Reversible and Irreversible Reactions

In an idealized irreversible reaction such as one done by expanding an ideal gas against zero pressure, no work will be done by or on the system so the:

w = 0

In the case of an ideal gas (whose molecules do not interact) there will be no change in internal energy either so:

ΔE = 0

since ΔE = q – w, in this irreversible reaction q = 0 also.

In a reversible reaction involving an ideal gas, ΔE still will equal zero, however, the pressure will be changing continuously and work (w) is a function of P, work done must be determined over the entire course of the reaction. This result in the following mathematical reduction:

w = RTln[V2/V1]

Since in this situation ΔE = 0, q = w. This demonstrates that some of the heat of the surroundings has to be absorbed by the system in order to perform the work of changing the system volume.

Reversible reactions differ from irreversible in that the former always proceeds infinitely slowly through a series of intermediate steps in which the system is always in the equilibrium state. Whereas, in the irreversible reaction no equilibrium states are encountered. Irreversible reactions are also spontaneous or favorable processes. Thermodynamic calculations do not give information as to the rates of reaction only whether they are favorable or not.