S C I E N C E O F T H E T O T A L E N V I R O N M E N T 3 9 2 ( 2 0 0 8 ) 1 4 5 – 1 5 6

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Control of sulfide in sewer systems by dosage of iron salts:Comparison between theoretical and experimental results,and practical implications

Dan Firer, Eran Friedler, Ori Lahav⁎

Faculty of Civil and Environmental Engineering, Technion, Haifa, 32000, Israel

A R T I C L E I N F O

⁎ Corresponding author. Tel.: +972 4 8292191;E-mail address: agori@tx.technion.ac.il (O.

0048-9697/$ – see front matter © 2007 Elsevidoi:10.1016/j.scitotenv.2007.11.008

A B S T R A C T

Article history:Received 31 May 2007Received in revised form29 October 2007Accepted 4 November 2007Available online 26 December 2007

Removal of sulfide species from municipal sewage conveyance systems by dosage of ironsalts is a relatively common practice. However, the reactions that occur between dissolvediron and sulfide species in municipal sewage media have not yet been fully quantified, andpractical application relies heavily on empirical experience, which is often site specific. Theaim of this work was to combine theoretical considerations and empirical observations toenable a more reliable prediction of the sulfide removal efficiency for a given dosingstrategy. Two main questions were addressed, regarding the dominant sulfur species thatresults from the oxidation of sulfide by Fe(III) and the dominant precipitation reactionbetween Fe(II) and sulfide species. Comparison of thermodynamic prediction obtained by anequilibrium chemistry-based computer program (MINEQL+) with experimental resultsobtained by dosing ferrous salts showed that the product of precipitation is FeS under alloperational conditions tested. Regarding the reaction between ferric salts and sulfidespecies, analysis of thermodynamic data suggested that the dominant product of sulfideoxidation under typical pe/pH conditions prevailing in municipal raw wastewater is SO4

2−.However, comparison between sulfide removal in laboratory experiments conducted withmultiple samples of raw municipal sewage with a varying composition, and the predictionof MINEQL+ showed the main sulfide oxidation product to be S0.In order to reduce sulfide in sewage to b0.1 mgS/l a minimal molar ratio of around 1.3 Fe to 1S should be applied when ferrous salts are used, as compared with a minimal ratio of 0.9 Feto 1 S required when ferric salts or a mixture of ferrous and ferric salts (at a 2 Fe(III) to 1 Fe(II)ratio) are used. It appears that the high Fe to S(-II) ratios often recommended in practice canbe reduced considerably by applying tight in-line control.

© 2007 Elsevier B.V. All rights reserved.

Keywords:Ferric ironFerrous ironSulfide removalTopic:WastewaterDosing strategy

1. Introduction

Odorous compounds in sewage systems are usually lowmolecular weight volatile compounds (30 to 150 g/mol) thatoriginate from oxidation of sulfur and/or nitrogen containingorganicmaterial or from the reduction of SO4

2− under anaerobicconditions. The most noticeable odorous gas produced ishydrogen sulfide (H2S), which gives septic sewage its typical

fax: +972 4 8228898.Lahav).

er B.V. All rights reserved

rotten-eggs smell. Other compounds (organic and inorganic),present to a lesser extent, include mercaptans (R–CH2–R–SH),volatile organic acids, alcohols, and ammonia (WERF, 2003).The presence of sulfide in wastewater is, by and large, theresult of anaerobic reduction of sulfate by sulfate-reducingbacteria such as Desulfovibrio or Desulfotamaculum, which takesplace in the submerged portion of combined and sanitarysewers, pumping stations wet wells, and forcemains (Elmaleh

.

146 S C I E N C E O F T H E T O T A L E N V I R O N M E N T 3 9 2 ( 2 0 0 8 ) 1 4 5 – 1 5 6

et al., 1998). Reduction of sulfate to sulfide is not favored ifdissolved oxygen (DO) or another electron acceptor morethermodynamically favored than sulfate (e.g., nitrate), ispresent in the aqueous phase (sewage). However, certainsulfate reduction invariably occurs in sewers, even whenoxygen is present in the bulk water because of the limitedpenetrability of O2(aq) into the thick, greasy biofilm thattypically develops on the sewer walls (Lahav et al., 2006).

The rate of sulfide generation depends on several factors:pH, temperature, nutrients, hydraulic retention time, pre-sence of biofilm on the pipe surface, and the oxidation–reduction potential (ORP) (Delgado et al., 1999). Following itsformation, H2S diffuses from the biofilm into the adjacentwastewater, where it dissociates into three species: hydrogensulfide (H2S), bisulfide (HS−), and sulfide (S2−), according to pH,as described in Eqs. (1) and (2) (Lahav et al., 2006).

H2S↔HS− þ H pK1 ¼ 7:05 ð1Þ

HS−↔S2− þ Hþ pK2 ¼ 12:89 ð2Þ

At the typical pH values of municipal sewage (6.5bpHb8.5)the dominant species areH2S(aq) andHS−while S2− is negligible.

H2S(g) has an extremely low odor threshold relative to mostother odorous compounds (around 0.5 ppb), along with hightoxicity to mammals and aquatic species. A part of the H2Sthat is generated in the aqueous phase of sewers is invariablyemitted to the atmosphere of manholes, head space of gravitysewers and pumping stationswetwells where it poses a risk tomaintenance personnel (WERF, 2003). When inhaled, H2S(g)inhibits an enzyme (cytochrome oxidase) which has animportant role in mitochondrial respiration. It also causeseye irritation at concentrations as low as 50–100 ppm, while300–500 ppm may result in severe poisoning, leading tounconsciousness and death. Since hydrogen sulfide is heavierthan air it tends to accumulate at the bottom of manholes andpumping stations. At concentrations higher than 100 ppm, theolfactory system is affected and the human nose cannot sensethe typical smell of rotten eggs, thus maintenance personalentering manholes may not be aware of the danger of lethalconcentrations. Indeed, between 1983 and 1992, more than 29H2S-related deaths and over 5500 exposures occurred in the U.S. (Snyder et al., 1995). Release of H2S to the atmosphere cancause a nuisance and health concern to the neighboringpopulation as well (Boon et al., 1998). Dissolved sulfide mayalso affect biological processes in wastewater treatmentplants (e.g. bulking problems, and inhibition of anaerobicdigestion), and may be toxic to fish in streams affected byoverflow events (Strom and Jenkins, 1984).

