monitoring and assessment of surface water acidification following rewetting of oxidised acid...

Monitoring and assessment of surface water acidificationfollowing rewetting of oxidised acid sulfate soils

Luke M. Mosley & Benjamin Zammit & Ann-MarieJolley & Liz Barnett & Rob Fitzpatrick

Received: 14 March 2013 /Accepted: 10 July 2013# Springer Science+Business Media Dordrecht 2013

Abstract Large-scale exposure of acid sulfate soils dur-ing a hydrological drought in the Lower Lakes of SouthAustralia resulted in acidification of surface water inseveral locations. Our aim was to describe the techniquesused to monitor, assess and manage these acidificationevents using a field and laboratory dataset (n=1,208) ofacidic to circum-neutral pH water samples. The medianpH of the acidified (pH<6.5) samples was 3.8. Significant(p<0.05) increases in soluble metals (Al, Co, Mn, Ni andZn above guidelines for ecosystem protection), SO4

(from pyrite oxidation), Si (from aluminosilicate dissolu-tion) and Ca (from carbonate dissolution and limestoneaddition), were observed under the acidic conditions. Thelog of the soluble metal concentrations, acidity and

SO4/Cl ratio increased linearly with pH. The pH, alkalin-ity and acidity measurements were used to inform aeriallimestone dosing events to neutralise acidic water. Fieldmeasurements correlated strongly with laboratory mea-surements for pH, alkalinity and conductivity (r2≥0.97)but only moderately with acidity (r2=0.54), which couldbe due to difficulties in determining the indicator-basedfield titration endpoint. Laboratory measured acidity cor-related well with calculated acidity (r2=0.87, acidity pres-ent as AlIII>>H+≈MnII>FeII/III) but was about 20 %higher on average. Geochemical speciation calculationsand XRD measurements indicated that solid phase min-erals (schwertmannite and jarosite for Fe and jurbanite forAl) were likely controlling dissolved metal concentra-tions and influencing measured acidity between pH 2and 5.

Keywords Pyrite . Acid mine drainage .Metalgeochemistry . Secondary oxyhydroxysulfateminerals . Metal speciation . Acid neutralisation

Introduction

Many submerged aquatic sediments or subaqueous soilscontain reduced iron sulfide minerals such as pyrite, par-ticularly in many coastal and inland water systems wheresulfate and organic matter are plentiful. The exposure ofthese sediments and soils to air has occurred globally as aresult of natural and human driven change including landreclamation, drainage for agriculture, land uplift or hydro-logical drought (Dent and Pons 1995; Cook et al. 2000;

Environ Monit AssessDOI 10.1007/s10661-013-3350-9

Electronic supplementary material The online version of thisarticle (doi:10.1007/s10661-013-3350-9) containssupplementary material, which is available to authorized users.

L. M. Mosley : B. ZammitWater Quality Branch, Environment Protection Authority(South Australia),GPO Box 2607, Adelaide, SA 5001, Australia

A.<M. Jolley : L. BarnettDepartment for Environment Water and Natural Resources,Adelaide, Australia

L. M. Mosley (*) :R. FitzpatrickAcid Sulfate Soils Research Centre, University of Adelaide,Adelaide, Australiae-mail: [email protected]

R. FitzpatrickCSIRO Land and Water,Adelaide, Australia

http://dx.doi.org/10.1007/s10661-013-3350-9

Fitzpatrick et al. 2009; Boman et al. 2010). Under theseaerated conditions, the pyrite minerals in these submergedmaterials react to form sulfuric acid. Sediments and soilscontaining iron sulfides (sulfidic material) or the productsof sulfide oxidation (sulfuric material) are commonlycategorised as acid sulfate soils (Pons 1973; Dent andPons 1995). Rewetting of oxidised acid sulfate soils withsulfuric material can mobilise the acidity into receivingwater bodies, and severe water quality and ecosystemimpacts can result (Sammut et al. 1996; Amaral et al.2012; Nystrand and Österholm 2013).

In similarity with acid mine drainage (Kirby andCravotta 2005a, b; Hedin 2006), the total acidity indrainage water from acid sulfate soils can be predomi-nantly comprised of soluble metals (FeII/III, MnII andAlIII ions and complexes, Cook et al. 2000) althoughH+ acidity can also be dominant at some locations(Macdonald et al. 2007). The H+ and FeII/III acidity inthe first instance result from pyrite oxidation reactions,while MnII, AlIII and trace metals (e.g. As, Co, Ni, Zn )commonly are released by subsequent acid dissolution ofsilicate, carbonate and oxide minerals. Metal acidity inthe form of soluble AlIII, FeII/III andMnII can subsequent-ly release substantial H+ acidity by oxidation (for FeII

and MnII) and hydrolysis reactions such as those shownbelow (from Hedin 2006):

Feþ2 þ 1�4 O2 þ 5=2H2O→Fe OHð Þ3 þ 2Hþ ð1Þ

Feþ3 þ 3H2O→Fe OHð Þ3 þ 3Hþ ð2Þ

A1þ3 þ 3H2O→A1 OHð Þ3 þ 3Hþ ð3Þ

Mnþ2 þ 1�2O2 þ H2O→MnO2 þ 2Hþ ð4Þ

The establishment of hydrolysis equilibria is general-ly very fast and there is a general tendency for metals tooxidise, hydrolyse and precipitate as pH and redox po-tential increases (Stumm and Morgan 1996; Kirby andCravotta 2005a). Water that initially has near-neutral pH(6–7) and contains dissolved metals can contain bothalkalinity and acidity and ultimately could become acidic(pH<4.5) after oxidation, hydrolysis and precipitation ofmetal acidity (Cook et al. 2000; Kirby and Cravotta2005a). Secondary oxyhydroxysulfate minerals such asjarosite and schwertmannite can also exert a controlling

influence on dissolved metal concentrations when theyprecipitate in acidic conditions (Acero et al. 2006;Hammarstrom et al. 2005; Sullivan and Bush 2004;Burton et al. 2006). Hence, an understanding of metalgeochemistry is very important in the assessment andmanagement of acid sulfate soil impacts onwater quality.

From 2007 to 2009, Australia's largest river catch-ment system, the Murray–Darling, experienced recordlow flows due to rainfall reductions and water overallocation in the Murray–Darling Basin (MDBA2010). This led to a major decline in water levels ofthe Lower Lakes at the end of the river system in SouthAustralia and severe deterioration in water quality withincreased levels of salinity, nutrients and turbidity(Mosley et al. 2012). The extremely low water levelsalso led to large areas of lake sediments and subaqueoussoils being exposed for the first time in a long period(>>100 years) and these submerged lake materialscontained large amounts of sulfidic (i.e. pH>4) material,which then oxidised to sulfuric (pH<4) material(Fitzpatrick et al. 2009, 2010). When these exposedsulfuric sediments were rewet either from local rainfallor winter inflows, surface water acidification occurred inseveral regions of the lakes (Fig. 1).

It was important that accurate and targeted monitor-ing and assessment of water quality in the Lower Lakesoccurred to enable an effective understanding andmanagement of acidity impacts. While there have beensome studies of the water quality impacts of acid sul-fate soils, these have been focussed mainly on ground-water and drainage channel chemistry (Sammut et al.1996; Cook et al. 2000; Burton et al. 2006; MacDonaldet al. 2007) which may have limited applicability forassessing lake water acidification impacts. Simpsonet al. (2010) undertook laboratory studies of acid sul-fate soil rewetting with surface water and found strongrelationships between dissolved metal release and pH,but it was unclear if these results translate to fieldconditions. Geochemical speciation models would po-tentially be very useful for assessment of surface waterimpacts from acid sulfate soils, but these have beenapplied in only a few studies of drain and swamp water(Sullivan and Bush 2004) and river water (Nystrand

Fig. 1 Map of the Lower Lakes study area indicating locations(numbered and dashed areas) that experienced surface wateracidification and sample sites (crosses). The right inset mapshows sample sites in Boggy Lake that are discussed further inthe text

b

Environ Monit Assess

and Österholm 2013). Techniques used to calculateacidity from pH and dissolved metal (Al, Fe and Mn)concentrations in acid mine drainage (Kirby andCravotta 2005b; Hedin 2006) would also potentiallybe useful for assessment of water quality impacts aris-ing from acid sulfate soils, but this has yet to bedemonstrated. There is also a need for improved guid-ance on the use of surface water quality monitoring tomanage acidification events arising from acid sulfatesoil rewetting although the general principles are wellestablished through research on the liming of acidiclakes (from acid rain deposition rather than pyriteoxidation, Sverdrup 1986; Olem 1991). Whether fieldmeasurements can provide accurate information forassessment and management of acidification is alsoan uncertainty. This is of particular importance whentreatment of acidity is being undertaken as the rapidinformation provided by field measurements is highlydesirable for environmental managers.

The aim of this paper is to describe the techniquesused to monitor and assess surface water acidificationfollowing rewetting of oxidised acid sulfate soils in theLower Lakes. Particular focus is placed on analysis ofwater quality changes following acidification, relation-ships between pH and other water quality parameters,the accuracy of field based measurements compared tolaboratory measurements, and examples of how geo-chemical speciation models and calculations can pro-vide useful information for assessment and manage-ment of acidified water bodies. The monitoring andassessment techniques we outline have broad rele-vance to other locations which are experiencing, ormay in the future experience, surface water acidifica-tion following acid sulfate soil exposure.

Methods

Description of study area

The Murray–Darling Basin has a total catchment area of1,061,469 km2 (equivalent to 14 % of Australia’s totalarea) and is a highly regulated river system with a seriesof locks, weirs and storages, the majority of which weredesigned and built in the middle of the last century toassist in the delivery of water for consumptive purposes.The study area in South Australia at the end of the riversystem comprises the two lakes, Lake Albert and LakeAlexandrina, collectively known as the Lower Lakes

(see Fig. 1). These lakes are large (821.7 km2 totalsurface area), shallow (mostly<3 m deep, see Fig. 1)and are predominantly freshwater, eutrophic and highlyturbid (Mosley et al. 2012). Under sufficient flow, waterexits through a series of barrages separating the lakesand Coorong (a coastal lagoon) and either into theCoorong itself or out the river mouth into the SouthernOcean (Fig. 1). The barrages are also used to regulatewater levels in the Lower Lakes when inflows are suf-ficient. The Lower Lakes are collectively recognised asone of Australia’s most significant ecological assets andhave been designated a wetland of international impor-tance under the Ramsar convention. The lakes alsosupport several regional townships, large irrigated agri-cultural areas, tourism and recreational activities, andthe region is of high cultural importance.

Sample sites, sampling and analytical methods

Surface water quality samples (n=1,208) were collect-ed from various locations (Fig. 1) in the Lower Lakesfrom 2007 to 2010. These locations consisted of re-gions that were acidic or were perceived to be at risk ofgoing acidic and included samples before and afterneutralisation of acidity. Samples were taken by shore-line, hovercraft or boat grab sampling depending onaccessibility and safety considerations due to the rap-idly declining water levels and very soft unconsolidat-ed sediments near the water’s edge. Field sampling wasundertaken in accordance with standard methods(APHA 2005). New polyethylene bottles, washed andrinsed with deionised water, were used to collect sam-ples for laboratory analysis of acidity, alkalinity, majorions and nutrients. Acid-cleaned bottles were used tocollect samples for metal analysis. Laboratory analyseswere undertaken by the Australian Water QualityCentre’s National Association of Testing Authorities(NATA) accredited laboratory using Standard Methods(APHA 2005, and EPA 200.8). NATA accreditationrequires maintenance and documentation of strict qual-ity control procedures.

The pH, specific electrical conductivity (micro-Siemens per centimetre at 25 °C), dissolved oxygenand oxidation–reduction potential were measured inthe field at the time of sample collection using a cali-brated instrument (YSI Pro Plus). Total alkalinity in thefield and laboratory was measured by titration to a pH4.5 end-point. The field titrations were performed usinga commercially available test kit (HACH model AL-


DT). Acidity in the laboratory was measured by titrationto pH 8.3 at 25 °C following hot peroxide digestion.Acidity in the field was measured using a commerciallyavailable test kit (HACH model AC-DT). Soluble(<0.4 μm filtered) and total (following EPA 200.8 refluxdigestion with nitric and hydrochloric acids) metals (Al,As, Co, Cr, Cu, Fe, Mn, Ni, Zn) were measured byinductively coupled plasma mass spectrometry (ICP-MS) (Agilent 7500 series). Major ions (SO4, Ca, Mg,Na, K) were also measured by ICP-MS and chloride(Cl) by a Ferricyanide method (APHA 4500-Cl- E),turbidity was measured by a nephelometer and nutrients(ammonia, NH4; total phosphorus, TP; total nitrogen,TN) by standard colorimetric methods. Specific electri-cal conductivity (EC, micro-Siemens per centimetre at25 °C) and pH were also measured in the laboratoryusing calibrated electrodes. Quality assurance and con-trol checks were carried out in both the field and labo-ratory and acceptable results were achieved for repli-cates (within 10 %), blanks and spikes (85–115 %recovery).

Statistical calculations

Statistical differences in water quality between acidi-fied (defined for this purpose as pH<6.5) and non-acidified (pH>6.5) conditions were determined usingthe non-parametric Mann–Whitney U test (Helsel andHirsch 2002). These tests were performed in theMicrosoft Excel™ add-in program XLSTAT™ withstatistical significance ascribed with an α of 0.05.Only significant values are reported. Non-parametricmeasures (medians and quartiles) were also used fordescriptive statistics.

Geochemical speciation calculations

Geochemical speciation and solubility calculations wereundertaken using the computer program PHREEQC (in-teractive version 3, Parkhurst and Appelo 2013). Themeasured pH, redox potential (pE), alkalinity, major ionsand soluble (<0.45 μm filtered) metal concentrations inselected samples were used as inputs. The model wasequilibrated with a constant partial pressure of oxygen(−0.67), similar to that observed in the field (near satura-tion with dissolved oxygen). Charge balancing was un-dertaken using Cl in the model but calculated chargebalance errors in the laboratory data were typically<2 %. The standard WATEQ4F database in PHREEQC

(Ball and Nordstrom 1991) was used for the calculationswith the following additional phase information forschwermannite [Fe8O8(OH)4.32(SO4)1.84+20.32 H+=Fe+3+1.84 SO4

−2+12.32 H2O, log k=10.5] obtained fromAcero et al. (2006). Solid solubility calculations wereperformed for schwertmannite, natrojarosite[NaFe3(SO4)2(OH)6,], goethite (α-FeOOH), amorphousiron oxide [Fe(OH)3(a)], jurbanite (AlOHSO4), gibbsite(Al(OH)3), basaluminite [Al4(OH)10SO4] and amor-phous aluminium oxide [Al(OH)3(a)]. These mineralswere individually equilibrated with a representative acid-ic water sample (from Boggy Lake site 31, 27May 2011,see insert in Fig. 1) using the “equilibrium phases” func-tion in PHREEQC. The total dissolved metal concentra-tion when the solid was in equilibrium with the aqueoussolution (saturation index (SI)=0) was used as the solu-bility concentration (Morel 1983). The pH was increasedin the model by simulating titration to a fixed pH withNaOH. The calculated solubility was compared to mea-sured solubility (soluble Fe, Mn and Al concentrations)using the wider dataset. Measured values that plot on thesolubility curve for a particular mineral indicate that themineral is in equilibrium with the solution and likelycontrolling soluble metal concentrations (Kirby andCravotta 2005a). The Supplementary Material containsan example of a PHREEQC input file.

