The air surrounding us is a mixture of gases, primarily nitrogen and oxygen. Partial pressure is the specific force exerted by any one individual gas within that mixture. If that single gas were to occupy the entire volume by itself, its partial pressure is the pressure it would exert. Partial pressure is the true measure of a gas’s activity and its ability to participate in physical and biological processes.
The atmosphere’s total pressure is the sum of all the individual partial pressures. At sea level, the air is composed of roughly 78% nitrogen and 21% oxygen. Each gas contributes a distinct portion of the total pressure. Understanding this contribution is important for fields like respiratory physiology, where the movement of gases into and out of the body is governed by these individual pressures.
Understanding Dalton’s Law of Partial Pressures
The behavior of gas mixtures is governed by Dalton’s Law of Partial Pressures. This law states that the total pressure exerted by a mixture of non-reacting gases equals the sum of the partial pressures of the individual component gases. Every gas in the mixture acts independently of the others.
The independence of gas molecules is why this law holds true. Because gas molecules are widely spaced and experience minimal attractive or repulsive forces, the presence of one gas type does not interfere with the pressure contribution of another. The total force measured is the cumulative result of all those individual collisions.
To determine the total pressure of a gas mixture, one only needs to measure the partial pressure of each gas and add them together. This concept allows analysis of complex gas systems by breaking them down into simpler, single-component pressures. The law applies widely, from chemical manufacturing processes to understanding the dynamics of the Earth’s atmosphere.
Calculating Partial Pressure and Common Units
Partial pressure is determined mathematically by relating the concentration of a gas to the total pressure of the mixture. The partial pressure of a specific gas equals the total pressure of the mixture multiplied by the fractional concentration of that gas. The fractional concentration is the proportion of that gas, often expressed as a decimal (e.g., 0.21 for 21% oxygen).
The formula is \(P_{gas} = P_{total} \times F_{gas}\), where \(P_{gas}\) is the partial pressure, \(P_{total}\) is the total pressure, and \(F_{gas}\) is the fractional concentration. At sea level, standard atmospheric pressure is approximately 760 millimeters of mercury (mmHg). Since oxygen makes up about 21% of the air, its partial pressure (\(P_{O_2}\)) is calculated as \(760 \text{ mmHg} \times 0.21\), equaling about 160 mmHg.
Partial pressure is measured in various units. The most common are millimeters of mercury (mmHg), also known as torr, atmospheres (atm), and kilopascals (kPa). For instance, 160 mmHg for oxygen at sea level corresponds to about 21.3 kPa. While the choice of unit depends on the context, mmHg is frequently used in medicine to discuss blood gases.
The Critical Role of Partial Pressure in Respiration
In biological systems, partial pressure drives the exchange of gases necessary for life. Gas molecules move spontaneously from an area of higher partial pressure to an area of lower partial pressure, known as moving down a pressure gradient. This principle dictates how oxygen enters the bloodstream and how carbon dioxide is removed.
When air is inhaled, it travels into the alveoli, the tiny air sacs in the lungs. In the alveoli, the partial pressure of oxygen (\(P_{O_2}\)) is high (around 104 mmHg). Deoxygenated blood arriving from the tissues has a much lower \(P_{O_2}\) (about 40 mmHg). This significant pressure difference creates a steep gradient, causing oxygen molecules to rapidly diffuse into the blood.
A reverse gradient exists for carbon dioxide (\(P_{CO_2}\)). Carbon dioxide-rich blood returning from the tissues has a high \(P_{CO_2}\) (approximately 45 mmHg). Conversely, the \(P_{CO_2}\) in the alveoli is lower (around 40 mmHg). This smaller, yet sufficient, gradient causes carbon dioxide to diffuse out of the blood and into the alveoli for exhalation. Without these partial pressure differences, gas exchange would halt.
Changes in atmospheric pressure affect the body. At high altitudes, the total atmospheric pressure decreases significantly. Although the air still contains 21% oxygen, the resulting lower total pressure means the partial pressure of oxygen is also lower. This reduced \(P_{O_2}\) lessens the driving force for oxygen diffusion into the blood, which can lead to hypobaric hypoxia, or altitude sickness.