|II. General Solution to Equilibria Problems|
|II-1. Introduction||II-2. Equilibrium Constant||II-3. Le Chatelier's Principle||II-4. Reaction Quotient||II-5. General Solution|
A chemical specie will always exist in equilibrium with other forms of itself. The other forms may exist in extremely small amounts, and might even be undetectable, but they are always present. These other forms arise due to the inherent disorder of nature that we call entropy (it's impossible to be perfect). As an example, pure water consists of the molecular compound, H2O, and the dissociated ions H+ and OH-. Pure water at room temperature contains equilibrium concentrations of H+ and OH- of 1.0x10-7 M. The equilibrium reaction is:
At equilibrium, the concentrations (or partial pressures of gases) of the reactants and products are in a steady state, that is, they are not changing. However, a very important point to remember is that on the molecular level reactant species (atoms, molecules, or ions) are still forming products, and product species are returning to reactants. At equilibrium, the rate at which reactants go to products is equal to the reverse reaction rate of products going to reactants. The figure to the right shows a precipitate of solid PbCrO4 in equilibrium with Pb2+ and CrO42- ions in solution. At equilibrium the concentrations of ions in solution are constant. Pb2+ and CrO42- ions continue to form solid PbCrO4, and solid PbCrO4 continues to dissolve. Because the rate of precipitation and dissolution are the same, there is no change in the concentrations of the ions in solution. This equilibrium is represented by the reaction:
PbCrO4 (s) Pb2+(aq) + CrO42-(aq)
The right and left double arrow in this reaction equation indicates that the reaction is still occuring, but that the concentrations have reached equilibrium, i.e., a steady state.