environmental science and technology

Influence of pH on the Reductive Transformation of Birnessite byAqueous Mn(II)Joshua P. Lefkowitz,† Ashaki A. Rouff,‡ and Evert J. Elzinga†,*†Department of Earth & Environmental Sciences, Rutgers University, Newark, New Jersey 07102, United States‡School of Earth and Environmental Sciences, Queens College, Queens, New York 11367, United States

*S Supporting Information

ABSTRACT: We investigated the effect of pH (5.5−8.5) on themineralogical transformation of hexagonal birnessite induced byreaction with aqueous Mn(II) (50−2200 μM), using batch sorptionexperiments, X-ray diffraction analyses, X-ray absorption and infraredspectroscopic measurements. Samples reacted at pH < 7.0 exhibiteddisrupted stacking of birnessite sheets, but no mineralogicaltransformation products were observed. At pH 7.0 and 7.5, reactionwith Mn(II) under anoxic conditions caused reductive transformationof birnessite into manganite (γ-MnOOH), whereas at pH 8.0 and 8.5,conversion into hausmannite (Mn3O4) occurred. Feitknechtite (β-MnOOH) is a major transformation product at low Mn(II) inputs atpH 7.0−8.5, and represents a metastable reaction intermediate that isconverted into manganite and possibly hausmannite during furtherreaction with Mn(II). Thermodynamic calculations suggest thatconversion into hausmannite at alkaline pH reflects a kinetic effect where rapid hausmannite precipitation prevents formation ofthermodynamically more favorable manganite. In oxic systems, feitknechtite formation due to surface catalyzed oxidation ofMn(II) by O2 increases Mn(II) removal relative to anoxic systems at pH ≥ 7. The results of this study suggest that aqueousMn(II) is an important control on the mineralogy and reactivity of natural Mn-oxides, particularly in aqueous geochemicalenvironments with neutral to alkaline pH values.

■ INTRODUCTION

Hexagonal birnessite type minerals have garnered much interestdue to their natural ubiquity, unique structural and reactiveproperties, and for their potential impact on the fate andtransport of a diverse range of chemical species in theenvironment.1−5 These phyllomanganates dominate Mnmineralogy in a variety of geochemical environments withreported occurrences in soils, marine Mn−Fe nodules, anddesert varnishes.1 Hexagonal birnessites are the main productderived from the biologically catalyzed oxidation of Mn-(II),2,4,6−8 and are structurally characterized by mineral sheetswith hexagonal layer symmetry and a significant proportion ofreactive anionic vacancy sites.9 They are further noted forsignificant Mn(IV) content, with reported average Mnoxidation states of 3.7−4.0,2,8,9 as well as for large specificsurface area,10 high redox potential,11 and low point of zerocharge.12 Their physical and chemical characteristics providethese minerals with high reactivity with respect to both sorptionand redox reactions, explaining their critical role in determiningthe speciation and distribution of trace element and pollutantspecies in the environment.2,13−19

Recent work has shown that hexagonal birnessites are subjectto structural and mineralogical changes during reaction withaqueous Mn(II), which suggests that the dissolved Mn(II)concentration represents an important control on the structure

and reactivity of Mn-oxides in aqueous geochemical environ-ments. Early work on interactions between aqueous Mn(II) andhexagonal birnessite focused on the adsorption of the aqueousmetal by the mineral substrate.20−22 Tu et al.23 andMandernack et al.24 demonstrated that under oxic conditions,reaction of Mn(II) with hexagonal birnessite yielded a variety ofMn-oxide mineral products dependent on pH, with formationof nsutite (γ-Mn(IV,III)(O,OH)2) and ramsdellite (Mn(IV)-O2) observed at pH 2.4, cryptomelane (K1.3−1.5Mn(IV,III)8O16) and groutite (α-Mn(III)OOH) at pH 4.0 and 6.0,respectively, and feitknechtite (β-Mn(III)OOH) and manganite(γ-Mn(III)OOH) at pH > 7. Bargar et al.25 observed theformation of feitknechtite during reaction of aqueous Mn(II)with biogenic hexagonal birnessite in oxic systems at circum-neutral pH, and attributed this to electron exchange betweenadsorbed Mn(II) and structural Mn(IV) yielding Mn(III). Arecent study by Elzinga26 demonstrated that reaction ofaqueous Mn(II) with hexagonal birnessite under anoxicconditions at pH 7.5 leads to bulk transformation of themineral into manganite through a reductive transformation

Received: May 10, 2013Revised: July 18, 2013Accepted: July 22, 2013Published: July 22, 2013

Article

pubs.acs.org/est

© 2013 American Chemical Society 10364 dx.doi.org/10.1021/es402108d | Environ. Sci. Technol. 2013, 47, 10364−10371

pubs.acs.org/est

process whereby the substrate is initially converted intofeitknechtite (β-MnOOH) through interfacial electron ex-change between adsorbed Mn(II) and structural Mn(IV),followed by Mn(II)-catalyzed conversion of β-MnOOH intothe more stable manganite (γ-MnOOH) phase. Other studieshave shown the importance of Mn(II) in inducing structuralchanges of hexagonal birnessite substrates through interfacialelectron exchange reactions with bulk Mn(IV) as well.8,27−31

The current study focused on the influence of pH on thereductive transformation of hexagonal birnessite by aqueousMn(II). Solution pH affects both the extent and mechanisms ofmetal adsorption onto mineral surfaces,32−35 and may thus wellinfluence the interaction of Mn(II) with the hexagonalbirnessite surface and resulting impacts on Mn-oxide structureand mineralogy. Here, we investigated Mn(II) reactivity withhexagonal birnessite in the pH range 5.5−8.5 with acombination of batch experiments and spectroscopic measure-ments to assess sorption trends and Mn-oxide reactionproducts under oxic and anoxic conditions. Our results indicatethat pH has a major impact on the pathways and products ofthe Mn(II)-induced conversion of hexagonal birnessite intolower valence Mn-oxide phases.

■ MATERIALS AND METHODSMn-Oxide Substrates. Preparation of hexagonal birnessite

(nominally MnO2; approximate full chemical formulaKMn5O10.5 (see Supporting Information (SI)), and referencefeitknechtite (β-MnOOH), manganite (γ-MnOOH) andhausmannite (Mn3O4) is described in the SI.Mn(II)−Birnessite Isotherm Experiments. Birnessite-

Mn(II) sorption isotherm experiments were performed at pHvalues in the range 5.5−8.5. Experiments were conductedmostly under anoxic conditions, using protocols to exclude O2described in the SI. Anoxic aqueous suspensions of birnessitewere prepared in 0.1 M NaCl and maintained at the desired pHusing 20 mM of 2-(N-morpholino)ethanesulfonic acid (MES;pH 5.5−6.5), 4-(2-hydroxyethyl)-1-piperazineethanesulfonicacid (HEPES; pH 7.0−8.0), or N-(2-hydroxyethyl)piperazine-N′-(3-propanesulfonic acid) (EPPS; pH 8.5) buffer dissolved inthe reaction electrolyte. For each isotherm experiment, sampleswere prepared from a 250 mL volume of a 0.05 g L−1 birnessitesuspension prepared in a polyethylene container. Twelve 20mL aliquots were pipetted from the birnessite suspension into30-mL opaque polyethylene tubes. The samples were spikedwith aqueous Mn(II) from a 0.05 M MnCl2 stock to achieveinitial concentrations in the range 50−2200 μM, and thensealed and equilibrated inside the glovebox for 8 days. Theinitial aqueous Mn(II) concentrations in all samples were belowsaturation with respect to any Mn(II) precipitates includingMn(OH)2 as determined from speciation calculations in VisualMINTEQ employing the MINTEQA2 database.36 Followingreaction, the samples were syringe-filtered through 0.22 μmnitrocellulose membranes, and the filtered solids were syringe-washed with 5 mL of anoxic DDI water, and then dried insidethe glovebox prior to IR analysis (see below). Filtered reactionsolutes were analyzed for dissolved Mn(II) using theformaldoxime method37 to determine Mn(II) sorption,calculated as the difference between the initial and final Mn(II)solution concentrations. Isotherm samples for XRD and EXAFSmeasurements (see below) were prepared using suspensionvolumes of 2 L to ensure sufficient Mn oxide product foranalysis. Control samples were run in parallel to the sorptionsamples, and consisted of birnessite suspensions identical to

those of the sorption experiments, except that no Mn(II) wasadded. Analysis of the control sample substrates showed noevidence for mineralogical transformation or modification ofthe birnessite substrate, confirming the sorbent to be stable inthe absence of aqueous Mn(II) in the pH range considered.To assess the influence of O2 on the solid phase partitioning

of Mn(II) in the pH range considered here, a series of oxicMn(II)-adsorption isotherm experiments was performed aswell. The oxic samples were prepared using the sameexperimental conditions and procedures as described abovefor the anoxic experiments, except that the experiments wereconducted outside the glovebox under ambient conditions.

Spectroscopic Analyses of Mn-Oxide Sorption Prod-ucts. Sample solids collected from the isotherm and kineticsorption experiments described above were characterized by X-ray diffraction (XRD) analysis, attenuated total reflectanceFourier transform infrared (ATR-FTIR) spectroscopy, and X-ray absorption spectroscopy (XAS). Details of these analysesare described in the SI.

■ RESULTS AND DISCUSSIONMn(II)−Birnessite Sorption Isotherms. Figure 1 shows

the results of the anoxic birnessite-Mn(II) isotherm experi-

ments performed at pH 6.0−8.5. The isotherms report Mn(II)sorption (i.e., the y-axis values) as the difference between theinitial and final aqueous Mn(II) concentrations, and thusquantify Mn(II) sorption as the reduction in the Mn(II)solution concentration resulting from Mn(II) interaction withbirnessite. We do not normalize Mn(II) removal to the mass ofbirnessite sorbent, since, as will be shown below, reaction withMn(II) causes bulk mineralogical transformation of thebirnessite substrate and increases the mass of solid phaseMn-oxides present in the samples.The isotherms presented in Figure 1 exhibit a clear pH trend,

with generally higher Mn(II) removal observed at higher pHvalues across the Mn(II) concentration range. These findingsare consistent with the notion that metal sorption onto mineralsubstrates increases with increasing pH.38 In addition to theextent of Mn(II) sorption, the patterns of Mn(II) removal asdefined by the sorption isotherms differ markedly with pH aswell. At pH 7.5, the isotherm exhibits high affinity behaviorwith a sorption plateau of approximately 400 μM Mn(II)removed from solution (Figure 1), as previously reported by

Figure 1. Mn(II) sorption in anoxic 0.05 g L−1 birnessite suspensionsat pH 6.0−8.5 following 8 days of reaction, plotted as a function of theremaining Mn(II) solution concentration.

