Copper(II) sulfate

Copper(II) sulfate
Structure of pentahydrate
Crystals of CuSO4·5H2O
Names
IUPAC name
Copper(II) sulfate
Other names
Cupric sulfate
Blue vitriol (pentahydrate)
Bluestone (pentahydrate)
Bonattite (trihydrate mineral)
Boothite (heptahydrate mineral)
Chalcanthite (pentahydrate mineral)
Chalcocyanite (mineral)
Identifiers
7758-98-7 YesY
7758-99-8 (pentahydrate) N
16448-28-5 (trihydrate) N
19086-18-1 (heptahydrate) N
3D model (Jmol) Interactive image
ChEBI CHEBI:23414 YesY
ChEMBL ChEMBL604 YesY
ChemSpider 22870 YesY
ECHA InfoCard 100.028.952
EC Number 231-847-6
E number E519 (acidity regulators, ...)
KEGG C18713 YesY
PubChem 24462
RTECS number GL8800000 (anhydrous)
GL8900000 (pentahydrate)
UNII KUW2Q3U1VV YesY
Properties
CuSO4 (anhydrous)
CuSO4·5H2O (pentahydrate)
Molar mass 159.609 g/mol (anhydrous)[1]
249.685 g/mol (pentahydrate)[1]
Appearance gray-white (anhydrous)
blue (pentahydrate)
Density 3.60 g/cm3 (anhydrous)[1]
2.286 g/cm3 (pentahydrate)[1]
Melting point 110 °C (230 °F; 383 K) decomposes (·5H2O)[1]
<560 °C decomposes[1]
1.055 molal (10 °C)
1.26 molal (20 °C)
1.502 molal (30 °C)[2]
Solubility anhydrous
insoluble in ethanol[1]
pentahydrate
soluble in methanol[1]
10.4 g/L (18 °C)
insoluble in ethanol
1.724–1.739 (anhydrous)[3]
1.514–1.544 (pentahydrate)[4]
Structure
Orthorhombic (anhydrous, chalcocyanite), space group Pnma, oP24, a = 0.839 nm, b = 0.669 nm, c = 0.483 nm.[5]
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°[6]
Thermochemistry
5 J K−1 mol−1
−769.98 kJ/mol
Pharmacology
V03AB20 (WHO)
Hazards
Safety data sheet anhydrous
pentahydrate
GHS pictograms
Harmful (Xn)
Irritant (Xi)
Dangerous for the environment (N)
R-phrases R22, R36/38, R50/53
S-phrases (S2), S22, S60, S61
NFPA 704
Flammability code 0: Will not burn. E.g., water Health code 2: Intense or continued but not chronic exposure could cause temporary incapacitation or possible residual injury. E.g., chloroform Reactivity code 1: Normally stable, but can become unstable at elevated temperatures and pressures. E.g., calcium Special hazards (white): no codeNFPA 704 four-colored diamond
0
2
1
Flash point Non-flammable
Lethal dose or concentration (LD, LC):
300 mg/kg (oral, rat)[7]
US health exposure limits (NIOSH):
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]
Related compounds
Other cations
 Iron(II) sulfate
Manganese(II) sulfate
Nickel(II) sulfate
Zinc sulfate
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
N verify (what is YesYN ?)
Infobox references

Copper (II) sulfate, also known as cupric sulfate, or copper sulphate, are inorganic compound with the chemical formula CuSO4(H2O)x, where x can range from 0 to 5. The pentahydrate (x = 5) is the most common form. Older names for this compound include blue vitriol, bluestone,[9] vitriol of copper,[10] and Roman vitriol.[11]

This salt exists as a series of compounds that differ in their degree of hydration. The anhydrous salt is a white powder in its pure form, whereas the pentahydrate (CuSO4·5H2O), the most commonly encountered salt, is bright blue. Copper (II) sulfate exothermically dissolves in water to give the aquo complex [Cu(H2O)6]2+, which has octahedral molecular geometry and is paramagnetic.

Preparation and occurrence

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 its 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.[12]

Commercial copper sulfate is usually about 98% pure copper sulfate, and may contain traces of water. Anhydrous Copper sulfate is 39.81 percent copper and 60.19 percent 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).[12]

Chemical properties

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).[13][14] Dehydration proceeds by decomposition of the tetraaquacopper(2+) moiety, two opposing aqua groups are lost to give a diaquacopper(2+) moiety. The second dehydration step occurs with the final two aqua groups are lost. Complete dehydration occurs when the only unbound water molecule is lost. At 650 °C (1,202 °F), copper (II) sulfate decomposes into copper (II) oxide (CuO) and sulfur trioxide (SO3).

Copper sulfate reacts with concentrated hydrochloric acid to give tetrachlorocuprate(II):

Cu2+ + 4 Cl− → CuCl2−
4

It also reacts with more reducing metals to give copper metal and the corresponding oxidized metal, e.g.

CuSO4 + Zn → ZnSO4 +Cu

Uses

As a fungicide

Copper sulfate pentahydrate is a fungicide.[15] However, some fungi are capable of adapting to elevated levels of copper ions.[16] By mixing a water solution of copper sulfate and a suspension of slaked lime one obtains the Bordeaux mixture, a suspension of copper(II) hydroxide Cu(OH)2 and calcium sulphate, which is used to control fungus on grapes, melons, and other berries.[17]

Another application is Cheshunt compound, a mixture of copper sulfate and ammonium carbonate used in horticulture to prevent damping off in seedlings. Its use as a herbicide is not agricultural, but instead to control invasive aquatic plants and the roots of plants that may be situated near pipes containing water. It is used in swimming pools as an algicide. A dilute solution of copper sulfate is used to treat aquarium fishes for parasitic infections,[18] and is also used to remove snails from aquariums. Copper ions are highly toxic to fish, so care must be taken with the dosage. Most species of algae can be controlled with very low concentrations of copper sulfate. Copper sulfate inhibits growth of bacteria such as Escherichia coli.

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 an additive to concrete to provide water resistance and to make it antiseptic. Copper sulfate can be used as a coloring ingredient in artworks, especially glasses and potteries.[19] 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.[20]

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 tested by dropping it into a solution of copper sulfate of known specific gravity – blood which contains sufficient hemoglobin sinks rapidly due to its density, whereas blood which does not sink or sinks slowly has insufficient amount of hemoglobin.[21]

In a flame test, its copper ions 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.[22] The anhydrous salt is used as a dehydrating agent for forming and manipulating acetal groups.[23] The hydrated salt can be intimately mingled with potassium permanganate to give an oxidant for the conversion of primary alcohols.[24]

Chemistry education

Copper sulfate is commonly included in children's chemistry sets. It is often used to grow crystals in schools and in copper plating 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, and is known as blue vitriol.[25] Copper(II) sulfate pentahydrate can easily be produced by crystallization from solution as copper(II) sulfate is quite hygroscopic.

In an illustration of a "single metal replacement reaction", iron is submerged in a solution of copper sulfate. Upon standing, iron reacts, producing iron(II) sulfate, and copper precipitates.

Fe + CuSO4 → FeSO4 + Cu

In high school and general chemistry education, copper sulfate is used as 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.[26]

Cu2+ + 2e− → Cu (cathode) E°cell=0.34V
Lowering a copper etching plate into the copper sulfate solution.

Medical and public health

Copper sulfate was used in the past as an emetic.[27] It is now considered too toxic for this use.[28] It is still listed as an antidote in the World Health Organization's Anatomical Therapeutic Chemical Classification System.[29] Copper sulfate was once used to fight malaria. For example, during the 1940s in Trinidad, a malaria epidemic was caused by an increase of mosquito habitat in bromeliads growing on newly imported immortelle (Erythrina micropteryx) trees. The epidemic was controlled by spraying dilute copper sulfate solution into these epiphytes, killing them and removing the mosquito breeding grounds.[30] Copper sulfate is used as a molluscicide to treat bilharzia in tropical countries.[19] Cupric sulfate is also used to assist with the treatment of cutaneous phosphorus burns; however, it is not recommended for this purpose due to its toxicity.[31]

Art

In 2008, the artist Roger Hiorns filled an abandoned waterproofed council flat in London with 75,000 liters of copper sulfate 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.[32] Since 2011, it has been on exhibition at the Yorkshire Sculpture Park.[33]

Etching

Copper sulfate is used to etch zinc or copper plates for intaglio printmaking.[34][35] It is also used to etch designs into copper for jewelry, such as for Champlevé.[36]

Dyeing

Copper sulfate can be used as a mordant in vegetable dyeing. It often highlights the green tints of the specific dyes.

