Pressure  Boyle's Law / Charle's Law  General Gas Law  Partial Pressure (Dalton's Law)  Henry's Law

Pressure In order to stay underwater we need a means of breathing in order to supply our body with the necessary oxygen. What makes this task difficult is pressure. Pressure is defined as the amount of force per unit area. At sea level we are exposed to 14.7 pounds of pressure per square inch (psi). This means that the air surrounding us is pushing on each square inch of our body with 14.7 pounds due to its weight. These 14.7 psi is what is called one atmosphere. The same thing happens in the water. Since water is much denser than air it takes a lot less sea water to excert enough pressure to total 14.7 psi. If one were to take a column of sea water one square inch in diameter, the column would only have to be 33 feet (10 meters) tall to excert one atmosphere of pressure. This means, that if we dive to a depth of 33 feet our body is exposed to 2 atmospheres of pressure. 1 Atmosphere due to the weight of the air and one atmosphere due to the weight of the water. At a depth of 66 feet we are exposed to a pressure of 3 atmospheres or 3 * 14.7 psi = 44.1 psi. The following is a little Java Script calculator giving you the amount of pressure per square inch at any given depth in fresh or sea water:   Salt Water Fresh Water depth in feet answer Due to this pressure on a diver's body, a diver needs to breath pressurized air as the lungs are not capable of expanding enough to intake air which is not under pressure. This is the reason why we can't just use a snorkel of the necessary length to explore underwater without the worries of running out of air. In SCUBA diving the diver breathes air which is at the exact pressure of his surrounding environment. This means his lungs do not have to fight against the surrounding pressure. This fact poses other problems that need to be taken into consideration every time a dive is undertaken with compressed air.  back to the top

Boyle's Law Charle's Law  The first rule learned when breathing compressed air underwater is to never hold your breath. The reasoning behind this is simple. Gases are composed of molecules and under pressure these molecules get packed more tightly together. This means that with a higher pressure more gas molecules will occupy a given volume. At depth of 33 feet the same amount of gas will occupy only have the volume it would occupy at sea level. At 66 feet it would occupy one third of the volume. If a diver were to fill his lungs at a depth of 33 feet and then ascend to the surface while holding his breath. The air in his lungs would expand to twice the volume. This of course would lead to major medical problems (air embolism). The above phenomenon is explained in Boyle's Law:  If the temperature is kept constant, the volume of a given mass of gas is inversely proportional to the absolute pressure.  An increase or decrease of pressure is not the only way that the volume of a given amount of gas can expand or contract. This can also be achieved through a rise or fall in the temperature. When a gas is heated the molecules making up the gas move faster. If you were to heat up the gas in a balloon, the gas molecules would hit the sides of the balloon more often and with more force causing the balloon to expand. By cooling the gas you gain exactly the opposite effect. This phenomenon is explained in Charles law:  If the temperature remains constant, the volume of a given mass of gas is inversely proportional to the absolute pressure.  back to the top

General Gas Law This leads us to the General Gas Law:  P1 x V1    P2 x V2 _______ = ________  T1        T2 where:  P1 = initial pressure (absolute) V1 = initial volume T1 = initial temperature (absolute) and:  P2 = final pressure (absolute) V2 = final volume T2 = final temperature (absolute)  When applying the gas laws:   use the same unit of measure throughout the calculations  where pressure is involved use only absolute pressure (This means you have to add 14.7 psi to the pressure gauge reading. This is due to the pressure of the atmosphere.)  where temperature is considered use absolute temperature. To find absolute temperature add 460 degrees to Fahrenheit readings and 273 degrees to centigrade readings.  The following calculator uses the General Gas Law to calculate pressure/volume relationships.   back to the top

Partial Pressure (Dalton's Law) The second thing we need to consider when breathing compressed air under pressure is the effect the individual gases making up this air have on our metabolism. The air we breath is made up of aproximatley 20% oxygen and 80% nitrogen. Since these gases have effects on our body when breathed under pressure, it useful to understand the concept of partial pressure. Dalton's Law states:  The total pressure exerted by a mixture of gases is the sum of the pressures that would be exerted by each of the gases if it alone were present and occupied the total volume.  This means, that if you were to fill a cylinder with air to a total pressure of 1,000 psi then the partial pressure of nitrogen in the cylinder would be 800 psi and the partial pressure of oxygen in the same cylinder would be 200 psi making a total of 800 psi + 200 psi = 1000 psi.  depth in feet Salt Water Fresh Water back to the top

Henry's Law Now that we've understood partial pressure there is one more concept that needs to be understood in order to appreciate the complexity and dangers involved in breathing compressed air. This concept is stated in Henry's Law:  The amount of gas that will dissolve in a liquid at a given temperature is almost directly proportional to the partial pressure of that gas.  For example the average man's body contains about 1 quart of dissolved nitrogen from the air he breathes at the surface. If he breathes air for a long enough time at 5 atmospheres his body will contain about 5 quarts of dissolved nitrogen. This is due to the fact, that the partial pressure of nitrogen in the lungs is much higher than the respective amount of gas in the blood. Therefore more nitrogen will move into the blood from the lungs until a state of equilibrium is reached. This is why dive tables are used in SCUBA diving to avoid decompression sickness. Another problem can also arise when the partial pressure of oxygen gets to high oxygen poisoning.  back to the top This information was obtained from The Dive Locker