Atmospheric pressure is the force exerted on every surface by the column of air stretching from the ground to the edge of space. This pressure is essentially the weight of all the air molecules in the Earth’s atmosphere being pulled down by gravity. At sea level, this force averages about 14.7 pounds per square inch (psi), establishing the physical conditions under which all terrestrial life evolved. This steady, external force is an indispensable component of several fundamental biological processes. Without this constant surrounding pressure, the mechanics of our breathing would fail, the liquid state of our body fluids would be compromised, and the delicate balance of gases dissolved in our blood would instantly unravel.
How Pressure Drives Breathing
The simple act of taking a breath is a direct consequence of the difference between the air pressure outside the body and the pressure inside the lungs. This mechanical process is governed by Boyle’s Law, which states that the pressure and volume of a gas are inversely related. To inhale, the body must actively create a pressure inside the lungs that is lower than the outside atmospheric pressure.
Inhalation begins when the diaphragm, a large dome-shaped muscle beneath the lungs, contracts and moves downward. This action, assisted by the muscles between the ribs, causes the volume of the chest cavity and the lungs to increase significantly. As the volume of the air inside the lungs expands, Boyle’s Law dictates that the internal pressure must immediately drop.
This temporary pressure differential creates a powerful gradient, as the internal lung pressure is now slightly negative compared to the higher atmospheric pressure outside. Air naturally flows from an area of higher pressure to an area of lower pressure, so the external air is effectively pushed into the respiratory passages and lungs. Exhalation is largely a passive process, occurring when the diaphragm relaxes, decreasing lung volume, which in turn raises the internal pressure and forces the air back out.
Preventing Body Fluids from Vaporizing
A primary role of air pressure is ensuring that the water comprising most of the body remains in a liquid state. This function is based on the thermodynamic principle that the boiling point of a liquid decreases as the surrounding pressure decreases. At sea level, water boils at 100° Celsius (212° Fahrenheit), safely above the human body temperature of approximately 37° Celsius (98.6° Fahrenheit).
If the surrounding pressure drops too low, the boiling point of water can fall to a temperature equal to body heat, a phenomenon known as ebullism. This threshold occurs at an altitude known as Armstrong’s Line, approximately 19,000 meters (61,000 feet) above sea level. At this low pressure, the vapor pressure of the body’s water content exceeds the ambient pressure, causing the water to turn to gas.
The immediate result is the rapid formation of water vapor bubbles within the bodily fluids, including the blood, saliva, and the moisture on the eyes. Although the skin and tissues prevent the body from simply exploding, the expansion of these bubbles would cause severe swelling and tissue damage. More dangerously, bubbles forming in the blood would stop circulation, leading to a condition called cardiac vapor lock. The constant force of atmospheric pressure is necessary to physically suppress the transition of water from liquid to gas, maintaining the structural and fluid integrity of all cells and tissues.
Stabilizing Gases Within the Bloodstream
Atmospheric pressure also determines the concentration of gases that can be physically dissolved within the liquid components of the body, particularly the blood. This relationship is described by Henry’s Law, which states that the amount of gas dissolved in a liquid is directly proportional to the partial pressure of that gas above the liquid. This law is primary for both the uptake of oxygen and the management of non-metabolic gases like nitrogen.
For oxygen, the external pressure ensures a sufficient amount dissolves into the blood plasma to be transported efficiently throughout the body. For nitrogen, which makes up about 78% of the air we breathe but is not metabolized by the body, the pressure keeps it dissolved and stable in the blood and tissues. This stability is immediately jeopardized by rapid pressure changes, such as those experienced by deep-sea divers ascending too quickly.
If the external pressure drops too fast, the solubility of the dissolved nitrogen decreases rapidly, causing it to come out of solution and form bubbles. This is the mechanism behind decompression sickness, commonly known as “the bends,” where nitrogen bubbles obstruct blood flow and cause debilitating pain or tissue damage. Stable atmospheric pressure is necessary to maintain the delicate equilibrium of dissolved gases, preventing them from physically bubbling out and disrupting physiological function.