Diving Physics and "Fizzyology"

Introduction

Like all animals, human beings need oxygen in order to survive. When we breathe, we extract oxygen from the air, and use that oxygen for metabolism, which is how we convert the food we eat into useable energy to do the things that we do. One of the by-products of metabolism is carbon dioxide; whenever we exhale, we are getting rid of the carbon dioxide that our bodies produce.

The main purpose of breathing, therefore, is to provide our bodies with oxygen, and rid our bodies of carbon dioxide. We humans are terrestrial (land-dwelling) mammals, and as such, our lungs are designed to breathe gas. Unlike fishes, we have no gills, so we cannot breathe water. Therefore, the first problem we must overcome to explore the underwater realm is a means to provide breathing gas.

 However, if this were the only barrier humans must overcome to enter the sea, we would have long-ago discovered most of the mysteries of the ocean. All we would need to remain underwater indefinitely would be a long tube going to the surface -- a huge snorkel -- through which we could breathe. Unfortunately, there is another problem we must over come when descending to the depths -- a problem with far more complex and difficult consequences. .

Pressure

Have you ever wondered why nobody makes snorkels that are ten, or twenty, or a hundred feet long? The answer becomes obvious as soon as you try to breathe through a snorkel when your body is more than two or three feet (~1 meter) beneath the surface of the water.

If you have ever tried this, you will know that it becomes extremely difficult to inhale under those circumstances. That's because the deeper underwater you go, the greater the pressure is. Think of pressure as a force pushing on you from all directions. At sea-level, we are exposed to about 14.7 pounds per square inch (psi) of pressure. This means that each square inch of our bodies has the equivalent force of about 14.7 pounds pressing on it.

The source of this pressure is actually a result of the weight of the air in Earth's atmosphere. Like all gases, the air around us is composed of molecules of different gases; in this case, about 21% of these molecules are oxygen, about 78% are nitrogen, and the rest is composed of assorted trace gases. These gas molecules have weight, which means that gravity is pulling them toward Earth.

As it turns out, if you took a column of air, one square inch in cross-section, extending from sea-level all the way to the edge of the atmosphere, all the gas molecules in that column of air would have a combined weight of about 14.7 pounds -- which leads to 14.7 psi of pressure at sea level. For convenience, physicists have defined the "atmosphere" (abbreviated "ATM") as a unit measurement of pressure equal to the pressure caused by Earth's atmosphere at sea level (14.7 psi).
 

Therefore, at a depth of 33 feet (10 meters) beneath the sea surface, the total ambient pressure is about 29.4 psi, or 2 ATM -- 1 ATM caused by the weight of the air in Earth's atmosphere, plus 1 ATM for the weight of 33 feet (10 meters) of seawater. To avoid confusion, when people discuss pressures underwater, the unit "ATA" (referring to "atmospheres absolute") is often used to represent the total, "absolute" pressure caused by both the water and the air above the water
 

As is illustrated in the diagram at right, the ambient pressure increases underwater at an almost linear rate with increasing depth*. For every 33 feet (10 meters) of depth in sea water, the ambient pressure increases by an an additional 14.7 psi (1 atm).

At a depth of 99 feet (30 meters), the ambient pressure is 4 ATA -- one ATM caused by the Earth's atmosphere, plus 3 ATM for every 33 feet (10 meters) of depth. Similarly, the ambient pressure 297 feet (90 meters) beneath the surface is 10 ATA.

The problem with the long-snorkel idea has to do with the fact that the muscles we use to expand and contract our lungs during breathing are not strong enough to overcome much pressure.

Even just a few feet beneath the surface, the pressure is great enough that we cannot expand our lungs against the water pressure to inhale a breath of air from the surface. Our bodies simply are not designed to do that.
 

Underwater, the increased ambient pressure acts on the hull of the submersible, not the diver inside. Thus, the person inside the submersible is protected from the ambient pressure at all times and has no difficulty breathing.
 

Another way to overcome the problem of breathing under pressure is to provide a pressurized breathing gas mixture to the diver. If the breathing gas supply is delivered at the same pressure as the surrounding ambient pressure, the diver's lungs do not have to work against the water pressure (i.e., the pressure in the surrounding water and the pressure in the inhaled gas supply are balanced).

However, when using this sort of "Ambient Pressure Diving" technology, the diver's body is directly exposed to the ambient pressure. More importantly, the gas inhaled into the diver's lungs is pressurized. To understand the physiological ramifications of this, it's important to understand the effects of increased pressure on gases.
 

As mentioned above, gas is composed of molecules. As the pressure of the gas increases, these molecules (which are in constant motion) get packed more closely together. At a higher pressure, a given number of gas molecules will occupy a smaller volume (or a given volume will be occupied by a larger number of gas molecules).

For example, if you had a balloon filled with gas at the surface, and then pulled that balloon underwater, the balloon would shrink in size. In fact, if you brought the balloon down to a depth of 33 feet (10 meters), it would be half the size it was at the surface. At 66 feet (20 meters) it would be one-third the size it was at the surface; at 99 feet (30 meters) it would be one fourth the size, and so on. If, at a depth of 99 feet (30 meters), you wanted to expand the balloon to its original size, you would have to fill it with four times as many gas molecules as were required at the surface.

