1.5.2Unit 1. Structure, Bonding and Introduction to Organic Chemistry
Periodicity – Thermal Trends — IAL Edexcel Linear (Chemistry) Revision Notes & Topical Questions
5 min read 5 parts Lesson 16 of 36 Unit quiz
Part 1 of 5Pattern Across Period 3
Melting and boiling points depend on the type of structure and the strength of the forces that must be overcome.
Sodium, magnesium, aluminium
- Melting point rises from about 98 °C to 650 °C to 660 °C as metallic bonding strengthens.
Silicon
- The highest melting point in the period, about 1414 °C, because it is a giant covalent lattice with strong bonds throughout.
Phosphorus, sulfur, chlorine
- These are small simple molecules (P₄, S₈ and Cl₂) with low melting points of about 44 °C, 115 °C and −101 °C.
Argon
- The lowest melting point, about −189 °C, because it exists as single atoms with very weak London forces.
Part 2 of 5Explaining the Metals
Explaining the Metals
Metallic bonding
- Positive metal ions sit in a lattice held by electrostatic attraction to a sea of delocalised electrons.
Charge increases
- Na⁺, Mg²⁺ and Al³⁺ release one, two and three delocalised electrons per atom, so the electron sea is denser.
Ion size
- The ions get smaller as charge rises, so the charge density increases and the attraction becomes stronger.
Result
- More energy is needed to break the lattice, so the melting and boiling points rise from sodium to aluminium.
Part 3 of 5Explaining the Non-metals
1
Silicon forms a giant covalent structure in which every atom is joined to four others by strong covalent bonds.
2
Melting silicon requires breaking many strong covalent bonds, which needs a very large amount of energy.
3
Phosphorus, sulfur and chlorine are simple molecular, so melting only overcomes weak London forces between molecules and not the covalent bonds inside them.
4
London forces increase with the number of electrons in the molecule, so S₈ is higher than P₄, which is higher than Cl₂.
5
Argon is monatomic with only eighteen electrons, so its London forces are the weakest and it has the lowest melting point.
Part 4 of 5Trends in Other Groups and Periods
Trends in Other Groups and Periods
Period 2
- Lithium and beryllium are metals, boron and carbon form giant structures with very high melting points,
- and N₂, O₂, F₂ and Ne are small molecules or atoms with very low values.
Group 1
- Melting points fall down the group because larger ions and a single delocalised electron weaken the metallic bonding.
Group 17
- Melting points rise down the group from F₂ to I₂ as London forces strengthen with more electrons.
Exam wording
- Always name the structure first, then state what is overcome during melting, then compare strength.
| Element | Na | Mg | Al | Si | P₄ | S₈ | Cl₂ | Ar |
|---|---|---|---|---|---|---|---|---|
| Structure | Metallic | Metallic | Metallic | Giant covalent | Simple | Simple | Simple | Atomic |
Part 5 of 5Exam Checklist & Tips
1.5.2 – Key Points for the Exam
- Periodicity A pattern in properties that repeats across each period.
- Metallic bonding Attraction between positive ions and delocalised electrons.
- Giant covalent lattice A network of atoms linked by covalent bonds throughout the structure.
- London forces Weak forces between molecules caused by temporary and induced dipoles.
- Delocalised electrons Electrons that are free to move through a metallic lattice.
Common Mistakes
- Saying covalent bonds break when simple molecular substances melt;
- only the weak intermolecular forces are overcome.
- Explaining the high melting point of silicon by ionic or metallic bonding.
- Ignoring ion charge and size when comparing Na, Mg and Al.
- Claiming S₈ has weaker London forces than P₄ even though it has more electrons.
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