Heterogeneous reactions control Cr(VI) release and sequestration in complex chemical mixtures of Cr, Fe, Cu, and organics

Abstract

Waters contaminated with toxic hexavalent chromium [Cr(VI)] can contain co-occurring metals and organic compounds, promoting simultaneous redox, complexation, and precipitation reactions. This study systematically investigates how Fe minerals, aqueous Fe(II) and Cu(II), and various low molecular weight organic compounds induce Cr(VI) adsorption, reduction, and precipitation reactions. Through batch and column experiments and microscopic and spectroscopic analyses, we determine how aqueous metals and organic compounds interact with Cr(VI) to promote sequestration or release of Cr adsorbed to iron (oxyhydr)oxides. Aqueous Cu precipitates from solution, which removes Cr(VI) from the water column, but may prevent some Cr(VI) from being reduced. With the addition of ascorbic acid, Cu can be mobilized as Cu(0) colloids with Cr also being released. Aqueous Fe(II) promotes Cr(VI) reduction, but also may mobilize Cr associated with reacted Fe (oxyhydr)oxides. The results identified in our study provide insights about overlooked reactions that can control the sequestration and release of Cr in contaminated waters containing complex mixtures of inorganic and organic chemicals which have relevant implications for remediation strategies and recovery of critical minerals.

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Environmental significance

Chromium contaminated waters from the electroplating, steel, and tanning industries often contain co-occurring heavy metals and organic compounds. These constituents may induce unanticipated chemical reactions that interfere with remediation strategies of toxic chemical mixtures, or critical mineral resource recovery. To improve these strategies and better understand these reactions, we systematically investigate how Cr(VI), Fe minerals, Cu(II), organic compounds, and Fe(II) induce adsorption, reduction, and precipitation reactions and analyze the resulting transformation products. In complex chemical mixtures, reducing conditions can mobilize Cr, which counters the typical behavior of simple Cr(VI)-containing waters.

Introduction

Critical minerals are essential to meet the technological needs of our society in energy, electronics, defense, aerospace, and industrial manufacturing.^1–3 For example, steel making requires iron ore with chromium often added to prevent corrosion.^3,4 Chromium (Cr) and copper (Cu) are critical materials for the flourishing electronics and energy industries.^2,3 Despite decades of research into the reactivity and mobility of Cr, Fe, and Cu, the variable chemistry of aqueous and solid mixtures complicates the waste treatment and potential recovery of these metals. Although necessary as a critical mineral, chromium is also a well-known toxic transition metal found in water and solids in natural systems. Chromium occurs in two primary oxidation states: toxic and mobile Cr(VI) and relatively insoluble and less toxic Cr(III).^5,6 The U.S. Environmental Protection Agency has set a maximum contaminant level (MCL) of 100 µg L⁻¹ for total Cr with California setting an MCL of 10 µg L⁻¹ for the more toxic form of Cr(VI).^7–9 Cr(III) typically precipitates as a Cr(III) hydroxide or oxide mineral or strongly adsorbs to solid phases. By reducing Cr(VI) to Cr(III), which occurs under weakly reducing conditions, chromium can be removed from water systems.

Chromium commonly co-occurs with heavy metals or organic compounds in complex mixtures from the electroplating, steelmaking, and tanning industries, which can alter transformation products and reduce resource recovery. Leather tanning waste contains high levels of Cr from Cr salts and organic matter from hide degradation.¹⁰ The iron and steel industry can generate significant wastewater containing Cr(VI) and Fe at near neutral pH.¹¹ Electroplating waste streams contain high levels (>10 g L⁻¹) of numerous heavy metals, such as Cr, Cu, Ni, and Fe as well as considerable amounts (>100 mg L⁻¹) of organics including oils, solvents, and highly toxic cyanide to complex with metals.^12–16 Before waste discharge, the acidic plating liquid may be neutralized by limestone or NaOH, which typically induces precipitation of metals, such as Fe(III) hydroxide.¹⁶ Resource recovery from these heavy industries reduces waste generation and improves resource utilization,^17,18 but the complex mixtures of Cr, heavy metals, and organics from heavy industry may alter the recovery efficiency of critical minerals. By improving understanding of the nature of these complex waste streams and their transformation products, treatment and recovery strategies may be improved.

Chromium is known to react with Fe, Cu, and organic compounds, but their combined effects on metal solubility are largely unknown. Cr(VI) adsorbs to Fe(III) solids that form in Fe-containing wastewater.^16,19,20 Cr(VI) adsorption to Fe(III) solids is pH-dependent with the negatively charged chromate ion (CrO₄²⁻ or HCrO₄⁻) adsorbing more strongly at low pH when iron (oxyhydr)oxides are positively charged.^20,21 Other anions, such as sulfate, can competitively adsorb to iron (oxyhydr)oxides as well,^21,22 limiting chromate adsorption. If pH or anion concentrations shift, such as during wastewater treatment and/or resource recovery, Cr(VI) may desorb from these iron solids, releasing Cr(VI) back into solution. Reduced Fe(II) can promote recrystallization of Fe(III) solids, releasing some Cr in the process.^15,23–26 However, Fe(II) also promotes Cr(VI) immobilization through reduction to Cr(III) with Cr(III) precipitating as Cr hydroxide or adsorbing to solid surfaces.^5,6,27,28 In the presence of organic compounds, reduced Cr(III) may be transported in the form of soluble Cr(III)–organic complexes.^29–31 Heavy metals, such as Fe and Cu, can act as an electron shuttle for Cr(VI) reduction and affect the transformation of iron oxides that contain Cr.^32,33 The presence of Cu may also induce co-precipitation of Cr through the formation of CuCrO₄ or adsorption of Cr(VI) to highly insoluble Cu(II) minerals, such as CuCO₃ or Cu(OH)₂.^12,13,34 Therefore, Cr, aqueous metals, Fe solids, and organic compounds have a complex interplay of redox, adsorption, and precipitation reactions, which may generate unexpected products in complex mixtures.

If complex Cr mixtures are released to the environment, Cr, Cu, and other metals further interact with anthropogenic and natural organic matter and metal chemistry evolves with changing redox and pH conditions. Groundwater contaminated with industrial Cr has been identified by co-occurrence with volatile organic carbon compounds,³⁵ suggesting that industrial Cr sources characteristically are associated with organic carbon. Cr in groundwater can also interact with natural organic matter (NOM), which can reduce Cr(VI) to Cr(III), removing it from solution.^6,36 NOM may lead to the formation of small Fe-NOM colloids that transport Cr through water systems.^36–40 Therefore, Cr, Fe, and organic compounds can form mobile transformation products even under reducing conditions. However, if complex mixtures also contain Cu or similar heavy metals, metal mobility in groundwater systems may decrease.

Complex mixtures of toxic Cr, Cu, Fe, and organics can produce overlooked redox and precipitation reactions that control metal mobility, which have relevant implications for processing, remediation, and resource recovery. Understanding the potential reaction pathways of Cr with organic compounds and other co-occurring metals in water is necessary to assess their sequestration and future release. In this study, we conduct numerous batch and column experiments of Cr(VI) solutions interacting with goethite or ferrihydrite in the presence and absence of organic compounds, Cu(II), and Fe(II) emulating industrial mixtures and environmental matrices to improve understanding of adsorption, redox, and precipitation reactions affecting Cr mobilization. The novelty of this study is the integration of laboratory experiments with microscopy and spectroscopy to elucidate how Cu(II), organics, and Fe(II) interact to promote sequestration and release of Cr adsorbed to iron (oxyhydr)oxides in complex chemical mixtures.

Methods

Batch experiments

Batch experiments were initially conducted to determine the individual and combined effects of Cu, Fe, and organics on Cr reactions with iron oxyhydroxides. These experiments at near neutral pH conditions determine metal solubility of complex Cr mixtures, such as those observed in environmental samples, neutralized plating solutions, and steel wastewater.^11,16,36 Iron oxyhydroxides (goethite and ferrihydrite) were added to Cr(VI) solutions with variable amounts of Cu, Fe, and organics. All experiments were run in triplicate using a solution containing 0.25 mM sodium chromate, 10 mM sodium sulfate as a competitive anion for chromate, and 5 mM HEPES buffer to keep the solution pH at 7.0 ± 0.1. Roughly 10 mM of goethite was added to each goethite experiment. Goethite was purchased from Aldrich (lot number BCBM8699V). 20 mM of ferrihydrite synthesized by the method of Schwertmann and Cornell (1991)⁴¹ was added to each ferrihydrite experiment.

We tested the effect of organics, Cu, and Fe by adding 1 mM of Cu(II) sulfate, Fe(II) sulfate, and organic compounds (citric acid, ascorbic acid, or cysteine). Citric acid, ascorbic acid, and cysteine were selected for their variable, but relatively simple chemistry. These aqueous organic compounds act as a proxy for natural organic matter, but do not precipitate under our experimental conditions. Citric acid does not significantly reduce Cr(VI), whereas cysteine moderately reduces Cr(VI) and ascorbic acid rapidly reduces Cr(VI).⁴² Initially, only one chemical addition to the Cr(VI) solution was tested before we experimented on the combined effect of Cu and a chemical reductant (either Fe(II), cysteine, or ascorbic acid). The solubility of Cr, Cu, and Fe species within these chemical mixtures was calculated using PHREEQC with the Minteq database.^43,44 For each experiment, the pH was adjusted following the chemical addition(s) to a pH of 7.0 ± 0.1. After this pH adjustment, goethite or ferrihydrite were added to the solution.

All experiments were mixed using a rotisserie rotating mixer for fifteen minutes before sampling. The Cr(VI) solutions were filtered from the iron oxyhydroxides using 0.45 µm Nylon filters with the filters being dried and stored for solid-phase analysis (see below). The resulting solutions were acidified to a pH of ∼2 using nitric acid for UV-Vis and ICP-OES analysis.

Column experiments

Two column experiments were conducted in triplicate to investigate the effects of added metal and organic reductants on Cr(VI) mobility in flowing system with Fe oxyhydroxides. By adding and removing additional chemical constituents over time, we observed how Cu(II), Fe(II), and organic compounds affected Cr(VI) sequestration and mobilization with iron oxyhydroxides. Plastic columns with an inner diameter of 1.3 cm and a length of 20 cm were packed with 6 g of Aldrich goethite and sealed with glass wool. Aqueous solutions were run through the columns using a peristaltic pump and rubber tubing. Flow through the columns was set at 2.5 mL minute⁻¹, and the tubing was quickly moved from solution to solution to keep the columns flowing throughout the experiment.

After running DI water through the columns, a Cr(VI) solution containing 0.25 mM NaCrO₄, 10 mM NaSO₄, 5 mM HEPES buffer, and 1 mM NaBr flowed through the columns. The next solution had the same aqueous chemistry, but added 1 mM Cu(II) sulfate (and used 1 mM NaCl, instead of NaBr). The third solution added 1 mM cysteine to the Cr(VI) and Cu(II) mixture (with 1 mM Br), and the fourth solution added 1 mM Fe(II) sulfate to the mixture of Cr(VI), Cu(II), and cysteine (with 1 mM Cl). Bromide and chloride were alternated between solutions to act as a tracer of flow through the columns. Each solution flowed through the goethite column for 30 minutes before the peristaltic tubing was moved to the next solution. The solutions were open to air as would be observed in industrial waste, but oxidation of solutes could occur. Approximately two hours into the experiment, the columns clogged due to mineral precipitation. Several solution vials that fed the columns contained precipitates, including a powdery yellow precipitate in the Cr(VI) and Cu(II) sulfate vial, a thick black precipitate in the Cr(VI), Cu(II) sulfate, and cysteine vial, and a thick brown precipitate in the Cr(VI), Cu(II) sulfate, cysteine, and Fe(II) vial (Fig. S1). Supplying the column with the thick brown and/or black precipitate likely clogged columns over time. If the solutions with precipitates were injected for less time or if precipitates were allowed to settle out, clogging may have been avoided. Once the columns clogged, the experiment was ended, and the goethite was collected from the columns.

