In the following reaction, $\begin{array}{c}\mathrm{xA}⟶\mathrm{yB}\\ \log \left[-\frac{\mathrm{d}\left[\mathrm{A}\right]}{\mathrm{dt}}\right]=\log \left[\frac{\mathrm{d}\left[\mathrm{B}\right]}{\mathrm{dt}}\right]+0.3\end{array}$ where, negative sign indicates rate of disappearance of the reactant. Thus, $\mathrm{x}: \mathrm{y}$ is

Rate of formation of ${\mathrm{SO}}_3$ in the following reaction:

$2{\mathrm{SO}}_2+{\mathrm{O}}_2\to 2{\mathrm{SO}}_3$, is $100 \mathrm{g} {\min }^{-1}.$ Hence, the rate of disappearance of ${\mathrm{O}}_2$ is:

$\mathrm{aA}+\mathrm{bB}⟶$ Product, $\mathrm{dx}/\mathrm{dt}=\mathrm{k}\left[\mathrm{A}{]}^{\mathrm{a}}\right[\mathrm{B}{]}^{\mathrm{b}}.$ If concentration of $\mathrm{A}$ is doubled, rate is four times. If concentration of $\mathrm{B}$ is made four times, rate is doubled. What is relation between rate of disappearance of $\mathrm{A}$ and that of $\mathrm{B}?$

What will be the order of reaction and rate constant for a chemical change having ${\mathrm{logt}}_{50\%}\mathrm{Vs}$. log concentration $\left(\mathrm{a}\right)$ curves as ?

A reaction follows the given concentration-time graph. The rate for this reaction at $20$ seconds will be: