Arsenic (As) and antimony (Sb) are potentially toxic trace elements that occur ubiquitously in the environment. The average crustal abundances in the lithosphere of As and Sb are 2 and 0.2 mg kg-1, respectively. From the lithosphere, they are released into ground- and surface water, soils, and the atmosphere by mineral weathering or during volcanic emissions. Thereby, As background concentrations are generally much higher than the corresponding values for Sb in soils and sediments (As ~6.0 mg kg-1vs.
Sb ~1 mg kg-1), while elevated concentrations are comparable.1
Arsenic is or was used industrially in the manufacture of numerous products like semiconductors, glass, paint as well as ceramics, and in the latter half of the 20th century as pesticide, herbicide and for wood preservation.1 Antimony also has a wide range of uses including the production of, among others, flameproof retardants, tracer bullets as well as automobile brake linings. Antimony is further used as catalyst in plastics synthesis.2 World production for As is considerably lower than that of Sb. In 2019 the estimated As mine output was 35,000 t compared to 140,000 t for Sb, whereby the latter steadily increased over the last 20 years.3
Environmental enrichments of both metalloids occur naturally in areas of geological mineralization, but also due to anthropogenic action.1,2 However, the majority of Sb contamination originates from mining and smelting as well as industrial emission sources including shooting ranges.4-7 In contrast, next to anthropogenic contamination, the best known exposure pathway of humans toward As is the consumption of As-rich drinking water from geogenic sources in South and Southeast Asian countries.8,9 Arsenic-contaminated drinking water is therefore considered the cause of “the largest poisoning of a population in history”.10 Based on large-scale groundwater surveys, up to 100 million people are exposed to drinking water with As levels higher than 10 μg As L-1, the threshold value set by the World Health Organization (WHO).9,11 The use of As-rich groundwater for the irrigation of rice paddies and the fact that rice is an efficient accumulator of As, with transfer of the metalloid up into the grains, adds an additional thread to millions of people.12 Further, the contamination of paddy soils from industrial and mining activity exposes people not only to a health risk of As but also of Sb.12,13
While it is known that long-term chronic exposure to As causes skin, kidney, lung, bladder cancer, and other severe diseases,14 to date no clear evidence exists15,16 for the carcinogenicity of chronic Sb exposure by the oral route.17 The derivation of guideline values for Sb is therefore somehow unclear with respect to As. The United States Environmental Protection Agency,18 Health Canada16 and the European Union19 use 6, 6, and 5 µg Sb L-1, respectively, while the WHO17 drinking water threshold for Sb is set to 20 µg Sb L-1. In general, for a thorough evaluation of the (eco-)toxicity of both elements, the exposure to specific As or
Introduction
Sb species in their respective redox state has to be considered, whereby trivalent species are generally considered more toxic than their pentavalent counterparts.2,20-24 Therefore, understanding the key factors and biogeochemical processes controlling the fate of As and Sb species in environmental systems is crucial for assessing the risks these metalloids pose to people living in contaminated regions around the world.
