Hydrogen bonding is a powerful, yet relatively weak, force that profoundly influences the world around us. This specific attraction between molecules is responsible for many everyday phenomena, from material structure to the processes that sustain life. It is considered an intermolecular force, meaning it occurs between separate molecules, unlike stronger covalent bonds which operate within a single molecule. Understanding which molecules can participate in this interaction is key to grasping its broad consequences in chemistry and biology.
Defining Hydrogen Bonding
Hydrogen bonding is a particularly strong form of dipole-dipole attraction between molecules. A dipole forms when electrons are shared unequally in a covalent bond, causing a partial separation of charge. For hydrogen bonding to occur, a hydrogen atom must be covalently linked to a highly electronegative atom. This linkage strongly pulls the shared electrons toward the electronegative atom, leaving the hydrogen with a significant partial positive charge (\(\delta+\)) and the partner with a partial negative charge (\(\delta-\)).
The partially positive hydrogen atom is then attracted to the lone pair of electrons on a neighboring electronegative atom. This attraction is the hydrogen bond—an electrostatic interaction much weaker than a covalent bond, but substantially stronger than typical intermolecular forces. Its strength typically ranges from 4 to 50 kilojoules per mole. This attraction acts like a molecular glue, causing individual molecules to stick together more firmly.
The Strict Molecular Requirements
For a molecule to participate in hydrogen bonding, two distinct components must be present: the donor and the acceptor. The donor requirement involves a hydrogen atom covalently bonded to one of three small, highly electronegative atoms: Nitrogen (N), Oxygen (O), or Fluorine (F). This N-H, O-H, or F-H bond is the hydrogen bond donor. The high electronegativity of N, O, or F strips electron density from the hydrogen, leaving it exposed with a partial positive charge.
The second requirement is the hydrogen bond acceptor on a nearby molecule. The acceptor must be a separate N, O, or F atom possessing at least one non-bonding lone pair of electrons. This lone pair, carrying a partial negative charge, targets the partially positive hydrogen atom of the donor molecule. Both the high electronegativity and the small atomic size of N, O, and F are necessary conditions, as they concentrate the negative charge density and allow the positive hydrogen to approach closely enough for a strong attraction to form.
A molecule can act as both a donor and an acceptor, or just an acceptor, depending on its structure. Molecules containing O-H or N-H groups, such as water or alcohols, can both donate a hydrogen atom and accept one via their lone pairs. Molecules like certain ethers or ketones have oxygen atoms with lone pairs but lack hydrogen bonded directly to N, O, or F, meaning they can only act as acceptors.
Common Examples in Everyday Life and Biology
The most ubiquitous example of a hydrogen-bonding molecule is water (\(\text{H}_2\text{O}\)). Water has two O-H donor sites and two lone pairs on the oxygen atom that act as acceptor sites. This arrangement allows each water molecule to potentially form four hydrogen bonds with its neighbors, creating an extensive, three-dimensional network. Ammonia (\(\text{NH}_3\)) is another common example, using N-H bonds as donors and the nitrogen lone pair as an acceptor.
In organic chemistry, alcohols, which contain the hydroxyl (-OH) functional group, are strong hydrogen bonders. The O-H group allows them to both donate and accept hydrogen bonds. This capability explains why simple alcohols, such as ethanol, readily mix with water, forming strong intermolecular bonds.
Biological Examples
Hydrogen bonds are fundamental to stabilizing the structure of biological macromolecules. They hold the double helix of DNA together, forming precisely between complementary base pairs. Adenine pairs with Thymine using two hydrogen bonds, and Guanine pairs with Cytosine using three. In proteins, hydrogen bonds between the \(\text{N-H}\) groups and \(\text{C=O}\) groups of the polypeptide backbone create regular secondary structures, such as the alpha-helix and the beta-sheet.
The Influence on Physical Properties
The presence of hydrogen bonding dramatically alters the physical properties of a substance compared to molecules of similar size that lack this capability. Because these attractions are strong, significant extra energy is required to overcome them and separate the molecules from the liquid phase into the gaseous phase. This results in compounds with hydrogen bonding, such as water, exhibiting abnormally high boiling and melting points.
For example, water remains a liquid at room temperature and boils at \(100\,^{\circ}\text{C}\). In contrast, hydrogen sulfide (\(\text{H}_2\text{S}\)), a chemically similar molecule without hydrogen bonding, is a gas well below freezing. The increased intermolecular attraction also plays a large role in solubility. Substances that can form hydrogen bonds with water will easily dissolve in it because they can replace existing water-water bonds with new, favorable solute-water bonds.
The arrangement of hydrogen bonds also gives solid water, or ice, the property of being less dense than liquid water. When water freezes, the molecules arrange themselves into a rigid, open crystalline lattice structure to maximize hydrogen bonds. This ordered arrangement pushes the molecules farther apart than they are in the constantly shifting structure of liquid water, causing the solid to float.