Aquatic Geochemistry 8: 15–36, 2002. © 2002 Kluwer Academic Publishers. Printed in the Netherlands.

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Kinetics of the Abiotic Reduction of Polymeric Manganese Dioxide by Nitrite: An Anaerobic Nitrification Reaction GEORGE W. LUTHER, III
and JEANNETTE I. POPP College of Marine Studies, University of Delaware, 700 Pilottown Rd., Lewes, DE 19958, U.S.A. Received 22 November 2001; accepted 10 April 2002 Abstract. Manganese oxides are strong environmental oxidants recently found to be involved in the nitrogen cycle. Of the several possible reactions with reduced nitrogen species, the reduction of MnO2 by nitrite has only received marginal attention. Yet, this reaction might explain why nitrification can occur in the absence of O2 , observed in both sediments and water columns. We have determined the stoichiometry of this reaction, as well as the chemical kinetics and the activation parameters, using a soluble polymeric form of MnO2 . The reaction rate decreases with increasing pH and decreasing temperature. The reaction is first order in each reactant with a second order rate constant (k) = 493 M−1 min−1 at 21.5 ◦ C and pH = 5.00. The energy of activation (Ea = 9.370 kJ/mole) and the entropy of activation (S ‡ = −169.5 J/mole) show the reaction to be associative and diffusion controlled, occurring via an inner-sphere mechanism, likely with O atom transfer from MnO2 to HNO2 . The reaction is proton assisted and slows down at pH ≥ 5.5 where NO− 2 and MnO2 (unprotonated and negatively charged) become the dominant species. In natural waters and sediments where anaerobic nitrification has been observed the pH is higher than this. Thus, the thermodynamically favorable reaction will likely proceed by microbial mediation. Key words: Manganese dioxide, nitrite, nitrification, kinetics

1. Introduction Manganese oxides are considered to be important environmental oxidants for organic matter (OM) decomposition (e.g., Stone and Morgan, 1984a; Stone, 1987; Luther et al., 1999a). The biogeochmemical paradigm for OM decomposition based on the oxidant yielding the most free energy, ranks the oxidants in the follow2− ing order; O2 , NO− 3 , MnO2 , Fe2 O3 , SO4 and CO2 reduction (Froelich et al., 1979). − MnO2 reduction occurs after NO3 reduction (denitrification), however a certain degree of overlap of these oxidants has been found in (hemi)pelagic sediments (Miyajima, 1994; Luther et al., 1997; Hulth et al., 1999). An important aspect of sediment Mn redox chemistry is the possibility of its involvement in the N cycle providing alternate pathways for N2 formation (Luther et al., 1997). Traditionally, the nitrification-denitrification process is largely described
Corresponding author; [email protected], (302) 645-4208

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GEORGE W. LUTHER, III AND JEANNETTE I. POPP

by bacterially mediated organic matter decomposition (Kaplan, 1983; Libes, 1992), which does not exclude the possibility of alternate pathways for N2 formation. Observations of deep-sea sediment pore waters (Schulz et al., 1994) and the Black Sea suboxic water column (Murray et al., 1995) suggest that N2 formation may 2+ also result from the reduction of NO− 3 by dissolved Mn . Luther et al. (1997) showed that N2 forms from MnO2 and ammonia. Sorensen et al. (1987) suggested chemodenitrification by Mn as a possible explanation for discrepancies found between measured rates of denitrification and estimated NO− 3 fluxes, such as those described in previous field studies. Luther et al. (1997) performed thermodynamic calculations that show MnO2 can react with NH3 to produce N2 and/or NO− 3 , although the production of N2 is more thermodynamically favorable. Their pore water data shows that NO− 3 forms in environments where O2 is not detected. Mortimer et al. (1998, 1999) and Hulth et al. (1999) also mention that continental margin sediments rich in organic matter and Mn, the N and Mn cycles may be closely coupled. Based on incubation experiments, Hulth et al. (1999) suggested that Mn oxides are the only known sedimentary constituents present with high enough oxidation potential and sufficient abundance to sustain reported nitrification rates. The primary conclusions from − these experiments are that NO− 3 and NO2 can be produced in the absence of O2 and that reduced-N oxidation takes place simultaneously with Mn oxide reduction. Hulth et al. (1999) stated that the introduction of Mn oxides from oxidized surface sediment into underlying anoxic regions during biogenic or physical mixing results − in anoxic nitrification and net production of NO− 3 and NO2 . Anoxic nitrification rates are proportional to the quantity of Mn oxide available in the solid phase. Vandenabeele et al. (1995) reported that nitrifying (ammonium- and nitriteoxidizing) bacteria play a role in the biological removal of Mn. In their batch culture experiments of the bacterium Nitrosomonas, NO− 2 reduced the manganese dioxides that were previously formed biologically. This indicates that autotrophic NH+ 4 oxidizers indirectly cause the reduction of MnO2 through the production of NO− 2. − Bartlett (1981) hypothesized that Mn oxides can convert NO− 2 to NO3 in soils and suggested that the importance of reactive Mn oxides as electron acceptors for reduced species in aerobic field soils may have been overlooked. Although thermo− dynamically possible, abiotic oxidation of NO− 2 to NO3 by atmospheric O2 does not occur for kinetic reasons in aseptic systems at near-neutral pH. However, this reaction occurs rapidly in most aerobic soils via microbial mediation. The oxida− 3+ and Mn4+ oxides is a tion of NO− 2 to NO3 accompanied by the reduction of Mn thermodynamically spontaneous reaction. Bartlett (1981) suggested the following reaction (Equation (1)): + 2+ + NO− NO− 2 + MnO2 + 2H → Mn 3 + H2 O.

