Q.Why is desorption important for a substance to act as a good catalyst?

A catalyst works by providing an alternative reaction pathway with lower activation energy — the minimum energy needed to reach the activated complex — letting far more reactant molecules react at a given temperature and thereby increasing the rate. Two theories explain the mechanics of this lowering.

The intermediate compound formation theory (mainly for homogeneous catalysis) has the catalyst C combine with a reactant to form a short-lived intermediate, which then reacts with the remaining reactant to give the product and regenerate C: for $A + B \rightarrow AB$, the catalysed route runs $A + C \rightarrow AC$ then $AC + B \rightarrow AB + C$, with BOTH catalysed steps carrying lower activation energy than the single uncatalysed step. Four textbook illustrations: the Friedel–Crafts reaction of benzene with methyl chloride via the $\text{AlCl}_3$-derived intermediate $[\text{CH}_3^+][\text{AlCl}_4^-]$; $\text{KClO}_3$'s thermal decomposition via the $\text{MnO}_2$-derived intermediate $\text{KMnO}_4$ (i.e. $\text{MnO}_3^-$-type species in the source reaction); $\text{H}_2$+$\text{O}_2$ forming water via the Cu-derived intermediate $\text{Cu}_2\text{O}$; and HCl's oxidation by air via the $\text{CuCl}_2$-derived intermediates $\text{Cu}_2\text{Cl}_2$ and $\text{Cu}_2\text{OCl}_2$. This theory explains a catalyst's specificity (only certain reactants can form the needed intermediate with it) and the rate's rise with catalyst concentration, but it fails to explain catalytic poisons/promoters and cannot account for HETEROGENEOUS catalysis, where catalyst and reactants aren't even in the same phase to form a chemical intermediate. …