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Arsenate
Arsenate
Arsenate
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Arsenate
Names
IUPAC name
Arsenate
Identifiers
3D model (JSmol)
ChemSpider
UNII
  • InChI=1S/AsH3O4/c2-1(3,4)5/h(H3,2,3,4,5)/p-3 checkY
    Key: DJHGAFSJWGLOIV-UHFFFAOYSA-K checkY
  • InChI=1/AsH3O4/c2-1(3,4)5/h(H3,2,3,4,5)/p-3
    Key: DJHGAFSJWGLOIV-DFZHHIFOAQ
  • [O-][As+]([O-])([O-])[O-]
Properties
AsO3−4
Molar mass 138.918 g·mol−1
Conjugate acid Arsenic acid
Hazards
Occupational safety and health (OHS/OSH):
Main hazards
 Extremely toxic, carcinogenic
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).

The arsenate is an ion with the chemical formula AsO3−4.[1] Bonding in arsenate consists of a central arsenic atom, with oxidation state +5, double bonded to one oxygen atom and single bonded to a further three oxygen atoms.[2] The four oxygen atoms orient around the arsenic atom in a tetrahedral geometry.[2] Resonance disperses the ion's −3 charge across all four oxygen atoms.

Arsenate readily reacts with metals to form arsenate metal compounds.[2][3] Arsenate is a moderate oxidizer and an electron acceptor, with an electrode potential of +0.56 V for its reduction to arsenite.[4] Due to arsenic having the same valency and similar atomic radius to phosphorus, arsenate shares similar geometry and reactivity with phosphate.[5] Arsenate can replace phosphate in biochemical reactions and is toxic to most organisms.[5][6]

Natural occurrence

[edit]
Adamite, a naturally occurring arsenate mineral.

Arsenates occur naturally, in hydrated and anhydrous form, in a variety of minerals. Examples of arsenate-containing minerals include adamite, alarsite, annabergite, erythrite and legrandite.[7] When two arsenate ions balance the charge in a formula, it is called diarsenate for example zinc diarsenate, Zn3(AsO4)2.

Uses

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Arsenate-based pesticides such as lead hydrogen arsenate were commonly used until their replacement by newer pesticides such as DDT and subsequent ban by multiple regulatory bodies due to health concerns.[8][9]

Transition metal arsenate compounds are often brightly coloured and have been used to make pigments. Copper arsenate was a minor compound used in the Egyptian blue pigment used by the ancient Egyptians and Romans.[10] Cobalt violet pigment was made from cobalt arsenate before its toxicity led to its replacement by cobalt phosphate.[11][12][13]

Chromated copper arsenate (CCA) has been a widely used wood preservative since the 1930s.[14] Safety concerns have led to the phasing out of CCA-treated wood for residential projects in many countries.[14] CCA remains a common and economical treatment choice for non-residential uses such as agriculture. [14][15]

Speciation

[edit]
Pourbaix diagram showing the distribution of arsenate and arsenite species in water. Oxygenated waters have a high pe value and arsenate species dominate. In deoxygenated water, with low pe, arsenite species dominate.[16][17]

Depending on the pH, arsenate can be found as trihydrogen arsenate (that is arsenic acid H3AsO4), dihydrogen arsenate (H2AsO−4), hydrogen arsenate (HAsO2−4), or arsenate (AsO3−4).[18] Trihydrogen arsenate is also known as arsenic acid. At a given pH, the distribution of these arsenate species can be determined from their respective acid dissociation constants.[17]

H3AsO4 + H2O ⇌ H2AsO−4 + [H3O]+ (pKa1 = 2.19)
H2AsO−4 + H2O ⇌ HAsO2−4 + [H3O]+ (pKa2 = 6.94)
HAsO2−4 + H2O ⇌ AsO3−4 + [H3O]+ (pKa3 = 11.5)

These values are similar to those of phosphoric acid. Hydrogen arsenate and dihydrogen arsenate predominate in aqueous solution near neutral pH.[17]

The reduction potential (pe) of a solution also affects arsenate speciation. In natural waters, the dissolved oxygen content is the main factor influencing reduction potential. Arsenates occur in oxygenated waters, which have a high pe, while arsenites are the main arsenic species in anoxic waters with a low pe.[16]

A Pourbaix diagram shows the combined influence of pH and pe on arsenate speciation.

