When scientists began studying the relationship between pressure, temperature, and volume of gas, they realized that all gases followed the same relationship. There are several gas laws that apply to human physiology.
The kinetic theory of gases makes the following assumptions:
The molecules in a gas are small and very far apart. The majority of volume that a gas occupies is empty space.
Gas molecules are in constant, random motion—just as many molecules are moving in one direction as another.
Molecules can and will collide with each other and with the walls of the container. Collisions with the walls account for the pressure created by the gas.
When collisions occur, the molecules lose no kinetic energy; that is, the collisions are said to be perfectly elastic. The total kinetic energy of all the molecules remains constant unless there is some outside force that acts on the system.
The molecules exert no attractive or repulsive forces on one another except during the process of collision. Between collisions, the molecules move in straight lines.
DALTON’S LAW OF PARTIAL PRESSURES
Dalton’s law of partial pressures states that the partial pressure of a gas in a mixture of gases is the pressure that gas would exert if it occupied the total volume of the mixture. Dalton’s law is
where Px is the partial pressure of gas X (mm Hg), PB is the barometric pressure (mm Hg), PH2O is the water vapor pressure at a given temperature (mm Hg), and F is the fractional concentration of gas. For example, the partial pressure of O2 (PO2) in dry inspired air at 37°C would be calculated as follows:
PO2 = (760 mm Hg − 0) × 0.21 = 160 mm Hg
where 760 mm Hg is the atmospheric pressure of dry air at 37°C and 0.21 is the percent oxygen composition of air. We can contrast this with air in the trachea that has been humidified by the nasal turbinates:
PO2 = (760 mm Hg − 47 mm Hg) × 0.21 = 150 mm Hg
where subtracting 47 mm Hg from the atmospheric pressure of 760 mm Hg corrects for the added water vapor pressure, causing the PO2 to be reduced by 10 mm Hg. Inhaled anesthetics will diffuse from the lungs to the blood until the partial pressures in the alveoli and blood are equal.
Dalton went on to further explain that the sum of partial pressures of all gases in a mixture equals the total pressure of the mixture as follows: