Copper(II) sulfate

From Wikipedia, the free encyclopedia
(Redirected from E519)

Page Template:Chembox/styles.css has no content.

Template:Chembox IndexlistTemplate:Chembox CompToxTemplate:Chembox GHS (set)Template:Chembox Datapage check
Copper(II) sulfate
Lua error in package.lua at line 80: module 'Module:InfoboxImage/data' not found.
Crystals of Page Module:Chem2/styles.css has no content.CuSO4·5H2O
Lua error in package.lua at line 80: module 'Module:InfoboxImage/data' not found.
Page Template:Legend/styles.css has no content.
  Copper, Cu
Page Template:Legend/styles.css has no content.
  Sulfur, S
Page Template:Legend/styles.css has no content.
  Oxygen, O
Page Template:Legend/styles.css has no content.
  Hydrogen, H

Portion of the structure of the pentahydrate
(sulfate links Page Module:Chem2/styles.css has no content.Cu(H2O)2+4 centers)
Lua error in package.lua at line 80: module 'Module:InfoboxImage/data' not found.
Unit cell of the crystal structure of Page Module:Chem2/styles.css has no content.CuSO4·5H2O
with hydrogen bonds in black[1]
Names
IUPAC name
Copper(II) sulfate
Other names
Page Template:Plainlist/styles.css has no content.
  • Cupric sulphate
  • Blue vitriol (pentahydrate)
  • Bluestone (pentahydrate)
  • Bonattite (trihydrate mineral)
  • Boothite (heptahydrate mineral)
  • Chalcanthite (pentahydrate mineral)
  • Chalcocyanite (mineral)
Copper Sulphate pentahydrate
Identifiers
Page Template:Plainlist/styles.css has no content.
3D model (JSmol)
 Page Template:Plainlist/styles.css has no content.
ChEBI Page Template:Plainlist/styles.css has no content.
ChEMBL Page Template:Plainlist/styles.css has no content.
ChemSpider Page Template:Plainlist/styles.css has no content.
EC Number Page Template:Plainlist/styles.css has no content.
8294
KEGG Page Template:Plainlist/styles.css has no content.
Page Template:Plainlist/styles.css has no content.
RTECS number Page Template:Plainlist/styles.css has no content.
UNII Page Template:Plainlist/styles.css has no content.
  • InChI=1S/Cu.H2O4S/c;1-5(2,3)4/h;(H2,1,2,3,4)/q+2;/p-2 checkY
    Key: ARUVKPQLZAKDPS-UHFFFAOYSA-L checkY
  • InChI=1/Cu.H2O4S/c;1-5(2,3)4/h;(H2,1,2,3,4)/q+2;/p-2
    Key: ARUVKPQLZAKDPS-NUQVWONBAI
  • [O-]S(=O)(=O)[O-].[Cu+2]
Properties
Page Module:Chem2/styles.css has no content.CuSO4 (anhydrous)
Page Module:Chem2/styles.css has no content.CuSO4·5H2O (pentahydrate)
Molar mass 159.60 g/mol (anhydrous)[2]
249.685 g/mol (pentahydrate)[2]
Appearance gray-white (anhydrous)
blue (pentahydrate)
Density 3.60 g/cm3 (anhydrous)[2]
2.286 g/cm3 (pentahydrate)[2]
Melting point 110 °C (230 °F; 383 K) decomposes

560 °C decomposes[2](pentahydrate)

Fully decomposes at 590 °C (anhydrous)

