THE FOUR GAS LAWS

Boyle's law: P1V1 = P2V2. As gases get compressed their volume shrinks. As gases decompress their volume expands. Taking a 1 cubic foot balloon down in the ocean would cause it to shrink to 1/2 its size at 33'. Breathing at depth, ascending, and holding your breath may cause serious lung over-expansion resulting in arterial gas embolism, etc.

Henry's law: A gas over a liquid will flow into the liquid until the pressure is equal. If you covered a lake with neon, the atoms of that gas would pass though the surface of the water until the neon in the water had the same pressure as what is left in the air.

Charles's law - When you heat a gas in a container that is closed the pressure of the gas will rise. Likewise, if you cool the gas the pressure will drop. Likewise, if you decrease the pressure on a gas it will cool. If you take your scuba tank to have it filled and they fill it fast, the temperature will rise. As time goes by the air and tank will cool down. When that happens the pressure in the tank will drop. So, pressure, temperature, and  changing gas volume are all related.

Dalton's law: When you have a mixture of gases, each gas contributes to the total pressure. If you have a container of air that is at 2 atmospheres, 21% of the 2 atmospheres (0.42 atm.) is due to the oxygen, and 78% (1.56 atm) is due to the nitrogen. Both gases exert their own pressure and the two are combined to find the total pressure.

I found an article in the January 2005 scuba magazine Dive Training entitled Phiddling With Physics: Understanding the Gas Laws. It was so well written by Alex Brylske, the Senior Editor of the magazine, that I ask for and received his permission to include it here:

Phiddling With Physics: Understanding the Gas Laws

By Alex Brylske

While most scuba instructors are careful to avoid using the term directly, a good many of the concepts we must understand to dive safely involve the dreaded p-word, physics. We almost never utter this term because, quite frankly, it scare the daylights out of most folks. After all, physics is that obtuse and impossibly complex subject that only people like Albert Einstein or Steven Hawking can comprehend. This fact was perhaps reinforced by the F you received when, by some misguided advice of a guidance counselor, you took the class in high school.

Of course, as often happens with science, misconceptions abound, and reality is sometimes a far cry from perception. So is the case with physics. Contrary to what you might believe, it was not invented to torture under-achieving high school students, or confuse common sense. At its root, physics is really the science explaining the most familiar experiences in daily life. The dictionary defines it as "the scientific study of matter, energy, force, motion, and the way they relate to each other." In a way, physics is the science of everything. But understanding what you need to know about the physics of diving is as easy as pie-or rather, pi.

Does it Matter?

The raw material of physics is matter, which is the ultimate catch-all word for anything with substance-virtually everything. Unless you snoozed through junior high school science, you probably already know that all matter is made up of elements that can be combined into compounds. At the tiniest level of any practical consequence, matter consists of molecules.

Furthermore, matter exists in three states-solid, liquid and gas. At relatively low temperatures, molecules assume the rigid state that we term solid. Although the molecules comprising a solid vibrate, they remain in fixed positions. As temperature increases, the fixed structure of the molecules disappears, and the matter "melts" into a liquid state. In a liquid, the molecules vibrate faster than in a solid, so they aren't held together rigidly. A good way to conceptualize a liquid state is to think of the rubber ball pits common at children's playgrounds like, for example, at a MacDonald's Playland. While the balls are close together, they easily slip and slide past each other. So, a child playing in a ball pit is somewhat like someone swimming in the world's most viscous fluid. At still higher temperatures the molecules move further apart, like the ping-pong balls you see in pneumatic lottery machines. Finally, they break free from the surface of the liquid, and become a gas.

For this discussion, we'll consider all solids and liquids incompressible. But gases, because of the extreme relative distance between their molecules, are highly compressible. As divers, we deal in an environment with rapidly changing pressure, so this is of enormous consequence to us.

Air is a gas comprised mainly of nitrogen, oxygen, argon and carbon dioxide. So, although it's invisible, air and all other gases have substance; they are matter. Although this idea had been around since the days of the ancient Greeks, it was left to the famous seventeenth century astronomer Galileo to prove it. (Frankly, when you think about it, the fact that air has substance is really a matter of common sense; it's evident every time winds blow.) Later, Galileo's secretary, a mathematician named Evangelista Torricelli, showed that the atmosphere exerted enough pressure to hold up 30 inches/76 centimeters of mercury within a vacuum tube. A French mathematician named Blaise Pascal proved that the weight of the atmosphere was equal to the pressure exerted by 33 feet/10 meters of sea water.

