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General chemistry instructors are faced with a dilemma when introducing the topic of equilibrium constants: These constants are correctly written in terms of activities, yet activity is a complex topic better treated rigorously in a physical chemistry course than superficially in a general chemistry course. Thus, to introduce equilibrium calculations, it is necessary to use approximate forms for equilibrium constants. However, I find that some commonly used textbooks provide incorrect arguments leading to approximate equilibrium constants for aqueous reactions. Water is a reactant or a product in many of the reactions encountered in a general chemistry. For acid dissociation as an example, a common textbook argument leading to the equilibrium constant Ka begins by saying that an acid reacts with water according to the equation HA(aq) + H2O(l)  H3O+(aq)+ + A?(aq) and that there exists an equilibrium constant K for this reaction given by K = [H3O+][A?]/([HA][H2O]) The argument continues by stating that because the concentration of water is effectively constant during the reaction, it is possible to define a new equilibrium constant Ka = [H3O+][A?]/[HA] where Ka = K[H2O] To the extent that activities of H3O+, A?, and HA can be successfully approximated by their molar concentrations (with units omitted), the form of the Ka given above is correct. However, the argument leading to that form is incorrect. The reason that water does not appear in Ka is that the activity of water is assumed to very nearly equal to 1. (This assumption is correct to the degree that the solution is dilute.) In other words, K and Ka are effectively the same constant. They do not differ by a factor equal to the numerical value of water’s molar concentration (about 55.6). I find that most students enter physical chemistry courses with an entrenched notion that equilibrium constants are in general expressed in terms of molar concentrations. These students at first tend to view activities as rather mysterious “effective concentrations”. But such an understanding reverses the roles of activities and concentrations. In fact, equilibrium constants are defined in terms of activities, not concentrations. Only sometimes are numerical values of molar concentrations reasonable substitutes for activities. In a general chemistry course, it is possible to introduce the forms of aqueous reaction equilibrium constants without introducing misconceptions. The essential points are these: - All equilibrium constants are correctly expressed in terms of activities.
- In general chemistry, molar concentrations are used as approximate activities for solutes. The validity of this sometimes rough approximation will be explored further in later courses, particularly physical chemistry.
- Importantly, water activity usually does not appear in equilibrium constant expressions for reactions in aqueous solutions because the activity of water is near to 1 unless the solution is quite concentrat
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