For a zero order reaction,will the molecularity be equal to zero? Explain.

For the reaction $Cl_{2(aq)} + H_2S_{(aq)} \to S_{(s)} + 2H^+_{(aq)} + 2Cl^-_{(aq)}$, the rate law is given by $\text{Rate} = K[Cl_2][H_2S]$. Which of the following mechanisms is consistent with this rate law?
$(A)$ $Cl_2 + H_2S \to H^+ + Cl^- + Cl^+ + HS^-$ (slow); $Cl^+ + HS^- \to H^+ + Cl^- + S$ (fast)
$(B)$ $H_2S \rightleftharpoons H^+ + HS^-$ (fast equilibrium); $Cl_2 + HS^- \to 2Cl^- + H^+ + S$ (slow)

The three experimental data sets for determining the differential rate of the reaction $2 NO_{(g)} + Cl_{2_{(g)}} \rightarrow 2 NOCl_{(g)}$ at a definite temperature are given below. (Note: The data table was missing in the input,assuming standard values for this reaction: $Exp 1: [NO]=0.1, [Cl_2]=0.1, Rate=0.18$; $Exp 2: [NO]=0.1, [Cl_2]=0.2, Rate=0.36$; $Exp 3: [NO]=0.2, [Cl_2]=0.1, Rate=0.72$).
$(a)$ Derive the differential rate law of the reaction.
$(b)$ Calculate the order of the reaction.
$(c)$ Calculate the value of the rate constant.

The rate law for the reaction $A + B \rightarrow C$ at $25^{\circ} C$ is given by $\text{rate} = k[A][B]^2$. Calculate the rate of reaction if the rate constant at the same temperature is $6.25 \ mol^{-2} \ dm^6 \ s^{-1}$ where $[A] = 1 \ M$ and $[B] = 0.2 \ M$.

If $a$ is the initial concentration of the reactant,the half-life period of the reaction of $n^{th}$ order is inversely proportional to

$A$ chemical reaction proceeds through the following steps:
Step-$I$: $2A \rightleftharpoons X$ (fast)
Step-$II$: $X + B \rightarrow Y$ (slow)
Step-$III$: $Y + B \rightarrow \text{Product}$ (fast)
The rate law for the overall reaction is: