Chapter 14: Aquatic Life and Oxidation/Reduction Previekuwata/classes/2009-10/Env Chem... · 1 Chapter 14: Aquatic Life and Oxidation/Reduction Preview We now turn to water as the

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Chapter 14: Aquatic Life and Oxidation/Reduction

Preview

We now turn to water as the medium that supports life. All organisms require water, and a large fraction of them make their home in rivers, lakes, and the oceans. Life started in the ocean and occupied dry land only later. Moreover, biological processes have a profound influence on the chemistry of natural waters, and indeed of the entire globe. Were it not for the evolution of photosynthetic organisms, first in the ocean, and then on land, the atmosphere would be devoid of oxygen. The profound influence of oxygen on the chemistry of the atmosphere was considered at length in Part II. O2 is also the dominant actor in the chemistry and biochemistry of the hydrosphere. The limited availability of O2 in water sets the boundary between aerobic and anaerobic life, with crucial consequences for water quality and the health of ecosystems. In this chapter we consider

• Redox energy and dissolved oxygen • Biological redox and the reduction potential • Linkage of redox with acid/base chemistry • Earth’s redox evolution • Biological CO2 pump • Overfertilization of surface waters: eutrophication • Redox and metal pollution • Ocean fertilization with iron

14.1 Redox Reactions and Energy Life is powered by redox reactions, chemical processes in which electrons are

transferred from one molecule to another, with the release of energy. Organisms have evolved machinery, made up of proteins and membranes, which channels this energy into the biochemical pathways that support vital functions.

In an aerobic environment, the most important biological redox process is respiraton,

(CH2O)+O2=CO2+H2O [14‐1] which we encountered previously as part of the global carbon cycle [p.??]. In this case carbohydrate molecules provide electrons for the reduction of dioxygen. All higher life forms obtain their energy via respiration. However, many other redox processes are utilized by bacteria. Indeed, bacteria have evolved to exploit just about any redox process that is available in nature. Anyplace where a supply of oxidizable molecules coexist with molecules capable of oxidizing them, it is a good bet that bacteria are present which can utilize the potential redox reaction. The oxidation of FeS2 by thiobacillus ferrooxidans in the above discussion of acid mine drainage is a good example.

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14.2 Biological oxygen demand Wherever oxygen is present, respiration provides life-supporting redox energy,

but in liquid water oxygen can easily become depleted. The solubility of O2 in water is only 9 mg/L (about 0.3 millimoles) at 20 °C, and less at higher temperatures. The oxygen supply can be replenished by contact with the air, as in rapidly flowing streams. But in standing water or in waterlogged soils, the diffusion of oxygen from the atmosphere is slow relative to the speed of microbial metabolism, and the oxygen is used up.

Given the centrality of oxygen to metabolism, a parameter called biological oxygen demand (BOD) has been defined to measure the reducing power of water containing organic carbon. BOD is the number of milligrams of O2 required to carry out the oxidation of organic carbon in one liter of water. Values for various industrial wastes and municipal sewage are given in Table 14.1.

WorkedProblem14.1:BODQ:WhatistheBODofwaterinwhich10mgofsugar[empiricalformula

CH2O]isdissolvedinaliter?HowdoesthiscomparewiththeO2solubilityat20oC?A:SinceeachmoleofCH2OrequiresonemoleofO2[equation14.1],we

divide10mgbythemolecularweightofCH2O[30g],toobtaintherequirednumberofmolesofO2[32g/mol],andthenmultiplybythemolecularweightofO2toobtainthenumberofmg:

BOD=10mgx32g/30g=10.7mg/lThisexceedstheO2solubility[9mg/l]byabout20%.

14S.1OxidationLevelsandWaterManyelementscanexistinmultipleoxidationstates,dependingonthe

numberofelectronsaddedtoorremovedfromthevalenceshelloftheatoms.Inanaqueousworld,thestabilityofthesedifferentoxidationstatesdependsonthepropertiesofwater.ThuswearefamiliarwithNa+,andMg2+ions,becausesodiumandmagnesiumhaveoneandtwoelectrons,respectively,intheirvalenceshells,whichareeasilyremovedwhenwatermoleculesareavailabletostabilizethe

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resultingions[Figure12.5].Allmetalsformpositiveionsinwater,andinthecaseoftransitionmetalsmultipleoxidationstatesareavailable;forexample,ironcanexistinwaterasFe3+orFe2+.

Nonmetals,beingelectronegativeelements,readilyattainnegativeoxidationlevels,dependingonthenumberofelectronsthattheirvalenceshellscanaccommodate.ThusthelowestoxidationlevelsattainablebyF,O,NandCare–I,‐II,‐III,‐IVrespectively;weuseRomannumeralstodenotetheoxidationnumbertodistinguishthemfromtheactualcharge.Thus,althoughCl‐ionsexistassuchinwater,O2‐ionsdonot.TheirprotonaffinitiesarehighenoughthattheyarecompletelyconvertedtoOH‐[ortoH2O,dependingonthepH].ThelowestoxidationlevelsforNandC,arerepresentedbyNH3[orNH4+]andCH4.

Positiveoxidationlevelsarealsoaccessibletothenon‐metalsbecauseofthestabilizationavailablethroughbondingtooxideions.ThusC,N,SandClareintheirmaximumoxidationstates,+IV,+V,+VIand+VIIrespectivelywhensurroundedbyoxide:CO2[orCO32‐],NO32‐,SO42‐,andClO4‐.Theactualchargesonthecentralatomsinthesemoleculesaremuchlessthan+4,+5,+6or+7,sinceelectronsaresharedinthepolarbutcovalentbondswiththeOatoms.Nevertheless,theoxidationstateiscrucialindeterminingthepossibilitiesforredoxchemistry.Forexample,eightelectronsmustberemovedfromNinordertoconvertNH3toNO32‐.Inthecaseoftherespirationreaction,[14.1],carbonin(CH2O)isintheoxidationstateO[therulesarethatOcountsfor–2,andHcountsfor+1indeterminingthe‘effective’charge,i.e.theoxidationstate,oftheremainingatoms];fourelectronsaretransferredtoO2inconverting(CH2O)toCO2.

WorkedProblem14.2:CalculatingtheOxidationStateandBalancingRedoxEquationsQ:WhatistheoxidationstateofNinthenitriteion,NO2‐?A:SinceOcountsas–2,andthereisanoverall–1charge,Nmusthavean

effectivechargeof+3.TheoxidationlevelisIII.

Q:WriteabalancedchemicalequationforthereductionofNO2‐toNH3byH2.A:Firstbalancethenumberofelectronstransferredfromoxidantto

reductant.SinceNchangesfromIIIto–III,sixelectronsaretransferred.HchangesfromOtoI,sosixHatoms,orthreeH2molecules,arerequiredtoreceivetheelectrons.

NO2‐+3H2=NH3Sincethereactionisinwater,itispermissibletoaddH2OorH+orOH‐to

eithersideofthereaction,asneeded.SeeingthatnitritehadtwoOatoms,webalancethesebyaddingtwowatermoleculestotherighthandside.

NO2‐+3H2=NH3+2H2O ThetotalHcountontherighthandsideisnowseven,whichwebalancebyaddingoneH+tothelefthandside.Thisalsobalancesthecharge.

NO2‐+3H2+H+=NH3+2H2O

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14.3 Natural sequence of biological reduction When water is depleted of oxygen, organisms that depend upon aerobic

respiration cannot survive, and anaerobic bacteria take over. These bacteria utilize oxidants other than O2. These alternative oxidants are less powerful than O2, and cannot produce as much energy. Nevertheless, bacteria are quite capable of surviving on lower energy processes; in doing so, they can fill ecological niches that are not available to aerobic organisms. The oxidizing power of anaerobic environments in the biosphere is mainly controlled by five molecules. In decreasing order of energy produced, they are nitrate (NO3

–), manganese dioxide (MnO2), ferric hydroxide (Fe(OH)3), sulfate (SO42–) and,

under extreme conditions, carbohydrate (CH2O) itself. The biological oxidation processes supported by these oxidants are described in Table 14.2. Microbial populations first use the oxidant that produces the most energy until it is depleted; only then does another agent become the dominant oxidant. Table14.2RedoxReactions,Products,andConsequences

Redoxreaction Reactionproducts/consequences1.O2+CH2O→CO2+H2O Theaerobiccondition,characterizedbythehighest

redoxpotential,occurswhenthereisanabundanceofO2,andtherelativeabsenceoforganicmatterowingtooxicsewagewastes,andthedecompositionoforganicmatternearthesurfaceofwell‐aeratedsoils.Theendproducts,CO2andwater,arenontoxic.

