giant ionic lattice, not a collection of separate NaCl molecules. Its structure can be represented in several ways, each showing different information. A dot-and-cross diagram shows electron transfer from sodium to chlorine, giving Na⁺ and Cl⁻ with full outer shells. A ball-and-stick model shows the regular three-dimensional arrangement of ions, with each Na⁺ surrounded by six Cl⁻ ions and each Cl⁻ by six Na⁺ ions. A space-filling model shows how the ions pack together and why the lattice is dense. A two-dimensional diagram or a formula unit such as NaCl shows the simplest ratio of ions, 1:1, not a molecule. Whichever representation is used, the key ideas are strong electrostatic attraction between oppositely charged ions, a giant lattice, and a high melting point because many strong bonds must be broken.
deduce that a compound is ionic from a diagram of its structure in one of the specified forms
An ionic compound diagram shows a giant lattice: a regular, repeating grid of ions held in fixed positions. To deduce that a compound is ionic, look for charged particles rather than neutral atoms or molecules. In dot and cross diagrams, each ion carries a charge label such as Na⁺ or O²⁻, and electrons are shown transferred from metal to non-metal, giving full outer shells. In ball and stick or space-filling models, the pattern repeats in three dimensions with each ion surrounded by oppositely charged neighbours, for example six around each ion in sodium chloride. The key clue is a giant regular arrangement of ions, not separate molecules. If the diagram shows shared pairs of electrons between neutral atoms, it represents a covalent substance instead.
describe the limitations of using dot and cross, ball and stick, two and three-dimensional diagrams to represent a giant ionic structure
Every model of a giant ionic structure is a simplification, so each has limitations. Dot and cross diagrams show electron transfer and charges clearly, but they usually show only a few ions, so they hide the huge repeating lattice and give no sense of scale or three-dimensional shape. Ball and stick models show the arrangement and relative positions of ions, but the sticks suggest directional bonds, which ionic lattices do not have, and the balls are far too small compared with the spaces between ions. Two-dimensional diagrams flatten a three-dimensional lattice, so they miss the way ions surround each other in all directions. Three-dimensional models are better but still cannot show the true scale, the strong electrostatic forces in every direction, or the fact that the lattice is effectively endless.
work out the empirical formula of an ionic compound from a given model or diagram that shows the ions in the structure.
An empirical formula gives the simplest whole-number ratio of ions in a compound. In an ionic lattice diagram, each ion is drawn with its charge, for example Na⁺, Mg²⁺, O²⁻ or Cl⁻. To work out the formula, count the ions of each type shown, or use the ratio in which they appear, then reduce that ratio to its simplest whole-number form. For example, a diagram showing two Na⁺ ions for every one O²⁻ ion gives Na₂O, because the 2:1 ratio is already simplest. If a diagram showed four Na⁺ ions and two O²⁻ ions, the ratio 4:2 simplifies to 2:1, so the formula is still Na₂O. The charges must balance overall: the total positive charge equals the total negative charge. Always write the cation first, then the anion, and use subscript numbers only where more than one ion is needed.
Students should be familiar with the structure of sodium chloride but do not need to know the structures of other ionic compounds.
Sodium chloride is the ionic compound whose structure you must know in detail. It forms a giant ionic lattice in which Na⁺ and Cl⁻ ions are arranged in a regular repeating pattern. Each Na⁺ ion is surrounded by six Cl⁻ ions, and each Cl⁻ ion is surrounded by six Na⁺ ions, giving 6:6 coordination. The ions are held together by strong electrostatic forces of attraction between oppositely charged ions, acting in all directions. This explains the high melting point of sodium chloride, because a large amount of energy is needed to overcome these forces. The regular arrangement also means the ratio of Na⁺ to Cl⁻ ions is 1:1, so the empirical formula is NaCl. You do not need to recall the detailed structures of other ionic compounds such as magnesium oxide or calcium chloride; you only need to apply the general ideas of ionic bonding to them.
Your focus
- Describe an ionic compound as a giant structure of ions.
- Explain that ionic bonding involves strong electrostatic forces of attraction between oppositely charged ions acting in all directions in the lattice.
- Relate the giant ionic lattice to observable properties such as high melting and boiling points.
Show all 18 objectives
- Draw and interpret a dot-and-cross diagram showing the formation of Na⁺ and Cl⁻ ions from sodium and chlorine atoms.
- Describe the giant ionic lattice of sodium chloride, including the 1:1 ratio and the six-fold coordination of each ion.
