Chapter 1 and chapter 2 were about gas as a gas: how it shrinks under pressure and how that pressure is shared in a mix. Henry’s law is about gas that goes into a liquid. Your body is mostly water. So this law explains why you take up gas at depth, and why that gas can form bubbles on the way up.
Henry’s law in one sentence
The harder a gas presses on a liquid, the more of it dissolves into the liquid.
A bottle of fizzy drink shows this. At the factory, carbon dioxide is pressed into the drink at high pressure, so a lot of it dissolves. The sealed cap keeps the pressure in. When you open the bottle, the pressure drops and the extra gas comes out as bubbles.
What the law says
What matters is the gas’s partial pressure: its share of the total pressure, as chapter 2 explained. For nitrogen it is written PN2. Double the partial pressure and, given time, twice as much gas dissolves. In a mix, each gas dissolves according to its own partial pressure.
This is why your body takes up more gas the deeper you go. The equation is short:
Planned with DiveLogicCheck the no-stop time at 30 m
Gas over water at 1 bar
A closed jar holds water with air above it at 1 bar, the pressure at sea level. Nitrogen moves between the air and the water in both directions. After a while the amounts moving each way are equal. The water is then saturated: it holds all the nitrogen that the PN2 above it, 0.79 × 1 = 0.79 bar, allows.
Raising the pressure
Now the air pressure above the water is raised to 4 bar, the pressure at 30 m. The PN2 is now 0.79 × 4 = 3.16 bar, so more nitrogen goes into the water than comes out. Divers call this on-gassing.
This does not happen at once. The amount dissolved rises quickly at first, then more slowly as it nears its new level. Chapter 10 covers how long this takes in the body.
Henry’s law
Once the water is saturated again, it holds 4 times as much nitrogen as at 1 bar, because the PN2 is 4 times higher: 3.16 is 4 times 0.79.
This is Henry’s law: at a steady temperature, the amount of gas a liquid holds when saturated is proportional to the partial pressure of that gas. On the chart it is a straight line through zero.
A fast pressure drop
Now the lid comes off and the pressure falls to 1 bar at once. The water still holds 4 times the nitrogen that 1 bar allows. It is supersaturated: it holds more gas than it can keep. The extra gas comes out faster than it can escape through the surface, so it forms bubbles in the water, just like opening a fizzy drink quickly.
A slow pressure drop
This time the pressure is lowered 1 bar at a time, with a wait at each step so gas can leave through the surface. The water never holds much more gas than its new pressure allows, and far fewer bubbles form.
The pressure falls by the same total both times. Only the speed is different.
Gas in your body
Henry’s law applies to your blood and the rest of your body. At depth, the gas in your lungs is at a higher partial pressure. Your blood picks up more of it as it passes through the lungs and carries it to the rest of your body. Given enough time, each part of your body holds dissolved nitrogen in step with the nitrogen partial pressure you breathe. After a long time on air at 30 m (4 ATA), the PN2 is about 0.79 × 4 = 3.16 bar. Your body then holds about 4 times as much nitrogen as at the surface.
Oxygen dissolves too, but your body uses it up. The gases that build up are the inert gases: gases your body does not use, such as nitrogen and helium. Each one dissolves in step with its own partial pressure. So a mix with less nitrogen, such as nitrox (air with extra oxygen), puts less nitrogen into your body at the same depth.
Taking up gas takes time
Your body does not fill up with gas at once. Different parts of the body fill at different speeds. Dive planners model them as tissue compartments, each with its own half-time: the time it takes to fill halfway to the new level. Blood and the brain fill in minutes. Fat and joints take hours. Most dives are too short for the slow tissues to reach saturation (full).
Think of it like filling a bath: it fills fast at first, then slower as it nears full. Chapter 10 explains how dive planners work this out.
Bubbles on the way up
On the way up, the pressure around you falls. A tissue that took up gas at depth now holds more than the new pressure allows. If you rise slowly enough, your blood carries the extra gas back to your lungs and you breathe it out. This is off-gassing. If the pressure falls too fast, the gas can form bubbles in your body and blood. Those bubbles cause decompression sickness, which chapter 5 covers. Every decompression plan exists to keep the pressure drop slow enough for your body to clear the gas.
Where the law comes from
William Henry (1774-1836) was a chemist in Manchester and a friend of John Dalton. He measured how much of several gases water would take up at different temperatures and pressures. The Royal Society published his results in 1803. Henry accepted Dalton’s idea that each gas in a mix acts on its own, and together the 2 laws give the rule divers use: each gas dissolves according to its own partial pressure.
What to remember
- Henry’s law: at a steady temperature, the amount of a gas a liquid holds when saturated is proportional to that gas’s partial pressure.
- In a mix, each gas dissolves according to its own partial pressure, which Dalton’s law gives.
- At depth your body takes up the inert gases you breathe, nitrogen and helium, in step with their partial pressures.
- Taking up gas (on-gassing) takes time: fast tissue compartments fill in minutes, slow ones in hours.
- When the pressure falls, dissolved gas must leave (off-gassing). If it falls too fast, the gas forms bubbles. A decompression plan controls how fast the pressure falls.
No-stop time at depth
DiveLogic’s no-deco limit tool gives the time you can stay at a depth before your body holds too much gas to come straight up to the surface. Try 30 m on air, then on nitrox. Nitrox has a lower nitrogen partial pressure, so your body takes up less nitrogen and the time is longer.
Sources
- Henry W. (1803). Experiments on the quantity of gases absorbed by water, at different temperatures, and under different pressures. Philosophical Transactions of the Royal Society of London 93: 29-42.
- Henry W. (1803). Appendix to Mr. William Henry’s paper, on the quantity of gases absorbed by water, at different temperatures, and under different pressures. Philosophical Transactions of the Royal Society of London 93: 274-276.
- Dalton J. (1802). Experimental essays on the constitution of mixed gases. Memoirs of the Literary and Philosophical Society of Manchester 5(2): 535-602.
- Sheffield P. J., Vann R. D. (eds) (2004). Flying After Recreational Diving Workshop Proceedings, May 2, 2002. Durham, NC: Divers Alert Network.