A little bit of gas theory

 

This chapter explains why we can't stay underwater for a long time even with an unlimited supply of air and why we should surface slow and follow the SSI diving tables to plan and do dives

 

Diving Physics

Water is denser and heavier than air because in an equal space, water contains more molecules than air. A cubic foot of salt water weights 64 pounds (a cubic foot of fresh water weights 62.4 pounds) and a cubic foot of air weights 0.0881 pounds. We know that a square inch the height of the entire atmosphere exerts the same pressure as a one-inch square column of water 33 feet deep into the ocean. Therefore a diver under water will feel the pressure of the atmosphere and the pressure of the water above his body. Water is much better heat conductor than air so a diver under water will loose more heat than in air. Air temperatures of 80 F will feel comfortable but under water it will seem cold. Water absorbs body heat 25 times faster than air. This "pulling away" of heat is due to the property of heat energy transmission known as conduction. Conduction can be slowed or stopped by placing a pour heat conductor between the body and the water. Convection and radiation are the other two methods of heat energy transmission but they have little or no effect on the divers. The pressure felt by a diver (absolute pressure) is the sum of the hydrostatic pressure (the pressure of the water which results from the weight of the water) and the atmospheric pressure. If you are under water at a depth of 33 ft. the hydrostatic pressure is 1 ATM, the atmospheric pressure is 1 ATM so the absolute pressure is 2 ATM. A depth gage will only measure the hydrostatic pressure so at 33 ft. the depth gage will measure 1ATM. Here is a chart that shows the comparison between the pressure in bars, ATM and psi.

ATM
PSI
BAR
DEPTH
1
14.7
1
Sea level
2
29.4
2
33 ft (10 m)
3
44.1
3
66 ft (20 m)
4
58.8
4
99 ft (30 m)
5
73.5
5
132 ft (40m)

If a gas if formed from a mixture of gases, the partial pressure of one of the component gases is defined as the pressure of that gas in the mixture.

 

Gas Laws

Dalton's law

John Dalton, an English chemist and physicist, lived from 1776 to 1884. In addition to discovering a law of physics, he also made the first scientific account of color blindness, mostly because he was afflicted with it. He formulated his law of partial pressure while trying to solve a problem with Newtonian physics. His law states:

"The total pressure exerted by a mixture of gases is the sum of the pressure that would be exerted by each of the gases if it alone where present and occupied the total volume."

For example at the sea level the atmospheric pressure is 14.7 psi. Air, which is a mixture of gases, consist of 78.08% nitrogen and 20.95% oxygen, 0.03% carbon dioxide and 0.94% other gases.
The partial pressure of the nitrogen at the sea level is 0.7808 * 14.7 = 11.48 psi
The partial pressure of the oxygen at the sea level is: 0.2095 * 14.7 = 3.08 psi

For divers this means that at any depth or pressure the proportion of the nitrogen and oxygen in air will remain constant. However, when the pressure increases the partial pressure increases in proportion to the change in pressure.

 

Boyle's law

Robert Boyle was an English scientist who lived from 1627 to 1691 and was a friend and contemporary to Isaac Newton. He developed his law of pressure to show that air consisted of particles that behaved like tiny, coiled springs. In 1662 he published his finding that states:

"For any gas at a constant temperature, the volume will vary inversely with the absolute pressure while the density will vary directly with the absolute pressure."

This means that if pressure increases volume decreases and if pressure decreases volume increases. The change in volume is predictable and can be calculated. For example an 80 cubic feet flexible container (at 1 ATM) will decrease to 40 cubic feet at 2 ATM. For divers, this is the explanation for the need of equalization on descend and on ascent and for the other pressure-related occurrences such as squeezes. It also explain why air is consumed faster at depth and the danger of the air expansion injuries.

 

Charles' law

Jacques Charles was a French scientist who showed the how the temperature influence the volume of a gas under pressure. His discovery elaborated on Boyle's work that assumed a constant temperature. His law states:

"For any gas at a constant pressure, the volume of the gas will vary directly with the absolute temperature. For any gas at a constant volume, the pressure of the gas will vary directly with the absolute temperature."

When air is cooled molecules come closer together, which makes the air denser. Dense air takes up less space so its volume decreases. If the air is inside of a balloon, heating the air would make the air to expand causing the balloon to increase in size. For divers Charles' law has implication in the SCUBA tank filling process. If a tank is filled at an ambient temperature of 70 F and after that the ambient temperature rise to 80 F the pressure inside the tank increases since the volume of the tank is constant.

 

Henry's law

William Henry, an English chemist and physician, lived from 1774 to 1836 and conducted concurrent investigations into the solubility of the gases with John Dalton, who was researching the partial pressure of gases. In 1802 he developed Henry's law, which states:

"The amount of any given gas that will dissolve in a liquid at a given temperature is a function of the partial pressure of the gas in contact with the liquid and the solubility coefficient of the gas in the particular liquid."

According to Henry's law, pressure will force a gas into solution. Conversely, if the pressure is removed, the gas will come out of solution and return to its original form. The classic example of Henry's law is the soda bottle. For divers the concept of a gas coming out of solution has major implications. During a dive, nitrogen is absorbed into the diver's tissues. On ascent, pressure decreases and the nitrogen will come out of the solution as bubbles if the pressure is decreased to quickly. If this occurs while the nitrogen is still in the diver's tissues, decompression sickness (DCS) results. The rat with which the nitrogen is coming out of the solution can be controlled by slow ascents (30 ft. per minute).

 

 

 


Updated July 1, 2002 3:49 PM by Vlad Pambucol