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The formation of silicate-stabilised passivating layers on pyrite for reduced acid rock drainage Rong Fan, Michael Short, Sheng-Jia Zeng, Gujie Qian, Jun Li, Russell Schumann, Nobuyuki Kawashima, Roger St.C. Smart, and Andrea R Gerson Environ. Sci. Technol., Just Accepted Manuscript • DOI: 10.1021/acs.est.7b03232 • Publication Date (Web): 23 Aug 2017 Downloaded from http://pubs.acs.org on August 27, 2017

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Environmental Science & Technology

The presence of silicate in the surface iron oxy(hydroxide) layers results in a pronounced change in surface morphology, increase in pyrite surface stability and decrease in dissolution. 86x46mm (150 x 150 DPI)

ACS Paragon Plus Environment


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The formation of silicate-stabilised passivating

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layers on pyrite for reduced acid rock drainage

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Rong Fana, Michael D. Shorta, Sheng-Jia Zenga, Gujie Qiana, Jun Lia, Russell C. Schumanna,b,

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Nobuyuki Kawashimac, Roger St.C. Smarta,d and Andrea R. Gersond* a

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School of Natural and Built Environments, University of South Australia, Mawson Lakes,

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South Australia 5095, Australia; b

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c

Levay & Co. Environmental Services, Edinburgh, South Australia 5111, Australia;

Future Industries Institute; University of South Australia, Mawson Lakes, SA 5095, Australia d

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Blue Minerals Consultancy, Middleton, SA 5213, Australia

*Corresponding author: Tel.: 61 422 112 516; Email andrea@blueminerals consultancy.com.au

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ABSTRACT

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Acid and metalliferous release occurring when sulfide (principally pyrite)-containing rock from

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mining activities and from some natural environments is exposed to the elements is

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acknowledged as a major environmental problem. Acid rock drainage (ARD) management is

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both challenging and costly for operating and legacy mine sites. Current technological solutions


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are expensive and focused on treating ARD on release rather than preventing it at source. We

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describe here a viable, practical mechanism for reduced ARD through the formation of silicate-

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stabilised iron hydroxide surface layers. Without silicate, oxidised pyrite particles form an over-

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layer of crystalline goethite or lepidocrocite with porous structure. With silicate addition, a

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smooth, continuous, coherent and apparently amorphous iron hydroxide surface layer is

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observed, with consequent pyrite dissolution rates reduced by more than 90% at neutral pH.

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Silicate is structurally incorporated within this layer and inhibits the phase transformation from

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amorphous iron (oxy)hydroxide to goethite, resulting in pyrite surface passivation. This is

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confirmed by computational simulation, suggesting that silicate-doping of a pseudo-amorphous

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iron oxyhydroxide (ferrihydrite structure) is thermodynamically more stable than the equivalent

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undoped structure. This mechanism and its controlling factors are described. As a consequence

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of the greatly reduced acid generation rate, neutralisation from onsite available reactive silicate

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minerals may be used to maintain neutral pH, after initial limestone addition to achieve neutral

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pH, thus maintaining the integrity of these layers for effective ARD management.

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Introduction

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The discharge of acid and associated dissolved metals from mine wastes and some natural

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environments has caused extensive waterway acidification and toxic metal contamination,

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leading to adverse water quality outcomes for natural, agricultural and domestic use1, 2. Pyrite

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(FeS2), the primary contributor to the acid release, is the most abundant sulfide mineral in the

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earth’s crust3; consequently its oxidation and leaching characteristics have been widely studied,

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frequently with a view to rate reduction4-6. The oxidant (ferric iron or dissolved oxygen),

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aqueous environment and the presence of catalysts (for example iron-oxidising bacteria), all

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significantly impact pyrite oxidation rates and have been examined in detail7. At neutral pH, as is

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relevant here, the dominant reactions are:

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FeS2 + 15/4O2 + 7/2H2O → Fe(OH)3 + 2H2SO4 → FeOOH + H2O + 2H2SO4