Following the clogging of the first set of columns, we attempted another column experiment using ascorbic acid, instead of cysteine. While the Cu(II) and Cr(VI) solution formed a yellow precipitate once more, the solution vial with Cr(VI), Cu(II), and ascorbic acid did not contain a thick black precipitate as seen with cysteine. Instead, it formed a thinner pink solid (Fig. S1). The solution with Cr(VI), Cu(II), ascorbic acid, and Fe(II) formed a brown precipitate again, but it appeared less thick than with cysteine. A similar protocol was followed as in the first experiment with running a Cr(VI) solution, then Cr(VI) and Cu(II) solution, then Cr(VI), Cu(II), and ascorbic acid solution, and finally a Cr(VI), Cu(II), ascorbic acid, and Fe(II) solution (Fig. 4). The solutions each flowed through the column for 30 minutes before moving to the next solution. At two hours into the experiment, the flow from the Fe(II) containing solution was switched back to the third solution of Cr(VI), Cu(II), and ascorbic acid. At every 30 minutes, the flowing solutions were switched in backward order until the initial Cr(VI) solution was run for a final time.

Aqueous chemical analyses

Filtered samples from batch and column experiments were acidified and compared to chromate standards of 10 µM, 25 µM, 50 µM, 100 µM, 250 µM, and 500 µM Cr(VI) on a Cary 50 Bio spectrophotometer using a wavelength of 370 nm.⁴² This UV-Vis method provides quantification of Cr(VI) concentrations in solution. Aqueous Cr, Cu, and Fe concentrations of filtered and acidified samples were also analyzed by ICP-OES on an Optima 5300DV. Using UV-Vis and ICP-OES in concert generates Cr(VI) and total Cr data respectively. By subtracting Cr(VI) concentrations from total Cr concentrations, aqueous Cr(III) concentrations may be estimated. However, this estimate may be skewed by uncertainties in chemical analyses or Cr behavior. Bromide and chloride concentrations were determined using Thermo Fisher/Dionex ICS 1100 ion chromatography. The solubility of Cr, Cu, and Fe species in batch and column experiments was calculated using PHREEQC version 3.3.12 with the Minteq database.^43,44

Solid chemical analyses

From the batch experiments, selected filters containing goethite and associated Cr were analyzed for Cr oxidation state using X-ray photoelectron spectroscopy (XPS) as well as X-ray absorption (XAS). The surface (5–10 nm) elemental composition and Cr oxidation state of filters was analyzed using a Kratos AXIS-UltraDLD XPS with CasaXPS used to fit and quantify spectra. The filters charged readily, so a considerable amount of tape was required to achieve proper charge neutralization. This tape required overnight degassing in an anaerobic chamber or pumpdown chamber to reach the necessary vacuum for XPS analysis. It was observed that during the pumpdown process and exposure to X-rays, Cr(VI) can be reduced to Cr(III).^45–48 Gold powder was added to the goethite and associated Cr samples for energy correction. Following a survey scan, the gold 4f_7/2 peak was scanned and used to correct Cr energies. Cr was scanned at the Cr 2p peaks with the oxidation state quantified using the larger 2p_3/2 peak.⁴⁹ Filters from batch experiments using only Cr(VI), Cr(VI) and Cu(II), Cr(VI) and cysteine, Cr(VI) and citric acid, and Cr(VI) and Fe(II) were analyzed for Cr oxidation state.

Cr K-edge X-ray absorption measurements were performed at the Soft X-ray Microcharacterization Beamline (SXRMB) of the Canadian Light Source in Saskatoon, Canada. SXRMB is a bending-magnet-based beamline that utilizes InSb(111) and Si(111) crystals for monochromatization to cover an energy range of 1.7–10 keV. Samples were mounted onto double-sided, conductive carbon tape and loaded into the vacuum chamber with a 10⁻⁷ torr vacuum. Cr(OH)₃ and Na₂CrO₄ samples were used as references and for energy calibration. A 7-element SDD detector was used to record the powder samples' fluorescence yield (FLY). The total electron yield (TEY) from the drain current of the sample was also recorded. The Demeter program suite was used for data analyses, including background removal, normalization, and X-ray absorption near edge structure (XANES) analyses.

Goethite collected from the column experiments was further analyzed for distribution and elemental composition of Cr and Cu using electron microprobe analyses (EMPA) and scanning transmission electron microscope (STEM). Elemental mapping of goethite collected from the two column experiments was conducted using a JEOL JXA8200 electron microprobe for EMPA. Albite, orthoclase, diopside, apatite, and InAs were used as calibration standards. Probe for EMPA was used to determine elemental weight percent using the Phi-Rho-Z method.

Microstructural and chemical analyses of selected samples were achieved using a probe aberration corrected JEOL NEOARM 200CF scanning transmission electron microscope (STEM), operated at 200 kV. The samples were mounted on Ni lacey carbon grids. Electron micrographs were obtained in STEM mode. Energy dispersive X-ray analyses (EDXS) were performed on the same instrument using a 100 mm² JEOL detector and a hard X-ray (HX) aperture to avoid spurious signals. The collected maps were analyzed using the Oxford Aztec software.

