Water covers much of the Earth’s surface and makes up a large percentage of living organisms. Its ability to sustain life and act as an almost universal solvent is directly linked to its unusual physical properties. These unique characteristics are a direct consequence of a specific type of attractive force that acts between its individual molecules. This force is known as hydrogen bonding, and it dictates how water behaves in its liquid, solid, and gaseous states.
The Polar Nature of a Water Molecule
A single water molecule consists of one oxygen atom bonded to two hydrogen atoms. While the molecule is electrically neutral overall, the sharing of electrons within its bonds is unequal. This imbalance is caused by electronegativity, which is an atom’s ability to attract shared electrons in a chemical bond.
Oxygen is significantly more electronegative than hydrogen, meaning it pulls the shared electrons closer to its nucleus. This unequal sharing causes the oxygen atom to acquire a slight negative charge (\(\delta^-\)), while the hydrogen atoms become slightly electron-deficient, resulting in a partial positive charge (\(\delta^+\)) on each one.
This separation of charge creates a molecular dipole, making the water molecule a tiny magnet with a positive end and a negative end. The bent geometry ensures that these partial charges do not cancel each other out. This polarity is the necessary precondition for the formation of the hydrogen bond.
Defining the Hydrogen Bond
A hydrogen bond is a special type of attractive force that occurs between separate molecules. This attraction arises when the partially positive hydrogen atom of one water molecule is drawn toward the partial negative charge on the oxygen atom of a neighboring molecule. The attraction is specifically between the electron-deficient hydrogen atom and a lone pair of electrons on the adjacent oxygen atom.
This electrostatic attraction is considerably stronger than other general intermolecular forces, but it is substantially weaker than the covalent bonds holding atoms together within a single water molecule. A hydrogen bond is estimated to have only about one-twentieth the strength of a water molecule’s internal covalent bond. In liquid water, these bonds are temporary, constantly breaking and reforming as the molecules move past one another.
The existence of a hydrogen bond requires a hydrogen atom covalently bonded to a highly electronegative atom like oxygen, nitrogen, or fluorine, which creates the strong partial positive charge on the hydrogen. Water is an ideal example because each molecule has the necessary partial positive hydrogen atoms and lone pairs of electrons to form an extensive network. This network links many water molecules together, giving liquid water more structure than most other liquids.
The Unique Characteristics Hydrogen Bonding Creates
The network of hydrogen bonds is directly responsible for many of water’s distinctive physical properties that are important for life on Earth. One effect is cohesion, which is the attraction of water molecules to each other, resulting from the continuous making and breaking of these bonds. This cohesive nature gives water a very high surface tension, allowing small insects to walk on the water’s surface.
Another significant property is water’s remarkably high specific heat capacity, meaning it requires a large amount of energy to change its temperature. Much of the added heat energy must first be used to break the extensive web of hydrogen bonds before the molecules can begin to move faster. This helps regulate the temperatures of large bodies of water and the organisms living within them, preventing rapid temperature shifts.
A highly unusual characteristic is that solid water, or ice, is less dense than liquid water, allowing it to float. As water freezes, the hydrogen bonds lock the molecules into a highly ordered, open, crystalline lattice structure. This fixed arrangement holds the molecules further apart than they are in the constantly shifting liquid state. This makes ice about eight percent less dense than liquid water at its maximum density. The floating layer of ice then insulates the liquid water below, preventing entire bodies of water from freezing solid.