Autocatalytic oxidation of manganese

An autocatalytic reaction is one promoted by its own reaction products. A good example in geochemistry is the oxidation and precipitation of dissolved Mn11 by C 2(aq). The reaction is slow in solution, but is catalyzed by the precipitated surface and so proceeds increasingly rapidly as the oxidation product accumulates. Morgan (1967) studied in the laboratory the kinetics of this reaction at 25 °C and pH 9. 

He added an Mn11 solution to a stirred beaker containing a pH buffer, through which a mixture of oxygen and nitrogen gas was being bubbled. 

The 02(aq) reacted with the Mn11, forming a colloidal precipitate of oxidized manganese, likely composed of a mixture of Mnm and MnIv. We can write this reaction in a simple form as, 

As we start to model the reaction, we note the llnl thermodynamic dataset thermo.dat does not contain a redox couple linking Mnm and Mn11, as required for our purposes, but this is easily remedied by adding to the database the coupling reaction, 

A second complication to modeling the reaction is that it is necessary to allow oxidation at the start of the experiment, so the catalyst can begin to form. The nature of the initial reaction is not known, but it may be promoted by small amounts of colloids or particles, perhaps MnC03, present at the onset. A simple strategy, in light of our lack of knowledge, is to set in the initial system a small amount of Mn(OH)3(s) to represent the system s initial capacity to catalyze the reaction. 

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Fig. 28.2. Variation in the concentration of Mn11 versus time in a simulation of the autocatalytic oxidation of manganese at pH 9.5 and 25 °C. Squares are results of a laboratory experiment by Morgan (1967). Broken line shows the result of a simulation of catalytic oxidation, assuming the catalyzing surface has a constant area.

Fig. 28.3. Concentration of Mn11 versus time in simulations of the autocatalytic oxidation of manganese at pH 9.0, 9.3, and 9.5, at 25 °C, compared to results of laboratory experiments (symbols) by Morgan (1967). Simulations made assuming rate law of a form carrying mMa++, rather than mMnn. Rate constant in the simulations is taken to be 1013 molal-4 s-1, and the initial catalyst mass is 0.6 mg (pH 9.0), 5 mg (9.3), and 6 mg (9.5).

This kinetic equation involves an additive member containing the product of oxidation (in contrast to Fe(II) oxidation). It is an autocatalytic reaction. Therefore, when removing manganese from water, employing so-called manganese filters, it is first necessary to form a layer of higher hydrated oxides of manganese on the carrier, which is usually sand. 

The formation of manganese oxide solid phases (Mn02) enhances the oxidation and thus the process is autocatalytic. The process is also enhanced by bacteria and organic chelates with carboxyl functional groups. The initial steps in the electron transfer process can be represented simply by... 

A partial explanation of the observations cited is available from facts known to analytical chemists under certain conditions potassium permanganate decomposes to manganese dioxide with evolution of oxygen. A solution 0.04 Al in sulfuric acid decompo.ses nearly 20 times as fast as a neutral solution. On the other hand, decomposition is accelerated also by alkali as well as by manganese dioxide. Thus decomposition of reagent in the course of an oxidation is an autocatalytic process. This reaction probably accounts for the excessive consumption of permanganate in the Attenburrow experiment. 

As in all conversions of this type, which are autocatalytic, the induction period is relatively long. Catalysts are used to shorten it. These catalysts are soluble salts of cobalt, chromium, vanadium or manganese, usually acetates. The oxidation rate rises with the number of carbon atoms in the hydrocarbon and with the extent to which the chain is linear. Thus, if it is l for ethane, it is as high as 100 for propane, 500 for n-butane and 1000 for n-pentane. 

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