Observe the following reaction: $2 A + B \longrightarrow C$. The rate of formation of $C$ is $2.2 \times 10^{-3} \ mol \ L^{-1} \ min^{-1}$. What is the value of $-\frac{d[A]}{d t}$ (in $mol \ L^{-1} \ min^{-1}$)?

For a reaction,$2 \ N_2O_{5(g)} \longrightarrow 4 \ NO_{2(g)} + O_{2(g)}$. The $N_2O_5$ disappears at a rate of $0.06 \ mol \ dm^{-3} \ s^{-1}$. What is the rate of $NO_{2(g)}$ formation (in $mol \ dm^{-3} \ s^{-1}$)?

In a reaction,$n_1 A + n_2 B \rightarrow m_1 C + m_2 D$,$5 \ M$ solution of reactant $A$ is allowed to react with $3 \ M$ solution of reactant $B$. After $5 \ s$,the concentration of $A$ was found to be $4 \ M$. The rate of decomposition of $A$ and the rate of formation of $D$ are respectively $:-$

For the reaction $2N_2O_{5(g)} \rightleftharpoons 4NO_{2(g)} + O_{2(g)}$ in a closed vessel,the concentration of $NO_2$ increases by $2.0 \times 10^{-2} \, mol \, L^{-1}$ in $5 \, s$. Calculate the rate of reaction.

If the rate of disappearance of $N_2O_5$ in the following reaction is $1.2 \times 10^{-5} \ mol \ L^{-1} \ s^{-1}$,the rate of production of $NO_2$ in $mol \ L^{-1} \ s^{-1}$ is:
$2N_2O_{5(g)} \longrightarrow 4NO_{2(g)} + O_{2(g)}$

For the reaction,$N_2 + 3H_2 \rightarrow 2NH_3,$ if $\frac{d[NH_3]}{dt} = 2 \times 10^{-4} \ mol \ L^{-1} \ s^{-1},$ the value of $-\frac{d[H_2]}{dt}$ would be .................. $mol \ L^{-1} \ s^{-1}$.