Which of the following statements is $NOT$ true for a reaction having rate law $r=k[H_2][I_2]$?

$A$ reaction is second order with respect to a reactant. How is the rate of reaction affected if the concentration of the reactant is $(i)$ doubled $(ii)$ reduced to half?

The possible mechanism for the reaction $2NO + Br_2 \to 2NOBr$ is:
$NO + Br_2 \rightleftharpoons NOBr_2$ (Fast)
$NOBr_2 + NO \to 2NOBr$ (Slow)
The rate law expression is:

The reaction between $X$ and $Y$ is first order with respect to $X$ and zero order with respect to $Y$.

$Experiment$	$[X] / (mol \ L^{-1})$	$[Y] / (mol \ L^{-1})$	$\text{Initial rate} / (mol \ L^{-1} \ min^{-1})$
$I$	$0.1$	$0.1$	$2 \times 10^{-3}$
$II$	$0.2$	$0.2$	$4 \times 10^{-3}$
$III$	$0.4$	$0.4$	$M \times 10^{-3}$
$IV$	$0.1$	$0.2$	$2 \times 10^{-3}$

Examine the data of the table and calculate the ratio of the numerical value of $M$ to $0.2$.

For a reaction $A + B \to$ product,it was found that the rate of reaction increases four times if the concentration of $A$ is doubled,but the rate of reaction remains unaffected if the concentration of $B$ is doubled. Hence,the rate law for the reaction is

Consider the reaction:
$Cl_{2(aq)} + H_2S_{(aq)} \rightarrow S_{(s)} + 2H^{+}_{(aq)} + 2Cl^{-}_{(aq)}$
The rate equation for this reaction is:
$\text{rate} = k[Cl_2][H_2S]$
Which of these mechanisms is/are consistent with this rate equation?
$A.$ $Cl_2 + H_2S \rightarrow H^{+} + Cl^{-} + Cl^{+} + HS^{-}$ (slow)
$Cl^{+} + HS^{-} \rightarrow H^{+} + Cl^{-} + S$ (fast)
$B.$ $H_2S \rightleftharpoons H^{+} + HS^{-}$ (fast equilibrium)
$Cl_2 + HS^{-} \rightarrow 2Cl^{-} + H^{+} + S$ (slow)