Chapter 7

The Gas Phase

Among the different phases of matter, the gaseous phase is the simplest to understand and to
model, since all gases, to a first approximation, display similar behavior and follow similar laws
regardless of their identity. The atoms or molecules in a gaseous sample move rapidly and are far
apart. In addition, intermolecular forces between gas particles tend to be weak; this results in
certain characteristic physical properties, such as the ability to expand to fill any volume and to take
on the shape of a container. Furthermore, gases are easily, though not infinitely, compressible.

The state of a gaseous sample is generally defined by four variables: pressure (P), volume (V),
temperature (T), and number of moles (n), though as we shall see, these are not all independent.
The pressure of a gas is the force per unit area that the atoms or molecules exert on the walls of the
container through collision. The SI unit for pressure is the pascal (Pa), which is equal to 1 newton per
meter squared. (The SI units are those that are based on the simple metric units of kilogram, meter,
second, et cetera. A newton, for example, is a kg•m/s^2 .) In chemistry, however, gas pressures are
more commonly expressed in units of atmospheres (atm) or millimeters of mercury (mm Hg or torr),
which are related as follows:

1   atm =   760 mm  Hg  =   760 torr    =   101,325 Pa

Volume is generally expressed in liters (L) or milliliters (mL). The temperature of a gas is usually
given in Kelvin (K, not ºK), and its value, also known as the absolute temperature, is related to the
temperature in degrees Celsius by the expression T(K) = T(ºC) + 273.15. Gases are often discussed in
terms of standard temperature and pressure (STP), which refers to conditions of 273.15 K (0ºC) and
1 atm.

Measurement of  Gas Pressures

Ideal   Gases

Kinetic Molecular   Theory  of  Gases

Descriptive Chemistry   of  Some    Common  Gases