Copper(II) sulfate is aninorganic compound with thechemical formulaCuSO4. It formshydratesCuSO4·nH2O, wheren 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 itsanhydrous form is white.[11] Older names for the pentahydrate includeblue vitriol,bluestone,[12]vitriol of copper,[13] andRoman vitriol.[14] Itexothermically dissolves in water to give theaquo complex[Cu(H2O)6]2+, which hasoctahedral molecular geometry. The structure of the solid pentahydrate reveals a polymeric structure wherein copper is again octahedral but bound to four water ligands. TheCu(II)(H2O)4 centers are interconnected by sulfate anions to form chains.[15]
Preparation of copper(II) sulfate by electrolyzing sulfuric acid, using copper electrodes
Copper sulfate is produced industrially by treating copper metal with hot concentratedsulfuric acid or copper oxides with dilute sulfuric acid. For laboratory use, copper sulfate is usually purchased. Copper sulfate can also be produced by slowlyleaching low-gradecopper 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 ofcrystal 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]
Copper(II) sulfate pentahydratedecomposes 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 copperaquo complex, since the sulfate is not bound to copper in such solutions. Thus, such solutions react with concentratedhydrochloric acid to givetetrachlorocuprate(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 suchsingle 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 forgalvanic 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 teenagechemistry sets and undergraduate experiments.[21] It is often used to grow crystals inschools and inCopper electroplating experiments despite its toxicity. Copper sulfate is often used to demonstrate anexothermic reaction, in whichsteel wool ormagnesium ribbon is placed in anaqueous solution ofCuSO4. It is used to demonstrate the principle ofmineral hydration. Thepentahydrate 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.[citation needed] Copper(II) sulfate pentahydrate can easily be produced by crystallization from solution as copper(II) sulfate, which ishygroscopic.
Copper sulfate has been used for control ofalgae in lakes and related fresh waters subject toeutrophication. It "remains the most effective algicidal treatment".[22][23]
A dilute solution of copper sulfate is used to treataquarium fishes for parasitic infections,[25] and is also used to remove snails from aquariums andzebra mussels from water pipes.[26] Copper ions are highly toxic to fish. Most species of algae can be controlled with very low concentrations of copper sulfate.
Copper sulfate is used to test blood foranemia. The blood is dropped into a solution of copper sulfate of knownspecific gravity—blood with sufficienthemoglobin sinks rapidly due to its density, whereas blood which sinks slowly or not at all has an insufficient amount of hemoglobin.[27] Clinically relevant, however, modern laboratories utilize automated blood analyzers for accurate quantitative hemoglobin determinations, as opposed to older qualitative means.[citation needed]
In aflame test, the copperions of copper sulfate emit a deep green light, a much deeper green than the flame test forbarium.
Copper sulfate is employed at a limited level inorganic synthesis.[28] The anhydrous salt is used as a dehydrating agent for forming and manipulatingacetal groups.[29] The hydrated salt can be intimately mingled withpotassium permanganate to give an oxidant for the conversion of primary alcohols.[30]
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.[31] Copper sulfate is also rarely used in firework manufacture as a blue coloring agent, but it is not safe to mix copper(II) sulfate with metal powders, or it or any copper(II) compound with chlorates; the sulfate and other copper(II) compounds are not allowed in chlorate containing mixtures in the US.[32][33]
Lowering a copper etching plate into the copper sulfate solution
Copper sulfate was once used to killbromeliads, which serve as mosquito breeding sites.[34] Copper sulfate is used as a molluscicide to treatbilharzia in tropical countries.[31]
In 2008, the artistRoger Hiorns filled an abandoned waterproofedcouncil 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 titledSeizure.[35] Since 2011, it has been on exhibition at theYorkshire Sculpture Park.[36]
Copper(II) sulfate is used to etch zinc, aluminium, or copper plates forintaglio printmaking.[37][38]It is also used to etch designs into copper for jewelry, such as forChamplevé.[39]
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 aschalcocyanite.[41] The pentahydrate also occurs in nature aschalcanthite. Other rare copper sulfate minerals includebonattite (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]
^Rumble, John, ed. (2018).CRC Handbook of Chemistry and Physics (99th ed.). CRC Press, Taylor & Francis Group. pp. 5–179.ISBN9781138561632.
^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.ISBN978-0962209741.
^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.
^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.
^Oxford University Press, "Roman vitriol", Oxford Living Dictionaries. Accessed on 2016-11-13
^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.PMID20024268.
^Zumdahl, Steven; DeCoste, Donald (2013).Chemical Principles. Cengage Learning. pp. 506–507.ISBN978-1-285-13370-6.
^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.
^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.
^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.
^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.
^Estridge, Barbara H.; Anna P. Reynolds; Norma J. Walters (2000).Basic Medical Laboratory Techniques. Thomson Delmar Learning. p. 166.ISBN978-0-7668-1206-2.
^Partin, Lee."The Blues: Part 2".skylighter. Skylighter.Inc. Retrieved12 May 2015.
^"Approved and Prohibited Fireworks Chemicals, v2"(PDF).phmsa.dot.gov. US Department of Transportation - Pipeline and Hazardous Materials Safety Administration. 5 November 2019. p. 1. Retrieved3 August 2025.Copper Sulfate...Color Agent...Prohibited if mixed with a chlorate
^Despommier; Gwadz; Hotez; Knirsch (June 2005).Parasitic Disease (5 ed.). NY: Apple Tree Production L.L.C. pp. Section 4.2.ISBN978-0970002778. Retrieved12 May 2015.
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