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Cobalt(II) sulfate
Cobalt(II) sulfate
Cobalt(II) sulfate
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Cobalt(II) sulfate
Cobalt(II) sulfate Xray
Cobalt(II) sulfate Xray
Names
IUPAC name
Cobalt(II) sulfate
Other names
Cobaltous sulfate
Identifiers
3D model (JSmol)
ChEBI
ChemSpider
ECHA InfoCard 100.030.291 Edit this at Wikidata
EC Number
  • 233-334-2
KEGG
RTECS number
  • GG3100000 (anhydrous)
    GG3200000 (heptahydrate)
UNII
  • InChI=1S/Co.H2O4S/c;1-5(2,3)4/h;(H2,1,2,3,4)/q+2;/p-2 checkY
    Key: KTVIXTQDYHMGHF-UHFFFAOYSA-L checkY
  • InChI=1/Co.H2O4S/c;1-5(2,3)4/h;(H2,1,2,3,4)/q+2;/p-2
    Key: KTVIXTQDYHMGHF-NUQVWONBAJ
  • anhydrous: [Co+2].[O-]S([O-])(=O)=O
  • hexahydrate: [OH2+][Co-4]([OH2+])([OH2+])([OH2+])([OH2+])[OH2+].[O-]S(=O)(=O)[O-]
  • heptahydrate: [OH2+][Co-4]([OH2+])([OH2+])([OH2+])([OH2+])[OH2+].[O-]S(=O)(=O)[O-].O
Properties
CoSO4·(H2O)n (n=0,1,6,7)
Molar mass 154.996 g/mol (anhydrous)
173.011 g/mol (monohydrate)
263.08 g/mol (hexahydrate)
281.103 g/mol (heptahydrate)
Appearance reddish crystalline (anhydrous, monohydrate)
pink salt (hexahydrate)
Odor odorless (heptahydrate)
Density 3.71 g/cm3 (anhydrous)
3.075 g/cm3 (monohydrate)
2.019 g/cm3 (hexahydrate)
1.948 g/cm3 (heptahydrate)
Melting point 735 °C (1,355 °F; 1,008 K)
anhydrous:
36.2 g/100 mL (20 °C)
38.3 g/100 mL (25 °C)
84 g/100 mL (100 °C)
heptahydrate:
60.4 g/100 mL (3 °C)
67 g/100 mL (70 °C)
Solubility anhydrous:
1.04 g/100 mL (methanol, 18 °C)
insoluble in ammonia
heptahydrate:
54.5 g/100 mL (methanol, 18 °C)
+10,000·10−6 cm3/mol
1.639 (monohydrate)
1.540 (hexahydrate)
1.483 (heptahydrate)
Structure
orthorhombic (anhydrous)
monoclinic (monohydrate, heptahydrate)
Hazards
GHS labelling:
GHS07: Exclamation markGHS08: Health hazardGHS09: Environmental hazard
Danger
H302, H317, H334, H341, H350, H360, H410
P201, P202, P261, P264, P270, P272, P273, P280, P281, P285, P301+P312, P302+P352, P304+P341, P308+P313, P321, P330, P333+P313, P342+P311, P363, P391, P405, P501
NFPA 704 (fire diamond)
NFPA 704 four-colored diamondHealth 2: Intense or continued but not chronic exposure could cause temporary incapacitation or possible residual injury. E.g. chloroformFlammability 0: Will not burn. E.g. waterInstability 0: Normally stable, even under fire exposure conditions, and is not reactive with water. E.g. liquid nitrogenSpecial hazards (white): no code
2
0
0
Flash point Non-flammable
Lethal dose or concentration (LD, LC):
424 mg/kg (oral, rat)
Safety data sheet (SDS) ICSC 1396 (heptahydrate)
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
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Cobalt(II) sulfate
Cobalt(II) sulfate heptahydrate

Cobalt(II) sulfate is any of the inorganic compounds with the formula CoSO4(H2O)x. Usually cobalt sulfate refers to the hexa- or heptahydrates CoSO4.6H2O or CoSO4.7H2O, respectively.[1] The heptahydrate is a red solid that is soluble in water and methanol. Since cobalt(II) has an odd number of electrons, its salts are paramagnetic.

