An ionic compound is a chemical substance formed when a metal and a nonmetal combine, resulting in a transfer of electrons that creates positively charged ions (cations) and negatively charged ions (anions). These oppositely charged particles are held together in a rigid, ordered structure called a crystal lattice by strong electrostatic attraction. Solubility is defined as the maximum amount of a substance, the solute, that can dissolve in a specific amount of a liquid, the solvent. Not all ionic compounds dissolve in water; their solubility depends entirely on the compound’s unique chemical identity and its specific interactions with the water molecules.
The Process of Dissolving
The dissolving of an ionic compound in water is a process of overcoming the strong electrostatic forces within the solid crystal structure. Water molecules are effective at this because they are polar, meaning the oxygen atom carries a partial negative charge, while the hydrogen atoms carry a partial positive charge. This separation of charge allows water to interact with both the positive and negative ions of the compound. The water molecules crowd around the ionic solid, initiating dissociation where the ions are pulled away from the crystal lattice.
The partially negative oxygen end of the water molecules surrounds the positive cation, while the partially positive hydrogen ends surround the negative anion. This process of the solvent molecules surrounding the dissociated ions is known as hydration or solvation. The strong attraction between the polar water molecules and the charged ions shields the ions from each other, preventing them from snapping back together. Once separated, the ions become dispersed uniformly throughout the water, forming a stable, homogenous solution.
For a solid to dissolve, the energy released during the hydration of the ions must overcome the lattice energy, which is the energy required to break the ionic bonds holding the solid together. If the attractive forces between the ions and water molecules are greater than the attractive forces holding the ions to each other, the compound will dissolve. Water’s high dielectric constant also reduces the electrostatic force between the separated ions in the solution, facilitating their free movement.
Predicting Solubility
Since not all ionic compounds dissolve in water, chemists use a set of solubility rules to predict the outcome of mixing specific ionic solids with water. These rules are based on patterns observed across many experiments and provide a way to determine if a compound is soluble or insoluble. One reliable indicator of solubility is the presence of certain ions, such as those from Group 1 of the periodic table, including sodium (\(\text{Na}^{+}\)) and potassium (\(\text{K}^{+}\)), or the polyatomic ammonium ion (\(\text{NH}_4^{+}\)). Compounds containing any of these ions are almost always soluble in water.
Another category of ions that confer solubility includes nitrate (\(\text{NO}_3^{-}\)), acetate (\(\text{CH}_3\text{COO}^{-}\)), and perchlorate (\(\text{ClO}_4^{-}\)). If an ionic compound contains any of these anions, it is likely to dissolve completely in an aqueous solution. The halides, which are chloride (\(\text{Cl}^{-}\)), bromide (\(\text{Br}^{-}\)), and iodide (\(\text{I}^{-}\)), are generally soluble, but they have a few exceptions. Halide compounds are insoluble when they are paired with silver (\(\text{Ag}^{+}\)), lead (\(\text{Pb}^{2+}\)), or mercury(I) (\(\text{Hg}_2^{2+}\)) ions.
Sulfates (\(\text{SO}_4^{2-}\)) are usually soluble in water, but this rule comes with a list of exceptions. Sulfates of barium (\(\text{Ba}^{2+}\)), lead (\(\text{Pb}^{2+}\)), and strontium (\(\text{Sr}^{2+}\)) are insoluble. Calcium sulfate (\(\text{CaSO}_4\)) and silver sulfate (\(\text{Ag}_2\text{SO}_4\)) are considered sparingly or slightly soluble, meaning they dissolve only to a small extent.
Conversely, many other types of ionic compounds are considered insoluble, such as carbonates (\(\text{CO}_3^{2-}\)), phosphates (\(\text{PO}_4^{3-}\)), and hydroxides (\(\text{OH}^{-}\)). The exceptions to these general insolubility rules are compounds formed with the soluble cations: Group 1 metals and the ammonium ion. For instance, while most hydroxides are insoluble, sodium hydroxide (\(\text{NaOH}\)) and potassium hydroxide (\(\text{KOH}\)) are highly soluble. These generalized rules allow for the prediction of precipitation reactions, where two soluble compounds mix and form a new, insoluble solid.
Factors That Influence Solubility
The ultimate determination of whether an ionic compound dissolves is an energetic calculation involving two thermodynamic factors: lattice energy and hydration energy. Lattice energy represents the energy input required to break the electrostatic attractions holding the ions in their fixed positions in the solid crystal. Hydration energy is the energy released when the individual ions are surrounded and stabilized by the polar water molecules. A compound will be soluble if the hydration energy is greater than or equal to the lattice energy, making the overall dissolution process energetically favorable.
External conditions also modulate the solubility of ionic solids, with temperature being a factor. For most solid ionic compounds, solubility increases as the temperature of the solvent rises because the added thermal energy provides the kinetic energy needed to break the lattice structure. However, there are exceptions, such as calcium sulfate (\(\text{CaSO}_4\)), whose solubility decreases as temperature increases.
The Common Ion Effect influences the maximum amount of solute that can dissolve in a solution. If a solution already contains one of the ions present in the solid compound to be dissolved, the solubility of that solid will be reduced. This phenomenon is explained by Le Chatelier’s principle, where the presence of the common ion shifts the dissolution equilibrium back toward the solid, undissolved form, thereby decreasing the overall solubility.