Contamination

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Arsenates, along with arsenites, are a significant source of contamination in some natural water sources and can lead to arsenic poisoning with repeated exposure.[19][20] Countries with high levels of arsenic minerals in sediment and rock, such as Bangladesh, are especially at risk of arsenate contamination.[21][20]

Arsenate poisoning

[edit]

Arsenate is harmful to humans and animals as it interferes with the normal functioning of glycolysis and the Krebs cycle. Arsenate replaces inorganic phosphate in the step of glycolysis that produces 1,3-bisphosphoglycerate from glyceraldehyde 3-phosphate. This yields 1-arseno-3-phosphoglycerate instead, which is unstable and quickly hydrolyzes, forming the next intermediate in the pathway, 3-phosphoglycerate. Therefore, glycolysis proceeds, but the ATP molecule that would be generated from 1,3-bisphosphoglycerate is lost – arsenate is an uncoupler of glycolysis, explaining its toxicity.[22][23]

As with other arsenic compounds, arsenate binds to lipoic acid, inhibiting the conversion of pyruvate into acetyl-CoA, blocking the Krebs cycle and therefore resulting in further loss of ATP.[23]

See also

[edit]

References

[edit]
Revisions and contributorsEdit on WikipediaRead on Wikipedia

Arsenate

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from Grokipedia
Arsenate is the AsO₄³⁻, the trianionic conjugate base formed by of (H₃AsO₄), a pentavalent arsenic species with tetrahedral geometry analogous to the phosphate ion (PO₄³⁻). Arsenate ions form salts with various metals, exhibiting properties similar to phosphates, and occur naturally in minerals such as adamite and in elevated concentrations in certain groundwaters due to geochemical processes. Historically, arsenate compounds like lead arsenate and calcium arsenate served as insecticides and herbicides in agriculture, particularly for and orchards, though their use has largely been curtailed owing to environmental persistence and health risks. In aqueous environments, arsenate's speciation depends on and conditions, as depicted in Pourbaix diagrams, influencing its mobility and in soils and waters. Biochemically, arsenate's phosphate mimicry allows competitive inhibition in energy-transfer reactions, such as substituting in ATP synthesis to form unstable arsenylated intermediates that hydrolyze spontaneously, thereby uncoupling and disrupting . This interference, coupled with induction of and , underpins arsenate's acute and chronic toxicity, including risks of cancer in , , and upon exposure.

Chemical Fundamentals

Definition and Structure

Arsenate refers to chemical compounds containing the arsenate anion, AsO₄³⁻, or the anion itself, which is derived from (H₃AsO₄) by the removal of three protons. The anion features a central atom in the +5 bonded to four oxygen atoms, resulting in a net charge of -3. This structure positions arsenate as the arsenic analog of the ion (PO₄³⁻), sharing similar ionic radii and coordination chemistry that enable it to substitute for phosphate in certain biochemical contexts. The of the arsenate ion is tetrahedral, with the atom at surrounded by four oxygen atoms at the vertices, exhibiting bond angles close to the ideal tetrahedral value of 109.5°. In crystalline arsenate salts, As-O bond lengths typically range from 1.65 to 1.70 Å, reflecting the high and small size of arsenic(V) which favors strong covalent-ionic bonding with oxygen. This tetrahedral configuration contributes to the stability of arsenate in aqueous solutions at neutral to basic , where the fully deprotonated form predominates, though protonated species like HAsO₄²⁻ and H₂AsO₄⁻ exist under acidic conditions.

Physical Properties

Arsenate salts, which incorporate the tetrahedral AsO₄³⁻ anion, typically manifest as colorless or white crystalline solids that are odorless and prone to upon exposure to air. These properties arise from the ionic nature of the compounds, with hydration states influencing stability and handling. Aqueous solutions of arsenate ions are colorless and transparent, lacking any distinctive taste or smell, consistent with the absence of volatile components under standard conditions. Key physical parameters for representative salts, such as sodium arsenate (Na₃AsO₄), include a of approximately 1.75–1.87 g/cm³ and high solubility exceeding 30 g/100 mL at ambient temperatures. occurs around 86°C for the anhydrous form, though hydrates often dehydrate or decompose below 150°C rather than yielding a true melt.
PropertyValue for Na₃AsO₄ (anhydrous)Source
Density1.75 g/cm³
86°C (decomposes above 150°C)
Solubility in Highly soluble (>30 g/100 mL)
Solubility in Slightly soluble
For disodium hydrogen arsenate heptahydrate (Na₂HAsO₄·7H₂O), a common variant, the density is 1.9 g/cm³, with solubility at 39 g/100 mL in water at 21°C and thermal decomposition initiating near 130°C. These characteristics render arsenate salts denser than water, causing them to sink in aqueous environments, and insoluble in non-polar solvents like diethyl ether.

