Originally posted by: Throwmeabone
As long as it is not opened it doesn't matter what you do to it because the pressure is mainted which keeps the gas dissolved in the liquid. Once its opened, it is better to refrigerate it because gases stay dissolved better in low temperatures.
No. Remember the ideal gas law, PV=nRT, where P=pressure, V=Volume, n= number of atoms of gas present, R= universal gas constant and T= temperature. Therefore P=nRT/V
It says, among other things that, all other things constant, if temperature goes down, pressure goes down, if temperature goes up, pressure goes up.
Anyways, I could've said all this myself, but here is the answer to your question: CUT AND PASTE FROM ANOTHER FORUM BEGINS HERE:
"Anyway, your question relates to the capacity of liquid to hold gasses in solution. There are many factors that affect just how much will dissolve in the first place - composition of the liquid, and type of gas are the most obvious.
A carbonated beverage is supersaturated with CO2. To explain this, I'll have to introduce you to a theory related to the behaviour of ideal gasses across a boundary. The physical law of interest is Henry's Law which states that a gas will permeate (go through) a boundary in order to equalize the pressure of the gas on each side of the boundary. We use this in carbonating beverages by sealing the liquid into a container through which the gass can't pass, but can sustain a high pressure - like steel, for instance. We then apply pressurized CO2 to the open air above the liquid in the container - called the headspace. The liquid, which had previously only seen atmospherric pressure of that gas now see's it greatly increased in the headspace. The CO2 will move through the top of the liquid (the boundary between it and the headspace) until the pressure of CO2 in the liquid is the same as that in the headspace. This saturates the liquid with CO2 at that pressure and temperature.
To understand why the cold liquid can contain more gas than a warm liquid, we look to the ideal gas law from thermodynamics: PV=nRT. The important terms to your work are P (pressure), V (volume - how big), and T (temperature). n is the number of molecules of the gas present and R is the ideal gas constant. The relationship of interest is that the pressure and volume of the system (the gas in its environment) is related to the temperature of that system. You can see that if pressure and volume are held constant, for the equality to be true, the number of gas molecules has to increase as temperature increases (remember: R is a constant). THis proves that as temperature decreases, the amount of gas that can be contained in the environment with pressure P and and volume V must increase.
A couple of more terms: Saturation is where, at a particular temperature, you can get no more of something into another something. For gasses in most liquids, solubility increases as temperature decreases. For solids into liquids, the opposite is true - solubility increases as temperature increases. Supersaturation is where, by some means, you change conditions so a saturated something (liquid in our case) cannot hold all of what's dissolved in it. This can be caused by decreasing the pressure (that's why your soda fizzes when you open it), increase the temperature, or agitate it (why your soda gushes out if you shake the bottle up) - by changing the conditions.
When you open your cold soda, it will quickly let go of the head pressure (the CO2 stored in the open space above the soda in the can or bottle). This is the hiss you hear. In a bottle, you may also see a rush of bubbles through the liquid. This is the CO2 in the soda rushing to reach saturation equilibrium with the atmosphere. As the temperature increases, this saturation equilibrium decreases, causing more gas to leave the solution. To understand this, we need to understand another implication of the ideal gas law: the relationship of pressure to temperature. We can solve PV=nRT to find this relationship: P=nRT/V. Holding n, R, and V constant, and varying T, we see that pressure is directly proportional to temperature. What does this say to us? This: since our soda is colder, initially, than the atmosphere when we open the bottle, the soda will have more molecules of gas in it and remain at pressure equillibrium. We can see this directly from the ideal gas law by solving for the amount of gas in solution at a given temperature - assuming pressure and volume are constant: n = PV/RT. Again, we hold P, V and R constant (for our purposes in thi discussion, we can ignore R as it is a constant in all cases) and find that as temperature DECREASES, the amount of gas, n, increases! Also, we can see from the same equality , that the amount of gas decreases as temperature increases!
Now it all comes together: the volume is the volume of the soda. It is contained by the walls of the glass, and the air-liquid boundary at the top. Henry's law tells us that the gas can pass through the air-liquid boundary in order to equalize the pressure of CO2 apparent on each sideof that boundary. The ideal gas law tells us that a quantity of gas at a low tmperature and constant pressure is greater than the quantity of gas at a high temperature and constant pressure. In your system, using all these concepts, the soda goes flat as the temperature rises because the amount of gas the soda can hold decreases as its temperature increases."
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