Question

The rate constant of a reaction at 500k and 700k are 0.02s and 0.07s respectively. Calculate the values of Ea and A.

Answer 1

$\huge\mathtt{\fbox{Solution:-}}$

$log \times \frac{ k_{2}}{ k_{1}} = \frac{ E_{a} }{2.303r} ( \frac{ T_{2} - T_{1} }{ T_{1} T_{2} } )$

$= > \: \: log \times \frac{0.07}{0.02} = ( \frac{ E_{a}}{2.303 \times 8.314 {jk}^{ - 1}mol } )[ \frac{700 - 500}{700 \times 500}$]

$= > \: \: 0.544 = E_{a} \times 5.714 \times {10}^{ - 4} \ /19.15$

$= > \: \: E_{a} = 0.544 \times 19.15 \ /5.714 \times {10}^{ - 4}$

$= > 18230.8j$

Since,

$k = {Ae}^{ - ea \ /rt}$

$0.02 = {Ae}^{ - 18230.8 \ /8.314 \times 500}$

$A = 0.02 \ /0.012$

$\huge\boxed{1.61}$

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