Jump to content

Copper(II) sulfate

From Wikipedia, the free encyclopedia
(Redirected from CuSO4·5H2O)
Copper(II) sulfate

Crystals of CuSO4·5H2O

  Copper, Cu
  Sulfur, S
  Oxygen, O
  Hydrogen, H

Portion of the structure of the pentahydrate
(sulfate links Cu(H2O)2+4 centers)

Unit cell of the crystal structure of CuSO4·5H2O
with hydrogen bonds in black[1]
Names
IUPAC name
Copper(II) sulfate
Other names
  • Cupric sulphate
  • Blue vitriol (pentahydrate)
  • Bluestone (pentahydrate)
  • Bonattite (trihydrate mineral)
  • Boothite (heptahydrate mineral)
  • Chalcanthite (pentahydrate mineral)
  • Chalcocyanite (mineral)
Copper Sulphate pentahydrate
Identifiers
3D model (JSmol)
ChEBI
ChEMBL
ChemSpider
ECHA InfoCard 100.028.952 Edit this at Wikidata
EC Number
  • 231-847-6
8294
KEGG
RTECS number
  • GL8800000 (anhydrous)
    GL8900000 (pentahydrate)
UNII
  • 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
CuSO4 (anhydrous)
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
GHS labelling:
GHS07: Exclamation markGHS09: Environmental hazard
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 1: Normally stable, but can become unstable at elevated temperatures and pressures. E.g. calciumSpecial hazards (white): no code
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
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 ?)

Copper(II) sulfate is an inorganic compound with the chemical formula CuSO4. It forms hydrates 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 [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 Cu(II)(H2O)4 centers are interconnected by sulfate anions to form chains.[15]

Preparation and occurrence

[edit]
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

[edit]

Copper(II) sulfate pentahydrate decomposes before melting. It loses two water molecules upon heating at 63 °C (145 °F), followed by two more at 109 °C (228 °F) and the final water molecule at 200 °C (392 °F).[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):

Cu2+ + 4 Cl− → [CuCl4]2−

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

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:

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]

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 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.[22] Copper(II) sulfate pentahydrate can easily be produced by crystallization from solution as copper(II) sulfate, which is hygroscopic.

Uses

[edit]

As a fungicide and herbicide

[edit]

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".[23][24]

Bordeaux mixture, a suspension of copper(II) sulfate (CuSO4) and calcium hydroxide (Ca(OH)2), is used to control fungus on grapes, melons, and other berries.[25] 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,[26] and is also used to remove snails from aquariums and zebra mussels from water pipes.[27] Copper ions are highly toxic to fish. Most species of algae can be controlled with very low concentrations of copper sulfate.

Analytical reagent

[edit]

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.[28] 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

[edit]

Copper sulfate is employed at a limited level in organic synthesis.[29] The anhydrous salt is used as a dehydrating agent for forming and manipulating acetal groups.[30] The hydrated salt can be intimately mingled with potassium permanganate to give an oxidant for the conversion of primary alcohols.[31]

Rayon production

[edit]

Reaction with ammonium hydroxide yields tetraamminecopper(II) sulfate or Schweizer's reagent which was used to dissolve cellulose in the industrial production of Rayon.

Niche uses

[edit]

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.[32] Copper sulfate is also used in firework manufacture as a blue coloring agent, but it is not safe to mix copper sulfate with chlorates when mixing firework powders.[33]

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.[32]

Art

[edit]

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

[edit]

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

[edit]

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

[edit]

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 CuSO4·5H2O and sulfuric acid (H2SO4) is often used for electrodeposition of copper.[40]

Other forms of copper sulfate

[edit]

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]

Forms of copper(II) sulfate
  • Anhydrous CuSO4
    Anhydrous CuSO4
  • Copper(II) sulfate monohydrate
    Copper(II) sulfate monohydrate
  • Copper(II) sulfate pentahydrate
    Copper(II) sulfate pentahydrate
  • The rare mineral boothite (CuSO4·7H2O)
    The rare mineral boothite (CuSO4·7H2O)

Toxicological effects

[edit]
See also: Copper toxicity

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

[edit]

