Kinetics of manganese adsorption, desorption, and oxidation in coastal

marine sediments

Dominique Richard,a,* Bjørn Sundby,a,b and Alfonso Mucci a

a GEOTOP and Department of Earth and Planetary Sciences, McGill University, Quebec, CanadabInstitut des sciences de la mer de Rimouski, Universite du Quebec a Rimouski, Quebec, Canada

Abstract

Ejection of excavated manganese (Mn)-laden particles from anoxic subsurface sediment onto an oxicsediment–water interface by infauna and by other physical disturbances triggers desorption and oxidation ofreduced manganese species. These competing reactions determine whether adsorbed Mn(II) will be desorbed andescape into the water column or be oxidized to Mn(IV) and retained as an insoluble component of the sediment.Consequently, the net flux of Mn(II) to the water column depends on the relative rates of the two reactions. Wedetermined the rates of Mn(II) adsorption, desorption, and oxidation by incubating slurries of natural anoxicsediment at 25uC and monitoring the concentration of manganese in the aqueous phase of the slurry. The kineticsof adsorption, desorption, and oxidation were fitted to first-order rate laws. The first-order rate constants foradsorption (k’ads 5 7.0 6 3.4 h21) and desorption (k’des 5 12.7 6 3.3 h21) are more than two orders of magnitudegreater than the first-order rate constant for oxidation (k’ox 5 0.020 6 0.006 h21). The combination of rapiddesorption kinetics and slow oxidation kinetics allows dissolved manganese to be dispersed into the bottom waterby turbulent diffusion and escape local reoxidation and retention at the sediment–water interface. Thus, advectionof anoxic sediment to the sediment–water interface contributes to the overall flux of manganese from the sedimentto the water column.

This study focuses on manganese, an important compo-nent of marine oxidation–reduction cycles. Manganese(Mn) exists in multiple redox states and is readily cycledbetween these states because of the ease with whichelectrons are transferred between manganese species. Thestrong scavenging ability of manganese oxides, tightlycouples the manganese cycle to the cycles of many othertrace elements (e.g., arsenic, molybdenum, lead) andinfluences the distribution of these elements in the watercolumn and in the sediment (Turekian 1977; Post 1999). Insediments, manganese cycles via burial of oxidized manga-nese below the redox boundary, reductive dissolution ofoxidized manganese to Mn(II), upward diffusion ofdissolved Mn(II), and reoxidation of Mn(II) to insolubleMn(IV) above the redox boundary. This cycle enriches theoxic surface sediment in manganese at the expense of theanoxic subsurface sediment (Sundby et al. 1981).

Mass balance considerations have shown that, in somecoastal environments, the subsurface solid manganeseconcentration, which represents the manganese that istruly retained by the sediment as it is buried, can beconsiderably lower than the manganese content of theparticulate matter brought into the system by rivers. Thisdifference implies a net export of manganese from thesediment to the overlying water and ultimately to the openocean (Sundby et al. 1981; Trefry and Presley 1982;McManus et al. 2012). Manganese may escape from thesediment via several mechanisms. These include diffusionof dissolved manganese (Mn(II) and/or Mn(III)) across theoxidized surface layer (McCaffrey et al. 1980; Sundby et al.1981; Madison et al. 2011) and diffusion of dissolvedmanganese across burrow walls followed by exchange of

burrow water with overlying bottom water (McCaffreyet al. 1980; Burdige 2006). The slow oxidation kinetics ofMn(II) by oxygen (Morgan 2000), the stabilization ofMn(III) by organic ligands (Madison et al. 2011), andepisodic excursions of the Mn(II)/Mn(IV) redox boundarytowards the sediment surface (Sundby 2006) favor trans-port of soluble manganese by diffusion.

It has been proposed that manganese may also bereleased from the sediment via the ejection of anoxicsediments by burrowing benthic organisms (Schink andGuinasso 1978). The proposed mechanism relies onadvective transport of sediment particles directly onto thesediment surface, which short-circuits the redox boundary,and involves desorption and escape of Mn(II) into theoverlying water. Mounds and miniature mud volcanoes atthe sediment surface can show where excavated material isdeposited, and sometimes there is a clear color contrastbetween surface sediment and freshly deposited ejectedsediment because of the different colors of reduced andoxidized forms of iron and manganese (Belley et al. 2010).The success of this mechanism depends on whetheradsorbed Mn(II) is desorbed and escapes into the watercolumn or is oxidized to Mn(IV) and retained as aninsoluble component of the sediment.