Preparation, and structure

[edit]

It forms by the reaction of metallic cobalt, its oxide, hydroxide, or carbonate with aqueous sulfuric acid:[1]

Co + H2SO4 + 7 H2O → CoSO4(H2O)7 + H2
CoO + H2SO4 + 6 H2O → CoSO4(H2O)7

The heptahydrate is only stable at humidity >70% at room temperature, otherwise it converts to the hexahydrate.[2] The hexahydrate converts to the monohydrate and the anhydrous forms at 100 and 250 °C, respectively.[1]

CoSO4(H2O)7 → CoSO4(H2O)6 + H2O
CoSO4(H2O)6 → CoSO4(H2O) + 5 H2O
CoSO4(H2O) → CoSO4 + H2O

The hexahydrate is a metal aquo complex consisting of octahedral [Co(H2O)6]2+ ions associated with sulfate anions (see image in table).[3] The monoclinic heptahydrate has also been characterized by X-ray crystallography. It also features [Co(H2O)6]2+ octahedra as well as one water of crystallization.[2]

Uses and reactions

[edit]

Cobalt sulfates are important intermediates in the extraction of cobalt from its ores. Thus, crushed, partially refined ores are treated with sulfuric acid to give red-colored solutions containing cobalt sulfate.[1]

Hydrated cobalt(II) sulfate is used in the preparation of pigments, as well as in the manufacture of other cobalt salts. Cobalt pigment is used in porcelains and glass. Cobalt(II) sulfate is used in storage batteries and electroplating baths, sympathetic inks, and as an additive to soils and animal feeds. For these purposes, the cobalt sulfate is produced by treating cobalt oxide with sulfuric acid.[1]

Being commonly available commercially, the heptahydrate is a routine source of cobalt in coordination chemistry.[4]

Natural occurrence

[edit]

Rarely, cobalt(II) sulfate is found in form of few crystallohydrate minerals, occurring among oxidation zones containing primary Co minerals (like skutterudite or cobaltite). These minerals are: biebierite (heptahydrate), moorhouseite (Co,Ni,Mn)SO4.6H2O,[5][6] aplowite (Co,Mn,Ni)SO4.4H2O and cobaltkieserite (monohydrate).[7][8][6]

Health issues

[edit]

Cobalt is an essential mineral for mammals, but more than a few micrograms per day is harmful. Although poisonings have rarely resulted from cobalt compounds, their chronic ingestion has caused serious health problems at doses far less than the lethal dose. In 1965, the addition of a cobalt compound to stabilize beer foam in Canada led to a peculiar form of toxin-induced cardiomyopathy, which came to be known as beer drinker's cardiomyopathy.[9][10][11]

Furthermore, cobalt(II) sulfate is suspected of causing cancer (i.e., possibly carcinogenic, IARC Group 2B) as per the International Agency for Research on Cancer (IARC) Monographs.[12]

[edit]

References

[edit]
Revisions and contributorsEdit on WikipediaRead on Wikipedia

Cobalt(II) sulfate

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Cobalt(II) sulfate is an with the chemical formula CoSO₄, typically occurring as the vibrant red heptahydrate form CoSO₄·7H₂O that readily dissolves in to form pink solutions. This salt, characterized by its rose- crystalline appearance and odorless nature, has a molecular weight of approximately 155 g/mol for the form and 281 g/mol for the heptahydrate, with in reaching about 38 g/100 mL at 25°C. It decomposes upon heating above °C, releasing toxic oxides, and is produced industrially by reacting (II) oxide or with dilute . Cobalt(II) sulfate finds extensive applications across various industries due to its and content. It serves as a key component in storage batteries, particularly lithium-ion types, and in baths to deposit coatings on metals. Additionally, it acts as a drier in paints, inks, and varnishes, accelerating oxidation and , and as a imparting blue colors in ceramics, , and glazes. Other uses include its role as a catalyst in chemical reactions, a soil additive to correct deficiencies in , and a supplement in animal feeds to support health. Despite its utility, cobalt(II) sulfate poses significant health and environmental risks, classified as a possible carcinogen with evidence from animal studies showing tumors in rats and mice. It is or inhaled, potentially causing respiratory , skin allergies, and genetic damage, with an oral LD50 in rats of 424 mg/kg. Environmentally, it is highly toxic to aquatic life, necessitating careful handling and disposal to prevent long-term ecological harm.