Anhydrous and rare mineral forms

The anhydrous form is a white solid. It can be produced by dehydration of the hydrate. It occurs as a rare mineral known as chalcocyanite. The hydrated copper sulfate also occurs in nature as chalcanthite (pentahydrate). Two other copper sulfates comprise the remaining of these rare minerals: bonattite (trihydrate) and boothite (heptahydrate).

Structure of anhydrous CuSO4.
Space-filling model anhydrous CuSO4.
The rare mineral boothite (CuSO4·7H2O)

Toxicological effects

Copper sulfate is an irritant.[37] The usual routes by which humans can receive toxic exposure to copper sulfate are through eye or skin contact, as well as by inhaling powders and dusts.[38] Skin contact may result in itching or eczema.[39] Eye contact with copper sulfate can cause conjunctivitis, inflammation of the eyelid lining, ulceration, and clouding of the cornea.[40]

Upon oral exposure, copper sulfate is moderately toxic.[38] According to studies, the lowest dose of copper sulfate that had a toxic impact on humans is 11 mg/kg.[41] Because of its irritating effect on the gastrointestinal tract, vomiting is automatically triggered in case of the ingestion of copper sulfate. However, if copper sulfate is retained in the stomach, the symptoms can be severe. After 1–12 grams of copper sulfate are swallowed, such poisoning signs may occur as a metallic taste in the mouth, burning pain in the chest, nausea, diarrhea, vomiting, headache, discontinued urination, which leads to yellowing of the skin. In cases of copper sulfate poisoning, injury to the brain, stomach, liver, or kidneys may also occur.[40]

Environmental toxicity

Copper sulfate is highly soluble in water and therefore is easy to distribute in the environment. Copper in the soil may be from industry, motor vehicle, and architectural materials.[42] According to studies, copper sulfate exists mainly in the surface soil and tends to bind organic matter. The more acidic the soil is, the less binding occurs.

