Henry's Law
Calculate gas solubility from pressure using Henry's Law: C = kH × P.
Explains soda carbonation, scuba decompression sickness, and oxygen in blood plasma.
The Formula
C = k_H × P
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. Higher pressure forces more gas into solution.
Variables
| Symbol | Meaning |
|---|---|
| C | Concentration of dissolved gas (mol/L) |
| k_H | Henry's law constant (mol/L/atm, specific to each gas and temperature) |
| P | Partial pressure of the gas above the liquid (atm) |
Example 1
CO₂ has k_H = 0.034 mol/L/atm at 25°C. Find dissolved CO₂ at 2 atm pressure.
C = 0.034 × 2
C = 0.068 mol/L
Example 2
O₂ has k_H = 0.0013 mol/L/atm. How much dissolves at sea level (P_O₂ = 0.21 atm)?
C = 0.0013 × 0.21
C = 0.000273 mol/L ≈ 2.73 × 10⁻⁴ mol/L
When to Use It
Use Henry's law when:
- Calculating gas solubility in beverages (carbonation)
- Understanding decompression sickness in scuba diving
- Modeling oxygen and CO₂ exchange in blood
- Designing gas absorption equipment in industry
Limitations
- Henry's law is accurate only for dilute solutions at low pressures — at high concentrations or pressures, solute-solvent interactions cause deviations
- The constant k_H is strongly temperature-dependent: gas solubility generally decreases as temperature rises, which is why warm soda goes flat faster than cold soda
- The law applies only to gases that do not react with the solvent — CO₂ partially reacts with water to form carbonic acid (H₂CO₃), requiring a modified equilibrium treatment
Key Notes
- Formula: C = k_H × P: C is the concentration of dissolved gas (mol/L or mg/L), k_H is Henry's law constant (specific to each gas-solvent pair and temperature), and P is the partial pressure of the gas above the liquid.
- Temperature dependence — opposite of most solids: Gas solubility decreases as temperature increases. Cold water holds more dissolved oxygen than warm water — critical for aquatic life. This is why warm rivers and lakes have lower oxygen levels in summer.
- Decompression sickness (the bends): At depth, divers breathe compressed air; higher pressure dissolves more nitrogen in blood per Henry's law. Rapid ascent reduces pressure suddenly — dissolved N₂ comes out of solution as bubbles in blood and tissues, causing serious injury.
- Carbonation: CO₂ is dissolved in beverages at high pressure (~3–4 atm). Opening a bottle or can drops the pressure to atmospheric, dramatically reducing CO₂ solubility — the excess CO₂ comes out as bubbles.
- Applications: Henry's law is used in wastewater treatment (stripping dissolved gases), blood-gas exchange modeling in the lungs, environmental fate modeling of volatile organic compounds (VOCs), and hyperbaric oxygen therapy design.