DescriptionThe properties and rules of chemical equilibria apply to other reaction classes. Dissolution and precipitation and acid-base reactions are examples of systems in equilibrium. A solution becomes saturated once the rate that the solute dissolves becomes equal to the rate that it reforms a precipitate; at that rate the concentration of dissolved solute cannot increase. Because saturated solutions are in equilibrium, the concentration of dissolved solute and undissolved solid can be changed by applying stress to the system, thus shifting the equilibrium to one side or the other, according to Le Chatelier's principle. When a solution contains more than one equilibrium relationship, the concentrations of dissolved solute, and therefore the solubility of the solutes, can be affected or changed.
At A Glance
- Ksp, the solubility product constant, is an equilibrium constant that can be used to calculate molar solubilities.
- Solubility can be shifted by the common ion effect, which is a method of forcing a solute to precipitate out of a solution by adding a salt that has an ion in common with the solute. Solubility can also be shifted by a change in the acidity of a solution.
- Selective precipitation is a technique used to separate ions with the same charge and sufficiently different solubilities out of solution using the same reagent.
- The reaction of a Lewis acid (an electron-pair donor) with a Lewis base (an electron-pair acceptor) can form a complex ion.
- Equilibrium can be calculated by using the formation constant Kf.
- If a solution contains multiple solutes, the solutes can affect each other's equilibria, causing a substance that is insoluble on its own to dissolve or causing a substance that is soluble to precipitate out of solution at a concentration below its saturation point.
- The difference in the solubilities of strong and weak electrolytes is what is harnessed to purify or separate compounds by shifting the equilibrium of a specific equation in a system of multiple equilibria.