October 25, 1999, Aberdeen, South Dakota – After running out of fuel, a Learjet Model 35, tail # N47BA spirals to the ground and crashes, killing all occupants of the plane. – What had happened? The plane, operated by Sunjet Aviation had been dispatched that morning in Orlando for a flight to Dallas and cleared by ATC for FL 390. Shortly thereafter all communications seized and could not be re-established. The airplane was intercepted by several U.S. Air Force and Air National Guard aircraft as it proceeded northwestbound. The military pilots in a position to observe the accident airplane at close range stated that the forward windshields of the Learjet seemed to be frosted or covered with condensation. The safety board concluded later that a loss of pressurization led to crew incapacitation before the pilots donned their oxygen masks. This is not the first – and probably not the last – accident caused by lack of oxygen. It just became famous because of the passenger on board: Golf legend Payne Stewart. Why is it, that a sudden drop in pressure can have so devastating consequences? I f we can hold our breath under normal circumstances for a minute easily – why can we not do it at altitude? – Instead we lose consciousness within seconds? When I taught students how to scuba dive, I regularly talked about the risk of “shallow water blackout”, a condition where a breath hold diver after an extended time at depth ascends and loses consciousness at a shallow depth before the diver can reach the surface – often leading to death by drowning. What is the common thread? What is it, that the “Human Factors in Aviation” book doesn’t tell us? Unfortunately we have to go back to some rules of physics in order to understand how our body deals with pressure change. But it’s worth it because we will get a good grasp of what to expect when we venture into thin air. First let’s define pressure for the purpose of this discussion. Let’s not talk about pascals, bars or inches of mercury. The simple term “atmosphere” (atm) is sufficient for our purposes. The variation in pressure that we encounter daily – which may amount to approximately +/- 5% and for which we compensate on our altimeters, is not relevant for today’s topic. The fact that we live in an environment of 1 atm of pressure as long as our feet are on the ground is sufficient to understand what pressure change does to us. Let me briefly mention Boyle’s Law which states what we intuitively know: The volume of a given mass of gas is inversely proportional to the absolute pressure (assuming no temperature change). In other words, an air mass at 0.5 atm pressure – as we will find at about 18,000’ altitude - will take up twice the volume of the air mass on the ground (our 1 atm environment). Twice the volume for the same air mass, i.e. for the same amount of gas molecules means there is less air to breathe at altitude – but we knew that already! Now we are getting more serious: The Law of Partial Pressures, observed by John Dalton in 1801 states that the total pressure exerted by a mixture of gases is equal to the sum of the pressure of each of the different gases making up the mixture – each gas acting as if it alone were present and occupied the total volume. As our air is comprised of about 21% oxygen and 81% of nitrogen (for our calculation today we simplify this to a 20/80 ration), Dalton’s Law states that the 1 atm pressure that we are exposed to on the ground, consists of a partial pressure (Pp,O2) of oxygen of 0.2 atm and a partial pressure of nitrogen (PpN2) of 0.8 atm totaling 1.0 atm of total pressure for the gas mix, which we call air. PTotal = PpO2 + PpN2 = 0.2 atm + 0.8 atm = 1.0 atm If the pressure increases, so do the partial pressures – this is a typical scenario in scuba diving. At a water depth of 10 metres, the total pressure doubles and air at this pressure contains oxygen at a partial pressure of 2 x 0.2atm =0.4 atm and a PpN2 of 2 x 0.8atm = 1.6atm. But since we are going up in the earth’s atmosphere, as pilots, we are only confronted with a decrease of pressure as we climb – along with a decrease of the partial pressure of the gases in the air. By the way: For the purpose of our discussion, we can assume that the actual mix between oxygen and nitrogen does not change with altitude – only the density and the number of gas molecules per volume decreases. The graph in figure 1, comprised of data from different sources incl. NOAA and others shows the pressure decrease with altitude and it applies equally to the total air pressure and – since the gas mix remains constant – it applies as well to the partial pressures. We will specifically focus on the partial pressure of oxygen. The graph shows that the air pressure drops to 0.5 atm at an altitude of approximately 18,000 feet. What will the partial pressure of oxygen be at that altitude? – You figured it out: It’s 0.5 x 0.2 atm = 0.1 atm! The decrease in atmospheric pressure is not a linear function – the air pressure gets cut in half again (0.25 atm) at about 33,000 feet. At 70,000 (we would have to trade in our C152 for a SR 71 or Concorde) the remaining air pressure is less than 0.1 atm. If all of the above was rather boring and you regret the time that you wasted on this article, bear with me for another minute, because here comes the reward for your patience: Scientists may argue that life is much more complicated, but the simplified fact is: The human body’s need for oxygen is driven by the oxygen partial pressure (PpO2) it encounters! On the ground the body is exposed to 0.21 atm, but the human body has an “operating range” from about 0.14 to more than 2.0 atm. Above this range (or after long term exposure) oxygen can become toxic to the body, below this range, oxygen deprivation will lead to reduced performance, unconsciousness, eventually death. Although the tolerance to low oxygen levels (partial pressures) may vary, we can assume that below an oxygen partial pressure of about 0.15 atm symptoms may be encountered. What does that mean? Let’s look at some examples:
- I once flew with a student at 13,000’ and after 10minutes (remember: we are allowed to stay there without supplementary oxygen for 30 minutes) he showed symptoms of anxiety and only the immediate supply of oxygen calmed him down). –What is the PpO2 at 13,000’? (Approx. 0.12 atm)
- The crew in Glen Payne’s Citation encountered a depressurization at 39,000’ without wearing oxygen masks. Due to the rapid pressure loss, they were suddenly exposed to an air pressure of about 0.2 atm - the PpO2 in the pilot’s body would immediately drop to 0.04 atm, thereby rendering him unconscious in just a few seconds.
- Had the pilots recognized the depressurization in time and donned their masks they would be breathing pure oxygen at the same surrounding pressure of 0.2 atm. The oxygen partial pressure would have been 1.0 x 0.2 atm = 0.2 atm and they would have been fine!
- The SR 71 pilot flying at 70,000’ altitude has a pressurization loss and encounters a surrounding pressure of less than 0.1 atm. Even breathing pure oxygen only provides him with a PpO2 of less than0.1 atm – not enough to survive. That’s one reason why pilots in that environment wear suits that will provide a pressurized environment even in case of a malfunction of the pressure hull.
- The free diver who encounters “shallow water blackout” hyperventilated before the dive to lower the urge for breathing (which is not triggered by lack of oxygen but by a high concentration of CO2). He started his dive with a PpO2 of 0.2 atm, doubling it by diving to a depth of 10 metres. As he stayed at 10 metres for a period of time during which oxygen was metabolized, the PpO2 may have been cut in half, from 0.4 to 0.2 atm. As the diver ascends and comes close to the surface the water pressure decreases and the PpO2 approaches 0.1 atm – not enough to stay conscious and no will power in the world can change that!