Question

Consider the following reaction and its equilibrium constant:

I2(g) + Br2(g) ⇌ 2 IBr(g) Kc = 1.1 × 10^2

A reaction mixture contains 0.49 M I2, 0.28 M Br2 and 3.5 M IBr. Which of the following statements is TRUE concerning this system?

a. The reaction quotient will decrease.
b. The reaction will shift in the direction of products.
c. The equilibrium constant will increase.
d. The reaction will shift in the direction of reactants.
e. The system is at equilibrium

Answer 1

Statement b) is true concerning this system. This conclusion can be reached by comparing the reaction quotient and the equilibrium constant.

To determine which is the correct answer (except c, which is obviously incorrect, the equilibrium constant isn't dependent on the concentrations of the reaction components), we need to calculate the reaction quotient (Q) and compare it to the equilibrium constant (K).

The expression for the equilibrium constant (and the reaction quotient) looks like this:

$Q = \frac{[IBr]^{2} }{[I_{2} ][Br_{2} ]}$

Q = (3.5 M)²/(0.28 M * 0.49 M)

Q = 89 = 0.89 x 10²

This means that the reaction quotient is lower than the equilibrium constant (so the system is not at equilibrium, making option e) incorrect), which means that it will shift in the direction that will cause it to be equal to reach equilibrium. So, the reaction quotient will increase (making option a) incorrect) and the system will shift towards the product because increasing the concentration of the product and reducing the concentrations of the reactants.

You can learn more about reaction quotient here:
brainly.com/question/24202150

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