Hydrogen bond

Individual atoms have the ability to bond together and form molecular compounds. In doing so, certain molecules create what are called hydrogen bonds. Hydrogen bonds are created whenever a hydrogen atom is covalently bonded to a highly electronegative atom (that is, an atom that strongly attracts electrons) in the vicinity of another electronegative atom containing a lone electron pair (that is, a pair of valence electrons that is not shared with another atom). The atom bonded to the hydrogen pulls the hydrogen atom’s sole electron away from its nucleus, creating a region of positive charge that attracts the lone electron pair of the second electronegative atom.

Water is a good example of this. Its molecular formula is H₂O, meaning that two hydrogen atoms are paired with a single oxygen atom to form one molecule of water. Because oxygen is more electronegative than hydrogen, the electrons are drawn closer to the oxygen atom, causing it to exhibit a slight negative charge. Conversely, the hydrogen atoms develop a slight positive charge, as the lack of electrons surrounding them exposes the positively charged proton in each atom’s nucleus. When more than one water molecule is present, the exposed hydrogen nuclei of one molecule are attracted to the increased negative charge of another molecule’s oxygen atom. This attraction between molecules is the hydrogen bond.

Brief History

In his 1939 work The Nature of the Chemical Bond and the Structure of Molecules and Crystals, American chemist Linus Pauling (1901–94) credited two men, Thomas Field Winmill (ca. 1888–1953) and Tom Sidney Moore (1881–1967), with being the first to describe hydrogen bonds in a 1912 paper in the Journal of the Chemical Society. However, the theory of hydrogen bonding was not formally proposed until 1920, when Kansas-born chemists Wendell Mitchell Latimer (1893–1955) and Worth H. Rodebush (1887–1959), while working together at the University of California, Berkeley, published a paper in the Journal of the American Chemical Society describing how hydrogen bonding works.

In their paper, Latimer and Rodebush cited the unpublished thesis of fellow scientist Maurice Loyal Huggins (1897–1981), who was working in their laboratory at the time. In the thesis, which he had written the previous year, Huggins used the theory of hydrogen bonding to account for why certain organic compounds behaved the way they did.

In 2011, a new, evidence-based definition of hydrogen bonding was published in Pure and Applied Chemistry, the official journal of the International Union of Pure and Applied Chemistry (IUPAC). The report was quite technical and shed further light on hydrogen bonding and its functions.

Overview

A hydrogen bond is the strong attraction between one atom that is covalently bonded to a hydrogen atom and another atom that has a lone pair of electrons. It is not a true atomic bond but rather a form of dipole-dipole attraction, which is the attraction between the positive end of one polar molecule, or dipole, and the negative end of another. Hydrogen bonds are much easier to break apart than covalent or ionic bonds, but they are also much easier to reform. Some hydrogen bonds are intramolecular, meaning that the bond is between atoms within a single molecule. When a hydrogen bond forms between different molecules, it is said to be “intermolecular.”

A hydrogen bond can form only if the atom bonded to the hydrogen is sufficiently electronegative. Electronegativity refers to an atom’s ability to attract electrons. An atom with low electronegativity will not be able to attract electrons as strongly as one with high electronegativity. Hydrogen bonds cannot form when there is not a large difference in electronegativity between the hydrogen atom and the atom to which it is bonded. For example, it would be impossible for a phosphorous atom bonded to a hydrogen atom to form a hydrogen bond with another atom because their electronegativities are too close together (2.1 for hydrogen, 2.19 for phosphorus) to create a dipole. This is why hydrogen bonds form so easily between water molecules: oxygen is very electronegative, and hydrogen is not. As a general rule, elements on the left-hand side of the periodic table have lower electronegativity, while elements on the right-hand side of the periodic table have higher electronegativity.

Hydrogen bonds are crucial parts of living organisms for a variety of reasons. Without hydrogen bonding, DNA, the blueprint for life, could not function the way it does. A DNA molecule consists of two strands of nucleotides entwined in a double helix; the two strands are held together and given their structure by the hydrogen bonds that form between the paired nucleotides of each strand.

Hydrogen bonding plays a similar role in protein folding. For a protein to function properly within an organism, it needs to achieve a certain molecular structure. Without hydrogen bonding, many of these structures would be difficult, if not impossible, to achieve. For example, hemoglobin is a protein responsible for transporting oxygen molecules throughout the body. If hemoglobin molecules were not able to form hydrogen bonds, then they would not be able to do their job effectively, and the body would be deprived of much-needed oxygen.

In general, without the help of hydrogen bonding, almost no protein molecules would be able to maintain their shape. Whenever a protein denatures, or loses its folded structure, it is no longer able to do its job. Hydrogen bonding provides the protein with a stable structure and helps it maintain that structure when exposed to heat or acidic environments.

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