The Foundations of Chemistry

dioxide, SO 2. Dipole–dipole interactions are illustrated in Figure 13-3. All dipole–dipole
interactions, including hydrogen bonding (discussed in the following section), are some-
what directional. An increase in temperature causes an increase in translational, rotational,
and vibrational motion of molecules. This produces more random orientations of mole-
cules relative to one another. Consequently, dipole–dipole interactions become less
important as temperature increases. All these factors make compounds having only dipole–
dipole interactions more volatile than ionic compounds.

Hydrogen Bonding

Hydrogen bondsare a special case of very strong dipole–dipole interaction. They are
not really chemical bonds in the formal sense.

Strong hydrogen bonding occurs among polar covalent molecules containing H and

one of the three small, highly electronegative elements—F, O, or N.

Like ordinary dipole–dipole interactions, hydrogen bonds result from the attractions
between 
atoms of one molecule, in this case H atoms, and the atoms of another
molecule. The small sizes of the F, O, and N atoms, combined with their high elec-
tronegativities, concentrate the electrons of these molecules around these atoms. This
causes an H atom bonded to one of these highly electronegative elements to become quite
positive. The 
H atom is attracted to a lone pair of electrons on an F, O, or N atom
other than the atom to which it is covalently bonded (Figure 13-4). The molecule that
contains the hydrogen-bonding 
H atom is often referred to as the hydrogen-bond donor;
the atom to which it is attracted is called the hydrogen-bond acceptor.
Recently, careful studies of light absorption and magnetic properties in solution and of
the arrangements of molecules in solids have led to the conclusion that the same kind of
attraction occurs (although more weakly) when H is bonded to carbon. In some cases,
very weak CXH,O “hydrogen bonds” exist. Similar observations suggest the existence

Figure 13-3 Dipole–dipole

interactions among polar molecules.

(a) Bromine fluoride, BrF. (b) Sulfur

dioxide, SO 2. Each polar molecule is

shaded with regions of highest

negative charge () darkest and

regions of highest positive charge

(
) lightest. Attractive forces are

shown as XAX, and repulsive forces

are shown as XRX. Molecules tend

to arrange themselves to maximize

attractions by bringing regions of

opposite charge together while

minimizing repulsions by separating

regions of like charge.

13-2 Intermolecular Attractions and Phase Changes 489

A

R

R

R

A

A

R

R

R

R

R

A

A

A

A A

A

A A

A A A A

A

A

A

A

A

A

A

shown as

Br

(a) (b)

F






 






























 




 



































  



shown

OOS as

See the Saunders Interactive

General Chemistry CD-ROM,

Screen 13.6, Hydrogen Bonding.