Properties

Physical Properties

Cobalt(II) sulfate occurs in several hydrated and forms, with the heptahydrate (CoSO₄·7H₂O) and hexahydrate (CoSO₄·6H₂O) being the most common. The form (CoSO₄) is a crystalline solid, while the hexahydrate consists of monoclinic prisms, and the heptahydrate forms to rhombic prisms or granular solids. The compound is odorless in all forms. Key physical properties of the common forms are summarized below:
PropertyAnhydrous (CoSO₄)Heptahydrate (CoSO₄·7H₂O)
Density (g/cm³)3.711.948
Melting point (°C)Decomposes at 735Loses water at ~100
Solubility in water (g/100 mL at 20 °C)36.236.2 (as CoSO₄ equivalent)
Densities sourced from chemical reference data; melting behavior from international chemical safety cards. Solubility values represent the amount of dissolved CoSO₄; the heptahydrate is highly soluble overall, with additional solubility in and slight solubility in , but insolubility in . The heptahydrate is hygroscopic and prone to , remaining stable only above approximately 70% relative at , below which it dehydrates to the hexahydrate. Upon heating, the heptahydrate undergoes stepwise , losing to form lower hydrates such as the monohydrate between 100–200 °C before reaching the form at higher temperatures. Cobalt(II) sulfate exhibits paramagnetic behavior attributable to the unpaired electrons in the Co²⁺ ions.

Chemical Properties

Cobalt(II) sulfate has the CoSO₄ for the form, with a molecular weight of 154.996 g/mol, while the heptahydrate form, CoSO₄·7H₂O, has a molecular weight of 281.103 g/mol. As an ionic compound, (II) sulfate dissociates completely in to produce Co²⁺ and SO₄²⁻ s. Aqueous solutions of (II) are acidic, with a around 4–5 for a 0.1 M solution, resulting from partial of the Co²⁺ . In , the Co²⁺ primarily exists as the octahedral hexaaqua complex [Co(H₂O)₆]²⁺, which undergoes limited according to the equilibrium [Co(H₂O)₆]²⁺ ⇌ [Co(H₂O)₅OH]⁺ + H⁺. The hydrate forms of cobalt(II) sulfate are stable in air under normal conditions, whereas the form is hygroscopic and readily absorbs atmospheric moisture to form hydrates. At high temperatures above 700 °C, the form decomposes in air to yield cobalt(II) oxide and , following the reaction: \ceCoSO4−>CoO+SO3\ce{CoSO4 -> CoO + SO3} The Co²⁺ ion in cobalt(II) sulfate displays redox activity, capable of oxidation to Co³⁺ in the presence of strong oxidizing agents such as hydrogen peroxide and reduction to metallic cobalt with strong reducing agents such as magnesium. Cobalt(II) sulfate is incompatible with strong bases, which cause precipitation of cobalt(II) hydroxide, and with reducing agents, which can reduce the Co²⁺ ion.

Structure and Synthesis

Crystal Structure

Cobalt(II) sulfate exists in several solid forms, each exhibiting distinct crystal structures determined primarily by . The anhydrous form, α-CoSO₄, crystallizes in the orthorhombic system with Pnma and parameters a = 6.120 , b = 7.710 , c = 5.130 . In this structure, Co²⁺ cations occupy sites with distorted octahedral coordination to six oxygen atoms from SO₄²⁻ tetrahedra, resulting in corner-sharing CoO₆ octahedra that form a three-dimensional framework with the groups. The most common hydrated form, the heptahydrate CoSO₄·7H₂O (bieberite), adopts a with P2₁/c and parameters a = 14.048 , b = 6.494 , c = 10.925 , β = 105.23°. This structure consists of discrete [Co(H₂O)₆]²⁺ octahedra, where the cobalt ion is coordinated to six water oxygen atoms with Co–O bond lengths ranging from 2.08 to 2.11 , linked by hydrogen bonds to isolated SO₄²⁻ tetrahedra. One water molecule exhibits disorder, occupying a trigonally planar position that influences the overall hydrogen bonding network. The red color of this form arises from d–d electronic transitions in the octahedral Co²⁺ environment. The hexahydrate CoSO₄·6H₂O (moorhouseite) is also monoclinic, belonging to C2/c with parameters a = 16.262 Å, b = 7.032 Å, c = 9.549 Å, β = 125.95°. It features [Co(H₂O)₆]²⁺ octahedra similar to the heptahydrate but connected through a distinct hydrogen bonding arrangement with sulfate anions, leading to a layered motif. No major polymorphs have been reported for the hydrated forms of (II) sulfate, though induces phase transitions, such as the conversion from the heptahydrate to the hexahydrate or lower hydrates, often accompanied by structural rearrangements due to loss of molecules from the .

