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Chemistry: Properties of Gases

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About this Lesson

  • Type: Video Tutorial
  • Length: 14:21
  • Media: Video/mp4
  • Use: Watch Online & Download
  • Access Period: Unrestricted
  • Download: MP4 (iPod compatible)
  • Size: 154 MB
  • Posted: 07/14/2009

This lesson is part of the following series:

Chemistry: Full Course (303 lessons, $198.00)
Chemistry: Gases (14 lessons, $20.79)
Chemistry: Gases and Gas Laws (6 lessons, $8.91)

This lesson was selected from a broader, comprehensive course, Chemistry, taught by Professor Harman, Professor Yee, and Professor Sammakia. This course and others are available from Thinkwell, Inc. The full course can be found at http://www.thinkwell.com/student/product/chemistry. The full course covers atoms, molecules and ions, stoichiometry, reactions in aqueous solutions, gases, thermochemistry, Modern Atomic Theory, electron configurations, periodicity, chemical bonding, molecular geometry, bonding theory, oxidation-reduction reactions, condensed phases, solution properties, kinetics, acids and bases, organic reactions, thermodynamics, nuclear chemistry, metals, nonmetals, biochemistry, organic chemistry, and more.

Dean Harman is a professor of chemistry at the University of Virginia, where he has been honored with several teaching awards. He heads Harman Research Group, which specializes in the novel organic transformations made possible by electron-rich metal centers such as Os(II), RE(I), AND W(0). He holds a Ph.D. from Stanford University.

Gordon Yee is an associate professor of chemistry at Virginia Tech in Blacksburg, VA. He received his Ph.D. from Stanford University and completed postdoctoral work at DuPont. A widely published author, Professor Yee studies molecule-based magnetism.

Tarek Sammakia is a Professor of Chemistry at the University of Colorado at Boulder where he teaches organic chemistry to undergraduate and graduate students. He received his Ph.D. from Yale University and carried out postdoctoral research at Harvard University. He has received several national awards for his work in synthetic and mechanistic organic chemistry.