Boiling point decomposes to cupric oxide at 650 °C
pentahydrate
316 g/L (0 °C)
2033 g/L (100 °C)
anhydrous
168 g/L (10 °C)
201 g/L (20 °C)
404 g/L (60 °C)
770 g/L (100 °C)[3]
Solubility anhydrous
insoluble in ethanol[2]
pentahydrate
soluble in methanol[2]
10.4 g/L (18 °C)
insoluble in ethanol and acetone
1330·10−6 cm3/mol
1.724–1.739 (anhydrous)[4]
1.514–1.544 (pentahydrate)[5]
Structure
Orthorhombic (anhydrous, chalcocyanite), space group Pnma, oP24, a = 0.839 nm, b = 0.669 nm, c = 0.483 nm.[6]
Triclinic (pentahydrate), space group P1, aP22, a = 0.5986 nm, b = 0.6141 nm, c = 1.0736 nm, α = 77.333°, β = 82.267°, γ = 72.567°[7]
Thermochemistry
5 J/(K·mol)
−769.98 kJ/mol
Pharmacology
V03AB20 (WHO)
Hazards
NFPA 704 (fire diamond) Page Template:NFPA 704 diamond/styles.css has no content.

Error: Image is invalid or non-existent.

2
0
1
Flash point Non-flammable
Lethal dose or concentration (LD, LC):
300 mg/kg (oral, rat)[9]

87 mg/kg (oral, mouse)

NIOSH (US health exposure limits):
PEL (Permissible)
 TWA 1 mg/m3 (as Cu)[8]
REL (Recommended)
 TWA 1 mg/m3 (as Cu)[8]
IDLH (Immediate danger)
 TWA 100 mg/m3 (as Cu)[8]
Safety data sheet (SDS) anhydrous
pentahydrate
Related compounds
Other cations
 Page Template:Plainlist/styles.css has no content.
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).

Template:Chembox Footer/trackingTemplate:Short description

Copper(II) sulfate is an inorganic compound with the chemical formula Page Module:Chem2/styles.css has no content.CuSO4. It forms hydrates Page Module:Chem2/styles.css has no content.CuSO4·nH2O, where n can range from 1 to 7. The pentahydrate (n = 5), a bright blue crystal, is the most commonly encountered hydrate of copper(II) sulfate,[10] while its anhydrous form is white.[11] Older names for the pentahydrate include blue vitriol, bluestone,[12] vitriol of copper,[13] and Roman vitriol.[14] It exothermically dissolves in water to give the aquo complex Page Module:Chem2/styles.css has no content.[Cu(H2O)6]2+, which has octahedral molecular geometry. The structure of the solid pentahydrate reveals a polymeric structure wherein copper is again octahedral but bound to four water ligands. The Page Module:Chem2/styles.css has no content.Cu(II)(H2O)4 centers are interconnected by sulfate anions to form chains.[15]

Preparation and occurrence

File:Synthesizing Copper Sulfate.jpg
Preparation of copper(II) sulfate by electrolyzing sulfuric acid, using copper electrodes

Copper sulfate is produced industrially by treating copper metal with hot concentrated sulfuric acid or copper oxides with dilute sulfuric acid. For laboratory use, copper sulfate is usually purchased. Copper sulfate can also be produced by slowly leaching low-grade copper ore in air; bacteria may be used to hasten the process.[16]

Commercial copper sulfate is usually about 98% pure copper sulfate, and may contain traces of water. Anhydrous copper sulfate is 39.81% copper and 60.19% sulfate by mass, and in its blue, hydrous form, it is 25.47% copper, 38.47% sulfate (12.82% sulfur) and 36.06% water by mass. Four types of crystal size are provided based on its usage: large crystals (10–40 mm), small crystals (2–10 mm), snow crystals (less than 2 mm), and windswept powder (less than 0.15 mm).[16]

Chemical properties

Copper(II) sulfate pentahydrate decomposes before melting. It loses two water molecules upon heating at Template:Cvt, followed by two more at Template:Cvt and the final water molecule at Template:Cvt.[17][18]

The chemistry of aqueous copper sulfate is simply that of copper aquo complex, since the sulfate is not bound to copper in such solutions. Thus, such solutions react with concentrated hydrochloric acid to give tetrachlorocuprate(II):

