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#### Question

Explain a graphical method to determine activation energy of a reaction.

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The rate constant for the first-order decomposition of H_{2}O_{2} is given by the following equation:

`logk=14.2-(1.0xx10^4)/TK`

Calculate E_{a} for this reaction and rate constant k if its half-life period be 200 minutes.

(Given: R = 8.314 JK^{–1} mol^{–1})

The rate of a reaction quadruples when the temperature changes from 293 K to 313 K. Calculate the energy of activation of the reaction assuming that it does not change with temperature.

The rate constant for the decomposition of hydrocarbons is 2.418 × 10^{−5 }s^{−1} at 546 K. If the energy of activation is 179.9 kJ/mol, what will be the value of pre-exponential factor.

The rate constant of a first order reaction increases from 2 × 10^{−2} to 4 × 10^{−2} when the temperature changes from 300 K to 310 K. Calculate the energy of activation (E_{a}).

(log 2 = 0.301, log 3 = 0.4771, log 4 = 0.6021)

The decomposition of A into product has value of *k *as 4.5 × 10^{3} s^{−1} at 10°C and energy of activation 60 kJ mol^{−1}. At what temperature would *k *be 1.5 × 10^{4} s^{−1}?