For a first-order reaction,the time required for $93.75\%$ completion is $4$ times the half-life of the reaction. What is the relationship between the time required for $93.75\%$ completion and the half-life $(t_{0.5})$?

$A$ first-order reaction is $50\%$ complete in $45 \text{ minutes}$. How many hours will it take for the reaction to be $99.9\%$ complete?

If the rate constant for a first order reaction is $k$,then find the time required for completion of $80 \%$ of the reaction.

$A_{(g)} \xrightarrow{\Delta} P_{(g)} + Q_{(g)} + R_{(g)}$ follows first-order kinetics with a half-life of $69.3 \ s$ at $500^{\circ}C$. Starting from the gas '$A$' enclosed in a container at $500^{\circ}C$ and at a pressure of $0.4 \ atm$,the total pressure of the system after $230 \ s$ will be (in $atm$)

Consider the following reaction,the rate expression of which is given below:
$A + B \rightarrow C$
$\text{rate} = k[A]^{1/2}[B]^{1/2}$
The reaction is initiated by taking $1 \ M$ concentration of $A$ and $B$ each. If the rate constant $(k)$ is $4.6 \times 10^{-2} \ s^{-1}$,then the time taken for $A$ to become $0.1 \ M$ is . . . . . . . . . . $sec$. (nearest integer)

$A_{(g)} \rightarrow 2B_{(g)}$. Initially,$2 \ mol$ of $A$ are taken in a $5 \ L$ vessel. After $20 \ min$,$[A]_t = \frac{[B]_t}{2}$. Find the half-life time of $A$ in the first-order reaction in $\min$.