We did not consider the potential impact of dissolvedorganic matter (DOM, e.g. humic and fulvic substances)on metal speciation in our measurements or model. Weexpect that this would have a trivial impact on calculatedspeciation in our acidic waters where the fraction ofmetals bound to organic matter is typically very low(Milne et al. 2003; Nystrand and Österholm 2013).However, there is likely some unaccounted for influenceof DOM on metal speciation in our higher pH samples.

XRD analyses of solid phases

The suspended material/precipitates in 1.5-L sub-samples of an acidified water body in the Lower Lakes(at Boggy Lake, see Fig. 1 and inset) was recovered byhigh-speed centrifugation and then dried at room temper-ature under vacuum. The dried precipitates were groundin an agate mortar and pestle. The resulting fine powderswere either gently back-pressed into stainless steel sam-ple holders or lightly front-pressed onto silicon low back-ground holders for X-ray diffraction analysis (XRD)analysis. XRD patterns of samples were collected witha PANalytical X'Pert Pro Multi-purpose Diffractometer


Tab

le1

Sum

marystatisticsandpvalues

(for

sign

ificantdifferencesbasedon

Mann–Whitney

Utest)forwater

quality

(con

ductivity,pH

,acidity,alkalin

ity,nu

trients,major

ions,

turbidity

)in

pH<6.5andpH

>6.5samples

intheLow

erLakes

Param

eter

Con

ductivity

pHAlkalinity

Acidity

TKN

NH4

NOx

TP

Ca

Cl

FMg

KSi

Na

SO4

Turbidity

Units

μS/cm,

25°C

mg/L

CaC

O3

mg/L

CaC

O3

mg/Las

Nmg/Las

Nmg/L

mg/L

mg/L

mg/L

mg/L

mg/ L

mg/ L

mg/L

mg/L

mg/L

NTU

pH<6.5data

Median

6,47

03.8

083

.12.43

0.68

0.01

90.06

918

81,69

00.23

184

27.3

10.5

1,06

01,33

07.6

75th Percentile

5,28

03.3

049

.62.29

0.64

20.00

80.06

110

71,30

00.19

140

217

806

819

3.8

25th Percentile

14,000

4.5

016

63.48

1.12

10.07

20.09

138

83,40

00.28

351

82.7

151,84

03,15

025

Num

ber

7777

7766

99

99

6977

4069

6960

6977

61

pH>6.5data

Median

9,47

08.1

110

13.4

1.99

0.00

90.00

50.06

415

02,99

00.37

248

62.4

31,79

063

322

75th Percentile

3,44

07.6

817.6

1.70

0.00

70.00

50.04

756

866

0.31

7122

.52

486

196

11

25th Percentile

16,600

8.6

156

18.6

2.28

0.01

60.01

60.1

221

5,29

00.56

400

113

62,95

51,10

033

Num

ber

407

415

432

9524

023

490

240

390

413

249

392

401

222

390

419

275

Pvalue

NS

<0.00

01<0.00

01<0.00

010.01

39<0.00

010.00

5NS

<0.00

01NS

<0.00

01NS

NS

<0.00

01NS

<0.00

01<0.00

01

TKNtotalKjeldahlnitrog

en,N

Sno

n-sign

ificant,


Tab

le2

Sum

marystatisticsandpvalues

(for

sign

ificantdifferencesbasedon

Mann–

Whitney

Utest)forwater

quality

(solub

le(sol.)andtotal(tot.)metals)in

pH<6.5andpH

>6.5

samples

intheLow

erLakes

Param

eter

Al(sol.)

Al(tot.)

As

(sol.)

As

(tot.)

Cr

(sol.)

Cr(tot.)

Co

(sol.)

Co(tot.)

Cu

(sol.)

Cu

(tot.)

Fe(sol.)

Fe(tot.)

Mn

(sol.)

Mn

(tot.)

Ni(sol.)

Ni

(tot.)

Zn

(sol.)

Zn(tot.)

mg/L

mg/L

mg/L

mg/L

mg/L

mg/L

mg/L

mg/L

mg/L

mg/L

mg/L

mg/L

mg/L

mg/L

mg/L

mg/L

mg/L

mg/L

pH<6.5

data

Median

10.9

12.6

0.00

30.00

30.00

30.00

50.06

90.08

50.01

00.01

01.65

3.41

3.34

3.67

0.14

40.14

60.04

00.06

5

75th Percen-

tile

2.7

6.55

0.00

20.00

20.00

20.00

20.02

90.02

90.00

50.01

00.50

1.74

2.34

2.41

0.05

30.05

20.02

50.03

3

25th Percen-

tile

26.2

28.7

0.00

30.00

40.00

50.01

00.10

20.10

00.01

00.011

4.95

16.4

6.91

7.70

0.19

20.18

60.15

30.16

3

Num

ber

7577

7771

88

99

9.00

09

7677

6668

88

98

pH>6.5

data

Median

0.01

00.26

30.00

30.00

30.00

10.00

10.00

50.00

50.00

50.00

50.02

60.96

50.011

0.10

60.00

80.00

90.00

50.00

6

75th Percen-

tile

0.01

00.08

4–

0.00

30.00

10.00

10.00

50.00

50.00

50.00

50.00

80.30

00.00

10.05

20.00

50.00

60.00

50.00

5

25th Percen-

tile

0.02

00.69

20.00

30.00

30.00

10.00

30.00

50.00

50.00

50.00

60.12

02.11

0.10

50.22

80.01

40.02

30.01

20.02

5

Num

ber

392

418

376

401

2828

2827

2828

384

418

256

256

2828

2828

Pvalue

<0.00

01<0.00

01NS

NS

0.00

05<0.00

01<0.00

01<0.00

010.01

80.00

2<0.00

01<0.00

01<0.00

01<0.00

01<0.00

010.00

02<0.00

01<0.00

01

WQG

0.05

50.01

30.03

30.00

140.01

5NV

1.9

0.12

00.08

7

WQG

Water

QualityGuidelin

e=ANZECC(200

0)WQG

triggervalue(TV)for95

%speciesprotectio

napplicable

tofreshw

aters.Valuesprov

ided

arewith

outhardness

correctio

nexceptforh

ardn

ess-adjusted

(500

mg/Las

CaC

O3used

asaverage)WQGsforC

r,Cu,Niand

Znapplicabletofreshwatersas

perA

NZECC20

00.O

nlysolublemetalconcentrations

are

comparedto

guidelinevalues.N

Vno

guidelinevalueapplies,NSno

n-sign

ificant


in “standard” configuration mode using iron filtered CoKα radiation, automatic divergence slit and X'CeleratorSi strip detector. The diffraction patterns were recorded insteps of 0.017° 2 theta with a 0.5-s counting time per step.The water subsample from site BL31 (see Fig. 1, inset)had the highest concentration of suspended material ofthe five samples received and the results from this sampleare presented and compared to mineral saturation indicescalculated for the same sample in PHREEQC.

Results and discussion

Water quality changes following acidification

Surface water acidification (total area, 21.7 km2) wasobserved at ten locations on the shallow margins of theLower Lakes (see Fig. 1 for locations). The waterquality results are summarised in Tables 1 and 2 foracidic (pH<6.5) and non-acidic (pH>6.5) samples.

Median pH for the acidified (pH<6.5) samples was3.8 which is similar to that often found in acid minedrainage studies where there is buffering by iron min-erals (Kirby and Cravotta 2005a; Cravotta 2008).There was an absence of alkalinity as expected in allthe pH<6.5 samples and instead acidity became present(median 83 mg/L as CaCO3). A small amount of acid-ity (median 13.4 mg/L as CaCO3) was also present inthe pH>6.5 samples which is commonly found in acidmine drainage situations when metals have not yetfully hydrolysed (Kirby and Cravotta 2005a, b).

Conductivity, Cl, Mg, K and Na showed no signif-icant change in concentration during acidified condi-tions (Table 1) which indicates they are relatively un-affected by acid sulfate soil processes. In contrast,sulfate (SO4) approximately doubled in median con-centration, which is a result of pyrite (FeS2) in sulfidicmaterial oxidising during the soil/sediment dryingphase to release SO4 via the following reaction(Stumm and Morgan 1996):

FeS2 þ 15=4O2 þ 7=2H2O→Fe OHð Þ3 sð Þ2SO−24 þ 4Hþ ð5Þ

The SO4 released in the soil can then be mobilised intosurface waters upon rewetting. The dissolution of second-ary oxyhydroxysulfate minerals such as jarosite may alsoresult in increased mobilisation of sulfate upon rewetting.

Silica (Si) tripled in concentration, which is likely dueto acid dissolution of phyllosilicates/alumino-silicates

(clay) in the soil matrix found in many locations. Therewas also an increase in calcium (Ca) concentration in theacidified waters and this may be due to both: (a) disso-lution of calcium carbonate (CaCO3) in the soil/sedimentfollowing acidification and/or (b) a result of dissolutionof limestone (CaCO3) added by aerial dosing at CurrencyCreek and Boggy Lake (see Fig. 1 for locations, BoggyLake results described further below). Fluoride (F) de-creased in median concentration under the acidic condi-tions, and this may be due to increased adsorption by thehigh amounts of Al and Fe oxyhydroxides (e.g. Murray1984) likely present in the submerged materials andwater following acidification and neutralisation.

Turbidity decreased under the acidic conditions,which may relate to complete dissolution of thesuspended clay matrix and/or coagulation of colloidsdue to high ionic strength and high multivalent counterion (Al+3, SO4

−2) concentrations (Mosley et al. 2003).Nutrients were mobilised under acidic conditions withTKN, NH4 and NOx showing significant (p<0.05) in-creases in surface water. The release of nutrients canresult from microbial activity following the rewetting ofdried sediments (Baldwin et al. 2005), but this requiresfurther specific research for acid sulfate soils. TP showedno significant change.

For soluble metals, all metals with the exception ofarsenic (As) showed a significant (p<0.05) increase in thepH<6.5 samples (Table 2). At pH<6.5, most of the totalmetal concentration was comprised of the soluble metalconcentration (see Table 2, median soluble metal concen-trations similar to median total metal concentrations).The metal increases resulted in 25–100 % of samplesexceeding water quality guidelines for protection ofaquatic ecosystems for dissolved Al, Co, Mn, Ni, Zn(bold values in Table 2). Al exceedances were particular-ly large as also observed in other studies (Sammut et al.1996). In general, our results highlight the risks of toxicmetal release from sulfuric acid sulfate soils, particularlywhen the overlying water body is unable to buffer theacid addition and maintain alkalinity.

Arsenic concentrations were very low (<3 μg/L) inall samples indicating this metalloid does not appear tobe of concern with regard to surface water impacts inthe Lower Lakes. Arsenic has been observed to be-come more mobile in acid sulfate soils during therewetting phase in other locations (Burton et al.2008) with its mobility increased by high carbonateconcentrations/alkalinity (Appelo et al. 2002). Hence,it is recommended that arsenic should still be included


routinely in the metal analysis suite for assessing acidsulfate soil impacts on surface water in other locations.

Relationships with pH

Figure 2 shows the relationship between key acidifica-tion parameters and pH. There was a strong inverselog–log relationship between hydrogen ion concentra-tion (pH) and acidity (r2=0.75), and the individualmetals (Al, Fe and Mn) comprising the metal acidity(r2=0.6–0.9). The regression lines in Fig. 2 can be usedto predict metal concentration and potential toxicity forfuture acidification events in the Lower Lakes. Similarnon-linear relationships have been found in other stud-ies (Förstner 1995; Jenne 1995), including laboratoryexperiments on acid sulfate soils (Simpson et al. 2010).We believe the key influencing factors in our contextare that metals become increasing desorbed off particle

surfaces as the pH lowers (and complete desorptionlikely below pH 4–5, Stumm and Morgan 1996), andthat the soil matrix begins to dissolve at lower pHs (e.g.Shaw and Hendry 2009, dissolution of exchangeableAl on clays between pH 3 and 5, and clay matrixbetween pH 1 and 3). In contrast to our findings,Cravotta (2008) found a poor correlation between pHand total concentrations of dissolved Fe and Mn inacidic coal mine drainage which implied a large frac-tion of the metals were present as reduced FeII andMnII

species. Although we did not measure FeII, given theshallow nature of the Lower Lakes (highly oxygenat-ed) and surface sediment source of acidity, it is lesslikely that any reduced metal acidity species are pres-ent. Hence, all the acidity is likely to be manifested inthe water near the soil source and this appears to be thecase with the acidification restricted to the marginallake areas (Fig. 1).

−4

−2

0

2log Al = −0.5957*pH + 3.012 (r2 = 0.75)

Al

log

Al s

ol. (

mg/

L)

−4

−2

0

2

4log Fe = −0.4595 * pH + 2.215 (r2 = 0.60)

Fe

log

Fe

sol.(

mg/

L)

−4

−2

0

2 log Mn = −0.6971 * pH + 3.581 (r2 = 0.64)

Mn

log

Mn

sol.(

mg/

L)

0

1

2

3log Alk. = 0.3354 * pH − 0.6524 (r2 = 0.45)

Alkalinity

log

Alk

.(m

g/L

CaC

O3)

2 4 6 8 10−2

0

2

4log Acid. = −0.2921 * pH + 3.212 (r2 = 0.75)

Acidity

log

Aci

d.(m

g/L

CaC

O3)

pH2 4 6 8 10

−4

−2

0

2log SO4/Cl = −0.1255 * pH + 0.3162 (r2 = 0.43)

log

SO

4/C

l (m

g/L)

SO4:Cl ratio

pH

Fig. 2 Dissolved metals (Al, Fe, Mn), alkalinity, acidity and sulfate/chloride ratio as a function of pH in the Lower Lakes samples. Thelinear regression (following log transformation) fits of the data and adjusted r2 values are also shown


The correlation with pH and alkalinity was weaker(r2=.45) but closer examination of Fig. 2 suggests that thecorrelation is stronger as the pH lowers. The droughtconditions at the same time as the acidification eventcreated widely varying alkalinity due toevapoconcentration (Mosley et al. 2012), which resultedin much more scatter at higher pH values.

The decoupling of sulfate behaviour from that of othermajor ions during acidification discussed above is furtherillustrated in Fig. 2 which shows a moderate (r2=0.45)inverse correlation between the SO4/Cl ratio and pH. TheSO4/Cl ratio has been found to be a useful indicator ofacid sulfate soil impacts in other locations and sulfur

isotopes can also be used as a potential early warningindicator of sulfide oxidation (Kilminster and Cartwright2011). The SO4/Cl ratio in theory could also be used todetermine if sulfate reduction is occurring, which woulddecrease the ratio, but this was difficult to discern in ourdata as the rewetting phase resulted in rapid dilution.

Use of monitoring for management of acidification

Surface water acidification was actively managed intwo areas of the Lower Lakes, Boggy Lake andCurrency Creek (see Fig. 1), using limestone addition.A time series to illustrate the use of monitoring to

2

4

6

8

10

pH

pH

0

100

200

mg/

L C

aCO

3

Alkalinity

0

5000

10000

uS/c

m a

t 25C

Conductivity

0

250

500

mg/

L C

aCO

3

Acidity

0

3000

6000

mg/

L

Sulfate

0

5

10

mg/

L

Mn (soluble)

Apr−10 Oct−10 Apr−110

150

300

mg/

L

Fe (soluble)

Apr−10 Oct−10 Apr−110

75

150

mg/

L

Al (soluble)

Boggy L. 1 Boggy L. 10 Boggy L. 31 Boggy L. 33 Boggy L. 39

Fig. 3 The pH, alkalinity, specific electrical conductivity at25 °C, acidity, sulfate and soluble metals (Mn, Fe, Al) duringan acidification event in Boggy Lake and its neutralisation via

aerial limestone dosing. See the inset of Fig. 1 for sample sitelocations. The three dosing events (28 May–3 June 2010, 22–24June 2010, 29–30 August 2010) are shown as the vertical lines


inform and assess the management of acidification inone area, Boggy Lake (northern Lake Alexandrina, seeFig. 1 inset), is shown in Fig. 3. During water leveldeclines from 2008 to 2009, Boggy Lake became dis-connected from the main lake water body. As a conse-quence, large areas of acid sulfate soils (predominantlycracking clays) with sulfidic material were exposedallowing the oxidation of pyrite to occur with theformation of large areas of acid sulfate soils comprisingmainly sulfuric material (Fitzpatrick et al. 2010, alsosee photo S1 in the supplementary material). DuringMay 2010, rainfall events and water inflows toLake Alexandrina progressively re-inundatedBoggy Lake. With rewetting, very acidic water

(pH 2–3 at multiple sites, see Fig. 3) was present,particularly in the western and northwestern mar-gins of the lagoon that were furthest away fromthe main lake water body. Very high acidity andsoluble metal (Al, Fe, Mn) concentrations wereobserved, along with sulfate release (Fig. 3).

A decision was made to add limestone to BoggyLake to neutralise the acidity and raise the pH tosuitable levels for aquatic ecosystems (pH 6.5 –9.0,ANZECC 2000). Monitoring was used to determinethe extent of limestone treatment required and its ef-fectiveness. When limestone (CaCO3) is added to neu-tralise an acidic water body, the general reaction oc-curring is:

CaCO3 sð Þ þ Hþ þ solublemetal acidity→HCO−3 þ Caþ2 þmetal precipitates ð6Þ

In this reaction, the solid limestone dissolves andreleases calcium (Ca) and carbonate (CO3

−2) to the waterbody. The CO3

−2 reacts with the H+ ions to producebicarbonate (HCO3

−) and the rise in pH drives the solublemetal hydrolysis reactions (Eqs. 1, 2, 3 and 4) whichsubsequently produce additional H+ which is also

neutralised by CO3−2 to produce HCO3

−. Metal precipi-tates (e.g. hydrous ferric and aluminium oxides) are alsocharacteristically formed following the metal hydrolysisreactions and rise in pH. Based on the reaction shown inEq. 6, the limestone treatment requirement (in tonnes)can be calculated via the following equation:

Limestone required t CacO3ð Þ ¼ Acidity mg.Las CaCO3

� ��Water Vol: Lð Þ

.109 � 1

.EF %ð Þ ð7Þ

Where the laboratory measured acidity (expressed asmilligrammes per litre as CaCO3, includes H

+ and solu-ble metal acidity) corresponds to the amount of limestonethat is required to be added per unit volume of water toneutralise the water body and EF is the (limestoneneutralisation) efficiency factor (per cent). An EF of100 % implies that all the limestone dissolve and reactto neutralise the acidity. In practise, lower EFs are oftenpresent as limestone dissolution and neutralisation capac-ity is influenced by several factors, including (1) waterpH: the lower the pH of the water to be dosed, the higherthe amount and rate of dissolution (Plummer et al. 1978).EFs in the range of 50–60 % have been observed at pH 3reducing to 10 % at pH 7 (Sverdrup 1986); (2) watertemperature: limestone solubility increases with decreas-ing temperature with highest solubility in freshwater at2 °C (Stumm and Morgan 1996). The water temperaturein Boggy Lake at the time of dosing was approximately12 °C which would have resulted in a lower EF than that

observed in colder climates, (3) limestone grainsize/surface area: we used a very fine 8-μm naturallyprecipitated limestone that would be expected to have ahigh EF (Sverdrup 1986) and, due to its large surfacearea, dissolve very rapidly (Plummer et al. 1978), (4)coating/passivation of limestone: when limestone isadded to acidic water, it often becomes coated (processcommonly termed passivation) with precipitates (e.g.metal hydroxides and gypsum) which tend to dramati-cally lower the EF and can prohibit dissolution(Hammarstrom et al. 2003). This could be a very signifi-cant factor reducing the EF in our study due to the high Ca,SO4, and soluble metal levels present (Tables 1 and 2).

Three aerial limestone dosing events (total of995 tonnes) were undertaken in Boggy Lake as shownby the vertical dashed lines on Fig. 3 (also seeSupplementary Material Fig. S2 of a plane dosing acid-ified water). An EF of 100 % was applied to calculatethe treatment requirement (using Eq. 7) for the first two


events (420 and 400 tonnes CaCO3 used respectively).This initial EF was deliberately conservative (assumedcomplete dissolution and neutralisation) to prevent anyexcess limestone additions to what is an ecologicallysensitive site (Ramsar-listed wetland). The first two dos-ing events substantially reduced the water acidity andsoluble metal levels but did not achieve neutralisationof the water body (Fig. 3). An EF of 50 % was then usedfor the last dosing event (175 tonnes CaCO3) whichresulted in the water body becoming neutral (pH>6.5)in early September 2010 (Fig. 3). The combination of atime series approach for treatment and monitoring (pre-,during and post-dosing) ensured minimal adverse im-pacts as a result of the management intervention.

Comparison of field versus laboratory water qualityresults

Both field and laboratory measurements were used tomonitor the acidification events in the Lower Lakes.Field measurements have the advantages of deliveringimmediate information on the state of a water body and

being low cost, while laboratory measurements enablemore complex sample processing and analysis andhave higher quality control. A comparison of measure-ments of key acidification parameters undertaken inthe field with measurements on the same samples inthe laboratory is shown in Fig. 4 (using data fromBoggy Lake, see Fig. 3). A very strong correlation(r2≥0.97) was found between field and laboratory mea-surements for pH, alkalinity and conductivity. Acidityonly showed a moderate correlation (r2=0.54) betweenfield and laboratory measurements with laboratorymeasurements recording on average lower values(slope of regression line=0.65 of field acidity). Weexpected the laboratory method to record higher acid-ity as it involved a hot peroxide digestion which resultsin improved oxidation and hydrolysis of soluble metalscompared to the field method which does not involvethese pre-treatment steps. Hence, we believe that theobserved result may be due to the indicator endpointbeing often hard to discern in field titrations due toformation of a large amount of coloured metal precip-itates (particularly in high acidity samples) so there

2 4 6 8 102

4

6

8

10

Field pH

Lab

pH

pH

Lab = 0.99 * Field − 0.10 (r2 = 0.97)

0 50 100 1500

50

100

150

Field Alk. (mg/L as CaCO3)

Lab

Alk

. (m

g/L

as C

aCO

3)

Alkalinity

Lab = 0.96 * Field − 2.8 (r2 = 0.97)

100 200 300 400 500

100

200

300

400

500

Field Acid. (mg/L as CaCO3)

Lab

Aci

d. (

mg/

L as

CaC

O3)

Acidity

Lab = 0.65 * Field + 7.7 (r2 = 0.54)

1000 2000 3000 4000 5000

1000

2000

3000

4000

5000

Field Cond. (uS/cm, 25C)

Lab

Con

d. (

uS/c

m, 2

5C)

Conductivity

Lab = 0.92 * Field + 207 (r2 = 0.98)

Fig. 4 Field measured versus laboratory-measured concentrations for pH, alkalinity, acidity and specific electrical conductivity (at 25 °C) inthe Boggy Lake region. The linear regression line fit of the data and adjusted r2 values are also shown


may have been a tendency to “overshoot” the titration.Hence, caution needs to be applied in the use of fieldacidity results, and further research is required to im-prove these methods (e.g. potentially use pH meter forendpoint determination rather than just an indicator).

Acidity form

While the total amount of acidity present in an acidifiedwater body is a key parameter to measure, it is also

important to understand the form of acidity as solublemetal acidity (Fe, Mn and Al) presents a completelydifferent geochemical behaviour and risk profile to H+

acidity (Kirby and Cravotta 2005a). Given the stoichio-metric relationships between soluble metal concentra-tions and eventual H+ acidity (Eqs. 1, 2, 3 and 4), aciditycan be calculated using the equation (Kirby andCravotta 2005b):

Acidity mg=Las CaCO3ð Þ ¼ 50� 10 3−pHð Þ� �

þ 2Fe=55:8þ 2Mn=54:9þ 3A1=27:0� �

ð8Þ

where soluble metal concentrations are in milligrammesper litre. The acidity calculated via this equation iscompared to that measured in the laboratory in Fig. 5.There is a strong relationship between measured andcalculated acidity (r2=0.87); however, the measuredacidity tends to be about 20% higher than the calculatedacidity (slope of linear regression line of 1.19 on Fig. 5).Using the median values in Tables 1 and 2, on average,the per cent contributions of the various acidity species

to total acidity (calculated using Eq. 8) are AlIII 78.1 %,H+ 10.2 %, Mn and FeII/III 3.8 %. The relatively lowproportion of the acidity represented by H+ acidity/pHhighlights the importance of understanding the amountand geochemical behaviour of soluble metals releasedfrom these acid sulfate soils comprising mainly sulfuricmaterial.

Deviations of measured acidity from the calculated1:1 line in Fig. 5 could potentially be due to some of the

100 200 300 400 500 600

100

200

300

400

500

600

Calculated Acidity (mg/L as CaCO3)

Mea

sure

d A

cidi

ty (

mg/

L as

CaC

O3)

Measured Acidity = 1.191 * Calculated Acidity + 7.699 (r2 = 0.87)

1:1 line

Fig. 5 Measured versus calculated acidity in the Boggy Lake region. The 1:1 line and linear regression line fit of the data and adjusted r2

value are also shown


assumptions made in Eq. 8 that (1) no other solublemetal or other ions contribute substantially to acidity,(2) none of the metals are in intermediate hydrolysisspecies with lower H+ acidity equivalent contributionsthan indicated in Eq. 8, (3) the Fe is present as reducedFeII species and makes only a 2 H+ equivalent contri-bution in Eq. 8 and (4) no solid phases are present thatcould contribute to acidity. Assumption 1 is supportedin our results, which shows comparably low other

metal concentrations (Tables 1 and 2). With regard toassumption 2, intermediate hydrolysis species werecalculated to be present in PHREEQC results (e.g.FeOH2

+ and AlOH2+ present predominantly at pH 6

which contribute only one H+ equivalent, input file inSupplementary Material can be executed to view the-se), but this would lead to the calculated acidity beinghigher than the measured acidity which is opposite towhat we observed. Assumption 3 is likely to be incor-rect as discussed above, but changing the Fe stoichio-metric acidity contribution from 2 to 3 in Eq. 8 onlyslightly improved the relationship (slope value was1.14 compared with 1.21 in Fig. 4). For assumption4, as the acidity samples were unfiltered, there mayhave been solid mineral phases present that coulddissolve to release metals to solution in the laboratorydigestion and titration. Orange and brown iron oxideprecipitates were observed suspended in many of theacidified water bodies (see Supplementary MaterialFig. S3 from Boggy Lake). The potential influence ofthese phases on metal solubility and measured acidityis assessed and discussed in the next section.

Solid phase speciation

Table 3 shows SIs for relevant minerals based on theaqueous speciation of an acidified sample of Boggy Lakewater (see Supplementary Material S4 for an example of

Table 3 Saturation indices calculated in PHREEQC for selectedminerals in an acidic sample from Boggy Lake (see “Methods” fordetails of calculations and Supplementary Material for PHREEQCinput file)

Phase name Formula SI

Al(OH)3(a) Al(OH)3 −7.47Anhydrite CaSO4 −0.91Birnessite MnO2 −0.28Fe(OH)3(a) Fe(OH)3 −1.75Gibbsite Al(OH)3 −4.71Goethite FeOOH 3.87

Gypsum CaSO4/2H2O −0.67Jarosite-K KFe3(SO4)2(OH)6 1.83

Natrojarosite NaFe3(SO4)2(OH)6 −0.15Jurbanite AlOHSO4 −0.62Quartz SiO2 0.53

Schwertmannite Fe8O8(OH)4.32(SO4)1.84 0.52

46- 1045 Quartz 6- 263 Muscovite

36- 425 Natrojarosite

9- 466 Albite 19- 932 Microcline

47- 1775 Schwertmannite

2-Theta Angle (degrees)10.0 20.0 30.0 40.0 50.0 60.0 70.0

3

6

9

12

15

Inte

nsi

ty (

Co

un

ts)

X 1

000

Fig. 6 XRD pattern of the <2-μm size fraction (vacuum dried) of suspended material from Boggy Lake. The numbers for each mineralrefer to reference patterns in the international Powder Diffraction File database


the PHREEQC input file). SIs≈0 suggest solution equilib-rium with the solid phase while SIs<<0 indicateunsaturation (mineral unlikely to form) and SIs>>0 indi-cate supersaturation (formation of mineral possible butmay be subject to kinetic limitations and mineral is not inequilibrium with solution). Schwertmannite, natrojarositeand quartz were calculated to be just over saturation(Table 3) and XRD analysis confirmed the presence ofthese minerals suspended in the water column of BoggyLake (Fig. 6). The goethite SI indicated supersaturation,and this mineral was not identified in the XRD results,which suggested it was not present or insignificant relativeto schwertmannite and jarosite. Jurbanite and birnessitewere slightly undersaturated in the sample, and these werealso not identified in the XRD for this sample.

The FeIII, AlIII and MnII solid solubilities and cor-responding concentrations of soluble Fe, Al and Mnbetween pH 2 and 5 for the Boggy Lake region isshown in Fig. 7. Dissolved Fe concentrations appearto be in equilibrium with schwertmannite andnatrojarosite in many samples. These results are con-sistent with previous research that has found theseminerals are dominant at low pH in acid mine drainagesamples (Bigham et al. 1996; Acero et al. 2006) andacid sulfate soil-affected waters (Sullivan and Bush2004; Burton et al. 2006). The dissolved Fe concentra-tions and solid solubility increase as pH decreases.Equilibrium to more soluble FeII minerals such assiderite can also occur when FeII aqueous species aredominant (Kirby and Cravotta 2005a; Burton et al.2006). Al was calculated to be at or near equilibriumwith jurbanite at low pH (Fig. 7). Research on acidmine drainage has also found Al in equilibrium withAl-hydroxy-sulfate phases at low pH (Uhlmann et al.2004). MnII solid (birnessite) solubility did not corre-late well with dissolved Mn concentrations (Fig. 7),and this is consistent with previous research that Mntypically persists as Mn+2 except in alkaline (pH>8.5)and oxic environments when oxidation is morefavourable (Davies and Morgan 1989).

Secondary oxyhydroxysulfate minerals such asjarosite and schwertmannite can dissolve quite rap-idly when diluted into a higher pH, less sulfateand iron-rich environment, and this can acidify thedilution water and release trace metals (Aceroet al. 2006; Hammarstrom et al. 2005). As acidicLower Lakes water was diluted with circum-neutral pHlake water or neutralised with limestone, it is likely these

solid phases dissolved. The same process likely causedthe measured acidity to be higher than calculated acidity(Fig. 4) as the acidity titration raises the pH to >8 whichwould dissolve these phases. Over time, more stableminerals can also form (e.g. schwertmannite to goethite,Bigham et al. 1996), even under acidic conditions,which would require more consideration for longer du-ration acidification events.

