the amount demonstrated. Work is now in progress to get more information as to the origin of these unknown chlorinated hydrocarbons.
Literature Cited (1) Lunde, G., Int. Revueges. Hydrobiol., 52,265 (1967). (2) Lunde, G., J . Amer. Oil. Chem. Soc., 45,331 (1968). ( 3 ) Lunde, G., Biuchim. Biophys. Acta, 304, 76 (1973). (4) Lunde, G., J . rimer. Oil Chem. Soc., 49,44 (1972). (5) Holden, A. V., PCB Conference I1 Stockholm, 1972, National Swedish Environment Protection Board, Allmanna Forlaget, p 23, 1973. (6) Hylin, J . W., Spenger, R. E., Gunther, F. A,, Residue Rev., 26, 127 (1969).
(7) Suida, J. F., DeBernardi, J . F., Lloydia, 36, 107 (1973). (8) Lunde, G., Steinnes, E., “Halogens and sulphur in marine oils,” Proc. from 7th Nordic Symposium in Lipid Chemistry, Rsros, Norway, June 1973 (in press). (9) Tananger, A., Jordal, O., “About the halogen content and the sulphur distribution in marine oils,” ibid. (10) Schmitt, R. A., Zweig, G., J. Agr. Food Chem., 10, 481 (1962). (11) Lunde, G., J. Amer. Oil Chem. Soc., 48,547 (1971). ‘(12) Bjerk, J . E . , Nord.-Vet., 24,451 (1972). (13) Bowen, H. I. M., p 19, Academic Press, London and New York, N.Y., 1966.
Received for review April 19, 1974. Accepted September 30, 1974. Work supported by the Royal Norwegian Council for Scientific and Industrial Research.
Mechanism of Autoxidation of‘Manganese in Aqueous Solution Michael A. Kessick” Department of Environmental Science and Engineering, Rice University, Houston, Tex.
James J. Morgan Environmental Engineering Science, W. M. Keck Laboratories, California Institute of Technology, Pasadena, Calif.
The removal of manganous manganese from aqueous solution by oxidation with dissolved molecular oxygen a t pH of about 9.0 has previously been shown to be autocatalytic in nature ( I ) . Since the reaction production is solid, the autocatalysis is heterogenous and exact kinetic analysis of the reaction is difficult. Nevertheless at constant pH and under constant partial pressure of oxygen, an autocatalytic rate expression of the type
reaction solution could be predicted from existing thermodynamic data ( I ) . On the basis of such data, no precipitation was expected to occur in a carbonate-free system containing 5 X 10-4 mol of Mn(I1) per liter at pH of approximately 9.3 or below at 25°C. Reactions were therefore carried out in 25°C ammonia buffer solutions of maximum pH 9.02. Colorimetric procedures were used to determine the number of oxidizing equivalents per mole of suspended manganese when removal of soluble manganese through reaction with oxygen was complete. This enabled the calculation of empirical formulas MnO,. A combined gravimetric/volumetric technique was also used to determine equivalent weights of filtered product, and finally-the possibility of continued oxidation of the precipitated product was also investigated.
was shown to be obeyed during the initial stages of the oxidation. In this rate expression MnO, is used to denote a general empirical formula for the oxidation product, so that the values of x of 1, 1.5, and 2.0 would correspond to Mn(II), Mn(III), and Mn(1V) oxides, respectively, without concern for the degree of hydration. Experimentally the products have varied in composition from Mn01.3 to Mn01.9, depending on the conditions of reaction (2). Such conditions involved, for instance, supersaturation of the reaction solution with respect to manganous carbonate during the use of buffer systems similar to those in natural waters, or supersaturation with respect to manganous hydroxide at high pH. The formation of these variable reaction products indicates that a mixed removal mechanism may have been occurring. To characterize the products of oxidation alone, a study was undertaken where conditions were such that no supersaturation with respect to any species in the initial
Experimental ACS-certified reagents were used in makeup of reaction solutions and as reagents for subsequent analysis, except for manganous perchlorate, obtained from the G. F. Smith Chemical Co. USP-grade oxygen was used to presaturate reaction solutions. Two colorimetric procedures ( 3 ) were used in the characterization of suspended oxidation products. Oxidizing equivalents of manganese were estimated by the o-tolidine method. Spectrophotometer readings were taken within 2 min of the mixing of sample with reagents to reduce error arising from color fade. Total manganese was determined by the formaldoxime method. Here the order of addition of reagents to samples was important for reproducibility. The procedure depends upon the formation of a colored complex between manganese and formaldoxime a t high pH. Analysis of a typical 10-ml sample of approximately 5 X lO-4M Mn2+ solution involves addition of 2 ml of reagent and 5 ml of 5 N NaOH solution, followed by dilution to 100 ml and colorimetric measurement at 450 mp. Consistent color formation could be obtained only if reagent was added after base addition to the sample.
