Physics and physiology · Lesson 11
Dalton’s and Henry’s Laws: Breathing Gases and Decompression
Dalton’s law describes pressures within a gas mixture. Henry’s law describes dissolved gas at equilibrium. Together they help explain why depth matters even when the breathing-gas percentages stay the same.
Copperband learning sheet
Three gas laws answer different questions
Do not use a correct equation for the wrong problem.
Logic sheet: connect the idea with the decision
Use these prompts to check your understanding before discussing practical decisions with your instructor.
Remember: Example only: 0.21 × 3 bar = 0.63 bar oxygen partial pressure. This is not a personal exposure limit.
Educational information only. This sheet does not replace official training, practical instruction or certification through a recognised scuba diving organisation such as PADI. It is not an emergency protocol, personal medical assessment or technical dive plan.
Dalton: the parts add to the whole
For an ideal mixture, total pressure equals the sum of its component partial pressures. A component’s partial pressure is its fraction multiplied by absolute pressure: pGas = fraction × Pabsolute. Ordinary dry air contains about 21 percent oxygen, with nitrogen making up most of the remainder.
For a classroom example, 21 percent oxygen at 3 bar absolute gives an oxygen partial pressure of 0.63 bar: 0.21 × 3. The oxygen fraction is still 21 percent. This arithmetic is not a statement that a particular depth, gas or exposure is suitable for you.
Why a fraction is not an exposure limit
The physiological effect of a gas depends on more than its percentage. Partial pressure, exposure time, workload and the gas involved matter. More oxygen is not automatically better: oxygen exposure has limits, and Nitrox training teaches gas analysis and planning within the appropriate scope.
A calculator cannot verify what is actually in a cylinder. Gas identification, analysis, equipment compatibility and computer settings need the procedures taught in the relevant course. Do not use an illustrative fraction as an instruction to mix gas or change gases underwater.
Henry: dissolved gas depends on pressure
At constant temperature and equilibrium, the concentration of a gas dissolved in a liquid is proportional to its partial pressure above that liquid. Increasing pressure can increase the equilibrium amount dissolved; reducing pressure changes that equilibrium.
A real diver’s tissues do not all reach equilibrium instantly. Uptake and elimination occur over time, with different tissues and changing pressure histories. Repetitive dives matter because returning to the surface does not instantly reset the body’s gas loading.
Go a little deeper
an equation is not a decompression model
Henry’s law describes equilibrium, while decompression planning must account for time and changing conditions. Computers approximate aspects of uptake and elimination with mathematical models. Their outputs are not direct measurements of every tissue and cannot guarantee that decompression illness will not occur. Boyle, Dalton and Henry address different questions: gas-space volume, mixture partial pressure and dissolved-gas equilibrium. Keeping those questions separate avoids using a correct formula for the wrong problem.
Check your understanding
Does a lower nitrogen percentage remove decompression risk?
No. Enriched air changes the breathing mixture and planning considerations; it does not remove gas-exposure limits or guarantee freedom from decompression illness.
Connect theory with training
For a Pattaya trip, tell the team your qualification, last-dive date and learning goals so the practical programme can be matched to you.
Sources and further reading
This is an independently written educational guide, not an official PADI course manual or certification assessment. Ask your instructor about unclear concepts before applying them in the water.