2.4NO3‐+5CH2O+4H+→5CO2+2N2+7H2O

Whenmolecularoxygenisdepletedfromthesoilorwatercolumn,aswouldbethecase,forexample,inwaterloggedsoilsandwetlands,availablenitrateisthemostefficientoxidant.DenitrifyingbacteriaconsumenitrateandreleaseN2.N2O,agreenhousegas,isalsoreleasedasaside‐product.Inagriculturalsoils,denitrificationcanleadtolossesofnitrogenfertilizeramountingtoasmuchas20%ofinputs.Denitrifyingbacteriaarealsoveryactiveinheavilypollutedriversorinstratifiedestuarieswhereorganicmatteraccumulates.Insomeestuarysystems,denitrificationmaysignificantlyaffectthetransferofnitrogentotheadjacentcoastalwatersandatmosphere.

3a.2MnO2+CH2O+4H+→2Mn2++3H2O+CO2

3b.4Fe(OH)3+CH2O+8H+→4Fe2++11H2O+CO2

Inaerobicenvironmentswherenitratesareinlowconcentrationandmanganeseandferricoxidesareabundant,themetaloxidesareasourceofoxidantformicrobialoxidation.Thismaybethecaseinnaturalsoils,andinthesedimentsoflakesandrivers.Theenvironmentalsignificanceofthesemetaloxidesisthattheyserveadualrole.Notonlyaretheyasourceofoxidantstomicroorganisms,theyarealsoimportantfortheircapacitytobind

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toxicheavymetals,deleteriousorganiccompounds,phosphates,andgases.Whenthemetaloxidesarereduced,theybecomewater‐solubleandlosetheirbindingability.Thislossmayresultinthereleaseoftoxicmaterials.

4a.½SO42‐+CH2O+H+→½H2S+H2O+CO2

SulfidicconditionsarebroughtaboutalmostentirelybythebacterialreductionofsulfatetoH2SandHS‐accompanyingorganicmatterdecomposition.Sulfatereductionisverycommoninmarinesedimentsbecauseoftheubiquityoforganicmatterandtheabundanceofdissolvedsulfateinseawater.Infreshwater,suchreactionsareimportantinareasaffectedbyacidicdepositionintheformofsulfuricacid.H2Sisanextremelytoxicgas.Sulfidesarealsoimportantinscavengingheavymetalsinbottomsediments.

4b.MS2+7O2+H2O→M2++2SO42‐+2H+

Conversionofaheavy‐metalsulfide(MS2)tosulfatemayalsooccurwhenanaerobicsedimentsareexposedtotheatmosphere,asinthecaseoftheraisingofdredgesoils.Itmayalsooccurwhenwetlandscontainingpyrites(FeS2)aredrainedforagricultureorincoal‐miningareasasacidminddrainage.Oneconsequencemaybeanincreaseinacidificationfromthegenerationofsulfuricacid;anothermightbethereleaseoftoxicmetals.

5.CH2O+CH2O→CH4+CO2

Underanaerobicconditionsataredoxpotentialofabout‐200mV,andinthepresenceofmethogenicbacteriaasmaybefoundinswamps,floodedareas,ricepaddies,andthesedimentsofenclosedbaysandlakes,partiallyreducedcarboncompoundscandisproportionatelyproducemethaneaswellasCO2.Thisreactionismoretypicalinfreshwatersystemsbecausesulfateconcentrationsaremuchlowerthaninmarineenvironments,averagingaboutoneone‐hundreththeconcentrationinseawater.Methaneisacriticalgasinthedeterminationofglobalclimate.Sincetheearly1970s,globalatmosphericmethanelevelshavebeenincreasingatarateof1%peryear.Althoughthereasonsforthisincreasearestillunderinvestigation,theexpansionofricepaddycultivationinsoutheastAsiahasbeencitedasacontributingcause.Seechapter3,pp.?.

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Source:W.M.Stigliani(1988).Changesinvaluedcapacitiesofsoilsandsedimentsasindicatorsofnonlinearandtime‐delayedenvironmentaleffects.EnvironmentalMonitoringandAssessment10:245‐307.

The oxidizing power of a molecule depends on the specific reaction being carried out, and is measured as the reduction potential associated with the reduction of the oxidant. These are listed in Table 14.3 for the environmental oxidants we are considering. Microbial populations first use the oxidant that produces the most energy until it is depleted; only then does another agent come the dominant oxidant. Thus, the redox potential of a body of water tends to fall in a stepwise pattern as BOD increases (Figure 14.1).

As oxidants are consumed in the conversion of reduced carbon to CO2, the

reduction potential falls to successively lower plateaus, corresponding to the successively lower potential redox couples O2/H2O, NO3

–/N2, MnO2/Mn2+, Fe(OH)3/Fe2+, SO4

2–/HS–, and CO2/CH4 . These couples do not give reversible potentials at electrodes, but the metabolic activity of the vast array of microbes in soils and in water ensure that electron transfer does occur on a time-scale of hours or days (Table 14.3). Consequently, all redox-active materials respond to the reduction potential established by the microbial activity.

Note, however, that, while there is a general correspondence with the Eh values of the half-reactions, the plateau potentials in Figure 14.1 deviate substantially from the numbers listed in Table 14.3. This is because conditions in the environment are far from the standard conditions which establish the Eh values. While the pH may be close to 7, the concentrations of other reactants and products are unlikely to be 1.0 M [or 1 atm, in the case of a gas].

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Figure14.1Sequenceofredoxreactionsinaqueousenvironments.O2innaturalwatersat20oCissufficienttooxidizeabout3.4mgoforganiccarbon(shownhereasCH2O)perliterofwater.WhentherateofreplenishmentofO2fromtheatmosphereisslowerthantherateofoxidationofCH2O,oxygenisdepletedandmicrobeswillselectthenextmostenergeticoxidantinthesequenceshown.Forsimplicity,onlymajorproductsandtheirvalencestatesareshown.SeeTable14.2forbalancedequations.Source:W.M.Stigliani(1988).Changesinvaluedcapacitiesofsoilsandsedimentsasindicatorsofnonlinearandtime‐delayedenvironmentaleffects.EnvironmentalMonitoringandAssessment10:245‐307.

14S.2ReductionPotentialsAllredoxreactionscanbedivided,atleastconceptually,intotworeduction

half­reactions,oneproceedingforwardandtheotherinreverse.Forexample,theoxidationofhydrogenbyoxygen,

2H2+O2=2H2O [14‐2]canbedividedinto O2+4e–+4H+=2H2O [14‐3]and 4H++4e–=2H2 [14‐4]

Subtractinghalf‐reaction[14‐4]from[14‐3]givestheoverallreaction[14‐2]. Thesehalf‐reactionscanactuallybecarriedoutattheelectrodesofahydrogen‐oxygenfuelcell,asdiscussedinChapter10(pp.??).Apotentialdifferenceisdevelopedbetweentheoxygenelectrodeandthehydrogenelectrode,allowingacurrenttoflowthroughtheexternalcircuit.Forthehydrogen‐oxygenfuelcell,thispotentialdifferenceapproaches1.24volts(V),atthestandardtemperatureof25oC,whenthegasesareat1atmospherepressure,andtheelectrodesbehavereversibly,

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thatis,whenthereactantsandproductsareatequilibriumwiththeelectrodes(implyingrapidelectrontransferrates).

Thepotentialdifference,ΔE,istheenergyoftheelectrochemicalcellperunitofchargedelivered.[Specifically1V=1J/C,whereV=volt,J=jouleandC(coulomb)istheunitofcharge].ΔEisrelatedtothefreeenergyofthecellreactionbytherelation

ΔG=–nFΔE [I4‐5]

whereF(theFaraday)istheamountofchargeinamoleofelectrons,[96,500C]andnisthenumberofelectronstransferredinthereaction.Thusinreaction[I4‐2], 4electronsaretransferredfrom2H2toO2,andΔG=–4e‐x96,500C/mole‐x1.24J/C=–479,000J,or–479kJ.[Recallthatthisvalue,incombinationwiththeentropyofthereactiongivesatheoreticalenergyconversionefficiencyof80%fortheH2/O2fuelcell–P.?]