Environmental Science & Technology Article

dx.doi.org/10.1021/es402108d | Environ. Sci. Technol. 2013, 47, 10364−1037110365

Elzinga.26 High affinity partitioning is also observed for theisotherms measured at pH 6.0 and pH 8.5, with sorptionmaxima of ∼100 and ∼800 μM, respectively; the pH 5.5isotherm (not shown) is similar to the pH 6.0 isotherm with asorption maximum near ∼100 μM. The isotherms measured atpH 7.0 and pH 8.0 appear to exhibit two sorption plateaus inthe aqueous Mn(II) concentration range considered. At pH 7.0,Mn(II) removal reaches an apparent plateau of 100 μM atdissolved Mn(II) concentrations of 250−500 μM, but thendistinctly increases at aqueous [Mn(II)] > 500 μM,approaching the same sorption maximum of ∼400 μMobserved at pH 7.5 (Figure 1). Similarly, at pH 8.0, Mn(II)removal appears to plateau at a level of ∼400 μM (the samemaximum as observed at pH 7.5) for aqueous Mn(II)concentrations up to ∼450 μM, but then increases at higherMn(II) solution concentrations toward a plateau of ∼800 μM,which is similar to the sorption maximum observed at pH 8.5(Figure 1).The isotherm data presented in Figure 1 indicate that both

solution pH and the aqueous Mn(II) concentration stronglyimpact Mn(II) sorption by birnessite. Elzinga26 reported thatbirnessite reacts with Mn(II) to produce Mn(III)OOH phasesat pH 7.5 at the aqueous Mn(II) concentrations employed inthe current experiments. The strong pH dependence of Mn(II)sorption seen in Figure 1 suggests that different or additionalMn(II) removal processes may occur at the other pH valuesconsidered here. The spectroscopic analyses presented nextaddress the mechanistic influence of pH and the aqueousMn(II) concentration on the processes involved in theinteraction between Mn(II) and birnessite.XRD, XAS, and IR Results of Mn-Oxide Sorption

Products. The XRD patterns collected for Mn(II)−birnessitesorption samples reacted in the pH range 5.5−8.5 and at initialMn(II) concentrations between 200 and 2200 μM arepresented in Figure 2 (e-u), where they are compared tothose of birnessite, feitknechtite, manganite and hausmannite(a-d). The sorption samples show distinct changes inmineralogy as a function of pH and the aqueous Mn(II)input. The XRD patterns of the samples reacted at pH 5.5 and6.0 and the pH 7.0 sample reacted at low Mn(II) only containthe XRD reflections characteristic of hexagonal birnessite(Figure 2, e-g). However, the intensities of the two peaks at12.3 2θ (7.2 Å) and 24.8 2θ (3.6 Å), which arise from 001 and002 reflections respectively,39 are notably reduced in thesesamples relative to that of the birnessite starting material (a).This indicates that the ordering of birnessite sheet stacking hasbeen disturbed in these samples, but that no bulk mineralogicaltransformation has occurred. In contrast, the XRD data of thesamples reacted at pH ≥ 7.5, and the pH 7.0 samples reacted atintermediate and high Mn(II) inputs (h−u) show the presenceof Mn-oxide phases other than birnessite, which indicates thatMn(II) caused reductive transformation of birnessite intosecondary Mn-oxide minerals in these samples.Elzinga26 studied Mn(II)−birnessite sorption at pH 7.5 and

reported that birnessite converts into manganite (γ-MnOOH)through a metastable reaction intermediate (feitknechtite; β-MnOOH) during reaction with Mn(II) at this pH, withconversion of the metastable feitknechtite phase into manganitepromoted by aqueous Mn(II). Consistent with these findings,the XRD data of the sorption samples reacted at pH 7.5 showthe presence of both feitknechtite and manganite, with theproportion of manganite increasing for samples reacted athigher Mn(II) concentrations (Figure 2, patterns j−m), an

observation further confirmed by IR analyses of the pH 7.5isotherm samples (see SI, Figure S1).The XRD data of the samples reacted at Mn(II)

concentrations levels ≥1000 μM at pH 7.0, and 200−1000μM at pH 8.0 and pH 8.5 also show evidence for the presenceof feitknechtite (Figure 2, patterns h,i, n−p, r and s) withadditional formation of both manganite and hausmannite seenfor the pH 8.0 and 8.5 samples reacted with a Mn(II)concentration of 1000 μM (p, t). For the samples reacted at thehighest Mn(II) concentration (2200 μM), the XRD resultsindicate the presence of manganite in addition to feitknechtiteat pH 7.0 and pH 7.5, and predominantly hausmannite at pH8.0 and 8.5 (i, m, q, and u). The XRD results qualitatively pointto a trend in the types of secondary Mn-oxide mineral productsformed during Mn(II)-driven reductive transformation ofbirnessite as a function of both pH and the level of Mn(II)input in these experiments, where feitknechtite is formed at lowand intermediate Mn(II) concentrations across the pH range,manganite is a major secondary phase in the near-neutral pHsamples (7.0 and 7.5) reacted at intermediate and high aqueousMn(II) levels, whereas hausmannite is the dominant secondaryphase at alkaline pH and high Mn(II) concentrations. Themineralogical transformations evident from the XRD data are

Figure 2. XRD patterns of Mn-oxide reference samples (a-d) andMn(II)−birnessite sorption samples (patterns e−u) reacted underanoxic conditions in the pH range 5.5−8.5. The birnessite suspensiondensity in the sorption samples was 0.05 g L−1; numbers indicatedalong the patterns represent the initial Mn(II) solution concentrationsin μM. The red symbols in pattern f mark peaks resulting from theXRD sample holder that also appear in other patterns. The coloredareas in patterns g−u locate XRD peaks characteristic of feitknechtite(blue), manganite (green), and hausmannite (orange) in the sorptionsamples at pH ≥ 7.0 where reductive transformation of birnessiteoccurs.



