Water, often called the universal solvent, possesses a unique molecular structure that allows it to dissolve a vast number of chemical compounds. Solubility describes the maximum amount of a substance (the solute) that can completely disperse within a specific amount of another substance (the solvent). To determine whether a substance will dissolve in water, the fundamental approach involves analyzing the physical and chemical properties of the solute at the molecular level. This predictive analysis relies on understanding the forces of attraction between the solute particles and the surrounding water molecules.
The Guiding Principle of Solubility
The most reliable principle for predicting dissolution is the maxim that “like dissolves like,” which relates directly to molecular polarity. Water is a highly polar molecule because its oxygen atom pulls electrons away from the two hydrogen atoms, creating partial negative and positive charges. This uneven distribution allows water molecules to attract other charged or partially charged substances. Polar substances readily dissolve in water because the attractive forces between the solute and water are strong enough to overcome the internal forces holding the solute together. Nonpolar molecules, such as oils, have a uniform charge distribution, and when mixed with water, the water molecules are more attracted to each other than to the nonpolar solute, preventing dissolution.
Predicting Solubility for Ionic Compounds
Ionic compounds consist of positive and negative ions held together by strong electrostatic forces in a crystal lattice structure. These compounds are highly polar and generally dissolve in water because water molecules can surround and separate the individual ions, a process called dissociation. The partially negative oxygen end of the water molecule attracts the positive ion, while the partially positive hydrogen ends attract the negative ion, pulling the crystal apart. Despite this general rule, specific exceptions are governed by Solubility Rules.
Solubility Rules for Ionic Compounds
Most compounds containing nitrate (\(text{NO}_3^-\)) or alkali metal ions (Group 1 elements like sodium and potassium) are consistently soluble.
Conversely, most compounds formed with carbonate (\(text{CO}_3^{2-}\)), phosphate (\(text{PO}_4^{3-}\)), or sulfide (\(text{S}^{2-}\)) ions are typically insoluble in water.
A few exceptions occur among the generally soluble halides (compounds containing chlorine (\(text{Cl}^-\)), bromine (\(text{Br}^-\)), or iodine (\(text{I}^-\))). While most halides dissolve easily, combinations with heavy metal ions like silver (\(text{Ag}^+\)), lead (\(text{Pb}^{2+}\)), or mercury (\(text{Hg}_2^{2+}\)) result in insoluble salts. These combinations form strong ionic bonds that require more energy to break apart than water molecules can provide, leading to precipitation.
Predicting Solubility for Molecular Compounds
Molecular compounds, formed by covalent bonds between non-metal atoms, rely on different attractive forces to dissolve in water. Substances like ethanol or table sugar are highly soluble because they possess groups of atoms containing hydrogen bonded to electronegative elements like oxygen or nitrogen. This arrangement facilitates a powerful intermolecular attraction known as hydrogen bonding. Water molecules form these strong hydrogen bonds with the solute, integrating them into the liquid structure. This extensive network of hydrogen bonds allows a large amount of sugar, for example, to dissolve before saturation.
Molecules that lack the ability to form these strong hydrogen bonds, such as methane or benzene, are nonpolar and considered hydrophobic. Fats, oils, and waxes are large, entirely nonpolar molecular compounds that are immiscible. These substances are held together by weak London dispersion forces, which cannot compete with the cohesive forces between water molecules. When mixed with water, they separate into distinct layers because water molecules preferentially interact with each other.
Changing Solubility
Although chemical structure is the primary determinant of solubility, external factors can influence the maximum amount of a substance that can be dissolved. For most solid solutes, increasing the temperature of the solvent increases the solubility because the added thermal energy helps break the bonds holding the solute together. This allows more solute particles to disperse throughout the solution, as seen when stirring sugar into hot tea versus cold water.
The solubility of gases, however, generally behaves in the opposite manner; an increase in temperature typically decreases the gas’s solubility in a liquid. This is due to the increased kinetic energy of the gas molecules, which allows them to escape the attractive forces of the solvent and return to the atmosphere.
Pressure is another factor that significantly affects the solubility of gases but has little effect on solids or liquids. Higher pressures force more gas molecules into the liquid, which is the principle behind carbonated beverages where carbon dioxide is dissolved under pressure.