The kinetic molecular theory (KMT) describes the behavior of ideal gases at the particle level. The five main postulates of the KMT are as follows: (1) the particles in a gas are in constant, random motion, (2) the combined volume of the particles is negligible, (3) the particles exert no forces on one another, (4) any collisions between the particles are completely elastic, and (5) the average kinetic energy of the particles is proportional to the temperature in kelvins. Created by Sal Khan.
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- I don't really get about the velocity of the particles being preserved when they bounce of the walls.
Could someone explain?(0 votes)
- When particles collide with each other, or with the sides of the container, they realistically transfer some of their kinetic energy. Kinetic energy again being the energy associated with motion and is directly proportional to a particle’s velocity. This transferred kinetic energy is transformed into other forms of energy like potential energy or heat. However in the kinetic theory of gases we assume that these collisions are what is known as elastic where the collisions do not result in the transfer of any kinetic energy and thus the particles would maintain their velocity. Now like I said realistically there is at least some kinetic energy transferred when gas particles collide so rarely are these collisions ever perfectly elastic. However when averaged over the many trillions of gas particles, these collisions can be assumed to be essentially elastic which makes the math for calculations simpler.
All these assumptions describe gases as ideal gases which obey simple equations like the ideal gas law. And while no gas is ever completely ideal, these assumptions allow us to make roughly accurate predictions on the behaviors of gases.
Hope that helps.(14 votes)
- How do the particles exert no force on one another? Newton's law of universal gravitational force proves the opposite. I mean, they must apply force on each other, even if it is tiny or negligible.(4 votes)
- What are particles made of?(1 vote)
- Which particles are you referring to? Many things can be referred to as particles, so you'll need to be more specific.(5 votes)
- Can't we measure gasses in 'microscopic level' or what if we try to do that?(3 votes)
- No, because gas is not just one atom or molecule it is multiple. For example Oxygen can't stand alone it is O2 which is more than one atom or molecule. By studying only one atom or particle we can't measure the behavior of the gas as a whole.(3 votes)
- So, according to Axiom 3, the London Dispersion Forces shouldn't be there. But that's not true. So, is this right ? If not, why so ?(1 vote)
- These axioms essentially assume that the gas behaves ideally which involves the gas particles having no attractions to each other. So using the ideal gas law: PV = nRT, you are doing so under this simplification. Of course we know things like London dispersion forces exist which does cause attraction which makes the ideal gas law only an approximation. This approximation holds though for most real world situations so the assumption is acceptable. Other gas laws like the Van Der Waals equation include corrections to account for these attractive forces.
Hope that helps.(6 votes)
- 11:00Nov 21, 2023
Something I've wondered about that bothers me. I know that NH3 is polar covalent and combines odorless gas of H and N....where does the smell cone from, esp when h20 is polar covalent also, combines odorless gas of H and O, but doesn't smell?(1 vote)
- [Instructor] In this video, we're gonna talk about something called kinetic molecular theory, which sounds very fancy. But as we'll see in the next few seconds, or the next few minutes, it actually helps build our intuition for what is actually going on with the gas or at least an approximation of what's going on with the gas. So first, let's think about the types of things that we know we can measure about a gas at a macro level. Now, what do I mean at a macro? I'm saying at a large scale, at a scale that's much larger than the scale of atoms or molecules. And we know the types of things that we can measure. We can measure pressure. How do we do you do that? Well, pressure is just force per unit area. So, you can do this. There's various contraptions you can use to measure pressure depending what you're using it for. Force, you can measure with springs and you can apply a certain forces to certain square areas. But these are all ways that you can measure pressure and we can measure the pressure of a gas in a container. You can measure volume of a container. That's actually pretty straightforward. You can imagine a container that looks something like this, it's volume. We know how to find the volume of a rectangular prism like this, or even if it was sphere or some other type of figure. There's many ways of measuring the volume without even being able to observe or even know that things like molecules exist. We know how to measure temperature, and we can do that in different scales. Kelvin is what we use 'cause it's more of an absolute scale, but you can use literally thermometers to measure temperature. And once again, you can measure temperature without knowing anything about atoms or molecules or whether they even exist. And you can also measure an amount of a substance. And in particular, we could say, you can measure the number of moles. Now you might say don't moles involve a certain number of a molecule or an atom. Well, they do, but the notion of a mole actually existed even before we knew exactly how many molecules, how many particles made up a mole. It was just viewed as an amount where people knew it must be some number of particles, but they didn't know exactly. So all of these things, we can measure at a macro level. And we know that we can connect them all with the ideal gas equation that tells us that pressure times volume is equal to the amount of the gas we're dealing with. And this is, of course, we're talking about an ideal gas and in future videos, we'll talk about how some gases approach being an ideal gas while some are less than ideal. But the amount we have measures the number of moles. You have your ideal gas constant that just helps us make all the units work out depending on our units for everything else. And then you have your temperature measured in Kelvin. And, scientists long before we were actually able to know about things like atoms or even observe atoms or molecules directly, or even indirectly, they were able to establish this relationship using these macro measurements. But how do these macro measurements and this relationship actually make sense at a molecular level? And that's what kinetic molecular theory provides us. It says, imagine the gas is being made up of a bunch of really, really the small particles. Those are really the gas molecules. And their collective volume is very small compared to the volume of the container. So, it's mostly empty space between those particles. Now, the pressure is caused by these particles bouncing into the sides of the container. Because at any given moment, you have enough particles bouncing off the side of any unit area that it's providing a force per unit area. It's providing a pressure. It assumes that those collisions are what's known as elastic, which we'll study in much more detail in a physics course, but it really says that your kinetic energy is preserved. You might already be familiar with the notion that kinetic energy is equal to one half times mass times velocity squared. And so the kinetic energy of these particles, when they bounce off, their mass doesn't change. The mass of the particles still there. And we're saying that the velocity is going to be preserved. So you have all of these really small particles, even their collective volume is small compared to the volume of the container. They're providing the pressure by having these elastic collisions with the side of the container. And temperature is related to the average kinetic energy of these particles. It would be proportional. The higher the temperature, the higher average kinetic energy. Now average kinetic energy is really important because some of these particles might be moving faster than others. And of course, N, the number of moles, tells us how many particles we're dealing with. We know that each mole has Avogadro's number of particles. So, if you just multiply the most times Avogadro's number, you have the number of particles. And what's cool about kinetic molecular theory, I know it's built as a theory, but this is fundamentally what chemists and physicists visualize when they imagine a gas in a container of some kind. And just to make it a little bit more clear, the axioms you could say of kinetic molecular theory, the assumptions of it, I'll give them here. And it's important to realize that these are assumptions and the real world, we have slight variation from it, but these assumptions get us a long way to explaining the behaviors of gases. So, we've already talked about it. Gas consists of particles in constant random motion. We've already talked about that. They're bouncing off the side of the container. The combined volume of the particles is negligible compared to the total volume in which the gas is contained. And that also matters when you talk about things like ideal gases, because if it stops becoming negligible, then you have to start thinking about the repulsive and attractive interactions, a little bit more. The particles exert no attractive or repulsive forces on each other. And that kind of builds into the last point I just made, which is if they did, then we're getting closer to being a less than ideal gas. And we'll talk about that in other videos. The collisions between the particles are completely elastic. So, they preserve kinetic energy and it's actually, they would also preserve momentum. And that the average kinetic energy of the particles is proportional to the Kelvin temperature. And we already talked about that, that the macro variable, the macro measurement of temperature is giving us an indication, it's proportional to the average kinetic energy of the particles.