References

  1. 1 2 3 4 5 6 7 8 Haynes, p. 4.62
  2. ↑ Haynes, p. 5.199
  3. ↑ Anthony, John W.; Bideaux, Richard A.; Bladh, Kenneth W.; Nichols, Monte C., eds. (2003). "Chalcocyanite". Handbook of Mineralogy (PDF). V. Borates, Carbonates, Sulfates. Chantilly, VA, US: Mineralogical Society of America. ISBN 0962209740.
  4. ↑ Haynes, p. 10.240
  5. ↑ 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.
  6. ↑ Bacon, G. E.; Titterton, D. H. (1975). "Neutron-diffraction studies of CuSO4 · 5H2O and CuSO4 · 5D2O". Z. Kristallogr. 141 (5–6): 330–341. doi:10.1524/zkri.1975.141.5-6.330.
  7. ↑ Cupric sulfate. US National Institutes of Health
  8. 1 2 3 "NIOSH Pocket Guide to Chemical Hazards #0150". National Institute for Occupational Safety and Health (NIOSH).
  9. ↑ "Copper (II) sulfate MSDS". Oxford University. Retrieved 2007-12-31.
  10. ↑ 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.
  11. ↑ Oxford University Press, "Roman vitriol", Oxford Living Dictionaries. Accessed on 2016-11-13
  12. 1 2 "Uses of Copper Compounds: Copper Sulphate". copper.org. Copper Development Association Inc. Retrieved 10 May 2015.
  13. ↑ Andrew Knox Galwey; Michael E. Brown (1999). Thermal decomposition of ionic solids. Elsevier. pp. 228–229. ISBN 0-444-82437-5.
  14. ↑ Wiberg, Egon; Nils Wiberg; Arnold Frederick Holleman (2001). Inorganic chemistry. Academic Press. p. 1263. ISBN 0-12-352651-5.
  15. ↑ Johnson, George Fiske (1935). "The Early History of Copper Fungicides". Agricultural History. 9 (2): 67–79. JSTOR 3739659.
  16. ↑ Parry, K. E.; Wood, R. K. S. (1958). "The adaption of fungi to fungicides: Adaption to copper and mercury salts". Annals of Applied Biology. 46 (3): 446. doi:10.1111/j.1744-7348.1958.tb02225.x.
  17. ↑ "Uses of Copper Compounds: Copper Sulfate's Role in Agriculture". Copper.org. doi:10.1111/j.1744-7348.1933.tb07770.x. Retrieved 2007-12-31.
  18. ↑ "All About Copper Sulfate". National Fish Pharmaceuticals. Retrieved 2007-12-31.
  19. 1 2 Copper Development Association. "Uses of Copper Compounds: Table A - Uses of Copper Sulphate". copper. Copper Development Association Inc. Retrieved 12 May 2015.
  20. ↑ Partin, Lee. "The Blues: Part 2". skylighter. Skylighter.Inc. Retrieved 12 May 2015.
  21. ↑ Estridge, Barbara H.; Anna P. Reynolds; Norma J. Walters (2000). Basic Medical Laboratory Techniques. Thomson Delmar Learning. p. 166. ISBN 0-7668-1206-5.
  22. ↑ Hoffman, R. V. (2001). Copper(II) Sulfate, in Encyclopedia of Reagents for Organic Synthesis. John Wiley & Sons. doi:10.1002/047084289X.rc247.
  23. ↑ Philip J. Kocienski (2005). Protecting Groups. Thieme. p. 58. ISBN 978-1-58890-376-1.
  24. ↑ 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". Org. Synth.; Coll. Vol., 9, p. 462
  25. ↑ "Process for the preparation of stable copper (II) sulfate monohydrate applicable as trace element additive in animal fodders". Retrieved 2009-07-07.
  26. ↑ Zumdahl, Steven; DeCoste, Donald (2013). Chemical Principles. Cengage Learning. pp. 506–507. ISBN 978-1-285-13370-6.
  27. ↑ Holtzmann, N. A.; Haslam, R. H. (July 1968). "Elevation of serum copper following copper sulfate as an emetic". Pediatrics. 42 (1): 189–93. PMID 4385403.
  28. ↑ Olson, Kent C. (2004). Poisoning & drug overdose. New York: Lange Medical Mooks/McGraw-Hill. p. 175. ISBN 0-8385-8172-2.
  29. ↑ V03AB20 (WHO)
  30. ↑ 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.
  31. ↑ Barqouni, Loai; Abu Shaaban, Nafiz; Elessi, Khamis; Barqouni, Loai (2014). "Interventions for treating phosphorus burns". Cochrane Database of Systematic Reviews (6): CD008805. doi:10.1002/14651858.CD008805.pub3.
  32. ↑ "Seizure homepage". Artangel.org.uk. Retrieved 2009-09-21.
  33. ↑ "Roger Hiorns: Seizure". Yorkshire Sculpture Park. Retrieved 2015-02-22.
  34. ↑ greenart.info, Bordeau etch, 2009-01-18, retrieved 2011-06-02.
  35. ↑ ndiprintmaking.ca, The Chemistry of using Copper Sulfate Mordant, 2009-04-12, retrieved 2011-06-02.
  36. ↑ http://mordent.com/etch-howto/, How to Electrolytically etch in copper, brass, steel, nickel silver or silver, retrieved 2015-05-2015.
  37. ↑ Windholz, M., ed. 1983. The Merck Index. Tenth edition. Rahway, NJ: Merck and Company.
  38. 1 2 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
  39. ↑ TOXNET. 1975–1986. National library of medicine's toxicology data network. Hazardous Substances Data Bank (HSDB). Public Health Service. National Institute of Health, U.S. Department of Health and Human Services. Bethesda, MD: NLM.
  40. 1 2 Clayton, G. D. and F. E. Clayton, eds. 1981. Patty's industrial hygiene and toxicology. Third edition. Vol. 2, Part 6 Toxicology. NY: John Wiley and Sons. ISBN 0-471-01280-7
  41. ↑ National Institute for Occupational Safety and Health (NIOSH). 1981–1986. Registry of toxic effects of chemical substances (RTECS). Cincinnati, OH: NIOSH.
  42. ↑ United State Environment Protection Agency. "Reregistration Eligibility Decision (RED) for Coppers" (PDF). www.epa.org. United State Environment Protection Agency. Retrieved 12 May 2015.

Bibliography

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