(1)

This reaction, which consumes protons, occurs without O2 , and all of the NO− 2 − appears to be oxidized to NO− without any loss of N. Oxidation of NO by 3 2

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Nitrobacter differs from this reaction by requiring atmospheric O2 and by not consuming protons (Kaplan, 1983). Much of the research involving MnO2 reduction has centered on solid phases (Bartlett, 1981; Stone and Morgan, 1984b; Stone, 1987; McArdell et al., 1998), which have a broad range of reactivities (Nelson et al, 2002). Little has been done with soluble or polymeric forms of Mn3+,4+ (Kostka et al., 1995; Perez-Benito et al., 1996; Luther et al., 1999), which Kostka et al. (1995) have shown to be more reactive than solid forms. Soluble Mn3+,4+ compounds could well be important at the interfaces of oxic, suboxic and anoxic zones, where they can be formed from Mn2+ oxidation, and in low salinity waters (Luther et al., 1999a; Luther et al., 1997; Kostka et al., 1995; Luther et al., 1991). Traditionally Mn has been determined by filtration, assuming that dissolved Mn is Mn2+ and that particulate Mn is composed mainly of Mn4+ (Xyla et al., 1992; Wilczak et al., 1993; Kostka et al., 1995). Soluble and polymeric forms of Mn3+ and Mn4+ have been shown to pass through 0.45 µm filters (Wilczak et al., 1993; Kostka et al., 1995) disproving the assumption that dissolved Mn is only Mn2+ . Clearly, these measurements have oversimplified Mn speciation, thereby affecting our understanding of Mn chemistry and its interactions with the cycles of other elements. Polymeric MnO2 is the most reactive phase of Mn4+ formed from Mn2+ oxidation. Stone and Morgan (1984b) and Laha and Luthy (1990) state that a small unspecified amount of polymeric MnO2 exists in the natural environment. The existence of soluble/polymeric Mn3+ (e.g., Kostka et al., 1995) and Mn4+ (e.g., Perez-Benito et al., 1989; Luther et al., 1999) in laboratory solutions has been documented. The emphasis of this research is to show MnO2 as a strong oxidant being re2+ and NO− duced abiotically by NO− 2 , forming Mn 3 . A polymeric form of MnO2 , which is measurable by UV-VIS spectroscopy, was used to test the stoichiometry presented by Bartlett (1981; Eq. 1), and to determine the empirical rate law and the activation parameters (activation energy {Ea}, enthalpy of activation {H ‡ }, entropy of activation {S ‡ }). Nitrite is an intermediate species both in denitrification and in nitrification. In natural environments, the reduction of solid MnO2 {MnO2 (s)} by nitrite has been demonstrated (Bartlett, 1981; Vandenabeele et al., 1995). In this study, the rate of reduction of polymeric MnO2 was measured as a function of pH (4.5–6) and temperature (2.0–21.5 ◦ C) in order to gain information on the reaction mechanism. 2. Methods and Materials Reagent grade materials (A.C.S.) from Fisher Sci. were used in all experiments. Polymeric MnO2 was made with a 5% excess of sodium thiosulfate following the stoichiometirc titration reaction (Equation (2)) as described by Perez-Benito et al. (1989), who determined that the oxidation state of the Mn product is IV. 2− 2− + 8MnO− 4 + 3S2 O3 + 2H → 8MnO2 + 6SO4 + H2 O.

