Rate of a Chemical Reaction

Consider a reaction $a\text{G}+\text{bH}\to \text{products}$.  When concentration of both the reactants $\mathrm{G}$ and $\mathrm{H}$ is doubled, the rate increases by eight times. However, when the concentration of $\mathrm{G}$ is doubled, keeping the concentration of $\mathrm{H}$ constant, the rate is doubled. The overall order of the reaction is:

Consider the chemical reaction:

${\text{N}}_2\left(\text{g}\right)+3{\text{H}}_2\left(\text{g}\right)⟶2{\text{NH}}_3\left(\text{g}\right)$

The rate of this reaction can be expressed in terms of time derivatives of concentration of  $N_2\left(\text{g}\right)\text{ , }{\text{H}}_2\left(\text{g}\right)\text{ and }{\text{NH}}_3\left(\text{g}\right)$ Identify the correct relationship amongst the rate expressions:

Hydrogen gas, ${\mathrm{H}}_2\left(\mathrm{g}\right)$, reacts with iodine gas, ${\mathrm{I}}_2\left(\mathrm{g}\right)$, to form hydrogen iodide, $\mathrm{HI}\left(\mathrm{g}\right)$:

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

The mechanism of the two-step reaction is considered to be:

step $1$: ${\mathrm{I}}_2\left(\mathrm{g}\right) \underset{{\mathrm{k}}_{-1}}{\overset{{\mathrm{k}}_1}{\leftrightharpoons }} 2\mathrm{I}\left(\mathrm{g}\right)$                        fast

step $2$: $2\mathrm{I}\left(\mathrm{g}\right) + {\mathrm{H}}_2\left(\mathrm{g}\right) \stackrel{{\mathrm{k}}_2}{\to } 2\mathrm{HI}\left(\mathrm{g}\right)$     slow

What is the rate equation for the overall reaction?

For a reaction $1/2 \mathrm{A}\to 2 \mathrm{B}$, rate of disappearance of $\mathrm{A}$ is related to rate of appearance of $\mathrm{B}$ by the expression

Equilibrium constants $\mathrm{NO}\left(\mathrm{g}\right)+1/2{\mathrm{O}}_2\stackrel{{\mathrm{K}}_1}{\to }{\mathrm{NO}}_2( \mathrm{g})$ and $2{\mathrm{NO}}_2\left(\mathrm{g}\right)\stackrel{{\mathrm{K}}_2}{\to }2\mathrm{NO}\left(\mathrm{g}\right)+{\mathrm{O}}_2( \mathrm{g})$ for the following equilibria

Ammonia and oxygen react at high temperature as;
$4{{\mathrm{NH}}_3}_{\left(\mathrm{g}\right)}+5{{\mathrm{O}}_2}_{\left(\mathrm{g}\right)}\to 4{\mathrm{NO}}_{\left(\mathrm{g}\right)}+6{\mathrm{H}}_2{\mathrm{O}}_{\left(\mathrm{g}\right)}$

In an experiment rate of formation $\mathrm{NO}$ is $3.6\times {10}^{-3} \mathrm{mol}{ \mathrm{L}}^{-1}{ \mathrm{s}}^{-1}$. Calculate rate of disappearance of ammonia.

_____ (Law of mass action/Rate law) is an equation which represents the ratio between the masses of products and masses of reactants whose power is raised according to their stoichiometric coefficients.

For a reaction
${\mathrm{NO}}_2+\mathrm{CO}\to {\mathrm{CO}}_2+\mathrm{NO}$ the possible elementary steps (below $\left.500 \mathrm{K}\right)$ are
Step $1:{\mathrm{NO}}_2+\mathrm{N}{\mathrm{O}}_2\stackrel{ \mathrm{slow} }{⟶}\mathrm{NO}+{\mathrm{NO}}_3$
Step $2:{\mathrm{NO}}_3+\mathrm{CO}\stackrel{ \mathrm{Fast} }{⟶}{\mathrm{CO}}_2+{\mathrm{NO}}_2$
its rate expression will be

For a given reaction $3\mathrm{A}\mathrm{ }+\mathrm{B}\to \mathrm{C}+\mathrm{D}$ the rate of reaction can be represented by

The term $\frac{-\mathrm{dc}}{\mathrm{dt}}$ in a rate equation refers to:

Thermal decomposition of ${\mathrm{N}}_2{\mathrm{O}}_5$ occurs as per the equation below :

$2{\mathrm{N}}_2{\mathrm{O}}_5\to 4{\mathrm{NO}}_2+{\mathrm{O}}_2$

The correct statement is

The rate constnat for the reaction

$2 {\mathrm{N}}_2{\mathrm{O}}_5⟶4{\mathrm{NO}}_2+{\mathrm{O}}_2$

is $3.0\times {10}^{-5}{\sec }^{-1}$. if the rate is $2.4\times {10}^{-5} \mathrm{mol} \mathrm{litre}$${}^{-1}\sec ^{-1}$, then the concentration of ${\mathrm{N}}_2{\mathrm{O}}_5$ (in $\mathrm{mol} \mathrm{litre}{ }^{-1}$ ) is

Rate constant depends upon:

The rate constant of the reaction $\mathrm{A}\to \mathrm{B} \mathrm{is} 0.5\times {10}^{-3}\mathrm{M} {\sec }^{-1}$. What is the concentration of B after 10 min from the start of the reaction?
 

For the reaction, $2{\mathrm{SO}}_2+{\mathrm{O}}_2\rightleftharpoons 2{\mathrm{SO}}_3$, the rate of disappearance of ${\mathrm{O}}_2$ is $2\times {10}^{-4} \mathrm{mol} {\mathrm{L}}^{-1} {\mathrm{s}}^{-1}$ . The rate of appearance of ${\mathrm{SO}}_3$ is

For a reaction: ${\mathrm{N}}_2+3{\mathrm{H}}_2\to 2{\mathrm{NH}}_3$, the rate of reaction with respect to ${\mathrm{NH}}_3$will be:

For the reaction, ${\mathrm{N}}_2+3{\mathrm{H}}_2\to 2{\mathrm{NH}}_3$, if $\frac{\mathrm{d}\left[{\mathrm{NH}}_3\right]}{\mathrm{dt}}=2\times {10}^{-4} \mathrm{mol}/\mathrm{L}/\mathrm{s}$, the value of $-\frac{\mathrm{d}\left[{\mathrm{H}}_2\right]}{\mathrm{dt}}$ would be

Nitrogen tetraoxide  $\left({\mathrm{N}}_2{\mathrm{O}}_4\right)$ decomposes as
                   ${\mathrm{N}}_2{\mathrm{O}}_4\left(\mathrm{g}\right)\to 2 {\mathrm{NO}}_2\left(\mathrm{g}\right)$

If the pressure of ${\mathrm{N}}_2{\mathrm{O}}_4$ falls from $0.50 \mathrm{atm}$ to $0.32 \mathrm{atm}$ in $30$ minutes, the rate of appearance of ${\mathrm{NO}}_2\left(\mathrm{g}\right)$ is:

What is a rate constant?