accompanied by significant modifications of the morphology ofthe Mn-oxide solids as observed with SEM microscopy (SIFigure S2).Quantitative estimates of the mineralogical compositions of

the sorption samples are provided by the results of the linearcombination fits of the Mn K edge EXAFS data. The k3-weighted χ functions of the sorption samples show distinctchanges with pH and Mn(II) concentration, and could be fittedas linear combinations (LCs) of the spectra of the birnessite,feitknechtite, manganite, and hausmannite endmembers, asshown by the match of the raw and fitted sorption χ spectra (SIFigure S3). The LC fit results are summarized in SI Table S1,and graphed in Figure 3, which plots the contributions of the

various Mn-oxide references determined from the fittingprocedure. Consistent with the qualitative trends observedfrom the XRD data (Figure 2), the LC fit results show aprogressive increase with [Mn(II)] in the amount of manganiteformed at the expense of feitknechtite at pH 7.0 and 7.5,whereas hausmannite precipitation is observed at pH 8.0 and8.5 and becomes increasingly important with increasingsolution Mn(II) levels at these pH values (Figure 3). Wenote that we consider the LC fit results in Figure 3 to representa semiquantitative estimate of the Mn speciation in the sorptionsamples, as fractional contributions summing to values <1indicated that the set of reference spectra used for fitting wasnot fully representative of the Mn species present (seediscussion in the SI).The mineralogical information obtained from the XRD and

XAS analyses presented above can be used to explain theMn(II) sorption patterns of the isotherm data shown in Figure1. Reductive transformation of birnessite (nominally Mn(IV)-O2) into feitknechtite and manganite by aqueous Mn(II) can besummarized as follows:

+ +

→ +

+

+

Mn Mn(IV)O 2H O

2Mn(III)OOH 2H

22 2

(1)

Reaction 1 defines a 1:1 stoichiometry between aqueousMn(II) and structural Mn(IV). The concentration of Mn(IV)in the 0.05 g L−1 birnessite suspensions of the isothermexperiments is calculated to be 445 μM based on the Mncontent of 48.9 wt % of the birnessite starting substrate (seeSI). This corresponds reasonably well with the sorptionmaximum of 400 μM Mn(II) removal observed for theisotherms measured at pH 7.5 and pH 7.0 (Figure 1), wherefeitknechtite and manganite are the dominant reductivetransformation mineral products as determined from theXRD and XAS analyses (Figures 2, 3). The macroscopicsorption isotherms measured at pH 7.0 and 7.5 are thusconsistent with the 1:1 stoichiometry between aqueous Mn(II)and structural Mn(IV) predicted from reaction 1.At pH 8.0 and pH 8.5, hausmannite is produced during

reaction of Mn(II) with birnessite (Figures 2, 3). Theconversion of birnessite into hausmannite during reactionwith aqueous Mn(II) is described by

+ +

→ +

+

+

2Mn Mn(IV)O 2H O

Mn(II)Mn(III) O 4H

22 2

2 4(2)

For this reaction, the stoichiometry of Mn(II) to Mn(IV) is 2:1,that is, double that of reaction 1. This agrees with theexperimental adsorption isotherms presented in Figure 1, whichshow that the Mn(II) sorption plateau at pH 8.5 and 8.0 is 800μM, double the amount observed for the pH 7.5 and 7.0isotherms, and is thus consistent with the operation of reaction2 at alkaline pH in the Mn(II) concentration range considered.The observed effects of pH on the secondary Mn-oxide

phases formed reflect both thermodynamic and kinetic controlson operational pathways of birnessite transformation byaqueous Mn(II). The equilibrium constants (Keq) of reactions1 and 2 can be calculated (assuming reversible thermodynamicequilibrium) from the standard Gibbs free energy of reaction(ΔGR

0) using ln Keq = −ΔGR0 /RT, where R is the gas constant

and T is temperature, and ΔGR0 is calculated from the standard

Gibbs free energies of formation of the reactant and products p e c i e s i n vo l v ed (ΔG R

0 = ∑ (ΔG f0 ) p r o d u c t s −

∑(ΔGf0)reactants).

24,26 Elzinga26 showed that application ofthese thermodynamic equilibrium constants to predict thresh-olds of Mn(II)-induced transformation of birnessite is limitedby uncertainty in thermodynamic parameters of the nano-particulate Mn-oxide minerals involved; nevertheless, theyprovide useful constraints on the relative stabilities of thevarious secondary Mn-oxide mineral products formed underthe experimental conditions applied here. Using tabulated ΔGf

0

values,40,41 the thermodynamic equilibrium constants arecalculated as Keq = 10−12.14 for reaction 1 producingfeitknechtite; Keq = 10−7.01 for reaction 1 producing manganite;and Keq = 10−17.94 for reaction 2, which produces hausmannite.Using these constants, the relation between the equilibriumMn(II) solution activity (aMn(II)) and pH can be derived foreach of the transformation reactions, yielding −log(aMn(II)) =2pH − 12.14 for reaction 1 producing feitknechtite; −log-(aMn(II)) = 2pH − 7.01 for reaction 1 producing manganite; and−log(aMn(II)) = 2pH − 8.97 for reaction 2, producinghausmannite. Figure 4 plots these equilibria over the pHrange of 5.0−9.0.

Figure 3. Results of the linear combination (LC) fits of the k3-weighted χ spectra of the Mn(II)- birnessite sorption samples,indicating pH- and [Mn(II)]-driven trends in the fractionalcontributions of the Mn-oxide phases present (black square =birnessite; blue circle = feitknechtite; green triangle = manganite;orange triangle = hausmannite). The raw and fitted χ spectra arepresented in SI Figure S3, and the fit results are tabulated in SI TableS1.