Boyling the Air

The granddaddy of gas experts was a renowned English scientist (No! No! According to a 3/21/14 email from Leigh de Paor in Wexford, Ireland Boyle was born in 1627 in the county of Waterford in Ireland.) named Robert Boyle (after he got famous, it was Sir Robert Boyle). Building on Toricelli's work, he discovered that when they are exposed to either increased or decreased pressure, gases' behavior are highly predictable. He expressed his prediction mathematically in what has come to be known as Boyle's Law. Without all the algebra and scientific mumbo-jumbo, Boyle's Law essentially says that if you increase the pressure on a gas-filled object-like a balloon-the volume gets smaller. Squeeze a balloon and it gets smaller. Obviously, this stuff isn't rocket science. Of course, Boyle's Law acts in the exact opposite way when the pressure is reduced-like when a diver ascends. According to Boyle's Law, if we release a sealed balloon at depth, it will expand as it ascends; and if it exceeds its maxim volume, it will burst. So, too, when reducing the pressure on ascent, our lung volume increases. If we do nothing about it-like exhale-we risk lung injury due to overexpansion.

Ironically, while Boyle discovered his "Law," he could never explain what was actually happening. That took another scientist, the Italian Daniel Bernoulli, who explained that as gas, molecules collide with the container walls (remember the ping-pong balls). The more they're agitated, the more collisions occur and the greater the force. Bernoulli concluded that when the density of a gas increases, the number of collisions will also increase. This is like adding more ping-pong balls to the container. The increase in the number and force of collisions is what causes the pressure increase.

Every diver is introduced to Boyle's Law in their entry-level training. But just in case you're a little rusty, recall that for every 33 feet/10 meters of descent, sea water pressure increases by 15 psi or 1 atmosphere or atm. (I know, it's really 14.7, but decimals only complicate things here, and 15 is close enough for government work.) This means that at 99 feet/30 meters, the pressure is four times the surface pressure; three contributed by the water and one by the atmosphere. (99/33 = 3 + the surface pressure of 1 atm. = 4 atm.).

Similarly, a volume of gas in a flexible container like a balloon will also decrease predictably according to the depth (pressure). For example, at 33 feet/10 meters (2 atm.) the balloon will decrease to 1/2 its surface volume. At 66 feet/20 meters (3 atm.) it decreases to 1/3, at 99 feet/30 meters (4 atm.) to 1/4 and so on. The reverse holds true on ascent. The volume of a balloon partly inflated at 99 feet/30 meters (4 atm.) will increase four-fold if it's taken to the surface.

This constant relationship gives divers a neat little way to estimate how gas volumes change as they ascend or descend. Let's try an example. Assume that a balloon containing 20 liters of air at the surface is taken to 100 feet or 4 atm. (yes, 4 atm. is actually 99 feet, but one foot isn't worth arguing over.) By expressing the problem as a fraction where the original surface quantity (20) is the top number-the numerator, for the less mathematically-challenged-and the pressure at 100 feet (4) is the bottom number or denominator, we have 20/4 or 5 liters (once the fraction is reduced).

You can solve even more complex problems with this fractional relationship. Let's assume a balloon contains 5 liters at a depth of 66 feet/20 meters. What will its volume be if it's taken to 33 feet/10 meters? First, let's see what happens to the balloon if it's taken to the surface. As 66 feet/20 meters is 3 atm., the balloon will expand three-fold when it reaches the surface and contain 15 liters (3 x 5 = 15). Now, simply use the same fractional relation from the previous problem. The surface volume (15) over the pressure (2), or 15/2. Reduce the fraction to 7.5. Thus, the balloon would change from a volume of 5 liters at 66 feet/20 meters to 7.5 liters at 33 feet/10 meters.

The only drawback to this quick and dirty method of calculation is that it only works if you use atmospheres as increments. For more precise results, or for problems expressed in terms other than atmospheres, you'll have to use the dreaded mathematics of algebra, and famous equation P1V1=P2V2. The point is that the implications of Boyle's Law are easily determined using nothing more than some logic, common sense and simple arithmetic. You didn't even have to use a calculator!

