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Volume (V), temperature (T), moles (n), and pressure (P) are four experimental parameters of gases that are related to each other by gas laws. Laws are generalized observations of experimental evidence, not explanations of why. The gas laws apply to all gases, regardless of identity.

Of these four parameters pressure is the least familiar. Pressure is the force per unit area. With gases the force comes from the gas molecules hitting the side of the container. Air pressure is measured with a barometer. Barometers compare the force being applied by air to force being applied by a column of liquid. The higher the column, the more force. Therefore one unit of pressure is mm Hg, which refers to the height of a column of mercury. Another name for the unit mm Hg is a torr, in honor of the barometer's inventor, Evangelista Torricelli. The average atmospheric pressure at sea level is 760 torr. This leads to another unit of pressure, atmospheres (atm), where 1 atm is exactly equal to 760 torr. Pascals (Pa) are the SI unit of pressure that is based on the definition (force/area) rather than an experimental measurement. 1 Pa = 1 N/m². A related unit is a bar, where 1 bar = 100 Pa. These two types of units are related by 101,325 Pa = 1 atm.

Changing one of these parameters can affect the others. If temperature and amount of gas are kept constant and pressure is increased, volume will decrease. This is Boyle's law, that pressure and volume are inversely proportional. Charles's law says that volume is proportional to temperature, when moles of gas and pressure are constant. These two laws can be combined into the combined gas law (Equation 8.8). Avogadro's law says that moles are proportional to volume with constant pressure and temperature. The conditions used for comparison of gases are called standard temperature and pressure (STP). Standard temperature is 0°C (273.25 K) and standard pressure is 1 atm (760 torr). The volume of 1 mole of gas at STP is called standard molar volume and has a value of 22.4 L.

Because pressure, volume, temperature, and moles are the only variables, if three of the variables are known, the other can be determined. The relationship between these variables is called the ideal gas law:

PV = nRT       (Equation 8.9)

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In this equation, R is the gas constant. Its value depends on the units used in the other variables (Table 8.1). By rearranging this equation, these experimental parameters can be related to mass, density (Equation 8.11), and molar mass (Equation 8.10).

Since the identity of the gas is irrelevant to the gas laws, the laws work as well for mixtures of gases as a single gas. Thus the total pressure is proportional to the total number of moles. Since all molecules of gas in a mixture must have the same temperature and volume, the gases are differentiated by the pressure. The pressure of each gas is called its partial pressure. The sum of the partial pressures is equal to the total pressure (Dalton's law, Equation 8.12). Another way to state the relationship is that the partial pressure is equal to the mole fraction (X = moles gas/total moles) times the total pressure (Equation 8.13).

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A gas formed by molecules escaping from a liquid is called a vapor. In a closed container, the molecules escaping the liquid and being captured by it reach a constant rate and the pressure exerted by the vapor depends only on the temperature. This pressure is called the vapor pressure. The vapor pressure of water is particularly important, since many gases are collected over water. When gases are collected over water, the gas collected also contains water vapor. However, Dalton's law can be used to differentiate the pressure of the water from the pressure of the gas. The vapor pressure of water at the appropriate temperature can be obtained from a table (Table 8.2) and subtracted from the total pressure, giving the partial pressure of the collected gas. The ideal gas law is used to relate partial pressure to moles.

The extent to which a gas dissolves in a liquid is proportional to its pressure (Henry's law, Equation 8.14). It also depends on the identity of the gas. The proportionality constants (Henry's constants) are listed in Table 8.3.

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Temperature is related to the energy or speed of molecules. Since there is no such thing as a negative speed, the Kelvin temperature scale is always used in gas law calculations. Higher temperatures give the molecules more energy. However, as it take more energy to move a larger object than a smaller one, the molar mass of the molecules will affect their speed. The rate at which gas molecules move can be determined by the rate of effusion (rate gas escapes through a pinhole) or diffusion (rate of spreading). Graham's law states that the rate of effusion (or diffusion) is inversely proportional to the square root of the molar mass (Equation 8.18).

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All the laws simply state the observed relationships between pressure, volume, temperature, moles, and so on. The theory that explains these relationships is the kinetic molecular theory. The kinetic molecular theory describes gases as molecules in independent, constant motion. The molecules are neither attracted to each other nor repelled from each other and do not take up a significant amount of the volume. In addition, the molecules have elastic collisions, meaning that energy is not lost in a collision with either the container or another gas molecule. Under these conditions, the gas laws work perfectly and the gas is called an ideal gas.

However, molecules of real gases are often attracted to each other. The attraction is greater at low temperatures (when the molecular speeds are slow) and high pressure (when there are lots of collisions). The molecules must take up some space as well. That space would be significant at low volumes. These gases are called real gases and the effects of molecular attraction and space taken up by the molecules are corrected for with the van der Waals equation for real gases (Equation 8.19).