Reaction between compounds (A) and (B) occurs as follows

A (g) + 2B (g) → 2C (g)

Following results were obtained

The correct rate- law for the said reaction is (r and k rate and rate constant respectively)

• r = k[A]²B

• r = k[A][B]²

• r = k[A][B]

• r = k[A]²[B]²

In the reaction 2NO₂ $\underset{{\mathrm{k}}_2}{\overset{{\mathrm{k}}_1}{\rightleftharpoons }}$ N₂O₄; the rate of disappearance of NO₂ is equal to

• 2$\frac{{\mathrm{k}}_1}{{\mathrm{k}}_2}$ [NO₂]²

• 2k₁[NO₂]² - 2k₂[N₂O₄]

• 2k₂[NO₂]² - k₂[N₂O₄]

• (2k₁ - k₂)[NO₂]

The rate constant for the raection 2N₂O₅ → 4NO₂ + O₂ is 3.0 × 10^-5 s^-1. If rate is 2.40 × 10^-5, then concentration of N₂O₅ (in mol/L) is

• 1.4

• 1.2

• 0.04

• 0.8

The concentrations of the reactant A in the reaction A → B at different times are given below

Concentration (M)	Time (minutes)
0.069	0
0.052	17
0.035	34
0.018	51

The rate constant of the reaction according to the correct order of the reaction is

• 0.001 M/min

• 0.001 min^-1

• 0.001 min/M

• 0.001 M^-1 min^-1

The half-life of two samples are 0.1 and 0.8 s. Their respective concentration are 400 and 50 respectively. The order of the reaction is

• 0

• 2

• 1

• 4

Acid hydrolysis of ester is first order reaction and rate constant is given by

k = $\frac{2.303}{\mathrm{t}} \log \frac{{\mathrm{V}}_{\infty } - {\mathrm{V}}_0}{{\mathrm{V}}_{\infty } - {\mathrm{V}}_{\mathrm{t}}}$

where, V₀, V_t and V_∞ are the volume of standard NaOH required to neutralise acid present at a given time, if ester is 50% neutralised then

• V_∞ = V_t

• V_∞ = (V_t - V₀)

• V_∞ = 2V_t - V₀

• V_∞ = 2V_t + V₀

Consider the following reaction,

The reaction is of first order in each diagram, with an equilibrium constant of 10⁴. For the conversion of chair form to boat form ${\mathrm{e}}^{-{\mathrm{E}}_{\mathrm{a}}/\mathrm{RT}}$ = 4.5 × 10^-8 m at 298 K with pre-exponential factor of 10¹² s^-1. Apparent rate constant (=k_A/ k_B) at 298 K is

• 4.35 × 10⁴ s^-1

• 4.35 × 10⁸ s^-1

• 4.35 × 10^-8 s^-1

• 4.35 × 10¹²s^-1

358.

The unit of rate constant for a zero order reaction is

• mol L^-1 s^-1

• L mol^-1 s^-1

• L² mol^-2s^-1

• s^-1

In the reversible reaction,

$2{\mathrm{NO}}_2 \underset{{\mathrm{k}}_2}{\overset{{\mathrm{k}}_1}{\rightleftharpoons }} {\mathrm{N}}_2{\mathrm{O}}_4$

the rate of disapperance of NO₂ is equal to

• $\frac{2{\mathrm{k}}_1}{{\mathrm{k}}_2}[{\mathrm{NO}}_2{]}^2$

• 2k₁[NO₂]² - 2k₂[N₂O₄]

• 2k₁[NO₂]² - k₂[N₂O₄]

• (2k₁-k₂)[NO₂]

A chemical reaction was carried out at 300 K and 280 K. The rate constants were found to be k₁ and k₂ respectively. Then

• k₂= 4k₁

• k₂=2k₁

• k₂=0.25k₁

• k₂=0.5k₁