Greetings. In this video I'll be discussing how gases behave under different conditions. And we'll be watching some interesting experiments that demonstrate that behavior. The first thing we need to talk about, is the kinetic molecular theory of gases. You can see in these pictures, that gases are particles in constant, random motion. Professor Trough's trying to show that motion by having little traces behind each gas particle. The particle can either be a single atom, as would be the case for helium or argon gases, or it can be a molecule. For example, the nitrogen and oxygen in our atmosphere are molecules. So, in all of these cases, the gases are in constant random motion, and they can bounce off surfaces, or they can bounce off the walls of containers. Gas particles do not attract or repel each other. And any collisions between the particles, or between the particles and the wall of a container, are elastic collisions. Third, the gas particles are far apart. This is another way of saying that the volume of the particles themselves, is very small, or negligible, compared to the volume of the container that they're in. Finally, the average kinetic energy, or speed of the particle, depends upon the temperature. We can use this kinetic energy of gases to provide power. We do that all the time with windmills. Another example of using gas molecules' motion to provide power is sailboats. One of the properties that hasn't come up before in this course, but is very important for gases, is the concept of pressure. So, the gases are expanding to fill the shape of their container. We saw that in the last video when we had the carbon dioxide gas, which was subliming from the dry ice, filling the container that was either a flexible balloon or a glove, or was an inflexible container that eventually lost its lid. That pressure comes from the molecules colliding with the surface. It's not much different than bouncing a tennis ball off the floor, except in this case, the atoms are much, much smaller than the tennis ball. Pressure can be measured in force per unit area. For example, you could measure the pressure that your pen exerts when you set it down on the table. Or you could measure the pressure that occurs when you push on a wall with your hand. The SI unit for pressure is the Pascal. At sea level, atmospheric pressure is 101,325 Pascal. Because I did a lot of high pressure chemistry in graduate school, I got very adapt at converting between psi, because we had pressure gauges that showed psi in the readout, and bar which we often reported in papers. Millimeters of mercury is often used in meteorology, they also use bar in meteorology. Atmosphere is the pressure unit that's often used by scuba divers. One atmosphere is the pressure of our atmosphere at sea level. You can imagine as you go down under water, the pressure exerted by the water becomes greater. You've probably experienced that if you've dove down under the water, even to a depth as shallow as 10 or 12 feet. You can feel the pressure on your eardrum. This is a diagram that illustrates pressure. So the force is perpendicular, to the surface that's experiencing the pressure. And we, as you can see, with units like psi, which stands for pounds per square inch, we are looking at how much force is exerted per unit area. Let's look at some other common units. So, the pascal and the kilopascal again are the SI units, and here the ways we can convert between those other units that we talked about earlier. Let's conduct demonstrations to look more closely at the property of gases. In the first experiment, we have put an air filled balloon in a bell jar chamber that is attached to a vacuum pump. Our normal ambient pressure is 1,001,325 pascal, which is about 1 atmosphere or 760 millimeters of mercury, at sea level. Again, pressure is the force per unit area, and the atmospheric pressure results from collisions of gas molecules from the air on the exterior of objects. Let's see what will happen when Dr. Lyle turns on the vacuum pump, to reduce the external pressure in the atmosphere surrounding the balloon. What do you observe happening? As you can see, as we lower the outside pressure by pumping air out of the chamber, the pressure exerted from the gas molecules inside the balloon causes it to get larger and larger. The pressure on the interior and exterior of the balloon is being equalized by expansion of the volume of the balloon. What will happen if we turn off the vacuum pump, and then switch the valve open to allow the air from the atmosphere to go back into the bell jar? Thank you for submitting your answer. Opening a valve to the atmosphere, will cause the pressure to go back up inside the chamber, so the balloon shrinks back to its original size. Recall that Boyle is the one who summarized this relationship between volume and pressure. Boyle's Law can be summarized as is shown here. If we have two different chambers with different volumes, so we have volume one, volume one which is relatively large, and volume two, which is relatively small. Assuming that we have the same number of moles in both of these containers, and that they are both at the same temperature, the pressure in container one times the volume in container one, should be equal to the pressure in container two times the volume in container two. So, here is a single container and we're just moving a piston up and down. Decreases in the volume forces the molecules into a smaller space. In that smaller space, the