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 (Fig. 27a)
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 (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 term fractional concentration (F) can therefore be used to denote the relative quantities of gases in any mixture; thus Fn2 is 0.79 in dry air, and Fo2 is 0.21. The partial pres- sure of each gas in a mixture is that part of the total (e.g. barometric) pressure that is exerted by that gas, and is directly proportional to the quantity. Thus, 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. Water vapour behaves like any other gas, and exerts a partial pressure. The maximum or saturated water vapour pressure (SWVP) depends on the temperature: 2.33 kPa at 20 °C and 6.3 kPa at 37 °C. Inspired air quickly reaches body tempera ture and 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 (Fig. 27a). The water vapour content of room air depends on the conditions (e.g. desert vs seaside); 40% humidity denotes 40% of the predicted SWVP for that temperature.
Standardization. From the above and Boyle’s and Charles’ laws, it should be clear that gas volumes and partial pressures cannot be com pared unless corrected to a standardized pressure, temperature and humidity. Two standards are commonly used: standard temperature and pressure, dry gas (STPD), corrected to 1 standard atm (10 kPa), 0 °C and dry gas; and body temperature and pressure (1 atm), saturated with water (BTPS).
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, so that water exposed to the same partial pressures of CO2 and O2 will contain 20 times as much CO2 as O2. 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. The concept of a partial pressure of gas dissolved in a fluid (e.g. Po2 of blood) is sometimes difficult to understand, but merely reflects the partial pressure that would be required to dissolve that amount of gas in the fluid, according to Henry’s law. From the above, it can be deduced that 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. Typical values for partial pressures in the airways and blood are shown in Figure 27b.
Diffusion across the alveolar–capillary membrane (Fig. 27c)
Diffusion is discussed in Chapter 11. The rate of gas flow across the alveolar–capillary membrane = permeability × area × (difference in partial pressures), where the permeability depends on the membrane thickness, gas molecular weight and its solubility in the membrane (Chapter 11). Although CO2 is larger than O2, it crosses the membrane faster because it is more soluble in biological membranes. For gas transfer across the lungs, the permeability and area are commonly combined as the diffusing capacity (DL) for that gas, a measure of alveolar–capillary membrane function. Thus, the rate of O2 transfer = DLo2 × (alveolar Po2 – lung capillary Po2), or DLo2 = O2 uptake from lungs/(alveolar Po2 – lung capillary Po2). DLo2 is sometimes called the transfer factor. DLo2 cannot be estimated directly, because capillary Po2 cannot be measured. However, the factors affecting O 2diffusion also affect carbon monoxide (CO) diffusion. CO binds extremely strongly to haemoglobin, and so, if low concentrations of CO are inhaled, CO diffusing into the blood is completely bound to haemoglobin and capillary Pco remains close to zero (Fig. 27c). Thus, DLco = CO uptake from lungs/alveolar Pco, and can be easily measured as an estimate of alveolar–capillary transfer function. DLco is reduced by a decrease in lung exchange area (e.g. emphysema) or an increase in alveolar–capillary membrane thickness (e.g. lung fibrosis, oedema).
Diffusion and perfusion limitation (Fig. 27c)
Because CO binds so avidly and rapidly to haemoglobin, at low concentrations its rate of transfer into the blood is not affected by the blood flow, because there is always plenty of haemoglobin, it is limited solely by its rate of diffusion across the alveolar–capillary membrane, i.e. transfer is diffusion limited. For a poorly soluble gas, however (e.g. the anaesthetic nitrous oxide, N2O), the partial pressure in the blood rapidly reaches equilibrium with alveolar air, preventing further diffusion. In this case, increased blood flow will increase the rate of transfer, i.e. transfer is perfusion limited. O2 transfer is normally perfusion limited.