Removal of Heavy Metal Ions by Ferrihydrite: an

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May 17, 2016 - ferrihydrite by the rapid hydrolysis of Fe(III) ions and investigates its .... amorphous or poorly crystalline depending on synthesis methods.

Water Air Soil Pollut (2016) 227:193 DOI 10.1007/s11270-016-2899-7

Removal of Heavy Metal Ions by Ferrihydrite: an Opportunity to the Treatment of Acid Mine Drainage Nuray Karapınar

Received: 24 February 2016 / Accepted: 17 May 2016 # Springer International Publishing Switzerland 2016

Abstract Ferrihydrite is often an initial precipitate resulting from the neutralization of Fe(III) solution, and it seems to be one of the products of acid mine drainage forming reactions. Since having the adsorption properties, ferrihydrite can be effective for the remediation of acid mine drainage. This study prepared fresh ferrihydrite by the rapid hydrolysis of Fe(III) ions and investigates its adsorptive behaviours toward Pb(II), Cu(II), and Zn(II). When the sorption data were presented in plot of percent sorbed versus pH, it was found that sorption is strongly dependent on the solution pH and increasing as expected at higher pH for all metal ions investigated. All the observed metal cation sorption began at pH values below zero point charge (ZPC) of ferrihydrite (pH = 7.8–8.0), and almost all removal are achieved at pH values lower than that related metal hydroxide obtained. Enhanced removal of metal ions, as the pH of the solution and initial metal ion concentration are increased, was attributed to surface precipitation of metal hydroxide. The existence of ferrihydrite and adsorption of metal ions onto surfaces are favouring surface precipitation of metal ions at lower pH values than that for metal ion only. Depending on the pH of the solution and initial metal ion concentration, more than one mechanism such as adsorption by complexation and surface precipitation was responsible for the removal.

N. Karapınar (*) General Directorate of Mineral Research and Exploration, Üniversiteler Mahallesi, Dumlupınar Bulvarı No:139, 06800 ÇankayaAnkara, Turkey e-mail: [email protected]

Keywords Acid mine drainage . Secondary minerals . Heavy metal ions . Sorption . Removal

1 Introduction Because of posing serious hazard to the ecosystem and human health, acid mine drainage (AMD) has been considered to be one of the major environmental challenges facing the mining industry worldwide. The causes of AMD relate to the natural weathering of mine wastes and rocks enriched in metal sulphide minerals. Although varying widely depending on the physical, chemical, mineralogical and microbiological properties of each site, the typical characteristics of AMD can be stated as high acidity, and high sulphate and metal concentrations. AMD, often containing toxic concentrations of Fe, Al, Cu, Zn, Cd, Pb, Ni, Co and Cr, can be produced from the mining of coal and metallic deposits. Values of pH for AMD can range from −3.5 to 5, but even circumneutral (pH = 7) mine waters can have high concentrations of As, Sb, Mo, U and F (Nordstrom 2011). The production of AMD involve the chemical and biological oxidation of metal sulphides contained in mine waste heaps, active or abandoned mine workings, or in tailings piles left over from the processing of sulphide ores. The iron sulphides such as pyrite (FeS2), pyrrhotite (Fe1 − xS, 0.7 < x > 1.0) and marcasite are possibly the most common sources of AMD production, because they are ubiquitous in metal sulphide ores and they generally are not the target of ore beneficiation processes (Stumm and Morgan 1996). Since

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depended on numerous variable factors such as the quantity and type of sulphide minerals, grain size distribution and grain morphology, bacterial activity, moisture content, and availability of dissolved oxygen or other oxidants, the chemistry and microbiology of AMD formation are rather complicated. However, it is generally represented as a pyrite oxidation by oxygen from the air in the presence of water. The following reactions can represent transformation of pyrite to soluble iron and sulphuric acid (Kuyucak 2000): FeS2 ðsÞ þ 7=2 O2 ðgÞ þ H2 O→Fe2þ þ 2SO4 2− þ 2Hþ

