What you'll learn
This revision guide covers giant ionic lattices and their properties as specified in the AQA GCSE Chemistry specification. You'll understand how ionic compounds form three-dimensional structures, why they have high melting points, and when they conduct electricity. These concepts are fundamental to Paper 1 and regularly appear in 4-6 mark extended response questions.
Key terms and definitions
Ionic bonding — the electrostatic force of attraction between oppositely charged ions formed by the transfer of electrons from metal atoms to non-metal atoms.
Giant ionic lattice — a regular three-dimensional arrangement of alternating positive and negative ions held together by strong electrostatic forces in all directions.
Electrostatic attraction — the force between oppositely charged particles; in ionic compounds, this acts between positive metal ions and negative non-metal ions.
Lattice energy — the energy required to separate one mole of an ionic compound into gaseous ions; a measure of ionic bond strength.
Aqueous solution — a solution where water is the solvent, allowing ionic compounds to dissolve and their ions to move freely.
Electrical conductivity — the ability of a substance to allow the flow of electrical charge through the movement of charged particles.
Delocalised electrons — electrons that are not associated with a particular atom or bond but are free to move (note: ionic compounds do NOT have delocalised electrons, unlike metallic structures).
Dissociation — the process by which an ionic compound separates into its individual ions when dissolved in water or melted.
Core concepts
Formation of giant ionic lattices
Ionic compounds form when metal atoms transfer electrons to non-metal atoms, creating oppositely charged ions. These ions arrange themselves into a giant ionic lattice structure.
The formation process:
- Metal atoms lose electrons to form positive ions (cations)
- Non-metal atoms gain electrons to form negative ions (anions)
- Oppositely charged ions attract each other with strong electrostatic forces
- Ions pack together in a regular, repeating three-dimensional pattern
- Each ion is surrounded by ions of opposite charge
Common examples include:
- Sodium chloride (NaCl) — formed from sodium metal and chlorine gas
- Magnesium oxide (MgO) — formed from magnesium metal and oxygen gas
- Calcium fluoride (CaF₂) — formed from calcium metal and fluorine gas
The regular arrangement maximises attractive forces between oppositely charged ions while minimising repulsive forces between like-charged ions. This creates the most energetically stable structure.
Structure and arrangement in ionic lattices
Giant ionic lattices have distinctive structural features:
Three-dimensional structure: Ionic bonds extend in all directions throughout the structure, not just between pairs of ions. In sodium chloride, each Na⁺ ion is surrounded by six Cl⁻ ions, and each Cl⁻ ion is surrounded by six Na⁺ ions.
Regular repeating pattern: The lattice forms a crystalline structure with ions occupying fixed positions. This gives ionic compounds their characteristic crystal shapes — sodium chloride forms cubic crystals, while copper(II) sulfate forms blue triclinic crystals.
Ratio representation: The chemical formula represents the simplest whole number ratio of ions, not individual molecules. NaCl means there is one Na⁺ ion for every Cl⁻ ion throughout the entire structure, not discrete NaCl units.
Bond strength factors:
- Greater ionic charge increases electrostatic attraction (Mg²⁺O²⁻ has stronger bonds than Na⁺Cl⁻)
- Smaller ionic radius increases attraction (the closer ions get, the stronger the force)
These structural features directly determine the physical properties of ionic compounds.
Melting and boiling points of ionic compounds
Ionic compounds characteristically have high melting points and high boiling points.
Why melting points are high:
- Strong electrostatic forces exist between oppositely charged ions
- These forces act in all directions throughout the giant structure
- Large amounts of energy are needed to overcome these forces
- Ions must be separated from their fixed positions in the lattice
Typical melting points:
- Sodium chloride: 801°C
- Magnesium oxide: 2852°C
- Calcium fluoride: 1418°C
Factors affecting melting point:
Higher ionic charge = higher melting point
- MgO (Mg²⁺ and O²⁻) melts at 2852°C
- NaCl (Na⁺ and Cl⁻) melts at 801°C
- The 2+ and 2- charges create much stronger electrostatic attraction
Smaller ionic radius = higher melting point
- NaF melts at 993°C
- NaCl melts at 801°C
- F⁻ is smaller than Cl⁻, allowing ions to get closer together
The energy required to melt or boil an ionic compound is directly related to its lattice energy — compounds with stronger ionic bonding require more thermal energy to break apart.
