Gases, Kinetic Theory and the Mole Concept
What this domain tests
Mole calculations, ideal gas behaviour, and applying kinetic theory to explain deviations from ideality.
Core rules
- pV = nRT with p in Pa, V in m3, T in K and R = 8.314 J K−1 mol−1.
- Molar gas volume is 22.4 dm3 at s.t.p. (273 K, 101 kPa) and 24.0 dm3 at r.t.p. (298 K, 1 atm).
- Kinetic theory assumes negligible molecular volume, no intermolecular forces, random motion and elastic collisions. Real gases deviate most at high pressure and low temperature, where molecules are close and attractions matter.
- Dalton's law: total pressure is the sum of partial pressures; partial pressure = mole fraction × total pressure.
- For solutions n = cV with V in dm3; dilution obeys c1V1 = c2V2.
Worked example
Burning 2.00 g of a hydrocarbon gives 6.29 g CO2 and 2.57 g H2O. Moles CO2 = 6.29/44.0 = 0.143 mol, so 0.143 mol C. Moles H2O = 2.57/18.0 = 0.143 mol, so 0.286 mol H. Ratio C:H = 1:2, empirical formula CH2. If Mr = 42, the molecular formula is C3H6. Mass check: 0.143 × 12.0 + 0.286 × 1.0 = 2.00 g.
Traps
Substituting °C into the gas equation; pairing dm3 with R = 8.314; treating 22.4 dm3 as the r.t.p. value; ignoring the limiting reagent; forgetting that the hydrocarbon may contain oxygen, which combustion data alone cannot rule out.