The Plot Thickens

While understanding the relationship of pressure and volume is both vital and obvious to all divers-it explains the most important rule of scuba diving-fewer understand why density is just as important. Like pressure and volume, though, understanding density is also a matter of common sense.

To begin, let's squeeze the balloon again. As you saw before, this reduces its internal volume. So, if the volume is reduced, the molecules within the balloon have to come closer together. (Think of a conductor pushing a crowd of people aboard a subway car.) In other words, the air gets "thicker," but the more appropriate term is denser. There is one important difference between pressure/volume and pressure/density relationships, however. Note that the former relationship is inversely proportional-if one factor goes up, the other goes down. In the latter, the relationship is directly proportional-if one factor goes up, so does the other.

Luckily, the change in the density of a gas is as predictable as the change in its volume. If, for example, you take a flexible container to 2 atm. ( 33 feet/10 meters) the volume will decrease by 1/2 but the density will double. In turn, at 3 atm the volume will be one-third and the density will triple; at 4 atm. the volume is now one-fourth and the density quadrupled and so on. This relationship is important because it explains why a diver consumes his air more rapidly the deeper he goes. Here's why.

In free diving (breath hold), as a diver descends, his lungs decrease in size according to Boyle's Law. But this isn't the case when using scuba because a diver fills his lungs completely with every breath. To fill the lungs, the density of the air inhaled must increase to equal the ambient (surrounding) pressure. In fact, as we saw earlier, the air in a breath taken at 33 feet/10 meters will be twice as dense as at the surface. Likewise, to get a lung full of air at twice the density of the surface, twice the number of molecules must be drawn from the tank. So, as the diver is drawing twice the number of molecules from the tank at 33 feet/10 meters, his air supply will last only half as long as it would at the surface. Similarly, if all other conditions are equal, the air supply will last only one-third as long at 3 atm (66 feet/20 meters), one-fourth as long at 4 atm. (99 feet/30 meters), and so on.

Knowing this relationship gives you an easy way to estimate how long your air supply will last. But be forewarned, this is only a rough estimate. Your actual air consumption rate varies drastically due to factors such as cold, heavy exertion and stress. Additionally, the estimate works only if you plan to remain at a constant depth throughout the dive-something we divers rarely do.

You're Getting Warmer

Sir Robert Boyle's work dealt solely with the effects of pressure and volume. He did not consider the effect of a third important factor-temperature. The influence of temperature on gas behavior was first explored by a French scientist named Jacques Charles and his colleague, Joseph Gay-Lussac.

Through experimentation, Charles and Gay-Lussac found that if they kept the pressure of a gas constant in a container, the volume of the gas would increase as the temperature increased. Conversely, the volume would decrease when the temperature decreased and the pressure remained constant. This is called Charles' Law, and it explains, for example, why, if you leave an inflated balloon in a hot car all day it will expand.

One way Charles' Law relates to diving is in how temperature affects the pressure in a scuba tank. As a general rule, for every change in temperature of one degree Fahrenheit, the pressure in a scuba tank changes about 5 psi. For every change in temperature of one degree Celsius, the pressure in a scuba tank changes about 9 psi. To do more precise calculations requires a little simple algebra; but at least you now understand why your tank's burst disk let go that time you forgot and left it in your trunk on a hot summer day.

Years later a researcher named Julius Mayer explained the phenomenon Charles and Gay-Lussac described. He suggested that heat is simply a result of molecular motion. Therefore, increasing the temperature was merely increasing the motion of the molecules. As the molecules increased in motion, they increased in their frequency of collision. This is measured as an increase in pressure. In 1843, the scientist James P. Joule proved Mayer's theory experimentally. Incidentally, if you're not terribly afraid of a little math, both Boyle's and Charles' Laws can be integrated into what's called the General Gas Law. This lets us predict gas behavior regardless if the factors that vary are volume, pressure or temperature.

Mixed Gas

So far, nothing we've discussed considers what happens when different gases are mixed together. While this may seem a trivial matter, it's anything but. In fact, not knowing how gases behave in a mixture can get a diver killed. The first person to explore this phenomenon was an English scientist, John Dalton. (In addition to his work with gases, Dalton is famous for his work describing matter as being comprised of atoms).

