Diving gas laws

The Gas Laws are particularly relevant to diving. Having a good understanding of the implications can help make you a safer diver. We frequently come across divers who seem not to be aware that the greatest proportional pressure change happens in the final 10 metres. Diving physics and the diving gas laws can seem a little dry during training but it is fair to say that your safety really does depend on you learning the lessons well.  Once you have a sound understanding of what is happening to you and your equipment underwater then you can concentrate on learning and practising the practical skills to keep you safe.

Boyle’s Law

At a constant temperature, the volume of a gas varies inversely with the pressure, while the density of a gas varies directly with pressure.

If the temperature is constant and air pressure increases, the density of the air increases also, while the volume decreases and vice versa.

As a diver, Boyles law affects you every time you enter the water. Air spaces in the body are subjected to pressure and volume change, in direct proportion to your depth.

Without doubt, understanding Boyle’s Law is very important in scuba diving.

Note that Boyle’s law also relates to gas density. This part of the law becomes particularly important on deep dives; inhaled air will become denser the deeper one goes. It follows that increased gas density increases gas absorption.

Some applications of Boyle’s Law in action:

  • Running out of air at depth (not taking into consideration the increased gas usage at depth)
  • Rapid ascent, caused by not adjusting buoyancy quickly enough to allow expanding air to escape the BCD/ Dry suit
  • Sinus, ear and mask squeeze on descent – sometimes resulting in bleeding. Sinus or inner ear pain may be experienced on ascent.
  • From 10msw to surface the pressure halves. If you breath hold, the air in your lungs will double in volume causing a ruptured lung.

Charles’ Law

At a constant volume, the pressure of gas varies directly with absolute temperature.

Given a constant volume of gas, the higher the temperature the higher the gas pressure, and vice versa.

Suppose you have a 10lt steel scuba tank holding 200 bars of air this equates to 10 x 200= 2000 litres of available air filled when the air temperature was 20°c . Now you take the tank into water that is 10°c. Before you take your first breath of that tank’s air, Charles’s law predicts that the tank pressure will be lower than 200 bar.

Since the water temperature is less than the air temperature the law predicts that the pressure now will be less than at 20°c. Dive shop owners know about Charles’s law, which is why they often fill tanks in water where the temperature is kept lower than the surrounding temperature.

So what problems can occur with Charles law?

  • probably only one that will significantly affect the diver which is having less gas volume in the cylinder at the lower water temperature than you expect.

Dalton’s Law

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


The total pressure of any gas mixture is the sum of it’s parts. For example, with air:


Gas Pressure on surface Pressure at 50msw
Air 1 bar absolute 6 bar absolute
Air made up of:
Nitrogen, 79% 0.79 bar absolute 4.74 bar absolute
Oxygen, 21% 0.21 bar absolute 1.26 bar absolute


In the case above, you can see that on surface the partial pressure of nitrogen is 0.79 bar absolute and oxygen 0.21 bar absolute, which relates directly to the percentage. At 50msw the percentages stay the same, but the partial pressure increases.

The amount of gas absorbed by the diver at depth is directly proportional to the partial pressure of the gases breathed.

So what problems can occur with Daltons law?

  • Nitrogen narcosis traditionally becomes a problem in the 20 to 30msw range, equivalent to a partial pressure of 2.37 to 3.16 bar absolute breathing air. The feeling of euphoria and well-being can lead to confusion and cause a risk to the diver. Building up dive depths at the start of the season can reduce this effect.
  • As depth increases, more inert nitrogen is absorbed in the body and longer decompression is required to release the diffused nitrogen from the body tissues.
  • A range of gas mixtures, involving various proportions of nitrogen, oxygen and helium, are used to reduce the effects of nitrogen narcosis, increase dive time and reduce decompression. Breathing the right gas at the right depth is essential to avoid problems.
  • Ensuring the partial pressure of oxygen does not exceed 1.2 to 1.6 bar absolute minimises the risk of in-water convulsions caused by breathing too high a partial pressure of oxygen. In chambers we can go much higher for medical treatments and of course there is no risk of drowning.

Henry’s Law

At a constant temperature, the amount of a given gas dissolved in a given type and volume of liquid is directly proportional to the partial pressure of that gas in equilibrium with that liquid.

When the ambient pressure is increased with depth, the partial pressure of oxygen and nitrogen in the body rises. There will be more molecules of each gas dissolved in the blood and tissues.

Dissolved gases will diffuse out via the lungs on ascent as ambient pressure decreases, until a new equilibrium is established. This will continue even after surfacing until all the dissolved gas is removed. A controlled ascent rate and the completion of decompression stops goes a long way to avoiding problems.

So what problems can occur with Henry’s law?

  • Insufficient decompression of dissolved gases.
  • Re-absorption of gas due to reverse profiles, or saw tooth profiles.
  • Avoid getting too cold as it effects solubility of gases.
The Diver Clinic

Written by  The Diver Clinic