H2S released from the liquid phase to the confined seweratmosphere is readily absorbed on the thin moisture layer onthe sewer crown, walls and surfaces above the water line.There, it can be oxidized to sulfate by autotrophic aerobicbacteria (Acidithiobacillus thiooxidans) residing in this moistur-ized layer (Okabe et al., 2007). Corrosion, induced by this acidicreaction can substantially reduce the life of concrete compo-nents of wastewater conveyance systems (WWCS), e.g., pipes,manholes, wet wells, thereby having profound economiceffects (Witherspoon et al., 2004; Hvitved-Jacobsen, 2002).USEPA (1985) estimated that replacement or repair of corrodedsewers cost around US$ 3.2 billion in one decade alone.

In the past three decades, many technologies havebeen studied with the aim of minimizing odor release fromWWCS. These methods can be roughly divided into pre-ventive methods (i.e. minimization of odorous compoundsformation) and removal methods (treating the gas after it hasformed). Treatment methods can be further subdivided intoaqueous-phase and gas-phase treatments. The methodsmost often described in the literature are summarized inTable 1, which also shows a preliminary assessment of theadvantages and disadvantages of each of the WWCS odorcontrol methods.

Out of the methods listed in Table 1, dosage of iron salts isthe most common technique implemented worldwide for theabatement of sulfide-associated problems in gravity sewers.Being specific to H2S, and the relatively low cost of iron saltsare the main advantages of this method. The first advantagecan also be considered a disadvantage, as iron salts are noteffective in removing any other odorous compounds. How-ever, despite the fact that this method has been implementedfor many years and in different parts of the world, thereactions that occur between dissolved iron species andsulfide in gravity sewers have not yet been fully characterizedand quantified and thus the practical application of themethod relies heavily on empirical experience, which isoften site specific. There is also a large inconsistency betweenthe recommendations that appear in various sources regard-ing the required ratio between iron salts and sulfide, whichoften suggests a very high dosage of iron salts, much higherthan the stoichiometric requirements. For example, in the LosAngeles sanitation district where FeCl2 was added in the early1990's to large diameter gravity sewers, a ratio of 7:1was foundnecessary for 90% sulfide removal for dissolved sulfideconcentrations higher than 4 mg S/l, and 15:1 and 100:1 fordissolved sulfide of between 1 and 4mg S/l, and less than 1mgS/l respectively (all ratios based on anhydrous FeCl2 todissolved sulfide on a weight basis) (Padival et al., 1995). Allthese ratios are much higher than the stoichiometric require-ment. Likewise, in the city of Mesa, Arizona, FeCl2 was dosedto interceptor sewers in a ratio that exceeded the stoichio-metric ratio by at least 50% (Jameel, 1989).

The problem of dosage control is exacerbated by the factthat the vague definition “municipal sewage” describes aheterogeneous medium that often changes considerablywithin the course of a single hour, let alone between seasons.As a result, amajor control problem exists when attempting tominimize sulfide concentrations in the sewage via theaddition of iron salts if the effect of various components inthe sewage (pH for example, but also other parameters) has tobemonitored in real time to decide on the dosing strategy. Theaim of the current work was thus to combine theoretical andexperimental considerations in order to quantify the phenom-ena that occur in raw municipal wastewater following theaddition of iron salts. Identification of the controlling reac-tions would, hopefully, allow developing a reliable recom-mendation on dosing strategy, which would aid in reducingthe excessive addition of iron salts that is currently recom-mended, and thus make the method more cost effective. Thewastewater used in the work was collected from a maingravity collector in an arbitrary fashion and its composition(apart from the sulfide concentration and EC) was neither

Table 1 – Advantages and disadvantages of current methods for odor and corrosion control in collection systems andpumping stations

Method References Advantages Disadvantages

Increasing DO concentrationby air (or pure oxygen) injectionto prevent anaerobic conditionsand expedite oxidation of thesulfide in the bulk sewage

Takatoshi et al. (1998);Derek (1995); USEPA (1985);Holder and Leow (1994);Tanaka and Takenaka(1995)

Air is readily available — no needfor transportation; no addition ofchemicals; no negative byproductformation; may improve quality ofsewage arriving at treatment plant.

Low solubility of oxygen in watermakes for a local effect and thusmultiple injection points arerequired; high energy andmaintenance requirements.

Dosage of nitrate salts (KNO3,Ca(NO3)2) to elevate redoxpotential — encouraginganoxic activity ratherthan anaerobic

Bentzen et al. (1995);Derek (1995); Tomarand Abdullah (1994)

Nitrate salts are highly soluble thushigh concentration may be attainedin the water, allowing for fullpenetration into sewer biofilm,with effective prevention ofanaerobic conditions; relativelyinexpensive chemical.

A preventative measure rather thana sulfide removal method. Sulfidegenerated before point of injectionwill not be treated; unwantedaddition of counter cations tosewage (e.g. Na+, K+, Ca2+); need forfrequent transport of chemicals toinjection point; possible negativeeffects on wastewater treatmentplant because of nitrate load; needfor a control system in the sewer tooptimize dosages.

Chemical oxidation of odorouscompounds by ozone, hydrogenperoxide, chlorine compounds,potassium permanganate etc.

Tomar and Abdullah(1994); Boon (1995);Sercombe (1995); Charronet al. (2004)

O3 — no salts addition; harmlessbyproducts; adds oxygen to water.

General disadvantages for allagents: toxic to transportation andmaintenance personnel, expensiveto purchase and handle, non-specificoxidizers thus above stoichiometricamounts are required. Cl2 and Cr2O7

2−

addition releases unwanted salts tothe water.