Several studies have examined the kinetics of manganeseoxidation in sediments (Edenborn et al. 1985; Canfieldet al. 1993; Thamdrup et al. 1994a). Homogeneous oxidationof Mn(II) by O2 is extremely slow (107 times slower thanoxidation of Fe(II) at pH 8 and 25uC; Morgan 2000), but thepresence of MnO2 and other mineral surfaces in naturalsediments, such as c-FeOOH, a-FeOOH, SiO2, Al2O3, orcalcite catalyzes Mn(II) oxidation (Sung and Morgan 1981;Davies and Morgan 1989; A. Mucci unpubl.). Manganeseoxidation can be microbially mediated, and biological* Corresponding author: dominique.richard@mail.mcgill.ca

Limnol. Oceanogr., 58(3), 2013, 987–996

E 2013, by the Association for the Sciences of Limnology and Oceanography, Inc.doi:10.4319/lo.2013.58.3.0987

987

Mn(II) oxidation has been shown to be faster than abioticMn(II) oxidation, including surface-catalyzed reactions(Hastings and Emerson 1986; Tebo et al. 1997; Morgan2000).

The heterogeneous oxidation rate of Mn(II) in theabsence of oxidized iron and manganese is not known.Likewise, little is known about the sorption kinetics ofmanganese in anoxic sediments. The few sorption experi-ments carried out to date using natural sediments containingvarious amounts of MnO2, expressed as the degree ofoxidation of manganese (Taylor 1987; Canfield et al. 1993),or sedimentary materials (quartz sand, clay, carbonate) (Vander Weijden 1975) show: (1) rapid sorption (completedwithin 1 to 3 h), (2) a linear relationship between theconcentrations of dissolved and adsorbed Mn(II), and (3)Mn(II) uptake proportional to the particle concentration.The rate constants for Mn(II) adsorption and desorption inanoxic sediment, in the absence of MnO2, have not beendetermined experimentally, but a single study has estimatedthe first-order desorption rate constant for Mn(II) by fittingpore-water and solid-phase manganese profiles from conti-nental margin sediments to a reaction–diffusion model (Vander Zee et al. 2001).

In this study, we measured the rates of Mn(II) adsorption,desorption, and oxidation by incubating slurries of anoxicsediment in the presence and absence of oxygen in order toevaluate the ejection of anoxic sediment as a mechanism forexporting dissolved manganese from sediments.

Methods

Sampling and storage—Sediment and seawater werecollected in the Lower St. Lawrence Estuary (LSLE) in July2009 and 2010 at Sta. 23, off Rimouski, Quebec. Undis-turbed sediment cores, composed of a brown-colored surfacelayer (, 1 cm thick) grading downwards to a homogeneousunstratified olive-gray mud, were recovered onboard the R/V Coriolis II with an Ocean Instrument Mark II box corer(0.12 m2). The lithology of these sediments, principallyderived from the mechanical weathering of the crystallinerocks and the Quaternary deposits of the Canadian Shield, isnearly invariant with depth (Loring and Nota 1973). The ,10–30 cm subsurface interval, chosen in order to recoveranoxic sediment well below the manganese reduction zone,was subsampled onboard (see Anschutz et al. [2000] forpore-water and solid-phase manganese profiles at Sta. 23).The sediment was transferred to large mason (glass) jars thatwere sealed with a layer of Parafilm (American National CanCo.) before being capped and stored at 4uC. In thelaboratory, the mason jars were brought inside a Coy Labvinyl anaerobic chamber and opened under an N2–H2

atmosphere. Inside the chamber, the anoxic sediment washomogenized by mixing in a ceramic bowl with a plasticspatula at room temperature (23–27uC). The homogenizedsediment was transferred to a large beaker, sealed withParafilm, and kept in the anaerobic chamber until needed foranalysis and slurry incubation experiments. The sedimentwas stored for 12 to 24 months between the time of samplingand the slurry incubation experiments.

Seawater was collected , 15 m above the sediment–water interface (, 335 m depth) using 12 liter polyvinylchloride Niskin bottles mounted on a conductivity–temperature–depth recorder (CTD) rosette. The seawatersalinity was 34.52, as determined by the calibratedconductivity probe of the CTD (SeaBird SBE 911). Theseawater was brought back to the laboratory in 25 literpolyethylene containers previously washed with Milli-Qwater. In the laboratory, , 10 liters of the seawater werefiltered through a 0.45 mm Millipore polycarbonate filter.The filtered seawater was stored at room temperature in aclean screw-cap polypropylene container until needed forthe slurry experiments. The dissolved manganese concen-tration of the seawater was , 1 mmol L21.

Sediment and pore-water analyses—A pre-weighed ali-quot of homogenized sediment was freeze-dried andreweighed to determine the water content. The absolutereproducibility of the measurement was better than 0.1%for a porosity of 0.776 6 0.001. The freeze-dried sediment,homogenized by grinding with an agate mortar and pestle,was used to determine the reactive manganese contentaccording to the method described by Leventhal andTaylor (1990) and Raiswell et al. (1994). Briefly, reactivemanganese was extracted for 24 h with a 1 mol L21 HClsolution at a solid : solution ratio of 1 : 50. Prior to analysis,the extracts were filtered through a 0.45 mm Milliporepolycarbonate filter and diluted with Milli-Q water. Thereactive manganese content of the sediment was 7.8 mmol g21

dry weight (dry wt). Before each slurry incubation experi-ment, pore water was extracted by centrifuging an aliquot ofthe homogenized wet sediment for up to 3 h at , 220 g, usinga Eppendorf micro centrifuge with a fixed-angle rotor at1400 revolutions per minute. To prevent oxidation artifacts,the centrifuge tubes were filled inside the anaerobic chamberand returned to the anaerobic chamber after centrifugation.The pore water was immediately syringe-filtered through a0.45 mm Millipore polycarbonate filter and acidified with a1% equivalent volume of concentrated HCl before beingstored at 4uC for later manganese analysis. The measuredpore-water manganese concentration (Table 1) was used tocalculate the initial dissolved manganese concentration inthe slurry experiments.