10−12

10−7

10−2

Fe

(mol

/L)

Fe solubility

SchwertmanniteJarosite−Naamorph. Fe(OH)3Goethite

10−6

10−4

10−2

Al (

mol

/L)

Al solubility

Gibbsite

Jurbanite

Basaluminite

Al(OH)3(a)

3 4 5 610

−6

10−4

10−2

pH

Mn

(mol

/L)

Mn solubility

Birnessite

Fig. 7 FeIII, AlIII and MnII solid solubilities (lines) and mea-sured concentrations (circles) of soluble Fe, Al and Mn, respec-tively between pH 2 and 5 for samples in Boggy Lake. See“Methods” for details of PHREEQC calculations


Conclusion

The surface water acidification events that occurred in theLower Lakes region due to rewetting of oxidised acidsulfate soils yielded very useful knowledge on monitoringand assessment techniques. The key water quality param-eters we recommend are measured in future in the LowerLakes and other locations, along with whether field and/orlaboratory-based measurements are appropriate, are

summarised in Table 4. We found significant increases insoluble metals (exceeded guidelines for ecosystem protec-tion for Al, Co, Mn, Ni and Zn), SO4, Si, Ca and nutrientsunder acidic conditions. The log of the soluble metalconcentrations, acidity and SO4/Cl ratio correlated stronglywith pH. Acidity values were used to calculate limestonetreatment requirements, which led to the successfulneutralisation of an acidified region. Field measurementscorrelated strongly with laboratory measurements for pH,

Table 4 Parameters recommended for inclusion in monitoringand assessment studies of surface water acidification eventsarising from acid sulfate soil exposure and rewetting. The

purpose of a particular parameter and whether the parametercan be reliably measured in the field and/or the laboratory is alsopresented

Recommended parameter Field Lab Purpose

pH Yes Yes Primary acid–base indicator (decrease indicates acid inputs),inverse correlation with metal solubility, toxicity of H+ ion,inform metal speciation calculations

Alkalinity Yes Yes Primary acid–base indicator of amount of acidneutralising capacity present in the water body(a decrease indicates acid inputs)

Acidity Lower reliability Yes Primary acid–base indicator of amount of acidity presentin water body (an increase indicates acid inputs),assist in determining treatment requirements

Metals comprising acidity (Al, Fe, Mn) No Yes Form of acidity present, geochemical speciation calculations,ecotoxicity assessment

Trace metals (As, Cd, Co, Cu, Ni, Zn, etc.) No Yes Ecotoxicity assessment

Sulfate (SO4) No Yes Indicator of sulfide oxidation to sulfate, used to calculate SO4/Clratio which is an indicator of release of sulfate fromacid sulfate soils

Chloride (Cl) No Yes Used to calculate SO4/Cl ratio

Other major ions (Ca, Mg, Na, K) No Yes Used in metal speciation calculations and hardness corrections formetal guideline values

Nutrients (NH4, NOx, FRP) No Yes Inform risk of nutrient toxicity (for ammonia) and/oreutrophication

X-ray diffraction of precipitates No Yes Assess what mineral precipitates are present to assist ingeochemical speciation calculations and risk assessments

Dissolved oxygen Yes No Can be depleted following oxidation of metals and monosulfides,used in speciation calculations

Oxidation–reduction potential Yes No Indicates redox (pE, Eh) state of water body for metalspeciation calculations

Electrical conductivity Yes Yes Can indicate acid dissolution of minerals, particularly undersoft water conditions

Turbidity/suspended solids Yes Yes Determine if precipitates are present in high concentrations orcomplete acid dissolution of particles has occurred

Optional parameters

Sulfur isotopes No Yes Early detection of acid sulfate soil impacts

Chemical analysis of precipitates No Yes Determine chemical composition of solid phases to assist ingeochemical and risk assessments

FeII Yes Yes Assist in determining dissolved Fe speciation and risk assessmentsrelating to Fe


alkalinity and conductivity but only moderately with acid-ity which is likely due to difficulties in determining thetitration indicator endpoint. Acidity was calculated to bepresent mostly as soluble AlIII, which is likely due toacid dissolution of the aluminosilicate clays found in theLower Lakes region. Measured acidity was slightlyhigher than calculated acidity and geochemical specia-tion calculations and XRD evidence indicated that sec-ondary oxyhydroxysulfate minerals (schwertmannite,jarosite and jurbanite) were present and controllingdissolved Fe and Al concentrations in acidic samplesfrom the Boggy Lake region. These solid phases likelyresulted in the measured acidity being higher than thatcalculated using the soluble (<0.4 μm) metal acidity.Further research is warranted on the nature and impor-tance of these solid phases for controlling solutionchemistry and influencing field acidity titrations in acidsulfate soil impacted waters.

Acknowledgments The assistance of EPA staff (David Palmer,Emily Leyden, Peter Mettam, Ashley Natt, Karl Fradley and JarrodSpencer) in sample collection and analysis is gratefully acknowledgedas is the project management assistance of staff from the Departmentof Environment,Water andNatural Resources.We thankMarkRavenand Peter Self from CSIRO Land and Water for XRD analyses. Thepart funding contribution of the South Australian Government’s Mur-ray Futures program funded by the Australian Government’s Waterfor the Future Initiative, and the Murray–Darling Basin Authority arealso gratefully acknowledged. We also appreciate the constructivecomments of an anonymous reviewer.

References

Acero, P., Ayora, C., Torrentó, C., & Nieto, J. M. (2006). Thebehaviour of trace elements during schwertmannite precipi-tation and subsequent transformation into goethite andjarosite. Geochimica et Cosmochimica Acta, 70, 4130–4139.

Amaral, V., Thompson, E. L., Bishop, M. J., & Raftos, D. A.(2012). The proteomes of Sydney rock oysters vary spatial-ly according to exposure to acid sulfate runoff.Marine andFreshwater Research, 63, 361–369.

ANZECC. (2000). Australian and New Zealand Guidelines forFresh and Marine Water Quality. Canberra: Australian andNew Zealand Environment and Conservation Council andAgriculture and Resource Management Council of Australiaand New Zealand.

APHA. (2005). Standard methods for the examination of water andwastewater (21st ed.). Washington, DC: American PublicHealth Association, American Water Works Association andWater Environment Federation.

Appelo, C. A. J., Van der Weiden, M. J. J., Tournassat, C., &Charlet, L. (2002). Surface complexation of ferrous iron andcarbonate on ferrihydrite, and the mobilisation of arsenic.Environmental Science and Technology, 36, 3096–3103.

Baldwin, D. S., Rees, G. N., Mitchell, A. M., & Watson, G.(2005). Spatial and temporal variability of nitrogen dynam-ics in an upland stream before and after a drought. Marineand Freshwater Research, 56, 457–464.

Ball, J.W., & Nordstrom, D.K. (1991). WATEQ4F—user’s man-ual with revised thermodynamic data base and test cases forcalculating speciation of major, trace and redox elements innatural waters: U.S. Geological Survey Open-File Report90–129, 185 p.

Bigham, J. M., Schwertmann, U., Traina, S. J., Winland, R. L., &Wolf, M. (1996). Schwertmannite and the chemical model-ing of iron in acid sulfate waters. Geochimica etCosmochimica Acta, 60, 2111–2121.

Boman, A., Fröjdö, S., Backlund, K., & Åström, M. (2010).Impact of isostatic land uplift and artificial drainage onoxidation of brackish-water sediments rich in metastableiron sulfide. Geochimica et Cosmochimica Acta, 74, 1268–1281.

Burton, E. D., Bush, R. T., & Sullivan, L. A. (2006).Sedimentary iron geochemistry in acidic waterways associ-ated with coastal lowland acid sulfate soils. Geochimica etCosmochimica Acta, 70, 5455–5468.

Burton, E. D., Bush, R. T., Sullivan, L. A., Johnston, S. G., &Hocking, R. K. (2008). Mobility of arsenic and selectedmetals during re-flooding of iron- and organic-rich acid-sulfate soil. Chemical Geology, 253, 64–73.

Cook, F. J., Hicks, W., Gardner, E. A., Carlin, G. D., &Froggatt, D. W. (2000). Export of acidity in drainagewater from acid sulphate soils.Marine Pollution Bulletin, 41,319–326.

