Dalton's Law of Partial Pressures
Dalton's law of partial pressures states that the total pressure of a mixture of gases equals the sum of the pressures that each gas would exert if it was present alone. The pressure exerted by one gas in a mixture of gases is the partial pressure of that individual gas. This law assumes that the gases do not react with each other, but that each gas is a separate component of the system.
Dalton's law of partial pressures can be mathematically expressed by the following. If Pt is the total pressure of a mixture of gases and P1, P2, etc. are the pressures that each gaseous component would exert by itself, then the total pressure is given by Pt = P1 + P2 + etc. Dalton's law of partial pressures tells us that each gas behaves independently of the other gases in the mixture, i.e., each exerts its own pressure (the partial pressure) providing the gases do not react.
The total pressure of a mixture of gases is a function of the number of moles of gas present, irrespective of what type of gas is considered. One mole of any gas has the same volume as one mole of any other gas, providing the temperature and pressure are the same. This is the basic theory behind Avogadro's law. One mole of gas under the same conditions as one mole of a second gas will has the same number of particles as the second gas. The same volume (number of moles) of two gases exerts the same pressure in a container because each gas has the same number of particles and hence the same number of collisions between the particles and the container walls. Similarly one mole of gas in a fixed volume exerts the same pressure as one mole of any other gas in an identical volume. If both gases are in the same container then the total pressure is twice the pressure exerted by either of the gases, providing the same number of moles of each gas was present. The pressure exerted by any one of these gases is the partial pressure.
The pressure of a gas is a measure of how frequently molecules of the gas hit the sides of the container the gas is held in. Gases have molecules that are widely separated due to their high energy levels. As a result, the molecules within a volume of gas have plenty of space between them and they move in an entirely random manner. Consequently if another gas is introduced into the same volume the molecules of this second gas are equally free to move in a random manner. There are no space limitations and the number of collisions between gas molecules is very small compared to the number of collisions between the gas molecules and the much larger walls of the container (which gives us the pressure). The molecules of each gas strike the walls of the container independently of each other. The total pressure exerted is the pressure exerted by all molecules hitting the container walls, irrespective of which gas they belonged to. Each gas behaves as if it is the only gas present in the system, and the partial pressure that each gas exerts is the same as if no other gases were present. Because the pressure exerted is a function of the speed of the molecules, kinetic theory states that the pressure will increase if the energy of the molecules increases. One way of increasing the kinetic energy of a gas is to increase its temperature. If both gases are heated, the overall pressure increases and the partial pressure of each component of the system also increases.
Atmospheric pressure is the sum of the partial pressures of all of the gases present in the atmosphere. The atmosphere is a mixture of several main gases and many other gases present in much smaller quantities, with a large amount of local variation. The gas with the greatest concentration in the atmosphere is nitrogen with 78%. Next comes oxygen with a concentration of 21%.Argon, carbon dioxide, neon, krypton, xenon, and radon are present in the atmosphere at lower concentrations. Each of these gases exerts a pressure as if the other gases were not present, i.e., its partial pressure. The total pressure of the air is the sum of the partial pressures of all of the gases present. Atmospheric pressure varies depending on the height above sea level. When standard temperature and pressure are discussed the pressure used is the pressure at sea level--this is known as one atmosphere.
John Dalton was an English schoolteacher who lived from 1766 to 1844. He was interested in chemistry and meteorology. His first contribution to chemistry was an early version of atomic theory. It was in 1801 while studying the properties of air that Dalton formulated his law of partial pressures.
Dalton's law of partial pressures works for low, normal, and moderate pressures, but it starts to fail at very high pressures. This failure at high pressure is due to the fact that the molecules become more densely packed. Because the molecules of gas are close together, they are no longer able to move freely in the system and collisions between molecules become more frequent. The other factor adding to the failure of the law of partial pressures at high pressure is the presence of weak, intermolecular forces. When the molecules are free to move widely separated from each other the effect of these attractive forces is negligible. As the molecules are pushed closer together, the effects of small forces, such as van der Waals forces, become apparent. This leads to a clumping of the molecules present. If the molecules are clumped together in an aggregate, they move more slowly because they are larger. The immediate effect of this is a slight reduction in pressure because the number of collisions of gas molecules with the walls of the container is reduced.
Dalton's law of partial pressures states that in a mixture of gases each component gas exerts a pressure that is totally independent of the other gases present in the mixture. The total pressure of the mixture is the sum of all of the pressures of the gases present. Dalton's law of partial pressures only works for gases that do not react with each other.
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