- Explain how the lattice structure and strong electrostatic attractions account for the high melting point of sodium chloride.
- Identify ions and their charges from a dot and cross, ball and stick or three-dimensional diagram.
- Explain how the regular repeating arrangement of ions shows a giant ionic structure.
- Compare an ionic diagram with a covalent diagram and justify which type of bonding is shown.
- Describe at least three limitations of common diagrams used to represent giant ionic structures.
- Explain why each limitation matters for understanding the real lattice.
- Evaluate which type of diagram best represents a giant ionic structure and justify the choice.
- Count each type of ion shown in a diagram of an ionic structure.
- Simplify an ion ratio to its simplest whole-number form.
- Write the empirical formula of an ionic compound with the cation first and correct subscripts.
- Describe the arrangement of ions in a sodium chloride lattice, including 6:6 coordination.
- Explain how strong electrostatic forces lead to the high melting point of sodium chloride.
- Apply general ionic bonding ideas to unfamiliar ionic compounds without recalling their specific structures.
Ionic compounds exam tips
Marking Points
- Ionic compounds form giant structures, meaning the lattice extends in three dimensions rather than consisting of small discrete molecules.
- Ions are held together by strong electrostatic forces of attraction between oppositely charged ions.
- The forces act in all directions in the lattice, so each ion is attracted to all nearby ions of opposite charge.
- The term ionic bonding describes this overall electrostatic attraction within the giant ionic lattice.
- A correct example, such as sodium chloride, shows Na⁺ and Cl⁻ ions arranged in a regular repeating pattern.
- The strength and number of these attractions explain the high melting and boiling points of ionic compounds.
- Sodium atoms lose one electron to form Na⁺ ions and chlorine atoms gain one electron to form Cl⁻ ions, giving full outer electron shells.
- The ions are held in a giant three-dimensional lattice by strong electrostatic attraction between oppositely charged ions.
- Each Na⁺ ion is surrounded by six Cl⁻ ions and each Cl⁻ ion by six Na⁺ ions in the regular lattice arrangement.
- A dot-and-cross diagram represents electron transfer and outer-shell electron configurations, while a ball-and-stick or space-filling model represents the three-dimensional lattice arrangement.
- The formula NaCl represents the simplest ratio of sodium ions to chloride ions, 1:1, and does not represent a discrete molecule.
- The strong electrostatic forces throughout the giant lattice explain the high melting point of sodium chloride.
- Identify that the diagram shows ions, indicated by charge labels such as Na⁺, Mg²⁺, Cl⁻ or O²⁻, rather than neutral atoms.
- Recognise a giant regular lattice that repeats in three dimensions, with each ion surrounded by oppositely charged ions.
- Link the transfer of electrons from a metal to a non-metal to the formation of full outer shells and charged ions.
- Distinguish ionic diagrams from covalent diagrams, where atoms share electron pairs and no ions are present.
- Use the ratio of ions in the diagram to support the deduction, for example one Mg²⁺ for every two Cl⁻ in magnesium chloride.
- Explain that dot and cross diagrams show only a small number of ions, so they do not convey the giant repeating nature of the lattice.
- State that ball and stick models use sticks that imply directional covalent bonds, whereas ionic attraction acts in all directions.
- Describe how two-dimensional diagrams flatten a three-dimensional lattice and so misrepresent the arrangement of ions.
- Note that models do not show the true scale, because ions and the distances between them are not drawn to scale.
- Recognise that even three-dimensional models cannot show the lattice as effectively infinite or fully represent the electrostatic forces.
- Count the number of each type of ion shown in the model or diagram accurately.
- Write the ion ratio in the order cation then anion, for example Na⁺ : Cl⁻.
- Simplify the ratio to the simplest whole-number terms by dividing all parts by their highest common factor.
- Check that the total positive charge equals the total negative charge in the simplified ratio.
- Write the empirical formula using element symbols and subscript numbers, for example MgCl₂, without showing the charges in the final formula.
- Describe sodium chloride as a giant ionic lattice of alternating Na⁺ and Cl⁻ ions.
- State that each ion is surrounded by six ions of opposite charge, giving 6:6 coordination.
- Explain that strong electrostatic forces of attraction act between oppositely charged ions in all directions.
- Link the strong forces to the high melting point of sodium chloride.
- State that the ratio of Na⁺ to Cl⁻ ions is 1:1, giving the formula NaCl.
