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Years ago in a lab where I was working someone had the excellent idea of cooling down their can of coke by pouring liquid nitrogen over it. The can exploded and resulted in quite a mess.

My question is twofold:

(a) Does $\ce{CO2}$ really come out of aqueous solution when water freezes?

(b) If so, why and under what conditions?

I would have expected it to form some kind of clathrate.

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  • $\begingroup$ It would be combination of expulsion of dissoved CO2 and volume expansion of ice. The clathrate seems to need high pressure > 45 bar. $\endgroup$
    – Poutnik
    Commented Sep 26, 2022 at 10:02
  • $\begingroup$ @Poutnik But why would the CO2 be expelled? And given that ice expands, surely there would be more space for CO2? $\endgroup$
    – Dirk Bruere
    Commented Sep 26, 2022 at 10:06
  • $\begingroup$ Space inside the ice structure is unsuitable for CO2. Very few things are soluble in solids. $\endgroup$
    – Ivan Neretin
    Commented Sep 26, 2022 at 11:17
  • $\begingroup$ @IvanNeretin Clathrates $\endgroup$
    – Dirk Bruere
    Commented Sep 26, 2022 at 12:01
  • 3
    $\begingroup$ Wikipedia - CO2 hydrate/clathrate- phase diagram $\endgroup$
    – Poutnik
    Commented Sep 26, 2022 at 12:56

1 Answer 1

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What happens is more a kinetic phenomenon than a thermodynamic one, and more a physical reaction than a chemical one. The carbon dioxide is first introduced into the drink under several atmospheres pressure, so it becomes a supersaturated solution. When the water is cooled and forms ice crystals, the ice crystals act as nuclei for the carbon dioxide gas to "boil". This converts its latent potential (fugacity) to actual pressure, which explodes the can.

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  • $\begingroup$ I suspect that the can explodes from the ice expansion rather than gas pressure? I'm fairly certain that would be the case if the freezing were done slowly. Not sure about a flash freeze. Similarly, in a slow freeze, the growing ice crystals (100% H2O) cause the concentration of all solutes to increase in the remaining liquid volume, thus increasing the equilibrium CO2 pressure in the gas phase. But again, I'm not sure what happens in a flash freeze where rapid amorphous solid formation might be more common than actual crystal growth. $\endgroup$
    – Andrew
    Commented Sep 26, 2022 at 14:55
  • $\begingroup$ Usually there is free space above the liquid in a soda can, so ice expansion has decreased effect $\endgroup$
    – Oscar Lanzi
    Commented Sep 26, 2022 at 15:11
  • $\begingroup$ Hmmm. It becomes supersaturated after the can ruptures or when enough ice is formed. Cooling down a closed can before freezing makes it undersaturated as solubility increases. $\endgroup$
    – Poutnik
    Commented Sep 26, 2022 at 20:38
  • $\begingroup$ It is supersaturated from the start. The carbon dioxide gas is introduced under pressure. Also note what happens when a chilled carbonated dring is poured onto ice (without the can). $\endgroup$
    – Oscar Lanzi
    Commented Sep 26, 2022 at 20:39
  • $\begingroup$ A CO2 bubble would be even better centre than ice particle, if it was supersaturated. But it is in equilibrium at the start at given pressure when closed, as gas solubility is proportional to gas pressure. Pouring is very different situation. It gets supersaturated when the can or bottle of carbonated drink gets opened, because pressure and therefore solubility suddenly drops. $\endgroup$
    – Poutnik
    Commented Sep 27, 2022 at 5:30

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