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Copper(II) sulfate
Identifiers
CAS number 7758-98-7 Yes check.svgY
7758-99-8 (pentahydrate)
PubChem 24462
ChemSpider 22870
EC number 231-847-6
RTECS number GL8800000 (anhydrous)
GL8900000 (pentahydrate)
SMILES
InChI
InChI key ARUVKPQLZAKDPS-NUQVWONBAI
Properties
Molecular formula CuSO4
Molar mass 159.61 g/mol (anhydrous)
249.68 g/mol (pentahydrate)
Appearance blue crystalline solid (pentahydrate)
gray-white powder (anhydrous)
Density 3.603 g/cm3 (anhydrous)
2.284 g/cm3 (pentahydrate)
Melting point

110 °C (−4H2O)
150 °C (423 K) (−5H2O)
< 650 °C decomp.

Solubility in water 31.6 g/100 ml (0 °C)
32 g/100 mL (20°C)
61.8 g/100 mL (60°C)
114 g/100 mL (100°C)
Solubility anhydrous
insoluble in ethanol
pentahydrate
soluble in methanol and ethanol
Refractive index (nD) 1.514 (pentahydrate)
Structure
Crystal structure Triclinic
Coordination
geometry
Octahedral
Thermochemistry
Standard molar
entropy
So298
109.05 J K−1 mol−1
Hazards
MSDS ICSC 0751 (anhydrous)
ICSC 1416 (pentahydrate)
EU Index 029-004-00-0
EU classification Harmful (Xn)
Irritant (Xi)
Dangerous for the environment (N)
R-phrases R22, R36/38, R50/53
S-phrases (S2), S22, S60, S61
NFPA 704
NFPA 704.svg
0
2
1
Flash point Non-inflammable
LD50 300 mg/kg
Related compounds
Other cations Nickel(II) sulfate
Zinc sulfate
 Yes check.svgY (what is this?)  (verify)
Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa)
Infobox references

Copper(II) sulfate is the chemical compound with the formula CuSO4. This salt exists as a series of compounds that differ in their degree of hydration. The anhydrous form is a pale green or gray-white powder, whereas the pentahydrate (CuSO4·5H2O), the most commonly encountered salt, is bright blue. The anhydrous form occurs as a rare mineral known as chalcocyanite. The hydrated copper sulfate occurs in nature as chalcanthite (pentahydrate), and two more rare ones: bonattite (trihydrate) and boothite (heptahydrate). Archaic names for copper(II) sulfate are "blue vitriol" and "bluestone".[1]

Contents

Preparation

Preparation of copper(II) sulfate by electrolyzing sulfuric acid, using copper electrodes

Since it is available commercially, copper sulfate is usually purchased and not prepared in the laboratory. It can be made by the action of sulfuric acid on a variety of copper(II) compounds, for example copper(II) oxide; this oxide can be generated with the addition of hydrogen peroxide to the acid. It may also be prepared by electrolyzing sulfuric acid, using copper electrodes. It can also be prepared by electrolysis of magnesium sulfate [Epsom salts] solution at moderate voltage with a copper anode: this reaction produces hydrogen, copper sulfate solution, and copper hydroxide precipitate.

Chemical properties

Copper(II) sulfate pentahydrate decomposes before melting, losing two water molecules at 63°C, followed by two more at 109°C and the final water molecule at 220°C.[citation needed] At 650 °C, copper(II) sulfate decomposes into copper(II) oxide (CuO) and sulfur trioxide (SO3). Its blue color is due to water of hydration. When heated in an open flame the crystals are dehydrated and turn grayish-white.[2]

Uses

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As a herbicide, fungicide and pesticide

Copper sulfate pentahydrate is a fungicide. Mixed with lime it is called Bordeaux mixture and used to control fungus on grapes, melons, and other berries.[3] 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 for control of invasive aquatic plants and the roots of plants near pipes containing water. It is used in swimming pools as an algaecide. A dilute solution of copper sulfate is used to treat aquarium fish for parasitic infections,[4] and is also used to remove snails from aquariums. Copper ions are highly toxic to fish, 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 E. coli.

