Life is a far-from-equilibrium phenomenon, so one might reasonably question the utility of equilibrium theory to biochemists. There are a few reasons why you should care about equilibrium. One is that we still do a lot of experiments in test tubes, and under those conditions, systems eventually go to equilibrium. Another reason is that some reactions, like acid–base reactions, equilibrate much faster than others, so they can be treated as being in equilibrium. There is another, more subtle reason: we saw earlier that reversible processes, those in which the system is constantly in equilibrium, are the most efficient ones possible. Equilibrium therefore sets a limit to the efficiency of chemical processes, an idea we will use when we discuss how one reaction can drive another forward.

What does ΔrGm mean?

The molar free energy of reaction, ΔrGm (as well as other Δ quantities in thermodynamics), is a funny quantity. We typically think of it as the difference in free energy between reactants and products, since that's how we calculate it. However, its meaning is a bit more subtle than that. To make things specific, imagine that we have a reaction

reactants → P + some other products

in a system held at constant temperature and pressure. ΔrGm represents the change in free energy when we convert one molar equivalent of reactants to one mole of P under specified conditions (concentrations, pressures etc.). In other words, the composition of the system should not change as our mole of P is made. However, if we carry out this conversion, we will change the quantities of reactants and products, and thus the composition.