Henry’s Law – Definition, Formula, and Applications

Henry’s law states that, at a constant temperature, the amount of a gas dissolved in a liquid is directly proportional to the partial pressure of that gas above the liquid. This simple relationship provides a powerful way to understand how gases behave in solutions, especially in environmental chemistry, physiology, and industrial processes.

Henry’s law applies most accurately to dilute solutions and gases that do not chemically react with the solvent. It connects measurable quantities, pressure and concentration, making it a foundational principle in physical chemistry and chemical engineering.

Key Takeaways Henry’s Law

• Henry’s law relates gas solubility to gas pressure above a liquid.
• As pressure increases, the amount of dissolved gas increases proportionally.
• The law applies best to dilute solutions and non-reactive gases.
• The proportionality constant is called the Henry’s law constant.
• The value of the constant depends on temperature and the gas–solvent pair.
• Gas solubility generally decreases as temperature increases.
• Henry’s law has applications in respiration, carbonated beverages, and environmental science.

History of Henry’s Law

William Henry, an English chemist, first formulated Henry’s law in 1803. He conducted experiments measuring how gases dissolve in water under varying pressures. His work demonstrated that the concentration of a dissolved gas increases linearly with pressure, provided temperature remains constant.

Henry’s observations built on earlier work by scientists studying gas behavior and solution chemistry. Over time, the law became integrated into thermodynamics and statistical mechanics, where it is now understood as a limiting case of gas–liquid equilibrium.

The Henry’s Law Equation

The most common mathematical expression of Henry’s law is:
$$C = k_H P$$

Where:

• $C$ = concentration of the dissolved gas (mol/L or similar units)
• $k_H$​ = Henry’s law constant
• $P$ = partial pressure of the gas (atm, bar, or Pa)

This form shows that concentration increases linearly with pressure.

An alternative form is:$$P = k_H’ x$$

Where:

• $x$ = mole fraction of the gas in solution
• $k_H’$​ = a different form of the Henry’s law constant

Different forms exist depending on how concentration is expressed.

Types of Henry’s Law Constants and Units

Henry’s law constants vary depending on how the relationship is written. The most common types include:

1. Concentration-Based Constant

$$C = k_H P$$

• Units of $k_H$​: mol·L⁻¹·atm⁻¹ or mol·m⁻³·Pa⁻¹

2. Pressure-Based Constant

$$P = k_H’ x$$

• Units of $k_H’$​: atm or Pa

3. Dimensionless Constant

$$H = \frac{C_{gas}}{C_{liquid}}$$

• No units

Important Note

The value of $k_H$​ depends strongly on the definition used. Always check the form of the equation before using a tabulated constant.

Table of Henry’s Law Constants for Common Gases

Henry’s law constants vary widely depending on the gas and temperature. The values below are representative at 25 °C for gases dissolved in water. Units are given in mol/(L·atm) for consistency with the form C = kH × P.

Notes

• Higher kH values indicate greater solubility in this formulation.
• Reactive gases (such as NH₃ and SO₂) show unusually high apparent solubility because they chemically interact with water.
• Values depend strongly on temperature and the specific definition of kH used.

Temperature Dependence of Henry’s Law

Temperature has a significant effect on gas solubility:

• As temperature increases, gas solubility typically decreases.
• This occurs because gas dissolution is often exothermic.

The relationship can be expressed using a van’t Hoff-type equation:
$$k_H(T) = k_H(T_0) \exp \left[ -\frac{\Delta H_{sol}}{R} \left( \frac{1}{T} – \frac{1}{T_0} \right) \right]$$

Where:

• $\Delta H_{sol}$ = enthalpy of solution
• $R$ = gas constant
• $T$ = temperature in Kelvin

This explains why warm soda loses carbonation faster than cold soda.

Molecular Interpretation of Henry’s Law

Henry’s law can be understood at the molecular level by considering the interactions between gas molecules and the solvent.

When a gas is in contact with a liquid, gas molecules continuously collide with the liquid surface. Some molecules enter the liquid and dissolve, while others escape back into the gas phase. At equilibrium, the rate of dissolution equals the rate of escape.

Increasing the gas pressure increases the number of gas molecules striking the liquid surface per unit time. This increases the rate at which gas molecules enter the liquid, resulting in a higher equilibrium concentration of dissolved gas.