Page Module:Chem2/styles.css has no content.Cu2+ + 4 Cl− → [CuCl4]2−

Similarly treatment of such solutions with zinc gives metallic copper, as described by this simplified equation:[19]

Page Module:Chem2/styles.css has no content.CuSO4 + Zn → Cu + ZnSO4

A further illustration of such single metal replacement reactions occurs when a piece of iron is submerged in a solution of copper sulfate:

Page Module:Chem2/styles.css has no content.Fe + CuSO4 → FeSO4 + Cu

In high school and general chemistry education, copper sulfate is used as an electrolyte for galvanic cells, usually as a cathode solution. For example, in a zinc/copper cell, copper ion in copper sulfate solution absorbs electron from zinc and forms metallic copper.[20]

Page Module:Chem2/styles.css has no content.Cu2+ + 2e− → Cu (cathode), E°cell = 0.34 V

Copper sulfate is commonly included in teenage chemistry sets and undergraduate experiments.[21] It is often used to grow crystals in schools and in Copper electroplating experiments despite its toxicity. Copper sulfate is often used to demonstrate an exothermic reaction, in which steel wool or magnesium ribbon is placed in an aqueous solution of Page Module:Chem2/styles.css has no content.CuSO4. It is used to demonstrate the principle of mineral hydration. The pentahydrate form, which is blue, is heated, turning the copper sulfate into the anhydrous form which is white, while the water that was present in the pentahydrate form evaporates. When water is then added to the anhydrous compound, it turns back into the pentahydrate form, regaining its blue color.[citation needed] Copper(II) sulfate pentahydrate can easily be produced by crystallization from solution as copper(II) sulfate, which is hygroscopic.

Uses

As a fungicide and herbicide

Copper sulfate has been used for control of algae in lakes and related fresh waters subject to eutrophication. It "remains the most effective algicidal treatment".[22][23]

Bordeaux mixture, a suspension of copper(II) sulfate (Page Module:Chem2/styles.css has no content.CuSO4) and calcium hydroxide (Page Module:Chem2/styles.css has no content.Ca(OH)2), is used to control fungus on grapes, melons, and other berries.[24] It is produced by mixing a water solution of copper sulfate and a suspension of slaked lime.

A dilute solution of copper sulfate is used to treat aquarium fishes for parasitic infections,[25] and is also used to remove snails from aquariums and zebra mussels from water pipes.[26] Copper ions are highly toxic to fish. Most species of algae can be controlled with very low concentrations of copper sulfate.

Analytical reagent

Several chemical tests utilize copper sulfate. It is used in Fehling's solution and Benedict's solution to test for reducing sugars, which reduce the soluble blue copper(II) sulfate to insoluble red copper(I) oxide. Copper(II) sulfate is also used in the Biuret reagent to test for proteins.

Copper sulfate is used to test blood for anemia. The blood is dropped into a solution of copper sulfate of known specific gravity—blood with sufficient hemoglobin sinks rapidly due to its density, whereas blood which sinks slowly or not at all has an insufficient amount of hemoglobin.[27] Clinically relevant, however, modern laboratories utilize automated blood analyzers for accurate quantitative hemoglobin determinations, as opposed to older qualitative means.[citation needed]

In a flame test, the copper ions of copper sulfate emit a deep green light, a much deeper green than the flame test for barium.

Organic synthesis

Copper sulfate is employed at a limited level in organic synthesis.[28] The anhydrous salt is used as a dehydrating agent for forming and manipulating acetal groups.[29] The hydrated salt can be intimately mingled with potassium permanganate to give an oxidant for the conversion of primary alcohols.[30]

Rayon production

Reaction with ammonium hydroxide can yield tetraamminecopper(II) sulfate and Schweizer's reagent (nonsulphurous), which is used to dissolve cellulose in the industrial production of Rayon.