The concentration of dissolved manganese in overlyingwater, pore water, and filtered extracts was determined byatomic absorption spectrometry (AAS; Perkin-Elmer AAnalyst-100) with an air-acetylene flame and an impact-bead burner.External aqueous AAS standards were prepared by appro-priate dilution of a 1000 mg mL21 certified solution(Plasmacal, National Institute of Standards and Technology[NIST]–traceable inductively coupled plasma [ICP]– atomicemission spectroscopy and ICP–mass spectroscopy standardprepared in 4% HNO3) in an NaCl solution of ionic strengthcorresponding to the samples and acidified with a 1%equivalent volume of concentrated HCl. The reproducibilityof the measurements was better than 5% and the detectionlimit was , 1 mmol L21.

Slurry incubation experiments—The rates of manganeseadsorption, desorption, and oxidation in natural sediment

988 Richard et al.

were determined by monitoring the dissolved manganeseconcentration over time in sediment slurries at 25uC. Theexperiments, designed to measure the rates of thesereactions individually, were conducted in duplicate inseparate reactors. Adsorption or desorption was initiatedby mixing wet anoxic sediment and seawater underanaerobic conditions. Once the adsorption or desorptionexperiments reached a steady state, the oxidation reactionwas initiated by exposing the slurry to ambient air. Thesolid : solution ratios at which the experiments wereconducted (26 and 43 g L21) provided sufficient solutionfor analysis and observable responses. The initial condi-tions of the slurry incubation experiments are presented inTable 1.

Desorption experiments were initiated by preparing aslurry of a known amount of homogenized wet sediment(Table 1) and oxygen-free filtered LSLE bottom seawatercontaining , 1 mmol Mn L21. Adsorption experimentswere carried out in a similar manner except that MnCl2was added to the seawater to obtain dissolved manganeseconcentrations between 115 and 130 mmol L21. Thesediment particles were maintained in suspension duringthe experiments by mounting the 400 mL water-jacketedglass reactors on a VWR Scientific orbital shaker set at160 revolutions per minute. The reactors were closed withrubber stoppers equipped with holes for the pH electrodeand gas delivery tubes, and the temperature was maintainedconstant at 25.0 6 0.1uC by circulating water through theglass jacket from a constant temperature bath. Theadsorption and desorption experiments were conducted ina glove bag continuously purged with a stream of acommercially prepared CO2–N2 gas mixture (PO2

, 0.1 Pa,PCO2

5 100 Pa, total pressure 5 1.01 3 105 Pa). This gasmixture also flowed through the reaction vessels. Thepartial pressure of the CO2 in the gas mixture was chosen toreproduce the in situ PCO2

and pH conditions of LSLEbottom water (Mucci et al. 2011). Prior to conducting theadsorption and desorption experiments, 300 mL of filteredLSLE bottom water were placed in each reaction vessel andpurged overnight with the N2–CO2 gas mixture. A fewmilliliters of the seawater were then transferred to acid-cleaned plastic bottles and acidified with a 1% equivalentvolume of concentrated HCl for later analysis of the initialdissolved manganese concentration. The dissolved manga-nese concentration in the slurries at time 5 0 was calculatedby assuming that the seawater and the pore water mixedconservatively. The adsorption and desorption experimentswere initiated by mixing several grams (, 18 or , 30) ofwet, anoxic sediment with the oxygen-free seawater inthe reactors. The homogenized, wet sediment used for theexperiments had previously been transferred inside the glovebag after being quickly weighed. Most of the adsorptionand desorption experiments were conducted at a solid : so-lution ratio of , 43 g L21. One desorption experiment wasconducted at a solid : solution ratio of 26 g L21 (seeTable 1). During the first minute of the experiments, theslurries were stirred manually with a plastic spatula tobreak up wet sediment clumps. The reactors were thencapped and mounted on the orbital shaker. At differenttime intervals over a period of 7 to 95 h, eight to ten 5 mL

Table

1.

Ch

ara

cter

isti

csa

nd

qu

an

tity

of

sea

wa

ter

an

dse

dim

ent

use

dfo

rth

esl

urr

yin

cub

ati

on

exp

erim

ents

.d

l5

det

ecti

on

lim

it.

Ex

per

imen

tN

o.

