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Mercury(II) sulfate
Mercury(II) sulfate
Mercury(II) sulfate
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Mercury(II) sulfate
Mercury(II) sulfate
Mercury(II) sulfate
Names
Other names
Mercuric sulfate, Mercurypersulfate, Mercury Bisulfate[1]
Identifiers
3D model (JSmol)
ChemSpider
ECHA InfoCard 100.029.083 Edit this at Wikidata
EC Number
  • 231-992-5
RTECS number
  • OX0500000
UNII
UN number 1645
  • InChI=1S/Hg.H2O4S/c;1-5(2,3)4/h;(H2,1,2,3,4)/q+2;/p-2
    Key: DOBUSJIVSSJEDA-UHFFFAOYSA-L
  • [O-]S(=O)(=O)[O-].[Hg+2]
Properties
HgSO4
Molar mass 296.653 g/mol
Appearance white monoclinic crystals
Odor odorless
Density 6.47 g/cm3, solid
450 °C (dec.)[2]
Decomposes in water to yellow mercuric subsulfate and sulfuric acid
Solubility soluble in hot H2SO4, NaCl solution
insoluble in ethanol, acetone, ammonia
−78.1·10−6 cm3/mol
Structure
rhombic
Thermochemistry
−707.5 kJ mol−1[3]
Hazards
GHS labelling:[4]
GHS06: ToxicGHS08: Health hazardGHS09: Environmental hazard
Danger
H300, H310, H330, H373, H410
P260, P262, P264, P270, P271, P273, P280, P284, P301+P316, P302+P352, P304+P340, P316, P319, P320, P321, P330, P361+P364, P391, P403+P233, P405, P501
NFPA 704 (fire diamond)
NFPA 704 four-colored diamondHealth 3: Short exposure could cause serious temporary or residual injury. E.g. chlorine gasFlammability 0: Will not burn. E.g. waterInstability 1: Normally stable, but can become unstable at elevated temperatures and pressures. E.g. calciumSpecial hazards (white): no code
3
0
1
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
☒N verify (what is checkY☒N ?)

Mercury(II) sulfate, commonly called mercuric sulfate, is the chemical compound HgSO4. It is an odorless salt that forms white granules or crystalline powder. In water, it separates into an insoluble basic sulfate with a yellow color and sulfuric acid.[3]

Structure

[edit]
Portion of structure of HgSO4 illustrating the distorted tetrahedral geometry at Hg (dark blue spheres).

The anhydrous compound features Hg2+ in a highly distorted tetrahedral HgO4 environment. Two Hg-O distances are 2.22 Å and the others are 2.28 and 2.42 Å.[5] In the monohydrate, Hg2+ adopts a linear coordination geometry with Hg-O (sulfate) and Hg-O (water) bond lengths of 2.179 and 2.228 Å, respectively. Four weaker bonds are also observed with Hg---O distances >2.5 Å.[6]

History

[edit]

In 1932, the Japanese chemical company Chisso Corporation began using mercury sulfate as the catalyst for the production of acetaldehyde from acetylene and water. Though it was unknown at the time, methylmercury is formed as a side product of this reaction. Exposure and consumption of the mercury waste products, including methylmercury, that were dumped into Minamata Bay by Chisso are believed to be the cause of Minamata disease in Minamata, Japan.[7]

Production

[edit]

Mercury sulfate can be produced by treating mercury with hot concentrated sulfuric acid:[8]

Hg + 2 H2SO4 → HgSO4 + SO2 + 2 H2O

Alternatively yellow mercuric oxide reacts also with concentrated sulfuric acid.[9]

Uses

[edit]

Denigés' reagent

[edit]

An acidic solution of mercury sulfate is known as Denigés' reagent. It was commonly used throughout the 20th century as a qualitative analysis reagent. If Denigés' reagent is added to a solution containing compounds that have tertiary alcohols, a yellow or red precipitate will form.[10]

Hydration reactions

[edit]
Conversion of 2,5-dimethylhexyne-2,5-diol to 2,2,5,5-tetramethylte-trahydrofuran-3-one
Conversion of 2,5-dimethylhexyne-2,5-diol to 2,2,5,5-tetramethylte-trahydrofuran-3-one

