When examining the forces that hold matter together, we typically categorize them into strong chemical bonds, such as covalent and ionic bonds, and weaker intermolecular attractions. Covalent bonds, which involve the sharing of electron pairs, are immensely strong, requiring significant energy to break. Positioned between these robust chemical linkages and the feeblest attractions is the hydrogen bond, a unique type of interaction. Although an individual hydrogen bond is considerably weaker than a true chemical bond, its power lies in its ability to act collectively and repeatedly. This intermediate strength allows hydrogen bonds to serve as the foundation for the dynamic structures and processes that define much of chemistry and biology.
Defining the Hydrogen Bond
A hydrogen bond is not the result of electron sharing or transfer, but rather a special case of a dipole-dipole attraction between molecules. This attraction begins when a hydrogen atom is covalently bonded to a highly electronegative atom, typically oxygen (O), nitrogen (N), or fluorine (F). Because these electronegative atoms pull the shared electrons closer, the hydrogen atom acquires a partial positive charge, effectively becoming a hydrogen bond donor.
This electron-deficient hydrogen atom is then attracted to a nearby electronegative atom, which must possess an available lone pair of electrons to act as the hydrogen bond acceptor. The resulting interaction is a weak electrostatic bridge between the two electronegative centers, often represented by a dotted line to distinguish it from a covalent bond. This interaction can occur either between two separate molecules (intermolecular) or between two different parts of the same large molecule (intramolecular). Hydrogen bonds are stronger than ordinary dipole-dipole forces but far less energetic than the bonds holding atoms together within a molecule.
Quantifying Hydrogen Bond Strength
Hydrogen bonds are quantified by their dissociation energy, which falls into a moderate range compared to other chemical forces. A typical hydrogen bond, such as the one between two water molecules, has an energy of approximately 21 kilojoules per mole (kJ/mol). The overall range for hydrogen bonds is broad, extending from very weak interactions near 1 kJ/mol up to strong bonds around 40 kJ/mol.
For context, a strong covalent bond, like the carbon-hydrogen (C-H) bond, requires hundreds of kJ/mol to break, making a hydrogen bond roughly 5% as strong as a typical covalent link. Conversely, hydrogen bonds are substantially stronger than the weakest intermolecular forces, known as Van der Waals forces. This moderate energy level makes hydrogen bonds functionally significant: they are strong enough to maintain temporary structures yet weak enough to be rapidly broken and reformed under biological conditions. The energy required to break them is often comparable to the thermal energy available at room temperature, enabling dynamic processes.
Factors Modifying Bond Strength
The strength of a hydrogen bond is not a fixed number but is highly dependent on the local physical and chemical environment, explaining the wide 1 to 40 kJ/mol energy range.
Physical Determinants
One determinant is the distance between the donor and acceptor atoms. A shorter distance between the two electronegative atoms results in a stronger attractive force and higher bond energy.
The geometry of the interaction is also important. Hydrogen bonds are strongest when the donor atom, the hydrogen atom, and the acceptor atom are aligned in a straight, linear arrangement. Deviation from this 180-degree angle weakens the attraction by reducing the optimal electrostatic overlap.
Environmental Factors
The surrounding medium, or local environment, plays a decisive role, particularly in systems like water. The presence of neighboring polar molecules can enhance or diminish the bond’s strength through cooperative effects, where the formation of one hydrogen bond makes the next one easier to form.
Functional Significance of Moderate Strength
The moderate and dynamic nature of hydrogen bonds enables the complex structures and rapid processes necessary for life. Their strength is balanced to stabilize large biological molecules without making them rigid or permanent.
This is demonstrated in the structure of deoxyribonucleic acid (DNA), where hydrogen bonds form the rungs of the double helix, linking the base pairs. These bonds are strong enough to maintain the helical structure under normal conditions, but weak enough to be easily “unzipped” by enzymes during DNA replication or transcription.
In proteins, hydrogen bonds form repeatedly along the polypeptide backbone, dictating stable secondary structures, such as alpha-helices and beta-sheets. The ease with which these bonds break and reform allows proteins to rapidly fold into their specific three-dimensional shapes and undergo conformational changes necessary for function. The collective power of hydrogen bonds is also responsible for the unusual properties of water, giving it a high boiling point and surface tension because constant energy is required to break the extensive network of attractions between water molecules.