The extent to which a gas dissolves depends on intermolecular forces:

• Weak interactions (for example, nonpolar gases in water) lead to low solubility
• Stronger interactions (for example, polarizable gases) increase solubility

Henry’s law works best when:

• Gas–solvent interactions are relatively weak
• No chemical reaction occurs between gas and solvent

If a gas reacts with the solvent, such as carbon dioxide forming carbonic acid in water, the observed solubility deviates from simple Henry’s law behavior.

From a thermodynamic perspective, Henry’s law represents a linear relationship that arises in the limit of low concentrations, where interactions between dissolved gas molecules are negligible.

Dependence on Ionic Strength (Salting-Out Effect)

Dissolved salts affect gas solubility in water:

• Increasing ionic strength decreases gas solubility.
• This effect is called the salting-out effect.

Ions in solution compete with gas molecules for interactions with water molecules, reducing the ability of gases to dissolve.

The effect is described by the Setschenow equation:
$$\log \left( \frac{C_0}{C} \right) = k_s [salt]$$

Where:

• $C_0$​ = solubility in pure water
• $C$ = solubility in salt solution
• $k_s$​ = Setschenow constant

Importance and Applications of Henry’s Law

Henry’s law plays a central role in understanding how gases interact with liquids in both natural systems and engineered processes. Because it directly links gas pressure to dissolved concentration, it provides a practical framework for predicting and controlling gas solubility under real-world conditions. This makes it especially valuable in fields ranging from environmental science and physiology to chemical engineering and industrial chemistry. Whether modeling oxygen levels in aquatic ecosystems, explaining gas exchange in the lungs, or designing processes that rely on gas absorption or release, Henry’s law offers a simple but powerful tool for analyzing gas–liquid equilibrium.

Environmental Science

• Predicts how gases like oxygen, carbon dioxide, and pollutants dissolve in water bodies.
• Helps model gas exchange between the atmosphere and oceans.

Physiology and Medicine

• Explains gas exchange in lungs and blood.
• Important in understanding decompression sickness (the bends).

Carbonated Beverages

• High pressure dissolves CO₂ in soda.
• When opened, pressure drops and gas escapes.

Industrial Applications

• Used in gas scrubbing and purification processes.
• Important in chemical reactors involving gas–liquid systems.

Real-World Case Study: Decompression Sickness (The Bends)

Decompression sickness provides one of the clearest and most important real-world examples of Henry’s law in action. It occurs when dissolved gases, primarily nitrogen, come out of solution in the body due to a rapid decrease in pressure.

When a diver descends underwater, the surrounding pressure increases. According to Henry’s law, the higher pressure causes more nitrogen from the breathing gas to dissolve into the blood and tissues. While at depth, this process reaches equilibrium, and the body safely contains the dissolved gas.

Problems arise during ascent. If a diver rises too quickly, the pressure decreases faster than the body can eliminate the dissolved nitrogen through respiration. As a result, nitrogen comes out of solution and forms bubbles within tissues and the bloodstream.

These bubbles can cause a range of symptoms, including:

• Joint pain (the classic “bends”)
• Dizziness and fatigue
• Breathing difficulties
• Neurological symptoms in severe cases

Prevention relies on controlling the rate of pressure change:

• Slow ascent allows nitrogen to leave the body gradually
• Decompression stops give time for safe gas elimination
• Breathing gas mixtures can reduce nitrogen uptake

This case study demonstrates Henry’s law in a physiological context: increased pressure increases gas solubility, while decreased pressure causes gas to escape from solution.

Simple Demonstration: Carbonated Beverage and Pressure

This demonstration clearly illustrates Henry’s law by showing how dissolved carbon dioxide depends on pressure.

Materials

• A sealed bottle or can of soda or sparkling water (chilled works best)
• A clear glass
• Optional: a balloon

Procedure

1. Observe the sealed container
   Look at the unopened soda. You should not see many bubbles forming inside.
2. Open the container
   Slowly open it and listen for the hiss. This sound is carbon dioxide gas escaping as pressure decreases.
3. Pour into a glass
   Pour the soda into a glass and observe the rapid formation of bubbles.
4. (Optional) Balloon variation
   Carefully stretch a balloon over the mouth of a freshly opened bottle. The balloon may inflate slightly as CO₂ escapes.

What’s Happening

• In the sealed bottle, carbon dioxide is under high pressure.
• According to Henry’s law, higher pressure increases the amount of gas dissolved in the liquid.
• When you open the container, the pressure above the liquid drops suddenly.
• The dissolved CO₂ is no longer as soluble, so it escapes as bubbles.

Key Takeaway

High pressure keeps gas dissolved, while lower pressure causes gas to come out of solution.

Optional Extension: Temperature Effect

Try the same experiment with:

• A cold soda
• A warm soda

You’ll observe:

• Warm soda loses carbonation much faster
• This shows how temperature also affects gas solubility, complementing Henry’s law

Limitations of Henry’s Law

Henry’s law is not universally applicable. Its limitations include:

• Valid only for dilute solutions
• Assumes no chemical reaction between gas and solvent
• Works best at low to moderate pressures
• Deviates at high gas concentrations
• Less accurate for highly soluble gases (e.g., ammonia, hydrogen chloride)

Worked Example Problems

Example 1: Gas Solubility at a Given Pressure

A gas has a Henry’s law constant k_H = 3.3 x 10^-2 mol/(L·atm).
Calculate the concentration when the gas pressure is 2.0 atm.

Solution:$$C = k_H P = (3.3 \times 10^{-2})(2.0)$$

C = 6.6 x 10^-2 mol/(L·atm)

Example 2: Finding Pressure from Concentration

A solution contains 0.015 mol/L of a dissolved gas.
The Henry’s constant is 1.5×10⁻² mol/(L·atm).
Find the pressure.

Solution:$$P = \frac{C}{k_H} = \frac{0.015}{1.5 \times 10^{-2}}$$$$P = 1.0 \, \text{atm}$$

Henry’s Law vs Raoult’s Law

Henry’s law and Raoult’s law are both important relationships in solution chemistry, but they describe different components and conditions within a solution. Understanding the distinction helps clarify when each law applies.

Henry’s law describes the behavior of a gas (or a solute present at low concentration), while Raoult’s law describes the behavior of a solvent in an ideal solution. Together, they provide complementary descriptions of vapor–liquid equilibrium.

Comparison Table: Henry’s Law vs Raoult’s Law

This comparison highlights that Henry’s law governs how gases dissolve, while Raoult’s law governs how liquids contribute to vapor pressure.

Common Misconceptions

• “Henry’s law applies to all gases equally.”
  The constant varies widely depending on the gas and solvent.
• “Temperature does not matter.”
  Temperature strongly affects solubility and the constant.
• “The law works at all concentrations.”
  It is only valid for dilute solutions.
• “Higher pressure always means more gas dissolves indefinitely.”
  Deviations occur at high pressures.

FAQs

What is Henry’s law in simple terms?

It states that increasing the pressure of a gas above a liquid increases how much of that gas dissolves in the liquid.

Why does soda go flat when opened?

Opening the container reduces pressure, causing dissolved CO₂ to escape.

Does Henry’s law apply to all gases?

No, it works best for gases that do not react chemically with the solvent.

Why does warm water hold less oxygen than cold water?

Higher temperatures reduce gas solubility, which is consistent with Henry’s law behavior.

What is the Henry’s law constant?

It is the proportionality constant that relates gas pressure to solubility. Its value depends on the gas, solvent, and temperature.

References and Further Reading

• Felder, Richard M.; Rousseau, Ronald W.; Bullard, Lisa G. (2004). Elementary Principles of Chemical Processes. Wiley. ISBN 978-0471687573.
• Gillispie, Charles Coulston (1960). The Edge of Objectivity: An Essay in the History of Scientific Ideas. Princeton University Press. ISBN 0-691-02350-6.
• Henry, W. (1803). “Experiments on the quantity of gases absorbed by water, at different temperatures, and under different pressures”. Phil. Trans. R. Soc. Lond. 93: 29–43. doi:10.1098/rstl.1803.0004
• Mackay, Donald; Shiu, Wan Ying (1981). “A critical review of Henry’s law constants for chemicals of environmental interest”. Journal of Physical and Chemical Reference Data. 10 (4): 1175–1199. doi:10.1063/1.555654
• Sander, R.; Acree, W. E.; De Visscher, A.; Schwartz, S. E.; Wallington, T. J. (2022). “Henry’s law constants (IUPAC Recommendations 2021)”. Pure Appl. Chem. 94: 71–85. doi:10.1515/pac-2020-0302