Init

ial

vo

lum

eo

fse

aw

ate

r(m

L)

Init

ial

[Mn

] dis

s

(mm

ol

L2

1)

[HE

PE

S]

bu

ffer

(mm

ol

L2

1)

Wet

sed

imen

tw

eig

ht

(g)

[Mn

] po

rew

ate

r

(mm

ol

L2

1)

To

tal

vo

lum

eo

fse

aw

ate

r(m

L)

So

lid

:so

luti

on

rati

o

42

85

,d

l0

30

.26

93

73

02

.04

3.4

52

85

,d

l0

29

.50

13

73

01

.54

2.4

62

85

11

00

30

.00

02

83

01

.64

3.1

72

85

10

50

30

.00

02

83

01

.64

3.1

82

91

,d

l0

18

.10

02

33

01

.02

6.0

92

91

,d

l0

18

.10

02

33

01

.02

6.0

10

28

5,

dl

53

0.0

00

21

30

1.6

43

.11

12

85

,d

l5

30

.00

02

13

01

.64

3.1

12

28

51

30

20

30

.00

02

13

01

.64

3.1

13

28

51

30

20

30

.00

02

13

01

.64

3.1

Manganese reactions in sediments 989

aliquots of the slurry were sampled. The aliquots werecollected with a pipette and immediately filtered through a0.45 mm Millipore polycarbonate filter. Filtered samples wereremoved from the glove bag and transferred to acid-cleanedplastic bottles, acidified with a 1% equivalent volume ofconcentrated HCl, and stored at 4uC until analysis.

The oxidation experiments were initiated by purging thereactors with ambient air after the completion of theadsorption and desorption experiments. Over the next 45to 179 h, eight to ten 5 mL aliquots of the slurries weresampled at different time intervals. Each aliquot wasimmediately filtered through a 0.45 mm Millipore polycar-bonate filter, acidified with a 1% equivalent volume ofconcentrated HCl, and stored at 4uC until analysis.

The pH was measured on the seawater scale (pHSWS) atregular intervals throughout the slurry experiments using acombination glass-reference electrode (model 6.0259.100Unitrode; Metrohm) and a pH-meter (model 713; Me-trohm). The orbital shaker was stopped during themeasurements. Before and after each measurement, theelectrode was calibrated using three NIST-traceable buffersolutions (4.01, 7.00, 10.00 at 25uC) and a Tris(hydrox-ymethyl)aminomethane (TRIS) seawater buffer solution ofsalinity 33.55 (Millero et al. 1993). The Nernstian responseof the electrode was obtained from the three NIST-traceable buffers and corrected to the seawater scale usingthe TRIS buffer.

In an attempt to minimize pH variations during theoxidation experiments, N-2-hydroxyethylpiperazine-N9-2-ethanesulfonic acid (HEPES) buffer was added to theseawater before preparing the sediment slurries (experi-ments [expts] 10–11, 12–13). The HEPES buffer does notinterfere with Mn(II) sorption onto dMnO2 nor does itsolubilize this mineral (Rosson et al. 1984). The bufferconcentrations used are given in Table 1.

Results

Desorption and adsorption—During each of the threedesorption experiments, the dissolved manganese concen-tration increased rapidly (within minutes), settled backdown over the next few hours, and reached a constant valueafter 4–6 h (Fig. 1; a detailed compilation of the data canbe found in Richard 2011). The final concentration

presumably represents a steady state between desorptionand adsorption. At the higher solid : solution ratio (43 g L21;expts 4–5 and 10–11), the steady-state concentration was 10–15 mmol L21. It was lower, 5.7–8.6 mmol L21, at the lowersolid : solution ratio (26 g L21; expts 8–9). The ratio of thesteady-state dissolved manganese concentration ([Mn(II)-

diss]ss) to the mass of solid sediment (dry wt) was nearlyidentical in the three experiments (0.30 6 0.05 mmol g21 drywt). This result may be interpreted as an equilibriumbetween a first-order adsorption reaction with respect to[Mn(II)diss] and a first-order desorption reaction with respectto the adsorbed Mn(II) concentration ([S;Mnads]).

Within the narrow range investigated (115–130 mmol L21),the dissolved manganese concentration decreased rapidly inboth sets of adsorption experiments (expts 6–7 and 12–13)and reached a steady state at 58–70 mmol L21 within the firsthour (Fig. 2). Like the desorption experiments, the dissolvedmanganese concentration initially evolved beyond thesteady-state value before settling to a time-invariant value.Given the initial (115–130 mmol L21) and steady-state(65 mmol L21) dissolved manganese concentrations and theexperimental solid : solution ratio, we estimate that 1.45 to1.70 mmol Mn g21 dry wt was adsorbed.

The steady-state values measured in adsorption experi-ments with and without the HEPES buffer (respectively,expts 12–13 and 6–7) are identical within 1 standarddeviation of their respective means. The steady-state valuesmeasured in desorption experiments carried out with andwithout the HEPES buffer (respectively, expts 10–11 and4–5) are identical within 2 standard deviations of theirrespective means. pHSWS remained stable at 7.7 6 0.05throughout the desorption and adsorption experimentsboth in the presence and absence of the HEPES buffer.