Numerouselectrodecombinationsarepossibleinelectrochemicalcells,anditisconvenienttospecifyastandardpotential,Eo,foreachelectrodebyreferencingittothehydrogenelectrode,whosestandardpotentialisdefinedaszero.ThusEo=1.24Vfortheoxygenelectrode,representedbyhalf‐reaction[I4‐3].ThestandardconditionsforEoareunitactivities(partialpressureormolarconcentration)ofthereactantsandproducts,at25°C.

Therearemanyhalf‐reactionswhoseelectrodepotentialcannotactuallybemeasured,becausetheelectrontransferreactionatanelectrodeistooslow.Thesepotentialscanneverthelessbecalculatedfromthefreeenergyofappropriateredoxreactions.Forexample,theformationofNOfromN2andO2,whosethermodynamicswasconsideredinChapter4(p.??),isaredoxreaction:

O2+N2=2NO [14‐6]whichcanbedividedintothehalf‐reactions

O2+4e–+4H+=2H2O [14‐7]and 2NO+4e–+4H+=N2+2H2O [14‐8]

Fromthefreeenergyoftheoverallreaction(p.?),173.4kJ,weobtainacellpotentialof–0.45V(usingequation[14‐5]).Then,knowingthatthestandardpotentialoftheoxygenelectrodeis1.24V,wecanreadilycalculatethatthestandardpotentialforhalf‐reaction[14‐8]is1.69V(1.24V–[‐0.45V]),eventhoughitisimpossibletomeasurethispotentialdirectlybecausetheelectrontransferbetweentheelectrodeandtheNOandN2moleculesistooslowtoestablishareversiblepotential.

14S.3ConcentrationDependenceofthePotential:pHandE0[w]Whathappenstothereductionpotentialwhenconditionsarenotstandard?

Asinallchemicalreactions,thedrivingforceforelectrochemicalprocessesdependsontheconcentrationsofreactantsandproducts.ThisdependenceisgivenbytheNernstequation:

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E=Eo–[RT/nF]lnQ [14‐9] whereEoisthestandardpotential,Risthegasconstant,nisthenumberofelectronstransferredinthereaction,andQistheequilibriumquotient,i.e.theconcentrationexpressionfortheequilibriumconstant.Inthefuelcellreaction[14‐2],forexample,Q=1/PO2PH2

2(thewateractivitybeingdefinedasunity),andn=4.ThereforeE=1.24–[RT/4F]x[–lnPO2–2lnPH2]

AconvenientformoftheNernstequationis

E=Eo–[0.059/n]logQ [14‐10]

where0.059isthevalueofRT/Fat25°C,multipliedbytheconversionfactorfromnaturaltobase‐tenlogarithms[ln10=2.303].Fortemperaturesotherthan25°C,thefactor0.059mustberaisedorloweredaccordingly.

TheNernstequationappliesequallytowholecellreactionsorhalf‐reactions.Thusthepotentialofthehydrogenelectrode(half‐reaction[14‐4])at25°Cis(afterdividingthroughbyn=4]

E=0–0.059{logPH21/2/[H+]} [14‐11]Fromthisweseethatthehydrogenelectrodepotentialbecomesmore

negativeas[H+]diminishes.ThusH2gasisamorepowerfulreductantinalkalinesolutionthaninacid.Efallsby–0.059VforeveryunitriseinpH.AtpH7,thehydrogenelectrodepotentialis–0.42[whenallotherconditionsarestandard].

LikewiseO2isalesspowerfuloxidantinalkalithaninacid,becauseprotonsareconsumedinthereductionhalf‐reaction,[14‐3].Theoxygenpotential(againafterdividingbyn=4)is

E=1.24–0.059log{1/PO21/4[H+]} [14‐12]Againthepotentialdrops0.059VforeveryunitriseinpHandis0.82VatpH

7.BecausepH7isclosertomostbiologicallyandenvironmentallyrelevantconditionsthanispH0,electrodepotentialsareoftencitedforpH7,astheyareinTable14.3.TheEo[w]valuesareEovaluesrecalculatedforpH7.

Evenifnoprotonsappearexplicitlyinahalf‐reaction,thepotentialmaybepH‐dependentbecauseofsecondaryacid‐basereactions.Forexample,thepotentialoftheFe3 + reductionhalf‐reaction

Fe3++e‐=Fe2+ [14‐13]

hasnoprotondependenceperse,buttheequilibriumquotient,[Fe2+]/[Fe3+],ishighlydependentonpHbecauseoftheacidiccharacterofFe3+.AtquitelowpH,itformsaseriesofhydroxidecomplexes,andprecipitatesasthehighlyinsolubleFe(OH)3(Ksp=10‐37).Incontrast,Fe2+formshydroxidecomplexesonlyathighpH,andFe(OH)2(Ksp=1015)ismoresolublethanFe(OH)3.Consequently,thereductionpotentialfallswithincreasingpH,because[Fe3+]declinesmorerapidlythan[Fe2+].

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WorkedProblem14.3:Eo[w]andKspofFe[OH]3Q:TheFe3+/2+standardpotential[equation[14‐13]is0.77V.Fromthis

valueandtheKspcalculateEo[w]forthereductionofFe[OH]3toFe2+[seeTable14.3].

A:Eo[w]istheFe3+/2+potentialatpH7,whenFe[OH]3iscertainlyprecipitated.ThispotentialcanbecalculatedfromtheNernstequation

E=0.77–0.059{log([Fe2+]/[Fe3+])}and[Fe3+]canbecalculatedfromKsp=[Fe3+][OH‐]3.SubstitutiongivesE=0.77–0.059{log[Fe2+]‐log(Ksp)+3log[OH‐]}AtpH7,[OH‐]=1.0x10‐7ME=0.77‐0.059{log[Fe2+]+37‐21}=‐0.17–0.059{log[Fe2+]}whichistheNernstequationforFe[OH]3reduction,withEo[w]=‐0.17V.

WorkedProblem14.4:EffectiveOxygenPotentialQ.ThefirstplateauinFigure14.1,correspondingtoO2reduction,isat0.5V,

whereastheEo[w]value[Table14.3]is0.812V.Whatmightaccountforthisdifference?

A:AssumingthattheenvironmentalpHis7,thedifferencemustarisefromtheO2concentrationdependence.ThepotentialdiminisheswithdecreasingO2concentration.Recallthat

E=1.24–0.059log{1/PO21/4[H+]} [14‐14]or,atpH7,E=0.812–0.059log{1/PO21/4} [14‐15]IfE=0.50,thensubstitutingintoequation[14‐15]gives:log{1/PO21/4}=[‐logPO2]/4=‐0.312/(‐0.059)=5.28orPO2=10‐21.1atm.Thismayseemabizarrelylowvalue,butitreflectsthe

factthatwhenmicrobesareactivelyrespiringinanaqueousmedium,theydrawdowntheO2toverylowlevelsintheirimmediatevicinity.

14S.4ElectronandProtonAffinitiesAreLinked:pEversuspH Mostreductionreactionsareaccompaniedbyprotonuptake,and

oxidationsgenerallyleadtoprotonrelease.Sinceaddinganelectronincreasesnegativechargewhileaddingaprotondecreasesit,thecouplingofelectronandprotontransfersisasimpleconsequenceofthetendencytolowertheenergyofthemoleculebyneutralizingcharge.ThiscouplingleadstoastrongdependenceonthesolutionpHformosthalf‐reactionpotentials.

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ThehydrogenelectrodepotentialispH‐dependentbecausetwoprotons

combinewithtwoelectronsinproducingH2.Fromreaction[14‐4]:pE=0–[1/2]logPH2+log[H+][14‐16]

IfpH2ismaintainedat1atm,thenpE=–pH.Thus,hydrogengasisfarmore

reducinginalkalinethaninacidsolution.Forexample,atpH8,pE=–8V,andthehydrogenelectrodepotentialis–0.47V,nearlyhalfavoltmorenegative(reducing)thanatpH=0.Atthesametime,oxygenishardertoreduceinalkalinesolution(or,conversely,O2ismoreoxidizing).Fromreaction[14‐3]:

pE=pEo+[1/4]logPO2+log[H+]=20.75–pH,atPO2=1atm[4‐17]

andtheelectrodepotentialatpH7is0.83V.BecausepH7isclosertomostbiologicallyandenvironmentallyrelevantconditionsthanispH0,electrodepotentialsandpEvaluesareoftencitedforpH7,asinTable14.3.ThepEo[w]valuesarepEovaluesrecalculatedforpH7.

