Covalent bonds are a fundamental chemical linkage where two atoms share one or more pairs of electrons. It is a common misunderstanding that the bond itself dissolves; rather, solubility describes the ability of the entire molecule (the solute) to uniformly disperse throughout a liquid medium (the solvent). The determining factor for a covalent molecule’s solubility in water is the distribution of electrical charge, which depends directly on the type of covalent bonds present.
The Nature of Covalent Bonds and Polarity
Covalent bonds are formed when two non-metal atoms share valence electrons to achieve a stable outer shell configuration. This sharing is not always equal, introducing the concept of polarity. The degree of sharing is governed by electronegativity, an atom’s tendency to attract a shared pair of electrons toward itself.
When two atoms of the same element bond, such as in an oxygen molecule (\(\text{O}_2\)), the electrons are shared equally, resulting in a nonpolar covalent bond. When atoms of different elements bond, one atom often has greater electronegativity, pulling the electron pair closer to its nucleus. This unequal sharing creates a separation of charge, giving the more electronegative atom a partial negative charge (\(\delta-\)) and the less electronegative atom a partial positive charge (\(\delta+\)).
This charge separation establishes a dipole moment, making it a polar covalent bond. A difference in electronegativity greater than approximately 0.4 on the Pauling scale is sufficient to classify a bond as polar. The presence of these internal dipoles dictates how a covalent molecule will interact with water.
Water’s Unique Structure and Solvent Properties
Water (\(\text{H}_2\text{O}\)) functions as an effective solvent because of its unique molecular structure. The molecule has a bent geometry, with the oxygen atom positioned at the vertex. Oxygen is significantly more electronegative than hydrogen, causing it to draw the shared electrons closer, which gives the oxygen atom a partial negative charge and the hydrogen atoms partial positive charges.
This distinct separation of charge makes water a strong molecular dipole. The general chemical principle governing dissolution is summarized as “like dissolves like.” Since water is a highly polar solvent, it readily interacts with other substances that also possess charge or polarity, making it an excellent solvent for polar molecules and charged ionic compounds.
When Covalent Compounds Dissolve in Water
Covalent compounds dissolve in water when they are polar enough to engage in favorable intermolecular forces with water molecules. A prime example is table sugar (sucrose), a large molecule rich with numerous polar hydroxyl (\(\text{-OH}\)) groups. The partial negative charge on water’s oxygen atom is strongly attracted to the partial positive charge on the hydrogen atoms of the sugar’s hydroxyl groups.
The strongest of these attractive interactions is the formation of a hydrogen bond. This occurs when a hydrogen atom bonded to an electronegative atom like oxygen is attracted to a lone pair of electrons on a neighboring oxygen atom. When a polar molecule like sugar is introduced, water molecules break their existing hydrogen bonds to form new, stable hydrogen bonds with the solute. This energetically favorable process pulls the solute molecule away from its neighbors and surrounds it entirely.
The water molecules form a dynamic cage around the dissolved solute molecule, known as a hydration shell. This shell isolates the solute molecule, allowing it to move freely throughout the solution. For dissolution to occur, the energy released from forming these new solute-solvent interactions must compensate for the energy required to break the solute-solute and solvent-solvent bonds.
When Covalent Compounds Do Not Dissolve in Water
Covalent molecules that lack significant polarity, such as oils and hydrocarbons like methane (\(\text{CH}_4\)), are insoluble in water. These compounds are composed of nonpolar bonds, like carbon-hydrogen bonds, where electrons are shared almost equally, resulting in no substantial partial charges. Consequently, these molecules cannot form the strong dipole-dipole attractions or hydrogen bonds necessary to interact favorably with water.
When a nonpolar molecule is introduced to water, it must push apart the strongly hydrogen-bonded water molecules, requiring a substantial input of energy. The weak, temporary London dispersion forces that form between the nonpolar molecule and water are not strong enough to repay this energetic cost. Water molecules energetically prefer to remain bonded to each other rather than interact with the nonpolar substance.
This preference forces the nonpolar molecules to cluster together, minimizing the surface area of contact with the water. The resulting phase separation, often seen when oil floats on water, is not caused by repulsion but by the water molecules maximizing their strong self-attractions. The system achieves the lowest energy state when water molecules form the greatest number of hydrogen bonds with one another, excluding the nonpolar substance.