Oxidation—A total of four oxidation experiments wereconducted. Two followed desorption experiments (expts4–5 and 10–11) and two followed adsorption experiments(expts 6–7 and 12–13), and all four were conducted withdifferent initial dissolved manganese concentrations. In allcases, the dissolved manganese concentration decreasedover the course of 50 to 150 h and ultimately becameundetectable (Fig. 3A).

Unlike the adsorption and desorption experiments, thepHSWS varied significantly during the oxidation experi-ments (Fig. 3B). During the first few hours of the oxidation

Fig. 1. Desorption expts 4–5 (43 g L21), 8–9 (26 g L21), and10–11 (43 g L21—HEPES 5 mmol L21).

Fig. 2. Adsorption expts 6–7 (43 g L21) and 12–13(43 g L21—HEPES 20 mmol L21).

990 Richard et al.

experiments, the pHSWS rose rapidly and then decreasedgradually until the end of the experiment. The initial pHincrease is caused by the decrease in the CO2 partialpressure (from 100 Pa to 38 Pa) as the reactors were purgedwith ambient air instead of with the N2–CO2 gas mixture.The subsequent decrease in pH is explained by the releaseof hydrogen ions during the oxidation of Mn(II) andprecipitation of MnO2 according to:

Mn2zz1=2O2zH2O?MnO2z2Hz ð1Þ

The largest pHSWS variations were observed in the twoexperiments carried out in the absence of the HEPES buffer(expts 4–5 and 6–7). In these experiments, pH varied by0.43 and 0.50, respectively. In the presence of 5 mmol L21

HEPES buffer (expts 10–11), pHSWS varied by 0.21,whereas at 20 mmol L21 HEPES (expts 12–13) it variedonly by 0.17. The presence of HEPES buffer did notinterfere with the manganese oxidation kinetics.

Discussion

In the following, we first develop the kinetic formalismappropriate for the conditions of our study. We then applythis formalism to our data and derive values for thecorresponding rate constants. Finally, these results are

discussed in the context of previous studies, and theirenvironmental significance is evaluated.

Adsorption and desorption kinetics—In both the adsorp-tion and the desorption experiments, the dissolved manga-nese concentration changed rapidly at the beginning of theexperiments and evolved initially (within minutes) beyondthe longer-term steady-state values (Figs. 1, 2). The way thedata fluctuate during this early stage of the experimentlikely reflects that experiments were carried out withnatural, heterogeneous sediments. We therefore chose tofit the data to simple first-order kinetics. The first-orderrate laws for desorption Eq. 2a and adsorption Eq. 2b are:

d

dtMndiss½ �~k0des S:Mnads½ �

~k0des Mndiss½ �ss{ Mndiss½ �� � ð2aÞ

and

d

dtMndiss½ �~k’ads Mndiss½ �ss{ Mndiss½ �

� �ð2bÞ

where [S;Mnads], [Mndiss]ss, and [Mndiss] are the adsorbed,steady-state, and dissolved manganese concentrations, andk9 is the first-order reaction rate constant.

Integration of Eqs. 2a and 2b gives:

{ lnMndiss½ �ss{ Mndiss½ �tMndiss½ �ss{ Mndiss½ �0

� �~k’t ð3aÞ

and

lnMndiss½ �t{ Mndiss½ �ss

Mndiss½ �0{ Mndiss½ �ss

� �~k’t ð3bÞ

where the subscripts 0 and t denote time zero and time t,respectively. If the desorption or adsorption reaction ratescan be described by first-order kinetics, a plot of ln(([Mndiss]ss

2[Mndiss]t)/([Mndiss]0 2 [Mndiss]t)) or ln(([Mndiss]t 2[Mndiss]ss)/([Mndiss]t 2 [Mndiss]0)) vs. time will give a straightline with a slope of k9 (k’des or k’ads for the desorption andadsorption reactions, respectively). However, dissolvedmanganese concentrations obtained beyond the steady-statevalues cannot be used in Eqs. 3a and 3b, and therefore ourexperimental results cannot be interpreted using theseequations. In order to obtain first-order adsorption anddesorption rate constants using our experimental data, weapplied the following first-order rate equation:

d

dtMndiss½ �~k’ Mndiss½ � ð4Þ

which, in terms of modeling, is assumed to represent theadsorption and desorption reactions. In Eq. 4, [Mndiss] is thedissolved manganese concentration and k9 is the first-orderreaction rate constant. The integrated solution to Eq. 4 is:

lnMndiss½ �tMndiss½ �0

~k’t ð5Þ

Fig. 3. Oxidation expts 6–7 (ads), 12–13 (ads—HEPES20 mmol L21), 4–5 (des), and 10–11 (des—HEPES 5 mmol L21).(A) shows the evolution of dissolved manganese concentrationsduring the course of the experiments. (B) shows the temporalevolution of pH during the course of the experiments. Thesolid : solution ratio was 43 g L21 for all experiments.

