AS Chemistry: Ideal Gases and Structure — Revision Notes

Condensed for the final weeks. For the full explanation, use the States of Matter: Ideal Gases and Structure study guide.

The ideal gas equation

pV = nRT        p in Pa
                V in m^3
                n in mol
                R = 8.31 J K-1 mol-1
                T in K

Conversions that cost marks: kPa × 1000 · cm³ ÷ 10⁶ · dm³ ÷ 1000 · °C + 273.

Combined with n = m/M, this gives a route to relative molecular mass from experimental gas data:

M = m R T / (p V)

Worked example. A 4.40 g sample of gas occupies 2.24 dm³ at 101 kPa and 273 K. Convert to SI: V = 2.24 × 10⁻³ m³, p = 101 000 Pa. n = pV/RT = (101 000 × 2.24 × 10⁻³) ÷ (8.31 × 273) ≈ 0.0997 mol. Mr = mass ÷ n = 4.40 ÷ 0.0997 ≈ 44 g mol⁻¹ (consistent with CO₂).

Assumptions of an ideal gas

1. Particles have negligible volume compared with the container.
2. There are no intermolecular forces between particles.
3. Collisions are perfectly elastic (no kinetic energy lost).
4. Particles are in constant random motion.

When real gases deviate

Deviation is greatest at high pressure and low temperature — precisely when the two key assumptions fail:

• High pressure → particles are close together, so their own volume is no longer negligible.
• Low temperature → particles move slowly, so intermolecular attractions become significant.

Gases behave most ideally at low pressure and high temperature. Gases with larger molecules and stronger intermolecular forces (e.g. NH₃, H₂O vapour) deviate more than small non-polar ones (He, H₂).

The four structure types

Structure	Particles	Forces broken on melting	m.p.	Conducts?	Solubility
Giant ionic	Ions	Strong electrostatic attraction	High	Only molten/aqueous	Often soluble in water
Simple molecular	Molecules	Weak intermolecular forces only	Low	No	Depends on polarity — “like dissolves like”: non-polar molecules dissolve in non-polar solvents, polar molecules (e.g. those that hydrogen bond) can dissolve in water
Giant covalent	Atoms	Many strong covalent bonds	Very high	No (except graphite)	Insoluble
Giant metallic	Ions + delocalised electrons	Attraction of ions to electron sea	High	Yes, solid and molten	Insoluble

Ice is a special case worth naming: its molecules are held together by hydrogen bonding — the strongest type of van der Waals’ force (a special case of permanent dipole–permanent dipole attraction) — rather than by id-id forces alone, giving it an unusually open structure — which is why ice is less dense than liquid water and floats.

The rule that answers most questions

Properties follow from what must be broken.

Melting a simple molecular solid overcomes only the intermolecular forces — the covalent bonds inside the molecules stay intact. That is why iodine melts at only 114 °C while diamond needs over 3500 °C, even though both contain covalent bonds (iodine also has a noticeable vapour pressure below this temperature, which is why it appears to sublime, but 114 °C is its true melting point).

Carbon allotropes

	Diamond	Graphite
Bonds per C	4	3
Structure	Tetrahedral 3D network	Hexagonal layers
Conducts	No	Yes — one delocalised electron per atom
Hardness	Hardest natural substance	Soft, layers slide

Deducing structure from properties

Worked example. An unknown white solid has a very high melting point, does not conduct as a solid, but conducts well once molten, and dissolves readily in water. A high melting point rules out simple molecular. Conducting only when molten (not as a solid) rules out giant metallic (conducts in both states) and giant molecular (except graphite, doesn’t conduct at all). Conducting when molten plus water solubility is the signature of mobile ions — the solid is giant ionic.

Exam traps

• Mixing units in pV = nRT — the single largest source of lost marks here.
• Saying covalent bonds break when a molecular solid melts.
• Claiming real gases deviate at high temperature — it is low temperature.
• Saying graphite conducts “because it is a metal”.
• Forgetting that ionic solids do not conduct until melted or dissolved.

Self-test

1. State the ideal gas equation with the correct units.
2. Give two assumptions of the kinetic theory of ideal gases.
3. Under what conditions do real gases deviate most, and why?
4. Why does iodine have a much lower melting point than diamond?
5. Explain why graphite conducts electricity but diamond does not.
6. A gas sample of mass 4.40 g occupies 2.24 dm³ at 101 kPa and 273 K. Calculate its Mr.
7. An unknown white solid has a very high melting point, doesn’t conduct as a solid, conducts when molten, and dissolves in water. Deduce its structure type.

Answers: 1. pV = nRT, with p in Pa, V in m³, n in mol, T in K, R = 8.31 J K⁻¹ mol⁻¹. 2. Any two: negligible particle volume; no intermolecular forces; perfectly elastic collisions; constant random motion. 3. High pressure and low temperature — particle volume becomes significant relative to the container, and intermolecular attractions are no longer negligible. 4. Iodine is simple molecular, so melting overcomes only weak intermolecular forces; diamond is giant covalent, requiring many strong covalent bonds to be broken. 5. Each carbon in graphite bonds to only three others, leaving one delocalised electron per atom free to move along the layers; in diamond all four outer electrons are used in covalent bonds. 6. n = pV/RT = (101 000 × 2.24 × 10⁻³) ÷ (8.31 × 273) ≈ 0.0997 mol; Mr = 4.40 ÷ 0.0997 ≈ 44 g mol⁻¹. 7. Giant ionic — a high melting point rules out simple molecular, conducting only when molten rules out giant metallic and giant molecular, and water solubility plus molten conductivity signals mobile ions being released.