Kinetic theory of matter[]
“But what causes this pressure?” I’m pretending to hear you ask. Well, matter is made up of tiny molecules. Just like my father after he gets drunk on fermented grain juice on Chinese New Year, they like to move around randomly and bump into things, only with significantly less vomiting. This action of bumping creates a force, the force we feel as pressure. Because these molecules are numerous in number and moving continuously around, there are approximately equal numbers moving in all directions, colliding with the walls of the container and with each other. We are assuming, of course, that the speed of a molecule before and after a collision is equal.
Pressure is affected by temperature[]
The magnitude of the pressure depends on how hard and how frequent molecules hit the wall. When temperature is increased and the molecules gain internal energy, the average kinetic energy of the molecules also increases, since temperature is directly proportional to the average kinetic energy. This causes an increase in pressure as the molecules hit the walls in greater force and frequency provided the volume and mass are constant. Hence, for a fixed mass of gas, pressure increases with increasing temperature.
Pressure is affected by volume[]
For a fixed mass of gas, decreasing the volume of the container would cause the number of molecules per unit area to increase. This causes to number of molecules hitting the walls per second to increase accordingly. Therefore, for a fixed mass of gas at constant temperature, gas pressure increases as volume decreases.
Volume is affected by temperature[]
When the temperature of a gas is increased, the molecules move around faster and hit the container walls more frequently and forcefully. This causes the internal pressure to increase. If the container is not rigid, the internal pressure will exceed the external pressure and force the container walls outwards so as to balance the internal and external pressure. Thus, the volume of a gas with fixed mass and pressure increases with temperature.
Temperature is affected by compression[]
When molecules hit a stationary wall, it rebounds at the same speed. However, if the molecules rebound off the walls of a moving piston that is compressing the gas, the molecule will rebound faster as the piston is moving in the same direction as the rebound. Thus, the average speed of the gas molecules (and hence the temperature) will increase as a fixed mass of gas is compressed.
The opposite is true as well. If the piston in moving in the opposite way, the temperature will decrease as the average speed of the molecules decrease due to slower rebounds.
Kinetic theory and change of state[]
The concept of everything being made up of tiny, randomly moving molecules can help explain the process of matter changing states. Normally, molecules are held together by strong intermolecular forces that keep them together in the solid and liquid states. The intermolecular forces between gas molecules are so small they’re essentially negligible.
Boiling/Melting[]
Since temperature causes molecules to gain internal energy and move around faster, it follows that eventually, they will gain enough energy to break these bonds and change states. Hence, you have melting and evaporation/boiling.
Evaporation[]
The random movement of molecules help explain why evaporation happens at all temperatures. Because of these random collisions between molecules, some molecules at the surface may obtain greater velocities. Should these molecules happen to possess enough kinetic energy in the upward direction, they can break free of their intermolecular bonds and exist independently as free and emancipated gaseous molecules.
Because only the fastest molecules evaporate, the molecules that are left behind tend to be amongst the slower moving ones. Because temperature is, once again, directly proportional to the average kinetic energy of the liquid molecules, evaporation causes cooling.
If there is wind above the liquid surface, evaporation takes place even faster. The liquid molecules that evaporate are immediately blown away and more liquid molecules can fill up this now empty space. Also, fewer molecules are able to return to the liquid, accelerating the rate of evaporation.
Because evaporation takes place only at the surface of a liquid, it follows that a greater surface area would allow for more liquid molecules to escape at any one time. Thus, evaporation would be faster with a larger surface area.
Proof of the kinetic theory[]
Hmm. I get it. You want proof that everything around us is made up of tiny particles randomly bumping around. Well, here’s the evidence:
Brownian motion[]
This is essentially the random and erratic motion of small particles (e.g. smoke particles) within a fluid. This is caused by the random bombardment of these molecules of the fluid. It is independent of container movement and varies directly with temperature and inversely with particle size. This occurrence can be observed by viewing smoke particles in a lighted box under a microscope. Specks of light will be seen to move around randomly.
Diffusion[]
Because these molecules move around, we have diffusion, the process in which molecules move from a region of high concentration to a region of lower concentration. This can be seen in liquids (drop a drop of food dye in a jug of water and see what happens) and gases (e.g. diffusion of brown coloured bromine gas). It happens in solids too but it takes so long, there’s no point bothering.