Chemical Reactivity and Stability

Arsenate ions (AsO₄³⁻) exhibit pH-dependent in aqueous solutions, predominating as H₃AsO₄ at pH below 2.2, H₂AsO₄⁻ between pH 2.2 and 6.9, HAsO₄²⁻ between pH 6.9 and 11.5, and AsO₄³⁻ above pH 11.5, reflecting the acidity constants of . This speciation influences reactivity, with dihydrogen and monohydrogen forms facilitating and adsorption reactions more readily than the fully deprotonated . Arsenate acts as a moderate , capable of accepting electrons to form (AsO₃³⁻) under reducing conditions, as evidenced by its role in transformations in environmental systems. Precipitation reactions are prominent, with arsenate forming insoluble salts such as silver arsenate under neutral to alkaline conditions and ferric arsenate compounds like scorodite (FeAsO₄·2H₂O) in acidic media, which serve for arsenic stabilization. These precipitates arise from coordination of arsenate's oxygen atoms to metal cations, often via inner-sphere complexation on oxide surfaces. In contrast, arsenate esters, formed by reaction with alcohols or in biochemical analogs to esters, display low stability, undergoing rapid in water; for instance, trimethyl arsenate has a of approximately 0.02 seconds at neutral and . Stability of arsenate species varies with environmental factors. Arsenate minerals and solids, such as hydrous ferric arsenate, maintain structural integrity between 3 and 8, but dissolve at lower due to proton-promoted release of and at higher from competition. Thermally, arsenate compounds like solutions remain stable at ambient temperatures, though elevated heat (e.g., 80–95°C) can induce transformation of amorphous ferric arsenate to crystalline scorodite in media at 1–2. In oxidizing environments, as depicted in Pourbaix diagrams, arsenate predominates over , underscoring its thermodynamic favorability in aerobic, neutral to alkaline conditions.

Sources and Production

Natural Occurrence

Arsenate occurs naturally in secondary minerals formed by the oxidation of primary arsenic sulfides, such as (FeAsS), during processes in the . These minerals typically develop in the oxidized zones of deposits and include scorodite (FeAsO₄·2H₂O), a hydrous iron arsenate common in altered ores and mine environments, as well as arseniosiderite and pharmacosiderite, which are iron-bearing arsenates associated with similar processes. Arsenate minerals like these precipitate under acidic to neutral conditions in arsenic-enriched settings, often alongside analogs due to structural similarities. In natural waters, arsenate (As(V)) dominates under oxidizing conditions, such as in oxygenated surface waters, marine environments, and aerated groundwaters, where it exists primarily as HAsO₄²⁻ or H₂AsO₄⁻ depending on . In seawater, arsenate concentrations increase below the due to the remineralization of arsenic-bearing , reaching levels up to several micrograms per liter in deep waters. In freshwater and , arsenate prevails in oxic regimes, though it coexists with (As(III)) in transitional zones; natural levels in such systems range from <0.5 μg/L to over 5,000 μg/L in geologically arsenic-rich areas like volcanic terrains or sedimentary aquifers. Soils derived from arsenic-bearing rocks contain arsenate adsorbed onto iron and manganese oxides, with concentrations varying by geological history—typically 1–40 mg/kg in uncontaminated soils but higher in mineralized regions. This adsorption enhances stability in oxic soils, though desorption can occur under changing environmental conditions like flooding or phosphate competition. Volcanic activity and geothermal fluids also contribute arsenate to surface and subsurface waters through high-temperature oxidation.