H2O2 — fairly specific to sulfide,no salts and byproducts addition;adds oxygen to water.Cl2 and Cr2O7

2− strong oxidizers,high removal efficiency for H2Sand other organic substances.

Addition of iron salts — eitherferric or ferrous or combinationof both (FeCl3, Fe(Cl)2, Fe(NO3)3,Fe2(SO4)3).

Poulton et al. (2002); Jameel(1989); Padival et al. (1995)

Specific and effective oxidationof sulfide by Fe3+, followed by FeSprecipitation (double effect); nottoxic; no harmful byproducts;relatively inexpensive; may helpin reducing soluble phosphorouscompounds concentration.

Does not oxidize or precipitate anyother odorous compounds apartfrom sulfides. May add undesirableanions to the water; may causeunsolicited flocculation and settlingin the sewer; may precipitate with Pcompounds thus demand willincrease beyond stoichiometry.

Treatment of malodorous gasesby adsorption (activated carbon)or biofiltration reactors.

Jensen (1990); Wang et al.(2002); LeBeau and Milligan(1994); Vaith et al. (1996);Brennan et al. (1996)

Effective in places where sulfidesare generated locally (likepumping stations). No chemicalsaddition to the water; notransportation requirements;much experience has beengathered with these systems inrecent years.

Treats only the odor problem —does not prevent corrosion. Bedstend to degrade with time; requiresclose maintenance.

Elevation of pH to above 8.5 byaddition of strong base which shiftsthe equilibrium of the dissolvedsulfide towards the non-volatilespecies (S2−, HS−).

Jensen (1990) May be effective in cases wherelocal odor abatement is required.

Effective only locally, since the pH isbound to decrease down stream;addition of unwanted salts. Highrequirement due to high bufferingcapacity of the sewage. Requirestight control.

Microbial supplementation toout-compete sulfate-reducingbacteria

Richman (1997) May also reduce BOD and solidsin WWCS.

Uncertainty in potential impacts ondownstream treatment systems.

Control of headspaceventilation rates; constructionof tall vents – higher thanthe neighboring buildings – inorder to emit the odorouscompounds above thedwelling level.

Olson et al. (1997a,b); Boon(1995); Boon et al. (1998)

Relatively simple solution; doesnot require maintenance;no operating expenses.

Treats only the odor problem —does not prevent corrosion; localeffect; many vents are required —expensive construction; notaesthetic; ineffective in certainwind patterns.

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monitored nor altered. The objective was to understand themain reactions and consequently to establish an appropriatedosing strategy for iron salts for usage in the case of a typical,though ever changing, municipal sewage.

The two dissolved iron species (ferric iron and ferrous iron)that have been typically used for sulfide elimination reactwithsulfide species in two different ways: ferrous iron (Fe2+) tendsto precipitate with sulfide species to form (according to theliterature) a variety of precipitants, while ferric iron (Fe3+)oxidizes sulfide species (to either sulfate or elemental sulfur)while being reduced itself into ferrous iron.

The most common way to describe the reaction betweenferrous iron and sulfide is the following simple precipitationreaction:

Fe2þ þ S2−→FeSðsÞ↓: ð3Þ

However, apart from Eq. (1) one can find in the literatureother possible reactions between Fe2+ and sulfide species. Forexample, Wei and Osseo-Asare (1995) reported that in the firstfew minutes of the reaction between S(-II) and Fe(II) thereaction proceeds via the formation of intermediate speciesbefore FeS precipitates:

Fe2þ þ nHS−→FeðHSÞn2−n ðrapid reactionÞ ð4Þ

FeðHSÞn2−n→FeSðsÞ↓ þ ðn−1ÞHS− þ Hþ ðslow reactionÞ: ð5Þ

This is an interesting observation since Drobner et al. (1990)hypothesized that one of these possible intermediates may beinvolved in the precipitation of solid pyrite (FeS2) as describedin Eqs. (6) and (7):

FeS þ H2S→FeðHSÞ2 ð6Þ

FeðHSÞ2→FeS2ðsÞ↓ þ H2: ð7Þ

Drobner et al. (1990) managed to generate FeS2(s) in the labby mixing S(-II) and Fe(II) under strict anaerobic conditions at100 °C for 2 weeks. However, it is unclear from their workwhether FeS2 precipitation is also possible under the typicalconditions prevailing in gravity sewers (25–35 °C).

Two other possible mechanisms for the precipitation ofpyrite were suggested by Padival et al. (1995):

Fe2þ þ 2HS− þ 0:5O2→FeS2ðsÞ↓ þ H2O þ Hþ ð8Þ

FeS þ S0ðsÞ→FeS2ðsÞ↓: ð9Þ

The actual occurrence of the reactions described in Eqs. (8)and (9) in sewersmay also be considered questionable becauseunder the typical pH range encountered in municipal sewagedissolved oxygen, if exists, is more likely to oxidize sulfiderather than to take part in a precipitation reaction. Moreover,typically, elemental sulfur is not found in sewage at a highconcentration.

However, the option that the reactions described in Eqs.(3)–(9) may occur in gravity sewers should not be overlookedbecause if they do occur, they may significantly affect the Fe2+

dosage required for dissolved sulfide elimination.Possible reactions between ferric iron (Fe(III)) and sulfide

species are also not entirely straightforward. Under the most

common description, Fe(III) may oxidize sulfide species toeither solid elemental sulfur or to sulfate:

2Fe3þ þ HS−→2Fe2þ þ S0ðsÞ↓ ð10Þ

8Fe3þ þ HS− þ 4H2O→8Fe2þ þ SO2−4 þ 9Hþ: ð11Þ

The ferrous ironproduced in theabove reactionsmay furtherreact with surplus sulfide via precipitation (if pH is sufficientlyhigh), thus contributing to theoverall removalefficiencyof ferricsalts. Frompractical and economic perspectives, the knowledgeof whether the product of the oxidation reaction is sulfate orelemental sulfur is important since the amount of ferric ionsrequired to fully oxidize sulfide to sulfate is four times higherthan theamount required for fullyoxidizing sulfide toelementalsulfur. Furthermore, elemental sulfur is an end product that isfar less problematic from the environmental perspective(although in many cases the sulfate background concentrationismuch higher than the concentration possibly generated fromsulfideoxidation). As shown later in thepaper ananalysis basedon equilibrium considerations suggests that the favorablethermodynamic oxidation product is sulfate under the pH andredox conditions typically encountered in municipal sewage.Despite this, it has been widely reported that the main productof sulfide oxidation (at least in tap water) is elemental sulfur.