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More correctly, the initially-formed amorphous iron hydroxide Fe(OH)3 transforms, possibly via

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crypto-crystalline ferrihydrite8 (formally 5Fe2O3.9H2O, by loss of 6H2O from 10 units of

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Fe(OH)3), to the crystalline oxyhydroxide (goethite or lepidocrocite) FeOOH (by loss of a further

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H2O)9, 10. For simplicity, we will refer to the initial Equation 1 oxidation product as amorphous

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iron hydroxide and the intermediate and final products as semi-amorphous and crystalline iron

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oxyhydroxide respectively. Where the nature of the product is uncertain, we use the terminology

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(oxy)hydroxide.

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Reducing the concentration of oxidant at the pyrite surface, which can be achieved by the

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formation of surface layers, can yield significantly decreased oxidation rates4, 5, 11-15. The

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formation of an iron (oxy)hydroxide-silicate layer, via precipitation onto framboidal pyrite

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surfaces, was observed to reduce the pyrite dissolution rate by one order of magnitude12, but this


(1)

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effect had limited duration. On the basis of long-term column leaching research and field studies

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of acid rock drainage (ARD) from waste rock dumps with limestone additions at the Grasberg

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mine (Indonesia), it was proposed that silicate, from silicate mineral dissolution, may stabilise

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the iron (oxy)hydroxide over-layer at neutral pH, reducing pyrite oxidation rates by 50–95%

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through restricted O2 and water access to pyrite surfaces16, 17. It has been found that these

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passivating layers can be preserved in a continuous, coherent and stable form by providing

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sufficient alkalinity to maintain the pH above 618. The form of this surface layer, as observed to

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date, has been described as an amorphous iron hydroxide-containing silica11, 18. The inclusion of

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silicate is known to significantly affect the structural characteristics of iron oxyhydroxide

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precipitates from solution (without pyrite) and the transformation of semi-amorphous iron

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hydroxide (ferrihydrite) to more stable iron oxyhydroxide and oxide phases (goethite and

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hematite (Fe2O3))19-21.

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These observations suggest that soluble silicate may have a significant influence on the nature of

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iron (oxy)hydroxide surface layers formed on pyrite during oxidation at neutral pH, and

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consequently, alter the efficacy of these surface layers in inhibiting further oxidation. However,

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the formation mechanism, structure and, importantly, stability of these passivating layers on

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pyrite has not been defined. This paper provides detailed characterisation of these surface layers

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and their formation mechanisms with the goal of understanding their potential role in control of

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acid and metalliferous drainage.

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Materials and Methods

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Sample and dissolution tests. Pyrite (FeS2) was supplied by Geo Discoveries (NSW, Australia).

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It contained 42.8 wt.% Fe and 51.5 wt.% S with minor trace impurities of 0.01 wt.% Al, 0.01


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wt.% Ca, 0.01 wt.% Cu, 0.01 wt.% Zn and 0.16 wt.% C. Powder X-ray diffraction indicated that

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the sample contained no mineral phases other than pyrite. The pyrite was crushed, ground, and

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screened to a particle size range of 38−75 µm. Sonication was applied to remove fine particles

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adhering to the larger ones. The sample was washed with 3 M HCl solution for a few minutes

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and subsequently washed with ethanol and dried in a vacuum oven overnight. The surface area

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was 0.35 m2 g−1 according to BET analysis.

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A concentrated sodium metasilicate (Na2SiO3) solution (D serial sodium silicate, pH 11-12,

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Na2O 14.7 wt.%; SiO2 29.4 wt.%; solid 44.1 wt.%; dissolved Si 22,000 mg L−1 confirmed by

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ICP-MS) supplied by PQ Corporation (Malvern, PA, USA) was used as the source of dissolved

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silicate. The concentrated sodium metasilicate solution was diluted in Milli-Q water to 0.8 mM,

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i.e. 22 mg L−1 Si, which is much less than the concentrations used in previous studies using

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silicate additions 12. With these initial silicate concentrations, the portion of polymeric silica that

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forms in the pH range from 3.5 to 7.5 is less than 20% of the total Si and most is present as

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silicate monomers 22. For simplicity and relation to Si assay, the silicate concentrations will be

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expressed as mM Si. The concentration was selected because the passivation effect becomes

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pronounced at about 50 mg L−1 sodium metasilicate (0.4 mM Si) according to our preliminary

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study. 0.01 M KCl was used as the background electrolyte to ensure sure the solution ionic

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strengths were approximately constant throughout the experiment 23.

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The natural concentration of silicate in drainage waters is highly dependent on local geological

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and hydrological condition with the Si concentration chosen for application in this study being of

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similar magnitude to Si concentrations reported in neutral mine drainage or natural waters at pH

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6–8 of 3.5–15 mg L−1 24, 25. A further study carried out by some of the authors of this paper has


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reported Si concentration of 55 mg L−1 at a Tasmanian mine site in waste rock dump seepage

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(pH 3.3) and 3.7 mg L−1 at neutral pH in the downstream catchment. This suggests that the

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concentration of silicate used may be achievable some ARD systems but in others silicate

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amendments may be required. The minimum concentration of silicate required to achieve the

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same behaviours as reported here has not been examined as yet.

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Experiments were conducted at three different solution pH values (3.0 ± 0.1, 5.0 ± 0.2 and 7.4 ±

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0.4) without and with addition of dissolved silicate (0.8 mM Si). The pH values chosen are

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relevant to natural and mine waste ARD environments and to remediated systems. Our objective

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was to minimise the amount of alkaline amendment required for effective remediation to provide

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an economically feasible strategy that is practical for large-scale implementation. For this reason

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alkaline systems were not investigated. The solution pH was manually adjusted daily using 1 M

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HCl or 1 M NaOH solution. For experiments at neutral pH, 1.66 g calcite was used to buffer the

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solution pH. The phase purity of calcite of 98.7% is estimated from inorganic C content in the

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bulk assay and also confirmed by powder X-ray diffraction. The calcite was kept separate from

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the pyrite in a Nylon mesh bag (mesh size 31 µm) so there was no physical contact. No gypsum

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(CaSO4.2H2O) was observed using scanning electron microscopy (SEM) for these experimental

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conditions. Speciation calculations (PHREEQC) also indicated that the system was

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undersaturated with respect to gypsum. One litre of solution with 2 g of pyrite added was used

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for each experiment; 10 mL was sampled periodically and membrane filtered (0.45 µm) for

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inductively coupled plasma optical emission spectrometry (ICP-OES) analysis. The aliquots

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removed were not replaced but totalled no more than 110 mL over the course of the experiments.


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A 1 L wide-mouth high-density polyethylene (HDPE) bottle was used for each dissolution

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experiment to exclude the possibility of silicate leaching from glass containers11. Each bottle was

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sealed using a HDPE lid with a 5 mm hole ensuring full atmospheric contact and minimal

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evaporation. Milli-Q water was added based on daily weighing to compensate for evaporation

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losses. All experiments were conducted at room temperature without agitation. Periodic solution

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samples were taken for solution redox (Eh, SHE) and pH measurements, as well as Fe, S and Si

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concentration analyses by ICP-OES. Experimental errors, less than 10% for one standard

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deviation, were estimated from uncertainty of solution analysis as relative standard deviation.

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Ultrafiltration of dissolution supernatant. Total concentrations of Fe, S, and Si were measured

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on a portion of each sample digested with nitric acid prior to analysis by ICP-OES. The

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remaining sample (not digested) was filtered through a 0.45 µm nylon membrane. A portion of

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the filtrate was analysed by ICP-OES and the remainder subjected to ultrafiltration.

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Ultrafiltration was carried out using a Model 8400 Amicon stirred ultrafiltration cell (Millipore

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Corporation). Each sample was fractionated using regenerated cellulose membranes (Millipore

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Corporation) with nominal molecular weight cut-off (1 kDa NMWCO, Filter Code PLAC07610).

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Fe, S and Si concentrations were measured in the filtrate and retentate to determine the mass

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fraction of each element passing 1.5 nm (1kD NMWCO).

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Surface analyses are described in Supporting Information S1 and statistical analyses in

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Supporting Information S2.

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Results and Discussion

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Pyrite dissolution

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Experiments were conducted at three solution pH values (3.0 ± 0.1, 5.0 ± 0.2 and 7.4 ± 0.4)

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without and with addition of dissolved silicate (0.8 mM Si) in quiescent conditions to mimic

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ARD control conditions. The pyrite dissolution rate, as moles of pyrite per unit surface area

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(Supporting Information Table 1), for each solution condition, was calculated from the measured

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solution total S concentration across 290 days (FIGURE 1Figure 1).

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At pH 3.0, there was no significant difference (p = 0.695) in the pyrite dissolution rate

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(Supporting Information Table 1) with or without the addition of silicate during 290 days (Figure

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1a). At pH 5.0, the pyrite dissolution rate in the presence of added silicate is significantly smaller

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than without silicate addition (FIGURE 1Figure 1b). At pH 7.4 in the absence of dissolved

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silicate (Figure 1c), the pyrite dissolution rate (Supporting Information Table 1) was essentially

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constant during 290 days leaching, with approximately 1.3 mmol·m−2 dissolved at 290 days. In

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contrast, the equivalent pyrite dissolution rate in the presence of silicate (Figure 1c) was

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effectively zero (p = 0.143), with only ≈0.03 mmol·m−2 pyrite dissolved after 290 days,

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representing a dissolution rate reduction of >97%. It is clear that the presence of dissolved

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silicate significantly affects pyrite dissolution at pH 7.4. To more clearly demonstrate this, the

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pyrite dissolved in the presence of silicate minus that dissolved in the absence of silicate, at each

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pH, is plotted as a function of time (Figure 1d).


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FIGURE 1. Pyrite dissolution as a function of time at pH 3.0 (a), pH 5.0 (b) and pH 7.4 (c).

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Pyrite dissolution for the system with added silicate (initially 0.8 mM Si) minus pyrite

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dissolution with no added silicate (d). Rate data fitted with linear regression lines ± 95%

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confidence bands and level of statistical significance for differences in rates (Supporting

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Information S3) given as either not significant (ns), significant at p < 0.001 (***) or p < 0.0001

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(****).

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The pyrite dissolution rate (5.00×10−11 mol m−2 s−1) at circum-neutral pH 7.4 without silicate is

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significantly greater than at pH 3.0 (H = 7.20; p < 0.05) (Supporting Information, Table 1). Long

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term acid rock leaching tests have shown that the greatest sulfate release rate is obtained in the

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presence of limestone in rock wastes12, 26. It has been proposed that pyrite oxidation rates are


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enhanced by the addition of bicarbonate/carbonate ions27 via a cycle of ferrous-pyrite/ferric-

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carbonate redox couples, resulting in increased rates of transferral of electrons from S in pyrite to

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dissolved O2. In ref. 27, 0.1–1.0 M carbonate was used, whereas our carbonate concentrations,

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equilibrated with air, would not exceed 0.001 M.

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The log rate of pyrite dissolution during the 290 days (y; mol m−2 s−1) without silicate when

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plotted as a function of pH (x; 3.0, 5.0; 7.4) (Supporting Information Table 1) is significantly

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linear (log y = (0.0669)x − 10.764; R2 = 0.995). This near-perfect linearity corresponds to the

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leach trend found for the pH region where pyrite dissolution rate is determined by dissolved O228.

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This suggests that the effect of the presence of dissolved carbonate at pH 7.4 is negligible and

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that the increased leach rate is due predominantly to pH effects.

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Williamson and Rimstidt28 have shown that at pH 3.0 and above, where ferric iron

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concentrations are very low due to precipitation, dissolved O2 is the dominant oxidant. In our

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quiescent conditions, the rate measured for all three pH conditions is around one order of

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magnitude slower than their rates at the same pH in their study where solution stirring was

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applied, suggesting that the reaction is to some extent bulk diffusion-controlled. A similar

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conclusion has been made from hydrologic simulation of a mine tailing waste29.

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Iron hydroxide formation and silicate adsorption

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Adsorption of silicate on iron oxyhydroxides has been demonstrated by observation of an FTIR

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absorption frequency in the range 950−1000 cm−1 assigned to Fe-O-Si19, 21, 30-33, with the extent

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of adsorption determined by the aqueous silicate concentration and pH. Silicate is preferentially

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adsorbed on ferrihydrite surfaces at pH 8−11. At silicate to iron mole ratios >1, or at pH values

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below 8 or above 11, silicate adsorption decreases34, 35. Transformation studies using iron


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oxyhydroxide precipitates showed that the proportion of semi-amorphous ferrihydrite in the final

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products, which also included crystalline goethite and/or hematite, after 24 h increased from 1%

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to nearly 100% as silicate concentration was increased from 10−5 to 10−3 M19. Hence, the

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incorporation of silicate into the structure of iron oxyhydroxide species inhibits the

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crystallisation of ferric-oxyhydroxides and consequently, the presence of amorphous ferric-

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(oxy)hydroxide may be maintained36.

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Recent research on iron precipitation in water samples from a legacy radioactive waste site has

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confirmed this mechanism of inhibition of Fe(II) oxidation derived from elevated silica

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concentrations in the circum-neutral pH range37. Their results have shown that, as Si:Fe ratios

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increase, the primary Fe(III) oxidation product transitions from lepidocrocite to a

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ferrihydrite/silica-ferrihydrite composite. Competitive desorption experiments suggested that

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Fe(II) was associated with more weakly bound, outer-sphere complexes on silica-ferrihydrite,

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rather than lepidocrocite, conferring decreased ability for Fe(II) to undergo surface-induced

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hydrolysis inhibiting the heterogeneous Fe(II) oxidation mechanism.

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Figure 2 shows the dissolution, precipitation and silicate adsorption behaviour over time for each

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pH condition. There is a strong anticorrelation between the solution pH and iron concentration

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without (r = −0.650; p = 0.002) and with silicate addition (r = −0.813; p < 0.0001), with iron

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concentration decreasing with increasing pH. During the first 10 days of pyrite oxidation at pH

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3.0, the ratio of S-to-Fe in solution was nearly two both in the presence and absence of silicate.

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Subsequently, this ratio increased to greater than two, indicative of Fe(oxy)hydroxide

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precipitation38. The ratio of S-to-Fe in solution at pH 5.0 and 7.4 was greater than two from

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commencement of leaching.


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At pH 3.0 there is slightly less aqueous iron in the presence of silicate than in the absence of

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silicate (Figure 2a). In contrast, at pH 5.0 the iron concentrations are greater in the presence as

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compared to the absence of silicate (Figure 2b), while the iron concentrations at pH 7.4 are less

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than the detection limit in most instances due to iron-containing precipitation. It has been

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observed previously that the iron concentration retained in solution increases with silicate to iron

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ratio36. This is believed to be due to the presence of ultra-fine, iron hydroxide colloidal particles,

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stabilised by the presence of silicate, resulting in increased apparent iron concentration39.

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However, others have quantitatively demonstrated formation of the aqueous FeSiO(OH)32+

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species in dilute acidic aqueous solutions using UV-visible absorption spectroscopy40. Our

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ultrafiltration measurements at pH 3.0 showed 99% of the iron was in species

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