Hydrogen bond
|
Snapshot from a simulation of liquid water. The four thin green lines from the molecule in the center of the picture represent hydrogen bonds. |
In
chemistry, a
hydrogen bond is a type of attractive
intermolecular force that exists between two
partial electric charges of opposite polarity. Although stronger than most other
intermolecular forces, the typical hydrogen bond is much weaker than both the
ionic bond and the
covalent bond. Within
macromolecules such as
proteins and
nucleic acids, it can exist between two parts of the same molecule, and figures as an important constraint on such molecules' overall shape.
As the name "hydrogen bond" implies, one part of the bond involves a
hydrogen atom. The hydrogen atom must be attached to one of the elements
oxygen,
nitrogen or
fluorine, all of which are strongly
electronegativeheteroatoms. These bonding elements are known as the hydrogen-bond
donor. This electronegative element attracts the electron cloud from around the hydrogen nucleus and, by decentralizing the cloud, leaves the atom with a positive partial charge. Because of the small size of hydrogen relative to other atoms and molecules, the resulting charge, though only partial, nevertheless represents a large charge density. A hydrogen bond results when this strong positive charge density attracts a
lone pair of electrons on another
heteroatom, which becomes the hydrogen-bond
acceptor.
The hydrogen bond is not like a simple attraction between point charges, however. It possesses some degree of orientational preference, and can be shown to have some of the characteristics of a covalent bond. This covalency tends to be more extreme when acceptors bind hydrogens from more electronegative donors.
Strong covalency in a hydrogen bond raises the questions: "To which molecule or atom does the hydrogen
nucleus belong?" and "Which should be labeled 'donor' and which 'acceptor'?" According to chemical convention, the donor generally is that atom to which, on separation of donor and acceptor, the retention of the hydrogen nucleus (or
proton) would cause no increase in the atom's positive charge. The acceptor meanwhile is the atom or molecule that would become more positive by retaining the positively charged proton. Liquids that display hydrogen bonding are called
associated liquids.
Hydrogen bonds can vary in strength from very weak (1-2 kJ mol
−1) to so strong (40 kJ mol
−1) so as to be indistinguishable from a covalent bond, as in the ion HF
2−. Typical values include:
* O—H
...:N (7 kcal/mol)
* O—H
...:O (5 kcal/mol)
* N—H
...:N (3 kcal/mol)
* N—H
...:O (2 kcal/mol)
The length of hydrogen bonds depends on bond strength, temperature and pressure. The typical length of a hydrogen bond in water is 1.97 Å (197 pm).
The most ubiquitous, and perhaps simplest, example of a hydrogen bond isfound between
water molecules. In a discrete water molecule, water has two hydrogen atoms and one oxygen atom. Two molecules of water can form a hydrogen bond between them; the simplest case, when only two molecules are present, is called the
water dimer and is often used as a model system. When more molecules are present, as is the case in liquid water, more bonds are possible because the oxygen of one water molecule has two lone pairs of electrons, each of which can form a hydrogen bond with hydrogens on two other water molecules. This can repeat so that every water molecule is H-bonded with up to four other molecules, as shown in the figure (two through its two lone pairs, and two through its two hydrogen atoms.)
Liquid water's high
boiling point is due to the high number of hydrogen bonds each molecule can have relative to its low
molecular mass. Water is unique because its oxygen atom has two lone pairs and two hydrogen atoms, meaning that the total number of bonds of a water molecule is up to four. For example, hydrogen fluoride - which has three lone pairs on the F atom but only one H atom - can have a total of only two bonds (
ammonia has the opposite problem: three hydrogen atoms but only one lone pair).
H-F
...H-F
...H-F
The exact number of hydrogen bonds in which a molecule in liquid water participates fluctuates with time and depends on the temperature. From
TIP4P liquid water simulations at 25 °C, it was estimated that each water molecule participates in an average of 3.59 hydrogen bonds. At 100 °C, this number decreases to 3.24 due to the increased molecular motion and decreased density, while at 0 °C, the average number of hydrogen bonds increases to 3.69 (
Mol. Phys. 1985,
56, 1381). A more recent study (
J. Chem. Phys 2005,
123, 104501) found a much smaller number of hydrogen bonds: 2.357 at 25 °C. The differences may be due to the use of a different method for defining and counting the hydrogen bonds.