Preparation Methods

Cobalt(II) sulfate is commonly prepared in the laboratory by reacting metal with dilute at , following the equation: Co+H2SO4→CoSO4+12H2\text{Co} + \text{H}_2\text{SO}_4 \rightarrow \text{CoSO}_4 + \frac{1}{2}\text{H}_2 This reaction proceeds readily, yielding near-quantitative conversion to the soluble cobalt(II) , with gas evolving as a . Alternatively, cobalt(II) oxide can be used as the starting material, dissolving in dilute according to: CoO+H2SO4→CoSO4+H2O\text{CoO} + \text{H}_2\text{SO}_4 \rightarrow \text{CoSO}_4 + \text{H}_2\text{O} This method also achieves high yields under ambient conditions and is suitable for small-scale synthesis. The resulting is then concentrated by gentle and cooled to induce of the heptahydrate form, CoSO₄·7H₂O, which appears as red monoclinic crystals. Another route involves the reaction of cobalt(II) carbonate with dilute sulfuric acid: CoCO3+H2SO4→CoSO4+CO2+H2O\text{CoCO}_3 + \text{H}_2\text{SO}_4 \rightarrow \text{CoSO}_4 + \text{CO}_2 + \text{H}_2\text{O} This effervescent reaction occurs efficiently at room temperature, providing a convenient alternative when the carbonate is available, and similarly leads to the heptahydrate upon crystallization. Purification of the crude product is typically accomplished by recrystallization from hot , with careful control to avoid excessive heating that might cause partial or . To obtain the form, the heptahydrate is dehydrated under at approximately 250 °C, ensuring complete removal of without . Historical methods for isolating cobalt(II) sulfate in the early 19th century relied on processing natural cobalt ores, such as arsenides, through roasting to form the oxide followed by leaching with sulfuric acid to extract the sulfate.

Occurrence and Production

Natural Occurrence

Cobalt(II) sulfate occurs naturally as a rare secondary mineral in the oxidation zones of cobalt-bearing ore deposits, where it forms through the weathering of primary sulfides and arsenides in arid or evaporative environments. These minerals are scarce due to the overall rarity of primary cobalt sulfides in nature. The primary mineral forms include bieberite (CoSO₄·7H₂O), a pinkish-red efflorescent crust; moorhouseite ((Co,Ni,Mg)SO₄·6H₂O), a hexahydrate; aplowite (CoSO₄·4H₂O), a rarer tetrahydrate; and cobaltkieserite (CoSO₄·MgSO₄·6H₂O), a mixed sulfate. These minerals develop in geological settings characterized by hydrothermal veins hosting cobalt arsenides, such as , which undergo alteration in near-surface oxidation zones. The process involves the interaction of sulfate-rich waters with released from primary minerals during , often in association with other evaporative sulfates. Major known occurrences are documented in regions with significant mineralization, including the Bou Azzer district in , the copper-cobalt belts in , the Katanga region in the of Congo, and historical sites like , , where bieberite has been reported in oxidized arsenide zones. The extraction history of natural cobalt(II) sulfate traces back to the , when ores were processed for smalt production—a blue glass pigment—yielding sulfate intermediates during . However, pure mineral forms like bieberite were not formally identified until the , with bieberite first described in 1845 from German localities. Later discoveries, such as moorhouseite and aplowite in 1965 from , and cobaltkieserite in 1992 from , highlight their ongoing recognition as rare species.