Results and discussion

Cr(VI) reduction after adsorption to goethite and ferrihydrite

Starting with simple interactions of Cr(VI) with Fe corrosion products as observed in industrial and environmental settings,^15,19,36,50 we conducted batch experiments with Cr(VI) and goethite or ferrihydrite. In batch experiments, when the 0.25 mM Cr(VI) solution was exposed to 10 mM goethite, an average of 13% of Cr(VI) adsorbed to the goethite (3.3 mmole of Cr(VI) adsorbed per mole of goethite) (Fig. 1A). Similarly, 3.6 mmole of Cr(VI) adsorbed per mole of ferrihydrite with 29% of Cr(VI) adsorbing to 20 mM ferrihydrite (Fig. 1B). Cr(VI) adsorbed slightly to both Fe oxyhydroxides without any chemical additions of organics or additional metals. Synthetic Fe oxyhydroxides, such as goethite and ferrihydrite, typically have a point of zero charge occurring approximately at pH 8.^51,52 Therefore, at pH 7.0, these Fe oxyhydroxides were likely slightly positively charged, attracting the negatively charged Cr(VI)O₄²⁻ oxyanion and causing some minor Cr(VI) adsorption. According to analyses of XANES spectra, the Cr on the Fe surface was primarily Cr(III) with Cr(III) representing 68% of total Cr (Fig. 3 and Table S2). Under the vacuum conditions we used for solid chromium speciation analysis, we observed Cr(VI) reduction to Cr(III), likely due to adsorption of adventitious carbon that naturally occurs in the air.⁴⁵ Due to the high redox potential of Cr, adventitious carbon can effectively reduce Cr(VI). Previous spectroscopic studies have also observed primarily Cr(III) in environmental samples that do not contain an obvious reductant, which has been ascribed to reduction with adventitious carbon and/or radiation damage.^47,53,54 Fortunately, Cr(VI) is not fully reduced by adventitious carbon and/or radiation damage, so Cr(VI) reduction by organic reductants or Fe(II) can still be detected. This unintentional Cr(VI) reduction reaction affects XPS and XANES measurements to a similar degree. One benefit of this Cr(VI) reduction reaction is that prevention of Cr(VI) reduction may be detected by a higher proportion of Cr(VI) than in a control Cr(VI) sample.

Fig. 1 Cr(VI) removal for various treatment combinations of Cr, Cu, Fe(II), cysteine, and ascorbate added for batch experiments with goethite (A) and ferrihydrite (B).

Influence of organic compounds on Cr(VI) reduction

To understand the influence of organic compounds on Cr(VI) uptake relevant to industrial Cr groundwater contamination,³⁵ we measured Cr(VI) uptake and reduction in presence of organic acids (citrate) and reductants (ascorbic acid and cysteine). The fraction and mechanism of Cr(VI) removal depended on the type of organic compound. When 1 mM citric acid was added, 18% of Cr(VI) adsorbed to goethite according to UV-Vis spectroscopy, within uncertainty of the reaction without citric acid. Citric acid does not strongly reduce or bind with Cr(VI),⁴² so little change in Cr(VI) sequestration occurred. Using XPS, we observed a similar Cr(VI) signal with the addition of citric acid as the sample with no citric acid (Fig. 2). The addition of ascorbic acid (96% removal of aqueous Cr(VI)), resulted in reduction and much greater sequestration of Cr(VI) with goethite (Fig. 1A). Ascorbic acid is a strong reductant, which can rapidly reduce Cr(VI).⁴² XPS analysis confirmed Cr(VI) reduction with most Cr occurring as Cr(III) (Fig. 2). Cysteine can also reduce Cr(VI) with 40% of Cr(VI) removed with goethite. XPS confirmed Cr(VI) reduction with cysteine as well (Fig. 2). A similar removal was observed using ferrihydrite as the mineral addition with 53% removal with cysteine. The similarity in Cr(VI) removal using goethite or ferrihydrite in the presence of an organic reductant suggests that Cr(VI) reduction and uptake may be consistent amongst Fe oxyhydroxides.

Fig. 2 XPS spectra of batch experiments with 250 µM Cr(VI) and goethite as well as 1 mM Fe(II), cysteine, Cu(II), and citric acid. Cr(VI) (yellow) is highest in Cu(II) sample, but also observed with Cr(VI) only, citric acid, and possibly cysteine samples. Cr(III) hydroxide (blue) is the primary Cr(III) compound observed. Cr(III) oxide (green) may also be present, particularly in the Fe(II) sample. A satellite peak (purple) was also accounted.

Reaction with Cu(II) and Fe(II) improves removal of Cr(VI)

Since heavy metals, such as Cr, Cu, and Fe, are found in plating and steel industry waste streams and contaminated waters are often co-contaminated with Cr and other heavy metals, we conducted experiments in batch reactors containing 1 mM Cu(II) or Fe(II). When 1 mM Cu(II) was added to the Cr(VI) solution, 58% of Cr(VI) was removed in the goethite experiment. Cu(II) cannot reduce Cr(VI), but has been shown to lead to recrystallization of Fe(III) (oxyhydr)oxides.³³ However, in the fifteen minute reaction time, the extent of recrystallization and further incorporation of Cr(VI) is likely minimal. Instead, Cu(II) is likely precipitating as Cu(II) hydroxide, which is supersaturated in this experiment according to aqueous speciation simulations using PHREEQC.⁴³ According to solubility calculations and previous studies, Cu(II) can precipitate as Cu(II) hydroxide and Cu(II) carbonate minerals, which control its concentration in the environment.^13,43,44

Cu(II) hydroxide likely precipitated in our experiment with Cr(VI) adsorbing to the precipitate. Cu(II) chromate can also precipitate in waters with elevated Cu and Cr concentrations,^12,13,34 so this process also may have removed Cr(VI) from solution, but Cu chromate is undersaturated according to PHREEQC calculations.⁴³ The Cr bound to the goethite surface as either Cr(VI) adsorbed to goethite or Cu(II) hydroxide contains a significantly higher percentage of Cr(VI). Both spectroscopic techniques signal that the addition of Cu(II) results in significantly more Cr(VI) contained on/within the solid (Fig. 2 and 3). While XPS suggests that a portion of Cr on the goethite surface is Cr(III) (Fig. 2), Cr(III) is nearly undetectable according to analyses of XANES spectra (Fig. 3 and Table S2). The Cr(VI) incorporated into the Cu precipitate appears to be less susceptible to reduction. Perhaps this Cr(VI) may be more fully contained in the Cu mineral structure, preventing most Cr(VI) reduction by adventitious carbon. Whereas Cr(VI) removal with goethite increased significantly with the addition of Cu(II), Cu(II) did not improve Cr(VI) removal with ferrihydrite (33% removal with Cu versus 29% without). Perhaps, ferrihydrite is preventing the formation of Cu(II) hydroxide, but further study would be required.

