Ionic compounds such as NaCl, CaBr 2 , and K 2 SO 4 exist as extended arrays of discrete
ions in the solid state. As we shall see in Section 13-16, the oppositely charged ions in
these arrays are quite close together. As a result of these small distances, d, the energies
of attraction in these solids are substantial. Most ionic bonding is strong, and as a result
most ionic compounds have relatively high melting points (Table 13-2). At high enough
temperatures, ionic solids melt as the added heat energy overcomes the potential energy
associated with the attraction of oppositely charged ions. The ions in the resulting liquid
samples are free to move about, which accounts for the excellent electrical conductivity
of molten ionic compounds.
For most substances, the liquid is less dense than the solid, but H 2 O is one of the rare
exceptions. Melting a solid nearly always produces greater average separations among the
particles. This means that the forces (and energies) of attractions among the ions in an
ionic liquid are less than in the solid state because the average dis greater in the melt.
However, these energies of attraction are still much greater in magnitude than the ener-
gies of attraction among neutral species (molecules or atoms).
The product qqincreases as the charges on ions increase. Ionic substances containing
multiply charged ions, such as Al^3 , Mg^2 , O^2 , and S^2 ions, usuallyhave higher melting
and boiling points than ionic compounds containing only singly charged ions, such as
Na, K, F, and Cl. For a series of ions of similar charges, the closer approach of
smaller ions results in stronger interionic attractive forces and higher melting points
(compare NaF, NaCl, and NaBr in Table 13-2).Dipole–Dipole Interactions
Permanent dipole–dipole interactions occur between polar covalent molecules
because of the attraction of the atoms of one molecule to the atoms of another
molecule (Section 7-9).Electrostatic forces between two ions decrease by the factor 1/d^2 as their separation,
d, increases. But dipole–dipole forces vary as 1/d^4. Because of the higher power of din
the denominator, 1/d^4 diminishes with increasing dmuch more rapidly than does 1/d^2.
As a result, dipole forces are effective only over very short distances. Furthermore, for
dipole–dipole forces, qand qrepresent only “partial charges,” so these forces are weaker
than ion–ion forces. Average dipole–dipole interaction energies are approximately 4 kJ
per mole of bonds. They are much weaker than ionic and covalent bonds, which have
typical energies of about 400 kJ per mole of bonds. Substances in which permanent dipole–
dipole interactions affect physical properties include bromine fluoride, BrF, and sulfurIonic bonding may be thought of as
both inter- and intramolecularbonding.
488 CHAPTER 13: Liquids and Solids
TABLE 13-2 Melting Points of Some Ionic CompoundsCompound mp (°C) Compound mp (°C) Compound mp (°C)NaF 993 CaF 2 1423 MgO 2800
NaCl 801 Na 2 S 1180 CaO 2580
NaBr 747 K 2 S 840 BaO 1923
KCl 770See the Saunders Interactive
General Chemistry CD-ROM,
Screen 13.4, Intermolecular Forces (2),
especially the subsection “Dipole–
Dipole Forces.”