Hydrogen bonds are not covalent. These two terms describe fundamentally different types of chemical interactions governing how atoms and molecules behave. Covalent bonds are strong, internal forces that hold atoms together to form single molecules, while hydrogen bonds are much weaker, external forces that attract one molecule to another. Understanding this distinction is necessary for grasping the structure of all matter, especially complex biological molecules.
Defining Covalent Bonds
Covalent bonds form when two atoms share one or more pairs of electrons, creating a powerful, stable link between them. This electron sharing is a strong, intramolecular force that binds atoms together to form a molecule, such as the water molecule (\(\text{H}_2\text{O}\)). The sharing allows each participating atom to achieve a more stable electron configuration, requiring significant energy input to break. Because the electrons are physically shared between the atomic nuclei, the resulting bond is represented by a solid line in chemical diagrams.
Defining Hydrogen Bonds
Hydrogen bonds are an attractive force that occurs between separate molecules, making them an intermolecular force. This attraction is based on electrostatic forces involving partial charges, or dipoles, rather than electron sharing. A hydrogen bond forms when a hydrogen atom, already covalently bonded to a highly electronegative atom (oxygen, nitrogen, or fluorine), is attracted to another nearby electronegative atom. The highly electronegative atom pulls the shared electrons closer, leaving the hydrogen atom with a partial positive charge (\(\delta+\)), which is then electrostatically drawn toward a neighboring atom carrying a partial negative charge (\(\delta-\)). The hydrogen bond is represented by a dotted or dashed line to signify this weaker attraction.
Key Differences in Bond Strength and Function
The most significant difference lies in their relative strengths, which dictates their function in a system. Covalent bonds are extremely strong, requiring hundreds of kilojoules per mole of energy to break, typically ranging from 100 to 1,100 \(\text{kJ/mol}\). Hydrogen bonds are substantially weaker, possessing only about 5% of the strength of a typical covalent bond, with energy falling between 5 and 50 \(\text{kJ/mol}\).
Covalent bonds hold atoms tightly together to create the molecule itself, determining its fundamental shape and stability. Hydrogen bonds, being weak and easily broken, influence how molecules interact with each other. They are responsible for collective properties, such as water’s ability to remain liquid or the folding of a long polymer chain into a specific three-dimensional structure. This difference in strength allows hydrogen bonds to be dynamic, forming and breaking rapidly as needed for biological processes.
The Importance of Hydrogen Bonds in Biological Systems
The relative weakness and dynamic nature of hydrogen bonds make them indispensable to life. They are responsible for water’s unique properties, such as its high boiling point and function as an excellent solvent. The continuous network of hydrogen bonds between water molecules creates cohesion, allowing for processes like the transport of nutrients in plants.
In large biological molecules, hydrogen bonds provide the necessary temporary stability for function. They hold the two strands of the DNA double helix together via specific base pairings (adenine to thymine, and guanine to cytosine). These attractions also dictate the intricate, folded three-dimensional shapes of proteins. The ability of these bonds to be easily broken and reformed allows DNA to be unzipped for replication and proteins to change shape for enzymatic activity, processes impossible if the connections were permanent covalent bonds.