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Kinetic Molecular Theory
The kinetic-molecular theory is a theory that explains the states of matter and is based on the idea that matter is composed of tiny particles that are always in motion. The theory helps explain observable properties and behaviors of solids, liquids, and gases. However, the theory is most easily understood as it applies to gases, and it is with gases that we will begin our detailed study. The theory applies specifically to a model of a gas called an ideal gas. An ideal gas is an imaginary gas whose behavior perfectly fits all the assumptions of the kinetic-molecular theory. In reality, gases are not ideal, but they are very close to being so under most everyday conditions.
The kinetic-molecular theory, as it applies to gases, has five basic assumptions:
- Gases consist of very large numbers of tiny spherical particles that are far apart from one another compared to their size. The particles of a gas may be either atoms or molecules. The distance between the particles of a gas is much greater than the distances between the particles of a liquid or a solid. Most of the volume of a gas is composed of the empty space between the particles. In fact, the volume of the particles themselves is considered to be insignificant compared to the volume of the empty space.
- Gas particles are in constant rapid motion in random directions. The fast motion of gas particles gives them a relatively large amount of kinetic energy. Kinetic energy is the energy that an object has because it is moving. The particles of a gas move in a straight line until they collide with another particle or with one of the walls of their container.
- Collisions between gas particles and between gas particles and the container walls are elastic collisions. An elastic collision is one in which there is no overall loss of kinetic energy. This means that the gas particles bounce off of one another and the container walls. Kinetic energy may be transferred from one particle to another during an elastic collision, but there is no change in the total energy of the colliding particles.
- There are no forces of attraction or repulsion between gas particles. It is assumed that the particles of an ideal gas have no such attractive forces. The motion of each particle is completely independent of the motion of all other particles.
- The average kinetic energy of gas particles is dependent upon the temperature of the gas. As the temperature of a gas is increased, its particles begin to move faster, resulting in an increase in their kinetic energies. If the temperature is decreased, the gas particles move slower. Not all particles in a given sample have the same speed, so the sample will contain particles with a range of different kinetic energies. However, the average kinetic energy of the particles in a sample is related to its temperature.
The postulates are summarized on the front page of the Gas Laws section, but are repeated below:
- Gas particles are very small - so small that we can count their volume as zero. A gas is mostly just empty space
- Gas particles are in contant random straight-line motion, obeying Newton's laws, and only changing direction when they impact other particles or the walls of their containers.
- Collisions between gas particles are perfectly elastic, meaning that no kinetic energy is lost in those collisions.
- Gas paricles do not interact with each other except when they collide. There are no attractive or repulsive forces between them.
- The temperature of a gas is directly related to the average kinetic energy of the gas particles. As the particles speed up, the average kinetic energy increases as does the temperature.