SinceboththehydrogenandoxygenelectrodeshavethesamepHdependence,thedifferencebetweenthemispH‐independent,reflectingthefactthatthereisnogainorlossofprotonsintheoverallreactionforhydrogenoxidationbyoxygen[reaction[14.2]].Thus,thepotentialofthehydrogen‐oxygenfuelcellisindependentofthepHofthecellcompartments,eventhoughtheindividualelectrodepotentialsarestronglyaffected.

Evenifnoprotonsappearexplicitlyinahalf‐reaction,thepotentialmaybepH‐dependentbecauseofsecondaryacid‐basereactions.Forexample,thepotentialoftheFe3+reductionhalf‐reaction[14‐13]potentialhasnoprotondependenceperse,buttheequilibriumquotient[Fe2+]/[Fe3+]ishighlydependentonpHbecauseoftheacidiccharacterofFe3+.AtquitelowpH,itformsaseriesofhydroxidecomplexes,andprecipitatesasthehighlyinsolubleFe(OH)3(pKsp=38).Incontrast,Fe2+formshydroxidecomplexesonlyathighpH,andFe(OH)2(pKsp=15)ismoresolublethanFe(OH)3.Consequently,thereductionpotentialfallswithincreasingpH,because[Fe3+]declinesmorerapidlythan[Fe2+].

TherelationshipbetweenpEandpHisconvenientlyillustratedinadiagramsuchasthatshownfortheFe3+/2+coupleinFigure14.2.Theregionsofthediagramarelabeledaccordingtothedominantchemicalspeciespresent,andthelinesshowthepE/pHdependenceattheedgesofthesestabilityfields.Thus,thehorizontallineatthetopleftofthediagramrepresentspE=13.2,thevalueexpectedforanequimolarsolutionofFe3+andFe2+intheabsenceofhydroxidereactions.Fe3+predominatesabovethisline,whileFe2+predominatesbelowtheline.TheverticallineatpH=3.0arisesbecauseoftheprecipitationofFe(OH)3.ThishappenswhentheKspisexceeded,whichdependsonthepHandon[Fe3+].Forthepurposesofillustration,theironconcentrationwassetat10–5MindrawingFigure14.2.From

Ksp=10–38=[Fe3+][OH–]3[4‐18]

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wecalculate[OH–]={10–38/[Fe3+]}1/3=10–11,givingpH=3.AbovethispH,Fe(OH)3precipitates,and[Fe3+]declinesinconformitytotheKspandthepH.TheeffectofthisdeclineonpEisseeninthelineslopingdownwardfrompH=3.Thislinehasaslopeof3.0becauseofthethreehydroxideionsperironinFe(OH)3.Rearrangingequation[4‐18],wehave

log[Fe3+]=–38+3pOH[4‐19]andsincepOH=14–pH

andpE=13.2–log{[Fe2+]/[Fe3+]}[4‐20]thedependenceofpEonpHisgivenby

pE=13.2–log[Fe2+]+log[Fe3+]=22.2–3pH(with[Fe2+]=10–5M)Abovethisline,Fe2+isoxidizedandprecipitatesasFe(OH)3,whilebelowthe

lineFe(OH)3dissolvesbyreductiontoFe2+.

Figure14.2pE/pHdiagramforaFe‐O‐Hsystem.

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Thesecondverticalline,atpH=9.0,arisesfromtheprecipitationofFe(OH)2.From

Ksp=10–15=[Fe2+][OH–]2[4‐21]

wecalculate[OH–]={10–15/[Fe2+]}1/2=10–5,givingpH=9.AbovethispH,Fe(OH)2precipitatesand[Fe2+]declines.ThelineslopingdownwardabovethispHrepresentsthephaseboundarybetweenFe(OH)3andFe(OH)2.Itsslope,minusone,isthedifferencebetweenthetwohydroxidesofFe(OH)2andthethreehydroxidesofFe(OH)3.Rearrangingequation[4‐21],wehave

log[Fe2+]=–15+2pOH[4‐22]Whenbothhydroxidesarepresent,equations[4‐19]and[4‐22]canbe

substitutedintoequation[4‐20]toobtainthedependenceofpEonpH:pE=13.2–log[Fe2+]+log[Fe3+]=–9.8+pOH=4.2–pH.ThetopandbottomdiagonallinesinFigure14.2representthepE/pH‐

dependenceofthehydrogenandoxygenreductionreactions,equations[14‐3]and[14‐4].Theyrepresentthestabilitylimitsforaqueoussolutions.Belowthebottomdiagonalline,waterisreducedtohydrogen,whileabovetheupperdiagonalline,waterisoxidizedtooxygen.

Figure14.2isnotacompletediagramoftheFe3+/2+systembecausesoluble

complexesoftheionshavebeenomittedfromconsideration.Hydroxidecomplexes,alreadymentionedabove,arealwayspresentinaqueoussolutions,buttheypredominateonlyinnarrowregionsofpHanddonotgreatlyaffecttheappearanceofthediagram.Othercomplexingagentscanhavesignificanteffects.Thenaturallyoccurringanionschloride,carbonate,andphosphatebindFe3+andcanlower[Fe3+]andthereforepE,ascanorganicconstituentsofsoils,especiallythehumicacids,whichcaneitherbindtheFe3+tosoilparticlesorformsolublechelateswiththeFe3+.Despitethesecomplexities,Figure14.2presentsthemainfeaturesoftheFe3+/2+system,whichisdominatedbythespeciesFe2+andFe(OH)3.OvermostoftheavailablepHrange,3–9,thesearetheonlysignificantspecies.

14.4 Biological oxidations Bacteria also catalyze oxidation of reduced substances by molecular oxygen,

even though such reactions can occur spontaneously in an aerobic environment. Thus HS – oxidation to sulfate is catalyzed by sulfide oxidizers. These bacteria manage to extract energy from the HS–/SO4

2– and O2/H2O redox couples. Another important oxidation process is nitrification, the conversion of NH4

+ to nitrate ion. Since plants take up and utilize nitrogen mainly in the form of nitrate, this is a key reaction in nature, especially in connection with the use of ammonium salts in fertilizers [see p. ?]. The process actually occurs in two steps, ammonium to nitrite, NO2

–, and nitrite to nitrate: NH 4++2H2O=NO2–+8H++6e– [14‐23] NO2–+H2O=NO3–+2H++2e– [14‐24]

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These half reactions are catalyzed by two separate groups of bacteria,

Nitrosomonas and Nitrobacter, each utilizing the oxidizing power of O2 to extract energy from the process.

In summary, redox potential can be considered as a kind of chemical switch in the aqueous environment, one that determines the sequence by which oxidants are utilized by microorganisms. Changes in redox potential can have important consequences for environmental pollution (Table 14.2).

14.5 Aerobic Earth O2 was not always a constituent of the atmosphere; it arose from the evolution

of life itself. The primitive Earth had an atmosphere derived from outgassing of the minerals in the interior. Once the surface cooled sufficiently to condense water, and with it acidic gases like HCl and SO2, the main atmospheric constituents would have been N2 and CO2.

Life arose quite early in the Earth’s history; microfossils resembling modern cyanobacteria have been found in 3.5 billion year-old rocks. How life started is unknown, and remains one of the great scientific issues of our time. It is known that simple organic molecules are common in the universe, and are present in meteorites, which would have bombarded the young Earth. Laboratory experiments show that they could also have been formed from inorganic precursors when subjected to electric discharges from lightning, or to ultraviolet irradiation. The ultraviolet flux would have been intense, since, in the absence of an oxygen atmosphere, the Earth would have lacked an ozone shield. Many of the organic building blocks of organisms could have been produced in this way. Alternatively, the building blocks might have been formed on the surfaces of sulfide minerals under the high pressures and temperatures found in hydrothermal vents on the sea floor. [These vents are found in regions where the crustal plates are being formed through upwelling from the Earth’s mantle]. Recent experiments show that complex organic molecules can be formed in this way. Still a third possibility is that the building blocks came from outer space, since complex organic molecules are found in some meteorites. How the building blocks were assembled into the first self-replicating organisms remains an unanswered question, although many ingenious proposals have been put forward.

The first organisms were heterotrophic, assimilating organic compounds from their environment. Since there was no O2, they must have obtained their energy from redox reactions other than respiration, similar to the modern anaerobic processes discussed in the preceding section. The splitting of simple organic molecules, such as acetic acid:

CH3COOH=CH4+CO2 [14‐25]

may have been the first of such processes; this reaction still provides energy for modern acetogenic bacteria.

However, photosynthesis evolved quite early on, probably in the cyanobacteria mentioned above, which survive today as photosynthetic organisms in the oceans. Photosynthesis made these organisms autotrophic, capable of synthesizing their own organic molecules from CO2. They had a strong selective advantage over heterotrophs. In addition to the fossil evidence mentioned above, carbon isotope measurements on the fossil organic carbon show photosynthesis to be at least 3.5 billion years old. The fossil

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carbon is found to be depleted in the stable 13C isotope, relative to 12C, as a result of the slightly slower diffusion of 13CO2 and its slower rate of capture by the CO2-fixing enzyme ribulose bisphosphate carboxylase.

O2 was a byproduct of the rise of autotrophic organisms. Because of the reactivity of O2, it would have been a toxic byproduct; most anaerobes are very sensitive to O2, and cannot survive in an aerobic environment. However, O2 did not become a significant constituent of the atmosphere for a long time after the advent of photosynthesis, because it was first consumed by oxidizable elements in the ocean and in the Earth’s crust, particularly iron and sulfur. The early ocean would have had a high concentration of Fe2+, which is abundant in silicate minerals of the mantle, and is quite soluble, in contrast to Fe3+. Photosynthetic O2 would initially have been used up by reaction with Fe2+ to produce precipitates of Fe[OH]3. Indeed ferric oxide begins to be seen in sedimentary rock that is about 3.5 billion years old, occurring in banded iron formations, in which Fe2O3 is interbedded with siliceous sediment. These formations reach a peak occurrence in rock which is 2.5 to 3 billion years old.

Once the oceanic Fe2+ was used up, the accumulating O2 attacked oxidizable minerals on land, principally FeS2 [pyrite], producing Fe[OH]3 and H2SO4 [the same chemistry that still produces acid mine drainage (p. ?)]. Evidence for this transition is found in the occurrence of red beds, deposits of Fe2O3 found in geologic layers of terrestrial origin, starting about 2 billion years ago, after the last of the banded iron formations were formed.

Finally, when the rate of O2 production exceeded its rate of consumption by exposed oxidizable material, the O2 concentration in the atmosphere began to rise, permitting the evolution of respiring organisms. Fossil evidence of eukaryotic organisms has been found in rocks that are 1.3-2 billion years old. Eukaryotes [in contrast to the more primitive prokaryotes] have mitochondria, organelles which are specialized for respiration. Some eukaryotes can survive on O2 at only 1% of the present concentration, suggesting that this level was attained over 1 billion years ago. O2 production would have accelerated with the evolution of chloroplasts in the eukaryotes, organelles which are specialized for photosynthesis. The rising O2 was also accompanied by the production of stratospheric ozone, which permitted life to colonize the continents, freed from the destructive effects of UV radiation. Fossils of multicellular organisms have been found in sedimentary rocks that are 680 million years old, but the rise of green plants, and with them the modern O2 atmosphere, dates to 400 million years ago.

The time-line for the course of O2 production is shown schematically in Figure 14.3. The present atmospheric reservoir accounts for only about 2% of the estimated cumulative production of O2, the rest having been used up in the oxidation of minerals. Interestingly, the O2 concentration seems to have stayed at about 20% of the atmospheric gases over the last 400 million years; this constancy suggests some sort of feedback control. As with any reservoir (see p. ?), the amount of O2 reflects the balance between the rate of production and the rate of consumption. Over geologic time, O2 consumption results from exposure and weathering of reduced carbon-bearing rock; this rate is set largely by the earth’s tectonic movements. O2 production results from the burial of reduced carbon, whose rate depends [among other things] on the total biomass. The biomass is limited, at least in part, by forest fires, and it is possible that feedback control arises from the dependence of fires on the O2 concentration. It is

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known that fires cannot be maintained when the O2 concentration is less than 15%, while even wet organic matter burns freely at a concentration greater than 25%.*

-----

*J.E. Lovelock (1974). Gaia: A New Look at Life on Earth. Oxford University Press: Oxford, U.K. -----

Figure14.3CumulativehistoryofO2releasedbyphotosynthesisthroughgeologictime.Ofmorethan5.1x1022gofO2released,about98percentiscontainedinseawaterandsedimentaryrocks,beginningwiththeoccurrenceofbandedironformationsatleast3.5billionyearsago(bya).AlthoughO2wasreleasedtotheatmospherebeginningabout2.0bya,itwasconsumedinterrestrialweatheringprocessestoformredbeds,sothattheaccumulationofO2topresentlevelsintheatmospherewasdelayedto400mya.Source:W.H.Schlesinger(1997).Origins.Biogeochemistry:AnAnalysisofGlobalChange(2nded.)(SanDiego:AcademicPress).

If carbon burial has balanced O2 accumulation for the last 400 million years, what accounts for the rising O2 level starting 4 billion years ago? A much larger carbon burial rate sees unlikely. It has recently been suggested† that UV photolysis of CH4 could have provided the driving force. Methane production would have been much higher when O2 levels were low; methane-producing anaerobes would have been abundant, and the methane would have escaped to the atmosphere without oxidation. In the absence of the ozone UV shield, the methane would have been exposed to photons energetic enough to break the C-H bonds. At the top of the atmosphere, the light H atoms would have escaped Earth’s gravitational field, and would have been lost to space. This removal of oxidizable H atoms from the earth-atmosphere system would provide a mechanism for O2 accumulation.

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--- †D.C.Catlingetal.Biogenicmethane,hydrogenescape,andtheirreversibleoxidationofearlyEarth(2001).Science293:839‐843.‐‐‐

14.6 Water as Ecological Medium

14.6a The euphotic zone and the biological pump Biological productivity depends on primary producers, organisms that fix carbon

via photosynthesis, and provide the food for the animal food chain. In water, the primary produces are cyanobacteria, phytoplankton and algae. Because of their dependence on sunlight, they are limited to the region near the surface, where sunlight can penetrate. This is the euphotic zone. Its depth depends on the clarity of the water.

Most biological activity takes place in the euphotic zone. The primary producers are eaten by animals or decomposed by bacteria, in a continuing cycle of photosynthesis and respiration. However, because of gravity, some dead organisms fall below the euphotic zone. In the deeper layers bacterial decomposition continues and the waters are enriched in carbon and the other elements of life. Because of thermal stratification, there is little physical mixing between the warmer surface layer and the cold deep layer. Consequently, there is a kind of ‘biological pump’, which transfers carbon, nitrogen, phosphorus, sulfur, etc., from the surface to the deep layers and the sediments. Figure 14.4 shows the effect of biological production on the depth profiles of nitrate and iron, as well as oxygen, in the north Pacific. O2 is high at the surface and diminishes sharply over the first few hundred meters. Nitrate and iron are drawn down at the surface, due to uptake by organisms, but increase sharply with depth as the organisms are decomposed; below the surface layer their concentrations remain at elevated levels.

In the oceans, the biological pump is responsible for increasing the carbonate concentration of the deep layers with respect to the surface layers. This drawing down of carbonate from the surface increases the rate of transfer of CO2 from the atmosphere. This is an important contribution to the global carbon cycle. It has been calculated that the atmospheric CO2 level would double in the absence of the biological pump.

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Figure14.4VerticaldistributionofFe,NO3,andO2inthecentralNorthPacificOcean.Source:J.H.Martinetal.(1989).VERTEX:Phytoplankton/ironstudiesintheGulfofAlaska.DeepSeaResearch36:649‐680.

14.6b Eutrophication in freshwater lakes. Because the supply of oxygen is restricted, the species that inhabit an aquatic

ecosystem are in a dynamic balance, one that is easily disturbed by humans. In water, the O2 concentration falls with increasing distance from the air-water interface. Thus, aerated soils support oxygen-utilizing microbes as well as higher life forms, while deeper in the soil, in the saturated zone where the soil pores are filled with water, anaerobic bacteria dominate and utilize progressively lower E0[w] redox couples. Likewise in lakes, the sediments are generally oxygen-starved and rich in anaerobic microorganisms, while in the water column above, the O2 concentration increases towards the surface. The concentration of O2 at the surface is increased not only because the surface is in contact with air, but because the surface waters support the growth of vegetation and algae, which release O2 as a product of photosynthesis.

The biological productivity of a temperate lake varies annually in a cycle (Figures 14.5 and 14.6). The onset of winter diminishes the solar heating of the surface. The thermal stratification disappears and the water’s density becomes uniform, allowing easy mixing by wind and waves, which brings nutrient- rich waters to the surface. In winter, the nutrient supply is high, but productivity is inhibited by low temperatures and light

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levels. Spring brings sunlight and warming, leading to a bloom of phytoplankton and other water plants. As plant growth increases, the nutrient supply diminishes and phytoplankton activity falls. Bacteria decompose the dead plant matter, gradually replenishing the nutrient supply, and a secondary peak of phytoplankton activity is observed in the autumn. Because the nutrient supply is limited in unpolluted waters, the BOD in the surface waters rarely outstrips the available oxygen.

Figure14.5Seasonalcyclingofnutrientsinlakes.EZ=thermoclineandendoftheeuphoticzone;stipplerepresentsphytoplanktongrowth;N→signifiesdirectionofnutrientflow;enclosedarrowsindicatecirculationofwaters.Thesolidlineattherightisthetemperatureprofileofthewatercolumn.

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Figure14.6Seasonalphytoplanktonproductivityasafunctionofsunlightandnutrientconcentration.Source:AdaptedfromW.D.Russel‐Hunter(1970).AquaticProductivity(NewYork:MacmillanPublishingCo.,Inc.).

This natural cycle can be disrupted, however, by excessive nutrient loading from human sources such as wastewaters or agricultural runoff. The added nutrients can support a higher population of phytoplankton, producing “algal blooms” (Figure 14.7). When masses of algae die off, their decomposition can deplete the oxygen supply, killing fish and other life forms. If the oxygen supply is exhausted, the bacterial population may switch from predominantly aerobic bacteria to mainly anaerobic microorganisms that generate the noxious products (NH3, CH4, H2S) of anaerobic metabolism.

This process is called eutrophication or, more accurately, cultural eutrophication. Eutrophication is the natural process whereby lakes are gradually filled in. Over time, an initially clear (oligotrophic) lake eutrophies, filling with sediment and becoming a marsh, and then dry land. This process normally proceeds over thousands of years because biological growth and decomposition in the euphotic zone are closely balanced — the surface layers remain well oxygenated, and only a small fraction of biological production is deposited as sediment. When this balance is upset by overfertilization of the water, the eutrophication process accelerates greatly.

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Figure 14.7 False-color LandSat image of cyanobacteria surface blooms in lakes Mendota and Monona, Madison, Wisconsin. (Image courtesy NorthTemperate Lakes Long Term Ecological Research Program, http://lter.limnology.wisc.edu.) (From Carpenter, S.R. PNAS (2008)105,11039.)

14.6c Nitrogen and phosphorus: the limiting nutrients. The slow pace of natural eutrophication reflects the nutrient dynamics of an

aquatic ecosystem (Figure 14.8).The nutrients are assimilated from the environment by the primary producers, which serve as food for secondary producers, including fish. Dead plant and animal tissues are decomposed by bacteria, which restore the nutrients to the water. The growth of the primary producers is controlled by the limiting nutrient, the element that is least available in relation to its required abundance in the tissues. If the supply of the limiting nutrient increases through over fertilization, the water can produce algal blooms, but not otherwise; conversely, management of the aquatic ecosystem requires that the supply of the limiting nutrient be restricted.

The major nutrient elements are carbon, nitrogen, and phosphorus, which are required in the atomic ratios 106:16:1, reflecting the average composition of the molecules in biological tissues. Numerous other elements are also required, including sulfur, silicon, chlorine, iodine, and many metallic elements. Because the minor elements are required in small amounts, they can usually be supplied at adequate rates in natural waters. On the other hand, carbon, the element required in the largest amounts, is plentifully supplied to phytoplankton from CO2 in the atmosphere. Phytoplankton outrun the supply of CO2 only under conditions of very rapid growth such as in some algal blooms. In these cases, the pH of the water can be driven as high as 9 or 10 through the required shift of the carbonate equilibrium

HCO3– + H2O = OH- + CO2 [14-26]

The increase in pH can in turn alter the nature of the algal growth, selecting for varieties that are resistant to high pH.

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Figure14.8Nutrientcyclinginanaquaticecosystem.

Normally, the limiting nutrient element is either N or P. Although nitrogen makes up 80% of the atmosphere, it is unavailable except through the agency of N2-fixing bacteria, living in symbiotic association with certain species of plants. On land, these species are rare enough to make nitrogen the limiting nutrient under most conditions. In water, however, N2-fixing algal species are common, and nitrate ions are often abundant because of runoff from the land. Consequently, nitrogen is not usually limiting, although it may be in some regions, especially the oceans, where nitrate concentrations are low.

This leaves phosphorus as the element that is usually limiting to growth, at least in fresh water. A 37-year study at the Experimental Lakes Area in Canada discovered that phosphorous is the main cause of lake eutrophication.‡ In this study, the experimental lake was fertilized with constant inputs of phosphorus and decreasing amounts of nitrogen, and then during the last 16 years, phosphorous alone was added. Nitrogen-fixing cyanobacteria were able to provide the nitrogen inputs necessary from the atmosphere to allow biomass production in proportion to the phosphorus added to the lake. The lake was highly eutrophic, despite no additional inputs of nitrogen. ‡ D.W. Schindler et al. (2008) Eutrophication of lakes cannot be controlled by reducing nitrogen input: Results of a 37-year whole-ecosystem experiment. Proceedings of the National Academy of Sciences 105:11254-11258.

Phosphorus has no atmospheric supply because there is no naturally occurring

gaseous phosphorus compound. Moreover, the input of phosphorus in runoff from

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unfertilized lands is usually low because phosphate ions, having multiple negative charges, are bound strongly to mineral particles in soils. In surface waters, most of the phosphorus is contained in the plankton biomass; the phosphorus availability depends on recycling of the biomass by bacteria.

Some of the phosphorus is lost to the deeper water and to the sediments when dead organisms sink. When a lake turns over in winter, the phosphorus in the deep waters is carried to the surface and supports the plankton bloom in the spring. Whether this phosphorus is available to the surface waters depends on conditions in the lake. At the bottom, phosphate ions may be adsorbed onto particles of iron and manganese oxide. However, when the sediment becomes anoxic, the metal ions are reduced to the divalent forms, the oxides dissolve, and the phosphate ions are released into solution (see notes on maganese and iron oxides in Table 14.2). Phosphate solubility is also increased through acidification since at successively lower pH values, HPO4

2–, H2PO4–, and

H3PO4 are formed (p. ?). The most notorious instance of phosphate-induced eutrophication was in Lake

Erie, which “died” in the 1960s. Excessive algal growth and decay killed most of the fish and fouled the shoreline. A concerted effort by the United States and Canada to reduce phosphate inputs was put into effect in the 1970s. Over $8 billion was spent in building sewage treatment plants to remove phosphates from wastewater, and the levels of phosphate in detergents were restricted. These efforts, along with other pollution control measures, succeeded in bringing the lake back to life. Commercial fisheries have revived, and the beaches are once again in use.

14.6d Anoxia and coastal marine ʻdead zonesʼ Enhanced nutrient loading is also affecting many coastal areas, creating ‘dead

zones’ where marine life is curtailed by oxygen depletion. Over 400 dead zones have been identified, their global distribution corresponding roughly to the ‘human footprint’(Figure 14.9).

The progression of marine hypoxia is illustrated in Figure 14.10. As nutrient input increases, there is an initial pulse of energy up the food chain, but then a steady decrease as higher animals die off, and microbes take over. Because seawater is rich in sulfate salts, the favored reaction under anaerobic conditions is sulfate reduction to hydrogen sulfide (H2S), a chemical that is extremely toxic to fish and humans. Although H2S is generally confined to the lower layers of seawater, during storms the deeper, anoxic layers can mix with surface layers, exposing aquatic life to the deadly gas.

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Figure 14.9 Global distribution of 400-plus systems that have scientifically reportedaccountsofbeingeutrophication‐associateddeadzones.ThemapiscolorcodedaccordingtothenormalizedhumaninfluenceSource:R.J.DiazandR.Rosenberg(2008).SpreadingDeadZonesandConsequencesforMarineEcosystems,Science321:926‐929.

Figure14.10Energyflowinamarineecosystemaseutrophicationprogresses.Inhealthywaters(green–‘normoxia’)mobilepredatorsfeedontheorganismsthatliveontheseabed(benthicorganisms).Asoxygenisdepletedinthewater(orange‐‘hypoxia’),ashortpulseofenergyisfollowedbyadeclineinmobilepredators.Whennooxygenisleftinthewatercolumn(red–‘anoxia’),microbesprocessalloftheenergyandformH2S.Source:R.J.DiazandR.Rosenberg(2008).SpreadingDeadZonesandConsequencesforMarineEcosystems,Science321:926‐929.

In the U.S., the most notorious dead zone, in the Gulf of Mexico (Figure

14.11). In a vast area at the mouth of the Mississippi River (about 18,000 km2 , which has more than doubled since 1980) the O2 concentration is too low to support aquatic life during the spring and summer. These are the seasons of great algal blooms,

25

resulting from overfertilization of the Gulf by the nutrients in the river outflow. More than 40% of U.S. commercial fisheries are located in the Gulf of Mexico, and these have been hard hit by the annual appearance of the dead zone.

The Mississippi drains the vast mid-continent farmlands, and delivers 1.5 million tons of dissolved nitrogen annually to the Gulf, agriculture accounting for 80% of the total. However, there has been great controversy over whether N or P is the main culprit in producing the dead zone. Because of the transition from fresh to salt water, it is likely that both are important, and both need to be controlled.

Figure14.11The“deadzone”intheGulfofMexicoduetonutrientenrichmentinthedrainagebasinoftheMississippiRiver.

The complexities of nutrient enrichment are illustrated by the Chesepeake Bay, in the eastern U.S. Here eutrophication has long been evident, and appears to involve both N and P. Levels of these nutrients in the estuary rise and fall annually in a seasonal pattern(Figure 14.12). In winter, cold temperatures and lack of biochemical activity allow the concentration of O2 to reach its annual maximum. At the same time, nitrogen enters in large amounts because winter is the period of maximum freshwater flow, with accompanying transport of sediment and runoff. Simultaneously, sedimentation is removing phosphorus from the water column, mainly through the precipitation of manganese and iron oxides, which absorb phosphorus efficiently and are insoluble under aerobic conditions. (Phosphorus is also removed during the settling of organic debris.) Beginning in the late spring and early summer, the oxygen levels decline due to increased biological activity. Nitrogen concentrations also decline because 1) nitrogen is incorporated into biomass and sinks as the organisms die; 2) little new nitrogen is introduced in runoff; and 3) nitrogen is depleted as increasingly anoxic conditions force a switch from oxygen to nitrate as oxidant.

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Figure14.12Oxygenconcentrationinwateroverlyingthesedimentswithmajorseasonalnetfluxesofnitrogenandphosphorus(insets)inthePatuxentRiverattheestuaryofChesapeakeBay.Source:C.F.D’Elia(1987).Toomuchofagoodthing:NutrientenrichmentsoftheChesapeakeBay.Environment29(2):6‐11,30‐33.

The opposite situation prevails for phosphorus. Under anaerobic conditions, phosphorus is liberated from the sediments, in large part due to the reduction of manganese and ferric oxides to Mn2 + and Fe2+. In the 2+ valence states, the metals are soluble and release the bound phosphorus formerly adsorbed to the insoluble oxides of the metals. The phosphorus is readily mixed with the surface layers given the mechanical turbulence of estuarine environments. Thus, as conditions cycle from aerobic to anaerobic and back, the phosphorus is continuously recycled between the surface waters and the sediments. During anaerobic periods, phosphates are released to the water column to be taken up by microorganisms; during aerobic periods, phosphates are returned to the sediments. The amount of phosphate trapped in this cycle is vast, much greater than the annual quantities entering the estuaries from sewage effluents or other sources; it represents the cumulative inputs of many years. Thus, even though Maryland and Virginia banned detergents with phosphates in the 1980s, phytoplankton productivity is still excessive. Now the limiting nutrient may well be nitrogen, but nitrogen inputs are very difficult to control. Chesapeake Bay receives some of the highest atmospheric NOx emissions in the world, mainly due to the density of traffic in

27

the adjacent areas. Part of the strategy for cleaning up Chesapeake Bay might include reducing NOx from vehicle exhausts, demonstrating once again the link between the atmosphere and the hydrosphere.

14.6e Wetlands as chemical sinks. Wetlands are typically anoxic and have large amounts of organic carbon; they

create a natural buffer zone for nearby fresh or marine waters by trapping nitrates. The nitrates enter the wetlands in runoff, but are utilized by bacteria to oxidize stored carbon via the reduction of nitrate to N2 or N2O, which are vented to the atmosphere (Figure 14.13a). By depleting the nitrates before they can enter the estuary, the surrounding wetlands limit the excessive growth of biomass and subsequent anoxic conditions in the estuary. Restoration of wetlands has been proposed in many areas as a means of reducing overfertilization from runoff.

Figure14.13(a)Abilityofwetlandstobufferagainstnitrateandsulfateinputstowaterbodies;(b)underconditionswherewetlandsbecomedry,noneoftheprotectivereducingreactionsoccur.Inaddition,accumulatedsulfidesmayoxidizetosulfateassulfuricacid,andleachintoadjacentriversorlakes.Source:W.M.Stigliani(1988).Changesinvaluedcapacitiesofsoilsandsedimentsasindicatorsofnonlinearandtime‐delayedenvironmentaleffects.EnvironmentalMonitoringandAssessment10:245‐307.

If the original wetlands are of marine origin, they are likely to contain high concentrations of sulfur in the form of reduced sulfide minerals such as pyrite. Under the redox/pH conditions prevalent in wetlands, these sulfides are highly insoluble and

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immobilized (Figure 14.13a). Draining the wetlands (Figure 14.13b) exposes these compounds to oxidizing conditions, producing a situation similar to acid mine drainage (pp. ?).

One example of this phenomenon occurred in a coastal area of Sweden near the Gulf of Bothnia, where wetlands were drained in the early 1900s for use as agricultural lands. As shown in Figure 14.14, draining the wetland shifted the E[w]/pH conditions diagonally to the upper left, from the values typical of waterlogged soils to conditions close to those of acid mine drainage. The draining exposed sulfides to the atmosphere, and their oxidation to sulfuric acid acidified the soil and nearby lakes. The pH in one of these lakes, Lake Blamissusjon, dropped from 5.5 or higher in the last century to a current value of 3. Even though agricultural activities ceased in the 1960s, the lake has not recovered; it is widely known as the most acidic lake in Sweden.

Figure14.14Eh/pHasafunctionofdifferentaquaticenvironments.Ovalenclosedbydashedlineindicatesregionofhighestsolubilityofheavymetals.Source:AdaptedfromW.Salomons(1995).Long‐termstrategiesforhandlingcontaminatedsitesandlarge‐scaleareas.InBiogeodynamicsofPollutantsinSoilsandSediments,W.SalomonsandW.M.Stigliani,eds.(Berlin:Springer‐Verlag). Recently,itwasdiscoveredthatwetlandsactuallystoremorecarbonthandoesreforestedagriculturalland,about3000vs100gramsofcarbonpersquaremeterperyear.*WetlandscapturecarbonbyabsorbingCO2fromtheatmosphereintonewplantgrowth,butoncetheplantdies,thematerialiscoveredbywaterandmudthatslowsreactionwithO2andslowsdecomposition.Organicpeatsoilsformedfromthisprocesshavebeenfoundthatare60ft.deepand7,000‐10,000yearsold.Microbescanuseironoxides,sulfate,orCO2insteadofoxygentoformenergyfromredoxreactions(p.?),butwhenCO2isused,methaneisproduced.MethaneisamorepotentgreenhousegasthanCO2,andinfreshwatermarshestheamountofmethaneproducedcancelsoutanycoolingeffectsofCO2absorbedbythe

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plantmaterial.Insaltwatermarshes,though,theconcentrationsofsulfatearesohighthatmicrobesdonothavetouseCO2asanelectronacceptor,sonegligibleamountsofmethaneareproduced.CO2absorptionbysaltwatermarshescouldbeonemethodforreducingCO2concentrationsintheatmosphere.‐‐‐*J.Pelley(2008).Canwetlandrestorationcooltheplanet?EnvironmentalScience&Technology,8994.‐‐‐

14.6f Redox effects on metals pollution. Changes in the redox potential can have important consequences for

environmental pollution, especially with respect to metal ions such as cadmium, lead, and nickel. In general, the solubility of heavy metals is highest in oxidizing and acidic environments (Figure 14.14). At neutral to alkaline pHs in oxidizing environments, these metals often adsorb onto the surface of insoluble Fe(OH)3 and MnO2 particles, especially when phosphate is present to act as a bridging ion. When the redox potential shifts to only slightly oxidizing or slightly reducing conditions as a result of microbial action, and the pH shifts toward the acidic range, Fe(OH)3 and MnO2 in soils and sediments are reduced and solubilized. The adsorbed metal ions likewise become solubilized and move into groundwater (or into the water column of lakes when there is Fe(OH)3 or MnO2 in the sediment). Conversely, if sulfate is reduced microbially to HS– metal ions are immobilized as insoluble sulfides. But as we have seen, if sulfide rich sediments are exposed to air through drainage or dredging operations, then HS– is oxidized back to sulfate, and the heavy metal ions are released.

