While in Europe on a work trip and grabbing a bite for lunch at a café, I grabbed a bottle of water on my way to pay without checking the label. Checking the label is important, because in Europe, "still" water and "sparkling" (carbonated) water are sold side by side—and I can't stand the taste of sparkling water. Halfway through eating lunch, I opened the water bottle to have a drink and it sprayed water all over my tray and my clothes.
I'd grabbed the wrong sort of water. Not only that, I'd obviously shaken it at some point.
Because I'd opened it, I couldn't return it for a bottle of still water, so I decided to de-sparkle the sparkling water, in the hopes that it would improve the taste. Fortunately, this requires no special equipment and can be done in a café, although it might draw some funny looks.
All I had to do was use Henry's Law to my advantage. This law is the equilibrium relationship between the concentration of a gas dissolved in water, and the partial pressure of that gas in the atmosphere above the water.
\[c = P * K_H\]Where:
- c = concentration of CO2 in the liquid
- P = partial pressure of CO2 in the gas
- KH = Henry's constant for CO2 in water, 0.035 mol/kg.bar
Carbonated beverages are super-saturated with CO2 and stored under pure CO2 to keep them that way. Over time, if the bottle is left open all the CO2 migrates out of the liquid and you end up with a flat drink. Since I wanted to get the CO2 out of the liquid completely and without waiting half a day, I went in the opposite direction of the conditions needed to keep it carbonated: I pulled a vacuum above the liquid. As you can see from the equation above, when the pressure above the liquid is low, the concentration in the liquid must be low at equilibrium. If it's not, CO2 will move into the gas phase until the P rises and the c drops enough to make that equation balance.
Actually, for CO2 it's slightly more complicated than that, because CO2 reacts with water once it's dissolved. Instead of this equilibrium:
CO2(g) <=> CO2(aq)
which is the simple case of a gas (g) dissolving in a liquid, in this case water (aq), it's actually:
CO2(g) <=> CO2(aq) <=> H2CO3(aq) <=> HCO3-(aq) <=> CO32-(aq)
which shows all the different forms of CO2 in solution with water. If any one of those compounds is increased or decreased, whether by changing the gas above the surface or adding a chemical to the water, the CO2 will shift through the forms until they're all in balance again.
The water was in a flexible plastic bottle so drawing a vacuum was easy. After drinking a swallow of the horrible-tasting carbonic acid solution to increase the headspace in the vessel—um, I mean water bottle—to something more usable, I squeezed the bottle until the liquid level was up in the neck then capped it tightly and released the squeeze. Instant vacuum. Then I shook the bottle until it returned to its original shape and had air (actually CO2, with a little O2 and water vapour) above it. Then I opened the bottle, squeezed the gas out, and re-sealed it for vacuum. At first this didn't take much shaking at all; I was just rapping the side of the bottle with my knuckles and it was reaching its original shape very quickly. If I had continued agitating it without opening the cap, it would have built up pressure in the bottle just as it had by accident before I opened it the first time. I had to go through quite a few iterations of this; it seems there's a lot of CO2 (and the various forms of carbonate) in sparkling water.
As the concentration of dissolved CO2 decreased and more intense agitation—um, I mean shaking—was required to pull enough CO2 out of solution to return the bottle to its original shape, I had to drink a few more swallows of (less-nasty tasting) water to give myself even more headspace to work with.
At this point, I didn't pull as high of a vacuum over the water because having a bit of air in the bottle not only helps shake up the liquid by giving it space to move into, but having the drops and bubbles form during shaking provides a lot more surface area for the mass transfer of CO2 out of the liquid and into the gas, which makes it reach equilibrium faster. (It doesn't remove more CO2 from solution, it just does what it was always going to do faster.)
This still works despite the extra air and higher starting pressure because Henry's Law links partial pressure of the gas in question, not total pressure of the mixed gases, with its solution concentration. I could have not squeezed the bottle at all and had CO2 come out of solution into the headspace filled with air just as effectively. The reason I didn't is because then the bottle would have ended up pressurized and would have sprayed me again when I opened it!
After several more repetitions of that, the bottle stopped showing signs of returning to its original shape but stayed crumpled with low pressure despite intense shaking, so I tasted it. It had changed to still water and tasted (almost) normal.
(In case anyone is wondering, no, I didn't do any equations in my head while sitting in the café. I did, however, think about Henry's Law, making a vacuum, surface area for mass transfer, and partial pressure vs. total pressure, but in qualitative terms and trends instead of numbers.)
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