Hydrogen bonding is a type of attractive force between molecules. This interaction is a weaker attraction between existing molecules or within a single large molecule, unlike strong chemical bonds. Despite being individually weaker than covalent or ionic bonds, the collective effect of numerous hydrogen bonds is profound. This widespread intermolecular attraction plays a significant role in shaping the characteristics of many substances. These bonds are important to various natural phenomena and biological processes.
The Mechanics of Hydrogen Bonding
Hydrogen bonds form when a hydrogen atom, covalently linked to a highly electronegative atom, is attracted to another electronegative atom. Electronegativity is an atom’s ability to attract shared electrons. When hydrogen bonds to a very electronegative atom like oxygen, nitrogen, or fluorine, electrons are pulled closer to the electronegative atom. This unequal sharing creates a partial negative charge on the electronegative atom and a partial positive charge on the hydrogen.
The partially positive hydrogen atom then becomes attracted to a nearby electronegative atom that possesses a lone pair of electrons. This creates an electrostatic attraction. The atom covalently bonded to the hydrogen is termed the “hydrogen bond donor.”
The electronegative atom with the lone pair that attracts the hydrogen is known as the “hydrogen bond acceptor.” Common hydrogen bond acceptor atoms are often oxygen, nitrogen, or fluorine, due to their high electronegativity and available electron pairs. This interaction is a special case of dipole-dipole attraction. While weaker than covalent bonds, hydrogen bonds are stronger than other intermolecular forces like van der Waals forces, with strengths ranging from 4 to 50 kJ per mole.
Ubiquitous Presence
Hydrogen bonds are common throughout nature, appearing in both inorganic and organic compounds. Water molecules provide a primary example, where hydrogen atoms from one water molecule are attracted to the oxygen atoms of neighboring water molecules. This attraction allows water molecules to associate extensively, forming a dynamic network. Each water molecule can participate in up to four hydrogen bonds, acting as both a donor and an acceptor.
In biological systems, hydrogen bonds are important to the structure of deoxyribonucleic acid (DNA). These bonds connect the two long strands of the DNA double helix, forming between complementary nitrogenous bases. Adenine (A) always pairs with thymine (T) through two hydrogen bonds, while guanine (G) pairs with cytosine (C) using three hydrogen bonds. This specific pairing maintains the characteristic double helical shape of DNA.
Proteins also rely on hydrogen bonds for their three-dimensional structures. These bonds help stabilize the protein’s secondary structures, such as alpha helices and beta sheets. Within an alpha helix, hydrogen bonds form regularly between the backbone atoms of amino acids, specifically between the hydrogen of an amide group and the oxygen of a carbonyl group. Beta sheets are similarly stabilized by hydrogen bonds between adjacent polypeptide segments.
Influence on Properties and Functions
The presence of hydrogen bonds significantly influences the physical and chemical properties of many substances, particularly water. Water’s relatively high boiling point of 100 °C, compared to similarly sized molecules, is a direct result of the energy required to break these numerous intermolecular hydrogen bonds before water can transition into a gaseous state. These bonds also contribute to water’s high specific heat capacity, meaning it can absorb or release substantial amounts of heat with only minor temperature changes, which helps regulate temperatures in living organisms and environments.
Hydrogen bonding also accounts for water’s cohesive and adhesive properties. Cohesion, the attraction between water molecules, leads to phenomena like surface tension, enabling small insects to walk on water. Adhesion, the attraction between water and other surfaces, is evident in capillary action, which allows water to move upwards in plant stems. Furthermore, the unique arrangement of hydrogen bonds in ice causes it to be less dense than liquid water, allowing ice to float and insulate aquatic life in cold climates.
In DNA, the specific hydrogen bonding between base pairs is important for both its stability and its ability to replicate accurately. While individually weak, the cumulative strength of numerous hydrogen bonds holds the two DNA strands together, forming a stable double helix. This balance allows the bonds to be strong enough to maintain structural integrity under normal conditions yet weak enough to be broken by enzymes during DNA replication. This allows the strands to “unzip,” providing templates for new complementary strands, ensuring genetic information is faithfully transmitted.
Hydrogen bonds are important to the complex functions of proteins, which are directly related to their precise three-dimensional shapes. Beyond secondary structures, hydrogen bonds contribute to the overall tertiary and quaternary structures of proteins, dictating how different parts of a single polypeptide chain fold and how multiple chains assemble. These intricate folds create specific binding sites and active regions, allowing proteins to perform diverse functions, such as catalyzing chemical reactions, transporting molecules, or providing structural support. Disruptions to these hydrogen bond networks, often caused by changes in temperature or pH, can alter a protein’s shape, leading to a loss of its biological activity.