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Tutorial: Chemical Equilibrium

Part 7- Le Châtelier's Principle III- Pressure

There are a number of movies available on this page. I've given you a choice of files in most cases since movies can be very large and take forever to download. If you have a good connection (you're at achool, ISDN, fast modem connection, etc.) then choose the "high res" file. If you have a slow connection then choose the "low res" file.

Pressure Changes and Chemical Equilibrium

There are two ways to change the pressure of a system that has gaseous reactants and/or products. One is to add/remove one of the gaseous reactants or products and the other is to change the volume. Adding/removing one of the species gives a similar result as concentration and temperature change. Add one and the stress in the system is relieved by shifting to the opposite side. Remove one and the stress in the system is relieved by shifting to that same side.

Volume needs to be handled a little differently. You need to first consider how the volume change will affect pressure and then consider how the change in pressure affects the chemical equilibrium. Let's look back at Boyle's Law which will give us a relationship between pressure and volume.

Pressure is inversely propertional to the volume:

And this gives us Boyle's Law:

The inverse relationship between volume and gas pressure is what we need. Increase volume and pressure is decreased. Decrease volume and pressure is increased. Let's use this to dispel a common myth about soda bottles and keeping the soda from going flat. It's a common belief that a soda's "fizziness" can be maintained by squeezing the bottle and evacuating the most of the air before tightly capping it.

Ever done this before? I have. It seems logical, doesn't it? The soda stays fizzy before you initially open it because of the high CO₂ pressure above the solution.

I poured out about 20% of the soda and then squeezed the bottle until the volume level for the liquid was approximately the same before pouring out the liquid. What's happening in the room temperature, time-lapsed video? Recall that carbonated water is carbon dioxide dissolved in water. The soda is bottle under pressure. It's the pressure above the solution that keeps the carbon dioxide as a dissolved gas. As soon as you open it you get the reverse of the solution process-

While I did initially decrease the volume above the solution by squeezing it, there's nothing preventing some of the carbon dioxide from coming out of solution. Normally, the pressure above the solution prevents much of anything happening but not in this case. The carbon dioxide that escapes from solution actually serves to expand the bottle back to its original shape! Granted, once the bottle finally repressurizes at the original expanded shape the gas release will stop, but by then it's likely to be too late. The soda's already going flat.

You still might be questioning the above statement since it may be a belief that you've held for quite a long time. You could be thinking that the effect in the video is simply due to a bad seal and it's air entering the bottle and not CO₂ from the solution that's expanding the bottle. That's cool. This is a very easy one to do yourself and you don't even need to make a video of it. Simply take a fresh soda in a plastic container (needs to be a pliable container to get the effect), pour out about 20 % of the soda, squeeze the bottle until the liquid is back to its original level, and tightly cap the bottle. Now vigorously shake the bottle and you'll see it pop back into shape!

So how the heck can this be explained using Le Châtelier's Principle? Well, first thing to do is to look at the previous example differently.

Although I have removed some of the liquid, this doesn't change the concentration of dissolved CO₂.

What's really happened is that I've suddenly increased the volume above the solution which has decreased the pressure. Le Châtelier's Principle states that the system will reestablish an equilibrium by minimizing the stress applied to it.

Once you've determined the pressure change, you then need to examine the chemical equation for any gaseous reactants and/or products. We can ignore liquids, solids, and even aqueous solutions since they are relatively incompressible. So, if the pressure increases, then the reaction shifts to the side that produces fewer gas molecules. If the pressure decreases, then the reaction shifts to the side that produces more gas molecules. Let's look at the carbonic acid equilbrium equation once again.

The soda stays fizzy as long as the CO₂ stays in solution or stays on the reactants side of the equation as written. If you open the soda, you suddenly decrease the CO₂ pressure above the solution and the reaction shifts to the side that produces more moles of gas to reestablish equilibrium. In this case, it shifts to the right side or the products side. This is why the squeezed soda bottle expands as you saw in the video.

We can also look at it in terms of Q_P. Here's the reaction quotient (remember, we leave out pure solids, liquids, and aqueous solutions and only include gaseous species as explained earlier)-

The removal of the carbon dixoide gas immediately makes Q_P less than K_P and so Q_P needs to increase to regain equilibrium. The only way to do that is to increase the pressure of CO₂ above the solution which eventually forces the squeezed container back to its original shape.

If you want to keep an opened bottle of carbonated soda fizzy as long as possible, then keep the bottle cold. You should still drink it as soon as possible.

I have altered this page because it was brought to my attention that something that was in the previous paragraph was being used to justify fizzy keepers. Nope. And while I have some thoughts on what I'd love to test if I had the equipment, I think I'll keep that to myself. Repeat to yourself- Henry's law... Henry's law... Henry's law...