Both elements are members of group 15 in the periodic table and are situated below phosphorus (P), thus similarly having five valence electrons in a s2p3 electron configuration and possible oxidation states between –III and +V. Arsenic and Sb are chalcophile and siderophile and further have a high affinity for binding to oxygen (O). Therefore, they form many oxide, sulfide and mixed As/Sb-Fe oxide minerals at ambient pressure-temperature conditions, like As2O3 (arsenolite), FeAsO4*2H2O (scorodite), Sb2O3
(valentinite) and FeSbO4 (tripuhyite) and the sulfides FeAsS (arsenopyrite), As2S3 (orpiment), and Sb2S3
(stibnite).20,25
In aqueous solutions with environmentally relevant pH conditions (pH ~3-10), As forms the trivalent oxoacid arsenite (HxAsIIIO3x-3, x = 0-3, pKa1 = 9.17) under moderately reducing to moderately oxidizing conditions and the pentavalent arsenate (HxAsVO4x-3, x = 0-3, pKa1 = 2.30, pKa2 = 6.99, pKa3 = 11.80) under oxidizing conditions.26 Antimony reacts similar to As in aqueous solutions, however, still some ambiguity on the exact hydrolysis reactions exist.27-30 In general, under reducing conditions, the trivalent oxoacid antimonite (HxSbIIIO3x-3, x = 0-3, pKa1 = 11.82) is predicted to prevail, while at oxidizing conditions, the pentavalent antimonate (HxSbV(OH)6x-1, x = 0-1, pKa1 = 2.72) dominates the aqueous speciation. The octahedral configuration of Sb(V) compared to the tetrahedral coordination of P(V) or As(V) with O is explained by its larger ionic radius and lower charge density.31 Abiotic redox reactions of As and Sb species, respectively, which are thermodynamically feasible, are often slow and most often kinetically catalyzed directly or indirectly by microorganisms in the environment.32
Under reducing and circumneutral to alkaline pH conditions, As and Sb thioacids can form in S-rich environments by reaction of arsenite or antimonite with reduced sulfur (S) species.33-37 Thereby, thioarsenates (HxAsVS-IInO4-nx-3, n = 1-4, x = 0-3) are proposed to form in two steps from arsenite. First, at conditions of excess SH- over OH-, ligand exchange leads to the formation of thioarsenites (HxAsIIIS-IInO3-nx-3, n = 1-3, x = 0-3) as unstable intermediates.38,39 Second, addition of zero-valent S transforms thioarsenites to thioarsenates.40 Only the formation of monothioarsenate (HxAsVS-IIO3x-3; x = 0-3; MTAs(V)) does not require excess sulfide, because it forms directly from arsenite and zero-valent S or zero-valent-S-containing species (hereafter denoted as S(0)-species) as, for example, colloidal elemental S or polysulfides40 (Figure 1). Also, reactions of surface associated S(0)-species with arsenite are proposed to form MTAs(V).41 Thioantimonates were suggested to form in a similar way,37 but their occurrence and formation pathways are much less studied. A recent study, for example, proposed that direct SH-/OH- ligand exchange at the pentavalent antimonate, with a maximum at neutral pH, is also possible.42 By direct comparison, for example in geothermal systems, ~80% of total aqueous As were thioarsenates,
Introduction
whereas only ~40% of total Sb accounted for thioantimonates, suggesting higher affinity of As toward reduced S in aqueous solutions.37 Microorganisms, algae, plants, and animals produce volatile and non-volatile methylated As and Sb species as well as more complex structures like arsenobetaine, arsenocholine and arsenosugars.20,43 Abiotic reactions of sulfide with methylated As oxoacids can even form methylated thioarsenates under slightly acidic conditions in environmental systems.44,45 To date no observations about methylated thioantimonates have been made.
Figure 1: Overview of arsenic and antimony oxoacids and simplified model for formation of inorganic thioarsenates (adapted from Planer-Friedrich et al.40). For simplicity, all species are depicted fully protonated. Abbreviations:
MTAs(V) = monothioarsenate, DTAs(V) = Dithioarsenate, TriTAs(V) = trithioarsenate, TetraTAs(V) = tetrathioarsenate.
In environmental systems with low amounts of natural organic matter (NOM), aqueous As and Sb are primarily retained by adsorption reactions to the surfaces of fine-grained fractions of primary and secondary minerals. Sorption on mineral surfaces has therefore been extensively studied in the last decades,46,47 and thus, only the most important findings will be summarized in the following.
The major sorbents in those systems are iron (Fe), aluminum (Al) and manganese (Mn) (oxyhydr)oxides, which possess reactive ≡OH surface sites with variable charges mostly depending on pH and ionic strength as well as an inherent point of zero charge (pHPZC).48 Due to primarily electrostatic (Coulomb) interactions in outer-sphere complexations, negatively charged oxoanions, for example arsenate, adsorb strongly to positively charged goethite ≡OH2+ surfaces at low pH and sorption rapidly decreases with increasing pH (Figure 2). Competition with other negatively charged ions can strongly influence outer-sphere sorption.
Additionally, arsenate can form inner-sphere complexes with goethite. Arsenite stays uncharged over a
Introduction
wide pH-range (pKa1 = 9.17) but still has a high sorption affinity, which can mostly be attributed to inner-sphere surface complexation49 resulting in energetically favorable strong covalent bonds. The same mechanisms apply for antimonate and antimonite sorption to goethite.50
Figure 2: Comparison of arsenate and arsenite sorption edges on goethite as presented by Dixit and Hering49. The total arsenic concentrations shown are 100 µM (circles) and 50 µM (squares). Open symbols represent arsenate and closed symbols arsenite. Copyright (2020), with permission from American Chemical Society.
