P1, P2 ............ Pn = partial pressure of each of the constituent gases — Page 45, Lesson 68

I want to walk you through a set of gas laws that are absolutely fundamental to understanding how your body handles pressure changes during flight. These laws explain everything from why you need supplemental oxygen at altitude to why divers can't fly immediately after a dive. Let's start with the notation you'll see throughout. When I write P1, P2, up to Pn, each of those represents the partial pressure of one of the constituent gases in a mixture. So if you have a mixture of oxygen, nitrogen, and carbon dioxide, P1 might be the partial pressure of oxygen, P2 of nitrogen, and so on. Partial pressure simply means the pressure that one specific gas would exert if it alone occupied the entire volume. Now, Henry's Law. It states that at equilibrium, the amount of gas dissolved in a liquid is proportional to the gas pressure. Let me unpack that. If you have a liquid like your blood or body tissues, and you expose it to a gas under pressure, the gas will dissolve into that liquid. The higher the pressure of that gas, the more of it will dissolve. This is the law that explains decompression sickness, what divers call "the bends." When a diver is deep underwater under high pressure, nitrogen dissolves into their body tissues. If they surface too quickly, that dissolved nitrogen comes out of solution and forms bubbles in their blood and tissues, causing severe pain and potentially fatal damage. The same principle applies to pilots who fly shortly after diving. Next, Fick's Law. This one governs how gases actually move across membranes in your body. It states that the rate of gas transfer is proportional to the area of the tissue and the difference between the partial pressures of the gas on the two sides, and inversely proportional to the thickness of the tissue. In plain terms, gas moves from where it has a higher partial pressure to where it has a lower partial pressure. How fast it moves depends on three things: how big the surface area is, how big the pressure difference is, and how thin the tissue is. This is the law that describes diffusion of oxygen from your lungs into your blood and diffusion of carbon dioxide from your blood back into your lungs to be exhaled. It also governs gas exchange at the cellular level. Now, Charles' Law. This one deals with temperature and volume. It states that the volume of a fixed mass of gas held at a constant pressure varies directly with the absolute temperature. If you heat a gas, it expands. If you cool it, it contracts, as long as the pressure stays the same. The mathematical expression is V1 over V2 equals T1 over T2. Where V1 is the initial volume, V2 is the final volume, T1 is the initial absolute temperature, and T2 is the final absolute temperature. And absolute temperature means you take the temperature in degrees Celsius and add 273. So T1 equals t1 plus 273, and T2 equals t2 plus 273. This is critical because the air in your aircraft cabin and in your body cavities expands and contracts with temperature changes. Then we have the Combined Gas Law. This brings pressure, volume, and temperature all together. It states that the product of the pressure and the volume of a quantity of gas divided by its absolute temperature is a constant. Mathematically, PV over T equals K, a constant. This means if you change any one of those three variables, at least one of the others must change to keep the equation balanced. This law governs the behaviour of gases in your aircraft's pressurisation system, in your fuel tanks, and in the air spaces inside your body. Now let's look at partial pressure more closely, specifically in the context of Dalton's Law and the atmosphere. You already know that total atmospheric pressure decreases as altitude increases. What's important here is that the proportion of oxygen in the air remains constant at about 21 percent. So if total pressure drops and the percentage of oxygen stays the same, the partial pressure of oxygen must also drop. That's the key point. In aviation, when we deal with pressures at various altitudes, we use a different unit than the hectopascals or millibars you might see in Meteorology or Instruments. Here, the unit of measurement is the millimetre of mercury, written as mm Hg. At sea level, the standard pressure is 760 mm Hg. Since oxygen makes up 21 percent of the atmosphere, the partial pressure of oxygen at sea level is twenty-one hundredths of 760, which gives you 160 mm Hg. As you climb, that partial pressure drops, and when it drops low enough, your body can't get enough oxygen into your bloodstream to function properly. That's why we have oxygen systems and pressurised cabins. That figure shows the air passages in the lungs and the composition of the standard atmosphere, which ties directly into how these gas laws apply to your respiratory system.