Manganese reactions in sediments 991

Accordingly, a plot of ln([Mndiss]t/[Mndiss]0) vs. time will givea straight line with a slope of k9 (k’des or k’ads for thedesorption and adsorption reactions, respectively).

First-order kinetics plots, based on the time evolution ofdissolved manganese concentrations in our adsorption anddesorption experiments, are presented in Fig. 4A,B. Thefirst few minutes of the reactions are well described by afirst-order rate law with first-order rate constants (k’des) of12.7 6 3.3 h21 (r2 5 0.48 [Fig. 4A]) and (k’ads) of 7.0 63.4 h21 (r2 5 0.52 [Fig. 4B]). The uncertainty of the rateconstants (k’des or k’ads) corresponds to the 95% confidenceinterval of the least squares regressions.

To our knowledge, this is the first time that manganeseadsorption and desorption rates on natural anoxic sedi-ments have been measured directly. Van der Zee et al.(2001) estimated a first-order rate constant for manganesedesorption in anoxic sediments from the Iberian continen-tal margin by fitting pore-water and solid manganeseprofiles to a reaction–diffusion model. Their estimatedfirst-order desorption rate constant (1.8 3 1025 to0.027 h21) is 2.5 to 6 orders of magnitude smaller thanour value (12.7 6 3.3 h21). It is not clear why the twoapproaches should generate such a large difference, but itmay reflect the fact that Van der Zee et al. (2001) assumedMn(II) to be adsorbed onto Mn-oxide or -oxyhydroxidesurfaces and released to pore waters in the Mn-oxidereduction zone. Hence, they isolate a desorption rateconstant from modeling of three concurrent processes:Mn-oxide reduction, Mn2+ desorption, and authigenicMn2+ mineral formation.

Oxidation kinetics—The concentration of manganesedecreased over time throughout our oxidation experiments(Fig. 3A). Given the autocatalytic nature of the Mn(II)oxidation reaction (Morgan and Stumm 1965), we canexpect the oxidation rate of Mn(II) to increase as theprecipitation of MnO2 progresses. Other factors that mayhave influenced the oxidation rate, given the complexity ofthe natural substrate used in our experiments, includebacterial mediation (Tebo et al. 2004) and catalysis by ironoxides and other particles (Davies and Morgan 1989).Whereas adsorption and desorption took place withinminutes (see Desorption and adsorption experiments), ourdata show that oxidation proceeded over several hours(Fig. 3A). In order to derive a rate constant for the un-catalyzed oxidation reaction for comparison with theadsorption and desorption rates on similar time scales, we

focused on the first few hours of the oxidation experiments.Thus, we assumed that the reaction is not autocatalyzed orinfluenced by Mn-oxidizing bacteria at the beginning of theexperiments. At the start of the experiment, the availabilityof MnO2 surfaces is limited, and the rate can therefore berepresented by pseudo-first-order kinetics (Eq. 4) wherek’ox is the first-order oxidation reaction rate constant. Afirst-order kinetics plot showing the temporal evolution ofdissolved manganese in our oxidation experiments ispresented in Fig. 4C (see Eq. 5 and description below).The first few hours of the oxidation reaction are welldescribed by a first-order rate law with a first-order rateconstant (k’ox) of 0.020 6 0.006 h21 (r2 5 0.52 [Fig. 4C]).The uncertainty of the rate constant (k’ox) represents the95% confidence interval of the least squares regression. Asthe experiments proceeded past the first 12 to 20 h,deviations from the log-linear fit suggest that the oxidationreaction became progressively autocatalyzed (Stumm andMorgan 1996).

Comparing rate constants with literature values—Whencomparing the first-order oxidation rate constant obtainedin the present study with values reported in the literature(Table 2), one should keep in mind that our rate constantdescribes only the first few hours of the reaction duringwhich the abundance of MnO2 surfaces and/or the numberof Mn-oxidizing bacteria had not yet reached sufficientlyhigh abundances to influence the oxidation rate. This mayexplain why the first-order oxidation rate constant wemeasured is lower than most other first-order rateconstants reported for coastal sediments. For example,Edenborn et al. (1985), who conducted slurry experimentswith sediments from the LSLE, found that the first-orderoxidation rate constant for Mn(II) decreased from 50 h21

to 2.74 h21 between a sample taken from the top fewmillimeters of a core and one taken 3 cm below the coretop. Concurrent with the lower rate measured at depth wasa lower number of viable bacteria and lower concentrationsof iron and manganese oxides (Edenborn et al. 1985). Thefirst-order oxidation rate constant we measured, 0.02 h21,is about two orders of magnitude smaller than the value(2.74 h21) they obtained with sediment taken 3 cm belowthe sediment–water interface. A possible explanation forthe discrepancy is that our reducing sediment (collectedfrom the 10–30 cm subsurface interval) provided even fewerMn-oxidizing bacteria and reactive surface sites forcatalytic oxidation.