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 also 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, floats or sinks less rapidly.[5]

In a flame test, its copper ions emit a deep blue-green light, much more blue than the flame test for barium.

Organic synthesis

Copper sulfate is employed in organic synthesis.[6] The anhydrous salt catalyses the transacetylation in organic synthesis.[7] The hydrated salt reacts with potassium permanganate to give an oxidant for the conversion of primary alcohols.[8]

Chemistry education

Copper sulfate is a commonly included chemical in children's chemistry sets and is often used to grow crystals in schools and in copper plating experiments. Due to its toxicity, it is not recommended for small children. 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 in school chemistry courses 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 copperas.[9]

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

Fe + CuSO4 → FeSO4 + Cu

Other uses

Lowering a zinc etching plate into the copper sulfate solution.

Medical

Copper sulfate was also used in the past as an emetic.[10] It is now considered too toxic for this use.[11] It is still listed as an antidote in the World Health Organization's ATC code V03.[12]

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 crystallise for several weeks, and the flat was drained, leaving crystal-covered walls, floors and ceilings.

The work is titled Seizure.[13]

Etching

Copper sulfate is also used to etch zinc plates for intaglio printmaking.[14][15]

Toxicological Effects

Copper sulfate is a strong irritant [16]. 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 [17]. Skin contact may result in itching or eczema [18]. Eye contact with copper sulfate can cause conjunctivitis, inflammation of the eyelid lining, ulceration, and clouding of the cornea [19] Upon acute oral exposure, copper sulfate turns to be only moderately toxic [20]. According to studies, the lowest dose of copper sulfate that had a toxic impact on humans is 11 mg/kg.[21]. 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 case of copper sulfate poisoning, injury to the brain, stomach, liver, kidneys may also occur [19]

References

  1. ^ "Copper(II) sulfate MSDS". Oxford University. http://ptcl.chem.ox.ac.uk/MSDS/CO/copper_II_sulfate.html. Retrieved 2007-12-31. 
  2. ^ Holleman, A. F.; Wiberg, E.. Inorganic Chemistry. San Diego&year= 2001&isbn= 0-12-352651-5: Academic Press. 
  3. ^ "Uses of Copper Compounds: Copper Sulfate's Role in Agriculture". Copper.org. http://www.copper.org/applications/compounds/copper_sulfate02.html. Retrieved 2007-12-31. 
  4. ^ "All About Copper Sulfate". National Fish Pharmaceuticals. http://www.fishyfarmacy.com/Q&A/all_about_copper.html. Retrieved 2007-12-31. 
  5. ^ Barbara H. Estridge, Anna P. Reynolds, Norma J. Walters (2000). Basic Medical Laboratory Techniques. Thomson Delmar Learning. p. 166. ISBN 0766812065. 
  6. ^ Hoffman, R. V. (2001). Copper(II) Sulfate, in Encyclopedia of Reagents for Organic Synthesis. John Wiley & Sons. doi:10.1002/047084289X.rc247. 
  7. ^ Hulce, M. Mallomo, J. P.; Frye, L. L.; Kogan, T. P.; Posner, G. H. (1990), "(S)-( + )-2-(p-Toluenesulfinyl)-2-Cyclopentanone: Precursor for Enantioselective Synthesis of 3-Substituted Cyclopentanones", Org. Synth., http://www.orgsyn.org/orgsyn/orgsyn/prepContent.asp?prep=CV7P0495 ; Coll. Vol. 7: 495 
  8. ^ 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., http://www.orgsyn.org/orgsyn/orgsyn/prepContent.asp?prep=cv9p0462 ; Coll. Vol. 9: 462 
  9. ^ "Process for the preparation of stable copper (II) sulfate monohydrate applicable as trace element additive in animal fodders". http://www.freepatentsonline.com/4315915.html. Retrieved 2009-07-07. 
  10. ^ Holtzmann NA, Haslam RH (July 1968). "Elevation of serum copper following copper sulfate as an emetic". Pediatrics 42 (1): 189–93. PMID 4385403. http://pediatrics.aappublications.org/cgi/content/abstract/42/1/189. 
  11. ^ Olson, Kent C. (2004). Poisoning & drug overdose. New York: Lange Medical Mooks/McGraw-Hill. p. 175. ISBN 0-8385-8172-2. 
  12. ^ V03AB20
  13. ^ "Seizure homepage". Artangel.org.uk. http://www.artangel.org.uk/projects/2008/seizure. Retrieved 2009-09-21. 
  14. ^ Bordeau etch
  15. ^ The Chemistry of using Copper Sulfate Mordant
  16. ^ Windholz, M., ed. 1983. The Merck Index. Tenth edition. Rahway, NJ: Merck and Company.
  17. ^ U. S. Environmental Protection Agency. 1986 Guidance for reregistration of pesticide products containing copper sulfate. Fact sheet no 100. Office of Pesticide Programs. Washington, DC.
  18. ^ 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.
  19. ^ a b Clayton, G. D. and F. E. Clayton, eds. 1981. Patty's industrial hygiene and toxicology. Third edition. Vol. 2: Toxicology. NY: John Wiley and Sons.
  20. ^ 1986. Guidance for reregistration of pesticide products containing copper sulfate. Fact sheet no 100. Office of Pesticide Programs. Washington, DC.
  21. ^ National Institute for Occupational Safety and Health (NIOSH). 1981- 1986. Registry of toxic effects of chemical substances (RTECS). Cincinati, OH: NIOSH.

