The time required for $90\%$ completion of a certain first order reaction is $1 \text{ hour}$. Calculate the time required for $99.9\%$ completion of the same reaction.

At $30^{\circ}C$,the half-life for the decomposition of $AB_{2}$ is $200\,s$ and is independent of the initial concentration of $AB_{2}$. The time required for $80\%$ of the $AB_{2}$ to decompose is $....s$ (Given: $\log 2 = 0.30; \log 5 = 0.70$)

The half-life of a first-order reaction $X \longrightarrow Y + Z$ is $3 \ minutes$. What is the time required to reduce the concentration of $X$ by $90 \%$ of its initial concentration?

The ratio of the time taken to complete $60\%$ and $20\%$ of a first-order reaction $[A \rightarrow \text{product}]$ is : $(\log 2 = 0.3)$

For a first-order reaction,if the half-life period is $10 \text{ minutes}$,how many minutes will it take for the concentration of the reactant to decrease from $0.08 \ M$ to $0.02 \ M$?

For the first$-order$ reaction,$T_{av}$ (average life),$T_{50}$ and $T_{75}$ in the increasing order are: