A flask is charged with 1.500 atm of N2O4(g) and 1.00 atm NO2(g) at 25°C, and the following equilibrium is established. 2 NO2(g) ⇌ N2O4(g) At equilibrium, the partial pressure of NO2(g) is 0.512 atm. Calculate the equilibrium constant Kp for this reaction.

The equilibrium constant, for the given reaction is 6.653

Explanation:

We are given:

Initial partial pressure of nitrogen dioxide = 1.00 atm

Initial partial pressure of dinitrogen tetraoxide = 1.500 atm

Equilibrium partial pressure of nitrogen dioxide = 0.512 atm

For the given chemical equation:

Initial:          1.00           1.500

At eqllm:   1.00-2x        1.500+x

Evaluating the value of 'x'

So, equilibrium partial pressure of dinitrogen tetraoxide = (1.500 + x) = [1.500 + 0.244] = 1.744 M

The expression of for above equation follows:

Putting values in above equation, we get:

Hence, the equilibrium constant, for the given reaction is 6.653

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