Gas laws explorer
Measure the pressure of an ideal gas from its wall collisions as you change n, V and T, find the partial pressures in a mixture, and see when real gases stop behaving ideally.
Learning goals
- Ideal Gas Model and pV = nRT
- Real Gases and Deviations
- Dalton’s Law and Partial Pressures
1.0 mol of nitrogen in 25 cubic decimetres at 300 K. The gauge is taking its first reading.
- Pressure
- — kPa
- Mass of gas
- — g
- Mr of the gas
- —
- nRT/V
- 99.8 kPa
- pV/nRT
- —
- p(N₂)
- — kPa
- p(He)
- — kPa
- Total pressure
- — kPa
- Mole fraction of N₂
- 1.00
- pV/nRT for He
- 1.09
- pV/nRT for N₂
- 0.98
- pV/nRT for CO₂
- 0.38
Try this
0 of 5 doneHalve the volume of the gas at the same temperature and let the pressure settle. (not done yet)
The pressure doubles. In half the volume each particle hits the walls twice as often: p ∝ 1/V at constant n and T.
Double the absolute temperature at the same volume and let the pressure settle. (not done yet)
The pressure doubles. Faster particles hit the walls harder and more often: p ∝ T in kelvin, as pV = nRT says.
Mix helium with the nitrogen, then change the amount of helium. Compare the partial pressures with the total. (not done yet)
Each gas pushes on the walls as if it were alone, so p = p(N₂) + p(He), and p(N₂) = x(N₂) × p. The light, fast helium exerts the same pressure per mole as nitrogen.
Measure an unknown gas at two volumes or temperatures and work out its Mr from each reading. (not done yet)
Both readings give the same n = pV/RT, so the same Mr = mass ÷ n. Use p in Pa with V in m³ (or kPa with dm³) and T in kelvin.
Compare the real-gas curves at 330 K or below and at 800 K or above. (not done yet)
Real gases are most nearly ideal at high temperature and low pressure. Hot molecules move too fast for attractions to matter, and far-apart molecules take up a negligible share of the volume.
Your readings
| # | V / dm³ | p / kPa | 1/V / dm⁻³ | T / K | Remove |
|---|---|---|---|---|---|
| No readings yet. Set up a measurement, then record it. | |||||