Industrial Production

Cobalt(II) sulfate is primarily produced through hydrometallurgical processes that recover from ores such as copper-cobalt sulfides or deposits, as well as from sources like spent lithium-ion batteries. The dominant method involves leaching of cobalt-bearing materials, often under atmospheric or high-pressure conditions, to dissolve the metal into solution. For sulfide ores, oxidative leaching with converts (CoS) to cobalt(II) sulfate, as represented by the reaction: 2CoS+O2+2H2SO4→2CoSO4+2S+2H2O2 \text{CoS} + \text{O}_2 + 2 \text{H}_2\text{SO}_4 \rightarrow 2 \text{CoSO}_4 + 2 \text{S} + 2 \text{H}_2\text{O} This step is followed by purification techniques including solvent extraction—typically using organophosphorus extractants like Cyanex 272—to separate cobalt from impurities, and subsequent precipitation or crystallization to yield the sulfate product. In nickel-cobalt laterite processing, crude cobalt sulfate is obtained as an intermediate after separation from via solvent extraction or , with the stream then refined to high purity. Global production of refined for chemical applications, where is the predominant form, reached 179,000 metric tons of contained in 2023, with major output from (78.5% of refined ), the of Congo (76% of mined ), and . In 2024, refined production grew to 222,000 metric tons, driven by demand for batteries. High-purity grades exceeding 99.5% content are targeted for battery precursors, achieved through additional to remove contaminants like and . Post-2020 developments have driven production growth to support supply chains, with annual output increasing alongside rising demand for materials. from end-of-life lithium-ion batteries contributed about 8% of total supply as of 2024 (22,000 metric tons secondary supply), involving acid leaching of followed by solvent extraction to produce battery-grade heptahydrate (CoSO₄·7H₂O). This recycled fraction is projected to rise to 13% by 2033, emphasizing sustainable sourcing amid diversification efforts.

Uses and Reactions

Industrial Uses

Cobalt(II) sulfate has been employed as a colorant in the production of pigments and ceramics since the , particularly in the form of smalt, a finely ground used to impart hues to , porcelains, glazes, and enamels. This historical application leverages the compound's ability to provide stable, transparent tones, preventing discoloration in ceramic finishes and serving as an affordable alternative to more expensive pigments like . In modern contexts, it continues to function as a pigment in and glassmaking, as well as a agent in inks and paints. In , cobalt(II) sulfate serves as a supplement in fertilizers applied to cobalt-deficient soils and pastures, enhancing quality for . It is widely used in formulations, particularly for ruminants such as and sheep, where it prevents cobalt deficiency by supporting microbial synthesis of ; recommended concentrations in feed range from 0.07 to 0.10 ppm to meet nutritional needs. Veterinary applications include its addition to salt licks and supplements at levels around 0.1-0.2% to address deficiencies, promoting formation and overall growth in . The compound plays a key role in and , acting as an in cobalt plating baths to produce wear-resistant coatings on metals, and as an additive in plating solutions to enhance deposit smoothness, brightness, hardness, and ductility. In , cobalt(II) sulfate is a critical precursor for synthesizing lithium oxide (NMC) cathodes in lithium-ion batteries, with battery applications accounting for approximately 64% of global demand in 2020, rising to over 75% as of 2025 largely due to production; however, growth is tempered by the adoption of cobalt-reduced chemistries like (LFP). Additional uses include its role in sympathetic inks and as a stabilizer for storing concentrated .

