If 0.1 mole of I2, is introduced into 1 litre flask at 1000 k, at equilibrium (Kc = 10-6), which one is correct?  Option 1) [ I2 (g) ] > [I (g)] Option 2) [I2 (g)] < [1 (g)] Option 3) [I2 (g)] = [ 1(g) ] Option 4) [I1 (g)] = $\frac{1}{2}[I (g)]$

As we learnt in

Law of Chemical equilibrium -

$A+B\rightleftharpoons C+D$

where  A & B are the reactants, C & D are the product in balanced chemical equations.

- wherein

$K_{c}=\frac{[C]\:[D]}{[A]\:[B]}$

Kc is the equilibrium constant.

The reaction takes place as follows

$I_{2}\left ( g \right )\leftrightharpoons 2I (g), K_{c}=10^{-6}$

Clearly, the equilibrium constant favours $I_{2}$

$\therefore \left [ I_{2} \right ]> \left [ I \right ]$

Option 1)

[ I2 (g) ] > [I (g)]

correct

Option 2)

[I2 (g)] < [1 (g)]

incorrect

Option 3)

[I2 (g)] = [ 1(g) ]

incorrect

Option 4)

[I(g)] = $\frac{1}{2}[I (g)]$

incorrect

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