Gases in air

There’s a bit of mixture of mathematics, chemistry and physics involved in understanding gas exchange in the lung. For some, this may just be a reminder of material covered in school (albeit in a unfamiliar context), whereas for others this may all be a bit of a shock to the system, so we’ll go through it gently.

Let’s start with air. If you analyse air, you’ll find that it’s made up of two main gases:

This is a slight approximation. There are tiny fractions of other gases, including CO2, that collectively make up less than 1%. Climate scientists worry about CO2 as a greenhouse gas – and they probably have a point – but they are talking about CO2 in the range of parts per million, not percentages (parts per hundred). A tiny change in CO2 might well affect global warming, but it still doesn’t register on the scale that we are considering, where CO2 for all purposes doesn’t exist in air. This becomes important later because whenever we find some CO2 anywhere, we know one thing for certain – it was produced by the body.

Now, if barometric pressure (PB) is 100 kPa (a fair assumption close to sea level), then each of the two main gases makes an individual contribution to the total (you may recognise this as Dalton's Law). The pressure exerted by each is its partial pressure:

Working in kPa (kilopascals) is quite handy because unless you climb a mountain (and therefore reduce PB), the partial pressures of gases in air are more or less equal to their contribution as a percentage. If we did climb up a mountain to a level where PB is, say, 75 kPa then the partial pressures would change, but the percentage composition remains the same. So:

So, when we say that the air is thinner at altitude, what we really mean is that the partial pressure of O2 is less, whereas the percentage of O2 in air is still the same.


Partial pressures of gases in solution

This is where things get a little tricky. We need to deal with the amounts of gases in the air that we breathe, but we also need to consider how those gases are dissolved in blood and transported around the body. Once in solution, gases don’t exert a pressure. Dissolved gases have no direct physical effect on blood pressure at all. So why do we talk about the partial pressures of gases in blood?

Because it saves a lot of bother.

When we say that a solution has a partial pressure of O2 equal to 21 kPa, what we are really saying is that whatever is dissolved in the water is equivalent to allowing it to equilibrate with 21 kPa of gaseous O2. The actual content of gas dissolved in water depends on the gas concerned. For example, carbon dioxide is 25 times more soluble than oxygen. So if you bubbled one beaker of water with 10% CO2 and another with 10% O2, the CO2 beaker would have substantially more CO2 dissolved than O2 in the O2 beaker. Being “correct” in physical chemistry terms just makes life hard. It’s easier to say that the CO2 beaker has a partial pressure of 10 kPa.

How is all this fuss about partial pressures, percentages and vagaries about the physics of water going to help us? In clinical practice, we often alter the gases that people inhale. Basic physiological principles allow us to predict what the effect of putting someone on 60% O2 should be. If the result isn’t what is predicted from first principles, we know something is wrong.




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