Henry’s law

Clinical relevance in anaesthesia

Explains how much gas dissolves in blood/tissues at a given partial pressure (e.g. volatile agents, CO2, O2, N2, N2O).
Underpins blood–gas partition coefficient and uptake/washout of inhalational agents.
Key to decompression illness and gas embolism: dissolved inert gas comes out of solution when ambient pressure falls.
Explains why hyperbaric oxygen increases dissolved O2 content beyond that carried by Hb.
Relevant to storage/handling: CO2 in soda water; also to blood gas analysis concepts (dissolved vs combined forms).

Statement of Henry’s law

At constant temperature, the amount (concentration) of a gas dissolved in a liquid is directly proportional to the partial pressure of that gas in contact with the liquid.
Form: C = kH × P (or P = C / kH depending on convention).
- C: concentration of dissolved gas in the liquid (e.g. mol·L−1).
- P: partial pressure of the gas above the liquid (e.g. kPa or mmHg).
- kH: Henry’s constant (depends on gas, solvent, and temperature; units depend on the chosen form).
For mixtures, each gas dissolves according to its own partial pressure (Dalton’s law provides partial pressures; Henry’s law provides dissolution).

Temperature and solubility

For most gases in aqueous solutions, solubility decreases as temperature increases (warming drives gas out of solution).
- Clinical link: warming IV fluids/blood reduces dissolved gas content; warming a fizzy drink causes CO2 to come out of solution.
Henry’s law applies at a fixed temperature; changing temperature changes kH.

Solubility coefficient, partition coefficients, and anaesthetic uptake

Solubility coefficient (often used clinically): volume of gas dissolved per unit volume of liquid at a stated temperature and partial pressure (definitions vary; always state conditions).
Blood–gas partition coefficient (λB:G): ratio of concentration in blood to concentration in alveolar gas at equilibrium; reflects solubility in blood.
- Higher λB:G → more agent dissolves in blood for a given alveolar partial pressure → slower rise in FA/FI → slower induction and recovery.
- Lower λB:G → less dissolves → faster alveolar partial pressure rises → faster onset/offset.
Tissue–blood partition coefficients similarly reflect solubility in tissues and influence context-sensitive washout (fat solubility).

Blood gases: dissolved vs chemically bound

Henry’s law relates partial pressure to dissolved gas only (e.g. dissolved O2 and dissolved CO2).
O2: most is bound to Hb; only dissolved O2 contributes directly to PO2 (and thus drives diffusion).
- Dissolved O2 content ≈ 0.003 mL O2·dL−1·mmHg−1 × PaO2 (at 37°C).
CO2: substantial is converted to bicarbonate/carbamino compounds; partial pressure reflects dissolved CO2, but total CO2 content includes combined forms.

Pressure changes: diving, decompression, and embolism

Increased ambient pressure increases partial pressure of inspired inert gases (e.g. N2) → more dissolves in tissues (Henry’s law).
Rapid decompression reduces partial pressure → dissolved gas comes out of solution forming bubbles (decompression sickness).
- Risk increased with high tissue inert gas load, rapid ascent, cold, exertion, dehydration, and right-to-left shunts (arterialisation of bubbles).
N2O can expand pre-existing gas spaces (not Henry’s law directly, but related concepts of partial pressure gradients and gas transfer).
- Expansion is better explained by diffusion and relative solubility (N2O more soluble than N2), but Henry’s law still describes dissolved amount at a given partial pressure.

Limitations and assumptions

Best applies to dilute solutions and ideal behaviour; deviations occur with high pressures, reactive gases, or when gas chemically combines with the solvent.
Must specify temperature; kH varies with temperature and solvent composition (e.g. blood vs water).

State Henry’s law and define the terms used.

A common primary FRCA viva prompt: give the law, then define concentration, partial pressure, and the constant.

At constant temperature, the concentration of a gas dissolved in a liquid is proportional to the partial pressure of that gas above the liquid.
C = kH × P (or P = C/kH depending on convention).
C = dissolved concentration; P = partial pressure; kH = Henry’s constant (gas/solvent/temperature dependent).

How does Henry’s law relate to the blood–gas partition coefficient and speed of induction with volatile agents?