Apart fromEqs. (10) and (11) the literature cites otherpossiblereactions between ferric iron and sulfide species. Under thetypical pH of municipal sewage Fe(III) hydrolyzes rapidly toform amorphous Fe(OH)3(s). Davydov et al. (1998) suggested thatFe(OH)3(s) reacts with S(-II) to form a variety of species:

2FeðOHÞ3ðsÞ þ 3H2S→2FeSðsÞ↓ þ S0ðsÞ↓ þ 6H2O ð12Þ

2FeðOHÞ3ðsÞ þ 3H2S→Fe2S3ðsÞ↓ þ 6H2O: ð13Þ

According to Davydov et al. (1998) Fe2S3 further transformsto produce the more stable species, FeS2(s) and Fe3S4(s):

2Fe2S3ðsÞ→FeS2ðsÞ↓ þ Fe3S4ðsÞ↓: ð14Þ

Padival et al. (1995) observed that a mixture of ferric andferrous iron salts (at a ratio of 2 to 1) exhibited better sulfideremoval efficiency than thatof either saltby itself. Theseauthorssuggested the following reaction to explain this observation:

Fe2þ þ 2Fe3þ þ 4HS−→Fe3S4ðsÞ↓ þ 4Hþ: ð15Þ

However, Padival et al. (1995) did not show a directevidence for the formation of either Fe2S3(s) or Fe3S4(s), andconsequently the reaction proposed in Eqs. (12)–(15) cannot beconsidered to be more than a hypothesis. Furthermore, whendosing a mixture of iron species with a molar ratio of 2 to 1(Fe(III) to Fe(II)) to a solution that contains sulfide in excess,2/3 mol of Fe(III) will oxidize one third mole of sulfide(assuming that S0 is the final oxidation product as shown inEq. (8)). The 2/3 mol of ferrous iron generated in this reaction(plus the 1/3 mol that was initially dosed as Fe2+) will thenprecipitate with a further 1 mol of S2−. The overall reaction inthis case is theoretically capable of removing 1.33 mol of S2−

per 1 mol of iron, which is identical to the stoichiometricremoval efficiency derived from Eq. (13). Thus, the reactionproposed by Padival et al. (1995) cannot unequivocally explainthe observation that a 2 to 1 ferric/ferrous mixture is

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advantageous in terms of sulfide removal efficiency than thecombined effect of either iron species alone.

The examples given above demonstrate the incompletenature of the information provided in the literature regardingsulfide removal by iron salts in amunicipal sewer environment.The aim of the current work was thus to combine theoreticalconsiderations and experimental observations in order to attaina more reliable picture that would enable prediction of thesulfide removal efficiency for agivendosingstrategyof ferric andferrous iron salts. The two main questions that were addressedin the work were: (1) which is the dominant sulfur species thatresults from the oxidation of sulfide by Fe(III) in a municipalsewage environment, and (2) what is the dominant precipitationreaction between ferrous iron and sulfide species in amunicipalsewage environment. Another question that was addressedempirically was related to the sulfide removal efficiency of acombination of ferric and ferrous salts as compared with that ofeither salt alone. First a theoretical analysis was carried out,based on equilibrium chemistry, using the MINEQL+ computerprogram (Environmental Research Software, 2001). Followingthis, a set of laboratory experiments were conducted to deter-mine the minimal dosage of Fe(II) and Fe(III) salts required toattain N95% sulfide removal efficiency. Following this, theexperimental results were compared with the theoretical onesand conclusions were drawn with respect to dosing strategy.

Fig. 1 –Schematic describing the laboratory experimentalprocedure.

2. Materials and methods

2.1. Analyses

Ferrous iron concentration was determined by themodified o-phenanthroline method proposed by Herrera et al. (1989).Absorbance was measured with a Genesys 10 spectrophot-ometer, Spectronics. Ferric iron concentrationwas determinedby the sulfosalicylic acid technique (Karamanev et al., 2002).Sulfide in “clean” water was analyzed using the Iodimetricmethod and in sewage using the Methylene Blue method(Methods 4500-S2− F and 4500-S2− D; Standard Methods, 2005).

2.2. Experimental procedure

Sulfide stock solution was prepared using crystalline Na2-S·xH2O which was standardized using the Iodometric method(Method 4500-S2− F in StandardMethods, 2005). Rawmunicipalsewagewas taken in themorningbeforeeachexperimentdirectlyfroman80 cmgravity sewer. The samplewas rigorouslymixed inthe lab for 45 min in order to reduce the sulfide concentration inthe sample via oxidation/stripping. The sulfide concentrationthat remained in the sample after 45minwas analyzed using theMethylene Blue method (Standard Methods, 2005).

Sixteen 63 ml bottles were dried, sealed and weighed andsewage was pored into each of the bottles which were thensealed again with rubber caps and left for 30–40 min until alldissolved oxygen was consumed. A volume of between 0.30and 0.315 ml sulfide solution was then injected through therubber caps directly into the sewage in order to attain aconcentration of 5 mg S/l. Strict measures were exercised toavoid sulfide oxidation or volatilization during injection.Ferrous, ferric, or 2:1 ferric to ferrous mixture salts were then

injected into the bottles. Four iron salt dosages were injectedto the 12 bottles (3 repetitions per dosage), two more bottleswere injected with sulfide only and to two more bottles onlyiron salts were added (control). The experimental procedure isdescribed schematically in Fig. 1.

