Gases, Behavior and Properties Of
Gases are the most diffuse form of matter. In a gas the molecules (the smallest particles of which the gas is made--for some gases these equate to atoms) are highly energetic and they can be widely separated from each other. The behavior of gases and their properties derive from these facts.
Sometimes the term vapor is used to describe a gas. Strictly speaking a gas is a substance at a temperature above its boiling point. A vapor is the gaseous phase of a substance that, under ordinary conditions, exists as a liquid or solid.
One of the more obvious characteristics of gases is their ability to expand and fill any volume they are placed in. This contrasts with the behavior exhibited by solids (fixed shape and volume) and liquids (fixed volume but indeterminate shape). Gases have an indeterminate volume and shape. This is due to the freedom of movement exhibited by the molecules comprising the gas. A gas can be compressed to a smaller volume or it can expand to fill a larger volume. Gases will form homogenous mixtures with each other. Providing no reaction occurs between the gases, the gases will disperse throughout the volume of the container they are kept in. The pressure inside the container is the sum of the pressures that each gas would exert if it were present alone. This is summarized in Dalton's law of partial pressures.
Gases are mostly space. For example, in the air that we breathe the actual molecules only account for 0.1% of the volume--the rest is empty space. The practical effect of this is that a molecule in a gas behaves as if it is in isolation, i.e., it does not react with the other molecules in the gas. This is why all gases have similar properties, irrespective of the type of gas.
Many substances are referred to as gases, since this is the state in which they are normally encountered. In reality many substances may exist only in the gaseous phase over a small range of conditions. Consequently when we refer to something as being a gas and we do not specify the conditions it is assumed that the condition is standard temperature and pressure (STP). Standard temperature and pressure is defined as a temperature of 32°F (0°C) and a pressure of 1 atmosphere.
The behavior of gases and their properties is explained by the kinetic molecular theory, which, in its current form, was first published by Rudolf Clausius in 1857. The kinetic molecular theory of gases states that gases are comprised of a large number of molecules that are in constant, random motion. The volume of all of the molecules of a gas is very small compared to the volume of the gas. Intermolecular forces are largely irrelevant within a gas due to the large intermolecular distances. While energy can be transferred between molecules during collisions, the average kinetic energy of molecules does not change (providing the temperature remains constant). The average kinetic temperature of the molecules of a gas is proportional to the absolute temperature (the temperature in Kelvin) and at any given temperature the molecules of all gases have the same kinetic energy.
The kinetic molecular theory allows us to understand the behavior of gases. The pressure exerted by a gas is a measure of how frequently (and how hard) the molecules of a gas strike the walls of the container in which it is being stored. The absolute temperature of a gas is a measure of the average kinetic energy of the molecules of the gas. Different gases at the same temperature have the same average kinetic energy. If the temperature of a gas is altered then the kinetic energy of the molecules is altered accordingly. If the absolute temperature of a gas is halved then the average kinetic energy of the molecules is also halved. The average kinetic energy is the same between all moleculesof a gas, but some individual molecules in a gas are moving faster and some slower than the other molecules. The distribution of energies within a body of gas follows a normal distribution, if the number of molecules is plotted against the energy of the molecules.
Effusion (the movement of gas molecules through a pinhole in a container) and diffusion (the uniform spread of gas molecules throughout a mixture of gases) are both related to the energy levels of the gas. The higher the energy levels the more rapidly the molecules are moving and the more rapidly the processes of effusion and diffusion take place. This is due to the fact that the molecules are more likely to hit the pinhole the more rapidly they are moving, and the more rapidly they are moving the quicker they will spread into an unoccupied space. The law governing the rates of effusion of two gases is called Graham's law. Both effusion and diffusion are faster for lighter gases (diffusion will also occur with solids but the rate of diffusion is infinitesimal when compared to gases). Related to the rate of diffusion is the mean free path of a molecule. This is the distance traveled between collisions. The greater the distance traveled between collisions the more rapid will be diffusion, since collisions make gas molecules move from one place to another in a staggered manner. The length of the mean free path is governed by the density of the gas--the more the molecules are packed together the greater is the likelihood of a collision between molecules. At sea level the mean free path distance for a gas molecule in air is about 100 nm. At an altitude of 62 mi (100 k m) the air density is much lower and the mean free path is some million times longer than at sea level. The mean free path in air is about 6 in (16 cm) at this altitude.
The universal gas law can be used to describe the behavior of a gas. Derived from Boyle's law, Charles's law, and the constant volume law, it can be written as pV/T = a constant, where p is the pressure, V is the volume, and T is the absolute temperature. Another way of writing this law is pV/T = pV/T where the subscript 1 represents the start conditions and the subscript 2 represents the end conditions. A third way of writing the universal gas law is as pV = nRT where n represents the number of moles of gas being dealt with and R represents the universal gas constant. The universal gas constant is 0.08204 liter atmosphere per degree, or 8.314 Jmol1 K-1 .
Avogadro's hypothesis states that at equal pressure and temperature, equal volumes of gases contain the same number of molecules. This hypothesis led to Avogadro's law, which states that the volume of a gas maintained at constant temperature and pressure is directly proportional to the number of moles of the gas. So if the temperature and pressure of the gas remain constant and the number of moles of gas present is doubled then the volume will also double.
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