Mercury sulfate, as well as other mercury(II) compounds, are commonly used as catalysts in oxymercuration-demercuration, a type of electrophilic addition reaction that results in hydration of an unsaturated compound. The hydration of an alkene gives an alcohol. The regioselectivity is that predicted by Markovnikov's rule. For an alkyne, the result is an enol, which tautomerizes to give the carbonyl.[11] At one time, this chemistry was employed commercially for the preparation of acetaldehyde from acetylene:[12]

C2H2 + H2O → CH3CHO

A related and specialized example is the conversion of 2,5-dimethylhexyne-2,5-diol to 2,2,5,5-tetramethyltetrahydrofuran using aqueous mercury sulfate without the addition of acid.[13]

Health issues

[edit]

Inhalation of HgSO4 can result in acute poisoning: causing tightness in the chest, difficulties breathing, coughing and pain. Exposure of HgSO4 to the eyes can cause ulceration of conjunctiva and cornea. If mercury sulfate is exposed to the skin it may cause sensitization dermatitis. Lastly, ingestion of mercury sulfate will cause necrosis, pain, vomiting, and severe purging. Ingestion can result in death within a few hours due to peripheral vascular collapse.[1]

It was used in the late 19th century to induce vomiting for medical reasons.[14]

Further reading

[edit]

References

[edit]
[edit]
Revisions and contributorsEdit on WikipediaRead on Wikipedia

Mercury(II) sulfate

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from Grokipedia

Mercury(II) sulfate is an inorganic compound with the chemical formula HgSO4, appearing as an odorless white crystalline powder or granules denser than water. Its molecular weight is 296.66 g/mol, and it decomposes upon heating rather than melting. The compound exhibits low solubility in water but dissolves in concentrated sulfuric acid and dilute nitric acid, forming acidic solutions useful in specific chemical analyses. Primarily employed as a catalyst in organic synthesis, such as the hydration of alkynes to produce ketones and the industrial production of acetaldehyde from acetylene, it also serves as an analytical reagent to precipitate chloride ions and in processes for extracting gold and silver. However, mercury(II) sulfate is highly toxic, causing fatal outcomes via ingestion, inhalation, or skin absorption, with additional risks of severe respiratory irritation, organ damage, and long-term environmental persistence that renders it hazardous to aquatic life.

Properties

Molecular structure and bonding

Mercury(II) sulfate, HgSO₄, is an ionic compound comprising Hg²⁺ cations and SO₄²⁻ anions. In the solid state, it forms a crystalline lattice where the mercury(II) ions exhibit coordination to oxygen atoms provided by the sulfate anions. The crystal structure is orthorhombic, belonging to the space group Pmn2₁, and is isomorphous with cadmium(II) sulfate. The Hg²⁺ cation adopts an irregular eight-coordinate geometry, bonded to eight O²⁻ atoms with Hg–O bond distances spanning 2.23–2.93 Å, reflecting the hemidirected bonding characteristic of d¹⁰ mercury(II) centers due to relativistic effects and lone pair repulsion. This coordination results in a three-dimensional network framework linking [HgO₈] polyhedra with [SO₄] tetrahedra via corner-sharing. The sulfate tetrahedra maintain nearly regular geometry, with S–O bond lengths typical for SO₄²⁻ (approximately 1.47–1.49 Å). The bonding nature is predominantly ionic, arising from the electrostatic attraction between the highly charged, polarizable Hg²⁺ (a soft Lewis acid) and the SO₄²⁻ anions, though significant covalent character is present in the Hg–O interactions owing to orbital overlap and charge transfer from oxygen lone pairs. This hybrid bonding aligns with Pearson's , where soft-soft interactions stabilize the structure despite the formal ionic formulation. Experimental refinements confirm the lattice parameters and atomic positions, with reliability factors indicating high structural accuracy (R ≈ 0.033 for HgSO₄).

