A scuba diver breathes compressed air at depth where nitrogen has a partial pressure of 4 atm. At the surface, atmospheric nitrogen partial pressure is approximately 0.8 atm. According to Henry's law, what happens to the nitrogen dissolved in the diver's blood when they ascend?
AIt remains dissolved — nitrogen is permanently trapped in the bloodstream once absorbed
BIt is rapidly exhaled as the lungs exchange nitrogen gas with the atmosphere
CIts equilibrium solubility drops to about 20% of the depth value, so excess dissolved nitrogen tends to come out of solution
DNothing changes — blood plasma is not an ideal solution, so Henry's law does not apply to biological fluids
Henry's law (C = kH × P) tells us the equilibrium dissolved concentration is proportional to partial pressure. At 4 atm, the equilibrium concentration is 5× higher than at 0.8 atm. When the diver surfaces, the equilibrium concentration drops to 20% of its at-depth value, so the blood is supersaturated — it holds far more dissolved nitrogen than the new pressure can sustain. If ascent is too fast, nitrogen comes out of solution as bubbles in tissues and blood, causing decompression sickness (the bends).
Question 2 Multiple Choice
A factory dissolves CO₂ into water at 5°C under elevated pressure. An engineer proposes switching to 50°C water to dissolve more CO₂ per unit time. What is the fundamental error in this reasoning?
ACO₂ reacts with warm water to form carbonic acid, which corrodes the equipment
BHigher pressure would be required at 50°C to achieve the same flow rate, increasing energy costs
CGas solubility decreases with increasing temperature, so warm water would dissolve less CO₂ than cold water at the same pressure
DAt 50°C, CO₂ transitions to a supercritical state and Henry's law no longer applies
Unlike most solids, gases become less soluble as temperature rises. Higher temperature gives dissolved gas molecules more kinetic energy, making them more likely to escape the liquid back into the gas phase. The equilibrium shifts toward less dissolved gas. This is the opposite of what the engineer assumed. Cold water is deliberately used in carbonation precisely because it holds more CO₂ at a given pressure — and why warm soda goes flat faster than cold soda.
Question 3 True / False
A gas that reacts chemically with water (such as HCl or NH₃) will dissolve in greater quantities than Henry's law alone predicts.
TTrue
FFalse
Answer: True
Henry's law predicts dissolved concentration based on equilibrium between gas-phase and dissolved-phase molecules. When a gas reacts with water (e.g., HCl → H⁺ + Cl⁻; NH₃ + H₂O → NH₄⁺ + OH⁻), the dissolved molecules are consumed by the chemical reaction and converted to ions, removing them from the dissolved equilibrium. This continuously pulls more gas into solution beyond what the pressure-based equilibrium would predict. The total amount dissolved (as ions plus molecules) greatly exceeds Henry's law predictions, which is why HCl is described as deviating from ideal Henry's law behavior.
Question 4 True / False
Henry's law states that gas solubility is proportional to total atmospheric pressure, so at high altitude (where total pressure is lower), most dissolved gases in water should be equally reduced.
TTrue
FFalse
Answer: False
Henry's law concerns the partial pressure of the specific gas in question, not total atmospheric pressure. At high altitude, total pressure is lower, but the key quantity for each gas is its own partial pressure (determined by its mole fraction times total pressure). The partial pressure of O₂ at altitude is lower, reducing dissolved oxygen in water — but this is because O₂'s partial pressure dropped, not because total pressure dropped uniformly. Gases whose partial pressures are maintained (e.g., by a pressurized system) are unaffected regardless of ambient atmospheric pressure.
Question 5 Short Answer
Explain using Henry's law why a sealed soda bottle stays carbonated at room temperature but goes flat within hours of being opened.
Think about your answer, then reveal below.
Model answer: In a sealed bottle, the CO₂ above the liquid is at high pressure (2–4 atm), and Henry's law predicts a correspondingly high equilibrium dissolved CO₂ concentration. The system is at equilibrium: the rate of CO₂ escaping the liquid equals the rate of CO₂ dissolving back in. When the bottle is opened, the partial pressure of CO₂ above the liquid plummets to atmospheric CO₂ partial pressure (~0.0004 atm). Henry's law now predicts a much lower equilibrium dissolved concentration — the liquid is supersaturated. CO₂ escapes as bubbles until the dissolved concentration reaches the new, much lower equilibrium value, at which point the soda is flat.
The key insight is that Henry's law governs equilibrium: the dissolved concentration adjusts to match whatever partial pressure exists above the solution. Changing the headspace pressure (by opening the bottle) immediately changes what the equilibrium concentration should be, and the system then relaxes toward that new equilibrium by releasing dissolved gas.