Comparison of the three thermodynamic equilibria showsthat manganite is the expected transformation product of thereductive transformation of birnessite by Mn(II) as it maintainsthe lowest aMn(II) across the pH range considered, whilefeitknechtite is the least stable (Figure 4). The presence offeitknechtite and hausmannite in our samples (Figures 2, 3, S3)thus indicates that thermodynamic considerations alone cannotexplain the experimental results. The pH dependence of theequilibria indicates that conversion of birnessite into lower-valence feitknechtite, manganite and hausmannite becomes lessfavorable (i.e., requires higher aqueous Mn(II) concentrations)with decreasing pH, which is due to the release of protonsduring transformation as defined in reactions 1 and 2. This isconsistent with the experimental observation of MnOOH andMn3O4 phases appearing in the samples reacted at pH ≥ 7.0but not at lower pH values (Figures 2, 3, S3). Feitknechtite hasbeen identified as a transient phase in the overall conversion ofbirnessite into manganite at pH 7.5, where transformation ofmetastable feitknechtite into stable manganite is accelerated byreaction with aqueous Mn(II).26 We attribute the presence ofsubstantial feitknechtite in the samples reacted at pH 7.0 andpH 7.5 using low Mn(II) inputs to the relatively slow kineticsof feitknechtite conversion into manganite under theseconditions, requiring reaction times longer than the 8 daysallowed here to go to completion.Kinetic factors are also likely in play for the pH 8.0 and 8.5

samples, where reaction of birnessite with Mn(II) produceshausmannite (Figures 1−3) despite the fact that conversion tomanganite is predicted to be thermodynamically more favorable(Figure 4). This suggests that the kinetics of the transformationof birnessite to hausmannite at these high pH values are fastrelative to those of the conversion to manganite, so that thereaction producing hausmannite (reaction 2) effectivelyoutcompetes the formation of manganite (reaction 1). Theabsence of hausmannite in the pH 7.5 and 7.0 samples (Figures2 and 3) indicates that hausmannite precipitation is either muchslower (relative to manganite formation) than at pH 8.0 and8.5, or not thermodynamically favorable under these con-ditions. Further thermodynamic and kinetic studies are needed

to resolve in more detail the factors controlling transitioningfrom manganite to hausmannite precipitation as the majorpathway of birnessite transformation by aqueous Mn(II) at pH> 7.5.The pathway of the conversion of birnessite into

hausmannite is not entirely clear from our data, althoughsome constraints can be defined. The presence of feitknechtitein the pH 8.0 and 8.5 isotherm samples reacted at low Mn(II)concentrations (Figures 2 and 3) suggests that the conversionof birnessite into hausmannite may proceed through ametastable feitknechtite intermediate, as in the conversioninto manganite at pH 7.5.26 However, a kinetic experimentconducted at pH 8.5 and an initial Mn(II) concentration of1000 μM (SI Figures S4 and S5), does not show evidence forthe presence of feitknechtite at any time point in the 10 dayexperimental time frame, although hausmannite forms withinhalf an hour. This indicates either that hausmannite formeddirectly in this experiment without crystallization of amineralogically distinct reaction intermediate, or that anyintermediate feitknechtite was quickly converted into haus-mannite and did not build up to a detectable extent. Resolvingthe role of feitknechtite in the conversion of birnessite tohausmannite at alkaline pH requires further study.Since manganite rather than hausmannite is the thermody-

namically predicted endproduct of birnessite transformation, itis useful to consider the stability of hausmannite againstconversion into manganite under the experimental conditionsapplied here. The thermodynamic equilibrium betweenhausmannite and manganite is described by

+

→ +

+

+

Mn(III) Mn(II)O 2H

2Mn(III)OOH Mn2 4

2 (3)

The equilibrium constant for this reaction is K = 1010.92 ascalculated from the ΔGf

0 values of the reactant and productspecies involved using the procedure described above. At pH8.5, the equilibrium Mn(II) solution activity maintained by thisreaction is calculated as 0.83 μM, which is 2−3 orders ofmagnitude smaller than the equilibrium Mn(II) solutionconcentration of ∼300 μM where the sorption plateaus isreached in the pH 8.5 isotherm (Figure 1). This suggests that,once formed, hausmannite is stable with respect to conversionto manganite under our experimental conditions, although, asnoted above, the utility of thermodynamic calculations toaccurately predict the behavior of the nanoparticulateexperimental systems studied here may be limited. ATR-FTIR analysis of a 6 month anoxic Mn(II)−birnessite sorptionsample reacted at pH 8.5 (SI Figure S5) indicates no manganiteand the persistence of hausmannite, which suggests long-termstability of hausmannite under these conditions consistent withthermodynamic prediction.Of note in Figure 1 are the “kick-ups” seen in the

experimental isotherms measured at pH 7.0 and 8.0, whereMn(II) removal increases notably from an apparent lower to ahigher sorption plateau within the experimental Mn(II)solution concentration range applied. These distinct increasesin Mn(II) sorption are correlated with the onset of feitknechtiteformation at pH 7.0 (Figures 2, 3, SI Figures S1, S3), and withthe onset of hausmannite precipitation at pH 8.0 (Figures 2 and3), and thus identify Mn(II) solution thresholds for thetransformation of birnessite into feitknechtite and hausmannite,respectively. At pH 8.0, the rise in Mn(II) sorption to a level of800 μM as the solution [Mn(II)] increases to 1 mM (Figure 1)

Figure 4. Theoretical −log(aMn(II)) versus pH stability lines of theMn(II)−birnessite equilibria with manganite, feitknechtite andhausmannite calculated based on thermodynamic data provided byBricker40 and Stumm and Morgan.41 See text for derivation of therelations. The Mn(II)−birnessite equilibrium with manganitemaintains the lowest aMn(II) across the pH range, indicating thatmanganite is the most stable transformation product of the reductivetransformation of birnessite by aqueous Mn(II), whereas feitknechtiteis the least stable, and hausmannite is intermediate.