Dalton found that even though a gas mixture is made up of several different constituents, each gas will continue to demonstrate its own individual behavior, as though the other gases didn't exist. Again, more common sense; if a gas mixture is made up of 80 percent nitrogen and 20 percent oxygen, then it stands to reason that 80 percent of the pressure is exerted by the nitrogen and 20 percent by the oxygen. However, this tidbit of common sense is so important that it's called Dalton's Law.

Furthermore, Dalton termed these individual pressures within a mixture partial pressures; and he found that each partial pressure is proportional to the number of molecules of that gas within the mixture. Let's see how this concept applies to us as divers.

Let's return again to our gas mixture of 20 percent oxygen and 80 percent nitrogen, and we'll assume it's at a pressure of 15 psi. This is similar to our own air, but ignores the trace gases, and working in round numbers makes the concept easier to understand. According to Dalton's Law, oxygen exerts 20 percent of the total pressure of the gas, while nitrogen exerts 80 percent. Another way of stating this is that of the 15 psi total pressure, the partial pressure of the oxygen exerts 3 psi (20 percent) while nitrogen partial pressure exerts the other 12 psi (80 percent).

Now let's double the total pressure to 30 psi and see what happens. (We'll assume the temperature remains constant at all times.) Each gas component continues to exert its partial pressure in proportion to the 80/20 mixture. Of the total 30 psi, nitrogen exerts 80 percent or 24 psi, and oxygen exerts 20 percent or 6 psi.

If the ambient pressure increases, the pressure inside the container (such as our lungs) also must increase for the container to maintain its original volume. This means more gas must be put in the container to increase the internal pressure. As we saw before, this is exactly what happens when a diver breathes from his regulator at depth; more gas molecules go into his lungs to maintain a constant lung volume.

Note that while the percentages of the gases that make up the mixture never change, the number of gas molecules increases to exert more pressure. This means more molecules of each gas reach the diver's lungs. For example, let's assume a lung volume at the surface contains 100 molecules of air. (Of course the real number would be vastly more than this, but a small, round number makes it easier to understand.) If we continue to assume an 80/20 nitrogen-oxygen mixture, then of the total 100 molecules, 80 will be nitrogen and 20 oxygen.

Now assume the diver descends to 132 feet or 5 atm. At that depth the ambient pressure is five times what it was at the surface. So, to maintain a normal lung volume, the diver takes in 500 molecules with each breath. This is five times the number he breathed at the surface. As this is still an 80/20 gas mixture, nitrogen accounts for 400 molecules or 4 atm. of pressure, while the oxygen component is responsible for 100 molecules or 1 atm. Notice that breathing this air mixture at 132 feet is physiologically equivalent to breathing pure oxygen at the surface-100 molecules at one atmosphere.

We can demonstrate the insidious nature of increasing partial pressures in yet another way. Suppose that while filling a scuba tank a one percent quantity of carbon monoxide enters the mixture. Using the previous example of 100 molecules to equal a full lung, carbon monoxide would now account for one of the 100 molecules inhaled with each breath while at the surface. But, if the diver descends to 99 feet (4 atm), his lungs now require 400 molecules to fill-four times the surface quantity. Of those 400 molecules drawn from the tank with each breath, carbon monoxide will now account for four molecules. The problem is, breathing these four molecules of carbon monoxide at 99 feet has the same physiologic effect as breathing a gas mixture of four percent carbon monoxide at the surface (which is toxic).

Thus, we see why an air supply that might contain tolerable levels of contaminants at the surface can quickly become toxic at depth even though the gas inside the tank never changes. Once the tank is full, the proportion of gases in the mixture remains unchanged. The problem results purely from breathing the mixture under higher pressure; and it's all a result of the good Doctor Dalton's Law.

Bubbles 101

As a science educator, one of the more difficult concept for my students is understanding that gases can dissolve into liquids. The reason is that the concept seems to belie reality. For example, to the naked eye a liquid doesn't appear to have any "space" in it to allow gas to enter. However, we all know better; gases can dissolve into liquids. We're reminded of this every time we pour a can of soda into a glass of ice. The profuse formation of bubbles is merely carbon dioxide being released from the liquid. The gas, although not visible in its dissolved form, remains in solution until something acts on it to make it come out. The other important concept, beside the fact that gas can be held inside a liquid, is that the gas also continues to exert pressure while in solution. This internal pressure is called gas tension.