Industrial Synthesis

Arsenic trioxide (As₂O₃), the primary intermediate in arsenic production, is obtained as a volatile byproduct during the smelting of copper, lead, cobalt, and gold ores, where it is captured from flue gases and purified by sublimation or condensation. This compound serves as the starting material for arsenate synthesis, as direct extraction of arsenates from ores is uneconomical. Global production of arsenic trioxide has historically exceeded 50,000 metric tons annually, with major contributors including China, which accounted for the largest share as of 2003, though output varies with mining activity. Arsenic acid (H₃AsO₄), the precursor to most arsenates, is industrially produced by oxidizing arsenic trioxide with concentrated nitric acid in an aqueous medium: As₂O₃ + 2HNO₃ + 3H₂O → 2H₃AsO₄ + 2NO. Alternative oxidants like hydrogen peroxide have been explored but pose explosion risks in concentrated solutions, limiting their commercial adoption. The reaction is typically conducted under controlled heating to 80–100°C to ensure complete oxidation to the pentavalent state, yielding arsenic acid solutions or crystals upon evaporation, with nitric acid recycled via NO oxidation. Purity levels reach 99% or higher for industrial grades used in downstream arsenate formation. Sodium arsenate (Na₃AsO₄ or Na₂HAsO₄ variants) is manufactured by neutralizing arsenic acid with sodium hydroxide or carbonate: H₃AsO₄ + 3NaOH → Na₃AsO₄ + 3H₂O. In processes utilizing high-arsenic dust from metallurgy, alkaline leaching with NaOH directly forms crude sodium arsenate, followed by filtration, purification via recrystallization, and drying to produce dodecahydrate crystals (Na₃AsO₄·12H₂O). This method achieves high yields, with arsenic recovery exceeding 99% under optimized conditions of pH 10–12 and temperatures around 80°C. Historically significant for pesticide production, such synthesis has declined due to regulatory restrictions, though it persists for niche chemical applications. Lead arsenate, once a major arsenate product for insecticides, was prepared industrially by reacting arsenic acid with litharge (PbO): 3PbO + 2H₃AsO₄ → Pb₃(AsO₄)₂ + 3H₂O, or via metathesis of sodium arsenate with lead acetate to precipitate the sparingly soluble basic form Pb₅(OH)(AsO₄)₃. Catalysts like acetic acid were sometimes employed to accelerate precipitation and control particle size for sprayable formulations. Production peaked in the early 20th century, with U.S. output tied to agricultural demand, but ceased commercially by the mid-20th century amid toxicity concerns and alternatives like DDT. Modern remnants involve similar precipitation for waste stabilization rather than active synthesis.

Environmental Dynamics

Speciation and Transformation

Arsenic speciation in environmental systems primarily involves inorganic forms arsenite (As(III)) and arsenate (As(V)), with the latter dominating in oxic aqueous environments due to its greater thermodynamic stability under oxidizing conditions. Speciation is controlled by redox potential () and , as illustrated in Pourbaix diagrams, where As(V) prevails at Eh > 100 mV and neutral to alkaline pH, while As(III) forms under reducing conditions (Eh < 0 mV). Methylated species, such as monomethylarsonic acid (MMA) and dimethylarsinic acid (DMA), arise from microbial transformations and constitute minor fractions in uncontaminated systems but increase in biologically active sediments. Redox transformations between As(III) and As(V) occur abiotically and biotically, with microbial mediation accelerating rates under natural conditions. Abiotic oxidation of As(III) to As(V) proceeds slowly via dissolved oxygen but is enhanced by manganese and iron oxides, with half-lives ranging from hours to days depending on mineral surfaces and pH (optimal at 4-7). Biotic oxidation, performed by As(III)-oxidizing bacteria such as those in periphyton communities, couples to chemolithoautotrophy and predominates in aerobic interfaces, reducing As mobility by forming less soluble As(V). Conversely, dissimilatory arsenate reduction to As(III) by anaerobic bacteria, using As(V) as a terminal electron acceptor, mobilizes arsenic in sediments and groundwater, with rates increasing with temperature from 0 to 40°C and linked to organic carbon availability. pH influences both speciation and transformation kinetics; acidic conditions (pH < 5) favor As(V) protonation to H2AsO4^- and sorption to minerals, limiting transformation, whereas neutral to alkaline pH promotes desorption and microbial activity. In soils and aquifers, coupled sulfur oxidation can drive As(V) reduction, exacerbating contamination in rice paddies where flooded conditions lower Eh to -150 mV, shifting speciation toward bioavailable As(III). Microbial consortia facilitate cooperative detoxification via reduction followed by methylation, though demethylation recycles inorganic forms, underscoring the dynamic biogeochemical cycling of arsenic.

Mobility in Ecosystems

Arsenate, the dominant arsenic species (As(V)) in oxic environments, exhibits low mobility in soils and sediments due to strong adsorption onto iron and aluminum oxides, manganese oxides, and clay minerals, which limits its transport and bioavailability. Adsorption is governed by ligand exchange mechanisms, with ferrihydrite and goethite showing particularly high affinity, often retaining over 90% of arsenate at concentrations below 10 mg/L in neutral pH soils. This immobilization is pH-dependent, peaking between pH 4-7 and declining above pH 8 due to increased electrostatic repulsion from deprotonated surface sites and arsenate anions. In groundwater and surface waters, arsenate mobility increases under reducing conditions prevalent in anoxic aquifers or flooded sediments, where microbial respiration reduces As(V) to more soluble and weakly adsorbing arsenite (As(III)), potentially mobilizing up to 50-100% of sorbed arsenic through reductive dissolution of iron oxides. Competing ions such as phosphate can desorb arsenate by site competition, enhancing leaching in agricultural soils amended with fertilizers, while organic matter may complex arsenate or indirectly promote reduction via electron donors. In oxic river systems, however, adsorption to suspended particulates facilitates sedimentation rather than long-range transport. Ecosystem-specific factors further modulate arsenate dynamics; in floodplain soils subject to periodic flooding, alternating redox cycles can release bioavailable fractions, with studies reporting elevated dissolved arsenic during anaerobic phases exceeding 0.05 mg/L. Bioaccumulation in aquatic biota is thus linked to desorbed arsenate in porewaters, though plant uptake in soils remains low (typically <1% of total soil arsenic) unless facilitated by high phosphate levels or low iron content.

