BepriD\e~ from ENVIRONMENTAL SCIENCE a TECBI~OLOGY. Vol x. t"~~CopYricht I> 1"2 by the AmeriC&A Cbemic~ Society aDd repriDced by permi8io8 of the cop;~t 0WDeI'.

St8f.n P81ffer...t MarII dO8 ~~nt08 AfonIO.* 88rnhIrd W8ld.t 8nd R... Gichtert

Lake, Researd'l LabCX'atory of E),WAG/ET1i. CH-6O47 Kastanlenbawn. Depto de C~ lnorgira. AI\aftlca y ~Fislca. Fa~d de CIencias Exa,ctas y Natl.-'8les. Ck.dad ~ PabeI6n II. 1.28 Buenos AWes. ~

.The initial reaction betweel:l hydrogen sulfide and thesurface of lepidocrocite was studied in the pH range be-tween 4 and 8.6 by monitorin;g the change of the emf ofa pH.,s sensor. The rate of H:~ oxidation is pseudo firstorder with respect to H.,s and shows a strong pH depen-den,~ with a maximum at p'H 7. Two rate laws werederived: Rg<PH<S> = ha[H+]-2[H.'S]!f1A and R!f<PH<8.6J -hb[H+P[H~]~A with ha = 1.5 X 10- M2 min- and hb -2 X 1(/' M- min-t. The pH IDBlximum of the reaction ratecan be explained by using a surface speciation modelsua:ested for the reductive disllOlution of hematite by H.,s(1): The oxidation rate of H~ is proportional to theconl:entration of inner-sphere surface complexes of HS-formed with either the neutrlll ferric oxide surface sites(>FeOH) or the protonated ferric oxide surface sites(>FeOH2+). The amount of plrotons consumed per moleof H~ suggests that polysulfides and elemental sulfurare the initial products formed during the experiments.

Intr'OductionF,erric oxides play an impo'rtant role in the cycling of

elements in lakes. This is due to their ability to scavengereactive species hy adsorptiOIJl and again to release themto tbe overlying water if the jrerric oxide is dissolved bya reductive procelB.

FeOOH + e- + 3H3C'+ -Fe2+ + 5H2O (i)

In particular, the cycling of phosphate is regarded to belinked to the redox cbemis~, of iron since the classicalItuclies of Einsele (2) and Mo~ltimer (3). Also, the fa~ oftJ'ace elements is 8tl'ongly 8B)Ciated with that of iron (4-7).Mo;reover. the reductive disscdution of ferric oxides is ofgreat importance for the biogec~~mistry of sulfur in lakes:On the one hand, it provides 'Fe2+ ions to form solid Fe8and thus fixes su1fide in the lBediments (8).

Fe2+ + H~ + 2H2O1 -FeS + 2H30+ (ii)

On the other hand, sulfate recluction prOOu~ alkAlinity:

2CH2O + 8042- -..H~ + 2HC03- (ill)

The acidity prOOuced in reaction ii is counierbalanced bythe proton-consuming reaction i. The coupling of these~ reactions may thereforE~ playa key role in neutral-

.To whom corrsponden~ IhOU:Id be addrelled; pr8ent addre18:Limnolocical ~ Station, Ulniverlity of Bayreuth, P .0. Box

101:!51, 8580 Bayreuth, FRG.

fEAWAG/ETH.'I~iudad Universitaria Pabell6D n.

izing lake acidification due to antbJ~pogenic su1furic acid(9). However, ,t;be question remains ,open 88 to which redoxpr0t*8 providM the electzoons for reaction i. When 1MJriedin sediments, ferric iron may be used 88 an electron ac-ceptor by miCJ~rgaDilmB (10,11). On the other hand, itaJao comM intA~ contact with (miCf(JIbiologically prod~)reductants sudb 18 H~ (12). In this context it is im~tto improve our underatanding of tile abiotic chemical in-teraction be~reen the strong redU(:tant hydrogen sulfideand ferric oDclea. The dissolution of metal ondM is gen-erally accepted to be a surface-co][ltrolled proceaa. Thedj8()lution ratAB is proportional to the surface concentrationof reactive sites (13, 14):

R -b.pfi (1)

where R is thei diIIolution rate (m~11 m-2 S-I), k is the rateconstant (S-I), %. is the mole fraction of diIIolution activesites, Pj is the probability of findiIag a specific lite in thecoordinative elrrangement of the precursor complex, andS is the surfa,ce concentration of Ilites (mol m-2).

For ferric OJtidea, detachment of the Fe center from thelurface a&r I1Jrface protonation and surface compleutionwith liganda is888UDled to be the nLte-contzolling step (14,15). The di8BC)lution rate is very much enhanced if ligandcoordination is accompanied by a lteduction of ferric ironby an appropJiate reducing ligand, 801., ascorbic acid (16,17). With ~)ect to diIBOlved sulfide, the diIBOlution rateW88 found to increase with surfa~! area and proton con-centration (18,19). Pyzik and SomJIDer (19) suggested thatHS- is the rellctive SpeciM, which reduces surface ferriciron after an I'xchange reaction with OH-. A subeequentprotonation ~If surface ferrous hydroxide would lead todiIBOlution of a surface layer. Eleu.ental sulfur was foundto be the proJlDinent oxidation pf100uct. Unfortunately,both BtudiM 'gere done only at pE[ ValUM between 7 and8. H Pyzik and Sommer's hypotbesia of HS- 88 the reactivesulfide apeci~1 is valid, a different loeaction pattern shouldbe expected 1~low pH 7, where ]fS- is a minor sulfideSpeciM. TbiB W88 in fact dOCUl]~ented by dOl SantOBAfOD80 and Stwnm (1), who studied the reductive di88Oo.lution of heu1&tite in the pH range 4-7 at various, butoonatant, par1t;ia1 prMl~ of H~. They observed a SUf-face-controlled reaction increasinl: with total sulfide con-centration and decreasing with pro.ton concentration. Thereaction dePImded on the surfaoB concentration of thesurface complexM >FeS- and >FeHS. The overall ratelaw for the diBBOlution W88 derivlld 88

R -k{>FeS-1 + k1>FeHSI (2)

I>FeS1 and {~~FeHSI are surface (XlDceDtratiODB (mol m-2).

0013-936X/92/0926-2.108$03.0010 C 1992 j'"*lcan CheITaI Society14011 EnWon. ScI. T~.. Vol. 28. No.12. 1992

i~,- ---; , ~--~~-J"."~- -'-"~,-~-~--, ~-- , ,-

~r\,l

pH 5.12

.

~-

I'

-7 I .I. .I. I. I I I .I

0 Z 40 ~ ~ 100 1ZI

tinmin

pH 6.94

-6-tl

~-

J

~-

.I

--

.. I I .I .I .I .I I I .I .I .I .I , .I .I .I .I I I I. .I c

O .6 ..10 O .6 ..~ .

tinmin tiDmin

FIgIn 1. Negative kIgertIV11 of rur*- of nxJI8s of ~. n(HtS). ~ cki8 1D ,.ctk)n .~..cIte 8t ~ ~ --.

bad to bE~ oounterbe.1anced by using a oombined ~Impu1so!JDat 614 and ~t 1;65 system, which IIkIWI pBstabilizatjon on a chosen set point. By use of thii pH.-system, interferences of 'Y-FeIOOH with pH buffen ~avoided. The volume of HCI Eldded to the so1utioo durqthe reaction was recorded likewise. The change m ..strength due to the addition is neglectable.

The re18ction vessel was WELShed with 2 M HCl blfmleach run in order to clean adsorbed 'Y-FeOOH fnD .gla88 surface. In some experimental runs lamPI. ~taken for iron determination. A 1-mL lample WM fi]t.-1through a membrane filter (0.2 p.m) into 3 mL of dihd8HNOs (S~uprapur), using a plastic syringe, and ana]y8dby atomiic absorption spectroscopy. We assume the .~ dissolved iron to OOD8ist mainly of Fe(Il) .peca..

