The rate constant for a first order reaction is $4.606\times {10}^{-3} s^{-1}$. The time required to reduce 2.0 g of the reactant to 0.2 g is:

1.	200 s	2.	500 s
3.	1000 s	4.	100 s

An increase in the concentration of the reactants of a reaction leads to a change in:

1. heat of reaction

2. threshold energy

3. collision frequency

4. activation energy

The half-life for a zero-order reaction having 0.02 M initial concentration of reactant is 100 s. The rate constant (in mol L^-1 s^-1) for the reaction is:

1. $1.0\times {10}^{-4}$

2. $2.0\times {10}^{-4}$

3. $2.0\times {10}^{-3}$

4. $1.0\times {10}^{-2}$

What does Z_AB represent in the collision theory of chemical reactions?

1. The fraction of molecules with energies greater than E_a

2. The collision frequency of reactants, A and B

3. Steric factor

4. The fraction of molecules with energies equal to E_a

For a reaction, activation energy $E_a=0$ and the rate constant at 200 K is    1.6 $X$ ${10}^6{s}^{-1}$. The rate constant at 400K will be [Given that gas constant, R=8.314 J $K^{-1}$ $mo{l}^{-1}$]

1.  3. 2 x 10⁴ s^-1
2. 1.6 x 10⁶s^-1

3. 1.6 x10³ s^-1

4. 3.2 x 10⁶ s^-1

A first-order reaction has a rate constant of 2.303$\times {10}^{-3}$ ${\mathrm{s}}^{-1}$. The time required for 40 g of this reactant to reduce to 10 g will be [Given that ${\log }_{10}2=0.3010$]

1. 230.3 s

2. 301 s

3. 2000 s

4. 602 s

For a reaction A$\to$B, enthalpy of reaction is $-4.2$ $\mathrm{kJ}\hspace{0.17em}{\mathrm{mol}}^{-1}$ and enthalpy of activation is $9.6$ $\mathrm{kJ}\hspace{0.17em}{\mathrm{mol}}^{-1}$. The correct potential energy profile for the reaction is:

1.		2.
3.		4.

The slope of Arrhenius Plot (ln k v/s $\frac{1}{\mathrm{T}}$) of the first-order reaction is $-5\times {10}^3$ $\mathrm{K}$. The value of E_a of the reaction is:

[Given R = 8.314 JK^-1 mol^-1]

For the chemical reaction ${\mathrm{N}}_2\left(\mathrm{g}\right)+3{\mathrm{H}}_2\left(\mathrm{g}\right)$ $\rightleftharpoons$ $2{\mathrm{NH}}_3\left(\mathrm{g}\right)$ the correct option is:

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

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

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

4. $-\frac{\mathrm{d}\left[{\mathrm{N}}_2\right]}{\mathrm{dt}}=\frac{1}{2}\frac{\mathrm{d}\left[{\mathrm{NH}}_3\right]}{\mathrm{dt}}$

If the rate constant for a first order reaction is k, the time (t) required for the completion of 99% of the reaction is given by:

1. t=2.303/k

2. t=0.693/k

3. t=6.909/k

4. t=4.606/k