Practical Applications

Historical and Agricultural Uses

Lead arsenate (PbHAsO₄), an arsenate salt, was first applied as an insecticide in U.S. apple orchards during the 1890s to control the codling moth (Cydia pomonella), a major pest damaging fruit crops. Its adoption expanded rapidly in the early 20th century, becoming the dominant pesticide for deciduous tree fruits such as apples and cherries, with widespread use persisting until the mid-1940s in regions like Washington state, where applications occurred from 1905 to 1947. By 1919, annual U.S. production of lead arsenate reached 11.5 million pounds, peaking at over 37 million pounds by 1931, reflecting its efficacy against chewing insects but also its accumulation in soils and residues on produce. Calcium arsenate emerged as another key arsenate-based pesticide around 1910, particularly for combating the cotton boll weevil (Anthonomus grandis) in southern U.S. agriculture, supplanting earlier arsenicals like Paris green due to lower phytotoxicity. It was dusted over cotton fields and other crops, with production scaling alongside lead arsenate to address insect outbreaks; by the 1920s–1950s, arsenates including calcium variants were institutionalized in European potato farming to protect against pests like the Colorado potato beetle. These compounds functioned by disrupting insect nervous systems upon ingestion, providing broad-spectrum control but requiring repeated applications that elevated arsenic levels in agricultural soils. Arsenates also served as herbicides and cotton desiccants historically, with sodium arsenate and related forms applied to defoliate crops pre-harvest, though their use declined sharply after the 1940s introduction of synthetic organics like DDT and subsequent organophosphates, which offered greater specificity and reduced persistence. By 1960, most lead and calcium arsenate applications in fruit and cotton production were phased out in the U.S. due to documented soil contamination, human poisoning incidents from residues, and regulatory scrutiny over chronic toxicity, though legacy residues persist in former orchard sites. In Mexico, arsenate pesticides targeted cotton pests as early as 1896, illustrating early hemispheric adoption for staple crop protection.

Modern Industrial Applications

In contemporary industry, arsenate compounds, particularly chromated copper arsenate (CCA), continue to be employed as wood preservatives for non-residential applications, including utility poles, marine pilings, and highway structures, where durability against fungal decay and insect damage is critical. This usage accounts for approximately 29% of global arsenic consumption, primarily in As(V) form within CCA formulations, despite voluntary phase-outs for residential lumber in the United States and Canada since 2003 due to environmental and health risks. Regulatory approvals persist for these industrial contexts, supported by data demonstrating efficacy in high-exposure environments, though alternatives like alkaline copper quaternary are increasingly adopted. Water-soluble arsenates, such as sodium arsenate, serve as intermediates in the synthesis of other arsenic-based chemicals, including those for metallurgical processes and specialty reagents. In analytical chemistry, they function as standards or precipitants for detecting metals, though at limited scales compared to historical volumes. Overall, arsenate applications have declined sharply since the mid-20th century, with production shifting toward less toxic substitutes amid stringent environmental regulations, yet niche industrial roles remain due to their unique oxidative and preservative properties.

Toxicological Profile

Biochemical Mechanisms

Arsenate (AsO₄³⁻), structurally analogous to phosphate (PO₄³⁻), disrupts cellular energy metabolism by competing for uptake via phosphate transporters and substituting in phosphate-dependent reactions. This interference primarily manifests in mitochondria and glycolytic pathways, leading to inefficient ATP production and cellular energy depletion. In oxidative phosphorylation, arsenate enters mitochondria through phosphate carriers and reacts with ADP to form unstable ADP-arsenate complexes in submitochondrial particles. Unlike ATP, ADP-arsenate hydrolyzes spontaneously without yielding usable energy, allowing electron transport and oxygen consumption to proceed uncoupled from ATP synthesis. This mechanism, observed in isolated mitochondria, results in accelerated respiration without corresponding energy conservation, contributing to arsenate's acute toxicity.69115-5/pdf) During glycolysis, arsenate substitutes for phosphate in the glyceraldehyde-3-phosphate dehydrogenase (GAPDH) reaction, producing 1-arseno-3-phosphoglycerate instead of 1,3-bisphosphoglycerate. The arsenate ester hydrolyzes rapidly and non-enzymatically to 3-phosphoglycerate and free arsenate, circumventing the substrate-level phosphorylation step catalyzed by phosphoglycerate kinase and forgoing ATP generation. This futile cycle depletes glycolytic intermediates and exacerbates energy deficits, particularly under phosphate-limiting conditions where arsenate uptake increases.

Health Effects and Exposure

Inorganic arsenic species, including arsenate (AsO₄³⁻), primarily enter the human body through ingestion of contaminated drinking water and food, with rice and seafood being notable dietary sources due to bioaccumulation. Occupational exposure occurs mainly via inhalation of dust or fumes in industries like mining, smelting, and pesticide production, while dermal absorption is minimal and rarely contributes significantly to systemic toxicity. Globally, groundwater contamination from natural geological sources accounts for the majority of widespread exposure, affecting over 140 million people as of 2022, particularly in regions like South Asia and parts of Latin America. Acute exposure to high levels of arsenate, typically from accidental or intentional ingestion exceeding 100-300 mg, manifests within hours as severe gastrointestinal distress including vomiting, abdominal pain, profuse watery diarrhea (often described as "rice water" stools), and dehydration, potentially progressing to hypovolemic shock, hepatic injury, and multi-organ failure if untreated. Neurological symptoms such as numbness, tingling in extremities, and muscle cramps may follow, with cardiovascular collapse and death possible in severe cases without prompt chelation therapy using agents like dimercaprol. Children exhibit heightened vulnerability, showing similar symptoms at lower doses due to lower body mass and immature detoxification pathways. Chronic exposure to arsenate at levels as low as 10-50 μg/L in drinking water over months to years leads to dermatological changes such as diffuse hyperpigmentation, spotted pigmentation (raindrop pattern), and hyperkeratotic lesions on palms and soles, which serve as early biomarkers. Systemic non-cancer effects include peripheral neuropathy characterized by symmetric sensory loss and painful paresthesias, increased risk of type 2 diabetes via interference with insulin signaling, and cardiovascular diseases such as hypertension and peripheral vascular disease (e.g., blackfoot disease in high-exposure areas like Taiwan). Arsenate's biochemical mechanism involves competitive inhibition of phosphate-utilizing enzymes, disrupting ATP synthesis and cellular respiration, which underlies much of the oxidative stress and genotoxicity observed. Epidemiological evidence establishes inorganic arsenic, including arsenate, as a human carcinogen, with dose-dependent increases in skin, lung, and bladder cancers; for instance, cohorts in Bangladesh exposed to 100-500 μg/L arsenic in water showed standardized incidence ratios for lung cancer exceeding 2.0 after 10-15 years. Liver and kidney cancers show associations but less consistent causation due to confounding factors like viral hepatitis. Pregnant women face risks of adverse birth outcomes, including low birth weight and spontaneous abortion, linked to placental transfer of arsenate. Overall, pentavalent arsenate is less acutely toxic than trivalent arsenite but contributes comparably to chronic effects after metabolic reduction in vivo.

Regulatory and Remediation Aspects

Exposure Limits and Standards

The permissible exposure limit (PEL) for inorganic arsenic, including arsenate compounds, in occupational settings is set by the (OSHA) at 10 micrograms per cubic meter of air (µg/m³) as an 8-hour time-weighted average (TWA). An action level of 5 µg/m³ triggers requirements for exposure monitoring, medical surveillance, and training for workers exposed at or above this threshold for 30 or more days per year. The (NIOSH) recommends a lower exposure limit of 10 µg/m³ as a 10-hour TWA, with an immediately dangerous to life or health (IDLH) concentration of 60 mg/m³. For drinking water, the U.S. Environmental Protection Agency (EPA) established a maximum contaminant level (MCL) of 10 µg/L (parts per billion) for in 2001, effective from 2006, based on cancer risk assessments and cost-benefit analysis. The World Health Organization (WHO) guideline value for in drinking water is also 10 µg/L, designated as provisional due to practical quantification limits and ongoing evaluation of long-term risks below this level. These standards apply to total inorganic , encompassing as the dominant species in oxygenated waters.
Agency/StandardMediumLimitBasis
OSHA PELWorkplace air10 µg/m³ (8-hour TWA)Carcinogenicity and respiratory effects
EPA MCLDrinking water10 µg/LLifetime cancer risk of 1 in 10,000
WHO GuidelineDrinking water10 µg/L (provisional)Health-based, with margin for analytical challenges
Food standards for inorganic arsenic are less prescriptive; the U.S. Food and Drug Administration (FDA) sets action levels, such as 100 ppb in infant rice cereals, to guide mitigation rather than enforce strict limits, reflecting variable dietary exposure risks. These limits derive from toxicological data indicating no safe threshold for arsenic carcinogenicity, prioritizing reduction where feasible.