Examiner Tips
- 💡Use the phrase 'giant ionic lattice' and state that the forces act in all directions to secure the structure mark.
- 💡When explaining a property such as a high melting point, link it to the many strong electrostatic forces throughout the lattice, not just one bond.
- 💡Draw or describe a two-dimensional slice of the lattice with alternating positive and negative ions to show the repeating pattern clearly.
- 💡When asked to represent sodium chloride, label the ions and their charges clearly, for example Na⁺ and Cl⁻, and state the 1:1 ratio.
- 💡If a question asks why sodium chloride has a high melting point, link the answer to many strong electrostatic attractions in the giant lattice, not to a single bond.
- 💡Use the representation named in the question; if it asks for a dot-and-cross diagram, do not substitute a ball-and-stick model.
- 💡Scan the diagram first for charge labels or square brackets around ion symbols, as these are the quickest evidence of ionic bonding.
- 💡State the deduction and then justify it with one feature of the diagram, such as the repeating lattice or the transferred electrons.
- 💡If the diagram shows neutral atoms sharing pairs, say so explicitly and explain that this indicates covalent bonding instead.
- 💡Name the specific model in each limitation, so the examiner can see which diagram you are criticising.
- 💡Contrast what the model does show with what it fails to show, for example charges shown but scale missing.
- 💡Use precise terms such as lattice, three-dimensional, scale and electrostatic forces rather than vague words like 'not accurate'.
- 💡Underline or list each ion type and tally the counts before writing any ratio.
- 💡After simplifying, multiply the charge on each ion by its subscript to confirm the positive and negative totals are equal.
- 💡If the diagram is large, work methodically across the lattice rather than guessing from a small region.
- 💡Use the terms giant ionic lattice, electrostatic attraction and 6:6 coordination when describing sodium chloride.
- 💡When explaining a property such as melting point, always link it back to the strength of the electrostatic forces and the energy needed to overcome them.
- 💡If asked about another ionic compound, apply the same general principles rather than trying to recall a specific structure you have not been taught.
Common Mistakes
- Describing ionic compounds as simple molecules: correct this by stating that they form a giant lattice of ions.
- Saying the forces act only between one pair of ions: correct this by explaining that each ion is attracted to all oppositely charged ions around it, in all directions.
- Confusing ionic bonding with covalent bonding or with intermolecular forces: correct this by linking ionic bonding specifically to electrostatic attraction between oppositely charged ions in a lattice.
- Writing NaCl as a small molecule with a single Na–Cl covalent bond; correct this by describing a giant ionic lattice with many Na⁺ and Cl⁻ ions.
- Drawing a dot-and-cross diagram with shared electrons instead of transferred electrons; correct this by showing Na⁺ with no outer electrons and Cl⁻ with eight outer electrons.
- Stating that sodium chloride conducts electricity when solid; correct this by explaining that ions are fixed in the lattice, so conduction occurs only when molten or dissolved.
- Treating a dot and cross diagram of ions as showing shared electrons; correction: in an ionic diagram the electrons are transferred, so each ion has a full outer shell and a charge label.
- Assuming any regular pattern means ionic; correction: check for charged particles, because some covalent substances also form regular giant structures.
- Reading a ball and stick model as separate molecules; correction: the sticks show attractions between ions in a continuous lattice, not covalent bonds between molecules.
- Saying ball and stick models show the ions touching; correction: the balls are much smaller than the spaces between ions, so the model is not to scale.
- Claiming dot and cross diagrams show the whole lattice; correction: they show only a few ions and omit the repeating three-dimensional pattern.
- Stating that the sticks in ball and stick models represent ionic bonds in fixed directions; correction: ionic attraction is non-directional and acts in all directions.
- Writing the ratio as it appears without simplifying, for example giving Na₄O₂ instead of Na₂O; correct by dividing all counts by their highest common factor.
- Including the ionic charges in the final formula, for example writing Mg²⁺Cl⁻₂; correct by writing the neutral formula MgCl₂ with no charge symbols.
- Reversing the order and writing the anion first, for example ClNa; correct by placing the cation first, as in NaCl.
- Describing sodium chloride as simple molecules rather than a giant ionic lattice; correct by referring to a regular arrangement of ions held by electrostatic attraction.
- Saying each ion is surrounded by four ions of opposite charge; correct by stating six, as in the 6:6 coordination of sodium chloride.
- Claiming that the ions are held by covalent bonds; correct by explaining that the bonding is ionic, with electrostatic attraction between oppositely charged ions.