Scuba diving is great fun, so do we really need to look at all the technical stuff? The answer to that is without doubt ‘Yes’, if for no other reason but to ensure your own safety. Other than being devoured by a mythical sea monster, your biggest danger is from the air that you’re breathing! A reasonably basic understanding of the various gas laws will go a good way to keeping you safe. Let’s take a brief look at these laws and how they affect the scuba diver.
All That Water!
More than a few student divers seem to be concerned about the increasing water pressure that comes with depth. The fact of the matter is: You can’t feel it! The reason for this is that around 60% of the human body is water and as such virtually incompressible; however, the various airspaces in our body are compressible and thus, subject to the gas laws.
What Air Spaces?
We all hear of various stories concerning burst eardrums and severe mask squeezes but these are very easily avoided and should not cause a problem at all. As we dive deeper the increasing water pressure pushes on our mask and eardrums. We feel this pressure because the air contained within the mask and our inner ear is compressible. All we need to do is to equalise or balance this pressure to maintain normality. The airspace within the mask can be equalised by just breathing out through the nose… That’s it! The inner ear is just as simple, you just pinch your nose shut through the mask and blow out gently through your nose while slowly moving your head from side-to-side. Because you have pinched your nose shut, the air travels up through the Eustachian tubes and into the middle ear, equalising the pressure. On surfacing, both of these airspaces will equalise automatically.
The Gas Laws
We need to look at the various gas laws to understand better the way in which gas pressure affects the diver. It’s this basic understanding that will help you to avoid pressure related injuries and also improve your buoyancy control.
Robert Boyle (1627-1691) extended the work of Evangelista Torricelli on the pressure exerted by the atmosphere. He was interested in what would happen to a quantity of gas if the pressure should alter. Boyle demonstrated that: If the temperature remains constant, the volume of a given mass of gas is inversely proportional to the absolute pressure. This simply means that if the pressure is increased, the volume will decrease proportionally; the reverse is also true. It’s this law that explains the effects on air spaces within the diver’s body.
Jacques Charles (1746-1823) took Boyles Law a little further. Charles determined that a third factor needed to be taken in consideration; this factor was temperature. Charles had determined that: The amount of change in either volume or pressure of a given volume of gas is directly proportional to the change in the absolute pressure. It’s this law that explains how the pressure in a scuba tank increases on a hot day or how it decreases when you get into cold water. That’s why they fill scuba tanks in cold water! That way you don’t get a ‘short fill’ when the heat generated by filling pressure dissipates. This law does not affect the diver as the body’s core temperature is kept extremely constant.
John Dalton (1766-1844) determined that gases made up of more than one component could not be encompassed by the existing laws. He determined that although the gas ‘mixture’ exerts a pressure calculable with the previous laws, they make no allowance for the calculation of pressure exerted by the mixture’s individual components. His law is known today as Dalton’s Law of Partial Pressures and is extremely important to divers as air itself is a ‘mixture’, made up of approximately 21% Oxygen, along with 78% nitrogen and 1% Argon. Dalton’s Law states that: The total pressure is equal to the sum of the partial pressures and that the magnitude of each of these partial pressures is directly proportional to their percentage within the mixture. This law is incredibly important as the partial pressures exerted in a breathing mixture at different depths can have dire consequences physiologically for the diver.
William Henry (1774-1836), a close associate of John Dalton, involved himself in the physics associated with a liquid’s capability to absorb a gas and the factors that may change this capability. Henry’s Law states that: The amount of gas that will dissolve into a liquid at a given temperature is almost directly proportional to the partial pressure of that gas. This law explains how and why we ‘on gas’ and ‘off gas’ when diving and how we can avoid becoming a victim of decompression illness; indeed, diving tables and diving computer algorithms are based upon Henry’s Law. For divers, Henry’s Law is probably the most important, although Dalton’s law is very closely related almost to the point that they are inseparable.
We have covered the main gas laws associated with scuba, albeit very, very basically. To cover them in the detail that they deserve, with all the calculations and illustrations would be far too long and, dare I say, boring to most. The intention of this article is only to give an insight into the way that a diver’s breathing gas can affect his or her physiology. Often, a little understanding is all that’s needed to allay the fears that spoil the initial pleasure, or indeed, put people off altogether from exploring a fantastic underwater world.