Chemical Reactions

Cobalt(II) sulfate in aqueous solution primarily exists as the hexaaquacobalt(II) ion, [Co(H₂O)₆]²⁺, which readily undergoes ligand exchange reactions to form coordination complexes. A prominent example is the reaction with excess ammonia, where water ligands are sequentially replaced to yield the hexaamminecobalt(II) ion: [\ceCo(H2O)62+]+6\ceNH3⇌[\ceCo(NH3)62+]+6\ceH2O[\ce{Co(H2O)6^2+}] + 6\ce{NH3} \rightleftharpoons [\ce{Co(NH3)6^2+}] + 6\ce{H2O} This equilibrium-driven process, often conducted in concentrated ammonia, is central to the synthesis of Werner complexes, which provided key evidence for coordination theory by demonstrating the octahedral geometry and stability of such species. Similarly, ethylenediaminetetraacetic acid (EDTA) forms a stable chelate complex with cobalt(II): [\ceCo(H2O)62+]+\ceEDTA4−⇌[\ceCo(EDTA)2−]+6\ceH2O[\ce{Co(H2O)6^2+}] + \ce{EDTA^4-} \rightleftharpoons [\ce{Co(EDTA)^2-}] + 6\ce{H2O} The formation constant for [Co(EDTA)]²⁻ is approximately 10¹⁶, indicating high stability suitable for analytical and sequestration applications. Precipitation reactions of cobalt(II) sulfate are commonly employed to isolate cobalt(II) hydroxide. When an aqueous solution of CoSO₄ reacts with sodium hydroxide, a rose-pink precipitate forms: \ceCoSO4+2NaOH−>Co(OH)2v+Na2SO4\ce{CoSO4 + 2NaOH -> Co(OH)2 v + Na2SO4} This double-displacement reaction proceeds quantitatively under neutral or slightly basic conditions, with the hydroxide precipitate serving as an intermediate in further processing or purification steps. Redox transformations highlight the +2 oxidation state's reactivity. Oxidation to cobalt(III) can be achieved using hydrogen peroxide, particularly in the presence of base, yielding cobalt(III) hydroxide or related complexes: 2\ceCo2+(aq)+\ceH2O2(aq)+2\ceH2O(l)→2\ceCo3+(aq)+4\ceOH−(aq)2\ce{Co^2+ (aq)} + \ce{H2O2 (aq)} + 2\ce{H2O (l)} \rightarrow 2\ce{Co^3+ (aq)} + 4\ce{OH^- (aq)} This reaction, often visualized by the color change from pink to brown, underscores the instability of Co(III) in aqueous media without stabilizing ligands. Conversely, reduction to metallic cobalt occurs via displacement with zinc metal in acidic sulfate solutions: \ceCo2++Zn−>Co+Zn2+\ce{Co^2+ + Zn -> Co + Zn^2+} This cementation process is industrially relevant for cobalt recovery from leach solutions. In , solutions are identified through color-forming reactions. Addition of (NH₄SCN) produces a deep blue tetrahedral complex, [Co(SCN)₄]²⁻, which is intensified by adding acetone to reduce coordination: \ceCo2++4SCN−−>[Co(SCN)42−]\ce{Co^2+ + 4SCN^- -> [Co(SCN)4^2-]} This test provides a sensitive qualitative around 1 ppm for ions. Thermal decomposition of cobalt(II) sulfate involves stepwise followed by sulfate breakdown. The heptahydrate loses water progressively: for instance, CoSO₄·7H₂O dehydrates to the monohydrate at approximately 110 °C: \ceCoSO4⋅7H2O−>[110∘C]CoSO4⋅H2O+6H2O\ce{CoSO4 \cdot 7H2O ->[110^\circ C] CoSO4 \cdot H2O + 6H2O} Further heating to 700 °C decomposes the form into (II) oxide and : \ceCoSO4−>[700∘C]CoO+SO3\ce{CoSO4 ->[700^\circ C] CoO + SO3} These stages are characterized by endothermic dehydration and exothermic decomposition, as confirmed by . Isotope exchange studies using radioactive (⁶⁰Co) have historically employed cobalt(II) sulfate solutions to investigate ligand exchange kinetics and transport in radiochemical labeling. For example, exchange between aqueous Co²⁺ and solid-phase cobalt compounds allows uniform labeling for tracer applications in and environmental tracing.

Safety and Environmental Impact

Health Effects

Cobalt(II) sulfate exhibits moderate via oral exposure, with an LD50 of 424 mg/kg in rats, indicating potential harm if ingested in significant quantities. Dermal exposure shows lower , with an LD50 greater than 2000 mg/kg in rats, suggesting it is less hazardous through contact under normal conditions. represents a more severe route, with an LC50 of approximately 0.05 mg/L over 4 hours in rats, leading to damage. These values highlight the compound's varying risks depending on exposure pathway, primarily affecting workers in industrial settings. Acute exposure to cobalt(II) sulfate primarily causes gastrointestinal distress, including and , alongside respiratory irritation such as coughing and . Skin contact may result in , manifesting as allergic or eczema in susceptible individuals. Chronic exposure, often occupational, is linked to , a serious heart condition involving weakened cardiac function and potential , as evidenced by the 1965 Quebec beer incident where cobalt sulfate added to beer for foaming led to 48 deaths among heavy drinkers consuming approximately 3–10 mg daily. Other chronic effects include thyroid enlargement (goiter) and potential , an increase in red blood cells. The International Agency for Research on Cancer (IARC) classifies soluble cobalt(II) salts, including cobalt(II) sulfate, as probably carcinogenic to humans (Group 2A), based on sufficient evidence of carcinogenicity in experimental animals and limited evidence in humans from occupational exposures in and production. Despite its essential role in human health— is a key component of , with adults requiring about 2.4 micrograms of B12 daily (providing trace levels)—excess intake beyond this disrupts function and exacerbates toxicity. Primary exposure routes are and dermal in occupational contexts like battery and , with accidental oral ingestion occurring via contaminated water or food. Additional 1960s case studies from industrial contamination in and reported similar outbreaks of and respiratory issues among exposed populations.

