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Henry’s Law states that the solubility of a liquid is directly proportional to the partial pressure of the gas above the liquid. The greater the pressure, the larger the quantity of gas that will be forced into the liquid. In other words, at higher pressures, our bodies will absorb more gases.
How is Henry’s Law important to divers? Atmospheric pressure increases as the diver descends, and the pressure forces some of the gas, nitrogen, for example, to dissolve into the blood and tissues. Some of the nitrogen becomes liquid in our body tissues. As I descend into deeper water, the pressure continues to increase, and so does the amount of dissolved gas in my body. But the nitrogen isn’t a problem in itself. The air we breathe is basically comprised of 78% nitrogen, 21% oxygen, and some trace amounts of other gases. Despite the fact that most of the air we breathe is nitrogen, the only gas our body actually consumes is oxygen, which gets used for metabolism. The nitrogen, however, is physiologically inert. In other words, we breathe in nitrogen and we exhale it and is hasn’t changed. The nitrogen content of air in our scuba tanks becomes very significant, however, when diving at greater depths, for longer periods of time, because it stays in our body and begins to build up in our blood and tissues. As we begin to ascend, the only safe way our body tissues can get rid of the buildup of nitrogen is for the nitrogen to come out of solution, back to a gas, return to our lungs and be exhaled. Remember, too, this gas is at the same pressure as that of the atmosphere we are in, so when we are drawing it into our lungs, it is in a much more concentrated form. At 99’, we are breathing in four times as much gas as if we were at the surface.
Nitrogen becomes a problem if the diver ascends too quickly. As the diver ascends, pressure drops, and, if ascending very fast, the pressure drops very fast, and the nitrogen that has dissolved into our tissues isn’t given time to be returned to the lungs, and exhaled. Rather, it comes out of solution from our blood and tissues, in the form of gas bubbles, which then expand and become trapped in various parts of our bodies. This is what is known as decompression sickness. Also, oxygen toxicity is a potential danger to divers, if oxygen is being breathed in high concentrations.
What is oxygen toxicity? The oxygen we breathe on land normally only dissolves into our blood in small amounts. Well, as we know, when diving, any given concentration of gas comes under higher pressure than atmospheric, which means we’re breathing a greater amount of it while diving. Greater than normal amounts of inhaled oxygen becomes poison. Fortunately, oxygen toxicity, is a not a major potential hazard for recreational diving, because we normally don’t descend beyond a hundred feet or so, and we’re not really down at greater depths long enough for the oxygen to build up to dangerous levels.
What we are concerned with are the effects of increased concentrations of dissolved gases in our blood and tissues, and how those increased concentrations affect us while diving.
DALTON’S LAW OF PARTIAL PRESSURES
Dalton's law states that the total pressure exerted by a mixture of gases is equal to the sum of the partial pressures of each of the individual gases present. In other words, one gas in a mixture will not compress more or less than another.
How is Dalton’s Law important to divers? It’s important in describing the partial pressure of gases at depth, as it is exposure to these gas partial pressures that that allow us to predict, and avoid, dangers such as nitrogen narcosis, oxygen toxicity, and decompression sickness. What we’re concerned with here is how the gases we breathe, while diving, affect us during these pressure changes. The concentration, or partial pressure, of oxygen, for example, increases with depth. As the concentration of a gas increases, the physiological and psychological effects of that gas may increase or change.
So, air in our tanks is roughly a mix of 21% oxygen, 78% nitrogen. As the diver descends, the pressure "concentrates" the gases, and the partial pressure of each gas increases. It is when these gases become highly concentrated, with increased pressure, that they can become dangerous, depending on how long they are breathed and at what depth. On land, that 21% oxygen we’re breathing translates to a partial pressure of about 3psi. We figure this out by multiplying .21 times 14.7. The 14.7psi being the weight of the atmosphere. However, at 99’, 4 atm, the partial pressure of 21% oxygen becomes 12.3psi.
So, two factors determine the partial pressure of a gas in scuba diving – the percentage of the gas in the breathing mixture and the depth (and therefore the ambient pressure) at which a diver breathes the gas.
Mathematically, Dalton’s Law can be stated as follows: Ptotal=P1+P2+P3+P4, etc., (N2, O2, Ar, He)
A 2.0 L container is pressurized with 0.25 atm of oxygen and 0.60 atm of nitrogen. What is the total pressure inside the container?