External links


Simple English

Copper(II) sulfate
File:Copper
File:Hydrating-copper(II)
General
Systematic name copper(II) sulfate
Other names cupric sulfate, copper sulfate, chalcanthite, blue vitriol, bluestone
Molecular formula CuSO4
Molar mass 143.61 g/mol
Appearance blue solid crystals when hydrated, white solid when anhydrous
CAS number 7758-98-7
Properties
Density and phase 3.603 g/cm³ (anhydrous), 2.284 g/cm³ (hydrated)
Solubility in water 31.6 g/100 ml (0°C)
Solubility in ethanol insoluble, both forms
Solubility in methanol hydrate is soluble
Melting point 150°C (423 K) dehydrates, 650°C decomp.
Structure
Coordination
geometry
octahedral
Crystal structure triclinic
Hazards
MSDS MSDS
Main hazards (Xn) Harmful
(Xi) Irritant
(N) Dangerous for the environment
NFPA 704

0
2
1
R/S statement R: R22, R36/38, R50/53
S: S2, S22, S60, S61
Related compounds
Other anions Copper(II) chloride, Copper(II) oxide
Other cations Sodium sulfate, Manganese sulfate, Iron(II) sulfate
Except where noted otherwise, data are given for
materials in their standard state (at 25 °C, 100 kPa)
Disclaimer

Copper(II) sulfate, also known as cupric sulfate, copper sulfate, blue vitriol,[1] or bluestone,[1] is a chemical compound. Its chemical formula is CuSO4. It contains copper in its +2 oxidation state. It also contains sulfate ions. It is a blue solid that can kill fungi. It is also used to purify copper metal. It is common in chemistry sets and chemistry demonstrations.

Contents

Properties

Physical properties

Copper(II) sulfate is a blue solid when hydrated (attached to water molecules). It is whitish when anhydrous (not attached to water molecules).[2] When hydrated, it normally has five water molecules attached to it. It can be dehydrated by heating it.[3][4] When water is added to it, it gets hydrated again. When it is in air, it absorbs water vapor and becomes hydrated, too.

Chemical properties

It is a weak oxidizing agent. It reacts with most metals to make copper and a metal sulfate.[5] For example, it reacts with iron to make copper and iron(II) sulfate.

