At a certain temperature the equilibrium constant, Kc, equals 0.11 for the reaction:
2 ICl(g) ? I2(g) + Cl2(g).
What is the equilibrium concentration of ICl if 0.45 mol of I2 and 0.45 mol of Cl2 are initially mixed in a 2.0-L flask?

a.) 0.17 M

b.) 0.27 M

c.) 0.34 M

d.) 0.14 M

The equilibrium concentration of ICl is 0.27 M

Explanation:

We are given:

Initial moles of iodine gas = 0.45 moles

Initial moles of chlorine gas = 0.45 moles

Volume of the flask = 2.0 L

The molarity is calculated by using the equation:

Initial concentration of iodine gas =

Initial concentration of chlorine gas =

For the given chemical equation:

As, the initial moles of iodine and chlorine are given. So, the reaction will proceed backwards.

The chemical equation becomes:

Initial:         0.225      0.225

At eqllm:   0.225-x    0.225-x     2x

The expression of for above equation follows:

Putting values in above equation, we get:

Neglecting the value of x = 0.668 because equilibrium concentration cannot be greater than the initial concentration

So, equilibrium concentration of ICl = 2x = (2 × 0.135) = 0.27 M

Hence, the equilibrium concentration of ICl is 0.27 M

consider the reaction below.

 no rx

at 500 k, the reaction is at equilibrium with the following concentrations.

[pci5]= 0.0095 m

[pci3] = 0.020

[ci2] = 0.020 m

what is the equilibrium constant for the given reaction?

0.042

0.42

2.4

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