Fig. 3 XANES spectra of batch experiments with 250 µM Cr(VI) and goethite as well as 1 mM Fe(II) and 1 mM Cu(II) versus Cr(OH)₃ and Na₂CrO₄ standards. The Cr(VI) peak was clearly observed in the Cu(II) sample with a smaller Cr(VI) signal in the Cr(VI) only sample.

The addition of Fe(II) results in improved sequestration of Cr(VI) with 97% removal with goethite and 99% removal with ferrihydrite (Fig. 1). Fe(II) can directly reduce Cr(VI).^5,6,27 Near complete Cr(VI) reduction was confirmed by XPS and XAS with predominantly Cr(III) observed and Cr(VI) undetectable by both techniques (Fig. 2, 3 and Table S2). The Cr(VI) removal may also slightly improve through recrystallization of Fe minerals, incorporating more Cr.^19,23–26 Since Cr(VI) was not detected by XPS or XAS, Cr(VI) appears to have been reduced more rapidly than incorporated into the Fe mineral structure. Some Cr(III) may have been sequestered into the Fe mineral structure, but incorporation was likely limited given the short (15 minute) reaction time.

When combinations of Cu(II) and either an organic reductant or Fe(II) were added to the Cr(VI) solution, the removal of Cr(VI) remained similar to its behavior with only the reductant added. In a solution of Cu(II) and ascorbic acid, 97% of Cr(VI) was removed with goethite (versus 96% for only ascorbic acid) (Fig. 1A). With Cu(II) and Fe(II), 93% and 94% of Cr(VI) was removed with goethite and ferrihydrite, respectively (versus 96% and 99% for only Fe(II)). With Cu(II) and cysteine, 63% and 54% of Cr(VI) was removed with goethite and ferrihydrite (versus 40% and 52% removal with only cysteine). While Cr(VI) removal percentages are similar across experiments with and without Cu(II), Cr removal mechanisms may shift with the addition of Cu(II). Reaction with Cu(II) may result in the incorporation of Cr(VI) within Cu precipitates versus greater Cr(VI) reduction to Cr(III) without Cu(II).

Flow through goethite columns

To explore how Cr, Cu, Fe, and organics interact in a simple flowing system, which is relevant to groundwater flow or waste processing, column experiments were conducted with pure goethite, which could be compared to non-mixed systems in batch experiments. Since the column that incorporated cysteine clogged prior to the completion of the experiment, we primarily focused on the column using ascorbic acid. During the ascorbic acid experiment, flow through the columns remained steady throughout most of the experiment according to our chloride and bromide tracers (Fig. S2). The inflowing solutions, which alternated between bromide and chloride tracers, were changed every 30 minutes with samples collected at 10 minutes and 25 minutes following each change in solution chemistry (Fig. 4). Due to these dispersed sampling times, we unfortunately did not detect most of the increasing concentration of tracers. The initial pulse of bromide was higher than expected, possibly due to some fluctuating adsorption/desorption reactions as the columns stabilized. During each subsequent change to a bromide-containing solution, bromide reached ∼90% and nearly 100% of the concentration added in the samples collected 10 minutes and 25 minutes, respectively, of changing solutions (Fig. S2). The concentration of chloride remained a bit lower at ∼75% and ∼80% of the concentration added in the 10 minute and 25 minute samples, respectively, which may have been due to minor adsorption to goethite.⁵⁵ The concentrations of both bromide and chloride decreased to ∼15% and ∼5% of their maximum values after 10 minutes and 25 minutes, respectively, of switching to the opposite anion (Fig. S2). Therefore, the residence time of the columns was less than 10 minutes when the first samples were collected following a change in solution chemistry. In retrospect, this result should have been expected since the volume of the columns without goethite was approximately 25 mL, and the flow rate was set at 2.5 mL minute⁻¹. While the tracers do not provide definite evidence of residence time, we estimate the residence time at five minutes, which is how long it took for the columns to start flowing when peristaltic pumping began. The residence time did not appear to change throughout most of the experiment, but it may have slowed slightly near the end of the experiment due to some partial clogging. When switching to the final solution of Cr(VI) 180 minutes into the experiment, the concentration of bromide increased less than expected and then continued to decrease as the experiment ended.

Fig. 4 Cu, Fe, and Cr concentrations (measured by ICP-OES) as well as Cr(VI) concentrations (measured by UV-Vis) of samples collected at the base of the columns. Every 30 minutes, the aqueous chemistry passed through the goethite columns was changed to progressively more heavy metals/organics and then shifted back to less of these constituents to illustrate their collective effect on heavy metal fate and transport in water systems. The simple 0.25 mM Cr(VI) solution partially adsorbed to goethite in the columns. Cu and Cr precipitation occurred primarily in the solution vials prior to injection into the columns with no aqueous Cu observed in the vials. With the addition of ascorbic acid, aqueous Cu concentrations increased in the solution vials and was transported along with Cr through the goethite columns. Aqueous Fe(II) reacted with goethite and appeared to release adsorbed Cr(VI).

Effect of chemical addition on Cr sequestration and release over time in columns

At the start of the experiment, when 0.25 mM Cr(VI) was added to the goethite columns, initial aqueous Cr concentrations were low with most Cr(VI) adsorbing to the goethite. As more Cr(VI) adhered to goethite surfaces, a higher proportion of Cr(VI) advected through the columns. The average Cr concentrations were 0.05 mM at 10 minutes (2 pore volumes) and 0.14 mM at 25 minutes (5 pore volumes) into the experiment (Fig. 4 and S3). The Cr(VI) concentrations measured by UV-Vis spectroscopy (0.06 mM at 10 minutes and 0.15 mM at 25 minutes) were consistent with the aqueous Cr concentrations measured by ICP-OES, suggesting that aqueous Cr was entirely Cr(VI). The pH of collected water remained steady at 7.0, which did not change with the addition of Cu(II).

