Why is partial pressure used as a metric for oxygen in blood?

Question

I know haemoglobin basically acts as an oxygen buffer. But when we talk about non-bound oxygen in blood (or water), we're referring to the amount of O₂ in solution (aqueous state). From a biochemistry point of view, the concentration seems like a good metric to me.

So why is it that medical papers report blood oxygenation in terms of partial pressure pO₂ instead of concentration?

What does a partial pressure inside a liquid even mean? I tried interpreting the arterial pO₂ value of 11% atm (i.e. about 11kPa). When using the formula P = CRT at 310K (body temp), I found an oxygen concentration of about 4.3 mM. This value is 20× higher than the maximum O₂ concentration for water according to Henry's law. For those more familiar with DO values in mg/L, it would be 137.6 mg/L.

According to Henry's Law there is a limit to the amount of O₂ that can dissolve into water at a given pressure and temperature. I know the lungs experience a slightly lower partial pressure than the atmosphere. So how can the oxygen concentration be so high? Does their metric include the oxygen bound by haemoglobin (which would be misleading at the least), or is my calculation off somewhere?

Or does a pO₂ value correspond to the steady-state equivalent concentration at that pressure, in which case the arterial concentration in my example would be 0.14 mM (based on this table in wikipedia).

This similar question didn't answer my question: What is the state of aggregation (gas, liquid) of oxygen in blood?

Answer 1

The equation P=CRT is the relationship for the oxygen concentration in the gas phase. To get the mole fraction of oxygen the the liquid phase (blood), you first use the Henry's law constant. Here is a link that can help: https://www.google.com/webhp?sourceid=chrome-instant&ion=1&espv=2&ie=UTF-8#q=henry%27s%20law%20constand%20for%20oxygen%20in%20blood

If, after applying Henry's law, you end up with a concentration in units of x ml per unit volume, this result implies that x is the volume that the mass of dissolved gas (solute) would have it were an actual gas at the same pressure as that in the gas phase.

Answer 2

My own attempt at an answer:

Based on the answer from @Chester Miller it seems my last suggestion was accurate: the pressure values actually do correspond to an equivalent concentration at that pressure. So in fact a solute in a solution doesn't really have a partial pressure, just an equivalent partial pressure (EPP). This EPP value corresponds to the partial pressure of the gas outside the liquid--or the theoretical partial pressure should the liquid not be in contact with a gas.

In other words, the partial pressure in liquid is a bit of a misnomer. When we say e.g. the partial pressure of CO2 in the blood is 5%, what we mean is: that is the theoretical CO2 pressure required to achieve that concentration of CO2 if the liquid medium was in a steady-state exposure to a gas containing CO2.

This implies you can just convert between a concentration and its equivalent partial pressure by using Henry's constants (if known).

Now for the reason: in models where there is exchange between a gas and a liquid, the partial pressure determines the chemical potential driving the diffusion from one phase to the other--not the concentration. So one can more easily compare EPP values of a solute to determine the diffusion exchanges. This is typically relevant with lung-blood gas exchanges, which explains most of the physiology science papers use of partial pressure.

Unconfirmed part:

Lastly, although the PP drives the gas-liquid diffusion, once the gas is dissolved in the liquid it diffuses in a concentration-dependent manner. This implies at the tissue level in humans diffusion should be concentration dominated, the solubility differences between e.g. lipids and water shouldn't play a major role.

In contrast, if a gas interfaces with both lipid and water compartments, then the partial pressure dominates and the solubility becomes important, as it could allow high and low concentrations to exist near each other.