Feþ2 þ 1=4 O2 ðgÞ þ Hþ →Fe3þ þ § H2 O

ð1Þ

ð2Þ

FeS2 ðsÞ þ 14 Fe3þ þ 8H2 O→15 Fe2þ þ 2SO4 2− ðlÞ þ 16 Hþ

Fe3þ þ 3H2 O→FeðOHÞ3 ðsÞ þ 3Hþ

ð3Þ

ð4Þ

Ferric iron can also oxidize sulphide minerals (reaction 3). The amount of acid generated as a result of iron sulphide oxidation is higher than that for oxidation by oxygen. Moreover, precipitation of aqueous ferric iron as ferrihydrite (an iron oxyhydroxide) contributes to acid generation (reaction 4). The overall sulphide to sulphate oxidation is summarized by FeS2 ðsÞ þ 15=4 O2 ðgÞ þ 7=2H2 O→ FeðOHÞ3 ðsÞ þ 2SO4 2− ðaqÞ þ 4Hþ

ð5Þ

One of the weathering products of primary sulphide minerals is secondary minerals. Depending on the aqueous conditions such as pH, redox, metal and sulphate concentrations, they may vary in a large range of minerals such as sulphates, oxides, oxyhydroxides, carbonates, sulphides and silicates (Jamieson 2011). As pH increases, aqueous metal species tend to precipitate as hydroxide, oxyhydroxide or hydroxysulphate minerals and, moreover, these secondary minerals play important

roles in attenuating contaminants from mine effluent by adsorption. Since solubility of iron depends on its redox state as well as the pH, when ferrous iron oxidized to ferric iron, it must hydrolyse to precipitate given in reaction 6. Fe3þ þ H2 O→FeOH2þ þ Hþ

; pK1 ¼ 2:2

ð6Þ

Therefore, dissolved ferric iron precipitates at pH values of 2–2.5 (Nordstrom 2011). Depending upon the site-specific pH conditions, iron oxide precipitates such as goethite (α-FeOOH), lepidocrocite (γ-FeOOH), schwertmannite (Fe 1 6 O 1 6 (OH)y(SO 4 )z•nH 2 O) and ferrihydrite (Fe5HO8•4H2O) can typically be found in mine drainage discharges (Schwertman and Fischer 1973; Chapman et al. 1983; Ferris et al. 1989; Schwertmann and Cornell 1993; Cornell and Schwertmann 2003). Amongst them, ferrihydrite and schwertmannite are poorly crystalline iron oxides (Nordstrom and Alpers 1999). Ferrihydrite is one of the sources of iron for the formation of more crystalline iron minerals such as goethite, lepidocrocite, and hematite in the weathering environment (Mohapatra and Anand 2010). Ferrihydrite, a naturally occurring material, can be synthesized by rapid hydrolysis of Fe(III) solution (Cornell and Schwertmann 2003) and the resultant precipitates often referred to as Bferric hydroxide^, Bhydrous ferric oxide^, Biron oxide hydrate gels^ or Bamorphous ferric hydroxide^ (e.g. Towe and Bradley 1967; Schwertman and Fisher 1973; Davis and Leckie 1978; Carlson and Schwertmann 1981; Schwertmann 1987; Liaw et al. 1989; Stumm and Morgan 1996; Mohapatra and Anand 2010). These precipitates are amorphous or poorly crystalline depending on synthesis methods. The term Bferrihydrite^ is often used to describe both 2- or 6-line ferrihydrites, which have either two or six identifiable broad reflections in a diffraction pattern. X-ray diffraction patterns were first obtained by van der Giessen (1966) and independently by Towe and Bradley (1967). Chukhrov et al. (1973) presented a detailed date for a natural ferric hydroxide. They proposed ferrihydrite as a new name. There are a number of formulae proposed; however, since they are essentially equivalent, they can be reduced to FeOOH 0.4 H2O (a hydrated iron oxyhydroxide formula) (Zhao et al. 1994). The morphology of ferrihydrite is spherical, and unlike other iron oxides, it exists only as nanocrystals resulting in high specific surface areas ranging from 100 to 700 m2/g

