Surely on a hot day we have gone to the refrigerator for a soft drink, and when we open the drink we hear the refreshing sound of the release of carbonic gas contained in the bottle, and while we keep the bottle uncovered we continue to observe that wonderful bubbling. Well, while the bottle is closed, the gas is dissolved in the liquid, but as soon as we uncover it, reducing the pressure on the liquid, the solubility of the gas decreases, which is why it escapes from the bosom of the liquid, producing the characteristic effervescence of carbonated drinks. This is due to the effect that pressure has on the solubility of gases.
The decrease in the solubility of carbon dioxide is what produces the effervescence of soft drinks. Source: edited image, original from pxhere.com.
There is a direct relationship between the pressure and the solubility of gases, the solubility being greater the greater the pressure exerted on the surface of the liquid, and there is a mathematical relationship that describes this relationship, and it is known as the law of Henry.
What does Henry's law tell us?
For practical purposes, external pressure has very little influence on the solubility of liquids and solids, but as we well know this is not the case for gases, many properties being affected by changes in pressure, including solubility.
And the quantitative relationship between the pressure and the solubility of gases is expressed by Henry's law, which states that the solubility of a gas in a liquid, at a given temperature, is proportional to the partial pressure of the gas on the solution. . This means that the higher the pressure on the liquid, the greater the tendency of the gas to remain dissolved in the liquid.
Where C is the molar concentration of the dissolved gas, P is the partial pressure of the gas (in atmospheres) above the surface of the liquid, and K is a proportionality constant known as Henry's law constant, which is characteristic of each gas. And it just depends on the temperature.
This simple mathematical expression tells us that the greater the pressure exerted by a gas on a liquid, the greater the amount of that gas that can be dissolved in said liquid, thus increasing its concentration.
The higher the pressure, the higher the concentration of dissolved gas. Source: image elaborated in powerpoint.
Henry's law constant K
Through the formula that describes Henry's law we can realize that, when the gas pressure is equal to atmospheric pressure, the concentration only depends on the constant K, so its value represents the solubility of a gas at the temperature of the gas. system when P = 1 atm.
Henry's law can be interpreted qualitatively as described by the kinetic molecular theory; that is to say, that the quantity of a gas that dissolves depends on the frequency of the collisions of the gas molecules with the surface of the liquid, and that they can condense remaining trapped in the bosom of the liquid.
So the constant K comes to express the degree of interaction between the gas molecules and the liquid molecules, so a large value of this constant refers to strong interaction between these molecules and the greater the solubility of the gas in the liquid.
Interpretation of Henry's Law
Suppose we have a gas trapped in a container and that it is in equilibrium with a solution, when the gas molecules collide with the surface of the liquid phase, some bounce back but others penetrate the surface and are retained in the liquid; the same happens with the gas molecules that are already dissolved, they can cross the surface again and escape to the gas phase.
At equilibrium, at any instant the number of gas molecules entering the solution is equal to the number of molecules escaping from solution, so the concentration of the solution remains constant.
Now, if the pressure exerted by the gas on the solution is greater, more molecules will dissolve, that is, enter the solution, since there are more molecules that collide with the surface of the liquid. And the process will continue until again the number of molecules entering the solution per second equals the number of molecules leaving it. And with this, when the new equilibrium is achieved, the concentration of the solution will be higher.
An increase in pressure forces more molecules into the liquid. Source: image made in powerpoint.
Limitations of Henry's Law
We must keep in mind that this expression is an ideal model that tries to describe the effect of pressure on the solubility of a gas, however, as long as the temperature of the system remains constant, most gases conform well to the law of Henry, although there are some exceptions, and these are the cases in which the dissolved gas reacts with the solvent.
For example, the solubility of ammonia in water is greater than that predicted by Henry's law because this compound reacts with water:
This affinity of ammonia for water increases the concentration of dissolved gas.
Application of Henry's Law
A common example of Henry's law is the effervescence of soft drinks, when the bottles of these drinks are sealed they are subjected to a mixture of CO2 and water vapor under pressure, and due to the high pressure of the gas in the bottle, more CO2 dissolves in the liquid than would dissolve at atmospheric pressure. Therefore, when the bottle is uncovered and the pressure inside is released, the solubility of the gas decreases, and the excess dissolved CO2 escapes, since once the bottle is opened, the concentration of the gas in the liquid is only determined by the normal partial pressure of the CO2, which is equal to 0.0003 atm.
The bubbles in carbonated drinks are produced thanks to this principle. Source: publicdomainvector.org.
Another use case is diving gases. The importance of this law for divers is vital, since by increasing the hydrostatic pressure on the diver, a greater concentration of the gases that make up the air we breathe are dissolved in the blood, and in this nitrogen, which constitutes 78% from the air, it is not consumed by the body like oxygen, and by increasing its concentration in the blood it has narcotic effects and is responsible for decompression sickness in divers. It is estimated that these effects begin to be felt at a depth of about 30 meters, where the depth is 4 atm.
Understanding this principle is vital for divers. Source: pixabay.com.
For example, let's say a diver plans to swim to a depth of 50 meters, where the pressure is close to 6 atm. What is the maximum concentration of nitrogen that he should breathe to avoid narcosis problems?
If we consider that the negative effects of breathing air in normal concentrations begin to appear at a pressure of 4 atm, we can deduce that at that pressure the maximum concentration of nitrogen that can be tolerated in the blood is obtained, which is given by the law from Henry:
And we will use this concentration as the maximum concentration of nitrogen that a diver can breathe in his air mixture at any depth, so at a depth of 50 m:
By equating the two expressions we can find the maximum fraction of nitrogen that a diver can breathe at that depth.
This means that a diver intending to descend to a depth of 50 meters must breathe an air mixture that does not contain more than 52% nitrogen.
Well friends, as we can see, this law has important applications, ranging from the preparation of bubbly drinks to protecting divers from the negative effects of nitrogen that dissolves in the blood.
I hope you liked the information, until next time.
References
Chang Raymond (2002). Chemistry. 7a edition McGraw-Hill, México.