Cravotta, C. A., III. (2008). Dissolved metals and associated con-stituents in abandoned coal-mine discharges, Pennsylvania,USA: 2. Geochemical controls on constituent concentrations.Applied Geochemistry, 23, 203–226.

Davies, S. H. R., & Morgan, J. J. (1989). Manganese (II)oxidation-kinetics on metal-oxide surfaces. Journal ofColloid and Interface Science, 129, 63–77.

Dent, D. L., & Pons, L. J. (1995). A world perspective on acidsulfate soils. Geoderma, 67, 263–276.

Fitzpatrick, R. W., Shand, P., &Merry, R. H. (2009). Acid sulfatesoils. In J. T. Jennings (Ed.), Natural history of theriverland and Murraylands (pp. 65–111). Adelaide, SouthAustralia: Royal Society of South Australia (Inc.).

Fitzpatrick, R.W., Grealish, G., Chappell, A., Marvanek, S.,Shand, P. (2010). Spatial variability of subaqueous andterrestrial acid sulfate soils and their properties for theLower Lakes South Australia, CSIRO Land and Waterreport. http://www.environment.sa.gov.au/Conservation/Rivers_wetlands/Coorong_Lower_Lakes_Murray_Mouth/The_environment/Acid_sulfate_soils/Research_projects/Acid_Sulfate_Soils_Research_Program_reports. Accessed22 July 2013

Förstner, U. (1995). Non-linear release of metals from aquaticsediments. In W. Salomons & W. M. Stigliani (Eds.),Biogeodynamics of pollutants in soils and sediments (pp.247–307). Berlin: Springer.

Hammarstrom, J. M., Sibrell, P. L., & Belkin, H. E. (2003).Characterization of limestone reacted with acid-mine drain-age in a pulsed limestone bed treatment system at theFriendship Hill National Historical Site, Pennsylvania,USA. Applied Geochemistry, 18, 1705–1721.


http://www.environment.sa.gov.au/Conservation/Rivers_wetlands/Coorong_Lower_Lakes_Murray_Mouth/The_environment/Acid_sulfate_soils/Research_projects/Acid_Sulfate_Soils_Research_Program_reports




Hammarstrom, J. M., Seal, R. R., II, Meier, A. L., & Kornfeld, J.M. (2005). Secondary sulfate minerals associated with aciddrainage in the eastern US: recycling of metals and acidity insurficial environments. Chemical Geology, 215, 407–431.

Hedin, R. S. (2006). The use of measured and calculated acidityvalues to improve the quality of mine drainage datasets.Mine Water and the Environment, 25, 146–152.

Helsel DR, Hirsch RM (2002) Statistical methods in water re-sources. In: ‘Techniques of Water-Resources Investigationsof the United States Geological Survey; Book 4,Hydrologic Analysis and Interpretation’, 510 pp. http://water.usgs.gov/pubs/twri/twri4a3/. Accessed 22 July 2013

Jenne, A. J. (1995). Metal adsorption onto and desorptionfrom sediments: I. Rates. In H. Allen (Ed.), Metal contam-inated aquatic sediments (pp. 81–110). Chelsea: Ann ArborPress.

Kilminster, K.,&Cartwright, I. (2011). A sulfur-stable-isotope-basedscreening tool for assessing impact of acid sulfate soils onwaterways. Marine and Freshwater Research, 62, 152–161.

Kirby, C. S., & Cravotta, C. A., III. (2005a). Net alkalinity andnet acidity in mine drainage 1: theoretical considerations.Applied Geochemistry, 20, 1920–1940.

Kirby, C. S., & Cravotta, C. A., III. (2005b). Net alkalinity andnet acidity in mine drainage 2: practical considerations.Applied Geochemistry, 20, 1941–1964.

Macdonald, B. C. T., White, I., Åström, M. E., Keene, A. F.,Melville, M. D., & Reynolds, J. K. (2007). Discharge ofweathering products from acid sulfate soils after a rainfallevent, Tweed River, eastern Australia. Applied Geochemistry,22, 2695–2705.

Milne, C. J., Kinniburgh, D. G., Van Riemsdijk, W. H., &Tipping, E. (2003). Generic NICA—Donnan model param-eters for metal-ion binding by humic substances.Environmental Science and Technology, 37, 958–971.

Morel, F. M. M. (1983). Principles of aquatic chemistry. NewYork: Wiley.

Mosley, L. M., Hunter, K. A., & Ducker, W. A. (2003). Forcesbetween colloidal particles in natural waters. EnvironmentalScience and Technology, 37, 3303–3308.

Mosley, L. M., Zammit, B., Leyden, E., Heneker, T. M., Hipsey,M. R., Skinner, D., et al. (2012). The impact of extreme lowflows on the water quality of the Lower Murray River andLakes (South Australia).Water Resources Management, 26,3923–3946.

Murray, F. (1984). Fluoride retention in highly leached disturbedsoils. Environmental Pollution (Series B), 7, 83–95.

Nystrand, M. I., & Österholm, P. (2013). Metal species in aBoreal river system affected by acid sulfate soils. AppliedGeochemistry, 31, 133–141.

Olem, H. (1991). Liming Acidic Surface Waters. LewisPublishers, 331 pp.

Parkhurst, D.L., & Appelo, C.A.J. (2013), Description of inputand examples for PHREEQC version 3—a computer pro-gram for speciation, batch-reaction, one-dimensionaltransport, and inverse geochemical calculations: U.S.Geological Survey Techniques and Methods, book 6,chap. A43, 497 p., http://pubs.usgs.gov/tm/06/a43.Accessed 22 July 2013

Plummer, L. N., Wigley, T. M. L., & Parkhurst, D. L. (1978). Thekinetics of calcite dissolution in CO2–water systems at 5 to60 °C and 0.0 to 1.0 atm CO2. American Journal of Science,278, 179–216.

Pons, L.J. (1973). Outline of the genesis, characteristics, classi-fication and improvement of acid sulphate soils. In: H. Dost(Ed.), Proceedings of the 1972 (Wageningen, Netherlands)International Acid Sulphate Soils Symposium, vol 1 (pp. 3–27). International Land Reclamation Institute Publication18. (Wageningen, Netherlands).

Sammut, J., White, I., &Melville, M. D. (1996). Acidification of anestuarine tributary in eastern Australia due to drainage of acidsulfate soils. Marine and Freshwater Research, 47, 669–684.

Shaw, S. A., & Hendry, M. J. (2009). Geochemical and miner-alogical impacts of H2SO4 on clays between pH 5.0 and−3.0. Applied Geochemistry, 24, 333–345.

Simpson, S. L., Fitzpatrick, R. W., Shand, P., Angel, B. M.,Spadaro, D. A., & Mosley, L. (2010). Climate-drivenmobilisation of acid and metals from acid sulfate soils.Marine and Freshwater Research, 61, 129–138.

Stumm, W., & Morgan, J. J. (1996). Aquatic chemistry—chemicalequilibria and rates in natural waters (3rd ed.). New York:Wiley.

Sullivan, L. A., & Bush, R. T. (2004). Iron precipitate accumu-lations associated with waterways in drained coastal acidsulfate landscapes of eastern Australia. Marine andFreshwater Research, 55, 1–10.

Sverdrup, H. U. (1986). The dissolution efficiency for differentstream liming methods. Water, Air, and Soil Pollution, 31,827–837.

Uhlmann, W., Bütcher, H., Totsche, O., & Steinberg, C. E. W.(2004). Buffering of acidic mine lakes, the relevance ofsurface exchange and solid bound sulphate. Mine Waterand the Environment, 23, 20–27.


http://water.usgs.gov/pubs/twri/twri4a3/

http://water.usgs.gov/pubs/twri/twri4a3/

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monitoring and assessment of surface water acidification following rewetting of oxidised acid...

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