hello folks, let me apologize if this was addressed before, but every time I try to search for partial pressure of oxygen and water, I get problems unrelated to mine. I also searched Wikipedia and Google and found their explanations a bit too obtuse. But it's possible I'm just doing a bad job of research, if so, please feel free to point me to the right link... otherwise, I hope you can help me here. When calculating the actual partial pressure of oxygen entering an alveoli, we take the following steps... We first calculate barometric pressure (PB): usually 760 mmHg at sea level We then calculate the vapor pressure of water in the lung (because the lung releases water in the respiratory space) at body temperature (PW): usually 47 mmHg at 37° C We subtract the water vapor pressure from barometric pressure to get dry gas pressure (PDRY): PB - PW = 713 mmHg = PDRY Only at this point can we calculate the partial pressure of oxygen (PO2): which is 0.21 times the pressure of dry gas... 0.21 x 713 = 150.0 mmHg ← PO2 If you did not account for the water vapor reduction, you'd get a higher value of PO2 (160.0 mmHg) and you'd be wrong. OK... I can do this much with my eyes closed. my problem is why the water behaves as such, and some potential paradoxes that arise. I don't like to do math that seems magicky and I don't still get the actual physics behind it. Dalton's Law of Partial Pressures: a simple equation that states PTOT = P1 + P2 + Pn so the total pressure is the sum of the partial pressures. Why is it that atmospheric pressure in the airway (or anywhere else) is reduced by water subtractively, but by oxygen fractionally? In other words, if PO2 in dry environment is 160 mmHg, why doesn't it remain this way when moistened and instead is reduced by its fraction? why aren't we just ADDING PW to the 760, or (if the pressure can't go up since the lung inspires and creates negative pressure) don't we just fractionate the water and all other gas pressures? is it the fact water vapor is "liquid"? (am I not correct in treating water vapor as a gas, and subject to Dalton's Law?) The pressure of gases at 100°C: at this temperature, water's partial pressure would be 760 mmHg, equal to atmospheric pressure. Now, forgetting that human tissue boils at that temperature (imagine we were calculating this for a machine that is similar to a lung physically except it can handle the heat instead), wouldn't water's vapor pressure be such that there would be NO GAS PRESSURE? by the rules above, it seems at 100°C, the volume would be all water vapor and no oxygen, nitrogen, or argon pressure would be present... doesn't seem right. Keep in mind though temperature rise produces pressure rise, we're working in a constant pressure environment, so high temperature won't produce higher pressure if it can simply dissipate. In short why do we subtract vapor pressure from atmospheric pressure? why is PO2 not added and instead multiplied/fractioned? at 100°C, it seems PGASES would all equal ZERO since PW = PB thank you so much!