Predict properties from bonding and structure
Relate lattice structure to melting and boiling points, electrical conductivity and solubility without confusing unlike solids.
AS Level; Cambridge International Chemistry 9701, 2025–2027 syllabus. Keep pressure, volume, amount and absolute temperature consistent in gas calculations. Written practice is self-checked
Before you begin
- Rearrange a formula and convert between common metric volume units.
Work on paper. Try each question before revealing help, and explain why your method works.
Quick readiness check
Why can molten NaCl conduct electricity?
What you will learn
- Describe and predict melting and boiling points from structure
- Explain conductivity and solubility trends using the stated structure
The particles and attractions determine melting and boiling behaviour
A simple molecular solid often melts and boils at relatively low temperatures because melting or boiling overcomes intermolecular forces between molecules; the covalent bonds inside molecules remain intact.
Giant ionic lattices usually have high melting points because many strong electrostatic attractions must be overcome. Giant covalent structures usually have very high melting points because many covalent bonds extend through the network. Metallic melting points vary between metals, so use given data when comparing named substances.
Conductivity and solubility depend on mobile particles and solvent interactions
Solid ionic crystals do not conduct because their ions are fixed; molten ionic compounds and suitable aqueous solutions can conduct because ions move. Metals conduct as delocalised electrons move. Simple molecular substances generally do not conduct because they lack mobile charged particles. Graphite is a giant covalent exception with delocalised electrons along its layers; diamond does not conduct.
Solubility depends on how the solute interacts with the solvent. Many ionic substances dissolve in water, but this is not universal; simple molecular solubility varies with polarity and the solvent. Giant covalent networks are generally insoluble in water. Use evidence and named examples rather than treating a structure label as a universal solubility guarantee.
Pause and explain
What is usually overcome when a simple molecular liquid boils?
Worked example
Explain why solid NaCl does not conduct but molten NaCl can conduct.
Show the worked solution
- In solid NaCl, ions occupy fixed lattice positions.
- Fixed ions cannot carry charge through the solid.
- When molten, ions can move and carry charge.
Answer Solid NaCl does not conduct because its ions are fixed. Molten NaCl conducts because its ions are mobile charge carriers.
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Test this topic →From guided practice to a new situation
Write your answer and reasoning first. Use a hint only if you are stuck. The model solution is for self-checking; your written work is not automatically marked.
Compare conductivity in solid NaCl, molten NaCl and copper.
Give me a hint
For each material, identify whether charged particles are mobile.
Compare my reasoning
- Solid NaCl ions are fixed, so it does not conduct.
- Molten NaCl ions can move, so it conducts.
- Copper delocalised electrons move through the metal, so it conducts.
Solid NaCl does not conduct; molten NaCl conducts through mobile ions; copper conducts through delocalised electrons.
Look for these in your work
- Solid ionic state correct
- Molten ionic state correct
- Metal charge carriers identified
Explain why iodine melts at a lower temperature than diamond, using their structures.
Give me a hint
Compare forces between I₂ molecules with bonds throughout diamond.
Compare my reasoning
- Solid iodine is simple molecular.
- Melting iodine overcomes intermolecular attractions between molecules.
- Diamond is a giant covalent network, so melting requires disrupting many covalent bonds.
Iodine generally melts at a lower temperature because its molecular lattice is held by intermolecular forces, whereas diamond has covalent bonds throughout a giant network.
Look for these in your work
- Structures correctly classified
- Iodine attraction identified
- Diamond network bonds identified
Compare graphite and diamond conductivity and explain why sharing the same structure class does not give identical behaviour.
Give me a hint
Both are giant covalent, but compare their electrons.
Compare my reasoning
- Graphite contains delocalised electrons that can move along its layers.
- Diamond's electrons are localised in covalent bonds.
- Their different electron arrangements explain their different conductivity.
Graphite conducts because it has delocalised electrons; diamond does not conduct because its electrons are localised in covalent bonds. Both are giant covalent structures with different arrangements.
Look for these in your work
- Graphite's mobile carriers stated
- Diamond's localised electrons stated
- Structure class and property distinguished
Common mistakes
- Saying ions are mobile in solid NaCl
- Claiming graphite and diamond have identical conductivity
- Saying simple molecular boiling breaks covalent bonds
- Treating water solubility as universal for ionic solids
What can you explain now?
Why can molten NaCl conduct electricity?
Compare with the explanation
Its ions can move and carry charge
The molten lattice allows mobile ions.
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Make it stick
For each structure, what must move or be overcome for it to conduct, dissolve, melt or boil?
If you needed the solution, close it and solve the problem again from a blank page. A correct answer is stronger when you can explain the reason for each step.