So I was watching CO2 bubble out of a glass of Sprite yesterday, and it occurred to me that there is a lot of gas dissolved in the liquid without appreciably increasing the volume of said liquid. Of course, we all learned the ideal gas law as young 'uns–the pressure of a given amount of gas is inversely proportional to volume, if temperature is held constant. But the gas is just hanging out, dissolved, in the Sprite, at drastically decreased volume, but at 1atm of pressure.
What gives? The gas exerts pressure after it comes out of solution, as everyone knows who has every opened a soda can after shaking it. But why doesn’t it exert any appreciable pressure while it is dissolved? I was not able to understand any explanation on the interwebz.
It’s not really correct, but you can think of it as behaving like a liquid when it’s solvated. Think of a liquid right around it’s boiling point in a closed vessel. As a liquid, the molecules are all tugging on each other instead of banging on the walls and exerting pressure. Heat it up a few degrees and the pressure spikes.
When a gas is dissolved [in a liquid or solid], it is stuck to the [molecules] of the [other liquid or solid]. “Stuck to” like surface tension.
“Surface tension” is just called ‘surface’ tension because the difference between inside and outside is what you can measure, and is often what you care about - water drops form drops because the inside is more stuck to the inside than it is stuck to the outside.
And because they stick, then can stack into a denser structure.
One example of this is water vapour - a gas. When disolved into water, it forms a compound called “water”, where the gas molecules are attracted to the water molecules enough to stick it all together into a denser structure. Yes, of course the gas molecules are exactly the same molecules as the water molecules, but it works exactly the same way with a mixture of 02 and H20 – there is an attraction between the molecules that enables the 02 to stack into a denser structure.
A ‘perfect’ gas is defined as one where there is no stickiness at all. The gas molecules just bounce off each other.
It does exert pressure while it’s dissolved. In a closed container, an equilibrium will exist between the pressure of the gas above the liquid, and the gas dissolved in the liquid. You can prove it:
grab a 2-liter bottle of soda. Observe that its skin is taut because it’s under pressure. That’s the equilibrium that I just described.
open the cap. The gas above the liquid escapes, achieving a new equilibrium with the atmosphere. There’s still a bunch of CO2 dissolved in the liquid; it wants to come out of solution, but it can’t do so very quickly because there’s not much free surface area.
close the cap again, and then either wait for a few hours or shake the hell out of the soda. Pressure will build again inside the bottle, and you will achieve a new equilibrium twixt the dissolved and undissolved CO2.
You can repeat steps 2-3 a number of times before there’s very little CO2 left in the bottle.
You’re right, you can shove more gas into a fixed volume of liquid and produce increased pressure. Having said that, I don’t know the relationship between dissolved-gas mass concentration and pressure, so I couldn’t say whether it’s a strict linear relationship like the ideal gas law. Hopefully a chemist will chime in here.
Co2 is a little bit unlike N2 or O2 or He in water because CO2 reacts to form carbonic acid. Henry’s law predicts quite well how much CO2 is dissolved in water for the common temperatures / pressures / concentrations you encounter in everyday life.
As a side remark - most aquatic life depend on oxygen dissolved in water to survive - so the law helps you get insight into that too.
The problem here is the use of the term “gas”. We have become accustomed to thinking of certain compounds as gases, because they exist as a gas in uncombined form at standard temperature and pressure. But they behave as gases only while in the gaseous state. They cease to be gases when they are dissolved, cooled, etc., and they stop behaving according to gas law when they no longer exist in a gaseous state.
For example, we think of Oxygen as a gas, but when it combines with Iron to form rust on your car door sills, it is no longer bound by gas law.
Some physical chemists are proposing Henrietta’s Law, which will require that the International Union of Pure and Applied Chemistry (IUPAC) maintain a publicly viewable registry of gases that display deviant solubility effects.
Note also that gases don’t always obey the ideal gas law either. As the name implies, it is an idealization. It is pretty accurate for many gases, but becomes especially inaccurate near the condensation temperature (a.k.a. boiling point).
This is helpful. I had read the entry on Henry’s Law already, but while it provided a law governing the behavior of dissolved gases, it didn’t provide an explanation for why they behaved that way. My (8th-grade) understanding of how gas exerts pressure on a container is that the pressure results from the collision of constantly-moving gas molecules against the walls of the container. So I was puzzled–do dissolved gases quit moving? Henry’s Law doesn’t explain that (or the Wikipedia entry doesn’t explain that.)
Signed,
Sophistry “Just Enough Knowledge to Sound Stupid” and Illusion