Why is oxygen and CO2 measured in pressure instead of mmoles/L like other ions?

Na, K, Cl...All measured in mmol/L in plasma. Why is oxygen and CO2 measured in pressure (mmHg)? (I guess it has something to do with them being bound to RBC but I don't know why the difference).

• By the ideal gas law, PV = nRT, you can find pressure by dividing both sides by volume, P = nRT/V. Since R is a constant, and you can usually assume T isn't changing very quickly for biological systems, you can pretend P = n/V, which is moles per unit volume, analogous to concentration. The reason the units are written in pressure is probably because gases are hard to weigh, but measuring pressure is easy. This also accounts for the affect of temperature, which shouldn't affect the soluble ions you mentioned, but would affect the "concentration" of gases. – user137 Jun 5 '16 at 11:49
• @user137 Good answer. Could you make it into an official answer – 360ueck Jun 5 '16 at 14:41

As you suspected, hemoglobin is the reason why O$_2$ and CO$_2$ in blood are expressed as a pressure (mmHg) and not a concentration. Most of the oxygen in blood is bound to hemoglobin; the concentration of free oxygen in plasma is low because oxygen does not dissolve well in water (which is of course the reason why hemoglobin exists in the first place).

When the oxygen content of blood is given as a pressure, it refers to the partial pressure that a gas (air) must have in order to be in equilibrium with the blood. This might seem a bit roundabout, but it captures the effects of both solubility and hemoglobin. In essence, this pressure describes how much oxygen the blood can release, which is the critical parameter for oxygenation of tissues.

The molar concentration of free O$_2$ in plasma is proportional to the partial pressure over the liquid, a fact known as Henry's law. The proportionality constant is the solubility of oxygen in the liquid, about 1.3 10$^{-6}$ M / mmHg for plasma. So at 100mmHg, the plasma O$_2$ concentration is about 130 $\mu$M. For comparison, the O$_2$ concentration in an ideal gas at 100 mmHg is 9 mM, or 70-fold higher.

See for example this book chapter for more information.

By the ideal gas law, PV = nRT, you can find pressure by dividing both sides by volume, P = nRT/V. Since R is a constant, and you can usually assume T isn't changing very quickly for biological systems, you can pretend P = n/V, which is moles per unit volume, analogous to concentration. The reason the units are written in pressure is probably because gases are hard to weigh, but measuring pressure is easy. This also accounts for the affect of temperature, which shouldn't affect the soluble ions you mentioned, but would affect the "concentration" of gases.

• No, this is not correct. The ideal gas law is not applicable here, since blood plasma is not a gas! Blood gas (partial) pressure refers to the pressure over a gas in equilibrium with the liquid. To find the concentration in blood, you must take in account O2 solubility in plasma and also the effects of hemoglobin. The actual plasma concentration is much smaller than the concentration P / RT in air. (And also you cannot pretend P = n/V; there is proportionality, but not equality.) – Roland Jun 6 '16 at 5:44
• @Roland I don't think OP was looking for a lecture in physical chemistry, but rather why the concentration of dissolved gases are expressed as pressure rather than a molarity. The ideal gas law should be a "close enough" demonstration that pressure is analogous to molarity, especially considering the practical aspects of taking measurements 100+ years ago when these unit conventions were put in place. If OP wanted to know how to determine the affinity of O2 for hemoglobin, that would be different, and the ideal gas law would be insufficient. – user137 Jun 6 '16 at 6:57
• Well, solubility is not a minor detail here, the concentration I get from the ideal gas law is 70x higher than the plasma concentration. And solubility + hemoglobin is the reason why concentrations are given in pressure, I think. See my answer. – Roland Jun 6 '16 at 10:28