Alright, folks. We now want to talk about respiration, the movement of oxygen and carbon dioxide in and out of the blood. But before we really get into it, we need to introduce another physics concept, and that is partial pressure. Now again, don't worry. We don't need to get into any advanced calculations here. We just really need to understand how this works conceptually. That's going to be because the movement of gases, and the gases we really care about here are oxygen and carbon dioxide, depends on pressure gradients. Right? That's how gases move in and out of your lungs during ventilation. And for respiration, that's how these gases are going to move in and out of the blood. Now remember that molecules are always going to move down a gradient sort of on their own, and that is a passive process. Now I say that it's a passive process. That's important because we don't need to use cellular energy to get these molecules to cross the membranes in and out of the blood. They do it on their own because of these pressure gradients. They're just going to naturally move down their pressure gradients. So as long as the pressure gradients are all aligned properly, these molecules will move in and out of the blood basically on their own. And when we're talking about this, we want to remember that for diffusion of gas molecules, we are going to care about partial pressure. We are always talking about pressure, not concentration. Now I know for me, that's a little not immediately intuitive because when I think of diffusion, I think of things moving from a high concentration and spreading out on their own until they're in sort of an equally low concentration everywhere. But when we are talking about gas molecules, we don't want to talk about concentration. We want to talk about partial pressure. Alright. So what is partial pressure? Well, Dalton's law of partial pressure, and that Dalton's law, that is a vocab word that you may need to know. Dalton's law of partial pressure says that in a mixture of gases, the total pressure is going to equal the sum of the individual pressures. So if we have a gas that's a mixture like the air around us that's made up of many different types of gas, well, the total pressure in that gas is going to be equal to the sum of gas a plus gas b plus gas c, etcetera. All those pressures added together will equal the total. Alright. Seems pretty straightforward, but let's see what we mean by actually looking at atmospheric pressure. So atmospheric pressure is 760 millimeters of mercury, and we have this table here that we can fill in. And we see on the left, we have the 4 major gases that make up the air around us, nitrogen, oxygen, argon, and carbon dioxide. And it shows here what their concentration is in the air. But, again, for the movement of these molecules, what we really care about is their partial pressure. So to figure this out, well, if 760 millimeters of mercury is the total pressure and about 78% of the gas is nitrogen. Well, to figure out the partial pressure of nitrogen, it's pretty straightforward. We're just going to take 78% of 760. So I'm going to do the math here. We do 0.7808×760 millimeters of mercury, and that gives us 593.4 millimeters of mercury is the partial pressure of atmospheric air due to nitrogen. Alright. Another way to think of that is that if you had a chamber that had atmospheric air in it and it was at 760 millimeters of mercury. If we took away all the other molecules, we just got rid of those molecules, but we kept the same number of molecules of nitrogen, what would the pressure be in that chamber? Well, if it was just the nitrogen in the air and we got rid of everything else, the pressure would be 593.4 millimeters of mercury. Alright. We can keep going down the list here. Next, we have oxygen. That's the next most common gas. You can see here it's about 21% of the air around us. Well, we can do this simple math here. So 21% of 760. So we do 0.2095×760, and that gives me 159.2 millimeters of mercury, is the partial pressure of oxygen in the atmospheric air around us. So again, if you just had the same number of molecules of oxygen in that chamber and got rid of everything else, what would the pressure be? It could be 159.2 millimeters of mercury. Alright. We can keep going. Argon. Argon. We don't really talk about that much. It is about 1% of the air around us, but it doesn't react physiologically. So we generally don't care about it in physiology, but we can do the simple math. So 0.93% of 760 is 0.0093×760, and that gives me 7.1 millimeters of mercury. Carbon dioxide, an incredibly important gas, but you can see it's actually pretty rare in the air around us. It's 0.0004×760 is going to give us this, and that comes out to, with a little rounding here, 0.3 millimeters of mercury is the partial pressure of carbon dioxide. Alright. Now Dalton's law of partial pressure says, now that I figured all those out, if I add them up, what should it equal? Well, I can take the sum, and I'll let you check my math here. It does work. I add all those up. It should equal 760 millimeters of mercury is the total pressure. But again, I can break it into those individual pressures. And why is that important? When we're talking about the diffusion of these molecules, I care about this column. I do not care about this column. I care about the partial pressures. Alright. We're going to see an example where that really makes itself clear why you care about partial pressures coming up. Check it out. I'll see you there.
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- 1. Introduction to Anatomy & Physiology5h 40m
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22. The Respiratory System
Law of Partial Pressure
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