The reaction between hexacyanoferrate(II) ion and hydrogen peroxide in slightly acidic aqueous solutions (pH 4.64–5.35), and in the presence of molybdate ion as catalyst, was followed by means of a spectrophotometric technique at 420 nm. The catalyst was found to be very active, since a concentration as low as a thousandth part of that of hydrogen peroxide was enough to observe a notable increase in the reaction rate. The method chosen to obtain the kinetic quantitative data was that of the initial rates. This method led to non-integer kinetic orders in the three reactants for the two reaction pathways, non-catalytic and catalytic. Both paths showed acid catalysis, a negative (inhibitory) saline effect and rather low activation energies (40 ± 5 kJ mol⁻¹ for the non-catalytic path and 35 ± 4 kJ mol⁻¹ for the catalytic one). A mechanism compatible with the available experimental information has been proposed for each reaction pathway. In the absence of molybdate ion, only the hydrated form of the reducing agent, pentacyanoaquaferrate(II) ion, is assumed to react with the oxidant, and protonation of an iron(II)-hydrogen peroxide complex leads to an inner sphere electron transfer, with formation of a hydroxyl radical in the rate determining step. On the contrary, in the presence of catalyst, both forms of the reducing agent (hydrated and non-hydrated) are believed to react with the protonated peroxomolybdate(VI) complexes, the rate determining steps being now outer sphere electron transfer processes, yielding a molybdenum(V) complex as reaction intermediate instead of any free radical.