Medium
From the rate expression for the following reaction,determine its order of reaction and the dimensions of the rate constant.
$H_2 O_2+3 I^{-}+2 H^{+} \rightarrow 2 H_2 O+I_3^{-} \text { Rate }=k\left[H_2 O_2\right]\left[I^{-}\right]$
$H_2 O_2+3 I^{-}+2 H^{+} \rightarrow 2 H_2 O+I_3^{-} \text { Rate }=k\left[H_2 O_2\right]\left[I^{-}\right]$
(N/A) Given Rate $= k[H_2O_2][I^{-}]$
The order of reaction is the sum of the powers of the concentration terms in the rate law expression.
Order $= 1 + 1 = 2$
Dimension of $k = \frac{\text{Rate}}{[H_2O_2][I^{-}]}$
$= \frac{\text{mol L}^{-1} \text{s}^{-1}}{(\text{mol L}^{-1})(\text{mol L}^{-1})}$
$= \text{L mol}^{-1} \text{s}^{-1}$
The order of reaction is the sum of the powers of the concentration terms in the rate law expression.
Order $= 1 + 1 = 2$
Dimension of $k = \frac{\text{Rate}}{[H_2O_2][I^{-}]}$
$= \frac{\text{mol L}^{-1} \text{s}^{-1}}{(\text{mol L}^{-1})(\text{mol L}^{-1})}$
$= \text{L mol}^{-1} \text{s}^{-1}$
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If the concentration of reactant $A$ is increased by $10$ times,the rate of reaction becomes $100$ times. What is the order of reaction if the rate law is,$\text{rate} = k[A]^{x}$?
For the reaction $2 A + B \longrightarrow D + E$,the following mechanism has been proposed: $A + B \longrightarrow C + D$ (slow) and $A + C \longrightarrow E$ (fast). Determine the rate law.
For the first order decomposition reaction of $N_2O_5$,it is found that -
$(a)$ $2N_2O_5 \rightarrow 4NO_{2(g)} + O_{2(g)}$ ; $-\frac{d[N_2O_5]}{dt} = k[N_2O_5]$
$(b)$ $N_2O_5 \rightarrow 2NO_{2(g)} + 1/2 O_{2(g)}$ ; $-\frac{d[N_2O_5]}{dt} = k'[N_2O_5]$
Which of the following is true?
$(a)$ $2N_2O_5 \rightarrow 4NO_{2(g)} + O_{2(g)}$ ; $-\frac{d[N_2O_5]}{dt} = k[N_2O_5]$
$(b)$ $N_2O_5 \rightarrow 2NO_{2(g)} + 1/2 O_{2(g)}$ ; $-\frac{d[N_2O_5]}{dt} = k'[N_2O_5]$
Which of the following is true?
Medium
View SolutionWhich of the following statements is $NOT$ true for a reaction having rate law $r=k[H_2][I_2]$?
The data for the reaction $A + B \to C$ is given below:
$Exp$ $[A]_0$ $[B]_0$ Initial Rate $1$ $0.012$ $0.035$ $0.10$ $2$ $0.024$ $0.035$ $0.80$ $3$ $0.012$ $0.070$ $0.10$ $4$ $0.024$ $0.070$ $0.80$
Determine the rate law for the reaction.
| $Exp$ | $[A]_0$ | $[B]_0$ | Initial Rate |
|---|---|---|---|
| $1$ | $0.012$ | $0.035$ | $0.10$ |
| $2$ | $0.024$ | $0.035$ | $0.80$ |
| $3$ | $0.012$ | $0.070$ | $0.10$ |
| $4$ | $0.024$ | $0.070$ | $0.80$ |
Determine the rate law for the reaction.
Difficult
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