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For the following reaction, \( K = 0.538 \) at 313 K:

\[ \text{N}_2\text{O}_4(g) \rightleftharpoons 2 \text{NO}_2(g) \]

If a reaction vessel initially contains an \(\text{N}_2\text{O}_4\) concentration of 0.0500 mol L\(^{-1}\) at 313 K, what are the equilibrium concentrations of \(\text{N}_2\text{O}_4\) and \(\text{NO}_2\) at 313 K?

Answer :

Final answer:

The equilibrium concentrations for the given chemical reaction at 313 K are calculated using the equilibrium constant and initial concentration. After solving the resulting quadratic equation, the equilibrium concentrations of N2O4 and NO2 at 313 K are found to be 0.016 mol L−1 and 0.068 mol L−1, respectively.

Explanation:

The given chemical equation can be written as N2O4(g)⇌2 NO2(g), which suggests that for every mole of N2O4 that decomposes, 2 moles of NO2 are produced.

Let's represent the initial concentration of N2O4 as [N2O4]i = 0.0500 mol L–1 and let's denote the change in concentration as x.

At equilibrium, the concentrations of N2O4 and NO2 will be [N2O4] = 0.0500 - x and [NO2] = 2x, respectively.

We can then use the equilibrium constant expression K = [NO2]^2/[N2O4], substituting the values we have.

Given K = 0.538 at 313 K, we set up the equation 0.538 = (2x)^2 / (0.0500 - x).

Solving this quadratic equation for x, we found x = 0.0340 mol L−1.

Substituting x into the expressions for [NO2] and [N2O4], we get [NO2] = 2x = 0.0680 mol L−1 and [N2O4] = 0.0500 - x = 0.016 mol L−1 at equilibrium.

Therefore, the equilibrium concentrations of N2O4 and NO2 at 313K are 0.016 mol L−1 and 0.068 mol L−1, respectively.

Learn more about Chemical Equilibrium here:

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Rewritten by : Brahmana