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(A) For a zero order reaction,the rate of reaction is independent of the concentration of the reactant.$1$. The rate law for this reaction is: $	ext{

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Rate of the reaction depends on concentration of $NH_3$

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(A) For a zero order reaction,the rate of reaction is independent of the concentration of the reactant.$1$. The rate law for this reaction is: $	ext{Rate} = k[NH_3]^0 = k$.$2$. Since the rate is equal to the rate constant $(k)$,option $B$ is true.$3$. At high pressure,the catalyst surface is fully saturated with $NH_3$ molecules. Therefore,increasing the pressure further does not increase the rate of reaction,making option $C$ true.$4$. Since the rate is constant,the rate of decomposition of ammonia remains constant,making option $D$ true.$5$. Option $A$ states that the rate depends on the concentration of $NH_3$,which is false for a zero order reaction. Thus,option $A$ is the correct answer.

Time $(sec)$	Rate $(mole \ litre^{-1} \ sec^{-1})$
$0$	$2.8 \times 10^{-2}$
$10$	$2.78 \times 10^{-2}$
$20$	$2.81 \times 10^{-2}$
$30$	$2.79 \times 10^{-2}$

At high pressure,the following reaction is zero order:
$2NH_{3(g)} \xrightarrow[Pt \ \text{catalyst}]{1130 \ K} N_{2(g)} + 3H_{2(g)}$
Choose the option which is not true with respect to the reaction,from the choices given below:

Similar Questions

The rates of a certain reaction $(dc/dt)$ at different times are as follows:
Time $(sec)$ Rate $(mole \ litre^{-1} \ sec^{-1})$
$0$ $2.8 \times 10^{-2}$
$10$ $2.78 \times 10^{-2}$
$20$ $2.81 \times 10^{-2}$
$30$ $2.79 \times 10^{-2}$

The reaction is:

In a zero-order reaction,if the initial concentration of the reactant is doubled,the time required for half the reactant to be consumed:

The rate constant for a zero order reaction is $2 \times 10^{-2} \ mol \ L^{-1} s^{-1}$. If the concentration of the reactant after $25 \ s$ is $0.5 \ M$,the initial concentration must have been $:-$ (in $M$)

For the reaction $A \rightarrow \text{products}$,the graph of $t_{1/2}$ versus $[A]_0$ is given below. The concentration of $A$ at $10 \ \text{minutes}$ is $.......... \times 10^{-3} \ \text{mol L}^{-1}$ $(nearest \ integer)$. The reaction was started with $2.5 \ \text{mol L}^{-1}$ of $A$.

The decomposition of $NH_{3}$ on a platinum surface is a zero-order reaction. What are the rates of production of $N_{2}$ and $H_{2}$ if $k = 2.5 \times 10^{-4} \, mol \, L^{-1} \, s^{-1}$?

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At high pressure,the following reaction is zero order:$2NH_{3(g)} \xrightarrow[Pt \ \text{catalyst}]{1130 \ K} N_{2(g)} + 3H_{2(g)}$Choose the option which is not true with respect to the reaction,from the choices given below: