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Hydrogenation of vegetable ghee at 25⁰C reduces pressure of H₂ from 2 atom to 1.2 atom in 50 minute. The rate of reaction in terms of molarity per second is:
1. 1.09 x 10^-6
2. 1.09 x 10^-5
3. 1.09 x 10^-7
4. 1.09 x 10^-8
 

Following mechanism has been proposed for a reaction,
2A+B $\to$D+E
A+B$\to$ C+D    ...(Slow)
A+ C$\to$ E         ...(Fast)
The rate law expression for the reaction is:
1. r = K[A]²[B]
2. r=K[A][B]
3. r= K[A]²
4. r= K[A][C]

For the reaction N₂ + 3H₂ $\to$ 2NH₃, the rate $\frac{\mathrm{d}\left[{\mathrm{NH}}_3\right]}{\mathrm{dt}}$= 2 x 10^-4 M s^-1 .Therefore, the rate $\frac{-\mathrm{d}\left[{\mathrm{N}}_2\right]}{\mathrm{dt}}$ is given as:
1. 10^-4 Ms^-1
2. 10⁴ Ms^-1
3. 10^-2 sM^-1
4. 10^-4 sM^-1

If 'a' is the initial concentration of a substance that reacts according to zero-order kinetics and k is the rate constant, then the time for the reaction to go to completion is:
1. a/k
2. 2/ka
3. k/a
4. 2k/a

In a reaction, the rate expression is, rate = K[A][B]^2/3[C]⁰, the order of the reaction is:
1. 1
2. 2
3. 5/3
4. zero

The rate of a reaction get doubles when the temperature changes from 7°C to 17°C. By what factor will it change for the temperature change from 17°C to 27°C?
1. 1.81
2. 1.71
3. 1.91
4. 1.76

For the elementary step, (CH₃)_3.CBr(aq) → (CH₃)₃C⁺ (aq) + Br^- (aq), the molecularity is:
1. Zero
2. 1
3. 2
4. Cannot be ascertained

When ethyl acetate was hydrolysed in pressure of 0.1 N HCl, the rate constant was found to be 5.40 x 10^-5 sec^-1 . But when 0.1 N H₂SO₄ was used for hydrolysis, the rate constant was found to be 6.25 X10^-5sec^-1. Thus, it may be concluded that:
1. H₂SO₄ is stronger than HCI
2. H₂SO₄ is weaker than HCl
3. H₂SO₄ and HCl both have the same strength
4. The data are not sufficient to compare the strength of H₂SO₄ and HCI

The half time of a second order reaction is:
1. Inversely proportional to the square of the initial concentration of the reactants.
2. Inversely proportional to the initial concentration of the reactants.
3. Proportional to the initial concentration of reactants.
4. Independent of the initial concentration of reactants.

The half-life period of a first order chemical reaction is 6.93 minutes. The time required for the completion of 99% of the chemical reaction will be (log 2 = 0.301):
1. 23.03 minutes
2. 46.06 minutes
3. 460.6 minutes
4. 230.3 minutes

A zero order reaction is one:
1. in which reactants do not react
2. in which one of the reactants is in large excess
3. whose rate does not change with time
4. whose rate increases with time

For A + B $\to \hspace{0.17em}$C + D, $∆$H = -20 kJ mol^-1 , the activation energy of the forward reaction is 85 kJ mol^-1. The activation energy for the backward reaction is…. kJ mol^-1.

For the elementary reaction M $\to$ N, the rate of disappearance of M increases by a factor of 8 upon doubling the concentration of M. The order of the reaction with respect to M will be:
1. 4
2. 3
3. 2
4. 1

The rate constant for a second order reaction is 8x10^-5 M^-1 min^-1 . How long will it take a 1M solution to be reduced to 0.5M?
1. 8.665 x 10³ minute
2. 8 x 10^-5 minute
3. 1.25 x 10⁴ minute
4. 4x10^-5 minute

The time for half-life of a first order reaction is 1 hr. What is the time taken for 87.5% completion of the reaction?
1. 1 hour
2. 2 hour
3. 3 hour
4. 4 hour

Which order of reaction obeys the relation t_1/2 = 1/Ka?
1. First
2. Second
3. Third
4. Zero

The chemical reaction, 2O₃ $\to$3O₂ proceeds as follows;
O₃ $\rightleftharpoons$O₂ + O .....(Fast)
O+O₃ $\to$ 2O₂ ....(Slow)
The rate law expression should be:
1. r = K[O₃]²
2. r = K[O₃]²[O₂]^-1
3. r = K[O₃][O₂]
4. unpredictable

