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# Pressure Effects On the Solubility of Gases

The solubility of gases depends on the varying change of pressure, therefore an increase in pressure would equal higher solubility, while a decrease in pressure would mean lower solubility. This statement can be proven by Henry's Law, which states that the solubility of a gas in a liquid is directly proportional to the pressure of that gas above the surface of the solution. This can be expressed in the equation: C = k*Pgas

### Introduction

Variable Represents
C the solubility of a gas in solvent
k the proportionality constant
Pgas the partial pressure of the gas above the solution

 Example

Example Using Henry's Law:
The aqueous solubility at 20 degrees Celsius of Ar at 1.00 atm is equivalent to 33.7 mL Ar(g), measured at STP, per liter of water. What is the molarity of Ar in water that is saturated with the air at 1.00 atm and 20 degrees Celsius? Air contains 0.934% Ar by volume. Assume that the volume if the water does not change when it becomes saturated with air.

STP molar volume: (22.414 L = 22,414 mL)

First Determine Molarity:
kAr = C/PAr
= ((33.7 mL Ar/ 1 L) x (1 mol Ar/ 22,414 mL)) / 1 atm pressure
= 0.00150 M/atm

Solve for Concentration:
C = kArPAr
= (0.0015 M/atm) x 0.00934 atm
= 1.40 x 10-5 M Ar

The illustration above shows:
(A) Low pressure, low concentration of gas solubility. Decreased pressure allows more gas molecules to be present, with very little being dissolved in solution.
(B) High pressure, high concentration of gas solubility. Increased pressure forces the gas molecules into the solution, relieving the pressure that is applied, causing there to be less
gas molecules present and more of it in solution.

Common examples of pressure effects on gas solubility can be demonstrated with carbonated beverages, such as a bottle of soda (above). Once the pressure within the unopened bottle is released, the result will be CO2(g) releasing from the solution as bubbles or fizzing.

### Deep Sea Divers and "the Bends"

In order for deep sea divers to breathe underwater, they must inhale highly compressed air in deep water, resulting in more Nitrogen dissolving in their blood, tissues, and other joints. If the diver returns to the surface too rapidly, the nitrogen gas diffuses out of the blood too quickly and causes pain and possibly death. This condition is also known as "the bends."

To prevent the Bends, one can return to the surface slowly so that the gases will diffuse slowly and adjust to the partial decrease in pressure or breathe a mixture of compressed Helium and Oxygen gas, since Helium is only one-fifth as soluble in blood than Nitrogen. To make this explanation easier to understand, we can think of our bodies under water as a soda bottle under pressure. Imagine dropping the bottle and trying to open it. In order to prevent the soda from fizzing out, you open the cap slowly to let the pressure decrease. On land, you breathe about 78% of nitrogen and 21% oxygen, but our bodies use mostly the oxygen. Under water however, the high pressure of water surrounding our bodies causes nitrogen to form in our blood and tissue. And like the bottle of soda, if we move around or come up from the water too quickly, the nitrogen will be released from our bodies too quickly and creates bubbles in our blood.

### Contributors

• Michelle Hoang (UCD)