The inflow was changed to the Cu(II) containing solution vial 30 minutes into the experiment, which contained a bright yellow precipitate (Fig. S1). The concentration of aqueous Cr in the solution vial itself was only 0.06 mM (76% removal), and aqueous Cu was not even detected in the vial due to Cu precipitation. Since no Cu was remaining in solution, the concentration of aqueous Cu remained undetectable in the goethite columns. With less Cr remaining in solution, the aqueous Cr concentrations passing through the Cr columns decreased from 0.14 mM prior to Cu(II) addition to 0.07 and 0.05 mM (72% and 80% Cr removal) following Cu(II) addition (Fig. 4). Therefore, most Cr was already removed through Cu precipitation before the solution entered the columns. The Cr(VI) concentrations measured by UV-Vis were nearly identical to aqueous Cr concentrations at 0.08 and 0.06 mM.

Following 30 minutes in the Cr(VI) and Cu(II) solution (60 total minutes), the inflowing solution was moved to the Cr(VI), Cu(II), and ascorbic acid solution. With the addition of ascorbic acid, the pH of the collected solutions decreased slightly to 6.8. During this period, the concentration of aqueous Cr measured by ICP-OES remained low at 0.06 and 0.05 mM (76% and 80% Cr removal) at 70 minutes and 85 minutes (Fig. 4), which was consistent with the Cr concentration measured in the inflowing solution (0.05 mM). Surprisingly, Cr removal from solution did not increase in either the solution vial or the columns with the addition of ascorbic acid, a strong reducing agent. However, the Cr(VI) concentrations measured by UV-Vis were considerably lower at 0.01 mM throughout the addition of ascorbic acid (Fig. 4). This discrepancy between the ICP-OES and UV-Vis data suggests that most aqueous Cr passing through the column existed as Cr(III), not Cr(VI) as measured by UV-Vis. This mobile Cr component, which is likely Cr(III), demonstrates that Cr(VI) reduction may not result in complete Cr immobilization.

While Cu(II) was completely removed due to Cu precipitation in the Cr(VI) and Cu(II) solution, the solution vial containing Cr(VI), Cu(II), and ascorbic acid contained 0.07 mM Cu. With the addition of ascorbic acid, the solid precipitate in the vial changed colors from yellow to pink, indicating possible formation of Cu(0) (Fig. S1). Ascorbic acid can reduce Cu(II) to Cu(0),⁵⁶ which likely resulted in the yellow Cu(II)-containing precipitate transforming to Cu(0) nanoparticles. These nanoparticles could pass through the 0.45 µm filters, resulting in higher measured Cu concentrations. Cu also passed through the goethite columns with samples collected at 70 and 85 minutes at 0.07 mM Cu, the same concentration as measured in the solution vial. The concentrations of Cu and Cr emerging from the columns was nearly identical to the inflowing water from the solution vial, which suggests that little of the Cr(III) or Cu(0) adsorbed to the goethite. These trace metals are likely uncharged, which allowed them to pass through the goethite columns with little adsorption occurring.

From 90 minutes to 120 minutes into the experiment, the inflowing solution was shifted to the Cr(VI), Cu(II), ascorbic acid, and Fe(II) solution. With the addition of Fe(II), the aqueous Cr concentration decreased to 0.02 and 0.03 mM at 100 and 115 minutes, which was the same concentration of Cr as observed in the solution vial. Interestingly, the Cr(VI) concentration, as measured by UV-Vis, increased during this period from 0.01 mM to 0.03 mM. The aqueous Fe concentrations increased to 0.09 and 0.14 mM at 100 minutes and 115 minutes, respectively (Fig. 4). This concentration is lower than the 0.38 mM measured in the solution vial, so Fe(II) was possibly reacting with goethite in the columns. Fe(II) may have also adsorbed to goethite or was oxidized by dissolved oxygen or chromate within the column. The pH during this time dropped from 6.8 or 6.9 down to 6.5, which in concert with the low aqueous Fe concentrations suggests oxidation of Fe(II).

At 120 minutes, the solution was changed back to the Cr(VI), Cu(II), and ascorbic acid solution. With the removal of Fe(II), the pH increased slightly to 6.7 and the Fe(II) concentrations returned to near zero. The aqueous Cr concentrations increased slightly to 0.04 mM, but the Cr(VI) concentrations decreased to 0.01 mM (Fig. 4), which is consistent with the Cr chemistry of the inflowing solution.

At 150 minutes, the inflowing solution was shifted to the Cr(VI) and Cu(II) solution. With the removal of ascorbic acid, the pH more fully recovered to 6.9. The Cu concentrations decreased back to near zero (Fig. 4), as had been observed previously with the Cr(VI) and Cu(II) solution. The Cr concentrations decreased slightly to 0.02 mM (Fig. 4), which was lower than measured when the Cr(VI) and Cu(II) were initially added (0.06 mM Cr). The lower Cr concentrations are likely due to increased Cr(VI) adsorption. The Cr(VI) and Cu(II) solution had previously been added to the goethite column directly after the addition of the 0.25 mM Cr(VI) solution, which adsorbed to the goethite. This addition of Cr(VI) and Cu(II) came after the Fe(II)-containing solution, which mobilized adsorbed Cr(VI). Cr(VI) could adsorb to open adsorption sites when Cr(VI) concentrations increased.

At 180 minutes, the solution was shifted a final time to the Cr(VI) solution. With the removal of Cu(II), Cr concentrations continued to increase with concentrations recovering back to 0.12 mM at 190 minutes and 0.14 mM at 205 minutes, which was similar to that measured at 25 minutes (0.14 mM) when Cr(VI) was initially added to the column (Fig. 4). The reduction phase did not appear to drastically change Cr behavior after removal of the reductants. Cr concentrations decreased to 0.09 mM at 220 minutes as tracer concentrations were decreasing, which may have been due to clogging of the column.

