The Foundations of Chemistry

of weak hydrogen bonds to chlorine atoms, such as OXH,Cl. However, most chemists

usually restrict usage of the term “hydrogen bonding” to compounds in which H is cova-

lently bonded to F, O, or N, and we will do likewise throughout this text.

Typical hydrogen-bond energies are in the range 15 to 20 kJ/mol, which is four to five

times greater than the energies of other dipole–dipole interactions. As a result, hydrogen

bonds exert a considerable influence on the properties of substances. Hydrogen bonding

is responsible for the unusually high melting and boiling points of compounds such as

water, methyl alcohol, and ammonia compared with other compounds of similar molec-

ular weight and molecular geometry (Figure 13-5). Hydrogen bonding between amino

acid subunits is very important in establishing the three-dimensional structures of proteins.

Figure 13-4 Hydrogen bonding
(indicated by dashed lines) in (a)
water, H 2 O; (b) methyl alcohol,
CH 3 OH; and (c) ammonia, NH 3.
Hydrogen bonding is a special case
of very strong dipole interaction.

490 CHAPTER 13: Liquids and Solids

(a) (b) (c)

Figure 13-5 Boiling points of some hydrides as a function of molecular weight. The

unusually high boiling points of NH 3 , H 2 O, and HF compared with those of other hydrides

of the same groups are due to hydrogen bonding. The electronegativity difference between

H and C is small, and there are no unshared pairs on C; thus, CH 4 is not hydrogen bonded.

Increasing molecular weight corresponds to increasing number of electrons; this makes the

electron clouds easier to deform and causes increased dispersion forces, accounting for the

increase in boiling points for the nonhydrogen-bonded members of each series.

GeH 4

SnH 4

Molecular weight

Boiling temperature (

°C)

20 40 60 80 100 120 140

 200

 150

 100

 50

0

50

100

150

Group VA

NH 3

Group VIA

H 2 O

Group IVA

H 2 S

H 2 Se

H 2 Te

SiH 4

CH 4

Group VIIA

HF

HCl HBr

HI

PH 3

AsH 3

SbH 3