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  • Hi. It's Mr. Andersen and this is chemistry essentials video 14. It's on

  • gases. When you see a hot air balloon flying, it looks pretty stationary. But if you're

  • to look at the molecules inside it, if we were to look at the heat inside it we'd see

  • there's going to be way more hot air at the top then at the bottom. It's less dense. And

  • that's why the thing is going to float. So when we're looking at gases, basically a real

  • gas is going to be made up of all of these independent particles that are all just shooting

  • around, moving around really really quickly. They're not connected together at all. And

  • so in science what we can do is we can model that using what's called an ideal gas. And

  • lots of times we'll use the ideal gas law. And this one, PV=nRT is one that you should

  • learn right away. And learn how to apply it over and over and over again. It's basically

  • built on two things. The kinetic molecular theory. This idea that molecules are bouncing

  • around and they have a huge amount of energy as we increase the temperature. And then the

  • Maxwell-Boltzmann distribution which means that even though they might collide every

  • once and awhile, the energy that they have is going to remain the same and their state

  • is going to remain the same. We can use the ideal gas law to figure out absolute zero.

  • But there are going to be some deviations. In other words times where this model just

  • doesn't work. And we have to start looking at real gases. And this usually occurs as

  • we get around condensation. That point at which that gas turns back into a liquid again.

  • And so this is going to be a real gas. Air, what's between me and the camera right now

  • is going to be a real gas. And an ideal gas is built on these two theories. First one

  • is kinetic molecular theory. It's that all of the molecules are bouncing around. Moving

  • off of each other. And as we increase temperature we're going to increase the molecular speed

  • of those molecules. And so that's a gas. Remember they're not connected together at all. They're

  • just moving around randomly. And then we have the Maxwell-Boltzmann distribution. And that

  • means that even though they might bounce into each other, they're not interacting with each

  • other. They're all going to be independent. And so this graph, what we're looking at are

  • a bunch of noble gases. And we're looking at their speed. And the probability of having

  • found molecules at that specific speed. And what we'll find is that real light noble gases,

  • like helium, are going to have a higher speed then those for example like argon. They're

  • going to have lower speed, but they're all independent of one another. And so when we're

  • looking at an ideal gas, we basically have these five properties. Volume, temperature,

  • pressure, n which is the number of moles. And then some kind of a gas constant. And

  • so you're going to have to look at your units to figure out what gas constant to use. But

  • if we use this model, which we use it three times, v is going to be the volume. So let's

  • say these are a bunch of gas molecules. Volume is going to be the size of the container that

  • they're in. T is going to be the temperature. How warm or cold it is. And that remember

  • is tied to the molecular motion. Pressure is going to be how much it's pushing out.

  • And if you push in on a balloon you can start to feel that pressure pushing out on you.

  • And then n is going to be the number of moles. How much stuff is going to be found inside

  • it. And so if we rearrange this PV=nRT, that's the relationship. Now you could learn all

  • of the different gas laws. But it's better to just learn the ideal gas law and then apply

  • it in different situations. And so we're going to use this simulation at phet. And we're

  • going to look at Boyle's law. And so that's the relationship between pressure and volume.

  • And so we're going to keep all those other constants, n, R and T are all going to remain

  • constant. What we're going to look at is how are volume and pressure related. So I'm going

  • to keep the temperature constant. And if we look at the pressure it's around 0.69 atmospheres.

  • Now let me decrease the volume. So I'm making the volume get smaller. And so what's going

  • to happen to the pressure? Well, we have to wait for a little bit. We have to wait for

  • that temperature to go back down to 300 kelvin again. But once it gets to 300 kelvin now

  • look, it was 0.69 atmospheres before, and now it's 1.73. And so as we decrease the volume,

  • pressure went up. Now let's increase the the volume. So we're making the volume go up.

  • And then let's look at what's going to happen to the pressure. And so we've got to wait

  • for that temperature to come back to 300 kelvin again. And so now we're going to see that

  • the pressure has decreased. So by increasing the volume, we decreased the pressure. And

  • so there's going to be an inverse relationship between pressure and volume. And you could

  • actually apply this in a problem. Let's say we had a balloon at 1.30 atmospheres. That's

  • going to be the pressure. It has a volume of 3.2 liters. What will the volume be if

  • the pressure is decreased to 0.62 atmospheres. And the temperature remains the same? Well,

  • we just have a before and after. So we have our pressure volume before, initial, and after.

  • So we'd put in our values before. Our after. And then we're going to solve for this final

  • volume. And so we could make sure our units are correct. And so what's going to happen

  • to the volume? The volume is going to increase as the pressure decreases as long as we keep

  • the temperature the same. And a weather balloon. As you send a weather balloon up, the weather

  • balloon is going to get bigger and bigger and bigger. There maybe changes in temperature,

  • but it's related to that relationship. Now let's look at what's called Charles' Law.