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Thinkwell
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Let's talk about gases. Gases are a state of matter like solid or liquid, but gases have a lot of unique properties that are a little different from solids and liquids. In particular, the first thing to note is that the molecules or atoms that make up the gas are really far apart. If you think about a solid or a liquid, typically you think about the particles actually touching each other, whereas in a gas, they're typically very far apart. Because they are very far apart, it turns out that gases are very compressible. That is you squeeze on them and you can decrease their volume quite a bit. Another property is that the particles fill all of the space. So if you're talking about a sample of gas, like this balloon I have here, we have to describe the volume of the balloon as where the gases are. And finally, the particles of gas are in motion. And some of them actually have really high velocities, and we'll talk about the distributions of the velocities later on. But the motion of these particles smashing in to the walls of the balloon is what gives the balloon its shape.
The first quantity that we have to define, and I mentioned this a second ago, is the volume. And the units of volume we're going use is liters. If you're not familiar with the liter, a liter is about a quart, and it's related to a geometrical size, a cubic centimeter. One liter is equal to 1,000 cubic centimeters.
Now another quantity that we have to define for gases is pressure. And you have some familiarity with pressure. If you've ever been on an airplane, so here's our airplane, and we've got the planet here, it turns out that the weight of the atmosphere, and weight is a force, pressing down on the surface of the planet gives rise to what we call pressure. And at the surface of the planet the pressure is about 14.7 pounds per square inch. That's to say that if you take a 15 pound bowling ball and you concentrate all of it's weight on a one square inch area, that's pressing down on us all of the time. And we'll see the manifestations of that a little bit later.
Now you've probably experienced that when you go on an airplane ride your ears pop. And the reason why your ears pop is because the pressure at the surface of the earth is not exactly the same as the pressure as you go up in an airplane to 36,000 feet. Even people in Denver, at 5,280 feet will notice a difference in atmospheric pressure. Why is that? The reason is if you think about pressure being the weight of the air divided by the area, there's less air above you if you're at 36,000 feet than there is if you're at the surface of the planet, and less air less weight, less weight less pressure. So that's why your ears pop. And then when you go back down when the airplane lands, after having been up in the air for a while and your ears have acclimated to the lower pressure, then the increased pressure from landing the airplane, you really notice that. And sometimes it's somewhat painful.
So we have this 14.7 pounds per square inch. And we'll convert that to a metric unit later on. Why aren't we crushed by all of the pressure? If there is all of this pressure crushing down on us, why don't we just collapse into nothing? And the answer is if you imagine that we have a drum or a tin can here, there's one atmosphere of pressure on the outside, but because it's open to the atmosphere, there's also one atmosphere of pressure on the inside. And those exactly counterbalance and so this drum isn't crushed.
But imagine what would happen if we took all of the air out of the inside. In other words, the outside has got 1 atmosphere, but the inside has 0 atmospheres. Then what you might imagine is that 14.7 pounds per square inch is going to cause this thing to collapse. And let's look at that on this video. What I have here is a heavy-walled, steel, solvent drum of the sort that we get solvents like acetone and ethyl acetate in. It's a 20-liter can, very heavy. And what I've done is I've put about a cup of water into an identical drum sitting over here. And I've boiled that water. And you can see that it's boiling pretty vigorously. And now what I'm going to do is stopper the can with this rubber stopper and take it off the heat. And now what I'm going to do is I'm going to throw some cups of ice on top to get it to cool down a little more quickly. Now let's see what happens. This looks like it might take a little while, so I think I'll go get a cup of coffee and come back and take a look at it.
Okay this baby is well on its way to being crushed. It's going to keep going for quite some time. It's still almost too hot to touch. What this illustrated is that there is atmospheric pressure pressing down on us all of the time. When we boiled the water what we effectively did was we filled the can with water vapor. The water vapor was at 1atmosphere because the water was boiling, and it displaced all of the air that was inside the can. Then when we stoppered it, we closed the system. Then when we took it off the heat, what we did was we started to allow the can to cool. The water cooled. The water vapor cools. And as it cools, the water vapor pressure goes down. Whereas there was 1 atmosphere of water vapor on the inside of the can when the water was boiling, and 1 atmosphere from all of the way to the atmosphere pressing on the outside of the can so the sides of the can were perfectly fine, as the water cools, the pressure on the inside drops. The pressure on the outside stays the same. What has to happen? The can has to collapse. It has to get crushed by atmospheric pressure. It's going to keep going for quite some time. You can see that it's well on its way to being crushed much, much, much more than I could have done by jumping on it or anything like that.
Have you ever wondered how a straw works, how you get liquid up from your glass up into your mouth? The reason why a straw works is because there's one atmosphere of pressure pushing down on the surface of the liquid, but as you suck on the end of a straw, you're removing some of the gas particles that are inside the straw. And in response to the fact that there's 1 atmosphere of pressure pushing on one side, and less than 1 atmosphere pushing on the other side, the liquid goes up the straw, so on the mouth side, pushing down here, less than 1 atmosphere.
