Explanation
Order of reaction with respect to O2:first
Order of reaction with respect to NO:second
Rate=k[NO]2[O]
Volume is decreased to half, then concentration of each reactants gets doubled.
Rate New=k[2NO]2[2O]
∴ Rate of reaction will increase to eight times of it’s initial value.
By using arrhenius equation, \implies \log k = \frac{-E_a}{2.303\space RT} + \log A
We get \implies \log A = \log(1.8 \times 10^{-5}) + \frac{94140}{2.303 \times 8.314 \times 313 } \implies (\log 1.8) - 5 + 15.7082) \implies 0.2553 - 5 + 15.7082 = 10.9635
\therefore \log A = (10.9634) = 9.914 \times 10^{10}
We get \implies \log A = \log(2.418 \times 10^{-5}) + \frac{179900}{2.303 \times 8.314 \times 546 } \implies (\log 2.418) - 5 + 15.7082) \implies 0.3834 - 5 + 17.209 = 12.6
\therefore \log A = (12.6) = 3.9 \times 10^{12}
In a first-order reaction the concentration of the reactant is decreased from 1.0{\text{ M to 0}}{\text{.25 M }} in 20 min.The rate constant of the reaction would be:
By first order kinetic rate constant,
k = \cfrac{2.303}{t} \log(\cfrac{a}{a-x})
a = 0.1 \space M
(a-x) = 0.05 \space M
t = 20 \space min
\implies k = \cfrac{2.303}{20} \log(\cfrac{0.1}{0.05})
\implies k = 0.0347 \space min^{-1}
Rate= \cfrac{dx}{dt} = k[A]^1
\implies 0.0347 \times 0.02
\implies 6.94 \times 10^{-4} \times M \space min^{-1}
For which order reaction a straight line is obtained along with x - axis by plotting a graph between half-life \left( {{t_{1/2}}} \right) and initial concentration 'a'.
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