Question

The reaction $2A+B+C\to D+E$ is found to be first-order in A, second in B and zero-order in C.

(i) Give the rate law for the reaction in the form of the differential equation.
(ii) What is the effect in the rate of increasing concentrations of A, B and C two times?

• A. (i) $\frac{dx}{dt}=k\left[A\right][B{]}^2$ (ii) rate increases by $8$ times
• B. (i) $\frac{dx}{dt}=k\left[A\right][B{]}^{-2}$ (ii) rate increases by $1/4$ times
• C. (i) $\frac{dx}{dt}=k\left[A{]}^{-1}\right[B{]}^{-2}$ (ii) rate increases by $1/8$ times
• D. None of these

Answer 1

The correct option is A (i) $\frac{dx}{dt}=k\left[A\right][B{]}^2$ (ii) rate increases by $8$ times

$2A+B+C\to D+E$

$(\frac{dx}{dt}{)}_1=K_1[A\left]\right[B{]}^2[C{]}^0$

if increases conc. 2 time

${\left(\frac{dx}{dt}\right)}_2=K_1\left[2A\right]\left[2B{]}^2\right[2C{]}^0=8{\left(\frac{dx}{dt}\right)}_1$

$\Rightarrow$ rate increase by 8 time

Hence, the correct option is $\text{A}$