R

Fig. 4. First-order kinetics plot based on the time evolution of dissolved manganeseconcentrations in slurries during (A) desorption, (B) adsorption, and (C) oxidationexperiments. (A) The first-order rate constant, k’des 5 12.7 6 3.3 h21, was obtained from theslope of the regression line (r2 5 0.48) for the combined expts 4–5 (43 g L21), 8–9 (26 g L21), and10–11 (43 g L21—HEPES 5 mmol L21). Only data obtained during the initial 15 min of eachexperiment were used to calculate the regression (see text for justification). (B) The first-order rate constant, k’ads 5 7.0 6 3.4 h21, was obtained from the slope of the regression line(r2 5 0.52) for the combined expts 6–7 (43 g L21) and 12–13 (43 g L21—HEPES 20 mmol L21).

992 Richard et al.

Only data obtained during the initial 7 min of each experiment were used to calculate theregression. (C) The pseudo-first-order rate constant, k’ox 5 0.020 6 0.006 h21, was obtained fromthe slope of the regression line (r2 5 0.52) for the combined expts 6–7 (ads), 12–13 (ads—HEPES20 mmol L21), 4–5 (des), and 10–11 (des—HEPES 5 mmol L21). Only data obtained in the initial12 h of each experiment were used to calculate the regression. The uncertainty on all least squaresregressions represents the 95% confidence interval.

Manganese reactions in sediments 993

In a study of manganese oxidation in the water column,Yeats and Strain (1990) derived a first-order oxidation rateconstant of 1.41 3 1023 h21, which is about one order ofmagnitude lower than the first-order rate constant wemeasured. Their value was derived from the Mn(II)residence time in a water column where suspendedparticulate matter concentrations were low (1–2 31023 g L21), the temperature was 1.5uC, and the pH was8. These conditions are not directly comparable with ourslurry incubations. Nevertheless, the two rate constantvalues can be reconciled by considering the temperatureeffect on the rate. The rate of homogeneous manganeseoxidation doubles every 11uC (Morgan 1967). In otherwords, the rate constant measured at room temperaturemay be four or five times larger than the rate constant at1.5uC, bringing the two estimates to within a factor of 2.

In our study, desorption and oxidation pseudo-first-order rate constants were calculated using data ([Mndiss])obtained during the early stages of the reactions, i.e.,during the first 15 min of desorption and the first 12 h ofoxidation. During these time intervals, the pHSWS in ourdesorption experiments (7.69 6 0.04) was identical to theLSLE bottom water in situ pH (pHSWS 5 7.69 at 25uC), butthe pHSWS in our oxidation experiments (7.9 6 0.1) wasslightly higher. pH has been shown to have a stronginfluence on Mn(II) oxidation rates (Davies and Morgan1989; Morgan 2000): as pH increases, the homogeneouspseudo-first-order rate constant for Mn(II) oxidationincreases (k’ox 5 k1[O2][OH2]2) (Morgan 2000). Therefore,the pseudo-first-order oxidation rate constant we obtained,k’ox 5 0.020 6 0.006 h21, is likely an upper limit of theoxidation rate constant at in situ pHSWS conditions (7.69),25uC, and 1.01 3 105 Pa atmosphere total pressure.

Is ejection of anoxic sediment particles a plausible escapemechanism for Mn(II) from sediments?—The plausibility ofan escape mechanism for dissolved manganese via theejection of anoxic sediment particles—a process thatbypasses the redox boundary—depends on desorption ofMn(II) competing successfully with oxidation of Mn(II) toinsoluble Mn(III/IV). The first-order rate constants wemeasured, k’des 5 12.7 6 3.3 h21 and k’ox 5 0.020 6

0.006 h21, show that desorption is 100 times faster thanoxidation under our experimental conditions (25uC, 1.01 3105 Pa, pHSWS 5 7.65 to 8). Although the experiments werecarried out at room temperature, desorption remains muchfaster than oxidation at in situ temperatures of 5uC. Aspointed out above, the oxidation rate constant at roomtemperature is four to five times smaller at the in situtemperature. Very few data are available on the effect oftemperature on adsorption and desorption reaction kinet-ics, but it is typically very small (activation energy on theorder of 10 kJ mol21; Yao and Millero 1996). This isequivalent to a , 30% decrease of the first-orderadsorption and desorption rate constant between 25uCand 5uC. Given the estimated four- to fivefold decrease inthe oxidation rate between 25uC and 5uC (see above), theadsorption and desorption rate will remain faster and maybe even faster than the oxidation reaction at 5uC than at25uC. Thus, desorption will be favored even more thanoxidation at the in situ temperature.

When ejected, anoxic sediment is abruptly exposed tooxygenated water, desorption and oxidation occur concur-rently but at different rates. With time, oxidized manganesebuilds up and the oxidation rate (no longer first order)increases. To demonstrate the net result of these twocompeting reactions acting on ejected sediment, we carriedout the following simulation (Fig. 5). The overall rateis obtained by subtracting the first-order oxidation rate(Eq. 4) from the first-order desorption rate (Eq. 2a):

R~d

dtMndiss½ �~k’des Mndiss½ �ss{ Mndiss½ �

� �{k’ox Mndiss½ � ð6Þ

Integration of Eq. 6 gives the concentration of dissolvedmanganese as a function of time:

Mndiss½ �t~ Mndiss½ �ss 1{e{kdes|t� �� �

{e{kox|t ð7Þ

where t is time, and k’des and k’ox are the measured first-order desorption and oxidation rate constants. Accordingto this simulation, the concentration of dissolved manga-nese increases and reaches a maximum of 12.3 mmol L21

during the first 30 min of exposing anoxic sediment tooxygenated bottom seawater at a solid : solution ratio of

Table 2. Compilation of the Mn(II) first-order oxidation rate constants taken from the literature and measured in this study.

SedimentFirst-order rateconstant (h21)

Original sediment depth (cm)and/or redox zone Reference

Coastal, Aarhus Bay 0.04 Diagenetic model result for oxidationin oxic sediment (0–2 mm)

Calculated from the work ofThamdrup et al. 1994b

Pelagic, eastern north Pacific(Manganese Nodule Program)

0.01 Diagenetic model result for oxidationin oxic sediment (0–10+ cm)

Dhakar and Burdige 1996

Coastal, Gulf of St. Lawrence 57 Diagenetic model result for oxidationin oxic sediment (0–5 mm)

Gratton et al. 1990

Saguenay Fjord 114 From 13 to 37 cm, anoxic Mucci et al. 2003Saguenay Fjord 1.4131023 Suspended particulate matter in

oxic water columnCalculated from the work of

Yeats and Strain 1990Coastal, Lower St. Lawrence

Estuary2.74 to 50 0.5 and 3 cm, oxic and suboxic Edenborn et al. 1985

Coastal, Lower St. LawrenceEstuary

0.02 10–30 cm, anoxic This study

994 Richard et al.

43 g L21. This value is nearly identical to the mean steady-state value (12.5 6 2.5 mmol L21) for desorption in expts 4–5 and 10–11 at a solid : solution ratio of 43 g L21. Afterreaching the maximum, the dissolved manganese concen-tration then decreases slowly over several hours asmanganese oxidation and precipitation proceed. Thisscenario agrees well with the results of an experimentcarried out by Saulnier and Mucci (2000), who resuspendedanoxic sediment (13–15 cm interval) from the SaguenayFjord in oxygenated seawater at a solid : solution ratio of35 g L21. In their experiment, a maximum dissolvedmanganese concentration of 11.8 mmol L21 was reachedwithin the first hour, after which the manganese concen-tration decreased over the course of several hours.

In natural aquatic environments, in contrast to theclosed-system experiments in this study and in Saulnier andMucci (2000), desorbed manganese would be dispersed intothe bottom water by turbulent diffusion and would not beavailable for local reoxidation and retention at thesediment–water interface. Combined with rapid desorptionand slow oxidation, dispersion by turbulent diffusionfavors release to the overlying water of manganese sorbedon excavated anoxic sediment deposited on the sedimentsurface. We conclude that bioturbation, in the form ofexcavation of burrows and galleries, followed by depositionof anoxic sediment at the sediment–water interface, is aplausible mechanism for exporting manganese from thesediment to the water column. The overall significance ofthis mechanism depends on the composition of the benthicassemblage and the intensity with which they rework thesediment column.

Caveat—The data from our adsorption, desorption, andoxidation experiments show a significant amount of scatter,which we attribute to the heterogeneous nature of sediments.When experimenting with natural sediments of complexmineralogical and chemical compositions, the solute of

interest (i.e., Mn(II)) may participate in a web of abiotic andmicrobially mediated reactions (sorption, precipitation,dissolution, redox) over which the experimentalist haslimited control. Furthermore, because our experiments werecarried out with St. Lawrence Estuary sediments, the resultsmay not apply directly to other sedimentary environments.Nevertheless, our interpretations of the relative reactionrates upon exposure of reducing sediments to oxygenatedoverlying waters are likely to be valid in most coastal marineenvironments.

AcknowledgmentsThis study was supported financially by Discovery grants and

Ship-time allocation grants from the Natural Sciences andEngineering Research Council of Canada to B.S. and A.M.Awards of J.W. McConnell Memorial, Alexander A. McGregor,and Robert Wares fellowships by the Department of Earth andPlanetary Sciences, McGill University, are gratefully acknowl-edged by D.R. We appreciate valuable comments by ChristianMarz and an anonymous reviewer on an early version of themanuscript.

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Associate editor: Bo Thamdrup

Received: 03 July 2012Accepted: 20 January 2013

Amended: 07 February 2013

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