how does liquid carbon dioxide relate to Chemical Systems in Equilibrium

Question

please I need explanation and long answers

Answer 1

Liquid carbon dioxide is one of the best teaching examples in the whole equilibrium unit, because it lets you watch an equilibrium rather than only calculate one. Here is the full picture.

1. Why liquid CO2 is unusual in the first place

CO2 has a triple point at 5.11 atm and -56.6 degrees C. That number matters: below about 5 atm, liquid CO2 cannot exist at any temperature. At ordinary room pressure solid CO2 goes straight to gas, which is why dry ice sublimes and never leaves a puddle.

So every time you meet liquid CO2 you are looking at a system held under pressure in a sealed container - and a closed container is exactly the condition an equilibrium requires.

2. The phase equilibrium itself

CO2(l) <==> CO2(g)

This is a genuine dynamic equilibrium. Molecules leave the liquid surface and rejoin it continuously; at equilibrium the two rates are equal, so the amounts stop changing even though nothing has stopped happening. The equilibrium constant is just

Kp = P(CO2)

because a pure liquid does not appear in the expression. That single fact gives you the most instructive demonstration in the topic.

3. The pressure gauge that refuses to move

Take a CO2 fire extinguisher or a beverage cylinder at room temperature. Inside is liquid CO2 with vapour above it, reading about 57 atm at 20 degrees C.

Now use half of it. The pressure is still about 57 atm. Use three quarters - still 57 atm. The gauge does not begin to fall until the last of the liquid is gone.

That is Kp = P(CO2) made visible. As long as any liquid remains the system restores the equilibrium vapour pressure, and how much liquid is left is irrelevant because it does not appear in K. It is also why you weigh a CO2 cylinder instead of reading its gauge to find out how full it is.

4. Le Chatelier applied to a physical equilibrium

Raise the pressure: the system shifts toward the side occupying less volume, which is the liquid. Compress CO2 gas hard enough at room temperature and it condenses - that is how the cylinder was filled.

Raise the temperature: vaporization is endothermic, so heating shifts the equilibrium toward gas and the vapour pressure climbs steeply. This is why a CO2 cylinder left in a hot car is genuinely dangerous, and why such cylinders carry burst discs.

Drop the pressure suddenly: the liquid flashes to gas, and because vaporization absorbs heat the remaining material cools sharply. Discharge a CO2 extinguisher and dry ice snow forms at the nozzle - the equilibrium shifting right and cooling its own surroundings past the freezing point.

5. Where the equilibrium stops existing

Above 31.0 degrees C and 72.9 atm, CO2 becomes supercritical. The meniscus between liquid and gas disappears and the two phases merge into one. There is no longer a liquid-gas equilibrium, because there is no longer a boundary.

That state matters industrially: supercritical CO2 dissolves things like a liquid but diffuses like a gas, and it is what decaffeinates coffee and dry-cleans clothes without chlorinated solvents.

6. The link to solution and acid-base equilibria

The same molecule gives you the other classic example. In a sealed fizzy drink:

CO2(g) <==> CO2(aq)

governed by Henry's law - the dissolved concentration is proportional to the pressure above the liquid. Open the bottle, the CO2 pressure collapses, the equilibrium shifts left, and the drink fizzes. Leave it open and it goes flat, because the system is no longer closed and can never reach equilibrium again.

Then it continues into a chemical equilibrium:

CO2(aq) + H2O(l) <==> H2CO3(aq) <==> H+(aq) + HCO3-(aq)

which is why carbonated water is mildly acidic and, at planetary scale, why rising atmospheric CO2 acidifies the oceans. Push more CO2 into the left of that chain and Le Chatelier drives it right, raising the hydrogen ion concentration.

The summary sentence for your write-up

Liquid CO2 shows that equilibrium is not only about reactions: any reversible process in a closed system reaches a dynamic balance, obeys an equilibrium constant, and responds predictably to changes in pressure and temperature. CO2 is the ideal case study because one substance demonstrates a phase equilibrium, a solubility equilibrium and an acid-base equilibrium - and because with a cylinder and a gauge you can see the constant with your own eyes.