Niche uses

Copper(II) sulfate has attracted many niche applications over the centuries. In industry copper sulfate has multiple applications. In printing it is an additive to book-binding pastes and glues to protect paper from insect bites; in building it is used as an additive to concrete to improve water resistance and prevent plant and mushroom growth. Copper sulfate can be used as a coloring ingredient in artworks, especially glasses and potteries.[31] Copper sulfate is also rarely used in firework manufacture as a blue coloring agent, but it is not safe to mix copper(II) sulfate with metal powders, or it or any copper(II) compound with chlorates; the sulfate and other copper(II) compounds are not allowed in chlorate containing mixtures in the US.[32][33]

File:Copper sulfate etching bath.JPG
Lowering a copper etching plate into the copper sulfate solution

Copper sulfate was once used to kill bromeliads, which serve as mosquito breeding sites.[34] Copper sulfate is used as a molluscicide to treat bilharzia in tropical countries.[31]

Art

In 2008, the artist Roger Hiorns filled an abandoned waterproofed council flat in London with 75,000 liters of copper(II) sulfate water solution. The solution was left to crystallize for several weeks before the flat was drained, leaving crystal-covered walls, floors and ceilings. The work is titled Seizure.[35] Since 2011, it has been on exhibition at the Yorkshire Sculpture Park.[36]

Etching

Copper(II) sulfate is used to etch zinc, aluminium, or copper plates for intaglio printmaking.[37][38] It is also used to etch designs into copper for jewelry, such as for Champlevé.[39]

Dyeing

Copper(II) sulfate can be used as a mordant in vegetable dyeing. It often highlights the green tints of the specific dyes.[citation needed]

Electronics

An aqueous solution of copper(II) sulfate is often used as the resistive element in liquid resistors.[citation needed]

In electronic and microelectronic industry a bath of Page Module:Chem2/styles.css has no content.CuSO4·5H2O and sulfuric acid (Page Module:Chem2/styles.css has no content.H2SO4) is often used for electrodeposition of copper.[40]

Other forms of copper sulfate

Anhydrous copper(II) sulfate can be produced by dehydration of the commonly available pentahydrate copper sulfate. In nature, it is found as the very rare mineral known as chalcocyanite.[41] The pentahydrate also occurs in nature as chalcanthite. Other rare copper sulfate minerals include bonattite (trihydrate),[42] boothite (heptahydrate),[43] and the monohydrate compound poitevinite.[44][45] There are numerous other, more complex, copper(II) sulfate minerals known, with environmentally important basic copper(II) sulfates like langite and posnjakite.[45][46][47] Lua error in package.lua at line 80: module 'Module:Plain text' not found.

Toxicological effects

Script error: No such module "Labelled list hatnote". Copper(II) salts have an LD50 of 100 mg/kg.[48][49]

Copper(II) sulfate was used in the past as an emetic.[50] It is now considered too toxic for this use.[51] It is still listed as an antidote in the World Health Organization's Anatomical Therapeutic Chemical Classification System.[52]

See also

References

Page Template:Reflist/styles.css has no content.

