Among the solubility rules previously discussed is the statement: Carbonates, phosphates, borates, and arsenates - except those of the ammonium ion and the alkali metals-are insoluble. (a) Write the expression for the equilibrium constant for the reaction represented by the equation \(\mathrm{CaCO}_{3}(s) \rightleftharpoons \mathrm{Ca}^{2+}(a q)+\mathrm{CO}_{3}^{2-}(a q) .\) Is \(K_{c} > 1, < 1,\) or \(\approx 1 ?\) Explain your answer. (b) Write the expression for the equilibrium constant for the reaction represented by the equation \(3 \mathrm{Ba}^{2+}(a q)+2 \mathrm{PO}_{4}^{3-}(a q) \rightleftharpoons \mathrm{Ba}_{3}\left(\mathrm{PO}_{4}\right)_{2}(s) .\) Is \(K_{c} > 1, < 1,\) or \(\approx 1 ?\) Explain your answer.

Understanding the equilibrium constant expression is pivotal when studying chemical equilibria of reactions. This expression, often represented as Kc for concentrations or Kp for partial pressures, quantifies the ratio of the concentrations of products to reactants, each raised to the power of its stoichiometric coefficient at equilibrium.

For example, in the dissolution of a salt like calcium carbonate, the solid phase does not appear in the expression because its concentration is constant. Thus, only the aqueous ions are included. The calcium carbonate equation, results in the equilibrium constant expression: .

The value of Kc gives insight into the favorability of a reaction. If Kc is much greater than 1, the reaction heavily favors the products. Conversely, when Kc is much less than 1, the reaction favors the reactants. Thus, a low Kc value in the context of solubility indicates that very little will dissolve and, in the case of calcium carbonate, Kc is indeed much less than 1, reflecting its poor solubility in water.

The solubility product constant, Ksp, is a special type of equilibrium constant that applies to the dissolution of sparingly soluble salts. Ksp is the product of the molar concentrations of the ions in a saturated solution, raised to the power of their coefficients in the dissolution equation.

This differs from Kc in that it pertains specifically to ionic compounds dissociating into their constituent ions in a solution. For the dissolution of calcium carbonate (), Ksp is defined as , illustrating how Ksp uniquely applies to the equilibrium between the solid phase and its dissolved ions. A small Ksp, like that found for calcium carbonate, signifies poor solubility. It becomes an invaluable tool for predicting whether a precipitate will form when two solutions are mixed.

Dissolution and precipitation are reverse processes that describe the dissolving of a solid in a liquid and the formation of a solid from a liquid, respectively. In dissolution, a solid like calcium carbonate breaks down into its ions, while precipitation refers to the formation of a solid compound from these ions.

When these processes reach a state of dynamic equilibrium in a saturated solution, the rate of dissolution equals the rate of precipitation. At this point, the solution is saturated, and any additional substance added will precipitate out. This behavior is predictable via the Ksp value: if the product of the ion concentrations exceeds Ksp, precipitation occurs; if it is less than Ksp, dissolution can continue. For instance, barium phosphate's low Ksp reflects its likelihood to precipitate from a solution.

Insoluble salts, like many carbonates, phosphates, borates, and arsenates (except those of the alkali metals and ammonium), have very low solubility in water. This means they will not significantly dissociate into ions when mixed with water.

Their solubility is so low that they are often considered as not dissolving at all in practical scenarios. This is critical information when predicting the outcome of chemical reactions, particularly in qualitative analysis. Salts like or are examples where the insoluble nature is directly related to their extremely small Ksp values. Understanding the principles behind insoluble salts assists chemists in designing separation strategies and predicting reaction products in aqueous solutions.