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Copper(II) sulfate

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Copper(II) sulfate
Copper(II) sulfate
Difference between anhydrous and hydrated copper sulfate
General
Systematic name copper(II) sulfate
Other names cupric sulfate, copper sulfate,
chalcanthite, blue vitriol, bluestone
Molecular formula CuSO4
Molar mass 159.60 g/mol (anhydrous)
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 Archived 2022-04-24 at the Wayback Machine
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, and practical lessons at high school.

Properties

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Physical properties

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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

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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 to a high temperature, 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]

Copper flame test

Occurrence

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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. 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

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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 or by adding copper to heated concentrated sulphuric acid:

CuO + H2SO4 → H2O + CuSO4
Cu(OH)2 + H2SO4 → 2 H2O + CuSO4
CuCO3 + H2SO4 → H2O + CuSO4 + CO2
Cu + 2H2SO4 → CuSO4 + 2H2O + SO2

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

Uses

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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

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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.

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References

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  1. ↑ 1.0 1.1 "Copper(II) sulfate MSDS". Oxford University. Archived from the original on 2007-10-11. Retrieved 2007-12-31.
  2. ↑ Holleman, A. F.; Wiberg, E. (2001). Inorganic Chemistry. San Diego: Academic Press. ISBN 0-12-352651-5.{{cite book}}: CS1 maint: multiple names: authors list (link)
  3. ↑ Andrew Knox Galwey, Michael E. Brown (1999). Thermal decomposition of ionic solids. Elsevier. pp. 228–229. ISBN 0444824375.
  4. ↑ Egon Wiberg, Nils Wiberg, Arnold Frederick Holleman (2001). Inorganic chemistry. Academic Press. p. 1263. ISBN 0123526515.{{cite book}}: CS1 maint: multiple names: authors list (link)
  5. ↑ Greenwood, Norman N.; Earnshaw, A. (1984), Chemistry of the Elements, Oxford: Pergamon, p. 451, ISBN 0-08-022057-6{{citation}}: CS1 maint: multiple names: authors list (link)
  6. ↑ "Another copper reaction". Arizona State University. Retrieved 2010-06-11.
  7. ↑ "Reaction video". Journal of Chemical Education. Archived from the original on 2010-06-10. Retrieved 2010-06-11.
  8. ↑ 8.0 8.1 "Copper". The University of North Carolina at Pembroke. Archived from the original on 2008-09-29. Retrieved 2010-06-11.
  9. ↑ "Decomposition". Cornell University. 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. JSTOR 3739659.
  12. ↑ Parry, K. E.; Wood, R. K. S. (1958). "The Adaptation of Fungi to Fungicides: Adaptation To Copper and Mercury Salts". Annals of Applied Biology. 46 (3): 446–456. doi:10.1111/j.1744-7348.1958.tb02225.x.
  13. ↑ "Uses of Copper Compounds: Copper Sulfate's Role in Agriculture". Copper.org. Archived from the original on 2013-09-07. Retrieved 2007-12-31.
  14. ↑ "All About Copper Sulfate". National Fish Pharmaceuticals. Archived from the original on 2014-08-10. 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.{{cite book}}: CS1 maint: multiple names: authors list (link)
  17. ↑ Holtzmann NA, Haslam RH (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. Archived from the original on 2010-06-16. Retrieved 2010-10-17.
  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. Archived from the original on 2009-07-31. 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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