[davedave]

I've had this argument so many times with so many people I've almost lost the will to argue my case anymore and am starting to doubt myself. But I am a little rusty on my chemistry, so here it goes.

I believe the best way to keep a drink fizzy (AFTER opening) is to keep it room temperature NOT put it in the fridge, and also to NOT squeeze it.

1) Lets say you have half a bottle of carbonated drink left. Am I right in saying that to not squeeze it means there is more gas in the bottle, if there is more gas then less aqueous CO2 needs to transfer to gaseous CO2 in order to reach equilibrium and prevent further "fizz loss"

2) Am I also right in saying that if you then put your bottle in the fridge, the pressure of gas would decrease, in turn resulting in the loss of equilibrium allowing more aqueous CO2 to transfer to gaseous CO2, causing "fizz loss"

If I am correct in what I am saying, can someone please help me by giving me a way of explaining this to people, because everyone I know (and I mean everyone) squeezes and puts it in the fridge and when I ask them why, they say "To keep the fizz!" No one will listen to me because they think I'm crazy as its something they were brought up doing

One more thing, If i am right, is anyone interested in starting a "Keeping The Fizz Awareness Programme". I really do think this is a major problem in today's society that needs to be addressed.

p.s. If I'm wrong, then I have a few apologies to make to certain people.

[DevaDevil]

1) Lets say you have half a bottle of carbonated drink left. Am I right in saying that to not squeeze it means there is more gas in the bottle, if there is more gas then less aqueous CO2 needs to transfer to gaseous CO2 in order to reach equilibrium and prevent further "fizz loss"

Henry's law: p_CO2 = k_H*C_CO2

Assume you have lots of CO₂-free gas above your drink, in the bottle. Then you need to evolve lots of CO₂ to reach equilibrium. (partial pressure is related to amount dissolved, lots of gas means partial pressure will be higher in absolute value, aka in quantity of molecules)

2) Am I also right in saying that if you then put your bottle in the fridge, the pressure of gas would decrease, in turn resulting in the loss of equilibrium allowing more aqueous CO2 to transfer to gaseous CO2, causing "fizz loss"

low pressure of gas = good.
Also, solubility of gases goes up when temperature goes down.



my advice: keep drink in appropriately sized container, in the fridge.

[Enthalpy]

And I believe gaseous drinks are out of equilibrium because gas pressure augments irreversibly in the closed bottle if you shake it.

So keeping the drink fizzy would not be a matter of partial pressure but of evolution speed, hence logics is much more complicated.

[azmanam]

The one correction I would make to your argument is in point 1.  DevaDevil already touched on it.  It's not the amount of gas above the liquid that's important - it's the amount of CO₂ in that gas that's important.

I always assumed the dissolved CO₂ was trying to reach equilibrium with the entire atmosphere once opened... which is why CO₂ continues to evolve as gas until the drink is flat.  Thus, imho, (with no maths involved) I would say the best way to maintain an acceptable amount of dissolved CO₂ would be to make the volume with which the gas is trying to equilibrate as small as possible (squeeze the bottle and put the cap back on). 

Also, if you wanted to be super cool, you could get a CO₂ tank or some other source of CO₂ and bubble it through your drink before capping.  This would minimally dissolve some more CO₂ and purge the air from the container and replace it with CO₂.  The converse of what devadevil was saying - you'll have lots of CO₂ gas above your drink and it won't need to evolve much to maintain equilibrium.

I'm willing to be wrong, though, because like I said no maths were involved in the making of this opinion.

[Jasim]

If you are squeezing the bottle, then aren't you creating a vacuum pressure?

Let's talk about pressures here. La Chatlier's principle states (from Wikipedia) "If a chemical system at equilibrium experiences a change in concentration, temperature, volume, or partial pressure, then the equilibrium shifts to counteract the imposed change and a new equilibrium is established."

I like to think of La Chatlier's in comparison to water flowing downhill or a balloon deflating. Let's consider the various possibilities...

* Removing the gas above the surface of the liquid...
The bottle can expand to a given volume and will do so if allowed. Gas bubbles in the liquid are already under pressure (when you first open the cap you see the bubbles form and come to the surface causing fizzing). Squeezing the bottle and then placing the cap on will remove most of the air above the surface of the liquid, but the bottle can still expand. This could create a slight negative pressure which may cause (La Chatlier's principle) the pressurized gas in the liquid to come out of solution to fill the void.

* Reducing the space above the surface of the liquid...
If less space is available for the gas to fill, then it will remain in solution. The best scenario would be to completely fill a container and seal it. In that case, no gas would be able to evolve from the container.

* Pressurize the container...
This is actually what they do during the manufacturing process. When you open a fresh bottle of coke it will fizz and foam. This occurs because you are releasing pressure inside the container. The pressure was keeping the dissolved CO2 gas inside solution. Obviously, the best solution to the problem (other than completely filling a container) would be to pressurize the container with CO2, but let's consider a more realistic option. They sell 2-liter bottle hand-pump air pressurizers. By pumping air into the space in the bottle above the liquid it creates a pressurized barrier for the evolution of CO2. According to La Chatlier's principle this will work.

* Reducing the temperature of the container...
Placing the container in the refrigerator versus leaving it out on the table will reduce the temperature of the system. Remember from general chemistry that higher temperature results in greater solubility - solubility in both gas and liquid. Since the CO2 gas is already under pressure to be in the liquid, it wants to be liberated gas above the liquid. Reducing the molecular movement (and thereby the solubility via reducing the temperature) should result in slowing the movement of the pressurized CO2 to the gaseous space above the liquid.


So, I would definitely put the container in the fridge. Squeezing the air out is based on a valid idea of reducing the space, but I would leave it alone or better yet put it in a smaller container for fear of creating a vacuum pressure.

The idea that CO2 in the gas is important is partially right, but one must consider if the other, non-CO2, constituents of the gas will exchange into the liquid. If they will, then the amount of CO2 above the liquid is relevant, but if they won't then the pressure will override species-specific concerns. Air will not go into solution, so squeezing the bottle will cause more harm than good.

[vmelkon]

If you are squeezing the bottle, then aren't you creating a vacuum pressure?

I'm going to have to agree with that. If you squeeze the bottle and close the cap, the plastic bottle wants to expand. Also, the CO2 in the drink has a tendency to dissipate. Thus, the bottle will expand until it reaches its maximum but CO2 will keep evolving until enough pressure builds up.

I use to squeeze the bottle when I was a kid. What a fool I was.