For a reaction $A$ $\xrightarrow{K_1} B$ $\xrightarrow{K_2} C$. If the rate of formation of $B$ is set to be zero,then the concentration of $B$ is given by:

The rate constant value for a reaction is $1.75 \times 10^2 \ L^2 \ mol^{-2} \ sec^{-1}$. The half-life period $t_{1/2} \propto$ . . . . . . .

For the reaction $2A + B \rightarrow C + D$,select the correct rate law based on the following data:
$1$. $[A] = 0.1, [B] = 0.1, \text{Initial Rate} = 7.5 \times 10^{-3}$
$2$. $[A] = 0.3, [B] = 0.2, \text{Initial Rate} = 9.0 \times 10^{-2}$
$3$. $[A] = 0.3, [B] = 0.4, \text{Initial Rate} = 3.6 \times 10^{-1}$
$4$. $[A] = 0.4, [B] = 0.1, \text{Initial Rate} = 3.0 \times 10^{-2}$

Half-life periods for a reaction at initial concentrations of $0.1 \ M$ and $0.01 \ M$ are $5$ and $50$ minutes,respectively. The order of reaction is

The rate law for the reaction between substances $A$ and $B$ is given by $\text{Rate} = k[A]^n[B]^m$. If the concentration of $A$ is doubled and the concentration of $B$ is halved,what is the ratio of the new rate to the initial rate?

In the reaction sequence $A$ $\xrightarrow{K_1} B$ $\xrightarrow{K_2} C$ $\xrightarrow{K_3} D$,where $K_3 > K_2 > K_1$,which step determines the rate of the reaction?