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### Course: AP®︎/College Chemistry>Unit 7

Lesson 6: Using the reaction quotient

# Using the reaction quotient

By comparing the reaction quotient to the equilibrium constant, we can predict the direction a reaction will proceed to reach equilibrium. If Q < K, the reaction will proceed towards the products. If Q > K, the reaction will proceed towards the reactants. If Q = K, the reaction is already at equilibrium and will not change. Created by Jay.

## Want to join the conversation?

• I don't see how you could get a different Qc value than your Kc value if you calculate them the exact same way. I guess sometimes the actual experimental Kc value is different than the Kc calculated? I just want to know why it would be different.....A big bucket of thanks to anybody who helps me with this! ;)
• So the reaction quotient, Qc, the equilibrium constant, Kc, are calculated using the same equation for any particular reaction, but the key is that we input different concentrations at different time periods for the two values.

The equilibrium constant is a single value (at a particular temperature) in which the forward and reverse reaction rates are equal. Realistically this is observed at constant concentrations for all chemical species involved in the reaction. So it's a single set of concentrations which yields a single unchanging value.

The reaction quotient meanwhile is taken at any instant of time and can have infinitely many possible values. The purpose of the reaction quotient is just to compare to the equilibrium constant to judge which reaction rate will be favored; the forward or the reverse.

Hope that helps.
• Why does each dot (blue/red) represent 0.1M? Is this arbitrary for explanation purposes?
(1 vote)
• Yeah you're right, it's arbitrary. Jay could have picked any value for each of the dots and the process would be identical.
• why did he change from 0.1 M to 1 M at the second and the third reaction
(1 vote)
• He simplified the math since you get the same ratios by counting the colored dots in the particulate diagrams as you ould if you actually used molarities. For exmaple, 0.5/0.3 = 5/3
(1 vote)
• Is it ever possible for Qc > Kc in the example with the red and blue dots from the video?
(1 vote)
• Well yeah, that simply means that we have more products than is allowed at equilibrium. Any reaction is capable of having too much of either reactants or products for equilibrium.
(1 vote)
• How do you know what direction a reaction will travel to reach equilibrium.?
Also, I heard that if Q < K, the reaction will proceed towards the products and that If Q > K, the reaction will do the same. Is this correct or did they miss something?

Cheers.

- THE WATCHER