The Gaseous State
Ideal-gas behaviour, partial pressures and deviations from ideality.
Before you begin
Ideal-gas behaviour, partial pressures and deviations from ideality.
Learning goals
- Ideal Gas Model and pV = nRT
- Dalton’s Law and Partial Pressures
- Real Gases and Deviations
Syllabus statements covered
- state the basic assumptions of the kinetic theory as applied to an ideal gas
- state and use the general gas equation pV = nRT in calculations, including the determination of Mr
- use Dalton’s Law to determine the partial pressures of gases in a mixture (see also Section 9)
- explain qualitatively in terms of intermolecular forces and molecular size: — the conditions necessary for a gas to approach ideal behaviour
- explain qualitatively in terms of intermolecular forces and molecular size: — the limitations of ideality at very high pressures and very low temperatures
Lessons
Work through them in order.
- Ideal Gas Model and pV = nRTState the ideal-gas assumptions and use pV = nRT, including to find relative molecular mass.
- Dalton’s Law and Partial PressuresSeparate a total pressure into component pressures before calculating.
- Real Gases and DeviationsIdentify which ideal assumption fails under extreme conditions.
Practise and check
Recommended nextThe Gaseous State topic check
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Topic reference
This topic covers the ideal gas equation, partial pressures and real gases. The lessons build the ideal gas model, apply it to mixtures, then explain when and why real gases deviate from it.
Be comfortable with:
- Mole and Avogadro Constant: amount of substance and molar mass.
- Gas Calculations with pV = nRT: reaction-linked gas stoichiometry once the gas amount is known.
- Intermolecular Forces and Properties: the attractions used in real-gas explanations.
Before you start, be able to convert °C to K, cm³ to dm³ or m³, and kPa to Pa.
Quick Reference
| Question evidence | First move | Essential check |
|---|---|---|
| one gas with p, V, T | use pV = nRT | T in K; p and V match R |
| mass and gas data | find n, then M = m/n | report Mᵣ without a unit |
| gas mixture | use pᵢ = xᵢ pₜₒₜₐₗ | ∑ xᵢ = 1 and ∑ pᵢ = pₜₒₜₐₗ |
| gas collected over water | subtract p_H₂O | use the dry-gas pressure in pV = nRT |
| low pressure, high temperature | gas approaches ideality | particles are far apart; attractions matter less |
| very low temperature | attractions become significant | pressure may be below the ideal prediction |
| very high pressure | molecular volume becomes significant | free volume is smaller than container volume |
Useful unit identity: 1 kPa dm³ = 1 J, so R = 8.31 is consistent with kPa and dm³ as well as Pa and m³.
Common Exam Traps
- Using °C instead of K.
- Mixing Pa with dm³ or kPa with m³ while retaining an incompatible value or unit for R.
- Treating Mᵣ as though it carries the molar-mass unit g mol⁻¹.
- Using total pressure for one component of a gas mixture.
- Forgetting to subtract the supplied water-vapour pressure from a wet-gas measurement.
- Claiming that real gases deviate most at high temperature and low pressure.
- Naming “non-ideality” without identifying the failed assumption and its effect on wall collisions or free volume.
- Importing this H2-only equation into an H1 8873 reacting-volume question.