Were the bond strengths more equivalent, one might instead find the atoms of two interacting water molecules partitioned into two
polyatomic ions of opposite charge, specifically
hydroxide (OH
−) and
hydronium (H
3O
+) (Hydronium ions are also known as 'hydroxonium' ions.)
H-O
− H
3O
+Indeed, in pure water under conditions of
standard temperature and pressure, this latter formulation is applicable only rarely; on average about one in every 5.5 × 10
8 molecules gives up a proton to another water molecule, in accordance with the value of the
dissociation constant for water under such conditions. It is a crucial part of uniqueness of water
Hydrogen bonding also plays an important role in determining the three-dimensional structures adopted by proteins and nucleic bases. In these macromolecules, bonding between parts of the same macromolecule cause it to fold into a specific shape, which helps determine the molecule's physiological or biochemical role. The double helical structure of
DNA, for example, is due largely to hydrogen bonding between the
base pairs, which link one complementary strand to the other and enable
replication.
In proteins, hydrogen bonds form between the backbone oxygens and amidehydrogens. When the spacing of the
amino acid residues participating ina hydrogen bond occurs regularly between positions
i and
i + 4, an
alpha helix is formed. When the spacing is less, between positions
iand
i + 3, then a 3
10 helix is formed. When two strands arejoined by hydrogen bonds involving alternating residues on eachparticipating strand, a
beta sheet is formed. Hydrogen bondings also play a part in forming the tertiary structure of protein through interaction of R-groups.(See also
protein folding).
Symmetric hydrogen bonds have been observed recently spectroscopically in
formic acid at high pressure (>GPa). Each hydrogen atom forms a partial covalent bond with two atoms rather than one. Symmetric hydrogen bonds have been postulated in ice at high pressure (ice-X). See references below (Goncharov, et al.)
The hydrogen bond can be compared with the closely related
dihydrogen bond, which is also an
intermolecular bonding interaction involving hydrogen atoms. These structures have been known for some time, and well characterized by
crystallography; however, an understanding of their relationship to the conventional hydrogen bond,
ionic bond, and
covalent bond remains unclear. Generally, the hydrogen bond is characterized by a proton acceptor that is a lone pair of electrons in nonmetallic atoms (most notably in the
nitrogen, and
chalcogen groups). In some cases, these proton acceptors may be
pi-bonds or
metal complexes. In the dihydrogen bond, however, a metal hydride serves as a proton acceptor; thus forming a hydrogen-hydrogen interaction.
Neutron diffraction has shown that the
molecular geometry of these complexes are similar to hydrogen bonds, in that the bond length is very adaptable to the metal complex/hydrogen donor system.
The hydrogen bond remains a fairly mysterious object in the theoretical study of
quantum chemistry and
physics. Most generally, the hydrogen bond can be viewed as a
metric dependent
electrostatic scalar field between two or more intermolecular bonds. This is slightly different from the
intramolecular bound states of, for example,
covalent or
ionic bonds; however, hydrogen bonding is generally still a bound state phenomenon, since the
interaction energy has a net negative sum. The question of the relationship between the covalent bond and the hydrogen bond remains largely unsettled, though the initial theory proposed by
Linus Pauling suggests that the hydrogen bond has a partial covalent nature. While a lot of experimental data has been recovered for hydrogen bonds in
water, for example, that provide good resolution on the scale of intermolecular distances and
molecular thermodynamics, the
kinetic and
dynamical properties of the hydrogen bond in
dynamic systems remains largely mysterious.
* George A. Jeffrey.
An Introduction to Hydrogen Bonding (Topics in Physical Chemistry). Oxford University Press, USA (March 13, 1997). ISBN 0195095499
*A New Intermolecular Interaction: Unconventional Hydrogen Bonds with Element-Hydride Bonds as Proton Acceptor Robert H. Crabtree, Per E. M. Siegbahn, Odile Eisenstein, Arnold L. Rheingold, and Thomas F. Koetzle
Acc. Chem. Res. 1996,
29(7), 348 - 354.
* Polymerization of Formic Acid under High Pressure Alexander F. Goncharov, M. Riad Manaa, Joseph M. Zaug, Richard H. Gee, Laurence E. Fried, and Wren B. Montgomery
Phys. Rev. Lett. 2005,
94, 065505.
