The reaction at hydrogen and iodine monochloride is represented by the equation is
$H_{2} (g) + 2 I C l (g) \to 2 H C l (g) + I_{2} (g)$
This reaction is first order in $H_{2} (g)$ and also first - order in ICI(g). which of these proposed mechanism can be consistent with the given information about this reaction ?
Mechanism I : $H_{2} (g) + 2 I C l (g) \to 2 H C l (g) + I_{2} (g)$
Mechanism II : $H_{2} (g) + I C l (g) s l o w \to H C l (g) + H I (g)$
$H I (g) + I C l (g) f a s t \to H C l (g) + I_{2} (g)$ :

I only

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II only

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Both I and II

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Neither I nor II

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Solution

The correct option is B II only
According to question the given reaction is first order in $H_{2} (g)$ and also first order in $I C l (g)$ so, the slowest (rate determining) step in the reaction should involve $H_{2} (g)$ & $I C l (g)$ one mole each.
The mechanism $I I$ is seeming consistent with the conditions stated as:-
$H_{2} (g) + I C l (g) ⟶ H C l (g) + H I (g)$ (slow)
$R a t e = K [H_{2}] [I C l]$
& $H I (g) + I C l (g) ⟶ H C l (g) + I_{2} (g)$ (fast)

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Similar questions

The reaction of hydrogen and iodine monochloride is given as:
$H_{2} (g) + 2 I C l (g) \to 2 H C l (g) + I_{2} (g)$

This reaction is of first order with respect to $H_{2} (g)$ and $I C l (g)$ , following mechanisms were proposed :
Mechanism A : $H_{2} (g) + 2 I C l (g) \to 2 H C l (g) + I_{2} (g)$

Mechanism B : $H_{2} (g) + I C l (g) \to H C l (g) + H I (g), s l o w$

$H I (g) + I C l (g) \to H C l (g) + I_{2} (g), f a s t$
Which of the above mechanism(s) can be consistent with the given information about the reaction?

Q. The following reaction has a second-order rate law:

H_{2} (g) + 2 I C l (g) \to I_{2} (g) + 2 H C l (g)

Rate

= k [H_{2}] [I C l]

Devise a possible reaction mechanism.

At 700 K, equilibrium constant for the reaction

is 54.8. If 0.5 molL^–1 of HI₍_g₎ is present at equilibrium at 700 K, what are the concentration of H₂₍_g₎ and I₂₍_g₎ assuming that we initially started with HI₍_g₎ and allowed it to reach equilibrium at 700 K?