The Physics of Pressure and Diving
Physics as a Survival Skill
The gas laws appear in every diving textbook, but they are often taught as a box-ticking exercise — memorise Boyle's Law, pass the exam, forget it in the water. That is a mistake. The physics of pressure and gas behaviour underpins every significant decision a diver makes, from how fast to ascend to why a wetsuit compresses at depth to why breathing from a scuba cylinder at 30 metres and then holding your breath on the way up will rupture lung tissue. Understanding the physics at the level of intuition, not just as equations, is part of what distinguishes a competent diver from a technically-trained one.
This is not complex physics. Three gas laws dominate: Boyle's, Henry's, and Dalton's. Each describes one aspect of how gas behaves under pressure, and each translates directly into something that happens to you underwater.
Pressure and Depth
Before the gas laws, the concept of absolute pressure. Pressure underwater increases with depth at a rate determined by the weight of the overlying water. Seawater weighs approximately 1025 kg per cubic metre, and the result is that absolute pressure increases by roughly 1 bar for every 10 metres of depth. At the surface, the pressure is 1 bar (atmospheric). At 10 metres it is 2 bar absolute. At 20 metres it is 3 bar. At 40 metres it is 5 bar.
This linear relationship is fundamental. Every gas-law calculation in diving starts from absolute pressure in bar (or in atmospheres, which are numerically equivalent for practical purposes). The mistake beginners make is working with gauge pressure — the pressure reading above atmospheric, as shown on most dive gauges — rather than absolute pressure. Boyle's Law, for instance, applies to absolute pressure. A gas compressed from 2 bar to 4 bar absolute has halved in volume; a gas compressed from 2 bar gauge to 4 bar gauge has gone from 3 to 5 bar absolute and the volume calculation is different.
Boyle's Law: Volume and Pressure
Boyle's Law states that at constant temperature, the volume of a gas varies inversely with the absolute pressure applied to it. As pressure increases, volume decreases proportionally. As pressure decreases, volume increases. The product of pressure and volume is constant: P × V = constant.
The practical consequences are everywhere in diving. Your lungs, BCD, drysuit, and wetsuit all contain gas spaces that obey Boyle's Law. When you descend from the surface to 20 metres, the pressure increases from 1 to 3 bar absolute, and any gas-filled space in your body or equipment that cannot be equalised to ambient pressure will compress to one-third of its surface volume.
Your wetsuit is compressed. A 7mm wetsuit at 30 metres has half the insulating thickness it has at the surface — the foam neoprene compresses under the fourfold increase in pressure. The thermal protection you relied on in shallower water is partly absent at depth. Your BCD contains less air at depth than it did at the surface for the same buoyancy contribution. Adding air to your BCD on descent compensates for the compression of its gas space; that same air will expand on ascent and must be vented to prevent an uncontrolled ascent.
The most dangerous application of Boyle's Law is pulmonary barotrauma. If a diver inhales compressed air at depth — say at 20 metres, where the air is at 3 bar absolute — and ascends to the surface without exhaling, that air will expand to three times its volume inside the lungs. Lung tissue cannot stretch that far. The gas will find another route out: through ruptured alveoli, into the chest cavity (pneumothorax), into the mediastinum (mediastinal emphysema), or into the arterial circulation (arterial gas embolism). This is why the cardinal rule of scuba diving — never hold your breath while breathing compressed gas — has a hard physical basis. It is not a procedural nicety.
Dalton's Law: Partial Pressures
Dalton's Law states that the total pressure of a mixture of gases equals the sum of the partial pressures of each gas in the mixture. Each gas in a mixture exerts a pressure proportional to its fraction of the total mix, regardless of the presence of the other gases.
Air is approximately 21% oxygen and 78% nitrogen (with trace amounts of other gases). At the surface, where total pressure is 1 bar, the partial pressure of oxygen (PO2) is 0.21 bar and the partial pressure of nitrogen (PN2) is 0.78 bar. At 40 metres, where total pressure is 5 bar, the PO2 is 1.05 bar and the PN2 is 3.9 bar.