(2)

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GEORGE W. LUTHER, III AND JEANNETTE I. POPP

This MnO2 solution remains stable, dark brown, and transparent for several months, and the polymeric particles are roughly spherical with a radius of approximately 500 Å or less. Also, this MnO2 species can pass through a 0.45 µm filter (Wilczak et al., 1993). The polymeric species has a negative electrostatic charge, which is responsible for its stability in solution. The addition of cations to the negatively charged MnO2 species will cause precipitation, which increases proportionally with the ionic radius and the oxidation state of the cation (Perez-Benito et al., 1989; Luther et al., 1999). The experimental procedure is similar to that described in Luther et al. (1999a) for the reduction of MnO2 by oxalate. To measure MnO2 , which has a broad absorption in the UV-VIS region (Perez-Benito et al, 1996), a 5 cm fiber optic dip probe with a 450 nm filter attached to a Brinkman model PC800 colorimeter was employed. No other chemical species [nitrite, nitrate and Mn(II)] has an absorption peak in this region. A Metrohm model 716 titrino titrator in pH stat mode was used to maintain the experiments at a constant pH of 4.5, 5.0, 5.5, and 6.0 (±0.03). For each of these values of pH, KNO2 was added to the reaction vessel containing MnO2 and the solution was adjusted to the desired pH by adding KOH. For pH 5.5 and 6.0, the pH of the KNO2 solution was also adjusted with KOH before addition to the reaction vessel. Potassium nitrite was then added to the MnO2 solution and pH maintained constant with the aid of the pH stat system. The initial KNO2 concentrations were five to twenty times larger than the initial MnO2 concentrations. The volume change over the entire reaction was less than 0.25% based on the acid added to maintain constant pH. − An HPLC system was used for the analysis of NO− 2 and NO3 . The HPLC was a SSI model 200 LC pump with a SSI model 210 Guardian. A SSI 500 Detector variable UV/VIS wavelength set at 210 nm was used as the detector. The sample loop volume was 100 µL. The anion column was a 4.6 mm ID × 100 mm Anion/S10 micron WESCAN purchased from Alltech
. The eluent was 1.5 mM H2 SO4 and the flow rate was 2 mL min−1 . Peak areas were determined with a Hewlett Packard HP3396A integrator. The NO− 2 peak was detected at 1.5 minutes and the peak at 8.5 minutes. The detection limit for this method is about 1 µM. Use NO− 3 of a Wescan anion/R 150 × 46 mm IC column (Alltech
) permits separation of these nitrogen species at nanomolar levels in basic perchlorate media (Rozan and Luther, 2002). Mn(II) was measured by the voltammetric method of Luther et al. (1994). Precision on triplicate analyses for all these methods is typically better than 5% and frequently better than 2%. With the measurements described above, the concentration of each manganese and nitrogen chemical species in Equation (1) can be determined. The total volume for each reaction was 75 mL and the concentration for polymeric MnO2 and KNO2 varied between 20 and 40 µM and 200 and 400 µM, respectively. Using a Fisher model 9101 Chiller, the temperature of the reaction vessel was kept constant at temperatures (2.00 ◦ C, 11.50 ◦ C, 21.50 ◦ C) normally found for sediments.

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2.1. KINETIC CALCULATIONS The rate of disappearance (Equation (4)) of a reactant (polymeric MnO2 ) was monitored to determine reaction kinetics (Connors, 1990). The overall rate equation for the oxidation of nitrite by MnO2 can be expressed by Equation (3) rate = −d[MnO2 ]/dt = k[MnO2 ]x [NO− 2 ]T , y

(3)

where k is the rate constant, [ ] symbolizes concentration for each species and the exponents x and y are the order of reaction for each species that must be − experimentally determined. [NO− 2 ]T is the sum of the nitrite species (NO2 and HNO2 ). When the experiments are performed with a constant excess of nitrite and varying the concentration of MnO2 , the rate equation reduces to rate = d[MnO2 ]/dt = kobs [MnO2 ]x ,

(4)

where kobs = k[NO− 2 ]T . y

(5)

The order x with respect to [MnO2 ] was determined by fitting the data to various rate equations. Plots of ln[MnO2 ] vs. time (e.g., Figure 1) produced straight lines indicating that the reaction order is first order in MnO2 . The order y with respect to [NO− 2 ]T in Equation (3) was determined by varying nitrite concentration at nearly constant [MnO2 ] and evaluating the rates for 50% reaction completion (e.g., Stone and Morgan, 1984a) at pH = 4.5. The order y is evaluated from separate reactions via the following expression (6) since the reaction is first order in polymeric MnO2 . − y (rate2 /rate1 ) = {[NO− 2 ]2 /[NO2 ]1 } {[MnO2 ]2 /[MnO2 ]1 }.

