Hydroxides are ionic compounds containing the hydroxide anion (\(\text{OH}^-\)) bonded to a positively charged ion (cation). Whether hydroxides dissolve in water is not a simple yes or no, but a spectrum determined by the specific metal ion involved. Solubility is the measure of a substance’s ability to dissolve in a solvent, typically water, to form a uniform solution. The extent to which a hydroxide dissolves depends entirely on the chemical identity of the cation.
The General Rules for Hydroxide Solubility
The general rule for hydroxides is that most are considered insoluble in water, but there are important and predictable exceptions involving Group 1 and Group 2 metals. All hydroxides of Group 1 alkali metals, such as sodium hydroxide (\(\text{NaOH}\)) and potassium hydroxide (\(\text{KOH}\)), are highly soluble, dissolving readily and completely in water.
The heavier elements from Group 2, the alkaline earth metals, are classified as sparingly or moderately soluble. This includes calcium hydroxide (\(\text{Ca}(\text{OH})_2\)), strontium hydroxide (\(\text{Sr}(\text{OH})_2\)), and barium hydroxide (\(\text{Ba}(\text{OH})_2\)). While they do not dissolve to the same extent as Group 1 hydroxides, a measurable amount will enter the solution. Calcium hydroxide is commonly known as limewater and forms a mildly basic solution due to its limited dissolution.
The general rule of insolubility holds true for the vast majority of other metal hydroxides, particularly those involving transition metals like iron, copper, and zinc. Iron(III) hydroxide (\(\text{Fe}(\text{OH})_3\)) and copper(II) hydroxide (\(\text{Cu}(\text{OH})_2\)) are classic examples of insoluble hydroxides. They form a solid precipitate when their ions are mixed in water, allowing chemists to predict solid formation accurately.
Chemical Forces Driving Solubility and Insolubility
The solubility of any ionic compound results from an energetic competition between two opposing forces: lattice energy and hydration energy. Lattice energy is the strong electrostatic attraction holding the cation and hydroxide anions together in the solid crystal structure. High lattice energy favors insolubility because it requires significant energy to break these bonds and separate the ions.
Hydration energy is the energy released when water molecules surround and stabilize the separated ions, pulling them out of the crystal lattice. Solubility occurs when the hydration energy is equal to or greater than the lattice energy, making dissolution energetically favorable. The size and charge of the cation significantly influence this balance.
The small size of the hydroxide ion (\(\text{OH}^-\)) makes the lattice energy sensitive to the size of the cation. Small, highly charged cations, such as iron(\(\text{Fe}^{3+}\)) or magnesium(\(\text{Mg}^{2+}\)), create a compact crystal lattice with extremely high lattice energy. The lattice energy often dominates the high hydration energy for these ions, resulting in insoluble hydroxides.
In contrast, the large, singly charged cations of Group 1 metals, like potassium (\(\text{K}^+\)), form less compact crystal structures with much lower lattice energy. The hydration energy easily overcomes this weaker lattice energy, leading to high solubility. Moving down Group 2, the increasing size of the cation causes the lattice energy to decrease faster than the hydration energy, explaining the trend of increasing solubility.
Quantifying Solubility: Understanding \(K_{sp}\) Values
Solubility is a continuous spectrum, quantified using the Solubility Product Constant (\(K_{sp}\)). The \(K_{sp}\) is an equilibrium constant applied to sparingly soluble salts, reflecting the balance between the solid compound and its dissolved ions in solution.
The \(K_{sp}\) is calculated by multiplying the molar concentrations of the dissolved ions, raised to the power of their coefficients from the balanced chemical equation. This value serves as a direct indicator of a compound’s solubility under standard conditions.
A very small \(K_{sp}\) value signifies extremely low solubility, such as aluminum hydroxide (\(\text{Al}(\text{OH})_3\)) which has a \(K_{sp}\) around \(3 \times 10^{-34}\). Conversely, a larger \(K_{sp}\) value indicates a higher concentration of dissolved ions and a more soluble compound.
Comparing magnesium hydroxide (\(\text{Mg}(\text{OH})_2\)), with a \(K_{sp}\) of \(7.2 \times 10^{-15}\), to calcium hydroxide (\(\text{Ca}(\text{OH})_2\)), with a \(K_{sp}\) of \(5.0 \times 10^{-6}\), illustrates this relationship. The nearly one-million-fold difference confirms that calcium hydroxide is substantially more soluble than magnesium hydroxide.