Explain Dalton's Law & How It Relates to Respiration

If you’ve heard of a therapeutic oxygen tent, you are seeing Dalton’s law of partial pressure put to work. Simply put, all of the gases in the Earth’s atmosphere exert individual pressures that add up to that 14.7 pounds per square inch (psi). Change the concentration of air’s gases, for example, making it more oxygen-rich, and you increase the pressure and amount of oxygen delivered to a patient with diminished lung capacity.

Dalton’s law of partial pressures states that the total pressure (Pt) of a gas mixture is equal to the sum of the partial pressures of the individual gases in the mixture. If a sphere contains 10 molecules of the gases X, Y and Z, each exerts a pressure proportional to the number of molecules of that gas that the sphere contains.

Dalton’s law may be stated as:

Pt = P1 + P2 + P3…

Atmospheric air is a mixture of several molecular components. Presuming the air is dry (such that water or H20 exerts no pressure), their concentrations are: Nitrogen (N2): 78.08 percent by volume Oxygen (O2): 20.95 percent Carbon dioxide (CO2): 0.03 percent Argon (Ar): 0.93 percent

Some trace components include neon, krypton, xenon, methane and nitrous oxide.

Those partial pressures may be expressed in barometric millimeters of mercury (mm Hg), which are as follows: Nitrogen: 593 mm Hg Oxygen: 159 mm Hg Argon: 7 mm Hg Carbon dioxide: 0.2 mm Hg

Within the lungs, total pressure may be expressed as:

Pt = P(CO2) + P(O2) + P(N2) + P(H20)

While the pressure of C02 is fairly negligible in the atmosphere, its concentrations are far higher within the lungs, as it is a product of respiration.

By increasing the percentage of any gas in air’s mixture, a higher partial pressure of that gas can be achieved, which is the basis of oxygen therapy. Dalton’s law is used in the practices of pulmonary physiology, ventilator care, medical gas administration, arterial blood gas and pulmonary pathophysiology, among other applications.

Individual gas pressures increase in hyperbaric conditions where the pressure is higher than that of the atmosphere (e.g., underwater and in hyperbaric chambers). Those pressures decrease in hypobaric conditions (e.g., at high altitudes), but remain in proportion to their concentration in the mixture.