Ptotal=PO2+PN2=0.25+0.60=0.85atm. The total pressure inside the container is 0.85 atm.
Let’s do one more. What is the partial pressure of 21% oxygen, in 60 feet of sea water? We multiply the percentage of the oxygen, by the ambient pressure. That’s it.
First we need to know the ambient pressure at 60 feet. Just standing on land, or sitting on the dive boat, we are subjected 14.7 pounds of pressure per square inch from the atmosphere. No matter what depth we are at under water, we need to factor in that 14.7 psi to the amount of pressure we are experiencing at depth, in order to know the ambient pressure. So, we know that salt water exerts .445 pounds of pressure per foot of depth, so the water pressure at 60 feet depth would be .445 x 60, or 26.7 psi. To this we add the atmospheric pressure of 14.7 to get our ambient pressure of 41.4, which is about 3atm. We can now figure out the partial pressure of the oxygen, by simply multiplying 41.4 times the 21% oxygen, to find that the partial pressure of oxygen, at a depth of 60 feet in sea water is 8.80psi.
Boyle’s Law states that pressure and volume, are inversely proportional. As pressure increases, gas volume decreases, and as pressure decreases, gas volume increases.
How is Boyle’s Law important to divers? Well, one of the first things we learn as scuba divers is never hold your breath, right? Because if I took a breath from a scuba tank at, say, 30 feet, 2 atm, and held my breath as I ascended back to the surface, the volume of the gas in my lungs would double, causing them to rupture.
We know that as the diver descends, the pressure increases at a rate of one atmosphere (ATA) every 33 feet, and 14.7lbs psi, every 33 feet.
Bone, muscle, blood and solid organs such as the kidney, heart, and liver are all non-compressible and therefore unaffected by water pressure. The compressible areas of our bodies, such as the lungs, middle ears, sinuses, nasal passages, interior of hollow organs (stomach and intestines), and any air pockets you may not know about (e.g., a tooth cavity), are compressible because they contain some air.
When I inhale air from a tank, the air that enters my lungs is at ambient pressure. Now, If I inhale from the tank on the surface, the pressure in my lungs will be at 1 atm. The lungs of an average adult can hold about 6 litres of air. At 66 feet, 3atm, if I took a breath of air and held it as I ascended to the surface, my lungs would attempt to expand to three times their initial volume, to 18 liters, and my lungs would rupture before I even reached the surface.
Also, the deeper I go, the more concentrated the air is that is coming out of my tank, so I’m actually consuming air faster. At 33 feet, 2 atm, I’m breathing in twice the amount of air that I would be at the surface. That is why, for example, (do the illustration) if I needed to cross an uninteresting sand bottom, to get to a drop-off for a wall dive, and the top of the wall is at a depth of say 66’, I wouldn’t go down to 66’ and then swim across. Instead, I would conserve my air supply by staying at 15’, swim across till I’m above the wall, and then dive down. The average beginner diver, in calm waters, runs a tank close to empty in around 1 hour, at 33’, 2atm, but will empty the tank in just a few minutes at 131’, 5atm.
GAY-LUSSAC’S LAW Pressure Proportional to Temperature
Gay-Lussac’s Law states when the pressure increases, the temperature increases. When the pressure decreases, the temperature decreases.
How is Gay-Lussac’s Law important to divers? It is most important in relation to the amount of breathable air in a tank. Scuba tanks made out of aluminum typically have a rated fill pressure of 3,000 psi. In filling a scuba tank, the faster the gas molecules are forced into the tank, the more heat is generated. If a tank is filled quickly, to 3000psi, its temperature can rise to as much as 150° F. However, when it cools, and its temperature and pressure drop, it may only have 2600psi of air in it.
Mathematically, Gay-Lussac’s Law is expressed as: P1 divided by T1 equals P2 divided by T2
Since all gas laws use absolute temperatures, the temperatures that we use in our calculations need to be converted to an absolute temperature of Kelvin by adding 273. So, 25°C, becomes 298 Kelvin. We convert to Kelvin because we don’t want to end up with any negative numbers, and Kelvin is the only temperature scale that has only positive numbers.