For a first-order reaction A $\to$ Products, the rate of reaction at [A] = 0.2 M is 1.0 x 10^-2 mol litre^-1 min^-1. The half-life period for the reaction will be:

1.	832 sec	2.	440 sec
3.	416 sec	4.	14 sec

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

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

In acidic medium the rate of reaction between (BrO₃)^- and Br^- ions is given by the expression

$-\frac{d\left({\mathrm{BrO}}_3^{-}\right)}{dt}=K\left[{\mathrm{BrO}}_3^{-}\right]\left[{\mathrm{Br}}^{-}\right]{\left[{\mathrm{H}}^{+}\right]}^2$

It means:

1. Rate constant of overall reaction is 4 sec^-1

2. Rate of reaction is independant of the concentration of acid

3. The change in pH of the solution will not affect the rate

4. Doubling the conc. of H⁺ ions will increase the reaction rate by 4 times

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

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:

For the reaction $A \stackrel{ }{\to }nB$, at the point of intersection of two curves show, the [B] is can be given by:

1. $\frac{n{A}_0}{2}$

2. $\frac{A_0}{n-1}$$$

3. $\frac{n{A}_0}{n+1}$

4. $\left[\frac{n-1}{n+1}\right]A_0$

The following mechanism has been proposed for the reaction of NO with Br₂ to form NOBr:

NO(g) + Br₂(g) $\rightleftharpoons$ NOBr₂(g)

NOBr₂(g) + NO(g) $\to$2NOBr(g)

If the second step is the rate determining step, the order of the reaction with respect to NO(g) will be:

1. 1

2. 0

3. 3

4. 2

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