Quan1~ification of Total Di.80lved Sulfide. ~slope of the pH~ electrode cell was determined by .externallcalibration procedure (24). In oontzast, the -.measure4i in experimental solutioD8 were relatsd ~ tMinitial blown H~ activity of ,each experiment. ThiI..done in order to prevent differences in rMponse })Er~standard calibration solutioDSI and the reaction Io1utMa.Thus, each pH~ value at a (~rtain time step t wu (8).culated E18pH~t = -Iog ([H~]tot/t-oao) -(emf t-o -emfJ / Sas (I)

where.ao is the protolysis ooefl5cient of H~ and Sat .tIllslope of It.he pH~ electrode cell determined by utema1

calibration (m V).

It is the objective of this Btudy to investigate the reactionof hydrogen sulfide itself with ferric oxide surfaces in thepH range of natural waters (pH 4-8.5) by monitoring theemf of a H~-sensitive electrode cell. Lepidocrocite ( 'Y-FeOOH) was used as a m,~el compound since it is formedduring reoxidation of fem~ iron under natural conditions(20) and is regarded to l>e reactive in sediments (21).

Materials and Methods

Kinetic batch experiml!nts with synthetic lepidocrociteand dissolved sulfide were performed in the pH range 4< pH < 8.7 at a constant ionic strength (1 = 0.1 M KCI)and at room temperature (26 = 1 °C). Dissolved H~ wasmeasured in situ using a 'pH~lectrode cell [AgO .AgII0.2M mlglasslsolutionIAg:~,AgOwith AgO and AgI in 0.2 MHI as the internal reference system, Ingold, Model Ag-275-85-6329 (22)}. The emf was continuously recordedusing an AID data aquisition system. Synthetic lepido-crocite was prepared by folloWing the method of Gupta(23). Surface area was determined as 57.26 m2 g-l (BETmethod). The samples produced electron and X-ray dif-fraction patterns corresI)Onding to lepidocrocite.

All reactions were conducted in a 200-mL glass vessel,into which various electJo<ldes could be placed by gas-tightscrew plugs. The soluti,ons were stirred with a Teflon-coated stirring bar at constant r.te.

The reaction solutions 'were prepared by deaerating 200mL of 0.1 M KCI with N2 and purified by bubblingthrough alkaline p~olsolution (40 9 of pyrogallol and60 9 of KOH into 200 mI. of deionized water) and addingappropriate amounts of reagent grade HCI and Na~ inorder to establish a precll~n pH. The Na~ stock solu-tion was iodometrically checked before each run. Allchemicals were reagent grade. Recording of the emf ofboth, pH and pH~ electrode cell, was started just beforean aliquot of deaerated 'Y-FeOOH stock solution was addedwith a syringe. When the, reaction with sulfide proceeded,protons were consumed 8J1d the su~uent increase of pH

Results

Beactilon Pattern and Reaction Order. In FiI1n 1.the \og8J~thm of the nwriber of moles of H~, q (,..

(H~), is plotted VI time for four different uj,.,-.aIpH val~B, ~ cxmstant surfao~ area ooncentzation d ~oxide (10 m2 L -1) and the same initial amount of ~au1fide (JlQ-4 M). The pattenlS are representative for anexperimtIDta perfonned in this study. At the lower aDd

EnWoo. ScI. TedVIOI.. V~. 26. No.12. 1"2 ..

-5.

-8

~I

.",

..

.."

pH 4

. ,)t"/

.026~-

1.010J;3 .ax,

js.8..:a

:

;

~

2

~/

~

Y

"xQ. I I. I' ., I. I I. , ...I I. I

O 8) 40 S) S) 100 J8)

8moface area [m2]

pH 5

O I ~. ...1 ..., ..I .I I I .I' I ,..1

OmlGJ8X)8X)1CXX)J8X)1.MX)tin.ec

, 4. ~tkIn of ~ ~ .t Jii 5.7.

in the alkaline legiOn. Between pH 5 and 8.6 an empiricalrate law can be derived as follows:

5<pH<6:;'

~J.:

y..-

/

/

£ ,..." """.,ii

..I

0.