Exactly how much gas a liquid will absorb, and what factors affect this absorption, was first studied by a colleague of Dalton's named William Henry. His experiments proved that the amount of gas that dissolves into a liquid at a given temperature depends on the partial pressure of the gas. Yep, you guessed it. This concept is called Henry's Law, and it's yet another phenomenon that we've all experienced. Why, for example, does a freshly tapped keg of beer taste better than the dregs at the end of the keg? Simple-because of Henry's Law. More gas-in this case, carbon dioxide-is dissolved in the liquid when it's first tapped. But over time, as the carbon dioxide dissipates, the contents "go flat." The decline in taste results from the reduced gas tension in the liquid.

Henry's Law further tells us that both pressure and temperature affect how gas dissolves into liquids. Let's see how this might happen. Imagine you have a beaker of water with absolutely no gas dissolved in it. The beaker is also inside a chamber where a perfect vacuum exists. This means that absolutely no gas is in contact with the water. (These conditions are impossible to achieve except perhaps in outer space, but it makes the concept easier to understand.) If air enters the chamber and comes in contact with the water, air molecules begin dissolving into the water. The gas in the water then exerts a pressure, or gas tension. In addition, the gas tension will obey Dalton's Law and each component gas continues exerting its own partial pressure.

Gas will attempt to dissolve into liquid until the gas tension equalizes with the air pressure in contact with the liquid. But this takes time. The difference between the air pressure in contact with the liquid and the gas tension within the liquid is called a pressure gradient. If the pressure gradient is high, gas will be absorbed into the liquid quickly. But, as the gas molecules continue to dissolve into the water, the gradient begins to decrease. Molecules dissolve into the water more slowly.

Over time the gas tension in the water equals the air pressure in contact with the liquid. Although molecules will continue to pass between the liquid/air interface, no net exchange of gas occurs. The water is now said to be saturated.

Let's now pressurize the vacuum chamber. This increases the pressure of the gas in contact with the water. This causes even more gas to dissolve into the water. Over time, gas will continue to enter the water until the gas tension in the liquid and the air pressure on the liquid are equal (saturation). As Henry's Law predicted, the more pressure the gas exerts on the water, the more gas dissolves into the water.

As you might expect, if we release the pressure in the chamber the phenomenon is reversed. With less pressure on the water, the gas dissolved in it has a greater gas tension than the air in contact with the water. Now the water contains more gas than it can keep in solution at that pressure.

Gas transfers out of the liquid until the gas tension is equal to the air pressure. If air pressure is reduced very slowly, if it isn't shaken, or if no foreign particles are present in the water, the gas transfer is undetectable-no gas bubbles form. But if the air pressure decreases too quickly, or the water is vigorously shaken, or if foreign particles are added to the water, the gas begins to escape more rapidly. So quickly that-like a shaken bottle of soda-the gas molecules will form visible bubbles.

In addition to pressure (Boyle's Law), temperature (Charles' Law) also affects gas absorption into liquids. Heat makes the molecules of a liquid vibrate more activity. This rapid movement leaves less space in the liquid for gas molecules to occupy. Now, fewer gas molecules can dissolve into or remain in the liquid. We see this take place anytime we boil water. As the water begins to heat, small air bubbles form and collect at the bottom of the container. They're caused by the speeded-up molecules pushing dissolved gas out of the water. It makes sense, then, that cooler liquids can hold more dissolved gas than warmer liquids. A cooler liquid contains slower molecules, allowing room for more gas molecules to occupy.

If you haven't already figured out why this phenomenon is so important, it's because gas dissolves into our blood and other tissues just as it dissolved into the beaker of water. But absorption isn't as big a concern as elimination; the real problem begins when pressure is reduced and the gas has to come out. Under certain circumstances, this could lead to significant bubble formation and decompression sickness. There are, however, some important differences between body tissues and a beaker of water, which makes understanding the mechanism of decompression sickness a bit more complicated than the oft-used "shaken bottle of soda" analogy. Still, whether it's carbon dioxide dissolving into soda water or nitrogen into the bloodstream, the physics is all the same.