Chemical kinetics practice · Medium

What fraction of a reactant showing first order remains after 40 minute if ${t_{1/2}}$ is 20 minute?

The reaction obey I order with respect to ${\mathrm{H}_2}$ and $\mathrm{ICl}$ both

${\mathrm{H}_2}\left( \mathrm{g} \right) + 2\mathrm{ICl}\left( \mathrm{g} \right) \to 2\mathrm{HCl}\left( \mathrm{g} \right) + {\mathrm{I}_2}\left( \mathrm{g} \right)$

Which of the following mechanism is in consistent with the given fact?

Mechanism$A:{\mathrm{H}_2}\left( \mathrm{g} \right) + 2\mathrm{ICl} \to 2\mathrm{HCl}\left( \mathrm{g} \right) + {\mathrm{I}_2}\left( \mathrm{g} \right)$

Mechanism$B:\left( \mathrm{i} \right){\mathrm{H}_2}\left( \mathrm{g} \right) + \mathrm{ICl}\left( \mathrm{g} \right)\xrightarrow{{Slow}}\mathrm{HCl}\left( \mathrm{g} \right) + \mathrm{HI}\left( \mathrm{g} \right)$

$\left( {\mathrm{ii}} \right)\mathrm{HI}\left( \mathrm{g} \right) + \mathrm{ICl}\left( \mathrm{g} \right) \to \mathrm{HCl}\left( \mathrm{g} \right) + {\mathrm{I}_2}$

The half life of a first order reaction is $69.3 \,{sec}$. The value of the rate constant of the reaction is:

If $'a'$ is the initial concentration of the reactant, the half-life period of the reaction of ${n^{th}}$ order is proportional to

The activation energy of exothermic reaction $A \to B$ is $80\;kJ\;\mathrm{mo}{\mathrm{l}^{ - 1}}$. The heat of reaction is $200\;kJ\;\;\mathrm{mo}{\mathrm{l}^{ - 1}}$. The activation energy for the reaction $B \to A$ (in $kJ\;\;\mathrm{mo}{\mathrm{l}^{ - 1}})$ will be

$50 \%$ of a first order reaction completes in 20 minute. $100 \%$ of same reaction shall complete in:

Which of the following is not a first order reaction?

The rate constant for the reaction $\mathrm{CO}_2+\mathrm{OH}^{-} \longrightarrow \mathrm{HCO}_3^{-}$ in aqueous solution is $4.0 \times 10^3 \mathrm{~L} \mathrm{~mol}^{-1} \mathrm{~s}^{-1}$ when $\left[\mathrm{CO}_2\right]=10^{-6} \mathrm{~mol} \mathrm{~L}^{-1}$ and $\left[\mathrm{OH}^{-}\right]=0.1 \mathrm{~mol} \mathrm{~L}^{-1}$ Number of moles of $\mathrm{HCO}_3^{-}$formed per second is $x \times 10^{-4} \mathrm{~mol} \mathrm{~L}^{-1}$. What is the value of $x$ ?

For the reaction,$C{l_2} + 2{l^ - } \to {I_2} + 2C{l^ - }$ the initial concentration of ${I^ - }$ was $0.20\;mol\;{L^{ - 1}}$ and the concentration after 20 min was $0.18\;mol\;{L^{ - 1}}$.Then the rate of formation of ${I_2}$ in $mol\;{L^{ - 1}}$ would be

The following mechanism has been proposed for the reaction of $\mathrm{NO}$ with $\mathrm{B}{\mathrm{r}_2}$ to form $NOBr$
 $\mathrm{NO}\left( \mathrm{g} \right) + \mathrm{B}{\mathrm{r}_2}\left( \mathrm{g} \right)\leftrightharpoons\mathrm{NOB}{\mathrm{r}_2}\left( \mathrm{g} \right)$
$\mathrm{NOB}{\mathrm{r}_2}\left( \mathrm{g} \right) + \mathrm{NO}\left( \mathrm{g} \right) \to 2\mathrm{NOBr}\left( \mathrm{g} \right)$
If the second step is the rate determining step, the order of the reaction with respect to $\mathrm{NO}\left( \mathrm{g} \right)$ is

In Arrhenius plot intercept is equal to

The rate constant of a reaction increases by 5% when its temperature is raised from $27^\circ C\;\mathrm{to}\;28^\circ C$. The activation energy of the reaction is

For reaction $A \to B,$ the rate constant ${k_1} = {A_1}{e^{ - E{a_1}/RT}}$ and for the reaction $X \to Y,$ the rate constant${k_2} = {A_2}{e^{ - E{a_2}/RT}}.$ If ${A_1} = {10^8},{A_2} = {10^{10}}$ and $E{a_1} = 600\;\mathrm{cal}/\mathrm{mol},\;E{a_2} = 1800\;\mathrm{cal}/\mathrm{mol},$ then the temperature at which ${k_1} = {k_2}\;is\;\left( {R = 2\;\mathrm{cal}/\mathrm{k }\;\mathrm{mol}} \right)$

Assertion: On increasing temperature rate of reaction always increases.