To show you that this is actually the way a straw works, because the atmosphere is finite it only has so much weight. There's only so much pressure that it can exert, what that suggests is that you should only be able to pull liquid up a straw a certain height and no higher. Even if you've got an absolute perfect vacuum at your mouth, because the atmospheric pressure is finite, you would only be able to push liquid up so far. And that's exactly what happens.
If we imagine a glass on the ground, and you had a really special straw, that special straw is 32 feet long, that's as high a column of water you could raise at sea level, where atmospheric pressure was pushing down on the top of the cup. So push down here, raise a column of 32 feet of water, and no higher. Where this comes into play is if you happen to be using well water. In order to pump the water out from the well, if you just pump it by sucking at the top of the well, you can only raise a column of water 32 feet. So if your well is 35 feet, you're out of luck.
32 feet is a little inconvenient for our laboratory so let's repeat this experiment using a more dense liquid instead of water, and then we'll get to a more reasonable size. And that dense liquid that we're going to use is mercury. Mercury you know is a metallic liquid. And if we repeat that experiment, here's a pan of mercury. Here's a glass column. We're going to create a perfect vacuum here up at the top. And the weight of the air pushing down, the atmospheric pressure, will allow us to raise a column of mercury 760 millimeters high. We're going to give the millimeter a special name. We're going to call it the Torr, after Evangelista Torricelli who invented the mercury barometer.
So this 760 atmospheric pressure is how meteorologists--when they talk about barometric pressure, the height here is going to change a little bit from day to day, because the weight of the air changes from day to day. And so by watching the height changes at the level of this mercury, we can learn whether or not--for instance, if the pressure is high, we're probably going to have really clear weather. And if the atmospheric pressure is dropping, so this column is going to drop, that's going to mean that we're going to have very changeable weather.
How are we going to relate this kind of device to a sample of gas that we'd like to measure for maybe doing an experiment? And the answer is we're going to use something called an open-end manometer. I've sketched a couple here. An open-end manometer is a glass tube. This is where our sample is going to go. We'll put some mercury in the bottom here. Since it's open to the atmosphere, there is 1 atmosphere of pressure pushing down at sea level. And if we imagine that the other side is open, then there is also 1 atmosphere of pressure pushing on the left-hand side, and the two levels of mercury are exactly balanced. If we imagine closing off this piece of glass on the left-hand side, if the two levels of mercury are still exactly the same, then what we're going to say is that we have 1 atmosphere of pressure of gas on the left-hand side. What if we have something other than 1 atmosphere? For instance, if we have a little more than 1 atmosphere, then this is what the open-end manometer is going to look like. Because we have more gas on the left-hand side, it's pushing down a little bit harder on the mercury so it unbalances the levels. What we can do is we can measure the height on the left-hand side, the height on the right-hand side, and relate them by this quantity h. And what we're going to say is that the pressure of the gas on the inside is equal to 760, that's our atmospheric pressure, plus h, where h is this height difference in Torr. And it's also the case that, if we expressed it in atmospheres, then the pressure of the gas on the inside is atmospheres.
Now Torr and atmospheres are two units that you're probably going to see a lot, but there are some other units of pressure that people talk about. Physicists in particular use SI units. And the SI unit of pressure is the Pascal. 1 atmosphere is equal to approximately 1.013 x 10^5 Pascals. So obviously a Pascal is a very small unit. And a unit related to that is the bar. And a bar is 10^5 Pascals. So one atmosphere is equal to approximately 1.013 bars. And then I said before that this atmosphere is also 14.7 pounds per square inch.
So we can easily solve a problem such as what is the pressure in Torr that corresponds to 2 atmospheres? And the reason why I chose 2 atmospheres is that that's approximately the pressure of carbon dioxide that's on top of a bottle of soda just before you open it. So you probably notice when you take the lid off, there's a sound of gas escaping from the bottle. That's the excess carbon dioxide. So we can relate our pressure in Torr to pressure in atmospheres by multiplying by one of our units of 1, a unit of 1, . Again we choose it this way because we want the units of atmospheres to cancel. So if we plug in our 2 atmospheres x , we get 1900 Torr.
So what have we talked about here? We've talked about the fact that gases are a little different from liquids and solids because they take up a lot of space. They are in motion, some with really high velocity. We described volume, which is a characteristic of a gas that we have to specify when we're describing our sample. And then we got into this concept of pressure. We talked about how the weight of the atmosphere creates pressure on all of us, enough to crush a big strong steel can. And then we talked about the open-end manometer, which allowed us to relate the pressure of a sample of gas that we might be interested in to the atmospheric pressure. And finally, we talked about several different units of pressure, and showed how we can, in this case, relate units of pressure simply by multiplying by units of 1.
Gases
Gases and Gas Laws
Properties of Gases Page [1 of 3]

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