Kinetic theory explains the properties of gases in terms of molecular composition and activity. It assumes that gases are made up of small particles that are constantly moving, and that pressure is the result of collisions between molecules. The theory defines pressure as the force exerted by gas molecules when they collide with the container wall. Kinetic energy varies in direct proportion to the absolute temperature of a gas.
Kinetic theory is a scientific theory of the nature of gases. The theory goes by many names, including kinetic gas theory, kinetic-molecular theory, collision theory, and kinetic-molecular theory of gases. It explains the observable and measurable, also called macroscopic, properties of gases in terms of molecular composition and activity. While Isaac Newton theorized that the pressure of a gas is due to static repulsion between molecules, kinetic theory holds that the pressure is the result of collisions between molecules.
Kinetic theory makes a number of assumptions about gases. First, a gas is made up of very small particles, each with a non-zero mass, that are constantly moving in a random way. The number of molecules in a gas sample must be large enough for statistical comparison.
The kinetic theory assumes that gas molecules are perfectly spherical and elastic, and that their collisions with the walls of their container are also elastic, meaning that they do not result in any change in velocity. The total volume of the gas molecules is negligible compared to the total volume of their container, which means that there is ample space between the molecules. Furthermore, the time during the collision of a gas molecule with the container wall is negligible in relation to the time between collisions with other molecules. The theory is also based on the assumption that any relativistic or quantum mechanical effects are negligible, and that any effects of gas particles on each other are negligible, except for the force exerted by collisions. Temperature is the only factor affecting the average kinetic energy, or energy due to motion, of gas particles.
These assumptions must be maintained for the equations of kinetic theory to work. A gas that satisfies all of these assumptions is a simplified theoretical entity known as an ideal gas. Real gases usually behave similarly enough to ideal gases for the rate equations to be useful, but the model is not perfectly accurate.
Kinetic theory defines pressure as the force exerted by gas molecules when they collide with the container wall. Pressure is calculated as force times area, or P = F/A. The force is the product of the number of gas molecules, N, the mass of each molecule, m, and the square of their average velocity, v2rms, all divided by three times the length of the container, 3l. Therefore, we have the following equation for the force: F = Nmv2rms/3l. The abbreviation, rms, stands for root-mean-square, an average of the velocity of all particles.
The equation for pressure is P = Nmv2rms/3Al. Since area times length equals volume, V, this equation can be simplified as P = Nmv2rms/3V. The product of pressure and volume, PV, equals two-thirds of the total kinetic energy, or K, allowing the derivation of macroscopic from microscopic properties.
An important part of kinetic theory is that kinetic energy varies in direct proportion to the absolute temperature of a gas. Kinetic energy is equal to the product of absolute temperature, T, and Boltzman’s constant, kB, multiplied by 3/2; K = 3TkB/2. Therefore, whenever the temperature is increased, the kinetic energy is increased and no other factors have an effect on the kinetic energy.
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