Environ. Sci. Technoi. 1992, 26, 1327-1333
Kinetic Modeling of Manganese( I I ) Oxidation by Chlorine Dioxide and Potassium Permanganate John E. Van Benschoten* and We1 Lln Department of Civil Engineering, State University of New York at Buffalo, Buffalo, New York 14260
Wllllam R. Knocke Department of Civil Engineering, Virglnla Polytechnic Institute and State University, Blacksburg, Virginia 2406 1
w Oxidation of Mn(1I) by potassium permanganate and chlorine dioxide was investigated for the environmental conditions of pH, temperature, and initial Mn(I1) concentrations typically observed in water treatment. For the conditions examined, oxidation reactions are rapid, with completion of the reactions within -30 s. A kinetic model based on solution-phase oxidation, adsorption, and surface oxidation mechanisms was developed that successfully simulates experimental data. The model was capable of simulating the experimental data (r2= 0.90 and 0.75 for KMn04 and CIOz, respectively) over a range of solution pH (5.5-8), temperature (2-25 "C), and initial Mn(I1) concentrations (0.4-1.25 mg/L). The kinetic model was solved numerically, and rate constants and activation energies (10-50 kJ/mol) were derived for the three mechanisms simulated. For reactive oxidants such as KMnO, and CIOz,adsorption of Mn(I1) to the oxide surface is the rate-limiting step. Because of rapid surface oxidation reactions, concentrations of adsorbed Mn(I1) were predicted to be low for the conditions studied. This result is in contrast to less reactive oxidants where surface adsorption is rapid relative to solution and surface oxidation reactions. Introduction Manganese exists in a number of oxidation states, but only the Mn(II), Mn(III), and Mn(1V) species contribute significantly to the aqueous chemistry of the metal (I). Under reducing conditions such as groundwaters and the bottom waters of stratified lakes and reservoirs, Mn(I1) is the most stable species. For oxygenated waters, Mn(II1) and Mn(IV) are the stable oxidation states, existing as the sparingly soluble oxide and hydroxide solid phases MnOOH(s) and MnOz(s), respectively. For potable water, a secondary maximum contaminant level for Mn(I1) has been set at 0.05 mg/L in order to reduce aesthetic problems such as discoloration of laundry and plumbing fixtures. Mn(I1) typically is removed by oxidation of soluble Mn(I1) to a colloidal precipitate followed by solid/liquid separation via sedimentation and filtration. Although removal of Mn(I1) from solution by oxidation with molecular oxygen has been studied extensively (2-5), the reaction rate at the pH of most natural waters is too slow to be of practical significance in most water treatment processes. Consequently, more reactive oxidants such as potassium permanganate and chlorine dioxide are used in treatment plants for Mn(I1) oxidation. While Mn(I1) oxidation with strong oxidant addition has been practiced in the water industry for many years, knowledge of reaction kinetics and removal mechanisms is limited. In this paper, experimental data and modeling results of Mn(I1) oxidation by KMn0, and CIOz are presented. Specific objectives of the paper are to (1) illustrate the dependence of Mn(I1) oxidation on variations in temperature, pH, and initial Mn(I1) concentrations commonly 0013-936X/92/0926-1327$03.00/0
encountered in potable water treatment; (2) present a kinetic model of Mn(I1) oxidation and test the model using experimental data; and (3) evaluate through model simulations the relative significanceof solution-phase oxidation, adsorption, and surface oxidation mechanisms for removal of Mn(I1). As used in this paper, solution-phaseoxidation is defined as the interaction of aqueous-phase Mn(I1) ions with oxidant to produce the manganese solid phase, Mn02(s). Surface oxidation is defined as a heterogeneous reaction involving oxidation of Mn(I1) that is adsorbed to the Mn02(s)surface. Assuming that colloidal manganese can be removed by sedimentation and filtration processes, both the solution-phase oxidation and adsorption pathways result in removal of Mn(I1) from the aqueous phase. Background The stoichiometry of Mn(I1) oxidation by KMn0, and CIOz based on balanced half-reactions is as follows: 3Mn2+ 2Mn04- 40H- = 5Mn02(s) 2Hz0 (1)
+
+
+
Mn2+ + 2C102 + 40H- = Mn02(s)+ 2C1O2-
+ 2H20