The rate of reaction becomes 2 times for every 10°C rise in temperature. How the rate of reaction will increase when temperature is increased from 30°C to 80°C?
1. 16
2. 32
3. 64
4. 128

What fraction of a reactant showing first order remains after 40 minute if t_1/2 is 20 minute?
1. 1/4
2. 1/2
3. 1/8
4. 1/6

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]

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  ${\mathrm{N}}_2<mspace\; width="1em"></mspace>\left(\mathrm{g}\right),<mspace\; width="1em"></mspace>{\mathrm{H}}_2\left(\mathrm{g}\right)$ and ${\mathrm{NH}}_3\left(\mathrm{g}\right)$.
The correct relationship amongest the rate expressions is: 
(1) Rate $=\frac{-\mathrm{d}\left[{\mathrm{N}}_2\right]}{dt}=-\frac{1}{3}\frac{<mspace\; width="1em"></mspace>\mathrm{d}\left[{\mathrm{H}}_2\right]}{dt}=\frac{1}{2}\frac{<mspace\; width="1em"></mspace>\mathrm{d}\left[{\mathrm{NH}}_3\right]}{dt}$
(2) Rate $=\frac{-\mathrm{d}\left[{\mathrm{N}}_2\right]}{dt}=-\frac{3<mspace\; width="1em"></mspace>\mathrm{d}\left[{\mathrm{H}}_2\right]}{dt}=\frac{2<mspace\; width="1em"></mspace>\mathrm{d}\left[{\mathrm{NH}}_3\right]}{dt}$
(3) Rate $=\frac{\mathrm{d}\left[{\mathrm{N}}_2\right]}{dt}=\frac{1}{3}\frac{<mspace\; width="1em"></mspace>\mathrm{d}\left[{\mathrm{H}}_2\right]}{dt}=\frac{1}{2}\frac{<mspace\; width="1em"></mspace>\mathrm{d}\left[{\mathrm{NH}}_3\right]}{dt}$
(4) Rate $=\frac{-\mathrm{d}\left[{\mathrm{N}}_2\right]}{dt}=-\frac{<mspace\; width="1em"></mspace>\mathrm{d}\left[{\mathrm{H}}_2\right]}{dt}=\frac{<mspace\; width="1em"></mspace>\mathrm{d}\left[{\mathrm{NH}}_3\right]}{dt}$

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

The activation energies of the forward and backward reactions in the case of a chemical reaction are 30.5 and 45.4 KJ/mol respectively. The reaction is
1. Exothermic
2. Endothermic
3. Neither exothermic nor endothermic
4. Independent of temperature

The first order reaction reaction \(\text{PCl}_5 (g) \rightarrow \text{PCl}_3 (g) + \text{Cl}_2 (g)\)  has a half-life (\(t_{0.5}\)​) of 20 minutes. What is the time required for the concentration of \(\text{PCl}_5\) to decrease to 25% of its initial concentration?

The activation energy of a reaction is zero. The rate constant of the reaction is :
1. increases with an increase in temperature
2. decreases with a decrease in temperature
3. decreases with an increase in temperature
4. is nearly independent of temperature

If the half-life is independent of its initial concentration, then the order of the reaction is:
1. 0 
2. 1
3. 3
4. 2

In a zero order reaction for every 10° rise of temperature, the rate is doubled. If the
temperature is increased from 10°C to 100°C, the rate of the reaction will become
1. 256 times
2. 512 times
3. 64 times
4. 128 times

Which one of the following statements for the order of a reaction is incorrect?
1. Order is not influenced by stiochiometric coefficient of the reactants.
2. Order of reaction is sum of power to the concentration terms of reactants to express the rate of reaction.
3. Order of reaction is always whole number.
4. Order can be determined only experimentally.
 