molecules, which are colliding with the walls of the container, will collide with the walls of the container more frequently, and that increases the pressure. We can say that pressure and volume are inversely proportional at constant number of moles of gas and temperature. In other words, if the container on the left has a 4 liter volume, and the pressure of the gas is 4 atmospheres, then that gives us 16 atmospheres per liter. Let's assume that, in the container on the right, we have the same number of moles of gas at the same temperature, but the pressure of the gas can now be observed with a pressure gauge to be 8 atmospheres. What must the volume be of container two? Please go ahead and try to answer that now. [BLANK_AUDIO] Thank you for your answer. Because pressure and volume are inversely proportional, as the pressure goes up to 8 atmospheres, the volume must go down, in this case to 2 liters. That gives us a constant, of 16 atmospheres per liter here. Next, we will cool some gas filled balloons using liquid nitrogen. First, some liquid nitrogen is being poured into a bowl. You can see fog from the water in the atmosphere condensing into little water droplets, because of the extremely cold temperature of the liquid nitrogen. Liquid nitrogen boils at minus 197 degrees Celsius. So, it is rapidly boiling at room temperature. Dr. Lyle is using tongs to fill a small flask, so that you can see this boiling better. He will pour it from the bath onto the counter in front, where it rapidly vaporizes. Liquid Nitrogen is a cryogenic fluid, which means it will rapidly freeze living cells. So, for example, you never put your hand in the liquid nitrogen, even if you are wearing gloves, because it will freeze and kill your body's cells. Let's determine what will happen if we lower the temperature of an air filled balloon, by submerging it in the liquid nitrogen bath. Mary Jane is being very careful not to put her hands in the liquid nitrogen, so she is using tongs. What happens to the volume of the balloon as it cools? Thank you for submitting your answer. When she lifts up the balloon, we can observe that the volume of the cold balloon is much smaller that the volume that it was at room temperature. The gas inside the balloon expands again as the balloon warms back up. Eventually, when the balloon gets back to room temperature, it will occupy the same volume that it did before the experiment. Now the experiment is being duplicated, but this time with a yellow balloon. Again, it's filled with air. Can you see the liquefied air in the bottom of the balloon, when she pulls it out of the liquid nitrogen bath? Since air is composed of about 78% nitrogen, what we are seeing is a solution inside the balloon, with liquid nitrogen as a solvent, and the other gasses dissolved in the liquid nitrogen. For example, the oxygen and carbon dioxide from our atmosphere are now dissolved in the liquid nitrogen that condensed when we cooled this balloon. Let's repeat this experiment with an air-filled balloon that is much more transparent and a different shape. We can really see the balloon shrinking this time. If you look closely, then you can see the liquid air again at the bottom of the balloon, when Mary Jane pulls it out of the liquid nitrogen bath. There is a dramatic change in the volume of the gas inside the balloon as it warms back up to room temperature. I hope you have enjoyed this initial look at some of the properties of gases at different temperatures and pressure. Again, the volume and temperature are summarized by Charles' law, which says, that temperature and volume are directly proportional, at constant number of moles and constant pressure of gas. K is also a constant that goes with Charles' Law. We can demonstrate this law graphically as shown on the left. In this case, the volume that the gas occupies is plotted as a function of temperature. Now, there are two x axes drawn here, only because the temperature is shown on two different scales. It's shown as Kelvin and also as degrees Celsius. We'll talk about how to convert between Kelvin and Celsius on the next slide. As you can see, there is a linear relationship between volume and temperature, at every amount of gas that was tested in this experiment. So there were four different amounts of gas used in this experiment. And each time the container with that amount of gas was gradually heated and the volume was measured at different temperatures. In all cases, there was a linear relationship. So again, Charles' Law tells us the relationship between volume and temperature. They are directly proportional. If we increase the volume, which in the numerator in this expression, then the temperature must also be increasing, which is in the denominator. We need to be careful about the type of temperature we use in this relationship. I always convert the temperature to Kelvin. Kelvin is the degrees Celsius plus 273.15, if you want to be more precise. So again, if you know the volume and you know the temperature, and then, for example, you decrease the volume, but you keep the pressure and the number of moles constant, then the temperature must also be decreasing. We saw that in the demonstration. So far I have outlined the kinetic molecular theory of gases. We have also learned that, for gases, volume and pressure are inversely proportional, and volume and temperature are directly proportional. Tune in to the next lecture to learn more about gas law, and also to see some worked examples.