reflects a kinetic effect where acceleration of hausmanniteformation with increasing [Mn(II)] increasingly outcompetesmanganite formation. The results of the pH 7.0 isotherm are ofinterest from a thermodynamic point of view, as the Mn(II)solution threshold evident from the isotherm can be used toestimate the equilibrium constant for reaction 1 (Keq = (H+)2/aMn(II)). Results from IR analyses of the pH 7.0 isothermsamples indicate first appearance of feitknechtite at an aqueousMn(II) concentration of 274 μM (SI Figure S1). Thiscorresponds to a free Mn(II) solution activity of ∼100 μMcalculated with the Davies equation and accounting for solutioncomplexation of Mn2+ with Cl−, yielding an estimated Keq =10−10.2 for reaction 1 producing feitknechtite based on thecurrent data. This compares to a theoretical value of Keq =10−12.14 calculated from available thermodynamic data asdiscussed above. The 2 orders of magnitude difference betweenexperimental and theoretical Keq underscores the need forimproved estimates of thermodynamic data pertinent to thesesystems. We note that we consider the experimental Keq value apreliminary estimate since our sorption experiments were notnecessarily designed to achieve sorption equilibrium; we arecurrently conducting long-term sorption experiments for moreaccurate estimates of such thermodynamic data from Mn(II)−birnessite sorption isotherm results.The influence of pH on Mn(II)−birnessite interactions is

significant under oxic conditions as well, as demonstrated inFigure 5, where the oxic and anoxic isotherms measured at pH

6.0−8.0 are compared. At pH 6.0, the oxic and anoxicisotherms overlap, indicating the absence of a significantinfluence of O2 on occurring sorption processes. No detectableoxidation of Mn(II) and associated formation of lower valenceMn phases occurs at this pH (Figure 2). In contrast, at pH 7.0and 7.5, and especially at pH 8.0, Mn(II) removal issubstantially higher under oxic than under anoxic conditions(Figure 5). The IR data from the pH 7.0, 7.5, and 8.0 isothermsample solids show a much larger fraction of feitknechtite in theoxic samples relative to equivalent samples reacted underanoxic conditions (Figure 6). We attribute these findings tosurface-catalyzed oxidation of Mn(II) by molecular oxygen

under oxic conditions, where sorbed Mn(II) is oxidized tofeitknechtite according to the reaction:

+ + → ++ +4Mn O 6H O 4Mn(III)OOH 8H22 2 (4)

The reaction represents a pathway of Mn(II) removal underoxic conditions that is operational in addition to the oxidationof Mn(II) by structural Mn(IV) defined by reaction 1, and thusleads to higher overall Mn(II) sorption in oxic experimentsrelative to anoxic systems. The release of protons in reaction 4explains the increased significance of this reaction withincreasing pH evident from the isotherm data presented inFigure 6. Formation of feitknechtite has also been shownduring autocatalytic oxidation of Mn(II) in oxic solutions, andduring surface-catalyzed oxidation of Mn(II) by O2 at thesurfaces of goethite and hematite,42,43 suggesting feitknechtiteas an important metastable Mn-oxide phase in geochemicalsystems exhibiting abiotic Mn(II) oxidation.

■ ENVIRONMENTAL IMPLICATIONSThe results from this study show that pH exerts strong controlon the extent and pathways of the reductive mineralogical andstructural transformation of hexagonal birnessite by aqueousMn(II). The findings are relevant to aqueous geochemicalenvironments with hexagonal birnessite in contact withsolutions containing appreciable dissolved Mn(II), a scenariocommon in, for example, suboxic riparian soils, the redox-clinesof stratified marine and lake water columns, and aquifersaffected by acid mine drainage. Solution pH in theseenvironments is an important variable, with marine systemstypically having pH values near 8, noncalcareous suboxic andanoxic soils pH values near 7.0, and acid mine drainage havingpH values <3.41,44 The results of the current study indicate thatsuch pH variations may significantly impact the extent and

Figure 5. Comparison of oxic and anoxic Mn(II)−birnessite sorptionisotherms at pH 6.0, 7.0, 7.5, and 8.0, measured using a birnessitesuspension density of 0.05 g L−1, and a reaction time of 8 days.

Figure 6. ATR-FTIR spectra of Mn-oxide references and Mn(II)−birnessite sorption samples (0.05 g L−1 suspensions) reacted underoxic and anoxic conditions at pH 7.5 and 8.0. Numbers indicated alongthe spectra represent the initial Mn(II) solution concentration in μM.Bands in the spectral region shown represent OH bending modesallowing identification of manganite and feitknechtite. Reacted solidsfrom oxic experiments (red lines) show additional formation offeitknechtite as compared to equivalent anoxic samples (blue lines)reacted at the same input of aqueous Mn(II).26



mechanisms by which dissolved Mn(II) influences the structureand mineralogy of solid phase Mn-oxides in these aqueousgeochemical environments, with reductive transformationexpected to be particularly relevant at neutral and alkalinepH, whereas effects in acidic systems are likely to be lesspronounced. Reductive transformation of hexagonal birnessiteby aqueous Mn(II) will radically alter the sorption and redoxreactivity of the Mn-oxide mineral fraction, and significantlyaffect the biogeochemical cycling of trace metal(loid)s thatinteract strongly with this mineral, including Co, Zn, Pb, Cr, As,and Se. Further studies are needed to define and quantify thethermodynamics and kinetics of the operational reductivetransformation pathways, and to assess associated impacts onthe fate of sorbed impurities; in turn, the effects of sorbedimpurities on the interaction of Mn(II) with the hexagonalbirnessite surface and resulting mineralogical transformationsrequire further study as well.

