The Kinetic Molecular Theory is a model used to explain the behavior of gases. It states that gases are made up of particles (molecules or atoms) that are in constant motion and that the temperature of a gas is directly related to the average kinetic energy of its particles.
The main assumptions of the Kinetic Molecular Theory are that gas particles are in constant, random motion; that there is no attraction or repulsion between the particles; and that the volume of the particles is negligible compared to the volume of the gas.
The Kinetic Molecular Theory explains gas pressure as the result of gas particles colliding with each other and with the walls of their container. The more collisions that occur, the higher the pressure of the gas.
The Kinetic Molecular Theory states that the temperature of a gas is directly proportional to the average kinetic energy of its particles. As the temperature increases, the particles move faster and have a higher average kinetic energy.
The Kinetic Molecular Theory explains that the differences in behavior between gases, liquids, and solids are due to the strength of the intermolecular forces between their particles. In gases, the particles are far apart and have weak intermolecular forces, allowing them to move freely. In liquids, the particles are closer together and have slightly stronger intermolecular forces, leading to more restricted movement. In solids, the particles are tightly packed and have strong intermolecular forces, resulting in a fixed shape and volume.