The decomposition of formic acid on gold surface follows first order kinetics- If the rate constant at 300 K is 1-0 - 10-3 s-1 and the activation energy Ea - 11-488 kJ mol-1- the rate constant at 200 K is - 10-5 s-1-Round off to the nearest Integer-

By using Arrhenius equation, logK_2K_1 = E_a2.303R[1T_1 1T_2 ] Where, K_2 is rate constant at T_2 ( 300 K) K_1 is rate constant at T_1 (200 K) R is univers ...

You visited us 1 times! Enjoying our articles? Unlock Full Access!

Byju's Answer

Standard XII

Chemistry

Zero Order Reaction

The decomposi...

Question

The decomposition of formic acid on gold surface follows first order kinetics. If the rate constant at $300 K$ is $1.0 \times 10^{- 3} s^{- 1}$ and the activation energy $E_{a} = 11.488 {kJ mol}^{- 1}$ , the rate constant at $200 K$ is $\times 10^{- 5} s^{- 1}$ .(Round off to the nearest Integer).

Open in App

Solution

By using Arrhenius equation,
$log \frac{K_{2}}{K_{1}} = \frac{E_{a}}{2.303 R} [\frac{1}{T_{1}} - \frac{1}{T_{2}}] Where, K_{2} is rate constant at T_{2} (300 K) K_{1} is rate constant at T_{1} (200 K) {R is universal gas constant = 8.314 J K}^{- 1} {mol}^{- 1} Substituting the values, we get log \frac{1.0 \times 10^{- 3} s^{- 1}}{K_{1}} = \frac{11.488 \times 1000}{2.303 \times 8.314} [\frac{1}{200} - \frac{1}{300}] log \frac{10^{- 3}}{K_{1}} = 600 \times \frac{3 - 2}{600} log \frac{10^{- 3}}{K_{1}} = 1 \Rightarrow 10 = \frac{10^{- 3}}{K_{1}} \Rightarrow K_{1} = 10^{- 4} s^{- 1} S o, x \times 10^{- 5} = 10^{- 4} \Rightarrow x = 10$

Suggest Corrections

Similar questions

Q. The decomposition of formic acid on a gold surface follows first order kinetics and specific rate constants are

5.5 \times 10^{- 4} s^{- 1}

and

9.2 \times 10^{- 3} s^{- 1}

140^{o} C

and

185^{o} C

respectively. Energy of activation is:

Q. Decomposition of

N H_{3}

on gold surface follows zero order kinetics. If rate constant

K

5 \times 10^{- 4} M

s^{- 1}

, rate of formation of

N_{2}

will be:

Q. At

400 K

, the rate constant of a chemical reaction is

3 \times 10^{- 4} s^{- 1}

and at

300 K

, the rate constant is

6 \times 10^{- 5} s^{- 1}

. Find the value of

E_{a}

i.e. activation energy required for the reaction.
Take:

{log}_{10} (0.2) = 0.7

Q. For the reaction,

a A + b B \to c C + d D,

the plot of

log K

1 / T

is given below:

The temperature at which the rate constant of the reaction is
$10^{- 4} s^{- 1}$ is ________ $K .$ (Rounded off to the nearest integer)

[Given: The rate constant of the reaction is $10^{- 5} s^{- 1}$ at $500 K .]$

The first order rate constant for the decomposition of ethyl iodide by the reaction:

$C_{2} H_{5} I (g) \to C_{2} H_{4} (g) + H I (g)$

at $600 K$ is $1.60 \times 10^{- 5} s^{- 1}$ . Its energy of activation is $209 k J m o l^{- 1}$ . The rate constant of the reaction at $700 K$ will be:

Join BYJU'S Learning Program

The decomposition of formic acid on gold surface follows first order kinetics. If the rate constant at 300 K is 1.0×10−3 s−1 and the activation energy Ea=11.488 kJ mol−1, the rate constant at 200 K is ×10−5 s−1.(Round off to the nearest Integer).

The decomposition of formic acid on gold surface follows first order kinetics. If the rate constant at $300 K$ is $1.0 \times 10^{- 3} s^{- 1}$ and the activation energy $E_{a} = 11.488 {kJ mol}^{- 1}$ , the rate constant at $200 K$ is $\times 10^{- 5} s^{- 1}$ .(Round off to the nearest Integer).