The three basic states of matter we encounter every day are solid, liquid, and gas (or vapor). These three states of matter can be distinguished by the arrangement of their constituent atoms; gas particles are well separated with no regular arrangement, while particles of solids and liquids are closer together. Because of the differences in particle arrangement, each state of matter displays different properties.
By the late 19th century, scientists had begun to accept the atomic theory of matter and related it to individual molecules. The kinetic molecular theory of gases comes from observations that scientists made about gases to explain their macroscopic properties. The theory outlines the basic behavior of an ideal gas:
- The volume occupied by the molecules of a gas is negligible compared to the volume of the gas itself.
- The molecules of an ideal gas exert no attractive forces on each other or their surroundings.
- The molecules are in a constant state of random motion and move in straight lines until they collide with another body.
- The collisions exhibited by a molecule are elastic; when two molecules collide, they change directions and kinetic energy, but the total kinetic energy is conserved.
- The average kinetic energy of the gas molecules is directly proportional to absolute temperature, which implies that all molecular motion would cease if the temperature were reduced to absolute zero.
Charles' Law states that at constant pressure, the volume of a gas increases or decreases by the same factor as its temperature. It can be written as:
According to kinetic molecular theory, an increase in temperature will increase the average kinetic energy of the molecules. As the particles move faster, they will likely hit the edge of the container more often. If the reaction is kept at constant pressure, they must stay farther apart, and the increase in particle collision with the surface of the container will be compensated for by an increase in volume.
At a given temperature, the pressure of a container is determined by the number of times gas molecules strike the container walls. If the gas is compressed to a smaller volume, then the same number of molecules will strike against a smaller surface area; therefore, the number of collisions against the container will increase, and therefore, the pressure will increase as well. Increasing the kinetic energy of the particles will increase the pressure of the gas.