Following the dosage the bottles were mixed thoroughlyand left to stand for 30–40min. Subsequently, 0.11ml of 100 g/lAlCl3 was injected to all bottles in order to allow for betterseparation between the solid and liquid phases. FollowingAlCl3 addition the bottles were mixed thoroughly, weighed,and centrifuged at 2800 rpm for 10 min until two distinctphases were visible. The bottles were then, one at a time,emptied slowly through the rubber cap with a 60 ml syringewithout disturbing the layer that contained the solids. In order

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to maintain a neutral pressure inside the bottle and minimizesulfide and ferrous oxidation N2(g) was simultaneouslyinjected into the bottle through the rubber cap. At the end ofthis step, only the solid phase remained in the bottle. 2.3 ml ofthe supernatant from each bottle was then put into two testtubes and analyzed for sulfide. The rest was injected into a60 ml test tube that contained 1 ml of 1 N HCl for ferric andferrous analysis. The acid was added a priori to prevent Fe(II)oxidation which takes place rapidly at high pH.

The now almost empty bottle with the solid phasesediment was reweighed, uncapped and connected to a zincsulfide trap. The trap consisted of two bottles in a row, eachcontaining 30 ml of a mixture of zinc acetate and NaOH(pHN11). The bottles were capped and N2 was injected into theempty bottle for 1 min to remove oxygen that entered thebottle while it was uncapped. Then, 5 ml of 6 N HCl solutionwas injected into the bottle in order to dissolve the solidphase. N2 was injected into the bottle for 2 min right after theacid injection in order to strip H2S(g) into the sulfide trap, inwhich it precipitated to form ZnS. The two sulfide traps werethen recombined, mixed thoroughly and analyzed for sulfideusing the Methylene Blue method. The original bottles(that now contained dissolved iron species) were weighedagain and the remaining liquid was filtered through a What-man filter #42 (2.5 μm) and the filtrate was analyzed for Fe(II)and Fe(III).

3. Results and discussion

3.1. Theoretical analysis

MINEQL+ ver. 4.5 (Environmental Research Software, 2001) isan equilibrium chemistry-based computer program whichstores equilibrium constants of over 2000 substances. Itreceives reactant concentrations in the aqueous phase asinput as well as operational conditions (pH, temperature, ionicstrength) and yields as output the solid and soluble productsthat exist in equilibrium. In the current work the programwasused to predict the residual sulfide concentration followingthe addition of Fe(II) and Fe(III) salts under the typical range ofIsraeli municipal sewage conditions. These conditions aredescribed in Table 2, which was compiled based on data fromthe National Water Company — Mekorot (1999) and fromGreater Haifa municipal area (HRATS, 2002).

Reactions between Fe(II) and S(-II) species were investi-gated under various values of pH (6–9), ionic strength (0.027–0.19 M), total alkalinity of 500–600 mg/l as CaCO3 (simulated inthe program by the total inorganic carbon concentration, CT,

Table 2 – Range of concentrations of parameters used in the MI

Parameter Units (mg/l) Parameter

[Cl−] 250–350 [Ca2+][SO4

2−] 37–210 [Mg2+][Na+] 100–230 [Zn2+][K+] 20–40 Sulfide (mg/l as S)[NH4

+–N] 20–100

calculated from the pH and alkalinity values) and temperature(19–31 °C). The aim of the runs was to assess the individualeffect of each parameter (and certain combinations thereof) onthe efficiency of the iron salts in removing sulfides. Testedsulfide concentrations ranged between 0.72 and 8.0 mg S/l.

Under all conditions tested, the MINEQL+ program pre-dicted FeS as the most significant precipitation product of thereaction between Fe(II) and S(-II). Other precipitants that havebeen reported in the literature (e.g. FeS2, Fe2S3, Fe3S4) werecalculated to exist at equilibrium only at negligible concentra-tions, and sometimes they were predicted not to exist at all.The two individual parameters found to mostly affect FeSprecipitation were pH and alkalinity. Changing the tempera-ture resulted in only a slight difference in results as did achange in ionic strength at the low pH range (pHb7.5). Athigher pH values, high ionic strength resulted in a lowerfraction of total S(-II) that precipitated as FeS.

As describe above, FeS precipitation is pH dependant. Outof the three soluble sulfide species, only themost basic species(i.e. S2−) reacts directly with Fe2+ to form FeS, thus at a higherpH more FeS is expected to form.

The effect of pH on FeS precipitation as calculated by theMINEQL+ model is shown in Fig. 2. When the pH is increasedfrom 6.5 to 7.5, the remaining sulfide concentration in solutiondecreases from 0.9 mg S/l to 0.2 mg S/l, or in other wordssulfide removal efficiency increases from 55% to 90% (initial S(-II) concentration=2 mg/l). Furthermore it is shown that inorder to reach an equilibrium concentration of 0.1 mg S/l therequired Fe(II) to S(-II) molar ratio at pH 6.5 is 3.5:1, at pH 7 it is1.6:1 and at pH 7.5 and pH 8.1 it is 1.1:1 and 1:1 respectively.Under these conditions the minimum attainable sulfideconcentration decreases by 50% (from 0.09 mg/l to 0.04 mgS/l) when pH increases from 7.0 to 7.5.

The total inorganic carbon concentration (CT) has twocontradicting effects on the equilibrium sulfide concentrationfollowing the addition of Fe(II) salts. On the one hand itprovides buffering capacity, i.e. it is instrumental in minimiz-ing a decrease in pH as a result of sulfide oxidation. On theother hand, when ferrous salts are added at a high Fe(II) to S(-II) molar ratio, precipitation of FeCO3 (siderite) may competewith the formation of FeS, resulting in a much higher residualsulfide concentration than at the same Fe(II) to S(-II) ratio butwith a lower inorganic carbon concentration. The effect of theinorganic carbonate concentration is shown in Fig. 3, usingtwo CT concentrations: 500 and 1200 mg/l as CaCO3. Fig. 3shows that under all three pH values tested (pH 6.5, 7.0 and 7.5)the precipitation of FeCO3 at the higher CT value results in ahigher residual sulfide concentration, but the effect is morepronounced at low pH values.

