The reaction $2 A + B + C \longrightarrow D + E$ is found to be first order in $A$,second order in $B$,and zero order in $C$. What is the effect of increasing the concentration of all reactants twice?

For the reaction $A + B \rightarrow C$,select the appropriate rate law based on the following data:
$1$. $[A] = 0.012, [B] = 0.035 \rightarrow \text{Initial Rate} = 0.10$
$2$. $[A] = 0.024, [B] = 0.070 \rightarrow \text{Initial Rate} = 1.6$
$3$. $[A] = 0.024, [B] = 0.035 \rightarrow \text{Initial Rate} = 0.20$
$4$. $[A] = 0.012, [B] = 0.070 \rightarrow \text{Initial Rate} = 0.80$

For a second-order reaction,the rate constant is $8 \times 10^{-5} \, M^{-1} \, \text{min}^{-1}$. In how much time will a $1 \, M$ solution decrease to $0.5 \, M$?

For the reaction $aA \to xP$,the rate is $2.4 \ mMs^{-1}$ when $[A] = 2.2 \ M$. When the concentration of $A$ is halved,the rate becomes $0.6 \ mMs^{-1}$. Determine the order of the reaction with respect to $A$.

Calculate the overall order of a reaction which has the rate expression:
$(a)$ $\text{Rate} = k[A]^{1/2}[B]^{3/2}$
$(b)$ $\text{Rate} = k[A]^{3/2}[B]^{-1}$

The rate for the reaction $A + B \rightarrow \text{product}$ is $1.8 \times 10^{-2} \ mol \ dm^{-3} \ s^{-1}$. Calculate the rate constant if the reaction is second order in $A$ and first order in $B$,given $[A] = 0.2 \ M$ and $[B] = 0.1 \ M$.