Physical and thermodynamic properties

Mercury(II) sulfate appears as an odorless white crystalline powder or granules. Its molecular formula is HgSO₄, with a molar mass of 296.65 g/mol. The compound has a density of 6.47 g/cm³ at standard conditions. The solid decomposes upon heating at approximately 450 °C, releasing mercury vapor, sulfur dioxide, and oxygen without exhibiting a distinct melting or boiling point. This thermal instability limits practical measurements of higher-temperature properties. Limited thermodynamic data are available; the standard enthalpy of formation is reported as -707.5 kJ/mol, though direct experimental verification in peer-reviewed thermochemical compilations is sparse due to the compound's reactivity and toxicity. No standard values for heat capacity or entropy at 298 K are widely documented in accessible databases, reflecting challenges in handling mercury compounds for calorimetry.
PropertyValueConditions
Density6.47 g/cm³Solid, room temp.
Decomposition temp.~450 °CHeating
Molar mass296.65 g/mol-
AppearanceWhite crystalline powder-
The crystal lattice is orthorhombic (rhombic), consistent with its granular form, though detailed lattice parameters require specialized studies not commonly tabulated in supplier data. Physical handling reveals it as denser than and non-volatile at ambient temperatures, with negligible (0 Pa at 25 °C).

Solubility and stability

Mercury(II) sulfate exhibits limited solubility in , undergoing rather than true dissolution to yield and a basic mercuric sulfate precipitate, such as HgSO₄·2HgO. It dissolves in concentrated , , hot dilute , and concentrated solutions, forming soluble complexes. The compound is insoluble in organic solvents including alcohol, acetone, and . Under standard ambient conditions, mercury(II) sulfate maintains chemical stability, showing no significant reactivity with air or moisture in the absence of hydrolysis triggers. Thermally, it decomposes in a single step between 450°C and 750°C, releasing toxic mercury vapors and sulfur oxides without forming stable intermediates. Hydrolytic instability in aqueous media limits its use in neutral or basic solutions, favoring acidic environments for handling.

Historical development

Early discovery and characterization

Mercury(II) sulfate was prepared through the oxidation of elemental mercury by hot concentrated , yielding the white crystalline solid, gas, and water. This straightforward synthesis, feasible after the production of concentrated in the , positioned the compound as a known in early modern chemical practice by the late . German chemist Johann Kunckel von Löwenstern documented its utility in producing mercuric chloride via sublimation of a with , underscoring its role as an established intermediate for mercury salt synthesis. Initial characterizations described mercury(II) sulfate as an odorless white powder or granules denser than , with noted via or , reflecting empirical observations from handling rather than quantitative analysis. Its insolubility in but reactivity in forming basic sulfates upon heating or dilution was observed, though systematic compositional verification awaited 19th-century gravimetric methods and .

Key industrial milestones

In 1932, Chisso Corporation in initiated industrial-scale production of from using mercury(II) sulfate as a catalyst in solution, implementing the Kucherov hydration process on a commercial basis for the first time.67944-0/fulltext) This marked a pivotal advancement, with initial output of 210 metric tons of that year, enabling efficient Markovnikov addition of water to under mild conditions (typically 60–80°C and 1–2 atm). The process converted (C₂H₂) to (CH₃CHO) via vinyl mercurinium intermediates, yielding up to 90% selectivity when optimized. By 1951, Chisso's capacity had expanded to 6,000 metric tons annually, reflecting broader adoption of the technology amid post-World War II demand for acetaldehyde-derived products such as , , and early plastics. Mercury(II) sulfate concentrations of 0.1–0.5% in 10–20% proved optimal for catalyst longevity and reaction rates, though mercury losses necessitated continuous replenishment. This route dominated acetaldehyde synthesis globally until the 1960s, when cheaper ethylene oxidation via the supplanted it due to petroleum feedstock availability and reduced concerns. Other industrial applications emerged concurrently, including mercury(II) sulfate's use as an in primary mercury batteries and in gold/silver extraction from pyritic ores, but these lacked the scale of acetaldehyde production until environmental regulations curtailed mercury use post-1970. The implementation highlighted both the compound's catalytic efficacy—stemming from mercury's soft Lewis acid coordination to alkyne π-bonds—and eventual drawbacks from effluent mercury accumulation.

