Water is the most abundant molecule in living organisms, often making up 60–70 percent of cellular content. Water’s chemical properties are entirely dictated by its molecular structure (\(\text{H}_2\text{O}\)). This unique arrangement gives rise to the substance’s distinctive features, allowing water to serve as the universal solvent and the medium for nearly all biological processes. Understanding these molecular traits is key to appreciating water’s profound role in sustaining life on Earth.
The Foundation: Polarity and Hydrogen Bonding
The unique chemical properties of water begin with the shape and electron distribution of the \(\text{H}_2\text{O}\) molecule. Water possesses a bent or V-shape, where the single oxygen atom is covalently bonded to two hydrogen atoms. This geometry, combined with the difference in electronegativity between oxygen and hydrogen, creates an uneven sharing of electrons within the molecule.
Oxygen is a strongly electronegative atom, meaning it pulls the shared electrons closer to its nucleus than the hydrogen atoms do. This unequal pull results in the oxygen side of the molecule acquiring a partial negative charge (\(\delta^{-}\)), while the hydrogen sides acquire partial positive charges (\(\delta^{+}\)). Although the water molecule has no net electrical charge, this separation of charge makes it a highly polar molecule.
This inherent polarity allows water molecules to form weak attractions with each other, known as hydrogen bonds. A hydrogen bond is a weak electrostatic force that forms between the partially positive hydrogen of one water molecule and the partially negative oxygen of a neighboring water molecule. While a single hydrogen bond is much weaker than the covalent bonds holding the \(\text{H}_2\text{O}\) molecule together, the sheer number of these bonds acting in concert gives water its extraordinary properties. These intermolecular attractions are the foundational force that influences water’s behavior in its liquid state.
Cohesion, Adhesion, and Surface Tension
The collective action of hydrogen bonds manifests in water’s ability to stick to itself and to other materials. Cohesion is the attraction between water molecules, driven by hydrogen bonds. This cohesive force is responsible for phenomena like the formation of spherical water droplets, as the molecules pull toward each other to minimize surface area.
Adhesion is the attraction between water molecules and molecules of a different substance. Water molecules are strongly attracted to other polar or charged surfaces, such as the cellulose walls in a plant stem. The combined forces of cohesion and adhesion enable capillary action, which is the ability of water to climb upward against the force of gravity in narrow spaces. This upward movement transports water and dissolved nutrients from the roots to the leaves of plants.
A consequence of strong cohesive forces is surface tension, the tendency of a liquid’s surface to resist external force. Water molecules at the surface are pulled inward and sideways by their neighbors, creating a taut interface. This force allows small insects, such as water striders, to walk across the surface without sinking.
Thermal Stability and Density Anomaly
Water exhibits an unusually high specific heat capacity, meaning it requires a large amount of heat energy to raise its temperature. When heat is absorbed by liquid water, the energy must first be used to break the numerous hydrogen bonds connecting the molecules. Only after these bonds are disrupted does the kinetic energy of the molecules increase, resulting in a temperature rise.
This characteristic allows water to absorb and store large amounts of heat with minimal temperature fluctuation. This thermal stability moderates the climate of coastal regions and stabilizes body temperatures in organisms. Likewise, water must lose a considerable amount of energy for its temperature to drop, preventing rapid temperature shifts that would be detrimental to biological systems.
Water also displays a peculiar density anomaly. Most substances become progressively denser as they cool and solidify, but liquid water reaches its maximum density at approximately \(4^\circ\text{C}\). Below this temperature, water begins to expand, and its solid form, ice, is less dense than the liquid water from which it forms.
As water freezes, the hydrogen bonds lock the molecules into a rigid, open, crystalline lattice structure. This organized arrangement spaces the molecules further apart than they are in the constantly shifting liquid state. The reduced density of ice causes it to float on liquid water, which insulates the water below and prevents large bodies of water from freezing solid, allowing aquatic life to survive winter.
Water as the Universal Solvent
Water’s polarity makes it an exceptional solvent, earning it the title of the “universal solvent” for its ability to dissolve more substances than any other liquid. Substances that dissolve readily in water are called hydrophilic (water-loving) and are typically polar molecules or ions. The charged regions of the water molecule are strongly attracted to the charges on the solute particles, allowing the water to pull them apart.
For ionic compounds, the partially negative oxygen end of the water molecule surrounds the positive ions, while the partially positive hydrogen ends surround the negative ions. This process is known as solvation or hydration. The water molecules form a structured layer around the solute particles called a hydration shell. This shell shields the ions from each other, preventing them from re-associating and keeping them dispersed in the solution.
In contrast, non-polar substances like oils and fats are hydrophobic (water-fearing) and do not dissolve in water. Because they lack charged or polar regions, water molecules cannot form hydrogen bonds with them. Instead, water pushes them together, forcing the non-polar molecules to aggregate and separate from the water. Water’s ability to dissolve and transport a vast array of polar and ionic substances facilitates the countless chemical reactions necessary for life within the cell.