A hydrogen bond is a special type of attractive force that occurs between molecules, representing the strongest variety of intermolecular attractions. This force is not a primary chemical bond, like a covalent or ionic bond, but rather a powerful interaction based on electrical polarity. It involves the electrostatic attraction between a partially positive hydrogen atom in one molecule and a partially negative atom in a neighboring molecule. The formation of these bonds governs the behavior of simple substances and underlies the complex structures of living systems.
The Necessary Chemical Conditions for Hydrogen Bonding
Hydrogen bonding occurs only under specific chemical circumstances involving highly polarized bonds. The interaction requires a hydrogen atom to be covalently linked to one of three small, highly electronegative elements: nitrogen (N), oxygen (O), or fluorine (F). These atoms strongly pull the shared electrons closer to their own nuclei, leaving the hydrogen atom with a significant partial positive charge, making it a hydrogen bond “donor.” The donor hydrogen atom then forms an attraction to a second, nearby electronegative atom that possesses a lone pair of electrons, called the hydrogen bond “acceptor.” This precise arrangement is why molecules like methane (\(\text{CH}_4\)) do not form hydrogen bonds, as carbon is not sufficiently electronegative to create the necessary degree of polarity with hydrogen. The stability of this interaction is generally weak compared to a covalent bond, but the cumulative effect of multiple hydrogen bonds can be substantial.
How Hydrogen Bonds Shape Water and Physical States
Water (\({\text{H}_2\text{O}}\)) serves as the most familiar example of a substance shaped entirely by hydrogen bonds, as its oxygen atom can act as two acceptors and its two hydrogen atoms can act as two donors. The collective strength of these bonds gives water a much higher boiling point than other chemically similar molecules, such as hydrogen sulfide (\({\text{H}_2\text{S}}\)), which is a gas at room temperature.
Hydrogen bonds are also responsible for water’s cohesive and adhesive properties. Cohesion, the attraction of water molecules to each other, creates surface tension, allowing small insects to walk across the liquid surface. Adhesion is the attraction of water to other polar surfaces, which enables phenomena like capillary action, helping water move through narrow tubes and plant vessels against the force of gravity.
Furthermore, hydrogen bonds explain the unusual density anomaly where solid ice floats on liquid water. In the liquid state, water molecules are constantly breaking and reforming hydrogen bonds, allowing them to pack relatively closely. When water freezes, the molecules settle into a more stable, open, crystalline lattice structure where each molecule is perfectly hydrogen-bonded to four neighbors. This open arrangement spaces the molecules farther apart in ice than in liquid water, making the solid form less dense.
The Role in Biological Macromolecules
In biological systems, hydrogen bonds are important for establishing and maintaining the precise three-dimensional shapes of large molecules. These bonds are numerous enough to provide stability but weak enough to be easily broken and reformed, which is necessary for biological processes. The structure of deoxyribonucleic acid (DNA) is stabilized by hydrogen bonds that link the two complementary strands together to form the double helix.
Specifically, the nitrogenous bases pair across the helix using these attractions: adenine (A) always pairs with thymine (T) via two hydrogen bonds, while guanine (G) pairs with cytosine (C) using three hydrogen bonds. This selective pairing maintains the genetic code’s integrity while allowing the bonds to be temporarily broken by enzymes during processes like DNA replication and transcription.
Hydrogen bonds also dictate the secondary structure of proteins, which are chains of amino acids. Within a single polypeptide chain, hydrogen bonds form regularly between the oxygen atom of one amino acid’s carbonyl group and the hydrogen atom of another amino acid’s amide group. This repeated pattern of bonding gives rise to stable, local structures, most notably the coiled \(\alpha\)-helix and the pleated \(\beta\)-sheet. These secondary structures are folding patterns that are foundational to a protein’s overall shape, which in turn determines its specific biological function.