Electrical conductivity in ionic compounds
The electrical conductivity of ionic compounds depends on whether the ions can move freely.
Solid ionic compounds do NOT conduct electricity:
- Ions are held in fixed positions in the lattice
- Ions cannot move to carry charge
- No flow of electrical current is possible
- The strong electrostatic forces keep ions locked in place
Molten (liquid) ionic compounds DO conduct electricity:
- When melted, the ionic lattice breaks down
- Ions become free to move throughout the liquid
- Mobile ions can carry electrical charge
- Positive ions move to the negative electrode (cathode)
- Negative ions move to the positive electrode (anode)
- This movement of charged particles constitutes an electric current
Dissolved ionic compounds DO conduct electricity:
- When dissolved in water, ionic compounds undergo dissociation
- Water molecules surround and separate the ions
- Ions become mobile in the aqueous solution
- Free-moving ions can carry charge through the solution
- The solution is called an electrolyte
This property distinguishes ionic compounds from covalent substances, which typically do not conduct electricity in any state because they do not contain charged particles.
Solubility of ionic compounds
Many ionic compounds dissolve in water, though solubility varies considerably.
Dissolution process:
- Water molecules are polar (slight positive and negative charges)
- Positive end of water molecules attracts negative ions
- Negative end of water molecules attracts positive ions
- Water molecules surround individual ions (hydration)
- Hydrated ions break away from the lattice and disperse in solution
Solubility patterns (required for GCSE):
Generally soluble:
- All Group 1 metal compounds (sodium, potassium, lithium salts)
- All nitrates (NO₃⁻ compounds)
- Most chlorides, bromides and iodides (except silver and lead halides)
- Most sulfates (except barium sulfate, calcium sulfate)
Generally insoluble:
- Most carbonates (except Group 1 and ammonium)
- Most hydroxides (except Group 1, barium hydroxide and calcium hydroxide which is slightly soluble)
Factors affecting solubility:
- Lattice energy (energy required to break apart the lattice)
- Hydration energy (energy released when ions are surrounded by water)
- If hydration energy exceeds lattice energy, the compound dissolves
- Compounds with very high lattice energies (like MgO) are generally insoluble
Not all ionic compounds dissolve equally — silver chloride is insoluble while sodium chloride is highly soluble, despite both being ionic.
Comparing ionic with other structures
Understanding ionic structures requires distinguishing them from metallic and covalent structures.
Ionic vs metallic:
- Both have high melting points
- Metals conduct electricity when solid; ionic compounds only when molten/dissolved
- Metals have delocalised electrons; ionic compounds have fixed ions
- Metals are malleable; ionic compounds are brittle
Ionic vs simple molecular (covalent):
- Ionic compounds have high melting points; simple molecular have low
- Ionic conduct when molten/dissolved; simple molecular do not
- Ionic are often soluble in water; simple molecular often are not
- Ionic form crystal structures; simple molecular form discrete molecules
Brittleness of ionic compounds:
- When force is applied, layers of ions may shift
- Like-charged ions align next to each other
- Repulsion between same charges causes the structure to shatter
- This explains why salt crystals break cleanly rather than bend
This brittleness distinguishes ionic compounds from metals, which can be hammered into shape because their structure can deform without breaking.
Worked examples
Example 1: Explaining properties (6 marks)
Question: Sodium chloride has a melting point of 801°C and does not conduct electricity when solid, but does conduct when molten. Explain these properties in terms of structure and bonding. [6 marks]
Mark scheme answer:
High melting point explanation (3 marks):
- Sodium chloride has a giant ionic lattice structure (1 mark)
- With strong electrostatic forces of attraction between oppositely charged Na⁺ and Cl⁻ ions (1 mark)
- Large amounts of energy are needed to overcome these strong forces throughout the structure (1 mark)
Conductivity explanation (3 marks):
- In the solid state, ions are held in fixed positions in the lattice (1 mark)
- Ions cannot move, so cannot carry charge (1 mark)
- When molten, ions become free to move and can carry electrical charge (1 mark)
Examiner note: This style of question requires you to link structure → bonding → property. Always mention the type of structure, the forces present, and how this explains the property.