Fe + CuSO4 → FeSO4 + Cu

It reacts with sodium hydroxide or potassium hydroxide to make copper(II) hydroxide.[6]

CuSO4 + 2 NaOH → Cu(OH)2 + Na2SO4

It reacts with sodium carbonate to make copper(II) carbonate.[7]

CuSO4 + Na2CO3 → CuCO3 + Na2SO4

It reacts with ammonia to make a dark blue solution.[8] This solution can dissolve fibers in cotton.

CuSO4 + 4 NH3 → Cu(NH3)4SO4

When it is heated very hot, it turns into copper(II) oxide and sulfur trioxide.[9]

CuSO4 → CuO + SO3

It makes a blue-green color when it is heated in a flame, like all copper compounds.[8]

File:Flametest--Cu.
Copper flame test

Occurrence

Copper(II) sulfate is found in the ground as chalcanthite. Chalcanthite is easily dissolved. It is only found in dry areas. When it is in air, it loses its bright blue color. It is found in very dry areas. Some minerals are tested by taste. Chalcanthite has a sweet metal taste. It should only be tasted carefully, as it is poisonous.[10] Its Mohs hardness is 2.5. It is the pentahydrate of copper sulfate. It is blue or green. Many people collecting minerals want it.

Preparation

File:Synthesizing Copper
Making copper sulfate by electrolysis of sulfuric acid with copper electrodes

Copper sulfate is not normally made in a small laboratory, because it is much easier just to buy it. There are some ways to make copper sulfate, however.

Copper(II) sulfate can be made by electrolysis of a solution of sulfuric acid with copper electrodes. Hydrogen is made, as well as copper sulfate solution.

Cu + H2SO4 → CuSO4 + H2

It can also be made by reacting copper(II) oxide or copper(II) hydroxide or copper(II) carbonate with sulfuric acid.

CuO + H2SO4 → H2O + CuSO4
Cu(OH)2 + H2SO4 → 2 H2O + CuSO4
CuCO3 + H2SO4 → H2O + CuSO4 + CO2

It can also be made by reacting copper with a mixture of nitric acid and sulfuric acid.

Uses

Copper(II) sulfate, as the most common copper compound, has many uses. It can be used to kill algae and fungi.[11] Some fungi can get resistant to copper sulfate, though. Then the copper sulfate does not kill them any more.[12] It can be mixed with lime to make a similar fungi killer.[13] It can be used to treat aquarium fish for infections.[14] It is also used to detect sugars. It turns into red copper(I) oxide when reduced by a sugar. It can be used in organic chemistry[15] as a catalyst and oxidizing agent. It is used to see whether blood is anemic.[16]

It is commonly found in chemistry sets. It is used to demonstrate a displacement reaction, where a metal reacts with copper sulfate to make copper and the metal sulfate. It is also used to demonstrated hydrated and anhydrous chemicals. It was used as an emetic in the past.[17] It is seen as too toxic now.[18]

It can be used to purify copper. A thin pure piece of copper and a thick impure piece of copper are placed in copper sulfate solution. The thin plate is connected to the negative wire and the thick plate to the positive wire. An electrical current is passed through them. The copper in the thick plate dissolves and plates on the thin plate. All of the impurities fall to the bottom, while the pure copper is made at the negative electrode.

Someone covered the walls of their apartment with copper sulfate crystals for decoration.[19]

Safety

Copper sulfate is somewhat toxic to humans.[20] It is very toxic to fish, though. In humans, it irritates skin and eyes.[21][22][23] It can cause nausea when eaten. It automatically makes one throw up when it is ingested. If too much is eaten, however, it can get into the stomach and cause many problems.

