The gas laws

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At **constant temperature** (and amount), the pressure of a gas is **inversely proportional** to its volume: $P_{1}V_{1} = P_{2}V_{2}$.

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Pressure is **directly proportional** to the **kelvin** temperature: $\dfrac{P_{1}}{T_{1}} = \dfrac{P_{2}}{T_{2}}$.

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$\dfrac{P_{1}V_{1}}{T_{1}} = \dfrac{P_{2}V_{2}}{T_{2}}$ — with T in **kelvin**. It is given in the data booklet.

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**T/K = θ/°C + 273** — always do this before a gas-law calculation.

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The particles have **no volume** of their own and there are **no forces** between them.

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At **high temperature** and **low pressure** — particles are far apart and fast-moving.

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At **low temperature** and **high pressure** — particle volume and intermolecular forces become significant.

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The pressure **halves** (Boyle's law: P ∝ 1/V).

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Only the **kelvin** scale starts at true zero (0 K), so only it gives the correct proportionality; °C would give wrong ratios.

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Because P ∝ kelvin T — at 0 K the particles would stop and the pressure would be **zero**.

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The **temperature** and the **amount** of gas; only P and V change.

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Use the combined gas law: $P_{2} = P_{1}\times\dfrac{V_{1}}{V_{2}}\times\dfrac{T_{2}}{T_{1}}$ (T in kelvin).