For the reaction:

[Cu(NH₃)₄]²⁺ + H₂O_{$\rightleftharpoons$}[Cu(NH₃)₃H₂O]²⁺ + NH₃

the net rate of reaction at any time is given by, net rate =

2.0x10^-4 [Cu(NH₃)₄]²⁺[H₂O] - 3.0x10⁵ [Cu(NH₃ )₃ H₂0]²⁺[NH₃]

Then correct statement is/are :

1. rate constant for forward reaction = 2 x 10^-4

2. rate constant for backward reaction = 3 x 10⁵

3. equilibrium constant for the reaction = 6.6 x 10^-10

4. all of the above

Consider the chemical reaction,

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

The rate of this reaction can be expressed in terms of time derivative of concentration of  ${\mathrm{N}}_2 \left(\mathrm{g}\right), {\mathrm{H}}_2\left(\mathrm{g}\right)$ and ${\mathrm{NH}}_3\left(\mathrm{g}\right)$.

The correct relationship amongest the rate expressions is: 

(1) Rate $=\frac{-\mathrm{d}\left[{\mathrm{N}}_2\right]}{dt}=-\frac{1}{3}\frac{ \mathrm{d}\left[{\mathrm{H}}_2\right]}{dt}=\frac{1}{2}\frac{ \mathrm{d}\left[{\mathrm{NH}}_3\right]}{dt}$

(2) Rate $=\frac{-\mathrm{d}\left[{\mathrm{N}}_2\right]}{dt}=-\frac{3 \mathrm{d}\left[{\mathrm{H}}_2\right]}{dt}=\frac{2 \mathrm{d}\left[{\mathrm{NH}}_3\right]}{dt}$

(3) Rate $=\frac{\mathrm{d}\left[{\mathrm{N}}_2\right]}{dt}=\frac{1}{3}\frac{ \mathrm{d}\left[{\mathrm{H}}_2\right]}{dt}=\frac{1}{2}\frac{ \mathrm{d}\left[{\mathrm{NH}}_3\right]}{dt}$

(4) Rate $=\frac{-\mathrm{d}\left[{\mathrm{N}}_2\right]}{dt}=-\frac{ \mathrm{d}\left[{\mathrm{H}}_2\right]}{dt}=\frac{ \mathrm{d}\left[{\mathrm{NH}}_3\right]}{dt}$

A reactant with initial concentration 1.386 \(\mathrm{mol} \text { litre }{ }^{-1}\) showing first order change takes 40 minute to become half. If it shows zero order change taking 20 minute to becomes half under similar conditions, the ratio, K₁/K₀ for first order and zero order kinetics will be:

1. 0.5 mol^-1 litre

2. 1.0 mol/litre

3. 1.5 mol/litre

4. 2.0 mol^-1 litre

In a first order reaction, the concentration of the reactant is decreased from 1.0 M to 0.25M in 20 minute. The rate constant of the reaction would be:

1. 10min^-1

2. 6.931 min^-1

3. 0.6931 min^-1

4. 0.06931 min^-1

The rate constant of a first-order reaction is\(4 \times 10^{-3} \mathrm{sec}^{-1}.\) At a reactant concentration of \(0.02~\mathrm{M},\) the rate of reaction would be:

If concentration of reactants is increased by 'X', the rate constant K becomes:

1. e^K/X

2. K/X

3. K

4. X/K

A graph plotted between log (t) 50% vs. log (a) concentration is a straight line. What conclusion can you draw from the given graph?

1. n=1, t_1/2 = 1/K.a

2. n=2, t_1/2 = 1/a

3. n=1, t_1/2 = 0.693/K

4. None of the above

The rate of a chemical reaction doubles for every 10°C rise of temperature. If the temperature is raised by 50°C, the rate of the reaction increases by about :

1. 10 times

2. 24 times

3. 32 times

4. 64 times

Consider the reaction:

Cl₂(aq) + H₂S(aq) → S(s) +2H⁺(aq) +2Cl^-(aq)

The rate equation for this reaction is rate = k[Cl₂][H₂S] Which of these mechanisms is/are consistent with this rate equation?

A. Cl₂ + H₂S → H⁺ + Cl^- +Cl⁺ + HS^- (slow)

cl⁺ + HS^- → H⁺ +Cl^- + S (fast)

B. H₂S $\iff$ H⁺ + HS^- (fast equilibrium)

Cl₂ + HS^- → 2Cl^- + H⁺ + S (slow)

1. A only

2. B only

3. Both A and B

4. Neither A nor B

For the reaction 2NO₂ + F₂ → 2NO₂F, following

mechanism has been provided,

 NO₂ + F₂   $\stackrel{slow}{\to }$  NO₂F+F

NO₂ + F   $\stackrel{fast}{\to }$ NO₂F

Thus, rate expression of the above

reaction can be written as:

1. r = K[NO₂]²[F₂]

2. r = K[NO₂ ][F₂]

3. r = K[NO₂]

4. r = K[F₂]