0 m 4D «J «J 100 1a>

8urface area [m2)

, 2. 0b8erv8d P881-»-fi"st-order rate constant b 1he ~ctkx1of ~ wIIh ~ 88 a fw1CtkJn of ~ oxkie .nace ar88.

~

-1~-

.

d[H~Jtot/dt -k.[H+J-2[H~JtatA (5)

7<pH<8.6

d[H~Jtot/dt -kb[H+r[H~JtatA -~ (6)n\.

with the rate constants k. -1.5 X 10-13 MI min-l and kb.2.1 X 10-6 M-i min-l f2r the acidic and alkaline ranges,respectively. A denotes ~e surface area concentration inm2 L -1. At pH <5 the slope log k. VI pH decreases again.The observed rate constants in the range 6 < pH < 7 maynot be determined precisely due to the fast kinetics. Itis POBBible that our values reflect only the lower limits ofthe fast reaction rates at neutral pH.

Formation of DialOlved Iron. Dissolved iron was onlymeasured at pH 5.7, 6.4, and 8. No increase with timecould be detected at the two higher pH values. At pH 5.7diBBOlved Fe2+ increased linearly with time (Figure 4).However, it corresponded to only 25 % of the total sulfideconcentration consumed during this experiment.

.-

~ -2

.g~ -(3.

.2

. ,~

-..

., ""c"' i.!

!I. ," 5 e 7 8 8

pH, 3. 0b8erYed pseld>-ht~~ rate constant tCX" ~ ~ctk)nof ~ w8I ~ ci:I8 8S a Mx:tkx1 of J*i (k ...rxm8iz8d k> ufa~8rea).

DiscU88ion

pH Maximum of Reaction Bate. Maxima of the feeaction rate can usually be found at a pH where the productof the con~ntration of two reactants is at maximwn (25).They have aJso been ot.rved in the reductive dissolutionof iron oxides by reductants other than sulfide (26-28). Ina study on the disaolution kinetics of magnetite in thio-glycolic acid, Bawngartner et al. (26) explained that pat-tern by the interaction of thyoglycolate anions with theprotonated OH groups located on the surfa~ of themagnetite particles. In other words, the change in con-~ntration of both, surface-bound and disaolved reactant,was responsible for the rate maximwn at pH 4.5. Incontrast, LaKind and Stone (28) argued that the changein surfa~ speciation alone leads to a maximwn reactionrate of quinones with the goethite surfa~ at pH 4.

According to the reaction scheme proposed by dosSantos AfOD8O and Stumm (1) for the reductive dissolutionof hematite surfa~ sites, the following reaction steps willtake pla~ at the ferric oxide surfa~ at pH values of >5(Table I): Surfa~ complex formation (reaction I) is fol-lowed by an electron transfer (reaction ll) and a subse-quent reversible release of the S- radical (reaction ill).In a further step, the reduced product Fe2+ will be rel~(reaction IV). Reactions I-ll are 8S8wned to be revemDleand the con~trations of the species > FeS-, >FenS, and>FenoH2+ to be in a steady state. We asswne an inner-aphere coordination of su1fide with Fe(Ill) followed by..-<Pv) to ...(dy.) electron transfer (29), which is expected tobe-/aster than an outer-sphere process. The S- radical is88SU1Ded to Ieact in fast steps with the ferric oxide surface

upper ends of the experimental pH range, the initial rateof oonsumption of dissolved sulfide is first order (pH 7.42and 5.12 in Figure 1). However, between pH 6 and' 7 asecond reaction is interfering with the pure flrst-orderpattern (pH 6.03 and 6.94 in Figure 1). A lag phase withinthe first minute can be observed at higher pH values (pH6.94 and 7.42). For the evaluation of data, only the linearrange was used, providing an experimental initial pseu-do-first-order rate oonstant kot.

R = -d[H~]/dt = k.[H~]tot (4)

The concentration of ferric iron surface sites remains ap-proximately constant during the experiment.

