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Behaviour of Gases

Our goal in this chapter is to explain the macroscopic properties of a gas - such as its pressure and its temperature in terms of the behavior of the molecules that make it up. But there is an immediate problem: which gas? Should it be hydrogen or oxygen, or methane, or perhaps uranium hexafluoride? They are all different. However, experimenters have found that if we confine 1 mole samples of various gases in boxes of identical volume and hold the gases at the same temperature, then their measured pressures are nearly though not exactly - the same. If we repeat the measurement as lower gas densities, then these small differences in the measured pressures tend to disappear. Further experiments show that, at low enough densities, all real gases tend to obey the relation
pV = nRT (ideal gas law) -----(i)

in which p is the absolute (not gauge) pressure, n is the number of moles of gas present, and R, the gas constant, has the same value for all gases, namely,
R = 8.31 J/mol. K

The temperature T must be expressed in Kelvin. Equation (i) is called the ideal gas law. Provided the gas density is reasonably low, this equation holds for any type of gas, or a mixture of different types, with n being the total number of moles present.
Using this law we can deduce many properties of the ideal gas in a simple way. Although there is no such thing in nature as a truly ideal gas, all gases approach the ideal state at low enough densities, that is, under conditions in which their molecules are fare enough apart. Thus the ideal gas concept allows us to gain useful insights into the limiting behaviour of real gases.

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