■ ASSOCIATED CONTENT

*S Supporting Information(1) Description of synthesis and characterization of the Mn-oxide substrates; (2) Experimental procedures applied toensure strictly anaerobic reaction conditions; (3) Figure withATR-FTIR data of the isotherm sample solids; (4) SEM imagesof Mn-oxide solids; (5) Figure, table, and discussion of thelinear combination fit results of the EXAFS χ spectra; (6)Results of the Mn(II)−birnessite kinetic experiment conductedat pH 8.5. This material is available free of charge via theInternet at http://pubs.acs.org.

■ AUTHOR INFORMATION

Corresponding Author*Phone: (973) 353 5238; e-mail: [email protected].

NotesThe authors declare no competing financial interest.

■ ACKNOWLEDGMENTS

This work was supported by start-up funds from RutgersUniversity to E.J.E. We thank Kumi Pandya (beamline X11A,NSLS) for support during EXAFS data acquisition, and threeanonymous reviewers whose comments helped improve thispaper.

■ REFERENCES(1) Post, J. E. Manganese oxide minerals: Crystal structures andeconomic and environmental significance. Proc. Natl. Acad. Sci. U. S. A.1999, 96, 3447−54.(2) Webb, S. M. M.; Tebo, B. M.; Bargar, J. R. Structuralcharacterization of biogenic Mn oxides produced in seawater by themarine bacillus sp. strain SG-1. Am. Mineral. 2005, 90, 1342−1357.(3) Tebo, B. M.; Bargar, J. R.; Clement, B. G.; Dick, G. J.; Murray, K.J.; Parker, D.; Verity, R.; Webb, S. M. Biogenic manganese oxides:Properties and mechanisms of formation. Annu. Rev. Earth Planet. Sci.2004, 32, 287−328.(4) Spiro, T. G.; Bargar, J. R.; Sposito, G.; Tebo, B. M. Bacteriogenicmanganese oxides. Acc. Chem. Res. 2010, 43, 2−9.(5) Cygan, R. T.; Post, J. E.; Kubicki, J. D. Molecular models ofbirnessite and related hydrated layered minerals. Am. Mineral. 2012,97, 1505−1514.(6) Grangeon, S.; Lanson, B.; Miyata, N.; Tani, Y.; Manceau, A.Structure of nanocrystalline phyllomanganates produced by freshwaterfungi. Am. Mineral. 2010, 95, 1608−1616.

(7) Villalobos, M.; Lanson, B.; Manceau, A.; Toner, B.; Sposito, G.Structural model for the biogenic Mn oxide produced by Pseudomonasputida. Am. Mineral. 2006, 91, 489−502.(8) Santelli, C. M.; Webb, S. M.; Dohnalkova, A. C.; Hansel, C. M.Diversity of Mn oxides produced by Mn(II)-oxidizing fungi. Geochim.Cosmochim. Acta 2011, 75, 2762−2776.(9) Villalobos, M.; Toner, B.; Bargar, J.; Sposito, G. Characterizationof the manganese oxide produced by Pseudomonas putida strain MnB1.Geochim. Cosmochim. Acta 2003, 67, 2649−2662.(10) Manning, B. A.; Fendorf, S. E.; Bostick, B.; Suarez, D. L.Arsenic(III) oxidation and arsenic(V) adsorption reactions onsynthetic birnessite. Environ. Sci. Technol. 2002, 36, 976−81.(11) Pal, S.; Bollag, J.; Huang, P. M. Role of abiotic and bioticcatalysts in the transformation of phenolic compounds throughoxidative coupling reactions. Soil Biol. Biochem. 1994, 26, 813−820.(12) McKenzie, R. M. Manganese oxides and hydroxides. In Mineralsin Soil Environments; Dixon, J. B., Weed, S. B., Ed.; Soil Sci. Soc. Am.Book Ser.: Madison, 1989; pp 439−465.(13) Manceau, A.; Lanson, B.; Drits, V. A. Structure of heavy metalsorbed birnessite. Part III: Results from powder and polarizedextended X-ray absorption fine structure spectroscopy. Geochim.Cosmochim. Acta 2002, 66, 2639−2663.(14) Peacock, C. L.; Sherman, D. M. Sorption of Ni by birnessite:Equilibrium controls on Ni in seawater. Chem. Geol. 2007, 238, 94−106.(15) Toner, B.; Manceau, A.; Webb, S. M.; Sposito, G. Zinc sorptionto biogenic hexagonal-birnessite particles within a hydrated bacterialbiofilm. Geochim. Cosmochim. Acta 2006, 70, 27−43.(16) Villalobos, M.; Bargar, J.; Sposito, G. Trace metal retention onbiogenic manganese oxide nanoparticles. Elements 2005, 1, 223−226.(17) Borch, T.; Kretzschmar, R.; Kappler, A.; Van Cappellen, P.;Ginder-Vogel, M.; Voegelin, A.; Campbell, K. Biogeochemical redoxprocesses and their impact on contaminant dynamics. Environ. Sci.Technol. 2010, 44, 15−23.(18) Manceau, A.; Lanson, B.; Schlegel, M. L.; Harge, J. C.; Musso,M.; Eybert-Berard, L.; Hazemann, J. L.; Chateigner, D.; Lamble, G. M.Quantitative Zn speciation in smelter contaminated soils by EXAFSspectroscopy. Am. J. Sci. 2000, 300, 289−343.(19) Weaver, R. M.; Hochella, M. F. The reactivity of seven Mnoxides with Cr3+(aq): A comparative analysis of a complex,environmentally important redox reaction. Am. Mineral. 2003, 88,2016−2027.(20) McKenzie, R. M. The adsorption of lead and other heavy metalson oxides of manganese and iron. Aust. J. Soil Res. 1980, 18, 61−73.(21) Pankow, J. F.; Morgan, J. J. Kinetics for the aquaticenvironment. Environ. Sci. Technol. 1981, 15, 1155−64.(22) Fendorf, S. E.; Zasoski, R. J.; Burau, R. G. Competing metal ioninfluences on chromium(III) oxidation by birnessite. Soil Sci. Soc. Am,J. 1993, 57, 1508−1515.(23) Tu, S.; Racz, G. J.; Goh, T. B. Transformations of syntheticbirnessite as affected by pH and manganese concentration. Clays ClayMiner. 1994, 42, 321−330.(24) Mandernack, K. W.; Post, J.; Tebo, B. M. Manganese mineralformation by bacterial spores of the marine Bacillus, strain SG-1:Evidence for the direct oxidation of Mn(II) to Mn(IV). Geochim.Cosmochim. Acta 1995, 59, 4393−4408.(25) Bargar, J. R.; Tebo, B. M.; Bergmann, U.; Webb, S. M.; Glatzel,P.; Chiu, V. Q.; Villalobos, M. Biotic and abiotic products of Mn(II)oxidation by spores of the marine Bacillus sp. strain SG-1. Am. Mineral.2005, 90, 143−154.(26) Elzinga, E. J. Reductive transformation of birnessite by aqueousMn(II). Environ. Sci. Technol. 2011, 45, 6366−72.(27) Perez-Benito, J. F. Reduction of colloidal manganese dioxide bymanganese(II). J. Colloid Interface Sci. 2002, 248, 130−5.(28) Lafferty, B. J.; Ginder-vogel, M.; Zhu, M.; Livi, K. J. T.; Sparks,D. L. Arsenite oxidation by a poorly crystalline manganese-oxide. 2.Results from X-ray absorption spectroscopy and X-ray diffraction.Environ. Sci. Technol. 2010, 44, 8467−8472.



