Enthalpy: thermodynamic background
The first law of thermodynamics may be considered as a statement of the principle of the conservation of energy. A consequence of this is that a concept termed ‘internal energy’ must be introduced, if the behaviour of a gas is to be explained with reasonable exactness during processes of heat transfer. Internal energy is the energy stored in the molecular and atomic structure of the gas and it may be thought of as being a function of two independent variables, the pressure and the temperature of the gas.
We can consider heat being supplied to a gas in either one of two ways: at constant volume or at constant pressure. Since the work done by a gas or on a gas, during a process of expansion or compression, is expressed by the equation: work done = p dV, it follows that if heat is supplied to a gas at constant volume, no work will be done by the gas on its environment. Consequently the heat supplied to the gas serves only to increase its internal energy, U. If a heat exchange occurs at constant pressure, as well as a change in internal energy taking place, work may be done.
This leads to a definition of enthalpy, FI:
H=U + pV (2.19)
The equation is strictly true for a pure gas of mass m, pressure p, and volume V. However, it may be applied without appreciable error to the mixtures of gases associated with air conditioning.
It is desirable that the expression ‘heat content’ should not be used because of the way in which enthalpy is defined by equation (2.19). This, and the other common synonym for enthalpy, ‘total heat’, suggest that only the internal energy of the gas is being taken into account. As a result of this, both terms are a little misleading and, in consequence, throughout the rest of this book the term enthalpy will be used.
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