Heat as motion: gases, temperature and the first law
Temperature is a measure of the average kinetic energy of the particles in a substance. For an ideal monoatomic gas the link is exact: the average translational kinetic energy per molecule is (3/2)kT, where k = 1.381 × 10⁻²³ J/K is Boltzmann's constant and T is absolute temperature in kelvin. Remember the offset: 0 °C = 273.15 K, and every gas-law calculation must use kelvin — a classic failure is putting 20 °C straight into Charles's law and getting nonsense.
The ideal gas law PV = nRT joins pressure (pascals, 1 Pa = 1 N/m²; standard atmospheric pressure is 101 325 Pa), volume (m³), amount (moles) and temperature, with R = 8.314 J/(mol·K). Each limiting law is a slice of it: Boyle's law (PV constant at fixed T), Charles's law (V/T constant at fixed P), and Avogadro's law (V/n constant at fixed P, T). Real gases deviate at high pressure and low temperature, where molecules occupy space and attract each other — the van der Waals equation adds a repulsion term a/V² and a attraction term b to repair the ideal picture.
From these relations, thermodynamics extracts state functions:
- Internal energy U — the total microscopic energy of the sample.
- Enthalpy H = U + PV — the heat exchanged at constant pressure.
- Entropy S — the number of ways energy can be arranged; it increases in every spontaneous process.
The first law, ΔU = Q − W, is conservation of energy: the change in internal energy equals heat added minus work done by the gas. The second law explains why engines can never convert all heat into work — a Carnot engine running between Tc and Th achieves at most η = 1 − Tc/Th. With Tc = 300 K and Th = 900 K that ceiling is 67%, and real engines sit well below it.