Contamination Mitigation

Arsenate, predominantly existing as the oxyanion AsO₄³⁻ under oxidizing conditions, is mitigated through adsorption processes utilizing iron oxide-based media such as granular ferric hydroxide (GFH), which achieves removal efficiencies exceeding 99% from contaminated water at influent concentrations around 10 μg/L, owing to ligand exchange mechanisms where arsenate binds to surface hydroxyl groups. Activated alumina sorbents similarly target As(V) with capacities up to 41.4 mg/g, performing optimally at pH 5-7 where electrostatic repulsion is minimized, though phosphate ions can compete and reduce efficacy by 20-50%. These adsorbents are regenerated via caustic or acid washes, but spent media requires disposal as hazardous waste due to concentrated arsenic residuals. Coagulation-flocculation with ferric chloride or alum salts precipitates arsenate as ferric arsenate (FeAsO₄), attaining 90-99% removal at pH 6-8 and coagulant doses of 20-40 mg/L Fe, particularly effective post-oxidation of any co-occurring arsenite to enhance overall arsenic capture below 10 μg/L. This method generates sludge volumes of 0.5-2% of treated water, necessitating dewatering and stabilization to prevent re-leaching, with full-scale applications demonstrating effluent levels under 0.010 mg/L in 19 of 68 evaluated systems. Sulfide precipitation provides an alternative for arsenic removal from low-concentration wastewater (0.8-1.5 mg/L), where addition of flocculants such as polyacrylamide (PAM) or coagulants destabilizes colloids and promotes settling to improve efficiency. Complementary strategies include iron or aluminum salt coagulation-adsorption following oxidation of As(III) to As(V), with Fe/As molar ratios of 5-10 enabling residual levels below 0.01 mg/L, or adsorption using activated alumina, iron-modified sorbents, or arsenic-specific resins. Ion exchange resins, such as sulfate-selective or cerium-impregnated types, selectively bind arsenate anions with capacities up to 53 mg/g and >99% , but suffer from by sulfates or silica, requiring frequent regeneration every 4 weeks and limiting applicability to low-turbidity waters. Membrane technologies like and nanofiltration reject arsenate via size exclusion and charge effects, yielding 93-99% removal for As(V) at pressures of 5-10 bar, outperforming for charged species compared to neutral As(III) without pretreatment. range $0.50-2.00 per 1,000 gallons treated, with operational challenges including 10-20% water rejection and mitigated by pre-coagulation. In soils, stabilization with amendments like zero-valent iron or phosphates induces arsenate precipitation as sparingly soluble phases (e.g., Ca₃(AsO₄)₂), reducing leachability to <5.0 mg/L via Toxicity Characteristic Leaching Procedure (TCLP) in 37 of 44 full-scale solidification projects using cement or lime binders at $60-290 per ton. Phytoremediation employs hyperaccumulators such as Pteris vittata ferns, achieving bioconcentration factors of 8-320 for soil arsenic, extracting up to 265 mg/kg dry biomass in shallow contaminated zones (<1 m depth), though limited by slow growth and biomass disposal needs at $60,000-100,000 per acre. Electrokinetic methods apply low-voltage fields to mobilize and extract arsenate from low-permeability soils, reducing concentrations from >250 mg/kg to <30 mg/kg in full-scale trials, at costs of $50-270 per cubic yard. Hybrid approaches, combining oxidation (e.g., permanganate dosing for >86% As(III) to As(V) conversion in 1 minute) with downstream adsorption or , optimize for mixed-valence , as demonstrated in treatment trains achieving <0.005 mg/L effluents across variable pH and co-contaminant matrices. Regulatory compliance drives adoption, aligning with U.S. EPA's 2006 maximum contaminant level of 10 μg/L for drinking water, though long-term efficacy depends on site-specific geochemistry and residuals management to avert secondary contamination.

