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Current time:0:00Total duration:7:57

Video transcript

in this video we're going to go through an example reaction that uses the shot Lee's principle so what we're going to do is we're going to apply the shot Lee's principle to look at various changes to this reaction when we perturb our reaction from equilibrium just as a reminder what do I mean when I say this reaction is at equilibrium so what that means is we have a reversible reaction we have the forward reaction which has some rate K forward the reverse reaction has some rate K backward and a dynamic equilibrium these rates are equal to each other so all of the concentrations are going to stay constant and then what we do is we decide to see what happens when we add some carbon dioxide gas if we add carbon dioxide gas the concentration of carbon dioxide will go up or you can think about it as a partial pressure going up and lashauwn leas principle tells us that if we had a reaction at equilibrium and then we perturbed it by adding more co2 it will shift to try to try to reduce the effect of that change so it will favor the reverse reaction so if we add co2 what happens is we favor our reactants another way we can see this is by looking at the equilibrium constant for this reaction so we can write our equilibrium constant K where this is a capital K it's kind of confusing but I will try to make this look like a capital K and we can write it in two ways we can write it in terms of the concentration the molar concentration and if we write KC the expression will be the product concentration so our co2 gas concentration and that's it because when you write out KC we write out concentrations of gases and we write out concentrations of solutions but we don't include solids so KC is just the concentration of co2 at equilibrium and I'm going to write an EQ there just to show that's the equilibrium concentration and I said you can also write it in terms of partial pressures so there's our fancy capital K with a P subscript which means that instead of concentrations we're writing everything for gases in terms of partial pressures so we have the partial pressure of co2 and again that's it because everything else is a solid so we don't include those in our equilibrium expression so writing these expressions out will be really helpful for our second condition so we're going to think about what happens when you increase when you increase the volume of our container we can rewrite the partial pressure actually in terms of the volume so if you use the ideal gas law the partial pressure of co2 is equal to the moles of co2 times RT divided by the volume and similarly we can rewrite molar concentration in terms of moles divided by volume if we increase the volume of our container increasing the volume since it's in the denominator we'll make our pressure go down so our partial pressure of co2 will go down and we won't be at equilibrium anymore and same for the concentration since our perturbation is decreasing the carbon dioxide concentration lilius principle says that our reaction will try to counteract that change it will try to get back to equilibrium and try to get the co2 concentration back up and so it'll have to favor products our reaction will favor the products to try to get the co2 moles of co2 back up so that we get back to the equilibrium concentration of co2 and the equilibrium partial pressure the third change we're going to look at is what happens when you add argon gas so argon gas is an inert gas we don't expect it to react with anything one thing that'll happen when you add the argon gas is it will increase the overall pressure of your container so it will increase the total pressure oops but that actually doesn't tell us what it does to our equilibrium let's look back at our equilibrium expressions KC and KP we can see that the partial pressure for KP only depends on the moles of our co2 and our volume since we didn't change the moles of co2 and we also didn't change the volume even though we increase P total the partial pressure of co2 stayed the same so that means we didn't perturb our reaction from equilibrium and since we didn't perturb it from equilibrium there'll be no shift we are still at equilibrium the concentrations will still stay the same what happens when we add more calcium carbonate so that's our starting material and it is a solid our go Librium expressions are determined by our co2 concentration so adding more calcium carbonate which is a solid isn't actually going to perturb our reaction from equilibrium our reaction is still going to be at equilibrium and we will get no shift in concentrations we're actually going to look at one more thing we're going to think about what happens when you add a catalyst let's say we want to speed up this reaction we can envision what's going on here when we add the catalyst by using an energy diagram so if you have an energy diagram we have energy on the y axis and we're looking at the difference in energy between our reactants or starting materials with our products and I just sort of made up these relative energies the way that I have this drawn here we can see that our product is lower in energy than our starting material and our forward rate KF which is up here KF is determined by the size of this activation barrier between our starting material in our transition state our backward rate K KB up here our backward rate is determined by the size of this energy barrier so the difference in energy between the product and our transition state if we add a catalyst to a reaction we can think about it as lowering the activation energy for for our reaction and that means we have a lower energy barrier for our forward reaction so our forward reaction is going to get faster but it's also going to lower the rate of our backward reaction so K backward or KB is also going to speed up and since it speeds up both the forward and the backward reactions adding a catalyst also won't perturb your reaction from equilibrium so adding a catalyst will result in no shift in concentrations so the main things to remember from this problem I think the things that I find most tricky anyway are that adding an inner gas it'll increase your total pressure but it won't actually change any of your partial pressures so it won't shift your reaction from equilibrium and the same thing is true for solids and catalysts so all of those three things in are gases solids and catalysts will not shift your reaction from equilibrium