Environmental Concerns

Cobalt mining and production, especially in the Democratic Republic of Congo (DRC), generate (AMD) that releases (SO42−SO_4^{2-}) and (Co2+Co^{2+}) ions into waterways, resulting in acidification (often < 4) and heavy metal contamination. This process occurs when minerals in ore are oxidized, mobilizing toxic metals into rivers and streams. For instance, in the DRC's Kafubu River near , mining effluents have elevated concentrations in sediments to levels exceeding 100 mg/kg, exacerbating downstream . These releases lead to bioaccumulation of in and aquatic biota, disrupting ecosystems by inhibiting and causing . In aquatic environments, accumulates in and benthic organisms, with thresholds for algal growth inhibition at an EC50 of approximately 0.1 mg/L . Terrestrial exposed to contaminated soils show reduced due to damage to structures and electron transport chains. Atmospheric emissions from and , including laden with particles, further contribute to deposition and long-term around sites. Globally, the DRC supplies about 70% of the world's , linking extraction to significant , with activities contributing to the loss of approximately 13,000 hectares of tree cover from 2001 to 2023, and through habitat loss and resource diversion. The post-2020 surge in battery demand has amplified e-waste generation, with improper disposal leading to leaching from batteries into soils and sediments, where CoSO4 persists due to low and binding to . efforts can mitigate these impacts by reducing energy use by up to 46% and water consumption by 40% compared to . offers a sustainable approach, employing plants like Alyssum murale to extract from soils at rates exceeding 100 mg/kg dry biomass. Contaminated water from these sources may indirectly affect human health through ecosystem exposure pathways.

Regulatory Status

Cobalt(II) sulfate is subject to stringent occupational exposure limits due to its potential health risks. In the United States, the (OSHA) has established a (PEL) of 0.1 mg/m³ for (as Co) averaged over an 8-hour time-weighted average (). The National Institute for Occupational Safety and Health (NIOSH) recommends a lower (REL) of 0.05 mg/m³ as a 10-hour to further protect workers from respiratory and other effects. Environmental regulations also impose limits on cobalt discharges to protect aquatic ecosystems. Under the European Union's , proposed environmental quality standards (EQS) for cobalt in freshwater typically set annual average concentrations at around 1 µg/L to prevent adverse effects on biota, though binding limits vary by member state. In the United States, the Environmental Protection Agency (EPA) enforces effluent limitations for industrial discharges containing cobalt, with technology-based standards under the Clean Water Act reaching as low as 0.11 mg/L in specific nonferrous metals manufacturing subcategories to minimize . In the , cobalt(II) sulfate is registered under regulation and classified as toxic to reproduction (Category 1B), a sensitizer (Category 1), a respiratory sensitizer (Category 1), mutagenic (Category 2), and carcinogenic (Category 1B via inhalation). Authorization is required for uses exceeding 1 ton per year due to its status as a (SVHC), ensuring risk management measures are in place. Global trade of cobalt(II) sulfate and related wastes is regulated under the , which controls transboundary movements of hazardous wastes containing to prevent environmental harm in developing countries. The 2023 U.S. provides tax incentives for ethical sourcing of critical minerals like , prioritizing supply chains free from forced labor and aligned with agreements to support clean energy production. Recent regulatory developments include the 2024 EU Battery Regulation (EU 2023/1542), which mandates a minimum of 16% recycled content in lithium-ion batteries placed on the market after August 18, 2031, to promote practices. In Volume 131 (2023, based on the 2022 meeting), the International Agency for Research on Cancer (IARC) classified soluble cobalt(II) salts, including cobalt(II) sulfate, as probably carcinogenic to humans (Group 2A), upgrading from the previous Group 2B classification based on sufficient evidence in experimental animals and limited human data. Under the Globally Harmonized System (GHS), cobalt(II) sulfate requires labeling with pictograms for (GHS07), health hazards including carcinogenicity (GHS08), and aquatic environmental toxicity (GHS09), along with the signal word "Danger" and hazard statements covering , , and long-term effects.