(6)

2.2. ACTIVATION PARAMETERS FROM TEMPERATURE DEPENDENCE Activation energy is the minimum energy required to initiate a chemical reaction (Connors, 1990). The relationship between the rate constant (k) of an elementary reaction and the absolute temperature (T ) is given by the Arrhenius Equation (7): k = Ae−Ea /RT ,

(7)

where A is the pre-exponential factor (units of k), R is the gas constant (8.314 J K−1 mol−1 ), T is the temperature (K) and Ea is the activation energy (kJ mol−1 ). Using an Arrhenius plot (ln k versus 1/T ), A and Ea can be calculated using Equation (8): ln k = ln A − Ea /RT .

(8)

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GEORGE W. LUTHER, III AND JEANNETTE I. POPP

If Ea < 25 kJ/mole, the reaction is considered diffusion controlled (Connors, 1990). To determine whether the reaction mechanism is associative or dissociative, the entropy of activation is needed (Atwood, 1985), which can be evaluated from an Eyring plot [ln k/T ) versus 1/T ] based on Equation (9): ln(k/T ) = (−H ‡ /RT ) + ln(k/ h) + (S ‡ /R),

(9)

where H ‡ is the enthalpy of activation, S ‡ is the entropy of activation, k is the Boltzmann constant (1.381×10−23 JK−1 ) and h is Planck’s constant (6.626×10−34 Js) , the slope (−H ‡ /RT ) provides the enthalpy of activation and the intercept {ln (k/h) + (S ‡ /R)} yields the entropy of activation. A large negative S ‡ indicates an associative reaction and a large positive S ‡ indicates a dissociative reaction (Atwood, 1985). 2.3. SEMI - EMPIRICAL MOLECULAR ORBITAL CALCULATIONS The energies of the molecular orbitals for HNO2 and NO− 2 were calculated at the AM1 level using the program HYPERCHEM from Hypercube, Inc. The AM1 method is generally the most accurate computational method in this program. The calculated values for the energies of the highest occupied and lowest unoccupied molecular orbitals agree with available experimental data (e.g., experimental HOMO energy of NO− 2 is given as −2.5 eV and the calculated value is −2.56 eV). 3. Results and Discussion 3.1. STOICHIOMETRY Bartlett (1981) showed in laboratory studies that MnO2 could oxidize NO− 2 , and proposed Equation (1) for this reaction. However, he did not actually verify it, because he used buffers to maintain constant pH. Our experiments confirm the stoichiometry of the reduction of MnO2 by nitrite using a polymeric form of MnO2 under pH stat conditions because all the reactants and products could be monitored to confirm Equation (1). Table I shows the results from these experiments. For pH values of 4.5, 5.0 and 5.5, the reaction is as written. At pH 5.5 and 6.0 the reaction does not go to completion even after an hour of reaction time. A one-to-one rela− tionship is shown with MnO2 , Mn2+ , NO− 2 , and NO3 . A two-to-one consumption − of protons for each MnO2 and NO2 reacted is also found. Thus, Equation (1) is verified. 3.2. REACTION RATE AND RATE LAW Using pseudo first order conditions with excess nitrite, the rate of disappearance for MnO2 was determined at various pH values (4.5–6.0). Figures 1–4 show that

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Table I. Reactants (∗ ) and species (∗∗ ) experimentally determined from the reduction of MnO2 by nitrite. Standard deviations based on triplicates. The negative H+ concentrations indicate its consumption.

1 2 3 4

pH

Mn(IV)∗ (µM)

∗ NO− 2 (µM)

H+∗ (µM)

∗∗ NO− 3 (µM)

Mn(II)∗∗ (µM)

