Consider the following reaction:

CO(g)+2H2(g)⇌CH3OH(g)

This reaction is carried out at a specific temperature with initial concentrations of [CO] = 0.27 M and [H2] = 0.49 M. At equilibrium, the concentration of CH3OH is 0.11 M. Find the equilibrium constant at this temperature.
Express the equilibrium constant to two significant figures.

I keep getting 1.69 but its not right. Please help!

To find the equilibrium constant (K) at a specific temperature for the given reaction, we can use the formula:

K = ([CH3OH]/([CO][H2]))

We are given the concentration of CH3OH as 0.11 M and the initial concentrations of CO and H2 as 0.27 M and 0.49 M, respectively.

Plugging these values into the formula, we get:

K = (0.11 / (0.27 * 0.49))

K ≈ 0.11 / 0.1323

K ≈ 0.831

Therefore, the equilibrium constant at this temperature, expressed to two significant figures, is approximately 0.83.

To find the equilibrium constant at a given temperature, you need to use the equilibrium concentrations of the reactants and products. In this case, the equilibrium concentrations are [CO] = 0.27 M, [H2] = 0.49 M, and [CH3OH] = 0.11 M.

The equilibrium constant, K, for the given reaction can be expressed as:

K = ([CH3OH] / [CO]) * ([H2] ^ 2)

Plugging in the given values, we have:

K = (0.11 / 0.27) * (0.49 ^ 2)
K = 0.407 * 0.2401
K = 0.09796

Rounding to two significant figures, the equilibrium constant at this temperature is approximately 0.10.

So, it seems that the correct answer is not 1.69, but rather 0.10.

Why didn't you show us how you came up with 1.69. If we don't know how you got there how can we tell where you're going wrong? I cam up with approx 1

9.43

You have to do the ICE rule. Where x=[CH3OH]