This has two major safety implications. The first is oxygen toxicity. Pure oxygen becomes hazardous to the central nervous system at partial pressures above 1.6 bar, which occurs at 6 metres when breathing 100% oxygen. For air, this limit occurs at depths far beyond recreational diving (below 60 metres for 21% oxygen). For enriched-air nitrox — say 36% oxygen — the maximum operating depth for a PO2 of 1.6 bar is only about 34 metres: (1.6 ÷ 0.36) − 1 × 10 = approximately 34 metres. Every nitrox diver must calculate their maximum operating depth and not exceed it, and this calculation is simply Dalton's Law applied.
The second implication is nitrogen narcosis. At elevated partial pressures, nitrogen produces an anaesthetic effect on the central nervous system — a feeling that can range from mild euphoria and reduced fine motor control to profound confusion and poor judgment. The onset is typically noticeable below 30 metres and becomes a real limitation below 40 metres for many divers. The physical cause is the increased PN2 — the same gas at higher partial pressure acting on neural membranes.
Henry's Law: Gas Dissolving in Liquid
Henry's Law states that the amount of gas dissolved in a liquid is proportional to the partial pressure of that gas above the liquid. More pressure means more dissolved gas.
At depth, your blood and tissues are exposed to breathing gas at elevated pressure. The partial pressure of nitrogen in the breathing gas is higher than at the surface, and so more nitrogen dissolves into the blood and eventually into the body's tissues — muscles, fat, central nervous system. This is the physical basis of decompression theory. The nitrogen absorbed during the dive must be eliminated on ascent, and the rate of elimination is limited by the capacity of the blood to transport it to the lungs. If the diver ascends too fast, the partial pressure of nitrogen in the tissues exceeds the local ambient pressure, and the dissolved nitrogen comes out of solution as bubbles — the mechanism of decompression sickness.
Henry's Law also explains why nitrox reduces decompression risk relative to air. When you breathe 32% nitrox instead of air at the same depth, the partial pressure of nitrogen in the breathing mix is lower (68% nitrogen versus 78% nitrogen). Less nitrogen is therefore driven into solution, the tissues absorb less decompression obligation per unit time, and no-decompression limits can be extended modestly. At 20 metres, the no-decompression limit on air is approximately 45 minutes according to most tables; on 32% nitrox it extends to around 95 minutes, because the nitrogen loading rate is slower.
At the other extreme, trimix — a blend of oxygen, nitrogen and helium used in deep technical diving — uses helium to displace nitrogen. Helium has favourable properties for deep diving: it diffuses in and out of tissues faster than nitrogen, eliminating more quickly on ascent, and it does not cause narcosis at recreational or technical recreational depths. The trade-off is that faster diffusion means faster saturation as well, which requires more careful decompression management rather than less, but the absence of narcosis and the more predictable decompression profile make it the preferred inert gas below about 50 metres.
The Interconnection of the Laws
In practice, the three laws operate simultaneously on every dive. When you descend, Boyle's Law compresses your gas spaces and your equipment. Dalton's Law determines the partial pressures of each gas you breathe at every depth, governing both your oxygen exposure and your nitrogen loading. Henry's Law drives the dissolution of nitrogen (and any other inert gas) into your tissues in proportion to those partial pressures.
Understanding how they connect — why a deeper dive on nitrox still has a decompression obligation even though it extends no-decompression limits; why breathing a higher-oxygen mix at depth increases convulsion risk while reducing narcosis risk; why ascent rate matters not just for decompression but also for the expansion of gas spaces in the body — is what separates a diver who can recite facts from one who actually understands what is happening to them underwater.
Open the map to find dive sites at varying depths around the world and consider how the physics of pressure changes the diving experience at each one.
The physics of pressure does not require mathematical fluency to be useful. It requires enough conceptual understanding that the consequences of depth and gas composition feel intuitive, so that when conditions change, when plans shift, or when something unexpected happens at depth, the response draws on genuine understanding rather than a half-remembered rule from a course taken years ago.