The gaseous state
Use H2 Chemistry ideal-gas and partial-pressure models, solve gas calculations and explain when real gases depart from ideal behaviour.
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Before you begin
This topic is about how gases behave. It starts with the ideal gas model and the equation pV = nRT, applies it to mixtures of gases through partial pressures, then explains when and why real gases behave differently from the model.
Be comfortable with: the mole, gas calculations, intermolecular forces, and converting °C to K, cm³ to dm³ and kPa to Pa.
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 pressuresUse Dalton’s law and mole fractions to calculate partial pressures in gas mixtures and correct for water vapour in collected gas.
- Real gases and deviationsExplain how intermolecular attractions and molecular volume cause real-gas deviations, and why high temperature and low pressure favour ideal behaviour.
Practise and check
Or choose
Topic reference
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 |
Core knowledge to remember
- An ideal gas has particles of negligible volume, no intermolecular forces, random motion and perfectly elastic collisions.
- pV = nRT, with T in kelvin and R = 8.31, using Pa with m³ or kPa with dm³.
- The relative molecular mass of a gas is found from M = mRT/pV.
- Dalton’s law: the total pressure of a mixture is the sum of the partial pressures, and pᵢ = xᵢ pₜₒₜₐₗ.
- Real gases deviate from ideal behaviour at low temperature, where attractions lower the pressure, and at high pressure, where molecular volume raises it.
- Gases are most nearly ideal at low pressure and high temperature.
Common mistakes
- Using the Celsius temperature instead of kelvin.
- Mixing Pa with dm³, or kPa with m³, with the same value of R.
- Giving Mᵣ the unit g mol⁻¹.
- Using the total pressure for one gas in a mixture, or forgetting to subtract the water-vapour pressure for a gas collected over water.
- Saying real gases deviate most at high temperature and low pressure. That is where they are most nearly ideal.
- Naming “non-ideal behaviour” without saying which assumption fails and how it changes the pressure.