  1. ^ Page Module:Citation/CS1/styles.css has no content.Varghese, J. N.; Maslen, E. N. (1985). "Electron density in non-ideal metal complexes. I. Copper sulphate pentahydrate". Acta Crystallographica Section B. 41 (3): 184–190. doi:10.1107/S0108768185001914.
  2. ^ a b c d e f g Haynes, p. 4.62
  3. ^ Page Module:Citation/CS1/styles.css has no content.Rumble, John, ed. (2018). CRC Handbook of Chemistry and Physics (99th ed.). CRC Press, Taylor & Francis Group. pp. 5–179. ISBN 9781138561632.
  4. ^ Page Module:Citation/CS1/styles.css has no content.Anthony, John W.; Bideaux, Richard A.; Bladh, Kenneth W.; Nichols, Monte C., eds. (2003). "Chalcocyanite" (PDF). Handbook of Mineralogy. Vol. V. Borates, Carbonates, Sulfates. Chantilly, VA, US: Mineralogical Society of America. ISBN 978-0962209741.
  5. ^ Haynes, p. 10.240
  6. ^ Page Module:Citation/CS1/styles.css has no content.Kokkoros, P. A.; Rentzeperis, P. J. (1958). "The crystal structure of the anhydrous sulphates of copper and zinc". Acta Crystallographica. 11 (5): 361–364. doi:10.1107/S0365110X58000955.
  7. ^ Page Module:Citation/CS1/styles.css has no content.Bacon, G. E.; Titterton, D. H. (1975). "Neutron-diffraction studies of CuSO4 · 5H2O and CuSO4 · 5D2O". Z. Kristallogr. 141 (5–6): 330–341. Bibcode:1975ZK....141..330B. doi:10.1524/zkri.1975.141.5-6.330.
  8. ^ a b c Page Module:Citation/CS1/styles.css has no content.NIOSH Pocket Guide to Chemical Hazards. "#0150". National Institute for Occupational Safety and Health (NIOSH).
  9. ^ Cupric sulfate. US National Institutes of Health
  10. ^ Page Module:Citation/CS1/styles.css has no content.Connor, Nick (2023-07-24). "Copper (II) Sulfate | Formula, Properties & Application". Material Properties. Retrieved 2024-02-03.
  11. ^ Page Module:Citation/CS1/styles.css has no content.Foundation, In association with Nuffield. "A reversible reaction of hydrated copper(II) sulfate". RSC Education. Retrieved 2024-02-03.
  12. ^ Page Module:Citation/CS1/styles.css has no content."Copper (II) sulfate MSDS". Oxford University. Archived from the original on 2007-10-11. Retrieved 2007-12-31.
  13. ^ Antoine-François de Fourcroy, tr. by Robert Heron (1796) "Elements of Chemistry, and Natural History: To which is Prefixed the Philosophy of Chemistry". J. Murray and others, Edinburgh. Page 348.
  14. ^ Oxford University Press, "Roman vitriol", Oxford Living Dictionaries. Accessed on 2016-11-13
  15. ^ Page Module:Citation/CS1/styles.css has no content.Ting, V. P.; Henry, P. F.; Schmidtmann, M.; Wilson, C. C.; Weller, M. T. (2009). "In situ neutron powder diffraction and structure determination in controlled humidities". Chem. Commun. 2009 (48): 7527–7529. doi:10.1039/B918702B. PMID 20024268.
  16. ^ a b Page Module:Citation/CS1/styles.css has no content."Uses of Copper Compounds: Copper Sulphate". copper.org. Copper Development Association Inc. Retrieved 10 May 2015.
  17. ^ Page Module:Citation/CS1/styles.css has no content.Andrew Knox Galwey; Michael E. Green (1999). Thermal decomposition of ionic solids. Elsevier. pp. 228–229. ISBN 978-0-444-82437-0.
  18. ^ Page Module:Citation/CS1/styles.css has no content.Wiberg, Egon; Nils Wiberg; Arnold Frederick Holleman (2001). Inorganic chemistry. Academic Press. p. 1263. ISBN 978-0-12-352651-9.
  19. ^ Page Module:Citation/CS1/styles.css has no content.Ray Q. Brewster, Theodore Groening (1934). "P-Nitrophenyl Ether". Organic Syntheses. 14: 66. doi:10.15227/orgsyn.014.0066.
