|Systematic name||copper(II) sulfate|
|Other names||cupric sulfate, copper sulfate,
chalcanthite, blue vitriol, bluestone
|Molar mass||143.61 g/mol|
|Appearance||blue solid crystals when hydrated,
white solid when anhydrous
|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,
|Main hazards||(Xn) Harmful
(N) Dangerous for the environment
|R/S statement||R: R22, R36/38, R50/53
S: S2, S22, S60, S61
|Other anions||Copper(II) chloride, Copper(II) oxide|
|Other cations||Sodium sulfate, Manganese sulfate,
|Except where noted otherwise, data are given for
materials in their standard state (at 25 °C, 100 kPa)
Copper(II) sulfate, also known as cupric sulfate, copper sulfate, blue vitriol, or bluestone, 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.
Properties[change | change source]
Physical properties[change | change source]
Copper(II) sulfate is a blue solid when hydrated (attached to water molecules). It is whitish when anhydrous (not attached to water molecules). When hydrated, it normally has five water molecules attached to it. It can be dehydrated by heating it. 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[change | change source]
- Fe + CuSO4 → FeSO4 + Cu
- CuSO4 + 2 NaOH → Cu(OH)2 + Na2SO4
- CuSO4 + Na2CO3 → CuCO3 + Na2SO4
- CuSO4 + 4 NH3 → Cu(NH3)4SO4
- CuSO4 → CuO + SO3
It makes a blue-green color when it is heated in a flame, like all copper compounds.
Occurrence[change | change source]
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. 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[change | change source]
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.
- Cu + H2SO4 → CuSO4 + H2
- CuO + H2SO4 → H2O + CuSO4
- Cu(OH)2 + H2SO4 → 2 H2O + CuSO4
- CuCO3 + H2SO4 → H2O + CuSO4 + CO2
- Cu + 2H2SO4 → CuSO4 + H2O + SO2
It can also be made by reacting copper with a mixture of nitric acid and sulfuric acid.
Uses[change | change source]
Copper(II) sulfate, as the most common copper compound, has many uses. It can be used to kill algae and fungi. Some fungi can get resistant to copper sulfate, though. Then the copper sulfate does not kill them any more. It can be mixed with lime to make a similar fungi killer. It can be used to treat aquarium fish for infections. 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 as a catalyst and oxidizing agent. It is used to see whether blood is anemic.
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. It is seen as too toxic now.
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.
Safety[change | change source]
Copper sulfate is somewhat toxic to humans. It is very toxic to fish, though. In humans, it irritates skin and eyes. 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[change | change source]
References[change | change source]
- "Copper(II) sulfate MSDS". Oxford University. http://ptcl.chem.ox.ac.uk/MSDS/CO/copper_II_sulfate.html. Retrieved 2007-12-31.
- Holleman, A. F.; Wiberg, E. (2001). Inorganic Chemistry. San Diego: Academic Press. ISBN 0-12-352651-5.
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- National Audubon Society, Field Guide to Rocks and Minerals, Alfred A. Knopf (publisher) (c) 1979, pg. 461
- Johnson, George Fiske (1935). "The Early History of Copper Fungicides". Agricultural History 9 (2): 67–79. https://www.jstor.org/pss/3739659.
- 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: 446. doi:10.1111/j.1744-7348.1958.tb02225.x.
- "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.
- "All About Copper Sulfate". National Fish Pharmaceuticals. http://www.fishyfarmacy.com/Q&A/all_about_copper.html. Retrieved 2007-12-31.
- Hoffman, R. V. (2001). Copper(II) Sulfate, in Encyclopedia of Reagents for Organic Synthesis. John Wiley & Sons. doi:10.1002/047084289X.rc247.
- Barbara H. Estridge, Anna P. Reynolds, Norma J. Walters (2000). Basic Medical Laboratory Techniques. Thomson Delmar Learning. p. 166. ISBN 0766812065.
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- "Seizure homepage". Artangel.org.uk. http://www.artangel.org.uk/projects/2008/seizure. Retrieved 2009-09-21.
- 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.
- Windholz, M., ed. 1983. The Merck Index. Tenth edition. Rahway, NJ: Merck and Company.
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