4.5 5.0 5.5# 6.0#

20.2 ± 0.1 34.8 ± 0.2 39.9 ± 1.5 19.9 ± 1.7

232.8 ± 9.2 193.2 ± 3.6 163.4 ± 17.6 187.7 ± 18.4

−51.5 ± 1.6 −69.3 ± 0.0 −26.0 ± 3.7 −6.5 ± 0.1

25.0 ± 0.6 32.8 ± 0.1 15.8 ± 1.6 4.7 ± 0.8

22.92 ± 1.4 34.9 ± 1.3 15.06 ± 0.6 9.6 ± 0.5

# indicates total time of experiment was 60 minutes.

the reaction obeys first order kinetics for MnO2 at a pH of 4.5 and nominal concentration ratios ([MnO2 ]/[NO− 2 ]T ) of 20/200, 20/400, 40/400 and 40/200. Thus, the order with respect to MnO2 is one. Evaluation of the order for nitrite gave values of 0.93 and 1.09 for the data in Table II using Equation (6) for the 20/200 and 20/400 reactions (Figures 1 and 2) and for the 40/200 and 40/400 reactions (Figures 3 and 4), respectively. Thus the reaction is first order in nitrite and the rate law is given in Equation (10). rate = −d[MnO2 ]/dt = k[MnO2 ][NO− 2 ]T .

(10)

3.3. P H DEPENDENCE Figure 5A shows the reduction of MnO2 for a [MnO2 ]/[NO− 2 ]T ratio of 20/200 over pH. The time of completion for the consumption of MnO2 increases substantially as pH increases. The experiments for pH values of 5.5 and 6.0 were terminated 60 minutes after the start of the reaction. This provided enough information to determine the initial reaction rates (10% completion, Table III). Figures 5A and B also show that there is a change in the slope of MnO2 disappearance with time for pH values of 5.5 and 6.0. The initial rate and kobs data were determined for the first 10% of the reactions at a pH of 5.5 and 6.0 and the initial rate data are plotted versus pH in Figure 6A. For these reactions at higher pH, there is also a break in the ln[MnO2 ] versus time plots so that two kobs can be calculated for the reaction (Figure 5B). The first or initial kobs (kobsi ) occurred within the first few minutes and was used to determine the initial reaction rate (Table III). The second kobs (kobst ) indicates a slowing in the reaction rate and may be due to a change in the MnO2 particle size with pH and/or Mn2+ adsorption as the polymeric MnO2 has a more negative charge with pH. The data for pH 5.5 and 6.0 were not used to calculate the order for the reaction. Figure 6A also shows that HNO2, and not NO− 2 , is a primary reactant and this is why the reaction rate decreases as pH increases (as shown in Table III). The effect

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GEORGE W. LUTHER, III AND JEANNETTE I. POPP

Figure 1. (A) The reduction of MnO2 (20.1 µM) by NO− 2 (232.8 µM) as a function of time at a pH of 4.5. (B) ln[MnO2 ] versus time plots showing pseudo first order conditions were achieved for the first 50% of the reaction; kobs is determined from the slope.

KINETICS OF THE ABIOTIC REDUCTION OF POLYMERIC MANGANESE DIOXIDE BY NITRITE

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Figure 2. A) The reduction of MnO2 (21.0 µM) by NO− 2 (358.8 µM) as a function of time at a pH of 4.5. (B) ln[MnO2 ] versus time plots showing pseudo first order conditions were achieved for the first 50% of the reaction; kobs is determined from the slope.

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GEORGE W. LUTHER, III AND JEANNETTE I. POPP

Figure 3. (A) The reduction of MnO2 (37.3 µM) by NO− 2 (194.7 µM) as a function of time at a pH of 4.5. (B) ln[MnO2 ] versus time plots showing pseudo first order conditions were achieved for the first 50% of the reaction; kobs is determined from the slope.

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Figure 4. (A) The reduction of MnO2 (34.4 µM) by NO− 2 (358.8 µM) as a function of time at a pH of 4.5. (B) ln[MnO2 ] versus time plots showing pseudo first order conditions were achieved for the first 50% of the reaction; kobs is determined from the slope.

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GEORGE W. LUTHER, III AND JEANNETTE I. POPP

Figure 5. (A) The reduction of MnO2 for a concentration ratio of 20/200 ([MnO2 ]/[NO− 2 ]) as a function of time for the pH range of 4.5–6.0. (B) The reduction of MnO2 as a function of time, and the two kobs for a pH of 6.0 at a concentration ratio of 40/200 ([MnO2 ]/[NO− 2 ]).