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(Cornell and Schwertmann 2003). Because of its poor crystallinity, the structure of ferrihydrite, however, is still a subject of controversy (Zhao et al. 1994). Being fine grained and having a high capacity for adsorption of potentially toxic metals, ferrihydrite can therefore limit the aqueous concentration of such elements by adsorption onto its surfaces. In this study, the sorption behaviour of Pb(II), Cu(II) and Zn(II) onto a freshly precipitated ferrihydrite has been investigated as a function of pH and metal concentration to identify the role of secondary iron oxide minerals in the remediation of AMD.

2 Experimental 2.1 Materials and Methods Analytical grade chemicals and deionized water which had a conductivity of 0.7 mS/cm were used to prepare stock solutions, and the pH adjustment was preceded by nitric acid and sodium hydroxide. All the experiments were performed at 25 ± 0.5 °C, using a suitable beaker (of 1000 ml volume) mixed by a mechanic stirrer at 230 rpm. Analysis of metal ions in solution was performed by atomic adsorption spectrophotometer. The stock solution of lead, copper and zinc was prepared by dissolving their nitrate salts, Pb(NO3)2, Cu(NO3)2.3H2O and Zn(NO3)2.6H2O, respectively, in 0.02 M HNO3 solution. A freshly precipitated ferrihydrite was chosen as the model adsorbent for the study. A convenient and widespread procedure that was also used in this work is to neutralize ferric solutions with excess of alkali to give a red-brown precipitate that is called ferrihydrite. Iron nitrate (Fe(NO3)3.9H2O) was used in the experiments to prepare the ferric solutions. Sodium hydroxide was added to the solution containing desired iron ions until a pH of 8.0 was attained and the resulting precipitate aged for 1 h. After aging, solid/liquid separation was done by centrifugation and the resulting precipitate is dried at 80 °C. Dried solid was characterized by X-ray diffraction (XRD), and the specific surface area and its density were measured by BrunauerEmmett-Teller (BET) method and pycnometer, respectively. The term Bferrihydrite^, as used here, corresponds to an amorphous Fe(III) oxyhydroxide.

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In sorption experiments, after aging of adsorbent at pH = 8.0, the solution pH was adjusted to the desired sorption pH and it was conditioned for 1 h, then metal ions added at desired concentration and pH of solution was kept at the pre-set sorption value by NaOH. After 40 min for sorption, unless stated otherwise, suspension was filtered and analysed. Sorption of metal ions onto ferrihydrite was investigated on the basis of pH and metal ion concentration. In this study, sorption term was used as a general term describing the attachment of metal ions from the solution to ferrihydrite. In solubility experiments of adsorbent, after aging, pH of the solution was gradually decreased by adding acid and, after 30 min conditioning for each stage, a sample taken was filtered and analysed for soluble iron. Metal ion speciation and saturation index values of solution with respect to metal hydroxide were calculated by using PHREEQC Version 2 which is a computer program written in C language to perform a variety of low-temperature aqueous geochemical calculations (Parkhurst and Appelo 1999).

3 Results and Discussion 3.1 Ferrihydrite Characterization The ferrihydrite was characterized by X-ray diffraction analysis, which showed a spectrum of typical amorphous structure (Fig. 1). The specific surface area of ferrihydrite was measured by the single-point BET method (nitrogen gas adsorption after 12 h of outgassing at room temperature) and found to be 140 m2/g. A pycnometrically measured density of the precipitate was found to be 3.8 g/cm3. The fresh ferrihydrite precipitate was insoluble in a wide pH range (Table 1). Iron oxides and related amorphous hydrates have been the subject of many studies regarding the preparation conditions and chemical and physical properties (e.g. van der Giessen 1966; Towe and Bradley 1967; Chukhrov et al. 1973; Davis and Leckie 1978; Eggleton and Fitzpatrick 1988; Zhao et al. 1994; Cornell and Schwertmann 2003). Therefore, it was emphasized to investigate the freshly precipitated ferrihydrite’s ability to remove heavy metals rather than its formation and structure.

