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Gas Laws and Scuba Diving
Gases behave in nature and in scuba diving according to several gas laws. Knowledge of these scuba diving gas laws is important to the diver because they influence the duration of the air supply and affect the gas containing spaces in the body such as the ears, sinuses and lungs. They also cause other scuba diving illnesses.
This defines the relationship between pressure and volume. It states that the volume of a given mass of gas varies inversely with the absolute pressure (if the temperature remains constant).
Stated simply, for a given amount of gas, if the pressure is increased, the volume is proportionally decreased and vice versa. This means that if the pressure is doubled, the volume is halved and vice versa.
Stated mathematically:
It follows that for a given amount of gas, the volume multiplied by the pressure always has a constant value.
i.e. P × V is constant.
So if a sample of gas has an original volume of V1 and an original pressure of P1, and either the pressure or volume are changed, the new volume V2 and the new pressure P2 will multiply out to the same value.
i.e. P1 × V1 = P2 × V2
This law can easily be demonstrated by a piston and cylinder such as a bicycle pump. If the piston is pushed into the cylinder half way, and the escape of gas prevented, the pressure in the cylinder will be found to have doubled. By this process, many liters of air can be crammed into a bicycle tire but at the cost of an increase in pressure in the tire (and hard work).
Compressors work in this way, squeezing 2000 or
more liters of air into a scuba cylinder – but at a high pressure. Since water pressure increases with depth, the consequent reduction in gas volume becomes very important to the diver because his body has numerous air spaces.
Descent Problems: The air in the diver’s middle ear and sinuses will contract in volume as the diver descends. If these volume changes are not compensated for by adding more air (“equalization”), then pressure damage (barotrauma) to the tissues will result.
For example:
In the same way, if a breath-hold diver takes a full breath at the surface and descends to 20 meters(3 ATA), the volume of air in his lungs may be reduced from 6 liters to 2 liters. The chest and lungs cope with compression better than distension. The limit for breath-hold diving is not known, but now has been shown to exceed 150 meters in certain individuals.
Ascent Problems. An average male diver’s lungs may contain about 6 liters of gas. If a diver takes a full breath at 20 meters (3 ATA) from his scuba set and returns to the surface (1 ATA) without exhaling, the volume of gas in his lungs will increase from the 6 liter total lung capacity to 18 liters (6 × 3 liters).
This can be easily calculated this way:
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The lungs would have to expand to 18 liters to accommodate this volume – well beyond their rupturing point, causing burst lung (pulmonary barotrauma of ascent). An important practical observation of Boyle’s Law is that the greatest volume changes take place near the surface. This means that the greatest danger from barotraumas is near the surface — and this applies with descent as well as ascent.
For example, if diver has a maximum of 4 liters of air in his lungs at 40 meters depth (5 ATA) and ascends 10 meters without exhaling (to 4 ATA), the volume in the lungs will increase to 5 liters:
Some people could possibly accommodate this expansion without lung damage. If the same diver started at 10 meters depth (2ATA), and then ascended 10 meters to the surface (the same ascent distance as before), without exhaling, the pressure would change from 2ATA to 1ATA. The air in the lungs would expand from 4 to 8 liters. This would rupture his lungs. Although the dives involved the same ascent distances, the volume change, and hence the danger, in response to Boyle’s Law, is much greater near the surface. Many divers are not aware of this and have a fallacious belief that if they confine their diving to shallow depths they will minimize the risk of barotrauma. Buoyancy compensator’s are similarly affected by depth changes in response to Boyle’s Law. Wet suits are also affected and lose their buoyancy and insulating properties with depth.
Charles’ Law
Most divers will have noticed that bicycle pumps and air compressors become hot during use. As the volume of gas is compressed, heat is produced. This is explained by Charles’ Law. This Law states that if the pressure remains constant, the volume of a given mass of gas varies directly with the absolute temperature (absolute temperature is obtained by adding 273 to the temperature in degrees Celsius). In other words, at a fixed pressure, if gas is heated it expands, and if gas is cooled its volume contracts.
Charles’ and Boyle’s laws can be combined into the General Gas Law : 
For the non-mathematically minded this means that for a given amount of gas, the pressure multiplied by the volume, divided by the temperature, always comes to the same value – so if one of these factors is varied, it has an effect on the other two.
Stated in another way; if a gas is compressed, its volume decreases and it gets hotter. If the gas is heated and the volume is prevented from expanding, the pressure rises. The consequence of this law has lead to the demolition of several perfectly good automobiles (and divers!) following the storage of full scuba cylinders in the boot (trunk) in hot weather. Similarly, inflatable dive boats are often pressurized to the maximum and are then left in the sun. As the temperature rises, the pressure of the contained air progressively increases and then suddenly reduces – when the volume increases and when the boat explodes. If gas is allowed to expand rapidly, it cools. Cooling from the expansion of previously compressed air, as it is breathed from a scuba cylinder, can lead to the regulator freezing up during cold water diving.
Dalton’s Law
With a mixture of gases, the total pressure exerted by the mixture, is the sum of the pressures that would be exerted by each of the gases if it alone occupied the total volume. That is, the total pressure is the sum of the partial pressures. As the overall pressure increases (with descent underwater), so the partial pressure of each constituent gas increases. e.g. if air contains approximately 21% oxygen (O2) and 78% nitrogen (N2), then in a sample of air at a given pressure, O2 will contribute 21% of the total pressure and N2 will contribute 78%.
To calculate the partial pressure of a gas, multiply the percentage of gas by the absolute pressure. This law is important when considering the toxic effect of gases at depth or the use of O2 for
treatment purposes.
Henry’s Law
This law describes the dissolving of gas in a liquid and states that the quantity of gas which will dissolve in a liquid at a given temperature is proportional to the partial pressure of gas in contact with the liquid. This means that if the pressure of gas exposed to a liquid increases, then more gas will dissolve in the liquid.
An example of this law can be seen whenever a fizzy soft drink bottle is opened. During the manufacture of these drinks, carbon dioxide is dissolved in the liquid under pressure and the lid on the bottle maintains the pressure. When the bottle is opened and the pressure released, the liquid will not allow as much gas to be dissolved and so the excess gas is released from solution in the form of bubbles. At sea level (1ATA) the human body contains approximately 1 litre of N2 dissolved in the tissues. Whenever a diver breathes compressed air at depth, more N2 will dissolve in the body because the partial pressure of N2 in the air being breathed is increased. This is the cause of nitrogen narcosis. Under certain circumstances, when the diver returns to the surface this N2 can come out of solution in the form of bubbles. These bubbles cause tissue injury which is the basis of decompression sickness (“bends”).
Henry’s Law 2
Henry’s Law 3
Diffusion of Gases
If a diver were to pass wind in a confined room, all the occupants of the room would soon be aware of the fact but, fortunately, not necessarily the source. This process of distribution of gas is termed diffusion. It is caused by the rapid random movement of gas molecules to all parts of a contained space. Gas molecules, being only single or small groups of atoms, are able to easily diffuse through watertight membranes such as
blood capillaries or cell walls. This process allows O2 and other gases to pass from the lungs to the blood and tissues, and then back.
Diffusion 1
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The majority of this text was contributed via the Free and Open Works of
Dr. Edmond's Diving Medicine- http://www.divingmedicine.info
please visit their site for a FREE Downloadable PDF of their entire works.
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