Reason: For any reaction equilibrium always shifts in forward direction on increasing temperature.

Assertion (A):Activation energy always increases the potential energy of reaction system.

Reason (R): The minimum kinetic energy that a molecule must possess in order to react on collision is called its activation energy.

Dinitropentoxide decomposes as follows:

$\mathrm{N}_2 \mathrm{O}_5(\mathrm{g}) \longrightarrow 2 \mathrm{NO}_2(\mathrm{g})+\frac{1}{2} \mathrm{O}_2(\mathrm{g})$

Given:

$\begin{aligned} & \mathrm{d}\left[\mathrm{N}_2 \mathrm{O}_5\right] / \mathrm{dt}=\mathrm{k}_1\left[\mathrm{N}_2 \mathrm{O}_5\right] \\ & \mathrm{d}\left[\mathrm{NO}_2\right] / \mathrm{dt}=\mathrm{k}_2\left[\mathrm{N}_2 \mathrm{O}_5\right] \\ & \mathrm{d}\left[\mathrm{O}_2\right] / \mathrm{dt}=\mathrm{k}_3\left[\mathrm{N}_2 \mathrm{O}_5\right] \end{aligned}$

What is the relation between $\mathrm{k}_1, \mathrm{k}_2$ and $\mathrm{k}_3$?

Consider the reaction,

$C{l_2}\left( {aq} \right) + {H_2}S\left( {aq} \right) \to S\left( s \right) + 2{H^ + }\left( {aq} \right) + 2C{l^ - }\left( {aq} \right)$ the rate equation for this reaction is , $rate = k\left[ {C{l_2}} \right]\left[ {{H_2}S} \right]$

Which of these mechanisms is/are consistent with this rate equation?

$\begin{array}{l} A.\;\;C{l_2} + {H_2}S\left( {aq} \right) \to {H^ + }C{l^ - } + C{l^ + } + H{S^ - }\left( {slow} \right)\\ \;\;\;\;\;C{l^ + } + H{S^ - } \to {H^ + }C{l^ - } + S\left( {fast} \right)\\ B.\;\;{H_2}S \leftrightarrow {H^ + } + H{S^ - }\left( {fast\;equilibrium} \right)\\ \;\;\;\;C{l_2} + H{S^ - } \to 2C{l^ - } + {H^ + }S\left( {slow} \right) \end{array}$

A catalyst is used in a reaction to

The rate constant of a zero order reaction is $2.0 \times {10^{ - 2}}\mathrm{mol}\,\,{\mathrm{L}^{ - \mathrm{1}}}{\mathrm{s}^{ - \mathrm{1}}}\mathrm{.}$ If the concentration of the reactant after 25 seconds is 0.5 M. What is the initial concentration?

Statement I : Catalyst can alter a reaction both ways means it can either decrease or increase rate of reaction
Statement II : A catalyst always decrease the activation energy of the reaction but does not alter Gibb's energy.

The half life period of a first order reaction is 6.93 min. The time required for the completion of 99 % of the chemical reaction will be

For the first-order reaction $t_{99 \%}=x \times t_{90 \%}.$ The value of '$x$' will be

The rate of reaction doubles by increase of $10^{\circ} \mathrm{C}$ temperature. If there is an increase of $50^{\circ} \mathrm{C}$ in temperature, how many times rate of reaction would increase?

A reaction is 50% complete in 2 hours and 75% complete in 4 hours. The order of reaction is

For the chemical reaction, $2 {O}_3 \rightleftharpoons 3 {O}_2$, the reaction proceed as follows

$\begin{aligned} & {O}_3 \rightleftharpoons {O}_2+{O} \text { (fast) } \\ & {O}+{O}_3 \rightarrow 2 {O}_2 \text { (slow) } \end{aligned}$

The rate law expression will be

For the reaction $\mathrm{X \longrightarrow Y+Z}$ the following were obtained.

$\begin{array}{cccccccccccccccccccc} {{\text{Time}}}&0&{900}&{1800} \\ {{\text{Conc}}{\text{. of X}}}&{50.7}&{19.6}&{7.61} \end{array}$

Find the order of the reaction.

For a reaction the initial rate is given as: $R_0=k[A]_0^2[B]_0$ by what factor, the initial rate of reaction will increase if initial concentration of $A$ is taken 1.5 times and of $B$ is tripled?

Assertion (A):Threshold energy of a reaction is dependent of temperature.

Reason (R):The energy of activation decreases with increase in temperature.

Consider the following statements.

a. For zero order reaction $t_{1 / 2}$ is directly proportional to the initial concentration of the reactants.

b. Inversion of cane sugar is an example of pseudo first order reaction.

c. Decomposition of gaseous ammonia on a hot platinum surface is a zero order reaction at high pressure.

The correct statement are

A first order reaction is half-completed in $45 \mathrm{~min}$. How long does it need for $99.9 \%$ of the reaction to be completed?