Equilibrium Calculation

Question

At a certain temperature, this reaction establishes an equilibrium with the given equilibrium constant, Kc.

3A ( g ) + 2B ( g ) ⇀ ↽ 4C ( g ) Kc = 1.73 × 10^23

If, at this temperature, 2.20 mol of A and 3.80 mol of B are placed in a 1.00 L container, what are the concentrations of A, B, and C at equilibrium?

[A] = 0.00100 M --- incorrect

[B] = 2.33 M ---- correct

[C] = 2.93 M ---- correct

I have calculated B and C correctly, but Sapling says that the concentration I have for A is not correct. Sapling says:

Note that the final concentration of A is approximately 0 M. It is initially assumed that the reaction went to completion; however, though this assumption was made, C will still dissociate, or backreact, to produce A and B in order to reach equilibrium. Therefore, set up an equilibrium table to calculate the equilibrium values for each species where the final concentrations in the previous table are now the initial values used in the equilibrium table.

Your help is greatly appreciated

Answer 1

If you know [B] and [C] are correct, than you know the value of "x" in your ICE table is 0.7325M. Using that information, [A] at equilibrium would be 2.20 - 3x, or 2.20 - (3)(0.7325) = 0.0025. Still very close to zero, but a different answer than what you got.

That being said, when you plug these values back into Kc you don't get the 1.73 x 10^23, are you sure this is the correct value of Kc?