Half-life period of a first order reaction is 1386s. The specific rate constant of the reaction
is
1. 5.0 x 10^-3s^-1
2. 0.5 x 10^-2s^-1
3. 0.5 x 10^-3s^-1
4. 5.0 x 10^-2s^-1

For the reaction, N₂ + 3H₂ $\to$2NH₃, if d[NH₃]/dt = 2x10^-4 mol L^-1s^-1, the value of -
d[H₂]/dt would be 
1. 3x10^-4 mol L^-1s^-1
2. 4x10^-4 mol L^-1s^-1
3. 6x10^-4 mol L^-1s^-1
4. 1x10^-4 mol L^-1s^-1

If a first-order reaction reaches 60% completion within 60 minutes, approximately how much time would it take for the same reaction to reach 50% completion? (log 4=0.60, log 5=0.69)
1. 50 min
2. 45 min
3. 60 min
4. 30 min

Consider the reaction
N₂(g) +3H₂(g) $\to$2NH₃(g)
The equality relationship between $\frac{\mathrm{d}\left[{\mathrm{NH}}_3\right]}{\mathrm{dt}}$ and $-\frac{\mathrm{d}\left[{\mathrm{H}}_2\right]}{\mathrm{dt}}$ is:
$$\begin{gathered} 1.<mspace\; width="1em"></mspace>+\frac{\mathrm{d}\left[{\mathrm{NH}}_3\right]}{\mathrm{dt}}<mspace\; width="1em"></mspace>=<mspace\; width="1em"></mspace>-\frac{1<mspace\; width="1em"></mspace>\mathrm{d}[\mathrm{H}{}_2]}{3<mspace\; width="1em"></mspace>\mathrm{dt}} \\ 2.<mspace\; width="1em"></mspace>+\frac{\mathrm{d}\left[{\mathrm{NH}}_3\right]}{\mathrm{dt}}<mspace\; width="1em"></mspace>=<mspace\; width="1em"></mspace>-\frac{2<mspace\; width="1em"></mspace>\mathrm{d}\left[{\mathrm{H}}_2\right]}{3<mspace\; width="1em"></mspace>\mathrm{dt}} \\ 3.<mspace\; width="1em"></mspace>+\frac{\mathrm{d}[\mathrm{NH}{}_3]}{\mathrm{dt}}<mspace\; width="1em"></mspace>=<mspace\; width="1em"></mspace>-\frac{3<mspace\; width="1em"></mspace>\mathrm{d}\left[{\mathrm{H}}_2\right]}{2<mspace\; width="1em"></mspace>\mathrm{dt}} \\ 4.<mspace\; width="1em"></mspace>+\frac{\mathrm{d}\left[{\mathrm{NH}}_3\right]}{\mathrm{dt}}<mspace\; width="1em"></mspace>=<mspace\; width="1em"></mspace>-\frac{\mathrm{d}\left[{\mathrm{H}}_2\right]}{\mathrm{dt}} \end{gathered}$$
 

Which of the following rate laws has an over all order of 0.5 for the reaction A + B + C $\to \hspace{0.17em}$product ?
1. R=k[A].[B].[C]                                            
2. R=k[A]^0.5[B]^0.5[C]^0.5  
3. R=k[A]^1.5[B]^-1[C]⁰  $<mspace\; width="1em"></mspace>$                                
4. R=k[A][B]⁰[C]^0.5

For a reaction of the type 2A+B $\to$2C, the rate of the reaction is given by $k{\left[A\right]}^2\left[B\right]$. When the volume of the reaction vessel is reduced to 1/4 th of the original volume, the rate of reaction changes by a factor of
(A) 0.25                                                          
(B) 16
(C) 64                                                             
(D) 4

The rate constant of a first-order reaction is typically determined from a plot of:
1. Concentration of reactant vs time t
2. Log (concentration of reactant) vs time t
3. \(\frac{1}{\text { concentration of reactant }} \text { vs time } t\)
4. Concentration of reactant vs log time t

The plot of log k versus $\frac{1}{T}$ is linear with a slope of
(1) $\frac{E_a}{R}$                                                    
(2) $\frac{-E_a}{R}$
(3) $\frac{E_a}{2.303R}$                                             
(4) $\frac{-E_a}{2.303R}$

The rate constant, the activation energy, and the Arrhenius parameter of a chemical reaction at 25°C are 3.0×10^-4 s^-1, 104.4 kJ mol^-1 and 6.0×10¹⁴s^-1 respectively.
The value of the rate constant as T → ∞ will be:
1. 2.0 × 10¹⁸ s^-1                                                  
2. 6.0 × 10¹⁴ s^-1
3. $\infty$                                                                   
4. 3.6 × 10³⁰ s^-1

87.5% of a radioactive substance disintegrates in 40 minutes. What is the half life of the substance ?
(1) 13.58 min                                                   
(2) 135.8 min
(3) 1358 min                                                    
(4) None of these