Synthesis and production

Laboratory synthesis

Mercury(II) sulfate is commonly prepared in the laboratory by heating elemental mercury with an excess of hot concentrated sulfuric acid, which oxidizes the mercury and yields the product along with sulfur dioxide gas and water:
Hg + 2 H₂SO₄ → HgSO₄ + SO₂ + 2 H₂O. This method requires careful control of temperature to ensure complete reaction, typically around 200–250 °C, and the product forms as white crystals that can be isolated by filtration after cooling and dilution. An alternative procedure involves dissolving yellow mercury(II) oxide in concentrated sulfuric acid, followed by evaporation to concentrate the solution, cooling to induce crystallization, and filtration to recover the solid. This approach avoids the production of sulfur dioxide gas and is suitable for smaller-scale preparations, though it presupposes access to mercury(II) oxide. Both methods produce anhydrous HgSO₄, which is sparingly soluble in water and must be handled under fume hood conditions due to the toxicity of mercury compounds and reaction byproducts.

Industrial-scale production

Mercury(II) sulfate is produced on an industrial scale primarily by the direct oxidation of elemental mercury with excess hot concentrated sulfuric acid, yielding the product alongside sulfur dioxide and water:
Hg + 2 H₂SO₄ → HgSO₄ + SO₂ + 2 H₂O. This exothermic reaction requires careful control of temperature and acid concentration to ensure complete conversion and minimize side products, with the mercury sulfate precipitating as white crystals that are filtered, washed, and dried. Historically, this method supported the preparation of mercury(II) sulfate as a catalyst for the large-scale hydration of to , a key process in early 20th-century chemical manufacturing, such as operations by the Corporation starting in 1932. The catalyst was typically generated or batch-processed in acidic media for direct use in reactors, enabling yields essential for downstream production of acetic acid and other organics before ethylene-based alternatives dominated post-World War II. Contemporary industrial production is minimal due to stringent regulations under the , which mandated phase-out of mercury catalysts in processes by 2018, driven by risks and environmental persistence. Residual output occurs via from spent catalysts, involving of mercury from waste followed by re-oxidation with or acid treatments to regenerate HgSO₄, as detailed in recovery s. Such approaches prioritize waste minimization but do not support broad commercial volumes, with global mercury supply now constrained to essential uses.

Chemical reactivity and applications

Reactions in analytical chemistry

An acidic solution of mercury(II) sulfate, designated as Denigés' reagent, facilitates qualitative identification of tertiary alcohols through dehydration to isoolefins followed by mercuration, yielding a characteristic white or colored precipitate of an organomercury compound; primary and secondary alcohols react more slowly or not at all under these conditions. This reagent, prepared by dissolving mercuric oxide in dilute sulfuric acid to generate HgSO4, was widely applied in 20th-century organic qualitative analysis to distinguish alcohol classes based on reaction kinetics and precipitate formation. In the for total determination, mercury(II) sulfate serves as a catalyst during digestion of organic samples, accelerating the conversion of organic to by enhancing oxidation rates without being stoichiometrically consumed. Typical procedures incorporate 0.5–1 g of HgSO4 per digestion flask alongside to elevate , enabling complete decomposition at 360–410 °C; post-digestion, mercury is sequestered with to prevent interference in and steps. For chemical oxygen demand (COD) assessment in wastewater via dichromate oxidation, mercury(II) sulfate complexes chloride ions as HgCl2, mitigating their interference by averting artificial oxygen demand from chloride oxidation. Standard protocols add HgSO4 at approximately 10 mg per estimated mg of chloride in the aliquot prior to acidification and reflux, ensuring selective quantification of organic oxidizable matter; this application persists in methods like EPA 410.3 despite toxicity concerns prompting mercury-free alternatives.

Applications in organic synthesis

Mercury(II) sulfate acts as a catalyst in the Kucherov reaction, facilitating the Markovnikov hydration of terminal alkynes to yield methyl ketones under acidic aqueous conditions. In this process, a terminal alkyne (RC≡CH) reacts with in the presence of HgSO₄ and dilute (typically 5-20% H₂SO₄) at elevated temperatures (around 60-100°C), producing the corresponding ketone (RC(O)CH₃) via enol tautomerization. The mercury(II) ion coordinates to the , promoting of and preventing rearrangements that could occur in uncatalyzed acid hydration. This arises from the formation of a vinyl mercurinium intermediate, followed by nucleophilic attack by at the more substituted carbon. The reaction's mechanism begins with π-complexation of Hg²⁺ to the alkyne, generating a mercurinium ion that directs addition to the internal carbon, yielding a vinyl mercury species. of this intermediate produces an , which rapidly tautomerizes to the . Yields are typically high (70-95%) for simple terminal alkynes, though steric hindrance or electron-withdrawing groups on R can reduce efficiency. For example, 1-hexyne hydrates to in quantitative yields under standard conditions. Historically, this application enabled industrial-scale production of from (CH≡CH + H₂O → CH₃CHO), a process commercialized in the early before ethylene-based routes dominated. In modern laboratory , HgSO₄ remains valuable for synthesizing ketones from alkynes where alternatives like hydroboration-oxidation (which favors aldehydes) are unsuitable. Immobilized variants, such as silica-supported HgSO₄, allow milder conditions (room temperature, shorter times) while minimizing mercury contamination, achieving comparable yields for diverse substrates. Internal alkynes react more slowly and often require higher catalyst loadings, yielding mixtures unless symmetrical. Due to mercury's , the reaction is increasingly supplemented by metal-free or less hazardous catalysts, though HgSO₄'s efficacy in regioselective hydration persists in targeted syntheses.

Electrochemical and other uses

Mercury(II) sulfate functions as an in select primary batteries, including reserve types and those designed for low-temperature operation, where its chemical stability supports performance in non-aqueous or organic systems. For instance, in experimental batteries for space probe applications, it has been incorporated as a component in slurries with organic liquids, enabling flat discs due to its compatibility and lack of reactivity with many solvents. Such uses leverage its electrochemical properties, though broader adoption is constrained by mercury's and environmental regulations. Beyond , mercury(II) sulfate finds application in metallurgical processes for extracting and silver from roasted pyrites, typically combined with to facilitate dissolution and recovery of the metals. It also serves as an intermediate in producing other mercury compounds, capitalizing on its in and reactivity to form derivatives like mercury fulminate or amalgams. Historically, limited medicinal uses have been noted, though these are obsolete due to health risks.

Toxicology

Mechanisms of toxicity

Mercury(II) sulfate, upon dissolution in biological fluids, releases Hg²⁺ ions, which serve as the principal toxic moiety due to their reactivity, while the sulfate component may contribute to local corrosivity through acidification. The core mechanism of Hg²⁺ stems from its exceptionally high affinity for sulfhydryl (-SH) groups in residues of proteins and low-molecular-weight thiols such as (GSH), forming stable mercuric-thiolate complexes with association constants on the order of 10¹⁵–10²⁰. This binding denatures proteins, inhibits critical enzymes (e.g., those in defense and pathways), and depletes intracellular GSH, thereby compromising cellular . Hg²⁺ uptake into cells occurs primarily as conjugates with thiols like cysteine or GSH, facilitated by amino acid transporters (e.g., system b⁰,⁺ or L-type) and organic anion transporters (e.g., OAT1/OAT3), enabling accumulation in target tissues such as the kidney proximal tubules, where up to 40% of an administered dose can localize within hours. Intracellularly, these interactions trigger oxidative stress by generating reactive oxygen species (ROS), promoting lipid peroxidation (e.g., elevated malondialdehyde levels), and impairing mitochondrial function, which collectively lead to membrane disruption, lysosomal instability, and activation of apoptotic pathways via caspase-3/7. Additionally, Hg²⁺ precipitates proteins on direct contact, inducing necrosis, and disrupts microtubule assembly while altering calcium homeostasis, further amplifying cytotoxicity. In the , limited penetration of inorganic Hg²⁺ is overcome in chronic exposures, where it binds synaptic proteins, inhibits neuromuscular transmission, and induces , contributing to functional deficits like reduced neuronal density. Renal arises from tubular reabsorption overload, exacerbating inhibition and ROS-mediated to yield cortical . These mechanisms underscore Hg²⁺'s non-specific interference with thiol-dependent processes, with absorption rates of 7–15% for oral inorganic mercury underscoring its despite poor .