(2) The number of studies examining oxidation of Mn(I1) by reactive oxidants under environmental conditions typical in water treatment is limited. A review of research involving reactions of Mn(I1) and KMnO, (6) indicates that experiments often have been carried out in acid media to avoid complications that arise from precipitation of manganese(I1) oxides and heterogeneous oxidation reactions. Previous studies of Mn(I1) oxidation under conditions typical of water treatment for both KMn04 (7,8) and C102 (8-10) have not examined the kinetics of the oxidation process. In contrast to more reactive oxidants, Mn(I1) oxidation by molecular oxygen has been studied extensively by Morgan and co-workers (2-5). Morgan (3) reported results of Mn(I1) oxidation kinetic experiments which examined the effect of oxygen partial pressure, pH, temperature, and solution chemistry on reaction rates. He found that the firstthree variables exerted a strong influence on oxidation rates while background ions such as borate, sulfate, and inorganic carbon species influenced reaction rates to a lesser extent. An autocatalytic model for Mn(I1) oxidation was proposed that included solution-phase and surface oxidation mechanisms. For oxidation of Mn(I1) by molecular oxygen, Morgan and co-workers assumed that adsorption of Mn(I1) to oxide surfaces was rapid relative to the two oxidation reactions. Experiments were conducted in the presence of excess dissolved oxygen and preformed MnO,(s) solids. The proposed autocatalytic model was as follows: -d[Mn2+]/dt = ko[Mn2+]+ k[Mn2+][Mn0,] (3) Evidence in support of the autocatalytic nature of the reaction included nonstoichiometric oxidation products
0 1992 American Chemical Society
Environ. Sci. Technol., Vol. 26,No. 7, 1992
1327
(i.e., MnO, with 1.3 C x C 1.9) and experimental results that demonstrated large sorption capacities of Mn(I1) at slightly alkaline solutions. At pH 7.5, Morgan (2)reported sorption of Mn(I1) in excess of 0.5 mol of Mn(I1) adsorbed per mole of MnOz(s). The capacity of the oxide surface at pH 9 was -2 mol of Mn(I1) adsorbed per mole of MnOz(s) (11). The dependence of the surface oxidation rate in eq 3 was found to be first order in PO, and second order with respect to [OH-] ( I I ) , a result consistent with the kinetics of Fe(I1) oxidation by molecular oxygen (12, 13). The oxidation of Mn(I1) in the presence of a-FeOOH, y-FeOOH, silica, and 6-A1203was investigated by Davies and Morgan (14). Kinetic data were accounted for by a rate expression written in terms of concentration of surface species on the oxide surface. d[Mn2+]
--=
k*
(>SOH)[Mn2+] APO,
(4) dt W+I2 In eq 4,(>SOH) is the surface concentration of hydroxyl sites, A is the mass solids concentration, and p 0 2 is the partial pressure of oxygen. The surface hydroxyl site concentration was hypothesized to be proportional to the concentration of surface Mn(I1) species. The surface coordination chemistry model described surface speciation in terms of acid/base and metal-binding equilibria at the oxide surface. Experimental Section Laboratory experiments were conducted using synthetically prepared water samples made by the addition of 0.25 mequiv/L Na2S04,1 mequiv/L CaCl,, and 1-4 mequiv/L NaHC03 to deionized water. The initial Mn(I1) concentration typically was 1mg/L, prepared by addition of MnSO, to synthetically prepared water (ionic strength approximately 0.01 M). Potassium permanganate stock solutions of 1mg/mL were prepared by dissolving KMnO, crystals in deionized water and stored in brown, groundglass bottles. All other solutions also were prepared using deionized water (Super-Q, Millipore Corp, Bedford, MA). Chlorine dioxide was generated using the sulfuric acid/sodium chlorite method (15). To increase the yield of chlorine dioxide and decrease the background concentrations of chlorite and chlorate, the initial reagent concentrations were doubled and the receiving solution was precooled (16). Glassware used for oxidation studies was soaked for several minutes in a solution of hydroxylamine sulfate to solubilize any oxidized Mn(I1) deposited on glass surfaces. The glassware was rinsed with tap water, washed in soap and water, and placed in a 20% sulfuric acid bath for at least 1 h. Following the acid soak, glassware was rinsed thoroughly with deionized water. Standard experimental procedures, analytical techniques, and quality assurance/quality control procedures employed in the study are described elsewhere (17). Kinetic Experiments. Experiments were conducted for conditions considered representative of potable water treatment. Independent variables in the experimental matrix were pH (5.549, temperature (2-25 “C), and initial Mn(I1) concentration (0.2-1.25 mg/L). Doses of ClO, and KMnO, were 125% and 105% of the stoichiometric requirement, respectively. In this study, commonly used kinetic techniques such as “flooding” methods (i.e., one reagent present in large stoichiometric excess) were difficult to apply. It has been recommended for flooding methods that a 10-fold excess of reagent be employed (18). For oxidation by KMnO,, for example, a large excess of oxidant would result in excess 1328