Related pages

References

  1. 1.0 1.1 "Copper(II) sulfate MSDS". Oxford University. http://ptcl.chem.ox.ac.uk/MSDS/CO/copper_II_sulfate.html. Retrieved 2007-12-31. 
  2. Holleman, A. F.; Wiberg, E. (2001). Inorganic Chemistry. San Diego: Academic Press. ISBN 0-12-352651-5. 
  3. Andrew Knox Galwey, Michael E. Brown (1999). Thermal decomposition of ionic solids. Elsevier. pp. 228–229. ISBN 0444824375. http://books.google.com/books?id=i9nyvTYBQtAC&pg=PA229. 
  4. Egon Wiberg, Nils Wiberg, Arnold Frederick Holleman (2001). Inorganic chemistry. Academic Press. p. 1263. ISBN 0123526515. http://books.google.com/books?id=LxhQPdMRfVIC&pg=PA1263. 
  5. Greenwood, Norman N.; Earnshaw, A. (1984), [Expression error: Unexpected < operator Chemistry of the Elements], Oxford: Pergamon, p. 451, ISBN 0-08-022057-6 
  6. "Another copper reaction". Arizona State University. http://www.public.asu.edu/~jpbirk/qual/qualanal/copper.html. Retrieved 2010-06-11. 
  7. "Reaction video". Journal of Chemical Education. http://jchemed.chem.wisc.edu/jcesoft/CCA/CCA4/MAIN/CUSO/PAGE1.HTM. Retrieved 2010-06-11. 
  8. 8.0 8.1 "Copper". The University of North Carolina at Pembroke. http://www.uncp.edu/home/mcclurem/ptable/copper/cu.htm. Retrieved 2010-06-11. 
  9. "Decomposition". Cornell University. http://pmep.cce.cornell.edu/profiles/extoxnet/carbaryl-dicrotophos/copper-sulfate-ext.html. Retrieved 2010-06-11. 
  10. National Audubon Society, Field Guide to Rocks and Minerals, Alfred A. Knopf (publisher) (c) 1979, pg. 461
  11. Johnson, George Fiske (1935). "The Early History of Copper Fungicides". Agricultural History 9 (2): 67–79. http://www.jstor.org/pss/3739659. 
  12. Parry, K. E.; Wood, R. K. S. (1958). [Expression error: Unexpected < operator "The Adaptation of Fungi to Fungicides: Adaptation To Copper and Mercury Salts"]. Annals of Applied Biology 46: 446. doi:10.1111/j.1744-7348.1958.tb02225.x. 
  13. "Uses of Copper Compounds: Copper Sulfate's Role in Agriculture". Copper.org. http://www.copper.org/applications/compounds/copper_sulfate02.html. Retrieved 2007-12-31. 
  14. "All About Copper Sulfate". National Fish Pharmaceuticals. http://www.fishyfarmacy.com/Q&A/all_about_copper.html. Retrieved 2007-12-31. 
  15. Hoffman, R. V. (2001). Copper(II) Sulfate, in Encyclopedia of Reagents for Organic Synthesis. John Wiley & Sons. doi:10.1002/047084289X.rc247. 
  16. Barbara H. Estridge, Anna P. Reynolds, Norma J. Walters (2000). Basic Medical Laboratory Techniques. Thomson Delmar Learning. p. 166. ISBN 0766812065. 
  17. Holtzmann NA, Haslam RH (July 1968). "Elevation of serum copper following copper sulfate as an emetic". Pediatrics 42 (1): 189–93. PMID 4385403. http://pediatrics.aappublications.org/cgi/content/abstract/42/1/189. 
  18. Olson, Kent C. (2004). Poisoning & drug overdose. New York: Lange Medical Mooks/McGraw-Hill. p. 175. ISBN 0-8385-8172-2. 
  19. "Seizure homepage". Artangel.org.uk. http://www.artangel.org.uk/projects/2008/seizure. Retrieved 2009-09-21. 
  20. U. S. Environmental Protection Agency. 1986 Guidance for reregistration of pesticide products containing copper sulfate. Fact sheet no 100. Office of Pesticide Programs. Washington, DC.
  21. Windholz, M., ed. 1983. The Merck Index. Tenth edition. Rahway, NJ: Merck and Company.
  22. 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.
  23. Clayton, G. D. and F. E. Clayton, eds. 1981. Patty's industrial hygiene and toxicology. Third edition. Vol. 2: Toxicology. NY: John Wiley and Sons.

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