The gas phase is unique among the three states of matter in that there are some simple models we can use to predict the physical behavior of all gases—independent of their identities. We cannot do this for the solid and liquid states. In fact, the development of this understanding of the behavior of gases represents the historical dividing point between alchemy and modern chemistry. Initial advances in the understanding of gas behavior were made in the mid 1600s by Robert Boyle, an English scientist who founded the Royal Society, one of the world’s oldest scientific organizations.

How is it that we can model all gases independent of their chemical identity? The answer is in a group of statements called the kinetic molecular theory of gases:

• Gases are composed of tiny particles that are separated by large distances.
• Gas particles are constantly moving, experiencing collisions with other gas particles and the walls of their container.
• The velocity of gas particles is directly proportional to the absolute temperature of a gas.
• Gas particles do not experience any force of attraction or repulsion with each other.

Notice that none of these statements relates to the identity of the gas. This means that all gases should behave similarly. A gas that follows these statements perfectly is called an ideal gas. Most gases show slight deviations from these statements and are called real gases. However, the existence of real gases does not diminish the importance of the kinetic theory of gases.

The kinetic theory also states that there is no interaction between individual gas particles. Although we know that there are, in fact, intermolecular interactions in real gases, the kinetic theory assumes that gas particles are so far apart that the individual particles don’t “feel” each other. Thus, we can treat gas particles as tiny bits of matter whose identity isn’t important to certain physical properties.

One of the statements of the kinetic theory mentions collisions. As gas particles are constantly moving, they are also constantly colliding with each other and with the walls of their container. There are forces involved as gas particles bounce off the container walls (Figure 8.9 “Gas Pressure”). The force generated by gas particles divided by the area of the container walls yields pressure. Due to the large amount of empty space between molecules in a gas, gases are compressible, so the pressure must be specified when other measurements are reported for gases.  Solids and liquids also experience and exert pressures,  but under everyday conditions, pressure does not affect their properties.

The basic unit of pressure is the newton per square meter (N/m²). This combined unit is redefined as a pascal (Pa). One pascal is not a very large amount of pressure. A more useful unit of pressure is the bar, which is 100,000 Pa (1 bar = 100,000 Pa). Other common units of pressure are the atmosphere (atm), which was originally defined as the average pressure of Earth’s atmosphere at sea level; and mmHg (millimeters of mercury), which is the pressure generated by a column of mercury 1 mm high. The unit millimeters of mercury is also called a torr, named after the Italian scientist Evangelista Torricelli, who invented the barometer in the mid-1600s. A more precise definition of atmosphere, in terms of torr, is that there are exactly 760 torr in 1 atm. A bar equals 1.01325 atm. Given all the relationships between these pressure units, the ability to convert from one pressure unit to another is a useful skill.