NEQL+ simulations

Units (mg/l) Parameter Range

40–220 pH 6–926.40 Ionic strength (M) 0.027–0.190.5 Alkalinity (mg/l as CaCO3) 500–6000.72–8

Fig. 2 –Theoretical sulfide concentration at equilibrium as a function of Fe(II) to S(-II) molar ratio applied at various pH values(MINEQL+ prediction for T=23 °C, total inorganic carbon concentration=500 mg/l as CaCO3, EC=1.6 dS/m and initial sulfideconcentration 2.0 mg S/l).

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3.2. Investigation of the reactions between Fe(III) and S(-II)

Based on thermodynamic analysis, sulfide oxidation by Fe(III)is expected to result in sulfate (SO4

2−) as the dominant productfor most pH and redox conditions. The thermodynamicreactions that link the species SO4

2−, S0, and H2S in the aqueousphase are listed in Table 3. Manipulating these equationsallows plotting lines on a pe–pH diagram that represent anequal concentration of each two species in the system. Withrespect to redox reactions, in the area above the line theoxidized species can be considered dominant while below theline the reduced species are dominant. By plotting all the

Fig. 3 –Residual sulfide concentration as a function of Fe(II) to S(-(MINEQL+ prediction for EC=1.6 dS/m and initial sulfide concent

redox and acid–base equilibrium equations it is possible togenerate a Pourbaix diagram which depicts the domains ofdominance of the various sulfur species for all pe and pHconditions (Fig. 4).

Fig. 4 shows, based on thermodynamic considerations, thatthe pe–pH domain in which S0 is expected to form is verynarrow. Above pH 7.0 it is hardly expected to form at all, andSO4

2− is the favorable thermodynamic product. The results ofthis analysis also suggest that in order to end up with S0 underthe conditions prevailing inmunicipal sewage (i.e. pH between6 and 8) the aqueous phase should be in a reducing state(pe=∼−3 or below). When ferric salts are dosed in excess

II) molar ratio applied, at various pH valuesration of 2.0 mg S/l).

Table 3 – Reactions used for plotting the Pourbaix diagram(source of equilibrium constants: Benjamin, 2002)

Reaction Equilibriumconstant

SO42− 8H++6e−⇔S0+4H2O pe0=6.03 Oxidation reduction

half reactionsHSO4−+7H++6e−⇔S0+4H2O pe0=5.66

SO42−+9H++8e−⇔HS−+4H2O pe0=4.21

S+2H++2e−⇔H2S pe0=2.89S+H+2e−⇔HS− pe0=−1.04H2S⇔HS−+H+ pK1=7.05 Weak acid reactionsHSO4

−⇔SO42−+H+ pK1=2.25

152 S C I E N C E O F T H E T O T A L E N V I R O N M E N T 3 9 2 ( 2 0 0 8 ) 1 4 5 – 1 5 6

the wastewater shifts to the oxidative state. According to Fig. 4SO4

2− should then become the dominant specie. Yet, it has beenoften reported (e.g. Padival et al., 1989; Lahav et al., 2004) thatS0 rather than SO4

2− is formed as the main product of thereaction between Fe2+ species and sulfides. Thermodynamicdata, therefore, appears not to give a good prediction of thereactions between ferric salts and sulfides. As stated before,knowledge of the expected sulfur product is of considerableimportance with respect to the dose of ferric salts required foroxidation of a given sulfide concentration.

3.3. Results of laboratory experiments

The results of the experiments in which ferrous salts wereadded to raw municipal sewage are shown in Fig. 5. Generallyit can be stated that the removal of sulfide with ferrous saltsshowed a pattern that was parallel to that calculated by theMINEQL+ program. The simulation lines (representing pH 7.0and pH 7.5) were attained from MINEQL+ by using thefollowing conditions (EC=1.6 dS/m, temperature 27 °C andCT=500mg/l as CaCO3). These conditions were also applied forsimulating the plots shown in Figs. 6 and 7. The measuredrange of conditions of the raw sewage samples used in theresearch was EC: 1.5–8.0 dS/m; pH: 7.0–7.5; CT=380–485mg/l asCaCO3; [NH4

+–N]=20–43 mg/l; and COD=700–1900 mg/l.

Fig. 4 –Pourbaix diagram of

In the low Fe(II) to S(-II) molar ratios (0.5:1 to 0.7:1) sulfideremoval in the laboratory was slightly better than thatanticipated by MINEQL+ (i.e. lower residual sulfide concen-trations). This observation can be attributed to one of the twofollowing reasons: (1) in real sewage sulfide undergoesfurther reactions, apart from FeS precipitation (e.g. precipita-tion with heavy metals that may be present at lowconcentrations); (2) certain spontaneous oxidation of sulfideby oxygen might have occurred despite the effort to createoxygen free environment in the experiments; (3) uncontrolledvolatilization of H2S, due to the fact that the system was notcompletely sealed at all times; and (4) natural analyticalerrors. The third reason is less likely to have occurredbecause no sulfide odor was sensed. When the Fe(II) to S(-II)molar ratio that was applied was higher than 0.9 to 1 at timesthe experimental results showed a slightly lower sulfideremoval (i.e. higher residual sulfide concentration) than thatpredicted by MINEQL+. This may perhaps be attributed touptake of Fe2+ by other species in the sewage (e.g. larger thanexpected FeCO3 precipitation due to CT higher than the500 mg/l as CaCO3 used in the simulation) or to the possiblepresence of substances in the sewage which may retard FeSprecipitation. The reader is reminded that the aim was to testthe practical average sulfide removal in typical, heteroge-neous municipal sewage, and thus the wastewater sampleswere not subjected to comprehensive characterization. Sincethe constituents in the sewage varied considerably betweensamples it was expected that the experimental results willnot precisely match the theoretical prediction which wascalculated based on a constant set of data (constant EC, CT,etc.). Taking this into account, the experimental results werevery close to the theoretical prediction, as much as can beexpected from a heterogeneous varying medium such asmunicipal sewage.