  20. ^ Page Module:Citation/CS1/styles.css has no content.Zumdahl, Steven; DeCoste, Donald (2013). Chemical Principles. Cengage Learning. pp. 506–507. ISBN 978-1-285-13370-6.
  21. ^ Page Module:Citation/CS1/styles.css has no content.Rodríguez, Emilio; Vicente, Miguel Angel (2002). "A Copper-Sulfate-Based Inorganic Chemistry Laboratory for First-Year University Students That Teaches Basic Operations and Concepts". Journal of Chemical Education. 79 (4): 486. Bibcode:2002JChEd..79..486R. doi:10.1021/ed079p486.
  22. ^ Page Module:Citation/CS1/styles.css has no content.Van Hullebusch, E.; Chatenet, P.; Deluchat, V.; Chazal, P. M.; Froissard, D.; Lens, P. N.L.; Baudu, M. (2003). "Fate and forms of Cu in a reservoir ecosystem following copper sulfate treatment (Saint Germain les Belles, France)". Journal de Physique IV (Proceedings). 107: 1333–1336. doi:10.1051/jp4:20030547.
  23. ^ Page Module:Citation/CS1/styles.css has no content.Haughey, M. (2000). "Forms and fate of Cu in a source drinking water reservoir following CuSO4 treatment". Water Research. 34 (13): 3440–3452. doi:10.1016/S0043-1354(00)00054-3.
  24. ^ Page Module:Citation/CS1/styles.css has no content.Martin, Hubert (1933). "Uses of Copper Compounds: Copper Sulfate's Role in Agriculture". Annals of Applied Biology. 20 (2): 342–363. doi:10.1111/j.1744-7348.1933.tb07770.x.
  25. ^ Page Module:Citation/CS1/styles.css has no content."All About Copper Sulfate". National Fish Pharmaceuticals. Retrieved 2007-12-31.
  26. ^ Page Module:Citation/CS1/styles.css has no content."With Zebra mussels here to stay, Austin has a plan to avoid stinky drinking water". KXAN Austin. 2020-10-26. Retrieved 2020-10-28.
  27. ^ Page Module:Citation/CS1/styles.css has no content.Estridge, Barbara H.; Anna P. Reynolds; Norma J. Walters (2000). Basic Medical Laboratory Techniques. Thomson Delmar Learning. p. 166. ISBN 978-0-7668-1206-2.
  28. ^ Page Module:Citation/CS1/styles.css has no content.Hoffman, R. V. (2001). "Copper(II) Sulfate". Copper(II) Sulfate, in Encyclopedia of Reagents for Organic Synthesis. John Wiley & Sons. doi:10.1002/047084289X.rc247. ISBN 978-0471936237.
  29. ^ Page Module:Citation/CS1/styles.css has no content.Philip J. Kocienski (2005). Protecting Groups. Thieme. p. 58. ISBN 978-1-58890-376-1.
  30. ^ Page Module:Citation/CS1/styles.css has no content.Jefford, C. W.; Li, Y.; Wang, Y. "A Selective, Heterogeneous Oxidation using a Mixture of Potassium Permanganate and Cupric Sulfate: (3aS,7aR)-Hexahydro-(3S,6R)-Dimethyl-2(3H)-Benzofuranone". Organic Syntheses; Page Module:Citation/CS1/styles.css has no content.Collected Volumes, vol. 9, p. 462.
  31. ^ a b Page Module:Citation/CS1/styles.css has no content.Copper Development Association. "Uses of Copper Compounds: Table A - Uses of Copper Sulphate". copper. Copper Development Association Inc. Retrieved 12 May 2015.
  32. ^ Page Module:Citation/CS1/styles.css has no content.Partin, Lee. "The Blues: Part 2". skylighter. Skylighter.Inc. Retrieved 12 May 2015.
  33. ^ Page Module:Citation/CS1/styles.css has no content."Approved and Prohibited Fireworks Chemicals, v2" (PDF). phmsa.dot.gov. US Department of Transportation - Pipeline and Hazardous Materials Safety Administration. 5 November 2019. p. 1. Retrieved 3 August 2025. Copper Sulfate...Color Agent...Prohibited if mixed with a chlorate
