Basically what the kinetic theory states- or what I think it states- is that solids, liquids and gases are all made up of particles which are generally hyperactive like me. In solids, they vibrate around fixed positions (unless the temperature is 0 degrees Kelvin in which case they're completely still). In liquids, they move a little more freely but are still loosely tethered together by intermolecular forces. As gases, they are much freer and move in random straight-line motion.
The most important part of the kinetic theory is the kinetic theory of gases. Here it is:
- Particles move in random straight-line motion until they collide with each other or another solid object (like the inside of a glass jar).
- Collisions are perfectly elastic- no change of energy results from a collision.
- The size of the particles is negligible.
- The bonding between particles is negligible.
- Average kinetic energy increases as temperature increases, according to that kinetic energy formula KE = (1/2)mv^2.
One thing to take in mind, however, is that the kinetic theory is not 100% accurate. There are bonds between particles. And particles don't have 0 size, otherwise they wouldn't exist. But that doesn't stop the kinetic theory from being good at explaining stuff.
For example, you can explain why gases spread through the kinetic theory idea of random straight-line motion. You can then explain why you can compress gases through the idea of the size of particles being negligible.
Now for the more equationy bits. It's time to talk about how the kinetic theory can explain the relationships between temperature, pressure and volume.
Volume and pressure are inversely proportional. That is to say, when volume decreases, pressure increases (provided that temperature remains constant). In a smaller space, the gas particles are squished closer together, so they collide more often with each other and with the walls of the container or whatever it is that they're confined in. In a larger space, the gas particles have more room to move and thus collisions are less likely.
Temperature and pressure are proportional. As temperature increases, so does pressure, provided that volume remains constant. (This does not work the other way around, i.e. increased pressure obviously won't increase the temperature.) This is because an increased temperature causes the particles to move with a greater velocity, making them more likely to collide with each other and with the walls of the container. When they collide, they will collide with a greater velocity. Opposite effects are achieved when the temperature is decreased.
Temperature and volume are proportional. As I said before, increasing the temperature also increases the pressure. However, if you want to keep the pressure constant you'd also have to increase the size of the container (i.e. the volume of the gas). The opposite happens when the temperature decreases.
One last thing. If you're ever given a question involving temperature and pressure or volume, convert the temperature to Kelvin first. For example if you get a question about what would happen to the pressure if the temperature increased from 30 degrees Celsius to 60 degrees Celsius, the answer would not be "the pressure doubles." When you convert to Kelvin you would find that the temperature isn't doubled, and therefore the pressure isn't doubled either.
I think I've covered the main bits of the kinetic theory. It's not a topic I'm comfortable with so I'll have to go back and revise this later...
Now for the more equationy bits. It's time to talk about how the kinetic theory can explain the relationships between temperature, pressure and volume.
Volume and pressure are inversely proportional. That is to say, when volume decreases, pressure increases (provided that temperature remains constant). In a smaller space, the gas particles are squished closer together, so they collide more often with each other and with the walls of the container or whatever it is that they're confined in. In a larger space, the gas particles have more room to move and thus collisions are less likely.
Temperature and pressure are proportional. As temperature increases, so does pressure, provided that volume remains constant. (This does not work the other way around, i.e. increased pressure obviously won't increase the temperature.) This is because an increased temperature causes the particles to move with a greater velocity, making them more likely to collide with each other and with the walls of the container. When they collide, they will collide with a greater velocity. Opposite effects are achieved when the temperature is decreased.
Temperature and volume are proportional. As I said before, increasing the temperature also increases the pressure. However, if you want to keep the pressure constant you'd also have to increase the size of the container (i.e. the volume of the gas). The opposite happens when the temperature decreases.
One last thing. If you're ever given a question involving temperature and pressure or volume, convert the temperature to Kelvin first. For example if you get a question about what would happen to the pressure if the temperature increased from 30 degrees Celsius to 60 degrees Celsius, the answer would not be "the pressure doubles." When you convert to Kelvin you would find that the temperature isn't doubled, and therefore the pressure isn't doubled either.
I think I've covered the main bits of the kinetic theory. It's not a topic I'm comfortable with so I'll have to go back and revise this later...
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