Kinetic Molecular Theory of Gases

Kinetic Molecular Theory of Gases Regina Marin Gas is called the state of matter in which, under certain conditions of temperature and pressure, its molecules cross-react only weakly with each other without forming molecular bonds, taking the shape and volume of their container and tending to separate , and expand, their best for their high kinetic energy. Gases are highly compressible fluids that experience large changes in density with temperature and pressure. Molecules constituting a gas almost are not attracted by each other, so that they move in space at high speed and quite separated from each other, thus explaining the properties: The gas molecules are virtually free, so that they are able to be distributed throughout the space in which they are contained. The gravitational attraction and forces between molecules are negligible compared to the rate at which the molecules are moving. Gases completely occupy the volume of their container. Gases have no definite shape, embracing the vessels containing them. Can easily be compressed, because there are large gaps between molecules, and other. At ambient temperature and pressure gases can be elements such as hydrogen, oxygen, nitrogen, chlorine, fluorine and noble gases, compounds such as carbon dioxide or propane, or mixtures like air. For the thermal behavior of particles of matter there are four measurable quantities that are of great interest: pressure, volume, temperature and mass of the sample material (or better amount of substance, measured in moles). Any gas is considered a fluid because it has properties that allow it to behave as such. Its molecules in constant motion, colliding elastically with each other and against the walls of the vessel containing the gas, against which exert a constant pressure. If the gas is heated, the heat energy is spent on kinetic energy of the molecules, that is, the molecules move more quickly, so that the number of collisions with the walls of the vessel increases in number and energy. As a consequence the gas pressure increases, and if the container walls are not rigid, gas volume increases. A gas tends to be chemically active because its molecular surface is also large, that is, to be its particles in continuous motion colliding with each other, this makes it easier the contact between a substance and another, increasing the rate of reaction in compared to liquid or solid. To better understand the behavior of a gas, where studies are conducted with respect to the ideal gas, although it never actually exists and its properties are: A pure gaseous substance consists of molecules of the same size and mass. A gaseous mixture is formed by different molecules in size and mass. Due to the large distance between molecules and other and that move at high speed, the forces of attraction between the molecules are considered negligible. The size of the gas molecules is very small, so that the volume occupied by the molecules is negligible compared with the total volume of the container. The density of a gas is very low. The gas molecules are in constant motion at high speed, so continuously collide elastically with each other and against the walls of their container. As part of the kinetic theory, the gas pressure is explained as the macroscopic result of the forces involved by collisions of gas molecules with the walls of the container. The pressure can thus be defined with reference to the microscopic properties of the gas. The kinetic theory of gases is a physical and chemical theory that explains the behavior and macroscopic properties of gases (ideal gas law), from a statistical description of the microscopic molecular processes. The kinetic theory was developed based on studies of physical and Daniel Bernoulli in the eighteenth century, Ludwig Boltzmann and James Clerk Maxwell in the late nineteenth century. This branch describes the thermal physical properties of the gases. These systems contain huge numbers of atoms or molecules, and the only reasonable way to understand the thermal properties based on molecular mechanics, we find certain dynamical quantities of average type and relate the observed physical properties of the system with t hese properties averaged molecular dynamics . Techniques to relate the overall macroscopic behavior of material systems with the average behavior of their molecular components are statistical mechanics. The main theorems of the kinetic theory are: The number of molecules is large and the average separation between them is large compared with their dimensions. Therefore occupy an insignificant volume when compared to the volume of the container and are considered point masses. The molecules obey Newton's laws, but individually they move randomly, each with different rates, but with an average speed that does not change with time. The molecules perform elastic collisions with each other, therefore both the linear momentum is conserved as the kinetic energy of the molecules. The gas is considered pure, in other words all molecules are identical. The gas is in thermal equilibrium with the walls of the container. As part of the kinetic theory of a gas pressure is explained as the macroscopic result of the forces involved by collisions of gas molecules with the walls of the container. The pressure can thus be defined with reference to the microscopic properties of the gas. It is generally believed that there is more pressure if the particles are in the solid state, if they are in liquid state is minimal distance between them and finally if you are in the gaseous state are far apart. Indeed, for an ideal gas with N molecules , each moving mass m with a random speed average content in a cubic volume V of the gas particles impacting with the wall of the container in a manner that can be calculated in a statistical manner exchanging momentum with the walls in each shock and effecting a net force per unit area that is the pressure exerted by the gas on the solid surface. The pressure can be calculated with this formula: The equation above states that the gas pressure is directly dependent on the molecular kinetic energy. The ideal gas law allows us to ensure that the pressure is proportional to the absolute temperature. These two statements allow one of the most important statements of the kinetic theory: The average molecular energy is proportional to temperature. The proportional constant is 3/2 is the Boltzmann constant, which in turn is the ratio of the gas constant R between the Avogadro number. So in a few words, the kinetic theory is a physical theory, based on a few facts: The density of the gas is very small. Individually molecules move randomly and at different speeds, which increases or decreases while the temperature and the movement causes them from hitting each other, increasing the pressure when striking more times. The cohesive forces or intermolecular forces in gases are almost nil. If all of the molecules forming the gas are identical, is said to be a pure gas. Bibliography: http://www.chm.davidson.edu/vce/kineticmoleculartheory/basicconcepts.htm http://en.wikipedia.org/wiki/Kinetic_theory http://www.sparknotes.com/chemistry/gases/kinetic/summary.html http://hyperphysics.phy-astr.gsu.edu/hbase/kinetic/kinthe.html