Environ. Sci. Technol., Vol. 26, No. 7, 1992
MnO,(s) when the quenching agent was added at the end of the experiment. Considering that as much as 0.5 mol of Mn(I1) can be adsorbed per mole of MnOz(s) at pH 7.5 (2),a 10-fold excess of KMnO, would result in sufficient solid phase to adsorb all of the Mn(I1) initially present in the experiments. In addition, given the rapid reaction rates observed in this study, flooding an experiment with an excess of KMnO, or ClO, would have resulted in oxidation rates too rapid to measure for many test conditions. Due to these problems, an alternative kinetic test procedure was used. Kinetic studies were conducted by mixing the oxidant with the test solution, allowing the reaction to proceed for a desired duration, and then terminating the reaction by addition of a reducing agent. Solutions of Mn(II), oxidant, and the quenching agent were all adjusted to the desired pH prior to the start of each experiment. At the end of each experiment, the final pH of the suspension was measured. Drift in pH was unavoidable and varied depending on the initial pH. The data reported in this paper are the initial pH measurements, which are considered most representative of the pH during these rapid kinetic reactions. The quenched test suspension was filtered using a 30K ultrafilter and the filtrate saved for Mn(I1) analysis by atomic absorption spectroscopy. The manganese concentration of the filtrate was assumed to represent residual soluble Mn(I1). Comparisons between ultrafiltration and 0.2-pm membrane filtration showed that efficient capture of precipitated manganese using 0.2-pm filters was dependent on flocculation of the suspension prior to filtration. To ensure better capture of precipitated manganese, ultrafiltration was used in the study. Two kinetic test procedures were employed. For test conditions resulting in rapid oxidation reactions (i.e., tl C 30 s), kinetic studies were conducted using a stoppedflow-type apparatus where pressurized ultrafiltration cells served as reservoirs for the Mn(I1) and oxidant solutions. Tubing from the cells was joined at a valve, and the mixed flow then passed through a tube and into a quench solution where the oxidation reaction was terminated. The system was pressurized with nitrogen to provide a constant pressure head and hence constant volumetric flow rates. The residence time in the tubing containing the mixed flow defined the reaction time, which could be varied by changing the length of tubing or adjusting the pressure in the system. Reaction times ranging from 0.1 to 30 s could be measured using this system. For cold-temperature studies, the apparatus was immersed in a constant-temperature water bath. For experimental conditions resulting in slower kinetic reactions (i.e., tl,z > 30 s), macro- and micropipets were used to inject oxidant and reducing agent into a beaker containing the stirred test solution. The final suspension then was filtered and the filtrate analyzed for reduced Mn(I1). Phenylarsine oxide (PAO; 0.005 64 N) was selected for use as a quenching agent when KMn0, was used as an oxidant. Testing indicated that, at reaction times as low as 0.1 s, P A 0 efficiently quenched the oxidation reaction and did not resolubilize previously formed Mn02(s). For tests involving ClO,, sodium sulfite was used as a quenching agent. Although Na2S03 rapidly quenched residual ClO, (PA0 was ineffective), preliminary testing showed that Na2S03also could reduce Mn02(s)to Mn(I1). To prevent resolubilization of Mn02(s),HOC1 (lox the stoichiometric amount of Na2S03)was added 2-3 s following NazSOBinjection. For the pH conditions being
,
hydroxyl ion concentration was important for the solution-phase reaction (k,) but not for the surface oxidation term (k,) for the pH range investigated. Thus, the final rate equations used for data fitting are shown in eqs 5-8.
Table I. Rate Equations for Mn(I1) Oxidation
-
[Mno~Mn]
0-4
OX]
2 S[Mn02]
(4
[Mn021+ [Mn2+]
kz
d [Mn2+] -- -a!kl[Mnz+][Ox][OH-]" - kz[Mn2+][MnO,] dt (5)
4
a[Mn02=Mn]
+
d[Mn2+l/dt = -akl[Mn2+l[Ox] - k2[Mn2+][Mn02] (d) d[MnOzl --
dt rkl[Mn2+l[Ox] - k2[Mn2+l[Mn02] + 6k3[Mn02=Mn] [Ox]
(4
d ( M n O ~ M n/1d t = k2[Mn2+][Mn02]- ak3[Mn02=Mn][Ox] (f)
d[Ox]/dt = -pkl[Mn2+l[Ox] - pk3[Mn02=Mn][Ox]
(9)
investigated (pH