  • We're just going to look at volume and temperature. And what happens to the volume as the temperature

  • changes? Let's use this simulation again. We're going to try to keep the pressure constant.

  • And the same number of moles inside there as well. And so now we're going to increase

  • temperature. So we're adding heat. And watch what happens to the volume. As we increase

  • the temperature, the volume got bigger. As long as we kept the pressure the same? Why

  • is that? As we're increasing temperature there's more kinetic movement. Now let's lower the

  • temperature. As we decrease the temperature, those molecules aren't moving as quickly.

  • And what happens? Not pushing out as much. And so the volume is going to go down. So

  • if we increase temperature volume goes up. If we decrease temperature, volume is going

  • to go down. And so there is a direct relationship between the two. And so hopefully you're learning

  • how to apply ideal gas law. If they're both on the same side, we had an inverse relationship.

  • But now since they're on opposite sides of the equal sign it's going to be a direct relationship.

  • What's an application of this? Well you could find absolute zero in a basic chemistry class.

  • So what we could do is take a balloon and we could measure the volume of the balloon

  • when the gas is around freezing. And so let's say we do that by measuring the circumference

  • around the balloon and then using a little geometry to figure out its volume. If we increase

  • the temperature, so let's say we bring it inside, and we really increase the temperature,

  • the volume is going to increase. And if we really lower the temperature, put it in a

  • freezer, then we're going to decrease the temperature. And what you'll find is a linear

  • relationship. And so you could extrapolate that line and we could get to negative 273.

  • Which is going to be absolute zero. That's when there's no molecular motion and everything

  • has kind of come to a stop. Now let's look at Avogadro's law. Now we're looking at volume

  • and n. n remember is going to be the number of moles of the gas that we put in. So let's

  • look at this simulation right here. Well there's no temperature and no pressure inside the

  • container. And that's because there's no gas. And so let me squirt a little bit of gas inside

  • there. As we increase that, now those molecules are going to start bouncing around. And so

  • we have a specific volume. What do you think is going to happen as I increase the number

  • of molecules inside there? The number of moles inside there? Well, we're going to try to

  • keep the pressure constant, try to keep the temperature constant. And what's going to

  • happen is look, the more molecules we put in there, now we're increasing that. We're

  • really increasing that volume. And so we've got a direct relationship. More moles, more

  • volume. And so this is for an ideal gas remember. PV=nRT only works if we have an ideal gas.

  • What is an ideal gas? It's a gas where these things are randomly moving around. And there

  • also is no attraction between those. Those molecules aren't pulling on each other. If

  • they start to pull on each other then this is going to fall apart. And also an ideal

  • gas only works when we have an infinite number of, an infinite volume, which we're not always

  • going to have. But let's say we do have an ideal gas. If I just solve for n, this should

  • always be equal to 1. In other words, PV over RT should always be equal to 1 for any kind

  • of a gas. And so let me just plot that for a specific gas. Let's say we're looking at

  • nitrogen right here. It's ideal gas at different pressures should stay exactly at 1.0. But

  • let me throw up some data. This is what nitrogen looks like at 1000 kelvin. This is what it

  • looks like at 500 kelvin. This is what it looks like at 200 kelvin. And so these are

  • real gases. One's that have been measured. And so what are we finding? Is that they don't

  • fit the ideal gas law. And why is that? Well, as we slow everything down, as we make it

  • cooler, we start to get attraction between those molecules. And also we are going to

  • have a finite volume. In other words it's in a container when we're measuring it. And

  • so what we have to do is we have to start to look at real gas interactions as we get

  • towards condensation. And what's condensation remember? That's the point at which the gas

  • becomes a liquid. And so this would be water that was in the air that's becoming liquid

  • water on this cool bottle right here. And as we get closer and closer to condensation

  • then we have to throw out some of those ideal gas laws and use real gas interactions. And

  • so did you learn the following? To use KMT, that's kinetic molecular theory and force

  • concepts to make macroscopic predictions of what's going on? Remember those are the two

  • things that guide ideal gas law. KMT, things are moving around. And then this Maxwell-Boltzmann

  • distribution. They interact with each other, but they're not going to influence each other.

  • Even though they might bounce into each other. And the idea that faster molecules are generally

  • going to be lighter in a gas. Also could you apply ideal gas law, PV=nRT. We did that in

  • a number of different simulations. And then finally could you use this to solve some problems?

  • And we did this for a basic pressure volume kind of a problem. And so those are gases.

  • They're everywhere. And I hope that was helpful.

Hi. It's Mr. Andersen and this is chemistry essentials video 14. It's on

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B1 US volume gas ideal gas temperature pressure ideal

Gases

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    Wayne Lin posted on 2015/04/04
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