Introduction

This chapter discusses a thermodynamic relationship that is a basis for explaining spontaneous chemical reactivity, chemical equilibrium, and the phase behavior of chemical compounds. This relationship involves a thermodynamic state function that allows us to predict the direction of a chemical reaction at constant temperature and pressure. These constraints might seem annoyingly restrictive, because any method of prediction of spontaneity based on them must be less general than the Second Law of Thermodynamics, but as far as we are concerned the gains will outweigh the losses. Why is that? One reason is at any given time an individual organism is practically at uniform pressure and temperature (an ‘exception’ is discussed in one of the exercises). Another is that constant T and p are the very conditions under which nearly all bench-top biochemistry experiments are done. Yet another is that, although the total entropy of the universe must increase in order for a process to be spontaneous, evaluation of the total entropy change requires measurement of both the entropy change of the system and the entropy change of the surroundings. Whereas ΔSsystem can often be found without too much difficulty, albeit only indirectly, ΔSsurroundingscannot really be measured. How could one measure the entropy change of the rest of the universe? The subject of the present chapter provides a way around this difficulty.

A particularly clear example of the inadequacy of ΔSsystem to predict the direction of spontaneous change is given by the behavior of water at its freezing point.