Chemistry - Chemical Kinetics Question with Solution | TestHub

Given:

Reaction: $A\to B+C$

It follows first order kinetics.

Initial concentration, $[A{]}_0=1M$

Final concentration, $[A]=0.25M$

Time = $138.6min$

For a first-order reaction:

$$k=\frac{2.303}{t}\log \frac{[A{]}_0}{[A]}$$

$$k=\frac{2.303}{138.6}\log \frac{1}{0.25}$$

$$\log 4=0.602$$

$$k=\frac{2.303\times 0.602}{138.6}$$

$$k=0.01mi{n}^{-1}$$

Now, the rate when $[A]=0.1M$:

For a first-order reaction:

$$Rate=k[A]$$

$$=0.01\times 0.1$$

$$Rate=1\times 10^{-3}Mmi{n}^{-1}$$