Topic 7 · page 5 of 14

Hydrogen bonds

A hydrogen bond needs a δ+ hydrogen on N, O or F pointing at a lone pair on a N, O or F atom of another molecule. It is the strongest intermolecular force and holds together liquid water, ammonia and hydrogen fluoride.

15 min read7.17.1(iii)7.2

Read first: Permanent dipole–permanent dipole interactions

What a hydrogen bond is 7.17.1(iii)

Nitrogen, oxygen and fluorine are small and very electronegative. When one of them is bonded to hydrogen, the bonding pair is pulled strongly towards it. A hydrogen atom has only one electron, so with that electron pulled away the hydrogen nucleus is left almost bare. The result is a hydrogen atom with a large δ+ charge, which can get very close to a neighbour. Each of N, O and F also has at least one lone pair of electrons, which a δ+ hydrogen can be attracted to.

Both parts are needed: the H must be directly bonded to N, O or F, and the other molecule needs a lone pair on N, O or F.

Two water molecules joined by a hydrogen bond. The left molecule has a delta plus hydrogen pointing at the right molecule. A dashed line runs straight from that hydrogen to a lone pair on the oxygen of the right molecule, which carries delta minus. Solid lines are covalent bonds. A key explains covalent bond, hydrogen bond, lone pair and partial charges.
A hydrogen bond between two water molecules: the dashed line is the hydrogen bond, the solid lines are covalent bonds.

A hydrogen bond is stronger than London forces and permanent dipole–dipole interactions, so it is the strongest of the three intermolecular forces. It is still much weaker than a covalent bond. Hydrogen bonds act in addition to London forces, which are still present.

MethodDrawing a hydrogen bond between two molecules

  1. Draw both molecules with their bonds and all their lone pairs (2 on O, 1 on N, 3 on F).
  2. Mark δ+ on the H atom that is bonded to N, O or F, and δ− on each N, O or F atom.
  3. Draw a dashed line from the δ+ H of one molecule to a lone pair on the N, O or F of the other molecule.
  4. Make the X–H···X line straight: the bond to H, the hydrogen bond and the lone pair all lie on one line.
  5. Label the dashed line "hydrogen bond".

Water, ammonia and hydrogen fluoride 7.2

The three small molecules H₂O, NH₃ and HF all have H bonded to N, O or F, and each has lone pairs. So each one forms hydrogen bonds with its own molecules in the liquid.

Pairs of molecules of water, ammonia and hydrogen fluoride, each pair joined by one hydrogen bond drawn as a dashed line from a delta plus hydrogen straight to a lone pair on the other molecule. Water has 2 hydrogen atoms and 2 lone pairs. Ammonia has 3 hydrogen atoms and 1 lone pair. Hydrogen fluoride has 1 hydrogen atom and 3 lone pairs.
One hydrogen bond in each of water, ammonia and hydrogen fluoride. Real ammonia is pyramidal and water is bent; the drawings are flat.

Each hydrogen bond uses one δ+ H and one lone pair. The numbers differ from molecule to molecule.

Molecule δ+ H atoms Lone pairs
H₂O 2 2
NH₃ 3 1
HF 1 3

Water has both: two H atoms to give and two lone pairs to receive. Ammonia has plenty of H but only one lone pair. Hydrogen fluoride has one H but three lone pairs. In ammonia and HF, one of the two runs out first.

One water molecule in the centre with four hydrogen bonds to four neighbours. Its two hydrogen atoms each point at a lone pair on a neighbouring molecule on the right, so it gives two hydrogen bonds. Its two lone pairs each receive a hydrogen atom from a neighbouring molecule on the left, so it receives two hydrogen bonds.
One water molecule can take part in four hydrogen bonds: two through its H atoms and two through its lone pairs.

A water molecule can take part in up to four hydrogen bonds. Count each bond once, and the average is two bonds per molecule: each bond needs one H and one lone pair, so the number of bonds per molecule cannot be more than the smaller of "H atoms" and "lone pairs". That is 2 for water, 1 for ammonia and 1 for HF. Water makes the most hydrogen bonds per molecule, which is why it behaves most unusually (pages 7 and 8).

Check yourself

  1. Define a hydrogen bond.
  2. Explain why an H atom bonded to N, O or F can form a hydrogen bond, but an H atom bonded to C cannot.
  3. Draw two molecules of ammonia, showing a hydrogen bond between them. Describe your drawing.
  4. For each of NH₃, H₂O and HF, give the number of δ+ H atoms and the number of lone pairs on the N, O or F.
  5. Hydrogen bonds are stronger than London forces. Are they stronger than a covalent bond?
  6. A student draws a hydrogen bond between two water molecules as a dashed line from the O of one molecule to the O of the other. What is wrong?

Answers cover these until you have tried the questions

  1. An attraction between a δ+ hydrogen atom bonded to N, O or F and a lone pair on a N, O or F atom of a different molecule.
  2. N, O and F are small and highly electronegative, so an H bonded to them is left strongly δ+. Each of them also has a lone pair that a δ+ H of another molecule can be attracted to. The C–H bond is almost non-polar (C and H have similar electronegativity), so the H on carbon is not left strongly δ+, and no hydrogen bond forms.
  3. Two NH₃ molecules drawn as in the ammonia pair in the figure above: each N has one lone pair and three N–H bonds. Mark δ− on each N and δ+ on each H. Draw a dashed line from one δ+ H of the first molecule straight to the lone pair on the N of the second molecule, with N–H···N in a straight line. Label it "hydrogen bond".
  4. NH₃: 3 H atoms, 1 lone pair. H₂O: 2 H atoms, 2 lone pairs. HF: 1 H atom, 3 lone pairs.
  5. No. A hydrogen bond is much weaker than a covalent bond. That is why molecules stay whole when water boils.
  6. A hydrogen bond must go from the δ+ H of one molecule to a lone pair on the O of the other. The line should start at an H atom, not at an O atom, and the O–H···O must be in a straight line, with the partial charges and lone pair shown.