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Fig. 1 XRD spectrum for amorphous Fe(III) oxyhydroxide

(a)

Table 1 Solubility of freshly precipitated ferrihydrite by pH (FeT = 0.5 × 10−2 M) pH

Dissolved iron, mg/l

Dissolved iron, %

3.69

0.220

0.08

4.76

0.074

0.03

5.51

0.044

0.02

6.65

0.044

0.02

8.18

0.040

0.01

100

Sorption, %

80 60 1

40 2

20

3

0 3

3.5

4

4.5

5

5.5

6

6.5

7

pH

(b) 100 1

80

Sorption, %

When the sorption data were presented in plot of percent sorbed versus pH, as expected, it was found that sorption is strongly dependent on the solution pH for all metal ions investigated (Fig. 2). pH dependence of sorption can be explained because of the changing of both the ferrihydrite surface properties (charge and potential) and the solution composition (the degree of ionization and speciation of the metal ion) with pH. With increasing pH, ferrihydrite surfaces will become more favourable to sorption of metal ions. The point of zero charge (pHZPC) of fresh iron oxyhydroxide prepared in this way is given as 7.9–8.1 (Davis and Leckie 1978). But results showed that all the observed metal cation sorption took place at pH values below pHZPC of ferrihydrite indicating that electrostatic adsorption could not be a dominant mechanism of sorption; it can be said that there is a specific sorption by ferrihydrite. This phenomenon is being explained by surface complex formation model indicating surface complex formation of metal ions with active sites on oxides (e.g. Davis and Leckie 1978; Sposito 1983; Dzombak and Morel 1990; Davis and Kent 1990; Stumm and Morgan 1996). In the presence of water, the surface of oxides is generally covered with surface hydroxyl groups. Unlike cations in the interior of the structure, cations at the surface are not fully coordinated. Therefore, they adsorb water molecule from the solution to complete their

2 3

60 40 20 0 3.5

4

4.5

5

5.5

6

6.5

7

pH

(c) 100 1 2 3

80

Sorption, %

3.2 Sorption Experiments

60 40 20 0 4.5

5

5.5

6

6.5

7

7.5

8

pH Fig. 2 Metal ion sorption onto ferrihydrite, FeT = 0.5 × 10−2 M. a Pb (1) 1.0 × 10−4 M, (2) 5.0 × 10−4 M, (3) 1.0 × 10−3 M. b Cu (1) 1.0 × 10−4 M, (2) 8.8 × 10−4 M, (3) 3.5 × 10−3 M. c Zn (1) 1.0 × 10−4 M, (2) 5.0 × 10−4 M, (3) 2.0 × 10−3 M

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Table 2 Hydrolysis and stability constants of ions studied (from phreeqc data base) Hydrolysis constant log Kh Pb −7.71 Cu −8.0 Zn −8.96

Stability constant of M(OH)2 log Ks 8.15 8.69* 11.50

*From Stumm and Morgan (1996)