Key Controversies

Claims of Biological Incorporation

In December 2010, researchers led by Felisa Wolfe-Simon reported the isolation of a halophilic bacterium, strain GFAJ-1 from Mono Lake, California, which purportedly substituted (AsO₄³⁻) for (PO₄³⁻) in its cellular biomolecules, including DNA, proteins, and metabolites, while growing in phosphate-depleted media supplemented with arsenate. The study, published in Science and funded by NASA, suggested that GFAJ-1 incorporated arsenic into its nucleic acids via techniques like inductively coupled plasma mass spectrometry (ICP-MS) and X-ray spectroscopy, implying potential alternative biochemistries beyond phosphorus-based life. These findings faced swift scrutiny from biochemists, who noted arsenate's chemical instability: arsenate esters hydrolyze spontaneously at rates up to 10⁷ times faster than phosphate analogs under physiological conditions (pH 7, 37°C), rendering stable incorporation into polymers like untenable without unprecedented evolutionary adaptations. Subsequent independent analyses in 2012, including gel electrophoresis, ICP-MS, and radiolabeling experiments on GFAJ-1 cultures, detected no arsenate in purified or major metabolites even under arsenate exposure; growth ceased without trace phosphate contamination, confirming reliance on phosphorus. Two key refutation studies, one by Reaves et al. and another by Erb et al., demonstrated that arsenate induced phosphate scavenging via ribosome degradation rather than substitution, with arsenic localized extracellularly or in unstable intermediates. The original paper's retraction by Science in July 2025 followed persistent calls from critics, including Rosemary Redfield, who highlighted methodological flaws like inadequate phosphate removal and unsubstantiated spectral interpretations; Wolfe-Simon maintained the data validity but conceded reanalysis needs. Broader biochemical literature affirms arsenate's role as a phosphate competitive inhibitor—disrupting glycolysis via arsenolysis of glyceraldehyde-3-phosphate and uncoupling oxidative phosphorylation—but no verified stable incorporation into DNA, RNA, or ATP equivalents in any organism, as arsenate-DNA duplexes exhibit weakened hydrogen bonding and rapid hydrolysis. Claims of arsenate-tolerant microbes, such as in contaminated soils, reflect efflux pumps and methylation detoxification rather than incorporation. Empirical consensus holds that phosphorus indispensability stems from its bond stability, with arsenate's mimicry limited to transient, toxic interactions.

Debates on Risk Assessment

Risk assessment for arsenate, the pentavalent form of inorganic arsenic predominant in oxygenated environments, often aggregates it with trivalent arsenite under total inorganic arsenic metrics due to metabolic interconversion in vivo, though speciation influences bioavailability and acute toxicity, with arsenate exhibiting lower potency. Debates center on the dose-response relationship for chronic effects, particularly bladder and lung cancers, where regulatory agencies like the U.S. EPA apply a linear no-threshold (LNT) model assuming proportional risk extrapolation from high-dose epidemiological data (e.g., Taiwanese studies with exposures >100 μg/L) to low environmental levels, justifying standards such as the 10 μg/L maximum contaminant level for . This approach, conservative for protection, faces criticism for lacking empirical support at sub-50 μg/L doses, where cohort studies (e.g., in and ) show no significant cancer elevation, suggesting a threshold mechanism driven by epigenetic alterations rather than direct DNA damage. Proponents of a cite arsenic's —requiring metabolic activation via and , which saturates at low doses—along with evidence of essentiality for growth in mammals at <1 μg/kg/day, implying no harm below physiological needs and potential hormetic benefits. Estimated thresholds for cancer risk range from 50–200 μg/L in , with meta-analyses indicating supralinear responses (steeper at high doses) incompatible with LNT assumptions. Critics of LNT, including toxicologists reviewing low-dose data, argue it overstates risks, inflating remediation costs (e.g., billions for U.S. post-2001 EPA rule) without commensurate health gains, while underemphasizing confounders like or in arsenic-exposed populations. Conversely, LNT advocates highlight precautionary principles, pointing to animal bioassays and as justification, despite inconsistencies with human where risks plateau or diminish below 100 μg/L. Speciation-specific debates underscore arsenate's reduced uptake via transporters compared to arsenite's aquaglyceroporin pathway, potentially warranting differentiated assessments; however, regulatory frameworks rarely disaggregate forms, leading to conservative arsenate inclusion despite its lower reactivity and immobilization. Ongoing research, including probabilistic modeling, supports threshold incorporation for non-cancer endpoints like dermatosis (threshold ~20 μg/L), but classification under LNT persists due to institutional inertia rather than mechanistic consensus. These disputes highlight tensions between empirical dose-response data favoring thresholds and policy-driven conservatism, with calls for Bayesian updates integrating recent cohorts over high-dose relics.

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