Solubility is the maximum amount of a substance (solute) that can dissolve in a given amount of another substance (solvent) at a specific temperature and pressure to achieve equilibrium. This measurement determines the concentration of a saturated solution, where the rates of dissolving and re-forming the solid are equal. Understanding and predicting solubility is a foundational concept in chemistry with broad applications, such as guiding drug formulation in pharmacology or determining pollutant behavior in environmental science.
The Foundation: Polarity and “Like Dissolves Like”
The primary principle for predicting the solubility of molecular substances is the concept of “like dissolves like,” which is based on molecular polarity. Polarity arises from the unequal sharing of electrons between atoms, creating a dipole moment with partial positive and partial negative charges within the molecule. Water is a distinctly polar solvent due to its highly polar bonds.
When predicting solubility, chemists consider the strength and type of intermolecular forces (IMFs) present in the solute and solvent. Polar substances use strong IMFs like hydrogen bonding, while nonpolar substances rely on weaker London dispersion forces. For a substance to dissolve, the forces of attraction between the solute and solvent molecules must be strong enough to overcome the existing attractions within both the pure solute and the pure solvent.
Following the “like dissolves like” rule, polar solvents readily dissolve polar solutes, and nonpolar solvents dissolve nonpolar solutes. For example, highly polar table sugar dissolves easily in polar water, whereas nonpolar oils separate stubbornly from water. The molecular structure, including the presence of highly electronegative atoms, determines a molecule’s polarity and its preferred solvent.
Applying Qualitative Rules to Ionic Compounds
While ionic compounds are composed of charged ions, their solubility in a polar solvent like water is determined by a set of empirical, qualitative rules. The dissolution of an ionic salt involves a complex balance between the energy required to break the crystal lattice and the energy released during hydration. Because predicting this energy balance is difficult, chemists rely on established guidelines for quick assessment.
A practical approach is to remember the ions that are almost always soluble in water, as these form the basis of the rules.
- All compounds containing alkali metal ions (e.g., \(\text{Na}^{+}\) and \(\text{K}^{+}\)) or the ammonium ion (\(\text{NH}_4^{+}\)) are soluble without exception.
- Salts containing the nitrate (\(\text{NO}_3^{-}\)), acetate (\(\text{C}_2\text{H}_3\text{O}_2^{-}\)), or perchlorate (\(\text{ClO}_4^{-}\)) ions are generally soluble.
The next step involves knowing the generally soluble ions that have predictable exceptions, such as the halides and the sulfate ion.
- Halides (chloride, bromide, and iodide) are typically soluble, but they form insoluble precipitates with lead (\(\text{Pb}^{2+}\)), silver (\(\text{Ag}^{+}\)), and mercury(I) (\(\text{Hg}_2^{2+}\)) ions.
- Sulfates (\(\text{SO}_4^{2-}\)) are also soluble, but they are insoluble when combined with barium (\(\text{Ba}^{2+}\)), strontium (\(\text{Sr}^{2+}\)), and lead (\(\text{Pb}^{2+}\)).
Conversely, most salts containing the following ions are considered insoluble:
- Carbonate (\(\text{CO}_3^{2-}\))
- Phosphate (\(\text{PO}_4^{3-}\))
- Sulfide (\(\text{S}^{2-}\))
- Hydroxide (\(\text{OH}^{-}\))
The only exceptions to these insoluble rules are the compounds formed with Group 1 elements and ammonium. For example, calcium carbonate is insoluble, but sodium carbonate is highly soluble due to the presence of the sodium ion.
How Temperature and Pressure Alter Solubility
Solubility is significantly influenced by external conditions, primarily temperature and, for gases, pressure. The effect of temperature varies depending on whether the solute is a solid or a gas. For most solid solutes dissolved in a liquid, an increase in temperature generally leads to an increase in solubility.
This occurs because heating provides energy to help break the solute’s structure and increase the movement of solvent molecules. However, this trend is not universal; some solids, such as cerium(III) sulfate (\(\text{Ce}_2(\text{SO}_4)_3\)), become less soluble as temperature rises. This happens because the dissolution process for these exceptions is exothermic, shifting the equilibrium to favor the undissolved solid when heat is added.
The solubility of gases in liquids consistently decreases as the temperature increases. When a liquid is heated, dissolved gas molecules gain kinetic energy and overcome the attractive forces holding them in the solution, causing them to escape. This phenomenon explains why a warm soda goes flat faster than a cold one and reduces dissolved oxygen for aquatic life in warmer waters.
Pressure has a profound effect on the solubility of gases but a negligible effect on the solubility of liquids and solids. The relationship between gas pressure and solubility is precisely described by Henry’s Law. This law states that the concentration of a dissolved gas is directly proportional to the partial pressure of that gas above the liquid, which is the mechanism used to carbonate beverages.