In Section 9.1 and again in Section 10.2B we discussed the characteristics of solids. These properties suggest that the particles (ions, molecules, or atoms) in a solid occupy fixed positions from which they cannot easily move. This orderly arrangement is called the crystal structure or crystal lattice of the solid. Strong attractive forces between the particles keep them in this arrangement. Even so, as suggested in Figure 10.9, each particle in a solid has some kinetic energy and is in

FIGURE 10.9 The crystal structure of a solid, showing the movement of the component ions, atoms, or molecules within their assigned space.

constant motion within its space in the solid (unless the solid is at a temperature of absolute zero, at which all motion ceases). The ions, atoms, or molecules of which the solid is composed vibrate, rotate, and even move around within their assigned space in the crystal structure.

As in gas and liquid samples, particles in a solid have a distribution of kinetic energies that depends on the temperature of the sample. At every temperature, some particles have enough energy to overcome the forces that hold them in place and escape from the solid as a vapor. Hence each solid, like each liquid, has a vapor pressure that increases as the temperature increases. The process by which atoms or molecules go directly from the solid state to the gaseous state is called sublimation. Solids that sublime have unusually high vapor pressures.

Dry ice (solid carbon dioxide) is a familiar example of a solid that sublimes. Solid water (ice) sublimes. If you live in a cold, dry climate, you may have noticed that ice disappears (actually it sublimes to a colorless gas) at temperatures below 0°C. Moth balls sublime; they disappear after protecting woolens during a hot summer.

As a solid is heated, its change in temperature depends on its specific heat. The solution of specific heat problems is the same whether for a liquid (Section 10.3B) or a solid (Section 2.5B).

A solid melts (changes to a liquid) when the average kinetic energy of its particles is high enough to overcome the attractive forces between them. By definition, the melting point of a solid is the temperature at which the liquid and solid states are in equilibrium at a pressure of 1 atm. We show this equilibrium in equation form by writing "solid" as a reactant and "liquid" as a product and connecting them with a double (equilibrium) arrow:

During melting, both solid and liquid are present until all the solid is converted to liquid. Conversely, during freezing, solid and liquid are present until all the liquid is converted to solid. The processes are indicated by writing "melting" over the forward equilibrium arrow and "freezing" under the reverse equilibrium arrow. Because freezing is the reverse of melting, the freezing point of a substance is the same as its melting point. The molar heat of fusion, H_fus, of a substance is the amount of heat that must be supplied to convert one mole of that substance from a solid to a liquid at its melting point. Calculations involving heats of fusion are similar to those involving heats of vaporization.

Example: The molar heat of fusion of carbon tetrachloride (CCl4) is 3.26 kJ/mol. How many jouls must be supplied to 34 g solid carbon tetrachloride at is melting point to change the sample to a liquid? Solution Wanted ?J Given 34 g CCl4 Conversion factors Equation Answer 0.72 kJ