Acute and chronic human effects

Acute exposure to mercury(II) sulfate primarily occurs via , of dust, or contact, resulting in severe corrosive effects on mucous membranes and rapid systemic absorption of mercuric ions. of as little as 1 gram has caused , , , and shock in a documented case of a 40-year-old male, necessitating emergency , with 2,3-dimercaptopropane-1-sulfonate, and for survival. Symptoms include intense , , , and gastrointestinal ulceration due to the compound's acidity and mercury's affinity for sulfhydryl groups in proteins, disrupting cellular function. may lead to irritation and pulmonary damage, while dermal exposure causes local burns and potential systemic through absorption. Fatal outcomes are possible without prompt intervention, as mercuric ions induce and enzyme inhibition. Chronic exposure, often occupational or from repeated low-level contact, manifests predominantly as , with mercuric ions accumulating in renal tubules and proximal cells, leading to , glomerular dysfunction, and over time. Neurological effects include tremors, , memory impairment, sensory disturbances, and , attributed to mercury's disruption of neuronal and systems. Additional symptoms encompass psychiatric disturbances such as and mood changes, increased salivation, , and potential fertility impairment in males due to gonadal toxicity. Dermatological issues like rashes and gum may arise, alongside cardiovascular strain from prolonged exposure. Long-term accumulation exacerbates multi-organ damage, including to the heart, lungs, and , underscoring the compound's classification as a cumulative .

Environmental impact

Persistence and transformation in ecosystems

Mercury(II) sulfate demonstrates low persistence as a distinct compound in ecosystems due to its reactivity with , hydrolyzing to produce insoluble basic mercuric sulfate (such as yellow mercuric oxysulfate) and , which limits its free dissolution and mobility in aquatic systems. This hydrolysis occurs rapidly upon exposure, releasing Hg²⁺ ions that bind avidly to particulate matter, sediments, and , thereby immobilizing mercury in soils and aquatic sediments where it can persist for decades to centuries without chemical degradation. In anaerobic environments, such as sediments and profundal lake bottoms, Hg²⁺ from hydrolyzed mercury(II) undergoes biotic transformation primarily through microbial by -reducing (SRB), which utilize the co-released as an , enhancing (MeHg) production rates by up to several-fold compared to sulfate-limited conditions. This process is - and organic carbon-dependent, with higher DOC levels promoting Hg for , while elevated inputs from the compound's breakdown amplify SRB activity and MeHg yields, as observed in freshwater systems where concentrations above 10 mg/L correlate with increased MeHg flux. Abiotic transformations include photoreduction to volatile mercury (Hg⁰) in oxic surface waters under , with quantum yields for Hg²⁺ reduction ranging from 0.01 to 0.1 depending on complexation, and precipitation as sparingly soluble sulfides (e.g., metacinnabar, HgS) in sulfide-rich anoxic zones, which further reduces but can remobilize under changing conditions. Overall, while the parent compound transforms quickly, the resultant inorganic mercury reservoir remains environmentally persistent, with transformation pathways favoring bioaccumulative MeHg in -influenced ecosystems, as evidenced by field studies linking anthropogenic deposition to elevated MeHg in biota.

Bioaccumulation and regulatory responses

Mercury(II) sulfate, upon environmental release, dissociates into Hg(II) ions and , with the inorganic mercury susceptible to microbial in anaerobic sediments by -reducing , forming highly bioavailable (MeHg). This process is facilitated by availability, as -reducing metabolize while incorporating Hg(II), enhancing MeHg production in ecosystems with elevated levels from sources like industrial discharges. MeHg exhibits strong , concentrating in aquatic organisms such as and before biomagnifying up the , with concentrations increasing by factors of 10^5 to 10^6 from water to top predators like . Organisms can bioconcentrate mercury from water up to 10,000-fold, leading to chronic exposure risks for higher trophic levels. Regulatory responses to mercury compounds, including mercury(II) sulfate, emphasize emission controls and phase-outs due to its persistence and toxicity potential. The , ratified by over 140 parties since 2013, mandates reductions in mercury releases to land and water, including from and , with specific thresholds for mercury waste management to prevent environmental entry. In the United States, mercury(II) sulfate is classified as a hazardous substance under the Clean Water Act (Section 311(b)(2)(A)), subjecting it to spill reporting and cleanup requirements, and its export has been prohibited since January 1, 2020, under the Mercury Export Ban Act to curb global circulation. The (ECHA) lists mercury(II) sulfate under REACH restrictions, prohibiting its use in mixtures where it poses risks to human health or the environment, with authorization required for specific industrial applications. These measures reflect empirical evidence of mercury's long-range transport and hazards, prioritizing protection over unrestricted industrial use.

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