Especially for inner-sphere complexation, each As and Sb species has its own specific sorption characteristics to the different minerals. For example, Fendorf et al.51 suggested a monodentate (1V), a bidentate binuclear (2C), and a bidentate mononuclear (2E) complex for inner-sphere complexation of arsenate to goethite, while for arsenite, Manninget al.52 found a bidentate binuclear (2C) complex by use of extended X-ray absorption fine-structure (EXAFS) spectroscopy. Arsenite and arsenate can also form inner-sphere complexes to other Fe (oxyhydr)oxides,53 Al (oxyhydr)oxides54 and Mn oxides, however, As(III) will be oxidized to As(V) on the Mn oxide surfaces before sorption.55,56 Similar observations were made for antimonate, with mono- and bidentate inner-sphere surface complexation to Fe, Al and Mn (oxyhydr)oxides and for antimonate on Fe and Al (oxyhydr)oxides.7,57-60 Figure 3 summarizes possible geometries of surface complexes of arsenate and antimonate on Fe (oxyhydr)oxide surfaces.
Only little is known about the sorption behavior of thioarsenates on Fe and Al (oxyhydr)oxides.
Coutureet al.62 found MTAs(V) and TetraTAs(V) bound via a monodentate inner-sphere complex to ferrihydrite (Fh) and goethite, while in the presence of Fe sulfides, both thioarsenates where unstable and transformed to arsenite. Sorption experiments of MTAs(V) to amorphous Al oxyhydroxide revealed a bidentate binuclear complexation at higher pH and a partial reduction to arsenite at low pH.63 In each case, the extent of adsorbed As was lower for thioarsenates compared to arsenate or arsenite, which suggests a higher mobility of these species in the environment.62-64
Introduction
Figure 3: Polyhedral representations of arsenate (AsVO4) (A) and antimonate (SbVO6) (B) inner-sphere sorption geometries on a goethite model (geometries are analogous for AsIII/SbIIIO3 complexes). Sorption complexes are:1V = monodentate,2C = corner-sharing, bidentate binuclear,2E = edge-sharing bidentate mononuclear,3C = tridentate over structural vacancies. Models were adapted from Shermanet al.61 after Foster and Kim46 (A) and Scheinostet al.7 (B).
Copyright (2020), with permission from Elsevier.
While sorption to (oxyhydr)oxides is dominant at mostly oxic conditions, reductive dissolution of for example Fe (oxyhydr)oxide minerals,65 induced by soil flooding and microbial activity,66 can lead to As and Sb release.67,68 Consequentially, the metalloids can be mobilized into the aqueous phase or repartitioned between not reduced Fe (oxyhydr)oxides and freshly precipitated secondary mineral phases. In sulfate-rich environments, microbially mediated sulfate reduction generates dissolved sulfide,69 which can lead to the formation of Fe sulfide minerals like mackinawite (FeS)70 and finally pyrite (FeS2), with capacities for As/Sb incorporation, surface complexation, or surface precipitation.71-74 High levels of Fe(II) and dissolved sulfide may induce homogenous precipitation of As with Fe and S, resulting in the formation of arsenopyrite or As-rich (arsenian) pyrite.75 In environments where Fe(II) is titrated out by Fe sulfides, the remaining sulfide can precipitate at acidic to neutral pH as amorphous As/Sb sulfides76,77 and subsequently form more stable phases like orpiment, realgar or stibnite.20,25 However, reaction of dissolved sulfide with Fe (oxyhydr)oxides leads in part to sulfide oxidation to form solid-phase associated S(0)-species, as well as to overall Fe mineral transformation.78,79 At circumneutral to alkaline pH, previously adsorbed As can then react with reduced S (S(0)-species and sulfide) to form thioarsenates, which can lead to overall As mobilization in such systems.34,80-82 For Sb, however, sulfidization of Sb-bearing Fh only led to minor formation and mobilization of thioantimonates.83