Introduction to the physiology of Diving - part 2

Partial Pressure To understand the physiological ramifications of breathing various gas mixtures under pressure, it is useful to understand the concept of partial pressure. The partial pressure of a particular gas constituent in a gas mixture is a representation of the portion of the total pressure of the gas mixture exerted by the particular constituent. If you add up all the partial pressures of all the different components of a gas mixture, their total would be equal to the total pressure of the mixture. As confusing as this may sound, partial pressures are actually quite easy to calculate: all you need to know is the fraction of each gas constituent in the mixture, and the total pressure of the gas mixture. For example, consider a person breathing air (a gas mixture containing approximately 80% nitrogen and 20% oxygen) at sea level. As discussed earlier, the ambient pressure at the sea surface is 1 ATA. Therefore, the pressure of the air which the person inspires is also 1 ATA. To get the partial pressure of nitrogen in the inspired air, simply multiply the fraction of nitrogen in the breathing mixture (80%) by the total pressure (1 ATA), and you calculate a nitrogen partial pressure of 0.8 ATA. Similarly, multiplying 20% oxygen times 1 ATA results in an oxygen partial pressure of 0.2 ATA. Now consider what happens when that same person descends to a depth of 99 feet (30 meters), where the ambient pressure is 4 ATA. In order for that person to be able to breathe at all, the inspired air pressure must be the same as the ambient pressure. Therefore, the inspired partial pressure of nitrogen is 80% times 4 ATA, or 3.2 ATA. The oxygen partial pressure is 20% times 4 ATA, or 0.8 ATA. At 99 feet (30 meters), the ambient pressure is four times greater than it is at the surface, and the partial pressures of each of the gases is also four times greater (although the percentages of each gas are the same in both cases). As discussed early, the gas molecules are more closely packed when under pressure; at a depth of 99 feet (30 meters), there are four times as many gas molecules (both nitrogen and oxygen) in a lung-full of air as there are at the surface. An easy way to think of partial pressures of gases is that the partial pressure represents an absolute concentration of that gas, regardless of depth or pressure. If a person breathed a gas mixture containing 80% oxygen at the surface, the oxygen partial pressure would be 0.8 ATA, which is exactly the same partial pressure of oxygen when breathing air at a depth of 99 feet (30 meters). In both cases (80% oxygen at the surface and air at 99 feet/30 meters), the concentration of oxygen molecules in the lungs (i.e., the total number of oxygen molecules in the lungs on each inhaled breath) is the same. Just a word on notation: in their gaseous forms, both oxygen and nitrogen are binary molecules; that is, they are bound in pairs of atoms. An oxygen gas molecule consists of two oxygen atoms bound together, and a nitrogen gas molecule consists of two nitrogen atoms bound together. The notation for oxygen is the letter "O", so oxygen gas is referred to as "O2"; the subscript "2" indicating two atoms of oxygen. Similarly, nitrogen gas is referred to as "N2", and carbon dioxide as "CO2". When discussing partial pressures of gases, the gas notation is usually prefaced by a capital "P". Thus, "oxygen partial pressure" is written as "PO2", and "nitrogen partial pressure" is written as "PN2". Khao Lak IDC - Instructor Training Center, Thailand

Henry's Law Henry's Law states that "The amount of any given gas that will dissolve in a liquid at a given temperature is a function of the partial pressure of that gas in contact with the liquid..." What this means for divers is that gas molecules will dissolve into the blood in proportion to the partial pressure of that gas in the lungs (as "warm-blooded" creatures, our core body temperature remains relatively constant). In the diagram at right, the top figure (1) represents a close-up of the interface between the lungs and the blood and tissues of a diver. At sea level, the dissolved gases in the blood and tissues are in proportion to the partial pressures of the gases in the person's lungs at the surface. As the diver descends underwater, the ambient pressure increases, and therefore the pressure of the gas inside the lungs increases correspondingly. Because the partial pressures of the gases in the lungs are now greater than the dissolved partial pressures of these gases in the blood in tissues, gas molecules begin to move from the lungs into the blood and tissues (represented by the blue and red arrows in the middle figure, 2). Eventually, the concentration of the dissolved gases in the blood and tissues will be proportional to the the partial pressures in the breathing gas (i.e., a state of equilibrium). The physiological complexities of "Ambient Pressure Diving" are a direct result of the effects of these increased dissolved concentrations of gases in the blood and tissues, and how those increased concentrations affect the way our bodies work. Khao Lak IDC - Instructor Training Center, Thailand