  34. ^ Page Module:Citation/CS1/styles.css has no content.Despommier; Gwadz; Hotez; Knirsch (June 2005). Parasitic Disease (5 ed.). NY: Apple Tree Production L.L.C. pp. Section 4.2. ISBN 978-0970002778. Retrieved 12 May 2015.
  35. ^ Page Module:Citation/CS1/styles.css has no content."Seizure". Artangel.org.uk. Retrieved 2021-10-05.
  36. ^ Page Module:Citation/CS1/styles.css has no content."Roger Hiorns: Seizure". Yorkshire Sculpture Park. Archived from the original on 2015-02-22. Retrieved 2015-02-22.
  37. ^ Page Module:Citation/CS1/styles.css has no content.Cedric Green (10 August 2014). "Bordeaux Etch: etching zinc and steel plates safely (Extract from GREEN PRINTS)". www.greenart.info. Cedric Green. Retrieved 3 August 2025.
  38. ^ Page Module:Citation/CS1/styles.css has no content.Semenoff, Nik (12 April 2009). "The Chemistry of using Copper Sulfate Mordant | New Directions in Printmaking". www.ndiprintmaking.ca. Nik Semenoff. Retrieved 3 August 2025.
  39. ^ Page Module:Citation/CS1/styles.css has no content.Moretti, Nick (2015). "How to Electrolytically Etch in Copper, Brass, Steel, Nickel Silver, or Silver". mordent.com. Mordent Design. Retrieved 3 August 2025.
  40. ^ Page Module:Citation/CS1/styles.css has no content.K. Kondo; Rohan N. Akolkar; Dale P. Barkey; Masayuki Yokoi (2014). Copper Electrodeposition for Nanofabrication of Electronics Devices. New York. ISBN 978-1-4614-9176-7. OCLC 868688018.{{cite book}}: CS1 maint: location missing publisher (link)
  41. ^ Page Module:Citation/CS1/styles.css has no content."Chalcocyanite". www.mindat.org.
  42. ^ Page Module:Citation/CS1/styles.css has no content."Bonattite". www.mindat.org.
  43. ^ Page Module:Citation/CS1/styles.css has no content."Boothite". www.mindat.org.
  44. ^ Page Module:Citation/CS1/styles.css has no content."Poitevinite". www.mindat.org.
  45. ^ a b Page Module:Citation/CS1/styles.css has no content."List of Minerals". www.ima-mineralogy.org. March 21, 2011.
  46. ^ Page Module:Citation/CS1/styles.css has no content."Langite". www.mindat.org.
  47. ^ Page Module:Citation/CS1/styles.css has no content."Posnjakite". www.mindat.org.
  48. ^ Windholz, M., ed. 1983. The Merck Index. Tenth edition. Rahway, NJ: Merck and Company.
  49. ^ Page Module:Citation/CS1/styles.css has no content.Guidance for reregistration of pesticide products containing copper sulfate. Fact sheet no. 100., Washington, DC: U.S. Environmental Protection Agency, Office of Pesticide Programs, 1986
  50. ^ Page Module:Citation/CS1/styles.css has no content.Holtzmann, N. A.; Haslam, R. H. (July 1968). "Elevation of serum copper following copper sulfate as an emetic". Pediatrics. 42 (1): 189–93. doi:10.1542/peds.42.1.189. PMID 4385403. S2CID 32740524.
  51. ^ Page Module:Citation/CS1/styles.css has no content.Olson, Kent C. (2004). Poisoning & drug overdose. New York: Lange Medical Mooks/McGraw-Hill. p. 175. ISBN 978-0-8385-8172-8.
  52. ^ Template:ATC

Bibliography

Lua error in package.lua at line 80: module 'Module:Navbox/configuration' not found. Template:Salts by element Lua error in package.lua at line 80: module 'Module:Navbox/configuration' not found. Lua error in package.lua at line 80: module 'Module:Navbox/configuration' not found. Lua error in package.lua at line 80: module 'Module:Authority control/config' not found.

Retrieved from "https://cosmopedia.co/index.php?title=Copper(II)_sulfate&oldid=294382"