Returning the re-inflated balloon back to the surface would cause it to grow to four times its original size. It doesn't require much imagination to understand what would happen to a diver's lungs if he or she took a full breath at depth, and held it while ascending to the surface. This is why the golden rule of diving is "never hold your breath". People who forget this rule (for example, in a state of panic), run the risk of suffering from ruptured lungs, allowing gas bubbles to directly enter the blood. It's called embolism, and it can lead to serious symptoms including paralysis and death.
 

A key point here for understanding other aspects of diving physics and physiology, is to realize that the greater the pressure, the more tightly packed (i.e., more highly concentrated) gas molecules are.
 

*Note: Because water is very slightly compressible, the relationship between depth and pressure is not exactly linear all the way to the bottom of the sea; however, for the purposes of diving technology, the deviation from a linear relationship is trivial.

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 earlier, 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".

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.

Oxygen

Oxygen is the only gas we really need to breathe in order to stay alive. If we don't breathe oxygen (or if we don't breathe enough oxygen), we soon die. Interestingly enough, too much oxygen can be a bad thing also. At right is a diagram illustrating the range of oxygen concentrations that we can breathe safely.

The green and red bar represents a scale of inspired oxygen partial pressures (PO₂), ranging from zero oxygen on the left, to a PO₂ of 2.0 ATA on the right. Through evolution, our bodies have become optimized to breathe oxygen at a partial pressure (PO₂) of 0.21 ATA.

If the inspired PO₂ is much less than this, our bodies begin to shut down -- leading to unconsciousness when the PO₂ drops below about 0.1 ATA (i.e., 10% oxygen at sea level). This is called hypoxia. Breathing more than 0.21 ATA oxygen is generally fine ... up to a point. If the inspired PO₂ is maintained above about 0.5 ATA for prolonged periods of time (many hours to days), people begin to suffer what is usually referred to as "pulmonary" or "chronic" oxygen toxicity.

The effects of this include a burning sensation or irritation in the lungs, and can affect breathing. Except for people who spend days under pressure at a time (e.g., commercial divers on oil rigs), this form of oxygen toxicity is not much of a problem for divers.
 

Although these convulsions do not appear to cause any sort of permanent damage by themselves, the problem of a diver experiencing such convulsions being able continue to hold a regulator in his or her mouth is obvious. More than a few divers have drowned underwater, apparently a result of oxygen-induced convulsions. This is perhaps the most serious and insidious of diving maladies, because it comes on unpredictably and without warning, and usually results in the death of the afflicted diver.
 

There is no clear understanding on the exact biochemical processes involved with CNS oxygen toxicity, nor is there a clear consensus on what the "safe" upper PO₂ limit should be. Convulsions have occurred in divers breathing an inspired PO₂ as low as 1.2 ATA, but such cases usually involve extenuating circumstances (such as medical conditions in the divers which pre-dispose them to convulsions).

Conversely, commercial divers in Europe have routinely breathed oxygen partial pressures as high as 1.9 ATA in the water, and hyperbaric chamber facilities regularly expose patients to 2.8 ATA of oxygen (or more) without difficulty. Amid the ambiguities, two trends seem very consistent. The first is that high levels of exercise (perhaps more specifically, high levels of CO₂ in the blood) appear to increase the probability of suffering from a convulsion.

Secondly, divers immersed in water have a lower tolerance to elevated concentrations of inspired oxygen than do divers kept dry in a hyperbaric chamber or undersea habitat. (this over and above the fact that divers in a dry habitat are much more likely to survive a convulsion than are divers immersed in water). Another unavoidable reality regarding oxygen toxicity is the extreme range of variation both between individuals, and within a single individual.

When immersed underwater, most divers regard a PO2 of 1.4 ATA as a safe upper limit during periods of physical exertion, and 1.6 ATA during periods of rest.

Nitrogen

Eighty percent of the gas molecules in air are nitrogen (N2). Our bodies do not need or use nitrogen for metabolism, so it serves little function in a breathing gas mixture, other than to dilute the concentration of oxygen. When high concentrations of nitrogen become dissolved in our bodies, however, nitrogen can affect our central nervous system.

Familiar to all divers is the effect known as "nitrogen narcosis". Cousteau called it "rapture of the deep", and its effects have been likened to alcohol inebriation. When breathed at high concentrations, nitrogen can impair our neurological abilities. The exact biochemistry is not known, but the symptoms include impaired judgement, loss of short-term memory, slowed response time, and sometimes euphoria. Obviously, just as one should not drive while intoxicated, diving with impaired mental abilities is at the very least unwise.

As with oxygen toxicity, there is a wide range of variation in susceptibility to nitrogen narcosis both between, and within individuals. There is some evidence that repeated exposure can lead to an "adaptation" effect, but this is a topic subject to continued debate. Some divers begin to notice the symptoms while breathing air as shallow as 90 feet (27 meters) or so, while other claim to suffer no incapacitation at depths in excess of 200 feet (61 meters).

Impairment likely occurs at lower PN2 levels than those at which divers begin to detect overt symptoms. In any case, the greater the inspired PN2, the more severe the narcosis. There is some evidence that oxygen also contributes to narcosis, but probably only at concentrations above which CNS oxygen toxicity would be of primary concern.

Nitrogen plays another important role in limiting conventional scuba diving: it's involvement with decompression sickness. This will be discussed in greater detail in the following section on decompression.