Hydrated Forms

Cobalt(II) sulfate exists in several hydrated forms, each characterized by distinct physical properties and stability conditions influenced by temperature and humidity. The form, CoSO₄, appears as a pale red powder and is prepared by dehydrating the hydrated salts at temperatures above 250 °C. It is hygroscopic and unstable in moist air, readily absorbing to form lower hydrates. The monohydrate, CoSO₄·H₂O, is a crystalline solid formed by heating the higher hydrates between 100 °C and 200 °C. It has a of 3.075 g/cm³ and melts at approximately 100 °C, making it suitable for certain analytical applications where a stable, lower-hydration source of (II) ions is required, such as in reagent preparations for and analysis. The hexahydrate, CoSO₄·6H₂O, consists of efflorescent red crystals that are stable under low humidity conditions (below 70% relative humidity at room temperature) but lose water in drier environments. It exhibits a transition point around 41.5 °C, above which it dehydrates further, and forms between 44 °C and 70 °C from aqueous solutions. The heptahydrate, CoSO₄·7H₂O, is the most common and stable form under typical laboratory conditions, appearing as red, deliquescent monoclinic prismatic crystals with a density of 1.948–2.03 g/cm³. It remains stable at humidities above 70% and room temperature but dehydrates at around 41 °C to the hexahydrate; phase stability is governed by a humidity-temperature diagram where higher water contents favor the heptahydrate below 44 °C. Interconversion among these forms follows a dehydration sequence starting from the heptahydrate: at 20–40 °C, it loses one molecule to form the hexahydrate; further heating to 70–100 °C yields the monohydrate; and complete to the form occurs above 250 °C. These transitions are reversible under appropriate conditions, with the and monohydrate forms rehydrating in moist air. Varying content significantly impacts properties such as and color intensity. The heptahydrate exhibits the highest , approximately 36.2 g/100 mL in at 20 °C, decreasing slightly for lower hydrates due to differences. Color intensity also increases with hydration level, from the pale red of the form to the deeper red of the hexahydrate and heptahydrate, attributable to octahedral coordination of ligands around the Co²⁺ .

Analogous Compounds

Cobalt(II) sulfate shares structural and chemical similarities with other transition metal sulfates, particularly those featuring divalent cations in hydrated forms. Nickel(II) sulfate hexahydrate (NiSO₄·6H₂O) is emerald-green and isomorphous with its cobalt analog, exhibiting a comparable monoclinic crystal structure based on octahedral [Ni(H₂O)₆]²⁺ units hydrogen-bonded to sulfate ions. Copper(II) sulfate pentahydrate (CuSO₄·5H₂O) appears as blue triclinic crystals and is employed similarly in pigment applications for paints and varnishes, though its pentahydrate form contrasts with the typical hexahydrate of cobalt(II) sulfate. Among cobalt halides, cobalt(II) chloride hexahydrate (CoCl₂·6H₂O) is notable for its role as a humidity indicator, undergoing a reversible color shift from blue (anhydrous or low humidity) to pink upon hydration due to changes in coordination environment. Double salts provide another point of comparison; Mohr's salt ((NH₄)₂Fe(SO₄)₂·6H₂O) enhances the air stability of iron(II) ions against oxidation compared to simple ferrous sulfate, whereas cobalt(II) sulfate maintains inherent stability without requiring such ammonium stabilization. Compounds in higher oxidation states of differ markedly in reactivity. Cobalt(III) sulfate (Co₂(SO₄)₃) exists primarily as a hydrated form (Co₂(SO₄)₃·18H₂O) but is unstable at , decomposing readily and acting as a strong in acidic solutions.[](https://www.researchgate.net/publication/243661056_The_preparation_of_cobaltIII_sulphate_and_its_alums_and_the_magnetic_spectroscopic_and_crystallographic_properties_of_the_CoH2O6_3_ ion) Related compounds include heptahydrate (, MgSO₄·7H₂O), which adopts an orthorhombic lattice, in contrast to the monoclinic structure of cobalt(II) sulfate heptahydrate, but both feature octahedral coordination geometries around the metal center and share the heptahydrate , though MgSO₄·7H₂O lacks the pink coloration and arising from the d⁷ configuration of Co²⁺, remaining colorless and diamagnetic due to Mg²⁺.

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