References

[edit]
  1. ^ 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. ^ Rumble, John, ed. (2018). CRC Handbook of Chemistry and Physics (99th ed.). CRC Press, Taylor & Francis Group. pp. 5–179. ISBN 9781138561632.
  4. ^ 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. ^ 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. ^ 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 NIOSH Pocket Guide to Chemical Hazards. "#0150". National Institute for Occupational Safety and Health (NIOSH).
  9. ^ Cupric sulfate. US National Institutes of Health
  10. ^ Connor, Nick (2023-07-24). "Copper (II) Sulfate | Formula, Properties & Application". Material Properties. Retrieved 2024-02-03.
  11. ^ Foundation, In association with Nuffield. "A reversible reaction of hydrated copper(II) sulfate". RSC Education. Retrieved 2024-02-03.
  12. ^ "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. ^ 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 "Uses of Copper Compounds: Copper Sulphate". copper.org. Copper Development Association Inc. Retrieved 10 May 2015.
  17. ^ Andrew Knox Galwey; Michael E. Green (1999). Thermal decomposition of ionic solids. Elsevier. pp. 228–229. ISBN 978-0-444-82437-0.
  18. ^ Wiberg, Egon; Nils Wiberg; Arnold Frederick Holleman (2001). Inorganic chemistry. Academic Press. p. 1263. ISBN 978-0-12-352651-9.
  19. ^ Ray Q. Brewster, Theodore Groening (1934). "P-Nitrophenyl Ether". Organic Syntheses. 14: 66. doi:10.15227/orgsyn.014.0066.
  20. ^ Zumdahl, Steven; DeCoste, Donald (2013). Chemical Principles. Cengage Learning. pp. 506–507. ISBN 978-1-285-13370-6.
  21. ^ 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. ^ "Process for the preparation of stable copper(II) sulfate monohydrate applicable as trace element additive in animal fodders". Retrieved 2009-07-07.
  23. ^ 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.
  24. ^ 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.
  25. ^ 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. Retrieved 2007-12-31.
  26. ^ "All About Copper Sulfate". National Fish Pharmaceuticals. Retrieved 2007-12-31.
  27. ^ "With Zebra mussels here to stay, Austin has a plan to avoid stinky drinking water". KXAN Austin. 2020-10-26. Retrieved 2020-10-28.
  28. ^ 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.
  29. ^ 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.
  30. ^ Philip J. Kocienski (2005). Protecting Groups. Thieme. p. 58. ISBN 978-1-58890-376-1.
  31. ^ 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; Collected Volumes, vol. 9, p. 462.
  32. ^ a b Copper Development Association. "Uses of Copper Compounds: Table A - Uses of Copper Sulphate". copper. Copper Development Association Inc. Retrieved 12 May 2015.
  33. ^ Partin, Lee. "The Blues: Part 2". skylighter. Skylighter.Inc. Archived from the original on 21 December 2010. Retrieved 12 May 2015.
  34. ^ 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. ^ "Seizure". Artangel.org.uk. Retrieved 2021-10-05.
  36. ^ "Roger Hiorns: Seizure". Yorkshire Sculpture Park. Archived from the original on 2015-02-22. Retrieved 2015-02-22.
  37. ^ greenart.info, Bordeau etch, 2009-01-18, retrieved 2011-06-02.
  38. ^ ndiprintmaking.ca, The Chemistry of using Copper Sulfate Mordant, 2009-04-12, retrieved 2011-06-02.
  39. ^ http://mordent.com/etch-howto/, How to Electrolytically etch in copper, brass, steel, nickel silver or silver, retrieved 2015-05-2015.
  40. ^ 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. ^ "Chalcocyanite". www.mindat.org.
  42. ^ "Bonattite". www.mindat.org.
  43. ^ "Boothite". www.mindat.org.
  44. ^ "Poitevinite". www.mindat.org.
  45. ^ a b "List of Minerals". www.ima-mineralogy.org. March 21, 2011.
  46. ^ "Langite". www.mindat.org.
  47. ^ "Posnjakite". www.mindat.org.
  48. ^ Windholz, M., ed. 1983. The Merck Index. Tenth edition. Rahway, NJ: Merck and Company.
  49. ^ 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. ^ 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. ^ Olson, Kent C. (2004). Poisoning & drug overdose. New York: Lange Medical Mooks/McGraw-Hill. p. 175. ISBN 978-0-8385-8172-8.
  52. ^ V03AB20 (WHO)

Bibliography

[edit]
  • Haynes, William M., ed. (2011). CRC Handbook of Chemistry and Physics (92nd ed.). Boca Raton, FL: CRC Press. ISBN 978-1439855119.
[edit]
Retrieved from "https://wiki.riteme.site/w/index.php?title=Copper(II)_sulfate&oldid=1264096091"