A random question that's bugged me for many years.
In the equilibrium constant of any aqueous solution, if water is involved, the activity of water is always treated as 1. This is by definition. However, it is the case that reactants or products within an aqueous solution can exceed a concentration/activity of 1. That doesn't logically make sense to me - the solvent surrounds all the molecules of reagent and product and it just doesn't seem right that the activity of a reagent can be higher than that of the solvent when both are involved in the reaction.
Say, for instance, the reaction of a 2M solution of Manganese (II) with sodium hypochlorite in hydroxide to yield sodium permanganate, chloride and water. If we remove the spectator ions, the conventional equilibrium constant approximating the activity to the concentration will be of the form K=[Cl-]^5[H2O]^3[MnO4-]^2/[ClO-]^5[OH-]^6[Mn2+]^2 (If i have not been careless)
Apologies for the formatting. The point is that we will conventionally take the activity of water as 1 and then remove it from the equation. But here, this implies that water is less active than Mn2+, the latter of which is orders of magnitudes more sparse within the reaction mixture. It thus makes little intuitive sense why water would be less active than Mn2+ in such a solution. The concentration of H2O is always 55.6 (iirc), and this factor is always ignored. Why?
I did first consider the issue to lie with the approximation of activity to concentration. However, activity coefficients still remain high in concentrations of solutions upwards of 10 molar. Meanwhile, water, despite surrounding literally all the reagent molecules, is arbitrarily capped at 1.
Stackexchange was pretty useless in answering this question. As such I come to /sci/ without much hope, considering this is fucking /sci/. But you guys did manage to solve some difficult math shit, I think some basic pedagogical equilibria chemistry cannot be especially difficult.