Example 2: Comparing compounds (4 marks)
Question: Magnesium oxide has a much higher melting point (2852°C) than sodium chloride (801°C). Explain why, despite both being ionic compounds. [4 marks]
Mark scheme answer:
- Both have giant ionic lattice structures (1 mark)
- Magnesium oxide contains Mg²⁺ and O²⁻ ions / ions with 2+ and 2- charges (1 mark)
- Sodium chloride contains Na⁺ and Cl⁻ ions / ions with 1+ and 1- charges (1 mark)
- Greater charges in MgO produce stronger electrostatic attraction / require more energy to overcome (1 mark)
Alternative acceptable point: Magnesium and oxide ions are smaller than sodium and chloride ions, allowing closer approach and stronger attraction (would replace mark 4).
Examiner note: Comparative questions require you to state the similarity first, then explain the difference. Specific ionic charges must be mentioned.
Example 3: Predicting solubility (2 marks)
Question: Predict whether barium sulfate is soluble in water. Explain your answer. [2 marks]
Mark scheme answer:
- Insoluble (1 mark)
- Because most sulfates are soluble except barium sulfate / barium sulfate is an exception to the general sulfate rule (1 mark)
Examiner note: GCSE questions on solubility test whether you know the general rules and exceptions. You must state both the prediction AND reference the rule.
Common mistakes and how to avoid them
Confusing ionic with metallic bonding: Ionic compounds do NOT have delocalised electrons — they have ions held in fixed positions by electrostatic attraction. Only mention delocalised electrons when discussing metals.
Saying "ionic bonds are broken" when dissolving: When ionic compounds dissolve in water, the ions separate but individual ions remain intact (Na⁺ stays as Na⁺). The electrostatic forces between ions are overcome, not the ions themselves broken apart.
Claiming ionic compounds conduct because electrons move: Ionic compounds conduct electricity through the movement of ions (charged particles), not electrons. Be precise about what carries the charge in different circumstances.
Writing molecular formulae for ionic compounds: NaCl is NOT a molecule — it represents the ratio of ions in the lattice. Avoid saying "a molecule of sodium chloride" or "NaCl molecules."
Forgetting to link structure to properties: Simply stating "strong bonds" is insufficient. You must explain: giant ionic lattice → strong electrostatic forces in all directions → high melting point (complete chain of reasoning).
Mixing up molten and aqueous: Both allow ionic compounds to conduct electricity, but the processes are different. Molten = heated until liquid; aqueous = dissolved in water. Both result in mobile ions but through different mechanisms.
Exam technique for "Bonding, structure and properties of matter: giant ionic lattices and properties"
Command word precision: "Explain" requires reasoning (because/so/therefore), while "describe" requires observation only. For ionic compounds, "explain conductivity" needs the mechanism (ions free to move carry charge), not just "it conducts when molten."
Six-mark answers: Extended response questions on ionic bonding require: structure type identified, bonding forces described, link between structure and property 1, link between structure and property 2. Write in clear paragraphs using scientific terminology — include "giant ionic lattice," "electrostatic forces," and "oppositely charged ions."
Comparison questions: Always state the similarity first ("both have giant structures with strong forces") before explaining differences (charge magnitude, ionic size). Award marks often depend on explicit comparison language ("whereas," "however," "in contrast").
Using data: If given melting points or conductivity data, reference specific values in your answer. "Compound A has a higher melting point (2852°C compared to 801°C) because..." scores better than vague statements about "higher melting points."
Quick revision summary
Giant ionic lattices are three-dimensional structures of alternating positive and negative ions held by strong electrostatic forces throughout. They have high melting points because significant energy overcomes these forces. Ionic compounds conduct electricity only when molten or dissolved because ions must be free to move to carry charge — solid lattices have fixed ions. Many ionic compounds dissolve in water through hydration, though solubility depends on lattice and hydration energies. Ionic lattices are brittle due to repulsion when like-charged ions align.