According to Eqs. (10) and (11) it transpires that in order tofully oxidize sulfide to either elemental sulfur or sulfate a 2:1and 8:1 Fe(II):S(-II) molar ratio is required, respectively. In theprocess of oxidizing sulfide Fe(III) is reduced to Fe(II) that, at

the SO42−/S/H2S system.

Fig. 5 –Residual sulfide concentration as a function of Fe(II) added: comparison of empirical and MINEQL+ predicted results(MINEQL+ simulation was carried out with EC=1.6 dS/m, temperature 27 °C and CT=500mg/l as CaCO3; ♦— experimental data).

153S C I E N C E O F T H E T O T A L E N V I R O N M E N T 3 9 2 ( 2 0 0 8 ) 1 4 5 – 1 5 6

the typical pH of municipal sewage, may precipitate with theremaining sulfate as FeS. Thus, 1 mol of ferric salts cantheoretically remove an overall amount of 1.5 (0.5 oxidizedplus 1 precipitated) and 1.125 (0.125 oxidized plus 1 precipi-tated) mol of sulfide when the oxidation product is elementalsulfur or sulfate, respectively. In both cases, the overall sulfideremoval efficiency of ferric salts is higher than that expected

Fig. 6 –Residual sulfide concentration as a function of Fe(III) addefor the two possible oxidation products (MINEQL+ simulation waCT=500 mg/l as CaCO3; ♦ — experimental data).

by the addition of an iron-equivalent amount of ferrous salts.Since the cost of ferrous and ferric salts (normalized per g Fe) isalmost identical this suggests that the use of ferric saltsshould be preferred.

The results of the thermodynamic analysis presented inFig. 6 suggest that themain product of sulfide oxidation underthe conditions prevalent inmunicipal sewage is sulfate. Direct

d: comparison of empirical and theoretical (MINEQL+) resultss carried out with EC=1.6 dS/m, temperature 27 °C and

Fig. 7 –Residual sulfide concentration as a function of the ratio (Fe(III)+Fe(II)) to S(-II). Theoretical prediction by MINEQL+assuming either S0 or SO4

2− as oxidation products at pH 7.5, compared with batch empirical observations (MINEQL+ simulationwas carried out with EC=1.6 dS/m, temperature 27 °C and CT=500 mg/l as CaCO3; ♦ — experimental data).

154 S C I E N C E O F T H E T O T A L E N V I R O N M E N T 3 9 2 ( 2 0 0 8 ) 1 4 5 – 1 5 6

analysis of the end product of sulfide oxidation by ferric saltsrequires themeasurement of either elemental sulfur or sulfateand since the sulfate background is high in Israeli sewage(25–40 mg S/l) and solid elemental sulfur is a problematiccompound to measure accurately in a heterogeneous med-ium such as raw sewage. Thus an indirect approach (i.e.comparison of laboratory results with theoretical resultspredicted by MINEQL+) was adopted for the aim of determin-ing the dominant end product of sulfide oxidation by ferricsalts.

The MINEQL+ program can be used to predict the results ofprecipitation reactions, but it does not cover redox reactions.In order to predict the sulfide concentration following itsoxidation by Fe(III), it was assumed that the oxidationcontinued until all Fe(III) was reduced to Fe(II) (i.e. sulfideconcentration was assumed to be in considerable excess). Themethod for calculating the residual sulfide concentrationfollowing this reaction was: For each Fe(III) to S(-II) ratioinvestigated, the oxidized sulfide concentration was calcu-lated according to either end product possible (i.e. 2:1 or 8:1 Fe(III):S(-II) molar ratio for oxidation to elemental sulfur orsulfate respectively). For example: for an initial sulfideconcentration of 5.0 mg S/l and the addition of ferric iron atamolar ratio of 0.4:1 Fe(III):S(-II) (i.e. 3.48mg Fe3+/l for 5mg S/l),20% of the initial sulfide would be oxidized to elemental sulfuror, alternatively, 5% of the initial sulfide would be oxidized tosulfate. Thus, the remaining sulfide concentration after theoxidation step would be either 4 mg S/l or 4.75 mg S/l,respectively. At the conditions tested in the lab it wasreasonable to assume that Fe(III) would be reduced completelyto Fe(II) and the remaining Fe(II) would precipitate withsulfide. At this point Fe(II):S(-II) molar ratio is either 0.5:1assuming oxidation to elemental sulfur or 0.42:1 assumingoxidation to sulfate. These two new molar ratios were now

inserted as inputs into the MINEQL+ program and the resultswere compared with the laboratory results obtained fromexperiments with raw wastewater to which sulfide and ferricsalts were added at varying concentration ratios, as shown inFig. 6. Since the empirical results were much closer to thetheoretical line that simulated the case in which S(-II) isoxidized to elemental sulfur it was concluded that, in alllikelihood, under all Fe(III):S(-II) molar ratios tested, thatsulfide is in fact oxidized to elemental sulfur rather than tosulfate.

A similar procedure was carried out to analyze the resultsof sulfide removal with a mixture of iron salts (at a ratio of 2:1between Fe(III) and Fe(II)). The effect of themixed dosages wastested following a report in the literature that mixtures wereobserved to result in more efficient sulfide removal (Padivalet al., 1995) than the equivalent dosage of each of theindividual species alone. Again, the experimental laboratoryresults were compared with MINEQL+ prediction. In this caseit was assumed that given an initial 1:1 molar ratio betweeniron species and sulfide, 33.33% of the initial sulfide would beoxidized if the oxidation product was elemental sulfur andonly 8.33% if the oxidation product was sulfate. Here again theresidual sulfide concentration (that remains following thereaction with all the ferric dosed) was assumed to precipitatewith Fe(II).

