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The enthalpy of an ideal gas is independent of its pressure or volume, and depends only on its temperature, which correlates to its thermal energy. Real gases at common temperatures and pressures often closely approximate this behavior, which simplifies practical thermodynamic design and analysis.
Ideal gases do not interact with each other (no intermolecular forces), so the enthalpy of an ideal gas is independent of pressure.
An ideal gas is a theoretical gas composed of many randomly moving point particles that are not subject to interparticle interactions. [1] The ideal gas concept is useful because it obeys the ideal gas law, a simplified equation of state, and is amenable to analysis under statistical mechanics.
For an ideal gas. ∂V ∂T = nR P. so Equation 9.4.2 becomes. (∂H ∂P)T = V − T(nR P) = 0. As we can see for an ideal gas, there is no dependence of H on P. 9.4: The Enthalpy of an Ideal Gas not declared. The enthalpy of an ideal Gas Is independent of pressure.
The ideal gas law is often written in an empirical form: where , and are the pressure, volume and temperature respectively; is the amount of substance; and is the ideal gas constant. It can also be derived from the microscopic kinetic theory, as was achieved (apparently independently) by August Krönig in 1856 [2] and Rudolf Clausius in 1857. [3]
When the volume of a system is constant, changes in its internal energy can be calculated by substituting the ideal gas law into the equation for ΔU. 5.3: Enthalpy. At constant pressure, heat flow (q) and internal energy (U) are related to the system’s enthalpy (H).
The internal energy of an ideal gas is therefore the sum of the kinetic energies of the particles in the gas. The kinetic molecular theory assumes that the temperature of a gas is directly proportional to the average kinetic energy of its particles, as shown in the figure below.