Analytical data indicate MnOOH to be the primary product resulting from manganese autoxidation in aqueous solution. Mechanisms for the formation of oxides of varying composition consistent with the initial formation of MnOOH are discussed.
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I I1
Table 11. Product Equivalent Weights Experimental
Figure 1. Glass reaction vessel, containing a ground rer with teflon blades Clamp A is on a short length of tygon tubing
glass
Calculated
pH 9.02
pH 8.79
M n02
Mn203
MnOOH
89.2 89.4
89.8 88.8
43.5
78.9
88.0
stir-
60 Table I. Empirical Formulas of Oxidation Products PH
8.98 8.88 8.79 8.69
10*[Mnltotai mol/l., A
101[Mn]oxidized equiv./l. B
Empirical formula,m MnO.
2.38 2.22 2.16
2.55 2.33 2.19
MnOl.64 MnOl.jz
1.15b
1.w
MnOl.51
MnOl.5i
+
x = 1 B/2A. bThese low values i,ndicate that much of the product had adhered t o the walls of, the reaction vessel, and was not in suspension at the end of the reaction.
Premixing base with reagent or addition of base after reagent to the sample led in each case to decreased color formation. For the equivalent weight determinations, oxide suspensions were filtered through 0.22-1 membrane filters. The filters and filter residue (5-8 mg) were allowed to dry to constant weight in a silica-gel desiccator a t room temperature over a period of three days. A dried sample was then divided into two parts carefully placed on small pieces of filter paper. Each piece of filter paper was weighed on a Mettler microbalance before and after addition of the solid product, the whole then being transferred to a titration flask for volumetric analysis. To the flask was added 1 gram of potassium iodide and 20 ml of deaerated 10% sulfuric acid. When there was no evidence of any solid residue remaining, the liberated iodine was titrated against sodium thiosulfate solution (0.99 x 10-2N) previously standardized against potassium iodate. From the results obtained, the equivalent weight, E, of the product as an oxidizing agent could be calculated from the formula x E = 0.99 x 10-22’ where y ml is the volume of thiosulfate solution equivalent to the iodine liberated by x mg of oxide product. The apparatus used for the experiments consisted of a 158
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t (min.) Figure 2. Continued oxidation of MnOr precipitate Time t is from point of addition of sand and increase in stirring rate
sealed, stirred reaction flask held in a thermostat bath a t 25°C (Figure 1).Samples were removed by opening clamp A and expelling sufficient solution under pressure of pure oxygen. These precautions were considered necessary to maintain the capacity of the ammonia buffer solution. A typical reaction solution was prepared by presaturating glass-distilled water at 25°C with oxygen, adding ammonia to 5 X 10-2M, and then adjusting the pH to the required value by addition of 70% perchloric acid. Addition of 10-IM manganous perchlorate to the stirred solution, followed by readjustment of the pH, initiated the reaction. No visible precipitation occurred during this preparation stage.
Results When a reaction had proceeded to a point when manganese could not be detected in the solution after filtration of the resulting suspension through a 0.22-p membrane filter, samples were obtained for colorimetric analysis to determine the product empirical formulas. Results of such experiments appear in Table I. At a similar point in two further reactions, samples were filtered for the determination of equivalent weights that are compared with those calculated for MnO2, Mn203, and MnOOH in Table 11. An additional oxidation reaction was carried out a t pH 8.98. When this reaction also had reached a point a t which no filterable manganese could be detected by the formaldoxime method, 10 grams of acid-washed 20-30 mesh beach sand was added to the resultant suspension, and the stirring rate was increased from a gentle agitation to the maximum possible for the apparatus (of the order of 4000 rpm) . Samples were withdrawn a t intervals during this period of violent agitation and analyzed, by the colorimetric procedures described above, to determine x in the
empirical formula MnO, of the product. Results are depicted in Figure 2 and indicate that continued oxidation of the solid product can occur under abrasive conditions.
