All About Hydrogen Bonding

By Melissa

From the Beginning

               Henry Cavendish, a British scientist, was known for his discovery of hydrogen which was described by him as “inflammable air”. Later, a scientist by the name of Antoine Lavoisier reproduced the works of Cavendish and gave hydrogen its name, hydro, meaning water and genes, meaning forming, due to the concept, and now the definition, of the formation of water upon combustion. Hydrogen is the most abundant chemical element in the universe, but is relatively rare to find in its diatomic form on Earth. In addition, as with most elements, hydrogen can exist as isotopes, but unlike most, hydrogen is the only element with named isotopes, the most notable of which include protium and deuterium.  In its pure state, hydrogen is the lightest of all gases.  Consequently, very little is found in the Earth’s atmosphere due to its ability to gain enough kinetic energy from collisions with other gaseous molecules to attain escape velocity from the Earth.

Bonding

           A chemical bond can be described as an attraction between two atoms caused by an electromagnetic force between opposing charges, such as an electron and a proton. Hydrogen bonding can be defined as the interaction between a covalently attached hydrogen and an electronegative atom that produces a permanent dipole.

               In order for a hydrogen bond to form, a hydrogen atom must be covalently bonded to an electronegative atom. This may seem counterintuitive, since hydrogen only has room in its outer orbital for 1 additional electron, provided by the electronegative atom to which it is bonded.  However, the high electronegativity of the atom bonded to hydrogen causes the electron cloud surrounding hydrogen to shift towards its electronegative partner, thereby exposing hydrogen’s positively charged nucleus.  This gives hydrogen a partial positive character, allowing for dipole-dipole interactions (hydrogen bonds) to occur.  Common functional groups that can participate in hydrogen bonding are alcohols, carboxylic acids, amides and amines. Common to all of these functional groups is hydrogen covalently attached to oxygen or nitrogen (both electronegative atoms).

Distances and Angles

          Bond length is the average distance between the nuclei of two bonded atoms. Bond length is inversely proportional to bond strength: the shorter the bond, the stronger it is. The distance of a bond depends on temperature, pressure, and orbital hybridization.  Pictured below is the average hydrogen bond length between water molecules.

The ideal bond angle depends on the nature of the donor (defined as the atom covalently bonded to the interacting hydrogen in question), but for the most part linear (180°) tends to be the accepted. As a rule, the more hybridized the molecule is the less likely it will show ideal hydrogen bond angle.

Enthalpy

           The exothermic process of bond breaking gives information about bond enthalpies, the required energy to break a chemical bond. Bonds with hydrogen vary from weak (around 1.2 kJ mol-1) to very strong (less than 155 kJ mol-1). Bond strength is dependent on bond angle, environment (liquid, solid, gas), temperature and pressure. As temperature rises the hydrogen bond enthalpy goes down. The critical temperature for the internal bond of hydrogen is 647.14K at all pressures.

Applications of H-Bonds

           Hydrogen bonding proves important not only in chemistry but biology too. The nucleic acids that hold together the DNA helix are connected by hydrogen bonds. The overall structure of a protein is partly determined by how well it can from internal hydrogen bonds. For example, in β-sheets, the hydrogen bonding that occurs between two lengths of amino acids determines how structurally sound the overall protein will be.

The study of hydrogen bonds is continuing and although more experiments have been done, there is still much that remains to be discovered. 

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