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Table II. MnO2 and nitrite concentrations for each concentration ratio and pH range used to ascertain rate expressions. Rate data are for the first 50% of the reaction pH

Nominal ratio

−(Rate) (µM/min)

NO− 2 (µM)

MnO2 (µM)

4.5

40/200

5.228 5.228 7.775 8.500 5.155 4.690 4.745 8.007 8.360 8.013

193.0 196.4 n/a 358.8 223.2 241.7 233.5 358.8 358.8 358.8 190.7 195.8 180.1 198.0 186.3 177.8 205.8 144.2 178.6 167.5 194.6 168.9 210.2 174.9 173.6 189.6

37.1 37.3 34.4 34.4 20.3 20.0 20.2 21.0 21.0 21.1 35.0 34.7 34.7 21.1 21.0 19.3 21.5 38.16 40.84 40.8 20.7 20.4 17.4 34.65 33.85 34.35

40/400 20/200

20/400

5.0

40/200

20/200 5.5

20/200 40/200

6.0

20/200

40/200

Note: n/a indicates data not available due to an instrument malfunction.

of pH on the oxidation of HNO2 can be shown via its fractionation, which is related through the following equations. Total nitrite ([NO− 2 ]T ) is expressed as Equation (11).

− [NO− 2 ]T = HNO2 + NO2 .

(11)

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GEORGE W. LUTHER, III AND JEANNETTE I. POPP

Table III. Calculated reaction rates for the first 10% of each ratio for each pH Ratio [MnO2 ]/[NO− 2 ]T

pH

−(Initial Rate) (µM min−1 )

20/200

4.5 5.0 5.5 6.0 4.5 5.0 5.5 6.0 4.5

8.67 2.77 0.978 0.412 10.8 3.11 0.431 0.303 9.54

40/200

40/400

The thermodynamic acid dissociation constant (K1 ; Equation (11)) for the ionization of HNO2 is given as 5.1 × 10−4 (pKa = 3.25; Skoog et al., 1992) K1 = ([H+ ] ∗ [NO− 2 ])/[HNO2 ].

(12)

Manipulating Equation (11) to solve for [NO− 2 ] and then substituting into Equation (12) yields [NO− 2 ]T = [HNO2 ] ∗ (1 + {K1 /[H+]}).

(13)

Rearranging Equation (13) to solve for [HNO2 ] gives [HNO2 ] = [NO− 2 ]T /(1 + {K1 /[H+]}).

(14)

The fraction of HNO2 (αHNO2 ) composing [NO2 −]T can be calculated by rearranging Equation (14) to Equation (15). + [HNO2 ]/[NO− 2 ]T = αHNO2 = 1/(1+ {K1 /[H ]}).

(15)

The nitrite fraction (αNO−2 ), can be determined by Equation (16). 1 − αHNO−2 = αNO−2 .

(16)

Thus, αNO−2 increases as the pH increases (Figure 6) and HNO2 is the primary oxidant for MnO2 . Substituting Equation (13) into Equation (10), the rate law can be modified to include nitrite speciation to give Equation (17). rate = k([HNO2 ] ∗ {1 + (K1 /[H+ ])}) [MnO2 ].

(17)

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Figure 6. (A) The effect of pH on kobs of MnO2 reduction by NO− 2 (first 10% of the reaction) at 21.5 ◦ C. As the pH increases, the reduction rate decreases and the αNO− increases. (B) 2 Fractionation factors for equations 19–22 versus pH. (C) Insert expanded to show data for reactions F1 (Equation (19)) and F3 (Equation (22)).

3.4. TEMPERATURE DEPENDENCE All of the experiments above were run at a temperature of 21.5 ◦ C. In order to easily determine activation parameters, all experiments were run at a pH of 5.0 and a concentration ratio of 40/200 for the temperatures of 2.0 ◦ C, 11.5 ◦ C, and 21.5 ◦ C, which are representative of most marine and freshwater sediments. Figure

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GEORGE W. LUTHER, III AND JEANNETTE I. POPP

Table IV. The reaction rates and second order rate constant, k, for the temperature range of 275.16–304.66 K (2.00–21.50 ◦ C) T (K)

Rate (M min−1 )

k (M−1 min−1 )

275.16 284.66 294.66

2.47 × 10−6 2.89 × 10−6 3.33 × 10−6

376 440 493

7A shows the reduction of MnO2 at these temperatures. As Table IV shows, the rate and k increased with higher temperature. Figure 7B shows the Arrhenius and Eyring plots, of which the slope and intercept are used to calculate Ea and S ‡ , respectively. Using Equation (8) and the slope from the Arrhenius plot, the calculated Ea is 9.370 kJ/mole, and the reaction is diffusion controlled if Ea is less than 25 kJ/mole (Connors, 1990). Therefore, this reaction is diffusion controlled. Using Equation (9) and the slope and intercept from the Eyring plot (Figure 7B), H ‡ is 6.989 kJ/mole and S ‡ is −169.5 J/(mole K). When the S ‡ is a large negative number, the reaction mechanism is associative (Atwood, 1985). Thus, an inner sphere mechanism, with both reactants binding in the transition state occurs. These data are consistent with known manganese chemistry.

