Q.State Law of Mass action.

In 1864, the Norwegian chemists Cato Maximilian Guldberg and Peter Waage generalised the experimental behaviour of many reversible reactions into the law of mass action:

"At any instant, the rate of a chemical reaction at a given temperature is directly proportional to the product of the active masses of the reactants at that instant."

For a single reactant with stoichiometric coefficient x, this reads $\text{Rate}\propto[\text{Reactant}]^x$, where the square brackets denote active mass -- molar concentration, $n/V$, in mol dm$^{-3}$ (or mol L$^{-1}$), with n the number of moles and V the container's volume.

For a general reversible reaction $xA+yB\rightleftharpoons lC+mD$, the law of mass action gives the forward rate $r_f=k_f[A]^x[B]^y$ and the backward rate $r_b=k_b[C]^l[D]^m$, where $k_f,k_b$ are the forward and backward rate constants. This is the mechanistic starting point for the entire equilibrium-constant framework: setting $r_f=r_b$ (the defining condition of equilibrium) and rearranging gives $K_C=k_f/k_b=\dfrac{[C]^l[D]^m}{[A]^x[B]^y}$ -- so the equilibrium constant is literally the ratio of the forward to the backward rate constant. Given any one of $k_f$, $k_b$ or $K_C$, the other two follow directly: for instance, if $k_f=2.5\times10^2$ and $K_C=50$, then $k_b=k_f/K_C=5$.

The law of mass action's picture of $K_C$ as an exact ratio of rate constants is, strictly, only an approximation -- more detailed treatment in chemical kinetics refines exactly how rate and concentration connect for a real reaction mechanism -- but the resulting equilibrium-constant expression itself, $K_C=\dfrac{[C]^l[D]^m}{[A]^x[B]^y}$, remains exactly correct as a description of the equilibrium concentrations, regardless of mechanism.