We know that the rate of a reaction $\ce{aA + bB \longrightarrow cC + dD}$, the rate of the forward reaction is given by

$\mathrm{Rate = k[A]^p[B]^q}$ where $\rm a\neq p$ and $\rm b\neq q$ according to Chemical Kinetics.

However when studying Chemical Equilibrium, $\ce{aA + bB<=> cC + dD}$ when we write the forward and backward reaction rates and equate them, we write them as

$\rm Forward Rate = k[A]^a[B]^b$ where the exponents are equal to the stoichiometric coefficients.

How is this the case?

We know that the rate of a reaction $\ce{aA + bB -> cC + dD}$, the rate of the forward reaction is given by $r_\mathrm f = k_\mathrm f[\ce A]^p[\ce B]^q$ where $ a\neq p$ and $ b\neq q$ according to Chemical Kinetics.

However when studying Chemical Equilibrium, $\ce{aA + bB <=> cC + dD}$ when we write the forward and backward reaction rates and equate them, we write them as $r_\mathrm f = k_\mathrm f[\ce A]^a[\ce B]^b$ where the exponents are equal to the stoichiometric coefficients.

How is this the case?

We know that the rate of a reaction $\ce{aA + bB \rightarrow cC + dD}$, the rate of the forward reaction is given by

$Rate = k[A]^p[B]^q$ where $a\neq p$ and $b\neq q$ according to Chemical Kinetics.

However when studying Chemical Equilibrium, $\ce{aA + bB<=> cC + dD}$ when we write the forward and backward reaction rates and equate them, we write them as

$Forward Rate = k[A]^a[B]^b$ where the exponents are equal to the stoichiometric coefficients.

How is this the case?

We know that the rate of a reaction $\ce{aA + bB \longrightarrow cC + dD}$, the rate of the forward reaction is given by

$\mathrm{Rate = k[A]^p[B]^q}$ where $\rm a\neq p$ and $\rm b\neq q$ according to Chemical Kinetics.

However when studying Chemical Equilibrium, $\ce{aA + bB<=> cC + dD}$ when we write the forward and backward reaction rates and equate them, we write them as

$\rm Forward Rate = k[A]^a[B]^b$ where the exponents are equal to the stoichiometric coefficients.

How is this the case?