Bate Dependency of Surface Area Concentration.In Figure 2, the first-order rate oonstant kot. is plotted asa function of the surface area of 'Y-FeOOH at pH 4 and5. A linear relation can be observed indicating a sur-, face-controlled reaction.

Bate Dependency of pH. Figure 3 reflects the influ-ence of pH on the pseudo-flrst-order experimental rateoonstant kot.. The surface area ooncentration of lepido-crocite is again 10 m2 L -I, and the initial amount of dis-solved sulfide lQ-.f M in each experiment. The rate oon-stant increases with pH up to pH -7 and decreases again

2410 Envron. ScI. Tea.o.. V~. 26. No.12. 1992

.16

.10

om

..'"--.

,.

Tule I. ProPO8&l for the B.edOD 8eqaeD~ of BS- with .1'8ITIc 0Dde Sarface AccordlDC to ref 1

revemDIe ;1dIOrptiU or HS-

At>FeOIJ + HS- ;::: >F~ + H,o

Ls(1)

reversible electron traDIfer

(U)

-q,l808- '..1---C ..-~reversible release of the oxidized prcxluct

(Ill)

detachment of Fe...

(IV)

.2

-g -4"'8b. -6

8) -8 ~-- '-.. --I.

--10 ~.. ~- ~ 1

2 4 6 8 10

pH

~ ,. (8. .) Ca--N.ItI, ~ ~~ IP8c88 >F8S n ~...~n of pH. caJa.8,t8d ~ b 8dIaptIon constanta deIer-"** by ~. SanIDI A.b'8O arwJ SUrm ( 1) b ~118. (b. ~)LOgIrItvn of b n.- fraclklrw of h ...~ .-.1eP6o-~" ,.c18n18 (~.~. .,., ~~+) to fam ..n.~ IP8ci8I >FeHS: ~ QrYe.,.ctanta ..-~ .,., F~+: k)w.' QrYe. ,.ctanta HS-.>FeOH. ~I HaO+; ~-4"' 1i b bolh pkJt8. S. -1.8 X 1o..e mm-2. [~]- .1~ M. I -0.1 M.

data is obtained if oDe plota tbe product of tbe reactantsV8 pH. Combining eq 11 witb eq8 7 8Dd 8 yields

d[HS-]-~~ -(kKF.s- + k'KFeHs[H30j+)I>FeOHI[HS-j

(12&)

Al~tively, HS- was allowed w react with the prown-ated ferric' oxide surface species FeOH2 +..which gives

drHS-]

to higher oxidized sulfur compounds. In Figure 5& weplotted the concentration of sulfide surface specieB 88 afunction of pH under the experimental conditioDB givenabove. The surface complf'..X formation coDBtaDtB for bothspecieB > FeS- and > FeSH

I>FeS-\/I>FeOH\[HS-] = KF.s-= 105.3 (7)

I>FeHSI/I>FeOH\[HS-][H3O+] = KFeHS = 101°.82 (8)

and the acidity coDBtaDt for the deprotonation of >FeSH

I>FeS-I[H3O+]/I>FeHS\- K. = 10-6.6 (9)

were taken from ref 1 and the acidity coDBtaDtB of thelepidocrocite surface (K.1 -10-6.25 mol L -1, Ka2 -= 1~.26mol L -1) from ref 23. At 1eut phenomenologically, the pHmaximum can be explained according to the pH depen-dency of the surface speciation of > FeSH. Application ofthe reaction scheme presented by ref 1 seems to be rea-BOnable, if one takes into account that the relative con-tribution of each surface BpecieB to the reaction rate isweighed by the individual rate CODBtantB k and k' (cf. 1)and that the adsorption CODI5taDtB of HS- to tbe ferric oxideBUrface will be different for hematite and lepid<X:rocite. Wetherefore conclude that the reaction rate of H~ with thelepidocrocite surface is also proportional to the surfacespecieB > FeS- and > FeHS. According to Table 1, thecoDBumption rate of HS- in terms of the surface specieB>FeS is derived 88

d [HS-] k.k21> FeS-\(10)

-T ..(kKF.s-/[H30]+ + k'KraHS>I>FeOH2+1[HS-]

