Gas is one of the three basic states of matter. The other two states are solid and liquid. These states differ from each other in the way they fill space and change shape. A solid, such as rock, always occupies a fixed volume (amount of space) and has a fixed shape. A liquid, such as water, always occupies a fixed volume. But it has no shape of its own, so it takes on the shape of its container. A gas, such as air, has neither a fixed shape nor a fixed volume. It fills any container that holds it and takes on the container’s shape. Like solids and liquids, gases have weight. But gases are thinner and lighter than solids and liquids.
Many gases, including the nitrogen and oxygen in air, have no color or odor. They can be identified by their chemical behavior, their weight, their ability to absorb heat, and their other properties. But some gases have a color, or an odor, or both. For example, nitrogen dioxide is brown. Hydrogen sulfide smells like rotten eggs.
Under special conditions, gases change into a fourth state of matter called a plasma. Plasmas are formed by heating a gas to an extremely high temperature or by passing an electric current through it. Matter exists in a plasma state in stars and the regions between stars.
How gases behave.
The behavior of gases is explained by what scientists call the kinetic molecular theory . According to this theory, all matter is made of constantly moving particles—atoms or molecules. An atom is one of the basic units of matter, and a molecule is a combination of atoms. The number of atoms or molecules of gas in a container the size of a pin-head is many millions of times as large as the number of people on the earth. But these gas particles are so small that they occupy only about one-thousandth of the space inside the container. The remaining space between the particles is empty.
Gas particles fly around in all directions at about the speed of sound. Their exact speed is determined by their weight and by the temperature of the gas. Gas particles move faster when the gas is hot than when it is cold. But light particles move faster than heavy ones at all temperatures. Each moving gas particle crashes into billions of other particles each second. Gas particles crashing into the walls of their container produce an effect called pressure.
A gas liquefies (changes to a liquid) when it is cooled to a temperature called its boiling point. At this temperature, the gas particles gather together to form a liquid. If the pressure of the gas is increased, it liquefies at a higher temperature. But pressure can raise the liquefying temperature only to a limiting value called the critical temperature. For example, oxygen under normal atmospheric pressure liquefies at its boiling point, -183 °C. But under a pressure of 5,171 kilopascals, oxygen liquefies at -119 °C, its critical temperature.
Gas laws.
Three laws explain approximately how the pressure, temperature, volume, and the number of particles in a container of gas are related. These laws are Boyle’s law, Charles’s law, and Avogadro’s law.
Boyle’s law
says that pressure increases as the volume of gas decreases. According to Boyle’s law, the product of the pressure (P) multiplied by the volume (V) remains constant if there is no change in the temperature or in the number of particles inside the container. This law is written: pV = constant
Boyle’s law says that the pressure doubles when a gas is compressed to half its volume at constant temperature.
Boyle’s law was first published by the Irish chemist Robert Boyle in 1662. But other chemists had discovered the law earlier. In 1660 and 1661, Richard Towneley and Henry Power of England experimented with air below atmospheric pressure. They found that the product of the air’s pressure and volume remained constant. At about the same time, Robert Hooke of England experimented with air above atmospheric pressure. Hooke’s findings agreed with those of Towneley and Power. Additional experiments by Boyle confirmed all these findings. In 1679, Edme Mariotte of France published the results of his own experiments with gases. Mariotte’s writings became well known in Europe. Thus, the law known today as Boyle’s law in North America and Great Britain is called Mariotte’s law in continental Europe.
Charles’s law
states that a gas expands by the same fraction of its original volume with each degree that its temperature rises. According to Charles’s law, the ratio between the volume (V) of a gas and its temperature (T) remains constant if the pressure does not change. The law is written: 
In this equation, T is the absolute temperature of the gas. It is usually measured in kelvins (Celsius degrees plus 273.15). Kelvin is abbreviated K. For example, when a gas is heated from 300 K (room temperature) to 600 K, its absolute temperature doubles. Doubling the temperature doubles the gas’s volume if the pressure does not change. See Absolute zero .
Charles’s law was discovered in 1787 by the French chemist Jacques Alexandre Cesar Charles. He found that carbon dioxide, hydrogen, oxygen, and nitrogen all expand at constant rates as their temperatures rise. Charles did not publish his findings, but explained his experiments to the French chemist Joseph Gay-Lussac. Gay-Lussac performed similar experiments and published his results in 1802. As a result, Charles’s law is sometimes called Gay-Lussac’s law.
Avogadro’s law
was first proposed in 1811 by the Italian scientist and philosopher Amedeo Avogadro. It says that equal volumes of different gases all contain the same number of particles if they all have the same pressure and temperature. It was later discovered that a volume of 22.4 liters of gas at 0 °C and atmospheric pressure contains about 602,000,000,000,000,000,000,000 (602 billion trillion) particles. This number is usually written 6.02 X 10 to the power of 23 and is called the Avogadro constant. This number of particles of any substance is called one mole of the substance. See Mole .
The universal gas law combines Boyle’s law, Charles’s law, and Avogadro’s law into a single statement. This law is written: pV = nRT
In this equation, P represents the pressure of the gas, V represents its volume, n represents the number of moles of gas, and T represents its absolute temperature. R is a constant called the universal gas constant. It has a value of 8.314 joules per kelvin per mole. According to the universal gas law, the pressure of a gas can be doubled in three ways: (1) the gas can be squeezed into one-half its original volume, (2) twice as much gas can be forced into the original volume, or (3) the absolute temperature can be doubled.
History.
During the early 1600’s, scientists began realizing that some matter can exist in a form that is similar to air. The word gas was first used to describe this form in the mid-1600’s in the writings of the Belgian chemist and physician Jan Baptista van Helmont. He invented the word gas by altering the Greek word chaos, meaning space. In this way, the word describes the ability of a gas to fill any amount of space.
Many gases were discovered and studied during the 1600’s and 1700’s. These gases include hydrogen, oxygen, and nitrogen.
The first successful attempts to liquefy many gases began in 1823 when the English scientist Michael Faraday liquefied chlorine. After heating chlorine hydrate in a sealed glass tube, Faraday noticed an oily-looking liquid inside the tube. When he tried to file the end off the tube to examine this liquid, the tube exploded. Faraday repeated the experiment, and concluded that the liquid was chlorine. The chlorine had been freed from the chlorine hydrate during heating, and had condensed under pressure inside the tube. The next day, Faraday liquefied hydrogen chloride in a similar tube. But when he tried to liquefy carbon dioxide by this method, the gas burst the tube without liquefying. Faraday later liquefied carbon dioxide and many other gases by cooling and compressing them. Today, all gases have been solidified as well as liquefied.