A particularly important instance of biological redox mediation of heavy-metal pollution occurs in the case of mercury. Inorganic mercury, in any of its common valence states, Hg0, Hg2

2+, and Hg2+, is not toxic when ingested; it tends to pass through the digestive system, although Hg0 is highly toxic when inhaled. But the methylmercury ion (CH3)Hg+ is very toxic, regardless of the route of exposure. The environmental route to toxicity involves sulfate reducing bacetria that live in anaerobic sediments. As pat of their metabolism these bacteria use methyl groups to produce acetate. When exposed to Hg2+ the bacteria transfer the methyl groups to the mercury, producing (CH3)Hg+; because methylmercury is soluble, it enters the aquatic food chain, where it is bio-accumulated in the protein-laden tissue of fish (see pp. ?).

14.6g Fertilizing the ocean with iron Although nitrogen and phosphorus are the limiting aquatic nutrients near land, it

has become evident that in large areas of open ocean, it is actually iron that limits biological production. Among the ‘trace metals’ essential for life, iron is required in largest amounts. Iron is utilized in many enzymes involved in electron transport, and in processing O2 and N2, as well as their reduction and (for N2) oxidation products. Thus all organisms require a steady supply of iron. Since iron is abundant in the Earth’s crust, iron limitation is not a problem for land plants, or for phytoplankton growing near land. However the concentration of iron in the ocean is extremely low (Table 13.1), because of the low solubility of Fe(OH)3 [p. ?] in the alkaline (pH = 8) seawater.

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In much of the oceans the settling of dust from the land provides phytoplankton with sufficient iron for growth. Prevailing winds blow sands from the Sahara and Gobi deserts far out over the Atlantic and Pacific oceans. Recent satellite measurements show a fairly good correlation between patterns of dust in the air and phytoplankton growth in the oceans below. However, there are large areas which are relatively dust-free, especially in the equatorial Pacific Ocean, and the waters ringing Antarctica at greater than 60o south latitude, called the Southern Ocean. These areas have less phytoplankton than could be supported by the available nitrogen and phosphorus. It has been known for some time that adding iron to samples of these waters stimulates phytoplankton growth in the laboratory, and a series of field experiments in the 1990’s showed that spreading iron over areas of nutrient-rich ocean produced phytoplankton blooms.

Iron limitation on biological productivity is an important ingredient in the carbon cycle, because phytoplankton take up CO2 and transport some of it to the deep ocean when they die. This is the mechanism of the ‘biological pump’ for CO2, discussed on p. ?. In iron-limited areas, adding iron to the oceans could increase the speed of the biological pump, drawing down the atmospheric CO2. Indeed it has been suggested that iron supplementation could offer a ‘geoengineering’ solution to the problem of rising atmospheric CO2. However, this solution has been set aside for several reasons:

a. The remedy would be very expensive, because the iron stimulation of phytoplankton blooms is a transient effect. The blooms quickly fade as the excess iron precipitates out of the photic zone. (The duration depends somewhat on the form of the added iron. Ferrous salts are soluble, but rapidly oxidize to insoluble Fe(OH)3. Ferric chelates are longer lived, but the chelating agents [see p. ?] would add to the expense). Consequently iron would have to be added continuously to have a permanent effect.

b. Modeling indicates that the maximum effect on the atmospheric CO2 concentration would be a ~60 ppm lowering, making a relatively small difference in the rising level.

c. There could be unforeseeable consequences to the biology of the oceans from such an intervention.

d. There would have to be international agreement on ocean alteration, particularly in the region of Antarctica, which is protected by international law.

However, the evidence that iron can fertilize the oceans, and that dust is an important source of iron, raises the possibility that changes in global dustiness may have contributed to the temperature changes that produced the ice Age. Data from ice cores and from deep sea sediments indicate that there was much more iron in ocean water during the ice ages. Thus the biological pump would have been stimulated; the ~60 ppm lowering in the CO2 level that might have been available from this mechanism corresponds approximately to the CO2 lowering which is also detected in ice cores [see p. ?]. The increased iron might have resulted from dust due to drying of the continents and expansion of deserts. However, as is usual in reconstructing the past, it is difficult to decide which factor is cause and which is effect.

Problems: 1.Calculatetheequilibriumpartialpressureofoxygeninawatersampleat

pH=7.0,whichcontainsequalconcentrationsofNH4+andNO3‐.[SeeTable3.9forreductionpotentials].

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2.WhatisthepEvalueofanacidminewatersamplehaving[Fe3+]=8.0x10‐3Mand[Fe2+]=4.0x10‐4M?

3.(a)Whatclassofmoleculesisresponsibleformostofthereducingpower

inaqueousenvironments?(b)Whatparameterisameasureofreducingpower?4. Fivehundredkgofn‐propanol(CH3CH2CH2OH)areaccidentally

dischargedintoabodyofwatercontaining108litersofH2O.ByhowmuchistheBOD(inmilligramsperliter)ofthiswaterincreased?Assumethefollowingreaction:

C3H8O+9/2O2 = 3CO2+4H2O5. Alakewithacross‐sectionalareaof1km2andadepthof50meters

hasaeuphoticzonethatextends15metersbelowthesurface.Whatisthemaximumweightofthebiomass(ingramsofcarbon)thatcanbedecomposedbyaerobicbacteriainthewatercolumnofthelakebelowtheeuphoticzoneduringthesummerwhenthereisnocirculationwiththeupperlayer?Thebacterialdecompositionreactionis:

(CH2O)n+nO2 = nCO2+nH2OThesolubilityofoxygeninpurewatersaturatedwithairat20°Cis8.9mg/l;1m3=1,000liters.

6. Assumethatalgaeneedcarbon,nitrogen,andphosphorusintheatomicratios106:16:1.Whatisthelimitingnutrientinalakethatcontainsthefollowingconcentrations:totalC=20mg/l,totalN=0.80mg/l,andtotalP=0.16mg/l?Ifitisknownthathalfthephosphorusinthelakeoriginatesfromtheuseofphosphatedetergents,willbanningphosphatebuildersslowdowneutrophication?

7. Namethesixmostimportantoxidantsintheaquaticenvironment,

andhowtheredoxpotentialregulatestheirreactivity.8.(a)IfalakecontainshighconcentrationsofdissolvedMn2+andFe2+,what

wouldbetheconcentrationofdissolvedNO3–andwhy?(b)WhatenvironmentaleffectmayaccompanyreductionofMnO2and

Fe(OH)3?9. Inanaerobicmarineenvironments,whattoxicgascanbegenerated

andbywhichreaction(namereactantsandproducts)?10. Explainthe“phosphatetrap”intheestuaryofChesapeakeBay.Why

wasalocalbanonphosphorusindetergentsnotparticularlyhelpfulinmitigatingeutrophicationintheestuary?

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11.(a)Explainwhyanaerobicfreshwaterwetlandswithhighconcentrationsoforganiccarboncanserveasnaturalbuffersagainstsulfatesandnitrogenoxides(givereactions).

(b)Whenotheroxidantsareabsentfromsuchwetlands,whichredoxreactionislikelytopredominate,andwhichproductswillbeemitted?

12. AnestuarinecreekinNewJerseycontainslargeamountsofmercury

boundassulfide(withK=10–52)undertheprevailingenvironmentalconditions(pH=6.8;Eh=–230mV).Environmentalscientistshavebeenaskedtoassessthepotentialimpactsofthepollutedsediments.Theyconcludethatthemercuryposesnodangerinitscurrentstate.However,theycautionagainstanyactionthatwouldexposeittoairandincreaseitsredoxpotential.Explainwhythescientistscometothisconclusion?