Explain Period 3 melting points and conductivity from structure
Connect the changing structures of Period 3 elements to their melting points and electrical conductivity.
AS Level; Cambridge International Chemistry 9701, 2025–2027 syllabus. Period 3 trends are qualitative unless supplied data are given. Equations show the stated products and conditions; observations and pH depend on concentration and experimental conditions. Written practice is self-checked
Before you begin
- Use electron configurations, oxidation numbers, balanced equations and ionic charges.
Work on paper. Try each question before revealing help, and explain why your method works.
Quick readiness check
Why do solid metals conduct electricity?
What you will learn
- Relate metallic bonding to Period 3 melting points
- Explain silicon's high melting point
- Compare charge carriers in metallic, giant-covalent and molecular elements
Metallic bonding strengthens from Na to Al
Na, Mg and Al form giant metallic lattices. Each contributes delocalised electrons to the metallic bonding model. Across these metals, the positive ions have increasing charge and the number of delocalised electrons per atom rises, strengthening attraction between the ions and electron sea; more energy is needed to melt the metal.
All three conduct in the solid state because their delocalised electrons can move through the lattice. Do not explain their conductivity by mobile ions; the ions remain at lattice positions.
Silicon and molecular non-metals have different structures
Silicon forms a giant covalent structure. Many strong covalent bonds must be overcome or rearranged for melting, so its melting point is high. It is a semiconductor rather than a metal; its conductivity is much lower than that of Na, Mg and Al and depends on conditions.
P₄, S₈ and Cl₂ are simple molecular substances. Melting mainly overcomes intermolecular attractions, not the covalent bonds within each molecule. They have low melting points and no mobile charged particles in the usual model, so they do not conduct electricity. S₈ has stronger intermolecular attractions than P₄ because its molecules are larger, giving sulfur a higher melting point than phosphorus.
Pause and explain
Why does silicon have a high melting point?
Worked example
Explain why silicon has a high melting point but does not conduct like aluminium.
Show the worked solution
- Silicon has a giant covalent network with many strong bonds.
- A large amount of energy is required to melt it.
- Unlike aluminium's delocalised metallic electrons, silicon has semiconductor behaviour and much lower conductivity.
Answer Silicon has a giant covalent structure, giving a high melting point, but it lacks the freely mobile metallic electrons responsible for aluminium's strong conductivity.
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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.
Explain the rise in melting point from Na to Al using structure and bonding.
Give me a hint
Compare ionic-core charge and delocalised electrons across the metals.
Compare my reasoning
- All three have giant metallic lattices.
- The metal ions have increasing charge and contribute more delocalised electrons.
- Attraction strengthens, so more energy is needed to melt the metal.
Metallic bonding strengthens from Na to Al as ion charge and delocalised-electron contribution increase, so melting point rises.
Look for these in your work
- Structure stated
- Across-metal change described
- Energy consequence linked
Compare the melting of Si with the melting of P₄ and explain the difference.
Give me a hint
One solid has a giant covalent network; the other contains separate molecules.
Compare my reasoning
- Si has a giant covalent structure with strong bonds throughout.
- P₄ is simple molecular; melting overcomes intermolecular attractions.
- The giant network requires much more energy to melt.
Si has a much higher melting point because strong covalent bonds extend through its giant structure; P₄ melting mainly overcomes weaker intermolecular attractions.
Look for these in your work
- Structures distinguished
- Bond types identified
- Melting-energy explanation correct
A student explains that solid magnesium conducts because Mg²⁺ ions move from atom to atom. Correct the model and contrast it with magnesium chloride solution.
Give me a hint
Separate charge carriers in a metal from those in an electrolyte.
Compare my reasoning
- In solid magnesium, metal ions stay in lattice positions.
- Delocalised electrons carry electrical current through the metal.
- In aqueous MgCl₂, mobile Mg²⁺ and Cl⁻ ions can carry charge.
Solid magnesium conducts by mobile delocalised electrons. MgCl₂ solution conducts by mobile ions; the charge carriers are different.
Look for these in your work
- Metal charge carrier corrected
- Solid ions fixed
- Electrolyte ions distinguished
Common mistakes
- Saying melting silicon breaks every covalent bond completely
- Calling all non-metals molecular
- Saying solid metal ions carry the current
- Treating silicon as either a typical metal or a perfect insulator
What can you explain now?
Why do solid metals conduct electricity?
Compare with the explanation
They contain mobile delocalised electrons
The electrons can move through the metallic lattice.
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Make it stick
How do structure and mobile charge carriers explain melting point and conductivity across Period 3?
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