Hydrogen Bonding in Molecules
A hydrogen bond is a strong attraction between a hydrogen atom covalently bonded to a very electronegative atom (N, O or F) and a lone pair on another N, O or F atom. In brief: hydrogen bonding is the reason why water boils at 100 °C although H2S, which is heavier, boils at −60 °C, and why ethanol (78 °C) boils so much higher than its isomer dimethyl ether (−25 °C).
Hydrogen bonds are weaker than covalent bonds but stronger than other intermolecular forces, and they play a decisive role in the properties of water, in the structure of proteins and in the pairing of DNA bases. This article explains what a hydrogen bond is, which molecules form them, how they affect boiling points, and what they do in biology.
What is a hydrogen bond?
In a polar bond such as O–H, the hydrogen atom has a considerable positive partial charge, since the electrons are pulled toward oxygen, and the small hydrogen atom has no inner electrons to shield its nucleus. This exposed proton-like end is strongly attracted to a lone pair on a neighboring electronegative atom. The result is written as X–H···Y, where X and Y are N, O or F, and the dots represent the hydrogen bond.
The energy of a hydrogen bond is typically 10–40 kJ/mol, with about 20 kJ/mol between water molecules. The bond is directional: it is strongest when X, H and Y are in a straight line. The O···H distance in liquid water is about 180 pm, compared with the covalent O–H length of 96 pm.
Which molecules form hydrogen bonds?
Two conditions must be met:
- The molecule needs an H atom directly bonded to N, O or F (the hydrogen bond donor).
- Another N, O or F atom with a lone pair (the acceptor) must be nearby.
Water, ammonia, hydrogen fluoride, alcohols, carboxylic acids, amines and amides all form hydrogen bonds with themselves. Molecules such as acetone and diethyl ether have acceptor atoms but no N–H or O–H group, so they cannot form hydrogen bonds with themselves, but they can with water. A C–H bond is generally too weakly polar to be a donor, except in special cases such as chloroform.
How do hydrogen bonds affect boiling points?
The boiling points of the hydrides of groups 14 to 17 show the effect clearly. In group 14 the boiling points rise smoothly with molar mass, as expected for dispersion forces. In groups 15, 16 and 17, the first member of each family (NH3, H2O and HF) boils much higher than the trend of the heavier members would predict.
| Period | Group 14 | Group 15 | Group 16 | Group 17 |
|---|---|---|---|---|
| 2 | CH4: −162 | NH3: −33 | H2O: 100 | HF: 20 |
| 3 | SiH4: −112 | PH3: −88 | H2S: −60 | HCl: −85 |
| 4 | GeH4: −89 | AsH3: −62 | H2Se: −41 | HBr: −67 |
| 5 | SnH4: −52 | SbH3: −17 | H2Te: −2 | HI: −35 |
Without hydrogen bonding, water would be expected to boil at about −80 °C, and would be a gas at room temperature. In HF each molecule can form only one strong hydrogen bond per H atom, but in water each molecule has two hydrogen atoms and two lone pairs, so it can form up to four hydrogen bonds, which is why water’s boiling point is higher than that of HF or NH3.
Why is water so unusual?
- High boiling point and heat capacity: a lot of energy is needed to break hydrogen bonds, which moderates the climate and the temperature of living organisms.
- High surface tension: 72 mN/m at 25 °C, enabling insects to walk on water.
- Ice floats: in ice each molecule is held by four hydrogen bonds in an open hexagonal lattice, and this structure is less dense (0.917 g/cm3) than liquid water (0.9998 g/cm3 at 0 °C). Lakes freeze from the top down, so aquatic life can survive the winter.
- Excellent solvent: water forms hydrogen bonds to many solutes, such as alcohols, sugars and ions.
Hydrogen bonding in organic molecules
Ethanol, CH3CH2OH, and dimethyl ether, CH3OCH3, have the same formula, C2H6O, and the same molar mass of 46.07 g/mol. Ethanol has an O–H group and forms hydrogen bonds, and it boils at 78.4 °C. Dimethyl ether has no O–H group and boils at −24.8 °C, which is a difference of more than 100 °C.
Carboxylic acids form hydrogen-bonded pairs, called dimers. Acetic acid (60 g/mol) boils at 118 °C, higher than propan-1-ol (60 g/mol, 97 °C), because its dimer has two hydrogen bonds. Even in the gas phase, acetic acid partly occurs as dimers.
What do hydrogen bonds do in biology?
- DNA: the two strands are held together by hydrogen bonds between base pairs: adenine–thymine (two bonds) and guanine–cytosine (three bonds). They are strong enough to keep the helix stable but weak enough to be separated when the DNA is copied.
- Proteins: hydrogen bonds between the C=O and N–H groups of the backbone stabilize α-helices and β-sheets.
- Cellulose and wool: hydrogen bonds between chains give strength to plant fibers.
- Enzymes: hydrogen bonds help to position substrates in the active site.
Who discovered the hydrogen bond?
Wendell Latimer and Worth Rodebush published the idea of a hydrogen atom bridging two atoms in 1920, to explain the properties of water and ammonia. Linus Pauling developed the concept in his book The Nature of the Chemical Bond (1939). In the 1950s, hydrogen bonds proved to be essential to the structure of proteins, as seen in the α-helix described by Pauling and Robert Corey in 1951, and to the double helix of DNA found by Watson and Crick in 1953.
Which mistakes are common?
- Calling it a true bond: a hydrogen bond is an intermolecular attraction, and not a covalent bond within the molecule.
- Assuming that any molecule with hydrogen can form hydrogen bonds: the hydrogen must be attached to N, O or F.
- Drawing the bond between two H atoms: the bond is between H and the lone pair of another atom.
- Saying that HCl forms hydrogen bonds: chlorine is not electronegative enough and too large.
- Ignoring the number of possible bonds: the number of donors and acceptors determines how extensive the network is.