I was recently given the question:
The equilibrium constant, K, for the following reaction:
N₂(g) + O₂(g) ↔ 2NO(g)
is 4.00 x 10⁻² at a very high temperature. The reaction is at equilibrium at this temperature with [N2] = [O2] = 0.100 M and [NO] = 0.0200 M in a 2.00 liter flask. If 0.120 mol of NO is suddenly added to the reaction mixture what will be the concentrations of all species when equilibrium is re-established?
I think filling out the ICE table would look like this:
N₂ | O₂ | 2NO | |
---|---|---|---|
I | .100 | .100 | .080 |
C | +x | +x | -2x |
E | .100+x | .100+x | .080-2x |
Which would make the the equilibrium formula: $$ 4.00\cdot 10^{-2}=\frac{(.080-2x)^{2}}{(.100+x)(.100+x)} $$
However, when solving for $x$ I get $x=\frac{8.2\pm\sqrt{7.84}}{198}$ or $.0\overline{5}$ and $.0\overline{27}$.
nvm I get it now but as this is my first time using stackexchange and this syntax took me too long I feel like posting it because the TeX looks nice to look at :)
noo why does the table only work in preview mode :(