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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 80º 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.
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ATM
|
PSI
|
BAR
|
DEPTH
|
|
1
|
14.7
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1
|
Sea level
|
|
2
|
29.4
|
2
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33 ft (10 m)
|
|
3
|
44.1
|
3
|
66 ft (20 m)
|
|
4
|
58.8
|
4
|
99 ft (30 m)
|
|
5
|
73.5
|
5
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132 ft (40m)
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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 70º F and after that the ambient temperature rise to 80º 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).
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