Identification of Cr association to iron oxyhydroxide solids after column experiments

Electron microscopy analyses using STEM and electron microprobe illustrated how Cr could be associated with the goethite surface. According to STEM energy-dispersive X-ray spectroscopy (EDXS) maps performed on several goethite particles that were in contact with ascorbic acid or cysteine, the presence of Cr was seen in all cases. Although the amounts of Cr (and also Cu) were below the quantification limit, the characteristic peaks of both elements located at around 5.4 and 8 keV, respectively, can confirm their presence on the analyzed particles (Fig. 5). No particular zoning or preferred locations for Cr or Cu were noted within the collected maps (Fig. S6). Therefore, results suggest that Cr and Cu could have a relatively uniform distribution on the goethite surface.

Fig. 5 Annular darkfield (ADF) scanning transmission electron microscopy (STEM) micrographs accompanied by corresponding energy dispersive X-ray spectroscopy (EDXS) maps of different goethite particles collected from experiments with (A) 1 mM ascorbic acid and (B) 1 mM cysteine.

According to electron microprobe analyses (EMPA), Cr was above the quantification limit in most measured locations of goethite collected following the ascorbic acid experiment (Table 1). Goethite contained low levels of Cr across two subsamples with an average concentration of 0.02 weight percent Cr. Cr is detected in more measured points (10/12) than Cu (6/12), which is consistent with the fact than Cr flowed through the columns throughout most of the experiment, whereas Cu was only detected when ascorbic acid was included in the inflowing solution. Locations that contained higher Cr concentrations typically had detectable Cu (Table 1 and Fig. S4). This association of Cr and Cu may suggest the co-transport or co-precipitation of Cu and Cr within the goethite columns. During the addition of the Cr(VI) and Cu(II) solution, no Cu flowed through the columns since Cu precipitated as a yellow solid. Cu only flowed through the column with the addition of ascorbic acid, which likely reduced Cu(II) to Cu(0) generating Cu(0) colloids along with surface-associated Cr. A small portion of the Cu appears to have deposited on the goethite along with Cr.

Table 1 Electron microprobe analyses (EMPA) of Fe, O, Cr, and Cu content (mg kg⁻¹ or weight%) of goethite collected from the primary column experiment using ascorbic acid (BQL = below quantification limit)

Sample	Cr (mg kg^−1)	Cu (mg kg^−1)	Fe (wt%)	O (wt%)
Ascorbate sample 1-1	BQL	BQL	69.4	31.6
Ascorbate sample 1-2	140	BQL	33.5	60.5
Ascorbate sample 1-3	210	360	63.8	34.5
Ascorbate sample 1-4	330	BQL	61.0	37.0
Ascorbate sample 1-5	230	440	59.6	38.4
Ascorbate sample 1-6	140	390	53.6	41.8
Ascorbate sample 1-7	170	BQL	53.8	43.3
Ascorbate sample 1-8	300	530	62.5	35.9
Ascorbate sample 2-1	BQL	BQL	38.9	55.8
Ascorbate sample 2-2	60	BQL	46.0	49.9
Ascorbate sample 2-3	180	280	54.3	42.7
Ascorbate sample 2-4	170	280	46.2	49.7

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The goethite used in the cysteine experiment that clogged was also collected and analyzed using STEM and EMPA, so comparisons could be made between the cysteine experiment and ascorbic acid experiment. Due to clogging, the cysteine experiment ended prior to the full protocol, but all solutions including Cr, Cu, cysteine, and Fe were injected. However, the sequential removal of Fe, cysteine, and Cu from solution could not be completed, so collected goethite could have shown greater signs of these constituents than the fully completed ascorbic acid experiment. Instead, goethite from both experiments is indistinguishable. EDXS using TEM detected Cr and Cu at similar levels to the ascorbic acid experiment with Cr and Cu below the quantification limit, but the characteristic peaks were observed (Fig. 5B). The relative size and composition of all studied goethite particles (clusters) were nearly the same, with larger grains having a length between 1 and 1.3 µm (Fig. S7). Electron microprobe analyses of goethite from the cysteine experiment detects Cr in more than half of spots measured (9/16) with an average of 0.05 weight percent Cr in those locations (Table 2 and Fig. S5). Overall, Cr and Cu were near the quantification limit for both EMPA and TEM, but we managed to measure Cr successfully in most analyses.

Table 2 Electron microprobe analyses (EMPA) of Fe, O, Cr, and Cu content (mg kg⁻¹ or weight%) of goethite collected from the column experiment using cysteine

Sample	Cr (mg kg^−1)	Cu (mg kg^−1)	Fe (wt%)	O (wt%)
Cysteine sample 1-1	BQL	BQL	57.5	40.2
Cysteine sample 1-2	BQL	BQL	58.0	39.8
Cysteine sample 1-3	740	480	59.8	37.9
Cysteine sample 1-4	800	BQL	53.1	43.4
Cysteine sample 1-5	690	1340	60.7	36.8
Cysteine sample 1-6	900	720	59.0	38.2
Cysteine sample 1-7	710	1260	61.9	35.2
Cysteine sample 1-8	530	BQL	57.0	35.2
Cysteine sample 2-1	BQL	BQL	37.2	57.4
Cysteine sample 2-2	150	BQL	38.9	55.9
Cysteine sample 2-3	BQL	BQL	37.9	56.7
Cysteine sample 2-4	140	BQL	38.0	56.6
Cysteine sample 2-5	BQL	BQL	32.8	61.0
Cysteine sample 2-6	180	BQL	35.2	58.9
Cysteine sample 2-7	BQL	BQL	34.2	59.8
Cysteine sample 2-8	BQL	BQL	40.2	54.7

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Mechanistic insights

Waters co-contaminated with heavy metals and organic compounds can react through unexpected mechanisms, which control their mobility. Chromium chemistry is influenced by organic compounds and other metals through adsorption, redox, and precipitation reactions.