http://pubs.acs.org

mailto:[email protected]

(29) Zhu, M.; Ginder-Vogel, M.; Parikh, S. J.; Feng, X. H.; Sparks, D.L. Cation effects on the layer structure of biogenic Mn-oxides. Environ.Sci. Technol. 2010, 44, 4465−4471.(30) Webb, S. M.; Tebo, B. M.; Bargar, J. R. Structuralcharacterization of biogenic Mn oxides produced in seawater by themarine bacillus sp strain SG-1. Am. Mineral. 2005, 90, 1342−1357.(31) Learman, D. R.; Wankel, S. D.; Webb, S. M.; Martinez, N.;Madden, A. S.; Hansel, C. M. Coupled biotic-abiotic Mn(II) oxidationpathway mediates the formation and structural evolution of biogenicMn oxixes. Geochim. Cosmochim. Acta 2011, 75, 6048−6063.(32) Scheidegger, A. M.; Sparks, D. L. A critical assessment ofsorption-desorption mechanisms at the soil mineral/water interface.Soil Sci. 1996, 161, 813−831.(33) Brown, G. E.; Henrich, V. E.; Casey, W. H.; Clark, D. L.;Eggleston, C.; Felmy, A.; Goodman, D. W.; Gratzel, M.; Maciel, G.;McCarthy, M. I.; Nealson, K. H.; Sverjensky, D. A.; Toney, M. F.;Zachara, J. M. Metal oxide surfaces and their interactions with aqueoussolutions and microbial organisms. Chem. Revi. 1999, 99, 77−174.(34) Morgan, J. J. Manganese in natural waters and earth’s crust: Itsavailability to organisms. Met. Ions Biol. Syst. 2000, 37, 1−34.(35) Diem, D.; Stumm, W. Is dissolved Mn2+ being oxidized by O2 inabsence of Mn-bacteria or surface catalysts? Geochim. Cosmochim. Acta1984, 48, 1571−1573.(36) Gustafsson, J. P. Visual MINTEQ v. 3.0 2010.(37) Burle, E.; Kirby-Smith, W. Application of formaldoximecolorimetric method for the determination of manganese in the porewater of anoxic estuarine sediments. Estuaries Coasts 1979, 2, 198−201.(38) Sparks, D. L. Toxic metals in the environment: The role ofsurfaces. Elements 2005, 1, 193−197.(39) Drits, V. A.; Silvester, E.; Gorshkov, A. I.; Manceau, A. Structureof synthetic monoclinic Na-rich birnessite and hexagonal birnessite. 1.Results from X-ray diffraction and selected-area electron diffraction.Am. Mineral. 1997, 82, 946−961.(40) Bricker, O. Some stability relations in the system Mn-O2-H2O at25° and one atmosphere total pressure. Am. Mineral. 1965, 50, 1296−1354.(41) Stumm, W.; Morgan, J. Aquatic Chemistry: Chemical Equilibriaand Rate in Natural Waters, 3rd ed.; Wiley: New York, 1996.(42) Kessick, M. A.; Morgan, J. J. Mechanism of autoxidation ofmanganese in aqueous solution. Environ. Sci. Technol. 1975, 9, 157−159.(43) Junta, J.; Hochella, M. Manganese (II) oxidation at mineralsurfaces: A microscopic and spectroscopic study. Geochim. Cosmochim.Acta 1994, 58, 4985−4999.(44) Bigham, J. M.; Schwertmann, U.; Carlson, L.; Murad, E. Apoorly crystallized oxyhydroxysulfate or iron formed by bacterialoxidation of Fe(II) in acid mine waters. Geochim. Cosmochim. Acta1990, 54, 2743−2758.



environmental science and technology

Documents