At an altitude of 10,000 ft, the partial pressure of oxygen in th...

Question

At an altitude of 10,000 ft, the partial pressure of oxygen in the lungs is about 68 mm Hg. What is the concentration in mg/L of dissolved O2 in blood (or water) at this partial pres- sure and a normal body temperature of 37 °C? The solubil- ity of O2 in water at 37 °C and 1 atm partial pressure is 1.93 * 10-3 mol>L.

Accepted Answer

Hey everyone, we're told that at 25 degrees Celsius, the partial vapor of carbon dioxide in a bottle of soda is 4.0 atmospheric pressure, calculate the concentration of carbon dioxide in grams per liter in the soda at 25 degrees Celsius. Given that the Henry's Law constant for carbon dioxide at 25 degrees Celsius is 3.3 times 10 to the negative two mole per liter times atmospheric pressure. To answer this question, we need to use the following formula, which is the solid ability of carbon dioxide is equal to our Henry's law constant times the partial pressure of carbon dioxide. First, we're going to solve for our scalability of carbon dioxide Plugging in our values, we know that our Henry's law constant for carbon dioxide is 3. Times 10 to the -2 Mole per leader times atmospheric pressure. And we were told that our partial vapor of carbon dioxide is 4.0 atmospheric pressure. Now, when we calculate this out and cancel out our units, We end up with a solid ability of 0.132 mole per leader. But since we want to sell for the concentration, We're going to go ahead and use dimensional analysis here, using carbon dioxide solar mass, we know that we have 44.01 g of carbon dioxide per one mole of carbon dioxide. So when we calculate this out and cancel out our units, we end up with a concentration of 5.8 g of carbon dioxide per leader. And this is going to be our final answer. Now, I hope this made sense and let us know if you have any questions.

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