The Kinetic Molecular Theory describes the behavior of ideal gases, which are gases that obey the ideal gas law, as follows:
- The particles of a gas are in constant, rapid, random motion
- Collisions between gas particles are elastic (there's no loss of energy).
- The volume of the particles themselves are insignificantly small compared to the space they occupy in moving around
- There are no forces of attraction of repulsion between gas molecules.
- The average kinetic energy of the particles in a gas sample is proportional to the Kelvin temperature.
The kinetic molecular theory explains the mathematical relationships between temperature, pressure, volume and the number of particles of a gas which is summed up in the ideal gas law PV = nRT. For example, the pressure and volume relationship is explained by the fact that if the volume of a gas is decreased the surface area of the container decreases so the particles will have more collisions per unit area of the container. This results in higher pressure. Most gases behave like ideal gases at temperatures and pressures close to STP (0 degrees C and 1.00 atm pressure).
The effect of intermolecular forces on gas behavior comes into play at lower temperatures and/or higher pressures. When gas molecules are closer together and have less speed intermolecular attractions become more important. Condensation from a gas to a liquid occurs at the temperature and pressure at which the gas molecules no longer have enough energy to overcome attractions to other molecules: or, in converse, condensation to liquid occurs when intermolecular attractions become stronger.
Molecules in a liquid are much closer together than in a gas, but they still move around and change position. Intermolecular attractions are stronger in the solid state than in the liquid state. The attraction between particles in a solid keep them locked into position and their only movement is vibrational. Neither liquids nor solids obey the Kinetic Molecular Theory.
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