The composition of the atmosphere is a mixture of gases and is subject to the same laws of physics as you would see in a physics lab. If you are to under- stand how and why these gases affect the body, you must also have a basic understanding of the classic gas laws.
Boyle’s Law
This law states that the volume of a gas is inversely proportional to the pres- sure (temperature remaining constant). This applies to all gases. V1/V2=P2/P1 (V1 is the initial volume of the gas, V2 is the final volume, P1 is the initial pressure on the gas volume, and P2 is the final pressure). In other words, if the pressure of the gas decreases with the temperature unchanging, then its volume increases and vice versa (Fig. 3-2).
In dealing with gas expansion in the body, a correction must be made for the ever-present water vapor; therefore, the formula now becomes: V1/V2 = (P2–47 mm Hg)/(P1–46 mm Hg) (Fig. 3-3). Water-vapor pressure at body temperature is 47 mm Hg. Such characteristics applied to the body explain
Figure 3-2
Figure 3-2 Comparative volumes of dry gases at increasing altitudes and decreasing pressure.
Gas laws 37
38 The atmosphere
the expansion of gases trapped within such moist areas as the middle ears, sinuses, stomach, and intestines. These are all actual or potential cavities within which moist air is present and can become trapped and expand like any other gas; hence, the physiological topic of “trapped gases,” which will be discussed in Chapter 5 regarding altitude physiology.
Charles’ Law
Charles’ Law states that the volume of gas is directly proportional to the tem- perature (pressure remaining constant). This applies to all gases. This law has no direct physiological significance because body temperature remains fairly constant. It does, however, explain the fact that pressure within sup- plemental oxygen containers will decrease if the ambient temperature sur- rounding the storage container decreases, even when no oxygen has been used, such as at altitude.
Dalton’s Law
Since the atmosphere is a mixture of gases, and each gas has its own pres- sure at any given temperature within a given volume, it is important to also be familiar with the physics of the combined pressures. Dalton’s Law states that the total pressure of a gas mixture is the sum of the individual pressure (also called partial pressure) that each gas would exert if it alone occupied the whole volume.
Figure 3-3
Figure 3-3 The effect of water on gas expansion.
Or expressed mathematically: PT=P1+P2+Ps+Pn; where PT is the total pres- sure of the mixture of gases and the P value is the partial pressure of each gas, which is determined by multiplying the percentage of the individual gas times the total pressure. Each kind of gas exerts its own pressure, depending on the percentage of that gas in the mixture; thus, even though the percentage of oxy- gen in the atmosphere is constant (about 21 percent), its partial pressure will decrease proportionately as atmospheric pressure (expressed as barometric pressure) decreases. Water vapor is considered a gas in the atmosphere.
The body is affected by the pressure of gases available. In other words, as we ascend, the percentage of each gas in the atmosphere remains the same, but there are fewer molecules at less pressure that the body can use (or get rid of). This decrease of available molecules of oxygen at a pressure required to pass to a blood cell, for example, is what leads to hypoxia in the body.
Henry’s Law
Henry’s Law states that the amount (or weight absorbed) of gas in solution, not chemically combined, varies directly with the partial pressure of that gas over the solution. P1A2=P2A1, where the A value is the amount of gas in solu- tion initially. In other words, when the pressure of a gas over a certain liquid decreases, the amount of that gas dissolved in the liquid will also decrease (and vice versa); therefore, when equilibrium is attained, the dissolved gas tension will equal the partial pressure of the gas in the atmosphere to which a solution is exposed. As the pressure “falls” (during ascent), the amount of gas that can be held in solution is reduced.
This is demonstrated every time you open a carbonated beverage. Once the seal is broken and the gas that has been under pressure escapes, gas that has been dissolved within the beverage begins to escape by forming bubbles on the side and rising to the surface (Fig. 3-4).
In the human body, however, additional factors influence and modify the process of gas uptake and elimination, such as the varying types of fluids in the body, the circulation rate and volume of the blood, and the amount of hemoglobin (where the gas is chemically attached), among others. These will all be discussed in subsequent chapters.
Graham’s Law
Graham’s law of gaseous diffusion states that a gas of high pressure exerts a force toward a region of lower pressure, and that if an existing membrane separating these regions of unequal pressure is permeable or semiperme- able, the gas of higher pressure will pass (or diffuse) through the membrane into the region of lower pressure. This process, which occurs in milliseconds, continues until the unequal regions are nearly equal in pressure.
For example, this law explains the transfer of oxygen from one part of the body to another as it passes within blood vessels and then through the vessel wall to the adjacent cell. Each gas behaves independently of the other gases in a solution and might move in opposite directions to another gas of differing partial pressures (Fig. 3-5). Figure 3-6 summarizes the gas laws.
Gas laws 39
40 The atmosphere Figure 3-4
Figure 3-4 The amount of gas dissolved in solution is directly proportional to the pres- sure of the gas over the solution.
Figure 3-5
Figure 3-5 Gases diffuse across membranes from a higher partial pressure to lower and vice versa.