D7.2 Molecular Formation of Solutions

Let us consider a solution of ethanol and water. Many common solutions contain these two components (usually with minor amounts of other substances as well). Ethanol and water are soluble in each other (what is known as “miscible”) in all proportions. For example, beer is typically about 3% alcohol (6% proof),114 wine about 6% alcohol (12% proof), and liquors such as whiskey or brandy are about 50% alcohol (100% proof). How do they dissolve into each other at the
molecular level, and why?

For a process to be thermodynamically favorable, the Gibbs (free) energy change (ΔG) associated with that process must be negative. However, we have learned that Gibbs energy change depends on both enthalpy (H) and entropy (S) changes in the system. It is possible to envision a wide range of situations – involving both positive and negative changes in H and S, and we have to consider the magnitudes of the enthalpy, the entropy and the temperature changes.

So what happens when we add a drop of ethanol to a volume of water? The ethanol molecules rapidly disperse and the solution becomes homogeneous. The entropy of the ethanol–water solution is higher than that of either substance on its own. In other words, there are more distinguishable arrangements of the molecules when they are mixed than when they are separate. Using simple entropic arguments we might, at least initially, extend the idea to include all solutions. Everything should be soluble in everything else, because this would to an entropy increase, right? Wrong. We know that this is not true. For example, oil is not soluble in water and neither are diamonds, although for very different reasons. So what are the factors influencing solution formation? We will see that some are entropic (involving ΔS) and some enthalpic (involving ΔH.)