If you halve the volume at constant T and n, what happens to pressure?

Boyle's Law: P₁V₁ = P₂V₂ → pressure doubles.

At 300 K, 1 mol of gas in 10 L: what is the pressure?

P = nRT/V = (1)(0.08206)(300)/10 = 2.46 atm.

Why do real gases deviate from PV=nRT at very high pressures?

At high P, particle volume is not negligible and intermolecular forces matter (van der Waals corrections).

Pro 🔒~25 min

Gas Properties

Pressure, volume, temperature, and the ideal gas law

How it works

The ideal gas law PV = nRT relates pressure, volume, amount, and temperature of a gas. At the molecular level, pressure arises from particles colliding with container walls. Temperature is proportional to average kinetic energy. Boyle's Law (P ∝ 1/V at constant T,n) can be observed by changing volume. Charles's Law (V ∝ T at constant P,n) relates volume to temperature. Gay-Lussac's Law (P ∝ T at constant V,n) shows pressure increases with temperature. Deviations from ideal behavior occur at high pressure (small volume) or low temperature, where intermolecular forces become significant.

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Step-by-step

Use the sliders to adjust temperature, container volume, and number of gas particles.
Watch particles bounce around — faster at higher temperatures, more wall collisions in smaller volumes.
The real-time graph plots pressure data so you can verify gas law relationships.

Key formulas

$PV = nRT$ Ideal Gas Law: P = pressure (atm), V = volume (L), n = moles, R = 0.08206 L·atm/(mol·K), T = temperature (K)
$KE_{\text{avg}} = \frac{3}{2} k_B T$ Average kinetic energy per particle is proportional to temperature. k_B = 1.38 × 10⁻²³ J/K

Frequently asked questions

If you halve the volume at constant T and n, what happens to pressure?: Boyle's Law: P₁V₁ = P₂V₂ → pressure doubles.
At 300 K, 1 mol of gas in 10 L: what is the pressure?: P = nRT/V = (1)(0.08206)(300)/10 = 2.46 atm.
Why do real gases deviate from PV=nRT at very high pressures?: At high P, particle volume is not negligible and intermolecular forces matter (van der Waals corrections).