During batch experiments, the fraction of Cr removed was similar with the addition of only cysteine versus both cysteine and Cu(II). However, different reaction mechanisms affect Cr removal in these experiments. According to XPS and XANES spectroscopy analyses, cysteine reduces Cr(VI) to Cr(III), whereas Cu(II) results in incorporation of Cr(VI) in Cu(II) precipitates.

Cr(VI) may be sequestered by precipitation reactions, but subsequently released by reduction reactions. A Cr(VI)-containing Cu precipitate was observed in Cr(VI) and Cu(II) solutions (Fig. S1). Collected waters from batch and column experiments saw a decrease in Cr concentrations with the addition of Cu(II). Given the low solubility of many Cu precipitates,^12,13,34,44 Cu solids may remain stable for extended periods of time allowing incorporated Cr(VI) to remain sequestered. However, the precipitate transformed with shifts in redox conditions. Under the presence of a reducing organic compound (ascorbic acid), Cu and Cr concentrations increased as the metals were mobilized. Goethite collected from the columns had co-location of higher Cu and Cr concentrations according to EMPA. These Cr- and Cu-rich spots suggest that Cr and Cu were transported through the column and then co-deposited within the column under reducing conditions. Cr-containing Cu(II) precipitates were possibly reduced to mobile Cu(0) nanoparticles with a portion depositing on goethite and a portion advecting through the column. The clogging of goethite columns in the presence of cysteine suggests that the type of organic compound is critical in controlling heavy metal mobilization. Further research is required to determine which reductants induce Cu and Cr mobilization.

Environmental implications

While Cr(III) is typically considered more insoluble than Cr(VI), Cr may be more mobile under reducing conditions in complex chemical mixtures containing metals and organics. The presence of Cu or other heavy metals can induce precipitation reactions that remove Cr under oxidizing conditions. With the addition of Cu(II) to batch experiments and particularly column experiments, Cr concentrations decreased considerably. As Cu hydroxide precipitates, it can incorporate Cr(VI).^12,13,34 Natural waters with high Cu concentrations will likely form Cu solids since Cu can precipitate at neutral to basic pH.^12,13,34 According to solubility calculations, other heavy metals, such as lead and cadmium, also tend to form precipitates, which could incorporate Cr.⁴⁴ In industrial wastewaters near neutral pH, precipitation of these heavy metals may form Cr-containing solids that can be removed for resource recovery. If complex chemical mixtures are released into the environment under aerobic conditions, precipitation reactions could largely remove toxic metals through natural attenuation. However, precipitation of Cu and Cr may be influenced to some extent by the water chemistry of complex chemical mixtures or natural waters. According to solubility calculations, common cations, such as calcium and sodium, should not affect Cu or Cr precipitation.^43,44 Common anions have a larger effect with carbonate inducing Cu carbonate precipitation,^12,13,34 while sulfate can form aqueous Cu complexes that may slow Cu solids from forming.^43,44

With addition of reducing metals or organic compounds, precipitated Cr can be mobilized, even as reduced Cr(III). Organic compounds, such as ascorbate and citrate, can generate aqueous Cr(III)–organic complexes or mobile colloids as the concentration of organic compounds shifts.^29,31,36,37 During column experiments as ascorbic acid was added to a mixture of Cu(II) and Cr(VI), Cr concentrations (as measured by ICP-OES) increased slightly while Cr(VI) concentrations (as measured by UV-Vis) decreased (Fig. 4). Therefore, this Cr was likely Cr(III), which was transported through the columns. While Cr(III) is often considered immobile, a significant portion of Cr(III) may continue to be transported through a contaminated aquifer by complexing to organic compounds or sticking to colloid particles.^29,36,37 In this study, we have simplified some of these reactions using low molecular weight organic compounds, but the resulting chemical reactions may be applied to more complex natural organic matter and organic contaminants. Other reducing compounds and metals may not complex with Cr(III), but can reduce or recrystallize Fe oxyhydroxides, releasing adsorbed Cr in the process. Overall, an onset of reducing conditions can induce Cr release from environments co-contaminated by organics and metals. While reductive permeable reactive barriers have been shown to remove Cr(VI) from contaminated groundwater,^57,58 these barriers may be less effective for complex Cr mixtures where Cr can remain mobile as Cr(III) complexes or colloids.

Conclusions

Complex chemical mixtures of Cr, Cu, Fe, and organics can induce redox, complexation, and precipitation reactions, which are relevant for the recovery of critical minerals and the remediation of wastewater and contaminated groundwater. Under oxidizing conditions, Cu precipitates and removes Cr from solution. However, with the addition of certain organic reductants, Cu and Cr can be mobilized. Fe(III) solids adsorb Cr(VI) from solution, but Fe(II) can release this Cr into solution. While reducing conditions are typically thought to reduce aqueous Cr concentrations, reducing conditions may mobilize Cr in complex chemical mixtures.

Conflicts of interest

There are no conflicts of interest to declare.

Data availability

The data supporting this article have been included as part of the supplementary information (SI). Supplementary information: Cr(VI) concentrations for batch experiments, Cr fitting for XANES, photos of Cr solutions used in column experiments, chemical data from column experiments, and microprobe images and TEM elemental maps of goethite particles. See DOI: https://doi.org/10.1039/d5em00786k.

Acknowledgements

Funding was provided through the Army Research Office (ARO), Chemical Sciences Branch, Environmental Chemistry Research Area under contract W911NF-21-1-0249, the National Institute of Environmental Health Sciences (Superfund Research Program Award 1 P42 ES025589), and the National Science Foundation (awards 2125298, and 1914490). Any opinions, findings, and conclusions or recommendations expressed in this publication are those of the author(s) and do not necessarily reflect the views of the Army Research Office, the National Institutes of Health, or the National Science Foundation. Electron microscopy was carried out in the Nanomaterials Characterization Facility at the University of New Mexico, a facility that is supported by the State of New Mexico, the National Science Foundation and the National Aeronautics and Space Administration. The acquisition of the JEOL NEOARM AC-STEM at the University of New Mexico was supported by NSF grant DMR-1828731 and NASA Emerging Worlds grant 80NSSC21K1757.

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