9701 Chemistry · Topic 4 · AS Level
States of Matter Cheat Sheet — A Level Chemistry 9701
Topic 4 is the bridge between bonding and physical properties: once you can name a structure you can predict a melting point, a solubility and a conductivity. This sheet covers the kinetic model, ideal gas behaviour and where real gases deviate from it, the four lattice types, and the structure–property arguments that earn full marks in explanation questions.
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What’s on this cheat sheet
States of Matter
01 · The three states
| State | Arrangement | Motion |
|---|---|---|
| solid | regular, touching | vibrate in place |
| liquid | close, disordered | slide past |
| gas | far apart, random | rapid, random |
Changes of state break or make intermolecular forces, not covalent bonds. Temperature stays constant during a change of state.
02 · Ideal gas equation
Units: p in Pa, V in m³, T in K. 1 dm³ = 1 × 10⁻³ m³, 1 cm³ = 1 × 10⁻⁶ m³, T(K) = °C + 273.
Rearranged with n = m ÷ M it gives M = mRT ÷ pV — the standard route to an Mr from gas data.
03 · Ideal vs real gases
Assumptions of an ideal gas: molecular volume negligible compared with the container, no intermolecular forces, collisions perfectly elastic, motion rapid and random, average kinetic energy proportional to T.
Deviation is greatest at high pressure and low temperature — molecules are close, so their own volume is significant and attractions pull them together, making the real volume smaller than predicted.
Gases with strong intermolecular forces or high Mr (NH₃, H₂O) deviate most; He and H₂ behave most ideally.
04 · Worked example
n = pV ÷ RT
n = (1.01 × 10⁵ × 1.52 × 10⁻⁴) ÷ (8.31 × 373)
n = 15.4 ÷ 3100 = 4.95 × 10⁻³ mol
M = 0.500 ÷ 4.95 × 10⁻³ = 101 g mol⁻¹
05 · Giant ionic lattices
NaCl — 6:6 coordination, each ion surrounded by six of the opposite charge. Strong electrostatic attraction throughout, so high melting point.
Brittle: a struck layer slips, like charges align and repel. Conducts only when molten or aqueous, when the ions are free to move.
06 · Giant covalent
Diamond — each C bonded to four others tetrahedrally, 109.5°. Very hard, very high melting point, does not conduct: no free electrons.
Graphite — hexagonal layers, each C bonded to three others at 120°, the fourth electron delocalised. Conducts along the layers; layers held by weak London forces so they slide — a lubricant.
SiO₂ — like diamond, each Si bonded to four O. High melting point, hard, insulating.
07 · Giant metallic
Cations in a sea of delocalised electrons. Bond strength rises with greater cationic charge and smaller radius: Na < Mg < Al.
Malleable and ductile — layers slide without breaking the bonding. Good conductor of heat and electricity in the solid.
08 · Simple molecular
Discrete molecules held by weak intermolecular forces — low melting and boiling points, non‑conductors, soluble in solvents of matching polarity.
Iodine forms a molecular lattice: strong I–I covalent bonds inside each molecule, weak London forces between them, so it sublimes readily.
09 · Comparing the four structures
| Type | m.p. | Conducts |
|---|---|---|
| giant ionic | high | molten / aqueous only |
| giant covalent | very high | no (graphite yes) |
| giant metallic | high | solid and molten |
| simple molecular | low | no |
Identify a structure from three clues together: melting point, conductivity in each state, and solubility.
10 · Intermolecular forces recap
Weakest to strongest: London (dispersion) in everything and stronger with more electrons; permanent dipole–dipole in polar molecules; hydrogen bonding where H is bonded to N, O or F.
These forces set the melting and boiling points of molecular substances. Hydrogen bonding explains water’s anomalous boiling point and why ice is less dense than water — the open lattice holds the molecules further apart.
11 · Worked example — gas mixture
total n = 0.50 mol
x(N₂) = 0.20 ÷ 0.50 = 0.40
p(N₂) = 0.40 × 150 = 60 kPa
p(O₂) = 150 − 60 = 90 kPa
Partial pressures always sum to the total.
12 · Changes of state
Melting and boiling need energy to overcome the forces holding particles together, so both are endothermic; freezing and condensing release it.
On a heating curve the plateaus are the changes of state: energy goes into separating particles, not into raising the temperature, so the kinetic energy is unchanged while the potential energy rises.
Boiling needs more energy than melting because the particles must be separated completely. Sublimation goes straight from solid to gas — iodine, solid CO₂.
13 · Worked example — gas density
Rearranging pV = nRT with n = m ÷ M:
ρ = m ÷ V = pM ÷ RT
ρ = (1.00 × 10⁵ × 0.0440) ÷ (8.31 × 298)
ρ = 4400 ÷ 2476 = 1.78 kg m⁻³
Denser than air (1.2 kg m⁻³), which is why CO₂ sinks and smothers a flame.
Marks lost here
— Using dm³ or cm³ in pV = nRT instead of m³, or °C instead of K.
— Saying covalent bonds break when a simple molecular solid melts.
— Saying real gases deviate at high temperature; it is low temperature and high pressure.
— Describing graphite’s layers as held by covalent bonds rather than London forces.
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States of Matter — Frequently Asked Questions
When does a real gas deviate most from ideal behaviour?
At high pressure and low temperature. At high pressure the molecules’ own volume is no longer negligible compared with the container, and at low temperature the intermolecular attractions are significant compared with the kinetic energy of the molecules.
Why does silicon(IV) oxide have a much higher melting point than carbon dioxide?
Silicon(IV) oxide is a giant covalent lattice, so melting means breaking strong covalent bonds throughout the structure. Carbon dioxide is simple molecular — only the weak instantaneous dipole–induced dipole forces between separate molecules need to be overcome.
What makes a metal conduct electricity and also malleable?
A sea of delocalised electrons carries charge through the lattice, and because the cations are all identical and non-directional the layers can slide over one another without breaking the metallic bonding.