Discuss ion The data presented indicate that the initial product of oxidation under the conditions described is an Mn(II1) oxide, and that it is probably MnOOH. X-ray investigations ( 4 ) have shown that y-MnzO3 will partially invert to y-MnOOH if kept in aqueous suspension for several months under these same conditions, indicating that the latter is the more stable form. On this basis, the formation of the compound Mn203eH20, which would also have an equivalent weight of 88.0, was considered unlikely. Characterization of the hydrate Mn203.Hz0 also does not appear to have been reported in the literature. The reaction conditions in this study were similar to those employed during the previous investigations ( 1 ) in which, as well as being autocatalytic in nature, the rate of manganous manganese removal from aqueous solution by autoxidation was shown to depend on PO, and [OH-]2, It is suggested then that the rate-determining step for this removal involves a one-electron transfer Mn(OH),(aq)
+
O2
-
MnOOH
+
HOz
for which the transition state is of the form HO--Mn(aq)---0-
--H- --O2
This would be consistent with the rate-determining step of a Haber-Weiss-type autoxidation mechanism ( 5 ) . The form of the transition state implies a hydrolysis pre-equilibrium, involving displacement of water molecules in the inner hydration shell by hydroxyl ions, prior to electron transfer. The probabilities of displacement of one or two water molecules by hydroxyl ions in the hydrated manganous species Mn OH+(aq) or Mn2+(aq), respectively, and simultaneous collision with an oxygen molecule must be considered low. Although thermodynamic data ( I ) suggest the existence of Mn(OH)z(aq) in free solution is unlikely in any appreciable concentration, a structure close to this could conceivably be formed in the double layer a t the surface of already precipitated product. Apart from the established heterogeneous autocatalytic nature of the reaction, it is also known that manganous ions are strongly absorbed to
the surface of manganese dioxide a t high pH, and that OH- ions are potential determining for this surface ( 3 ) . Further oxidation of the MnOOH precipitate under conditions designed to increase its surface area has been shown to occur. A similar phenomenon has been observed where Mn304 can be converted to an oxide as high as MnOl.dr on prolonged oxidation in an aqueous environment ( 4 ) . The results of this study are also in agreement with thermodynamic data which show a higher oxide, MnOz, to be the stable manganese compound under the reaction conditions imposed ( 1 ) . This possibility of further oxidation of the initial solid Mn(II1) precipitate demonstrates that rate of formation relative to rate of accretion may be an important factor in determining the nature of an oxidation product. An oxide precipitate, forming slowly a t low pH and probably only on the surface of existing product might be expected to be oxidized beyond Mn01,5 before it is effectively removed from contact with the aqueous phase by further deposition. Such could be the case, for instance, with the formation of the 6-MnOz sheets observed in the structure offerromanganese nodules found on the ocean floor (6). Under conditions of high pH, initial oxidation becomes very rapid with respect to coagulation, and again, the possibility of further oxidation of the colloidal precipitate prior to removal from contact with the aqueous phase by coagulation must be considered good. Highly oxidized products do, in fact, result under these conditions ( 1 ) . Where the rate of coagulation is close to that of initial oxidation, empirical formulas approximating MnOl 5 can be expected. The compounds, MnOx, where x is less than 1.5, must be considered as resulting from coprecipitation under appropriate conditions with manganous hydroxide or carbonate.
Literature Cited ( I ) Morgan, J . J., “Chemical Equilibria and Kinetic Properties of Manganese in Natural Waters,” in “Principles and Applications of Water Chemistry,” S. D. Faust and J . V. Hunter, Eds., Wiley, New York, N.Y., 1967. ( 2 ) Stumm, W.. Morean. J . 3.. “Aauatic Chemistry. An Introduction Emphasizing Chemical Eqiilibria in Katuial Waters,” Wiley, New York, N.Y., 1970. (31 . . Stumm. W.. Morgan. J . J.. J Amer Water Works Ass.. 57. io7 (i965j. (4) Bricker, O., Amer. Mineral., 50,1296 (1965). (5) Fallab, S., AngeLc. C h e n . Internat. Edit.,6,496 (1967). (6) Arrhenius, G., “Pelagic Sediments,” in “The Sea,” Vol. 111, M . K.Hill, Ed., Wiley Interscience, New York, S . Y . , 1963. Y
Received for review June 10. 1974. Accepted October22, 1974.
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