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Offline Araconan

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Question Regarding another Equilibrium Problem
« on: July 17, 2012, 10:54:36 PM »
I am currently attempting the following problem:

Nitric oxide and bromine react accordingly to the following reaction:

2NO(g) + Br2(g) <----------> 2NOBr(g)

The equilibrium constant at 300K is 134 atm-1

What would be the partial pressures of all species if NO and Br2, both at an initial partial pressure of 0.3 atm, were allowed to come to equilibrium at this temperature?

My attempt:

Because the equilibrium constant is relatively large, I'm going to assume that the forward reaction goes towards completion

Creating an ICE table:
                           2NO(g) + Br2(g) <----------> 2NOBr(g)
Initial:                 0.3           0.3                       0
Change:             -0.3          -0.15                    +0.3
Equilibrium:         0              0.15                     0.3

Considering the reverse reaction:
If I reverse the reaction, the equilibrium constant of the reverse reaction would be equal to 134-1

Creating an ICE table:
                           2NOBr(g) <----------> 2NO(g) + Br2(g)
Initial:                 0.3                             0             0.15
Change:             -2x                              +2x         +x
Equilibrium          0.3 - 2x                       2x          0.15 + x

The equation would be:

(2x)2(0.15+x)/(0.3-2x)2 = 134-1

Because the equilibrium constant is very small, I'm going to assume that x is negligible, and thus the equation becomes:


Solving for x gives me 0.033, while the textbook, using successive approximations (it didn't consider the reverse reaction, and just solved the equation from the original reaction) got 0.026.

I'm wondering, did I do something wrong? Or is my answer more accurate than the textbooks?  ;D

Thank you in advance!

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