TBR Chem 1 - Equilibrium

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mchung7

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Hey everyone. I have a quick question.

TBR gives an example question about equilibrium (K & Q):

When the reaction quotient is greater than the equilibrium constant, which of the following is NOT true?

A. The system has too many products and too few reactants.
B. The reaction is displaced from equilibrium.
C. The reaction must shift in the forward direction to reach equilibrium.
D. The reverse reaction rate is greater than the forward reaction rate.


The answer is choice C, which I understand. However, I am a bit confused about choice D. If the reverse reaction rate is GREATER than the forward reaction rate, shouldn't choice D be a valid answer as well? If K < Q as this question indicates, there is too much product and too little reactant. If the reverse reaction rate is greater than the forward, wouldn't that mean that there should be less product and more reactant?

I feel that my understanding of the word "greater" may be incorrect, which is the only possibility that I see. I am assuming that the word greater means "faster". When they refer to a reaction rate being greater, do they actually mean that the numerical value of kr is larger, and thus slower?

I just wanted to make sure that I misunderstood their use of "greater", and that this isn't a mistake.


Thanks!

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Like you said, when K<Q, there's too much product. At that moment, what does the reaction "want" to do? It wants to return to Q=K. The only way it can accomplish this is for the reverse reaction to dominate. Otherwise, you will keep building up even more product, and Q will get larger and larger, which would obviously not restore the reaction mixture to equilibrium.
 
My question with answer D is .. Wouldnt one think though that the forward rate is dominating because of K being < Q?

Or is answer D is saying which rate is dominating in an attempt to get equilibrium?
 
The forward rate dominated up to that point. We want to look at the equation at that point and onward, so the reverse reaction will dominate to return the reaction to equilibrium.

D is saying the reverse reaction will dominate in order to return the reaction to equilibrium.
 
Here is thought experiment to help you understand why D is NOT the answer.

Imagine a reaction that is at equilibrium. Next, let's add extra product and check the reaction status before it had a chance to re-equilibrate itself. BTW, we have reached the situation in the problem above.

Did we affect the forward rate? Did we affect the reverse rate? The answer is 'No' to both questions. We simply added more product. Thus, D is not the answer because we cannot conclude anything about the relative rates of the forward and reverse reactions.
 
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