In this example, it tells us that when climbers summit Mount Everest, they often use an oxygen mask to increase the amount of O₂ they inhale with each breath. Then it says the table below gives the concentration of oxygen and atmospheric pressure under three conditions. Use the table to calculate the partial pressure of oxygen under each condition, then answer the questions below. Alright. So we have this table here, and our columns are for sea level, Everest with supplemental oxygen, and Everest without supplemental oxygen.
And then our rows here are the concentration of oxygen, the total atmospheric pressure, and then what we need to calculate, the partial pressure of oxygen. Alright. So let's just do these one by one. We're going to go through first here for sea level. So at sea level, the concentration of oxygen is 20.9%, and the total atmospheric pressure is that familiar 760 millimeters of mercury.
Alright. So how do we figure out the partial pressure from those two numbers? Well, we multiply them. Alright. So I'm going to take 0.209 and multiply that by 760 millimeters of mercury. I don't have room in this box for my units, but we should write them in if we can. And that is going to give me right about 159 millimeters of mercury. That is my partial pressure of oxygen at sea level. Alright. Let's do the same thing for Everest with supplemental oxygen.
So it says here the concentration of oxygen that someone's breathing into that mask is going to be as high as 50%. I'll just note that's kind of an unrealistically high percentage at the top of Mount Everest even with an oxygen mask, but for this problem, we're going to go with it. So, you know, they're breathing in 50% oxygen, but the total atmospheric pressure is only 235 millimeters of mercury. Alright? A lot less.
So what's our partial pressure of oxygen in that scenario? Well, what we're going to do is we're going to take 50%, so 0.5, and multiply that by that 235, and that gives us 117.5 millimeters of mercury of oxygen that those people would be breathing in with that supplemental oxygen. Alright. Next, Everest without supplemental oxygen. Alright.
So just up there on Everest with no oxygen with you. Well, the concentration of oxygen, it's the same at sea level, the same percentage, 20.9% oxygen up there on top of Everest. But here the total atmospheric pressure is again way lower. We're there at that 235 millimeters of mercury. So how do we figure it out?
Well we multiply them together. 0.209 times 235. That's going to give me right about 49 millimeters of mercury. Alright. So those are my partial pressures. I have 159, 117.5, and 49 there. So moving on. Next question we have here, it says, using the information from the table, under which two scenarios would you expect the amount of oxygen dissolved in the blood to be most similar? Alright. Well, look at those numbers.
Which one do you think the amount of oxygen dissolved in the blood would be most similar? Well, I can sort of tell just by looking at it, but I'm actually going to do the math out here. So I'm just going to write it down, 159. We have 117.5, and we have 49. And so the difference between these is what I can do.
The difference between these two is 41.5, and the difference between these two is 68.5. Remember, we learned that what dissolves into the blood is dependent on the partial pressure. So we're just looking at here which two partial pressures are the closest. So that means that sea level and Everest with O₂, that's what I would expect the concentration of oxygen dissolved in the blood to be most similar because those partial pressures are closest to each other. Alright.
Final questions here. It says, what law allowed you to calculate the partial pressures? Okay. So in that first part when we're calculating partial pressures, what law was that? Well, that was Dalton's Law of Partial Pressures. Dalton's Law of Partial Pressures, that's what tells us that we can take the concentration of a gas, multiply it by the total pressure, and that will give us the partial pressure for that gas. Alright. Then we have what law allows you to predict how much oxygen would be dissolved in the blood. Well, what's the name of that law? That is Henry's Law.
Henry's Law lets us take those partial pressures, and it tells us that how much will dissolve is going to be dependent on those partial pressures. And we had a little memory tool to remember these to remember. We said that Dalton divided the pressure and Henry hydrated it. Dalton, you can break up that total pressure into partial pressures. Henry, it tells you how much is going to go into a liquid.
Alright. So that hopefully helps you better understand a little bit how to use these partial pressures. And you also know if you climb Mount Everest, bring that oxygen with you.