I was reading the paper Relaxation Kinetics of Ferric Thiocyanate (Goodall et. al, 1972) and I came across a passage which read
$\space$ Reaction (1) is the simplest representation of the equilibrium between ferric and thiocyanate ions and their complex. The ratio $k_{12}/k_{21}$ is the equilibrium constant $K_c$, which has the value $1391{\text {mole}^{-1}}$ at 298°K and ionic strength 0.5 mole $\text{kg}^{-1}$ ($\textit 8$) $$\ce{[Fe(H_2O)_6]^{3+}}+\ce{SCN^{-}}\underset{k_{21}}{\stackrel{k_{12}}{\rightleftharpoons}} \ce{[Fe(H2O)5SCN]}^{2+}+\ce{H2O} \space\space\space\space\space\space(1)$$
However, the paper later makes clear that reaction (1) cannot be single step
Reaction (1) is insufficient to explain the mechanism of the equilibration, since the relaxation rate decreases with increase in the concentration of hydrogen ion in the solution. This is explained by introducing eqns. (2)-(4). Most of the reaction proceeds through the hydroxo complex, which is present in amount determined by the pH and the acid dissociation constant of the hexaquo complex, $\ce{K_{OH}} = 1.89 \times 10^{-3} \text{mole}^{-1}$ at 298 $^\circ$K and ionic strength 0.5 mole $\text{kg}^{-1}$ ($\textit{10}$). $$\ce{[Fe(H_2O)_6]^{3+}} \xrightleftharpoons{\text{fast}} \ce{[Fe(H_2O)_5OH]^{2+}}+\ce{H^+}\space\space\space\space\space\space(2)$$ $$\ce{[Fe(H_2O)_5OH]^{2+}}+\ce{SCN^-}\underset{k_{43}}{\stackrel{k_{34}}{\rightleftharpoons}}\ce{[Fe(H_2O)_4(OH)SCN]^{+}}\space\space\space\space\space\space(3)$$ $$\ce{H^+}+\ce{[Fe(H_2O)_4(OH)SCN]^{+}\xrightleftharpoons{\text{fast}}\ce{[Fe(H2O)5SCN]}^{2+}}\space\space\space\space\space\space(4)$$
How then can it be possible that $K_c = k_{12}/k_{21}$ for reaction (1), since it is not elementary?
