A certain reaction ${\mathrm{B}}^{\mathrm{n}+}$ is getting converted to ${\mathrm{B}}^{\left(\mathrm{n}+4\right)+}$ in solution. The rate constant of this reaction is measured by titrating a volume of the solution with a redox agent which reacts only with ${\mathrm{B}}^{\mathrm{n}+}$ and ${\mathrm{B}}^{\left(\mathrm{n}+4\right)+}$. In the process it converts ${\mathrm{B}}^{\mathrm{n}+}$ to ${\mathrm{B}}^{\left(\mathrm{n}+2\right)+}$ and ${\mathrm{B}}^{\left(\mathrm{n}+4\right)+}$ to ${\mathrm{B}}^{\left(\mathrm{n}-1\right)+}$. At $\mathrm{t}=0$, the volume of reagent consumed is $25\mathrm{mL}$ and at $\mathrm{t}=10$ minute, the volume used is $35.5\mathrm{mL}$. Calculate the rate constant in ${\mathrm{hr}}^{-1}$ for the conversion of ${\mathrm{B}}^{\mathrm{n}+}$ of ${\mathrm{B}}^{\left(\mathrm{n}+4\right)+}$ assuming it to be a first order reaction. [Assume that the redox reagent has the same $n-$factor with both the species]
[Given: ${\log }_{10}2=0·30,{\log }_{10}3=0·48,{\log }_{e}x=2·3{\log }_{10}x$]

For the reaction: $2\mathrm{A}+{\mathrm{B}}_2+\mathrm{C}\to {\mathrm{A}}_2\mathrm{B}+\mathrm{BC}$, the rate law expression has been determined experimentally to be $\mathrm{R}=\mathrm{k}\left[\mathrm{A}{]}^2\right[\mathrm{C}]$ with $\mathrm{k}=3.0\times {10}^{-4} {\mathrm{M}}^{-2} {\min }^{-1}$.

Determine the rate after $0.04$ moles per litre of $\mathrm{A}$ has been reacted. If initial concentration was $\left[\mathrm{A}\right] =0.1\mathrm{M}, \left[\mathrm{B}\right] = 0.35\mathrm{M}, \left[\mathrm{C}\right] =0.25$.

If answer is $\mathrm{y} \times {10}^{-7}\mathrm{M}/\min . \mathrm{Then} \mathrm{y} \mathrm{is}.$

The decomposition of ${\mathrm{N}}_2{\mathrm{O}}_5$ in ${\mathrm{CCl}}_4$ solution at $318 \mathrm{K}$ has been studied by monitoring the concentrration of ${\mathrm{N}}_2{\mathrm{O}}_5$ in the solution.. Intially the concentration of ${\mathrm{N}}_2{\mathrm{O}}_5$ is $2.32\mathrm{M}$ and after $184$ minutes, it is reduced to $2.08\mathrm{M}$. The reaction take place according to the given equation.
$2{ \mathrm{N}}_2{\mathrm{O}}_5⟶4{\mathrm{NO}}_2+{\mathrm{O}}_2$

What is the average rate of production of ${\mathrm{NO}}_2$ in milli moles ${\mathrm{\ell t}}^{-1}{\sec }^{-1}$ ?

From the following data for the reaction between $\mathrm{A}$ and $\mathrm{B}$ :

If the pre-exponential factor is $\mathrm{a}\times {10}^{17}$, find the value of $\text{'}\mathrm{a}\text{'}$ up to the nearest integer.

From the following data for the reaction between $\mathrm{A}$ and $\mathrm{B}$ :

Calculate the energy of activation in $\mathrm{kJ}.$

Answer correct up to one place of decimal.

For the reaction $\mathrm{A}\to \mathrm{B}+\mathrm{C}$ the following data were obtained: 

$\begin{array}{lccr}t\left(\mathrm{s}\right)& 0& 900& 1800\\ \left[\mathrm{A}\right]& 50.8& 19.7& 7.62\end{array}$

What is the order of the reaction?

Two reactions $\left(\mathrm{i}\right)\mathrm{A}\to \mathrm{P}$ and $\left(\mathrm{ii}\right)\mathrm{B}\to \mathrm{P}$ Follows first-order kinetics. The rate of reaction $\left(\mathrm{i}\right)$ is doubled when the temperature is raised from $300 \mathrm{K}$ to $310 \mathrm{K}$. The half-life for this reaction at $310 \mathrm{K}$ is $30$ Minutes. At the same temperature, $\mathrm{B}$ decomposes twice as fast as $\mathrm{A}$. If the energy of activation for the reaction $\left(\mathrm{ii}\right)$ is half that of reaction $\left(\mathrm{i}\right)$, calculate the rate constant of the reaction $\left(\mathrm{ii}\right)$ at $300 \mathrm{K}.$