In the results presented in Fig. 7 it can be seen that theresidual sulfide concentrations obtained in the experimentswere indeed lower than the results predicted by the MINEQL+program for both sulfide oxidation to elemental sulfur and tosulfate (for Fe:S molar ratios of between 0.5:1 and 0.7:1).Nevertheless, the predicted line that was generated underthe assumption that sulfide is oxidized to elemental sulfurismuch closer to the experimental observations. These resultsstrengthen the conclusion derived from the previous observa-

155S C I E N C E O F T H E T O T A L E N V I R O N M E N T 3 9 2 ( 2 0 0 8 ) 1 4 5 – 1 5 6

tions (Fig. 6) that sulfide oxidation with Fe(III) produceselemental sulfur rather than sulfate as the dominant endproduct under the conditions prevailing in municipalwastewater.

The experimental results presented in Fig. 7 suggest, ashypothesized in the literature (Padival et al., 1995), that acombination of ferrous and ferric salts is more effective forsulfide removal in comparison with either salt alone. In orderto understand the theoretical basis behind this observationtwo possible reactions, involving iron species and sulfide, thatare capable of removing sulfide from solution at a ratio higherthan 1.33:1 (Fe to S2−) were considered in addition to thereactions presented in Eqs. (12) and (13) (Davydov et al., 1998),andwhich result in a 2 (Fe) to 3 (S) theoretical ratio. The formerreactions, that involve precipitation of FeS2, are shown in Eqs.(16) and (17) (Drobner et al., 1990; Padival et al., 1995):

Fe2þ þ 2HS− þ 0:5O2→FeS2ðsÞ↓ þ H2O þ Hþ ð16Þ

FeS þ H2S→FeðHSÞ2→FeS2ðsÞ↓ þ H2: ð17Þ

Eq. (16) is unlikely to occur in raw sewage since it requiresthe presence of dissolved oxygen and, furthermore, FeS2(s) is amineral that has been reported to be formed only underconditions that are not encountered in municipal sewage(high temperature and/or pressure). The same reasoningapplies for discounting the reaction depicted in Eq. (17).

Theoretically, Eq. (13) results in sulfide removal at a Fe:S(-II)ratio of 1:1.5 and can thus be used to explain the better sulfideremoval observed. However, since Eq. (13) involves only Fe(III)ions, Fe(II) was not oxidized to Fe(III) under the experimentalconditions, and a mixture of 2:1 Fe3+:Fe2+ was added, theconditions of the mixture under the assumption that thereaction depicted in Eq. (13) takes place allow only for atheoretical removal ratio of 1 (Fe) to 1.33 (S) i.e. the sameremoval ratio that is achieved under the assumption thatsulfide is oxidized to elemental sulfur by ferric salts alone.Therefore, Eq. (13) cannot be used to explain the slightly bettersulfide removal observed both in our study and in Padival et al.(1995). However, it should be noted that this slightly “better”removal efficiency may also be a result of natural experi-mental deviation not unlike the deviation observed in Fig. 6(addition of ferric salts alone) and Fig. 5 (addition of ferroussalts alone).

4. Summary and conclusions

Interactions between iron and sulfide species in sewagemedia are complex and despite the fact that the subject ofdosage of iron salts for the purpose of minimizing sulfideconcentration in wastewater has been investigated for manyyears, the data that appears in the literature is incomplete,and at times even contradicting. In particular, the mainprecipitation product of ferrous iron and sulfide, as well asthe main oxidation product of sulfide oxidation by ferric ironin wastewater medium, has not been unequivocally ascer-tained. Comparison of thermodynamic prediction obtainedby the MINEQL+ program with experimental results obtainedby using ferrous salts to remove sulfide showed that the

product of precipitation is very likely to be FeS under all theoperational conditions tested. With respect to the reactionbetween ferric salts and sulfide, analysis of thermodynamicdata (Pourbaix diagram) suggests that the dominant productof sulfide oxidation under typical pe/pH conditions prevailingin municipal raw wastewater is SO4

2−. However, comparisonbetween sulfide removal efficiency in laboratory experimentsconducted with multiple (different) samples of raw municipalsewage (with both ferric salts and a mixture of ferric andferrous salts) and the prediction of the MINEQL+ programshowed that the main sulfide oxidation product is in factelemental sulfur.

No evidence was found for the existence of other iron andsulfide precipitation products (apart from FeS) that weresuggested in literature.

The results obtained in this study are significant from thepractical and economic standpoints since it was demon-strated that when sulfide containing wastewater is dosedwithferric salts the main product is elemental sulfur and notsulfate, therefore the amount of ferric salts required forachieving a desired sulfide target in the aqueous phase ismuch lower than in the case that the main oxidation producthad been sulfate.

The results show that in order to reduce sulfide toconcentrations lower than around 0.1 mg S/l with a safetyfactor amolar ratio higher than around 1.3 (Fe) to 1 (S2−) shouldbe applied when ferrous salts are used, and N0.9 to 1 whenferric salts and a mixture of salts (at a 2 Fe(III) to 1 Fe(II) ratio)are applied. The decision between the alternatives should betaken based on local cost of chemicals, however, in manyplaces this cost (per Fe unit mass) is almost identical andtherefore the use of ferric salts seems to bemore cost effective.The results also suggest that the very high ratio between Feand S(-II) that appears in certain recommendations in theliterature may not be justified from a chemical standpoint. Itwould appear that a better control on the in-line sulfideconcentrations in the wastewater and sewage dischargewould allow minimizing the iron salts dosage and wouldenable high sulfide removal efficiency at a cost effectivefashion. Such control can be attained, for example, by usingsulfide-specific electrodes available on the market or byapplying polarographic H2S gas sensors to which samplesare pumped from the sewer at predetermined time intervals(see for example http://www.analyticaltechnology.com).

Acknowledgements

The financial support of The Grand Water Research Institute,Technion and The Pinhas Sapir Fund for Applied Research inthe Municipal Government (Pais – Israel National Lottery) isgreatly acknowledged.

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