3.5. MECHANISM The empirical rate law is given in Equation (10) and shows that the reaction is first order in both reactants. At 21.5 ◦ C, k = 493 (±4.9) M−1 min−1 . This value of k is similar to that of 436 M−1 min−1 found by Yao and Miller (1993) for the reduction of solid phase MnO2 by sulfide. With increasing pH, the reaction rate for Equation (1) decreases (Figure 6A). There is a sharp decrease in the rate between 4.5 and 5.0 as the reaction slows and levels off at a pH of 5.5. 3 ) in octahedral symmetry. Thus, it does Mn4+ is an inert metal cation (d 3 , t2g not undergo ligand exchange easily because inertness retards reactivity (Luther 1990; Shriver et al., 1994). We discuss two mechanistic possibilities for redox transformation in this reaction. First, nitrite donates two electrons to Mn4+ and accepts an O2− from MnO2 to form nitrate. The second process is loss of an O atom from MnO2 directly to nitrite to form nitrate and retention of two electrons from the Mn–O bond in MnO2 . We discuss these possibilities by looking at the possible precursor complexes in the transition state. The possible chemical species for nitrite are HNO2 and NO− 2 and for MnO2 the species are Mn–OH and Mn– − O . The polymeric Mn species are similar to that for the mineral because the first

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Figure 7. (A) Reduction of MnO2 as a function of temperature (2.0–21.5 C) at a pH of 5.0 for a concentration ratio of 40/200 ([MnO2 ]/[NO− 2 ]). (B) the Arrhenius and Eyring plots for the data in (A).

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GEORGE W. LUTHER, III AND JEANNETTE I. POPP

formed solid product must be similar to the species in solution prior to precipitation – the Ostwald step rule (Casey and Morse, 1988; Luther et al., 1999b). Mn–OH and Mn–O– are related by the second conditional acidity constant (Equation (18); Yao and Millero, 1993) for MnO2 by S >Mn–OH → MnO− + H+∗ Ka2 = 10−4.6 .

(18)

The acidity constant shows that MnO2 will be more protonated as pH increases than nitrite, which has a pKa = 3.25. The possibilities for reaction in the transition state are given in Equations (19)– (22), and the product of their fractions at a given pH is F (αMnO2 × αNO2), the fractionation factor (Figure 6B). The decrease in reaction rate with increasing pH is a function of Equation (22), which shows that two anionic species should repel each other and not react in the transition state. Equation (19) shows that two neutral species come together whereas Equations (20) and (21) show that a neutral species and an anion come together in the transition state. Thus, these chemical species in these reactions (19)–(21) should permit the oxidation of nitrite by MnO2 . >OMn–OH + HNO2 (two neutral species) → Mn-ONO2 + 2H+ ,

(19; F1)

>OMn–OH + NO− 2 (a neutral species and an anion) →Mn–ONO2 + H+ ,

(20; F2)

>OMn–O– + HNO2 (an anion and a neutral species) →Mn–ONO2 + H+ ,

(121; F3)

>OMn–O– + NO− 2 (two anion species) → Mn–ONO.

(22; F4)

Figure 6A shows that the initial rate decreases by a factor of 6 from pH 4.5 to 5 for the reactions with the stoichiometric ratios of 20/200 and 40/200. Figure 6B shows that reactions 20 and 22 do not follow that trend whereas Figure 6C (insert in Figure 6B) shows that only reaction 19 follows the trend exactly. Thus, the predominant chemical species involved with the oxidation of nitrite by MnO2 are HNO2 and Mn–OH (Equation (19)). These chemical species account for the stoichiometry of the reaction as well as the reaction kinetics. The loss of two protons in the reaction (Equation (1)) indicates that another O atom is lost from MnO2 later in the reaction and binds the protons to form water. Reduction of MnO2 requires the gain of two electrons. However, HNO2 is an acid, and semi-empirical molecular orbital calculations (AM1 level) indicate that it is a poor electron donor [the highest occupied molecular orbital (HOMO) energy is −11.24 eV; the more negative the HOMO energy, the less likely electron donation occurs]. Nitrite is an excellent donor (HOMO energy is −2.56 ev), but it is not the direct reactant based on our data. Because the MnO2 orbitals (LUMO; eg∗ ) that accept electrons are totally unoccupied and on the bond axis, loss of an O atom (an excellent electron acceptor)