13 12 11

-Log Kh

coordination shell. The protons attached to the adsorbed water molecules then tend to redistribute themselves that form surface –OH groups, various and not fully structurally and chemically equivalent. In surface complexation model, cation adsorption by hydrous oxides involves some form of OH−–M+2 interaction on the surface. Thus, the concept of surface complexation has been considered as adsorption equilibria in the same way as equilibria in solutions and typically categorized by relative Bstrength^ of interaction between the adsorbate (species in solution) and the surface or adsorbent. If solvating water molecules are positioned between the cation or anion and the surface, the adsorption complex is referred to as outer sphere and is considered to be weak. Conversely, if upon adsorption the adsorbate loses waters of hydration such that there are no water molecules positioned between the cation or anion and the surface, the adsorption complex is referred to as inner sphere and is considered to be strong. The extent to which the cation will sorb to secondary precipitates as outer sphere or inner sphere complexes will vary as a function of the ion species, the secondary precipitate, pH, particle size and surface area, and presence of other sorbing species that may compete for adsorption sites. From the adsorption-pH curves, it was observed that there is a selectivity sequence between the metal ions. The selectivity sequence (figures in parentheses indicate pH values for 50 % sorption, pH50) derived from the adsorption-pH curves at 1.0 × 10−4 M metal concentration for ferrihydrite is Pb(4.2), Cu(5.0) and Zn(6.5) (the lower the pH, the higher selectivity). When the observed metal ion sorption is compared to the solubility and hydrolysis data (Table 2), a linear relationship between pH50 and solubility product of metal hydroxide and hydrolysis constant was found. Although there is a direct relationship between adsorption and hydrolysis constant of the cation, it was seen from Fig. 3 that

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10

Zn(6.5)

9

Cu(5.0) Pb(4.2)

8

< %1 < %10 < %50

7 6 5 2

3

4

5

6

7

8

9

10

pH Fig. 3 Relationship between metal ion sorption (i.e. the pH50) by ferrihydrite and first hydrolysis constant, −log Kh. Metal ion concentration is 1 × 10−4 M

adsorption occurred at pH values (i.e. the pH50) at which only a small fraction of ions was hydrolysed. Early models such as the ion-solvent interaction model (James and Healy 1972) and the adsorption hydrolysis model (Matijevic et al. 1960; Matijevic 1967) attempted to explain the adsorption in terms of fundamental ion-surface interactions and the dominant species (e.g. M+2 and MeOH) adsorbed, indicating that hydrolysis reduces the charge of the adsorbing ion and hence lowers the solvation barrier to adsorption and adsorption of hydrolysable metal ions is directly related to the presence of hydrolyzed species. They postulate that hydroxyl groups in the coordination sphere of the hydrolysis products of the ion may act as bridging ligands between the adsorbed metal ion and adjacent surface groups. Additionally, the results of the study undertaken by Crawford et al. (1997) support the hypothesis that hydroxide ligand is essential to metal ion adsorption, possibly due to hydrogen bonding. In surface complexation model, coordination chemistry of the metal ion reactions with surface hydroxyl groups is most important and, therefore, cations that form strong complexes with OH− in water also bind strongly to hydrous oxides. With both hypotheses, this sequence can simply be explained on the basis of their complex formation stabilities with OH− ligand and it follows the sequence of first hydrolysis constants. Therefore, ionic properties determining its complex formation stabilities with OH− ligand such as charge, radius and electronegativity will affect the metal’s affinity to be adsorbed by oxides. According to these models, sorption is restricted by the amount of active sites on oxides so it must be expected that saturation of surfaces is attained after some adsorption. However, results of this study showed that when the initial metal ion concentration is

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increased, sorption continued with the increase of pH (Fig. 2). Therefore, the extremely high capacities of ferrihydrite cannot be explained by simple adsorption by surface complex formation model. It is quite possible that more than one mechanism is most probably responsible for the removal process. In order to gain a better insight into the removal, a batch removal study was conducted with and without ferrihydrite. The same procedure was applied in these experiments, but only separation of ferrihydrite was carried out by sedimentation. After 10 min sedimentation time, samples taken below the surface were acidified and analysed for metal ion. Results are given in Table 3. For the better explanation of the obtained sorption results for metals, saturation indexes (SI) of related metal hydroxide are presented in Fig. 4. Stability and first hydrolysis constants used in calculation are summarized in Table 3. For all three metal ions studied, the metal was removed from solution at lower pH than that required for hydroxide precipitation (Figs. 2, 3 and 4 and Table 3). Calculated SI of related metal oxides is equal or lower than zero under conditions of pH

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