(12b)

CombiDin& eq8 12a and 12b with eq 4., one obtains (Jettingk « k')

(138)k018 -k'KreHS[HaO-41SJP.oHaHS-

and

koba -k'K,..Hs8t0t.8,.eOHt+aHs- (13b)

P'.eOH and tJ,.eOHt+' respectively, denote the fraction of the~LW"lXinding ~ sita; ~ .the prorolysis ooefficientfor HS-. All three variables depend on the pH, while k ,

and K,.eH8 are constants. Th\J18 koba is proportional to[HsO+JaH8"~.OH and a~.oH,+, respectively. In Figure5b we plotted the decadic loganthm of both the products88 a function of pH, 888uming Jl OH- ion per square na-nometer of lepid<x:l'Ocite for the calculation of the surfacecharge. InCleed, a pH pa~m siJnilar to the experimentalone can be observed, independent of the sorbing species(>FeOH2+ or >FeOH). Note that the upper curve [Le.,lag (a~p'..oHt+)J, lib the experimental rate constant koba(Figure 3), ranges over 4 orders of magnitude in the ex-perimental pH interval (pH 4-8.6). In comparison to theaperimentlal data, the maximum of the product of the tworeactants is ~mewhat ~ to a higher pH value. It maytherefore bel concluded that >F~;-- is of minor importance

=dt ~3 + k8t(k3 + k-~[S-)

A eimi1ar derivation can be done for the species > FeHS.Assuming lteady-stau ooncentration for S-f proportion-ality between the rate and the surface ooncentration of> FeS- and > FeSH is obtained:

-d[HS-)/dt -kl>FeS-1 + k'I>FeSHI (11)

In oomparison to the pH dependency of the change insurface speciatiOD, an even better fit of the experimental

envroo. ScI. T8dm., 1iol. 26. No.12, 1992 2.11

U). The ra~ constant remains the lame. However, theblack color disap:a>eared again toward the end of the u-perimenta, which means that the greatest part of thesulfide stored in FeS was also oxidized and onJv a smallportion of the sulfide may have remA~ed as FeS. This isnot surprising. since we worked with excess ferric oxidein contrast to Pyzik: and Sommer'sltudy (19). where FeScould a~ula~ during the experimenta. Rediseolutionof FeS W(Xl]d, however. also imply an A1RA1inity gain W) thatagain the ratios prMented in reactions V-IX (Table U)should be observed. We therefore conclude that the fer-mIl iron iel~ after rediIIOlution of FeS adeorl:8 to theferric oxide surface and forme a surface complex > FeO-Fe+ (reaction XI in Table U). A combination d the ~proceseee (formation of polyeulfides and/or elementalsulfur. precipitation ofFeS. and adeorption ofFe2+) leadsto the low AH+ / AH~ ratios observed at pH >6. At lowerpH values FeS does not form and adeorption of Fe2+ de-creases. providing an increased AH+ / AH~ ratio.

~I

.AH.S.

3.&

~

IF.00H + 4HS- + 2HsO -s... + 6Fei+ + (V)140H-

IF.00H + 6HS"" + SHsO -s... + 8Fei+ + (VI)190H- ..

fFeOOH + HS-+ h,O -SO + 2Fei+ + 5OH- (VU)8F.00H + ~+ 3HsO -BtOa.. + 8Fei+ + (VnI)

160H-8F.00H + HS- + 3HsO -SO... + 8Fei+ + (IX)

150H-

3.8

6.0

8.0

16.0

FOnDaUOD of F~6FeOOH + 10HS- -6F~ + S4t- + 8OH- + (X) 0.8

4HtO6FeOOH + 8>FeOH + 4 HS- -6>FeO-Fe+ (XI) 2.0

+ S.t- + M>H- + 4HtO

.Whi1e 1.ctioDI V-IX m- pnlducta intD lOlution, F~ iifm'med in reactiODI X 8Dd XL

s

I

t

2

1

Conclu8ion

The reaction of hydrogen sulfide with "Y-FeOOH is a fastaurface-controlled plO(8S. The~ is strong evidence thatthe reaction mecbAniam proposed for hematite (1) ~ alsobe applied to our data. This model perfectly explains theol.erved pH dependency of the reaction rate. if only thesurface species > FeSH is oonsidered to be relevant for thee1ectloD-tJo8nsfer p~ Surface ~pic tecbniQU8will be needed to further elucidate the fate of oxidizedsulfur oompounda and ferrous iron at the ferric oxidesurface at pH values higher than 6.5.

