RESPIRATION PHYSIOLOGY: ALVEOLAR GAS EXCHANGE
REVIEW OF PARTIAL PRESSURE
Gases are transported across the alveolar wall between the alveoli and pulmonary capillary blood by the process of simple passive diffusion. Generally, the driving force for movement by diffusion is concentration difference or concentration gradient. However, when discussing the movement of gases, concentration is not the appropriate measure of driving force. This is because, when dissolved in body fluids, much of the gas may be immobilized by binding (e.g., O2 to hemoglobin) or transformed into another chemical form (e.g., C02 to bicarbonate). Thus the use of concentration would give an erroneously high force for diffusion, since concentration (or content) includes bound molecules which are not free to diffuse. When discussing movement of gases, we require some more accurate measure of tendency to diffuse; such a measure is given by partial pressure.
Definition. The partial pressure of a given gas is defined as the pressure exerted by that gas alone. For example, consider the gaseous mixture of oxygen, nitrogen, carbon dioxide, and water vapor, the typical mixture found in alveolar gas. If this mixture was enclosed in a sealed container, it would develop a pressure on the walls of the container, by the mechanism of collision between the gas molecules and the container walls. The force or pressure developed by all the molecules of the mixture as they bounce off the container walls is the total pressure exerted by the gas. The partial pressure of any component is the pressure developed by the molecules of that component acting alone. That is, the partial pressure of oxygen is the pressure developed by the oxygen molecules, the partial pressure of nitrogen is the pressure developed by nitrogen molecules, and similarly for the other components of the mixture. Clearly, the sum of the partial pressures of the individual components must equal the total pressure on the container walls.
We could measure the partial pressure of any component by removing the other components and determining the remaining pressure. Thus, the partial pressure of oxygen could be measured by removing nitrogen, carbon dioxide, and water vapor, and determining the pressure resulting from the remaining oxygen molecules. This is a tedious procedure, which we can circumvent, either by the use of other types of measuring instruments (special electrodes, etc.), or by calculation using the following rules.
CALCULATION OF PARTIAL PRESSURES
Gas mixtures. In a mixture of gases, the partial pressure of a given gas is directly proportional to the volume that gas contributes to the total volume. That is, if we define the fraction of any gas, Fx, as the ratio of the volume of that gas, Vx to the total volume of all gases:
Fx = Vx/Vtotal (1)
then the partial pressure contributed by that gas, Px, is simply that fraction multiplied by the total pressure, or
Px = Fx ·Ptotal (2)
For example, dry air is composed of approximately 21% oxygen (actually 20.9%), and 79% nitrogen (actually, nitrogen is closer to 78%, with the remaining 1% consisting mainly of argon; for purposes of respiration, we usually lump the inert gases with nitrogen). Assuming 21% O2 and 79% N2 and assuming the standard barometric pressure of 760 mmHg, the partial pressure of these gases is
PO2 = 0.21 · 760 = 160 mmHg
PN2 = 0.79 · 760 = 600 mmHg
Note that it follows from equations (1) and (2) that the sum of the partial pressures must equal the total barometric pressure (in the above example, 160 + 600 = 760).
Vapors. When a gas is in contact with a liquid, and is in equilibrium (saturated) with the liquid, the partial pressure of the gas is a function of temperature. The one gas to which this applies in a normal respiration is water. The lungs and airways are always moist, and inspired gas is rapidly saturated with water vapor in the upper segments of the respiratory system. The temperature in the airways and lungs is almost identical with deep body temperature (approximately 37°C); at this temperature water vapor has a partial pressure of 47 mmHg. (Note that the gaseous form of a liquid frequently is termed a "vapor").
Using the value of 47 mmHg, we can calculate partial pressure of oxygen and nitrogen in inspired air, after the gas mixture becomes saturated with water vapor in the upper airway (so-called tracheal air):
Ptotal = 760 mmHg
PH20 = 47 mmHg
713 mmHg for remaining inspired gases (21% O2 and 79% N2)
PO2 = 0.21 · 713 = 150 mmHg
PN2 = 0.79 · 713 = 563 mmHg
That is, since water vapor partial pressure must be 47 mmHg in a saturated gas mixture at 37°C, the total pressure remaining for the inspired gases is only 760-47 or 713 mmHg. The composition of this remaining gas is 21% O2 and 79% N2, giving the partial pressures indicated above.
Dissolved gases. When a gas is dissolved in a liquid, its partial pressure is determined by the components of the solution which bind or combine with the gas, and the total content or concentration of the gas in the liquid. Such a relation between partial pressure and content is termed a saturation or dissociation curve. Each gas has a particular saturation curve determined by the physical and chemical properties of the solution. When discussing the transport of gases in the blood, one must consider in detail the saturation curves for oxygen and carbon dioxide.
Summary. In a gas mixture, the partial pressure of each component is determined by the composition of the mixture. For a gas in equilibrium with its liquid state, partial pressure is determined by the characteristics of the liquid and its temperature. The partial pressure of each gas in solution is related to the content or concentration of that gas by the particular dissociation curve.
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© AC Brown 2004