This is a frequent link question: translate ‘solubility’ into ‘alveolar partial pressure rise’.

Henry’s law describes how much anaesthetic dissolves in blood at a given partial pressure.
Blood–gas partition coefficient (λB:G) reflects solubility in blood: higher λ means more dissolves for the same alveolar partial pressure.
Higher solubility → greater uptake into blood → slower rise in alveolar fraction (FA/FI) → slower induction and recovery.
Lower solubility → less uptake → faster rise in alveolar partial pressure → faster onset/offset.

A patient is treated with hyperbaric oxygen. Using Henry’s law, explain what happens to dissolved oxygen content and why it matters clinically.

Examiners want: increased PO2 increases dissolved O2 linearly; dissolved O2 can become clinically significant at high PaO2.

Increasing inspired/ambient pressure increases alveolar and arterial PO2.
By Henry’s law, dissolved O2 in plasma increases in proportion to PaO2.
Although most O2 is Hb-bound at normal pressures, at very high PaO2 the dissolved fraction can contribute meaningfully to total O2 content and tissue delivery (e.g. CO poisoning, problem wounds).

Explain decompression sickness using Henry’s law. What factors increase risk?

Classic physics-to-clinical viva: dissolution under pressure, bubble formation on decompression, and risk factors.

At depth, increased ambient pressure increases partial pressure of inert gas (mainly N2) → more dissolves in tissues (Henry’s law).
On rapid ascent, ambient and alveolar partial pressures fall → tissues become supersaturated → gas comes out of solution as bubbles.
Risk factors: depth and duration of dive, rapid ascent/inadequate decompression stops, cold, exertion, dehydration, obesity (fat stores), and right-to-left shunt (arterial gas embolisation).

What is the relationship between partial pressure and dissolved gas, and why does partial pressure drive diffusion rather than total gas content?

They are probing that only dissolved gas contributes to partial pressure; bound forms do not directly contribute.

Partial pressure reflects molecules in solution that are free (dissolved) and able to exert pressure; Henry’s law links this to dissolved concentration.
Diffusion across membranes is driven by partial pressure gradients, not by total content (which may include bound/combined forms).
Example: O2 content is mostly Hb-bound, but PaO2 reflects dissolved O2 and determines diffusion into tissues.

How does temperature affect gas solubility in liquids? Give a clinical or everyday example.

Often asked as a quick extension: direction of change and an example.

For most gases in water/blood, solubility decreases as temperature increases (kH changes with temperature).
Examples: warming a carbonated drink releases CO2; warming blood/fluids reduces dissolved gas content.

In blood gas analysis, what does PaO2 represent in terms of Henry’s law? How would you estimate dissolved oxygen content?

A common applied physiology/physics crossover.

PaO2 reflects the partial pressure of dissolved O2 in plasma (not Hb-bound O2).
Dissolved O2 content ≈ 0.003 mL O2·dL−1·mmHg−1 × PaO2 (at 37°C).
Total O2 content = Hb-bound (1.34 × Hb × SaO2) + dissolved component.

A viva question: ‘Does Henry’s law explain why nitrous oxide expands a pneumothorax?’ Discuss.

Examiners like a nuanced answer: Henry’s law describes dissolution; expansion is mainly diffusion driven by partial pressure gradients and relative solubility.

N2O expansion of closed gas spaces is primarily due to diffusion of N2O into the space faster than N2 can leave, driven by partial pressure gradients and higher blood solubility of N2O compared with N2.
Henry’s law is relevant in describing how much N2O is carried dissolved in blood at a given partial pressure, but it does not by itself describe the net volume change of a trapped gas space.
A good answer links Henry’s law + Fick’s law concepts: dissolved amount (Henry) enables transport; diffusion down gradients (Fick) causes expansion.

What are the assumptions/limitations of Henry’s law that might matter in clinical contexts?

Often asked to test deeper understanding beyond rote statement.

Applies best to dilute solutions and ideal behaviour; at high pressures or non-ideal conditions, proportionality may deviate.
Temperature must be constant; kH is temperature dependent.
If gas chemically reacts/combines (e.g. CO2 hydration, O2 binding to Hb), Henry’s law only describes the dissolved fraction, not total content.