KINETICS OF THE ABIOTIC REDUCTION OF POLYMERIC MANGANESE DIOXIDE BY NITRITE

33

Figure 8. Possible inner sphere complex formation of the reactants in the transition state and break up to form products (NO− 3 and Mn–OH2 ) with O atom transfer and electron transfer to 2+ Mn. As Mn forms and is released into solution, water molecules will satisfy its coordination geometry. Arrows indicate the direction of electron and proton movement in the transition state, and brackets indicate bond breaking. The reaction is pictured as a concerted reaction and shows the affect of acidity (protons) in the intermediate.

directly to HNO2 would result in direct reduction of MnO2 to Mn2+ . O atom transfer is well known in Mo and W enzyme systems where Mo6+ –O transfers an O atom to nitrite (sulfite) to form nitrate (sulfate) and Mo4+ (Cotton et al., 1999). Thus, unpairing and donation of electrons from the HNO2 to Mn4+ is not required during Mn4+ reduction (the first possibility mentioned above). For O atom transfer, an inner sphere complex where nitrogen in HNO2 donates electrons to one O atom as a proton from HNO2 is donated to the other O atom in MnO2 (Mn– OH species) is possible (Figure 8). The O atoms in MnO2 push electrons to Mn in charge transfer bands to satisfy the +4 oxidation state in Mn (Luther, 1991). Protonation of one of the O atoms weakens that Mn–O bond and enhances the electron transfer to Mn from the other nonprotonated O atom, which results in its transfer as an atom to N in nitrite. Thus reduction of Mn4+ occurs as O atom transfer to HNO2 occurs leaving the electrons in the original Mn–O bond with the Mn. The result is oxidation of nitrite to nitrate.

4. Geochemical significance During diagenesis, manganese oxides are important and recently have been suggested as an alternate pathway in the N-cycle (Luther et al., 1997). One such pathway participating in alternative pathways in the N-cycle in nitrification is the oxidation of NO− 2 . Hulth et al. (1999) suggested that anoxic nitrification rates in their experiments were proportional to the quantity of Mn oxide available. Vandenabeele et al. + (1995) state that NO− 2 produced by autotrophic NH4 oxidizers was instrumental in Mn oxide reduction. Thus reaction (Equation (1)) appears to be microbially mediated at higher pH in nature since our results show that the abiotic reaction

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slows as pH increases. Bartlett (1981) also determined that MnO2 (s) in soils could oxidize NO− 2 abiotically. 5. Conclusions The empirical rate law (Equation (10)) for the reduction of MnO2 by nitrite over the pH range of 4.5 to 6.0 (determined using soluble polymeric MnO2 rather than solid material as in previous studies e.g., Xyla et al., 1992) is first order in each reactant. The reaction is diffusion controlled and associative. This information predicts that the entering HNO2 or NO− 2 group will likely form a bond with the MnO2 before the Mn–O bond cleaves. The reaction rate for the reduction of MnO2 by nitrite slows with increasing pH S (pHzpc or pH at zero-point of charge) of MnO2 , and also because due to the ∗ Ka2 of the increase in the NO− 2 fraction at higher pH values (Equation (22)). As the pH increases beyond the pHzpc , the MnO2 becomes negatively charged and NO− 2 ions become the predominant species in solution. Thus repulsion between the two negatively charged reactants makes the reaction kinetically inhibited at higher pH. Polymeric MnO2 is a strong oxidant in the environment, easily transported because of its soluble nature. The use of soluble and polymeric MnO2 allows for direct analysis of MnO2 as a reactant in experiments. The distributions of O2 and NO− 3 are not consistent with nitrification exclusively in the presence of oxygen (Hulth et al. 1999; Luther et al., 1997; Murray et al., 1995). The oxidation of nitrite by MnO2 to form nitrate should play an important role in the anaerobic production of NO− 3 in the environment, which explains in part why anaerobic nitrification can occur, especially at lower pH. However, in the environment, microbial mediation can be expected to be more important than the abiotic reduction of MnO2 at higher pH values, such as the pH of seawater (8.0) and of porewaters (6.5–8.0).

Acknowledgements This work was supported by grants from NSF (OCE-0096365) and NOAA Office of Sea Grant (NA16RG0162-03). We would also like to thank Dave Ruppel for his help with the experimental set up. This manuscript improved because of the comments made by B. Sundby and three anonymous reviewers.

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