[box type=”download”] — Partial pressure as the contribution to barometric pressure exerted by that gas — Normal partial pressure values of oxygen and nitrogen in ambient room air — Henry’s law as the determinant of the quantity of gas dissolved in a fluid — Oxygen and carbon dioxide partial pressures in inspired, alveolar + expired air [/box]
Dry air contains 78.1% N2 and 21% O2; other inert gases account for the balance (0.9%), but are normally pooled with N2 (i.e. N2 = 79%). The small amount of CO2 in air (<0.04%) is usually ignored.
Partial pressures and fractional concentrations
The volume of a fixed amount of gas is inversely proportional to the pressure (V ∝ 1/P; Boyle’s law) and proportional to the absolute temperature (V ∝ T; Charles’ law).
An ideal gas occupies 22.4 L per mole at 1 atm pressure (101 kPa, 760 mmHg) and 0 °C (273 K), and thus the volume of each gas in a mixture is directly proportional to the quantity of that gas in moles.
The fractional concentration (F) of N2, Fn2 is 0.79 in dry air, and Fo2 is 0.21.
According to Dalton’s law, the partial pressure of O2 (Po2) in dry air is Fo2 × barometric pressure (PB), e.g. 0.21 × 101 kPa = 21.2 kPa.
At the summit of Everest, PB is ∼34 kPa, but the relative proportions of gases are the same as at sea
level, and so PO2 is 0.21 × 34 kPa = 7.14 kPa.
Water vapour pressure.
Inspired air quickly becomes fully humidified (100% saturated) in the airways.
Water vapour dilutes the other gases, so that PN2 and PO2 will be lower than in dry air.
Thus, PO2 will be 0.21 × (PB – saturated water vapour pressure) or, under these conditions, 0.21 × (101 − 6.3) = 19.9 kPa.
It is clear that gas volumes and partial pressures cannot be compared unless corrected to a standardized pressure, temperature and humidity.
Gases dissolved in body fluids
The quantity of gas dissolving in a fluid is described by Henry’s law: dissolved gas concentration = partial pressure of gas above fluid × solubility of that gas in that fluid.
The solubility tends to decrease with a rise in temperature, and varies significantly between gases.
For example, CO2 is 20 times more soluble than O2 in water.
Henry’s law describes an equilibrium – increasing the partial pressure of a gas will cause more to dissolve in the fluid until a new equilibrium is reached.
So, the movement of gases between gas and fluid phases (e.g. alveolar air and capillary blood) will be dependent on the difference in partial pressures rather than the concentration.