# Gases, Properties of

## Definition

The fundamental physical properties of a gas are related to its temperature, pressure, and volume. These properties can be described and predicted by a set of equations known as the gas laws. While these laws were originally based on mathematical interpretations for an ideal or perfect gas, modern atomic and kinetic theory of gases has led to a modified expression that more accurately reflects the properties of real gases.

## Description

Current understanding of gas properties came as a result of study of the interaction between volume, pressure, and temperature. Robert Boyle was the first to describe the relationship between the volume and pressure of a gas. In 1660 he learned that if an enclosed amount of a gas is compressed to half its original volume while the temperature is kept constant, the pressure will double. He expressed this mathematically as PV = constant, where P stands for pressure, V stands for volume, and the value of the constant depends on the temperature and the amount of gas present. This expression is known as Boyle's law.

The second fundamental property of gases was defined by Jacques Charles in 1787. He found that the temperature and volume of a gas are directly related. Charles observed that a number of gases expanded equally as heat was applied and the pressure was kept constant.

Charles's ideas were expanded upon in research by others in the field, most notably Joseph Gay-Lussac, who also studied the thermal expansion of gases. The volume/temperature relationship is known as Charles's law.

The third property of gases was described by Gay-Lussac who, in addition to his work with volume and temperature, researched the connection between pressure and temperature. In 1802, he formulated an additional law. These three laws can be combined into one generalized equation that expresses the interrelation between pressure, temperature and volume. This equation, called the ideal gas law, is written as PV = nRT.

While the ideal gas law works very well in predicting gas properties at normal conditions, it does not accurately represent what happens under extreme conditions. Neither does it account for the fact that real gases can undergo phase change to a liquid form. Modern atomic theory helps explain these discrepancies. It describes molecules as having a certain freedom of motion in space. Molecules in a solid material are arranged in a regular lattice such that their freedom is restricted to small vibrations about lattice sites. Gas molecules, on the other hand, have no macroscopic spatial order, and they can move about their containers at random. The motion of these particles can be described by the branch of physics known as classical mechanics. The study of this particulate motion is known as the kinetic theory of gases. It states that the volume of a gas is defined by the position distribution of its molecules. In other words, the volume represents the available amount of space in which a molecule can move. The temperature of the gas is proportional to the average kinetic energy of the molecules, or to the square of the average velocity of the molecules. The pressure of a gas, which can be measured with gauges placed on the container walls, is a function of the particle momentum, which is the product of the mass of the particles and their speed.