The ol.erved muimum of the reaction rate at pH 7correaponda to a pH to which anoxic. sulfide-containingaediment pore waters are usually buffered (30). In addi-tion. the formation of FeS is favored under these condi-tions. Polysulfides may be expected to be at least anintermediate product of the reaction (19). Therefore therequirements for the formation of pyrite from FeS andpolysulfides under natural conditions would be fulfilled(31-33). Upon further oxidation of the intermediate sulfurcompounda to sulfate with ferric oxides. the anaerobicoxidation of sulfide may also play an essential role in thecycling of sulfur in sediments (34). The tonclusion istherefore allowed that abiotic oxidation of H~ by (re-active) ferric oxidM will playa much more important rolein sediment diagenesis than believed 80 far.

x xx Xx

x x x x xx xx x

Acknowledgments

We thank A. T-er and three anonymous reviewen fortheir helpful and constructive comments and criticisms.

Becistry No. Hz8. 7783-08-(; lepidocrociu. 12022-37-6.

for the reaction of H~ with lepidocrocite. In summary,both the pH dependency of the surface speciation of thetwo > FeS complexes and the pH dependency of theproduct of surface-bound and diBBOlved reactant stronglyaupport the applicability of the model proposed in ref Ito our data.

Formation of Products: Stoichiometry of the Re-action. Since we did not measure the oxidation productsof H~, we estimated them by forming the ratio of con-aumed protons per mole of total sulfide conswned. InTable n, the reaction of H~ with a ferric oxide is for-mulated for different oxidation products of HS-. Due tothe wide span of ratios ranging from I to 15 it should bepossible to clearly differentiate between these products.Figure 6 shows that the measured rati~ range between 0.5.00 3.5 and reveals a distinct pH dependence. Sulfate doesnot appear to be a major product in these initial rate ex-periments. This is in (X)Dtrast to the findings of ~ San~AfoD80 and Stumm (1), who measured mainly sulfate,tbi~ulfate, and tl'aces of sulfite as oxidation products ofM~. Since our study ref1ecta more the experimentalconditions as met in Pyzik and Sommer's study (19), wemay also a8ume elemental sulfur or polysulfides (S,2- andS52-) to be the main products, which may be further oxi-dized to sulfate on a longer time scale. .

At pM valuM of >6.5, a black color appeared during theexperiments, indicating the formation of FeS. The lagphase in the H~ consumption, which can be observedwithin the first minute at higher pH values (d. pH 6.94and at pH 7.42 in Figure I), may be explained by thisphenomenon: The reaction rate is increased by the num-her of Fe2+ ions to react with surplus HS- to FeS (e.g., 6Fe2+ in the case of S,2- as a product, reaction X in Table

Literature Cited

(1) Doe Santoe AfOD80, M.; Stumm, W. Langmuir, in pfM8.(2) EiDlele, W. Arch. Hydrobiol. 1936, 29, 664.(3) Mortimer, C. H. J. Ecol. 1941, 29, 280.(.) Sholkovitz, E. R.; Copland, D. Geochim. COImochim. Acta

1182, 46, 393.(5) JohDIOD, A. C. Geochim. COImochim. ActG 1186, 50, 2433.(6) Te8ier, A.; Carignan, R.; Dubreui1, B.; Rapin, F. Geochim.

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Reaervoira; EUer. L A., FAl; ACS AdV8DC8 in ChemjstrySeri.; American Chemical Society: W88hiDgtoD. DC. inp'... .

Received for reuielw February 12, 1992. Revued manu.cnptreceived July 2; 1992. Accepted July 13" 1992.