Thermodynamics governs the relationships among heat, work, temperature, and energy. The Zeroth Law establishes the concept of temperature by stating that if two systems are each in thermal equilibrium with a third, they are in thermal equilibrium with each other. The First Law expresses conservation of energy: for a closed system, the change in internal energy equals the heat added minus the work done, \(\Delta U = Q - W\). The Second Law has two commonly cited statements. The Clausius statement says that heat cannot spontaneously flow from a colder body to a hotter body without external work, while the Kelvin-Planck statement says that no heat engine can convert all absorbed heat into work; some heat must always be rejected to a cold reservoir. The Third Law states that the entropy of a perfect crystal approaches zero as the temperature approaches absolute zero.
Two derived properties are particularly useful. Enthalpy, defined as \(H = U + PV\), combines internal energy with flow work and is convenient for analyzing open systems such as turbines and heat exchangers. Entropy is a measure of molecular disorder or energy dispersal, defined for a reversible process as \(dS = \delta Q_{rev}/T\). The Carnot efficiency, \(\eta = 1 - T_{cold}/T_{hot}\) (with temperatures in Kelvin), gives the maximum possible efficiency of any heat engine operating between two thermal reservoirs.
Thermodynamic processes are often idealized. An adiabatic process involves no heat transfer with the surroundings (\(Q = 0\)), while an isothermal process occurs at constant temperature. Specific heat capacity, \(c = Q/(m\Delta T)\), specifies the heat required to raise the temperature of a unit mass by one degree. For ideal gases, the behavior is captured by the ideal gas law, \(PV = nRT\), where \(P\) is pressure, \(V\) is volume, \(n\) is the number of moles, \(R\) is the universal gas constant, and \(T\) is absolute temperature.