A chemical solution is a homogeneous mixture where one substance, the solute, is evenly dispersed within another substance, the solvent. For example, when table salt dissolves in water, the salt is the solute and the water is the solvent. Solution equilibrium is reached when the concentration of the dissolved substance stops changing over time. This state represents the maximum amount of solute that can remain dissolved in the solvent under specific conditions.
Defining the Dynamic Nature of Equilibrium
Solution equilibrium is a dynamic process, not a state of rest. Two opposing actions occur continuously and simultaneously at the same rate. The first action is dissolution, the forward process where solute particles break away from the solid mass and move into the solvent. This involves solvent molecules surrounding and separating the solute particles.
The opposing action is crystallization, or precipitation, where dissolved solute particles collide with the undissolved solid and return to the solid state. This is the reverse reaction. Equilibrium is established when the rate of dissolution becomes exactly equal to the rate of crystallization. At this point, the total amount of dissolved solute remains constant, even though individual particles are constantly moving between the solid and dissolved forms.
In a solution at equilibrium, the concentrations of the dissolved ions or molecules appear fixed. However, the process of dissolving and reforming the solid is ongoing at the molecular level.
States of Solution and Saturation
The amount of solute dissolved defines the solution’s state relative to its equilibrium point, based on the maximum capacity of the solvent at a specific temperature. An unsaturated solution contains less than the maximum possible amount of dissolved solute. In this state, the solution is not at equilibrium, and more solute can be added and dissolved.
A saturated solution represents the state of solution equilibrium. This solution contains the maximum amount of dissolved solute the solvent can hold under those conditions. If additional solute is introduced, it will not dissolve and will instead settle at the bottom of the container.
The third state is a supersaturated solution, which is inherently unstable. This solution holds more dissolved solute than is theoretically possible for the solvent at that temperature. Supersaturated solutions are created by dissolving excess solute at a high temperature and then carefully cooling the solution. Adding a small seed crystal to this unstable solution triggers the rapid crystallization of the excess solute until the solution returns to the stable saturated state.
Quantifying Equilibrium Using the Solubility Product
While a saturated solution describes equilibrium qualitatively, the Solubility Product Constant, or \(K_{sp}\), provides a quantitative measure for the solubility of ionic compounds. \(K_{sp}\) is a specific type of equilibrium constant for the reaction where a solid ionic substance dissolves into its component ions in water. This value is calculated from the concentration of the dissolved ions in a saturated solution.
The \(K_{sp}\) value reflects the extent to which a slightly soluble salt will dissolve before the solution becomes saturated. A high \(K_{sp}\) indicates a highly soluble substance that dissolves a large amount of ions before reaching equilibrium. Conversely, a low \(K_{sp}\) indicates a sparingly soluble compound that dissolves only a minimal amount. The \(K_{sp}\) provides a precise standard for comparing the solubility of different ionic substances under the same conditions.
External Factors That Shift Equilibrium
Solution equilibrium is sensitive to environmental changes, and the system responds to a disturbance by shifting to establish a new balance. This behavior is governed by the principle that a system at equilibrium will shift to counteract any external change imposed on it. Two primary factors that influence solution equilibrium are temperature and the introduction of a common ion.
Changing the temperature affects the energy balance of the dissolution process. If dissolving the solute releases heat (exothermic), increasing the temperature shifts the equilibrium toward the undissolved solid, decreasing solubility. If the dissolving process absorbs heat (endothermic), increasing the temperature shifts the equilibrium toward the dissolved state, which increases solubility.
The Common Ion Effect describes how the solubility of a sparingly soluble salt decreases when a soluble compound containing one of the same ions is added. The addition of this “common ion” increases the concentration of one of the products of the dissolution reaction. To relieve this stress, the equilibrium shifts in the reverse direction, favoring the formation of the solid and reducing the amount of dissolved solute.