View source for Copper(II) sulfate

{{Chembox
| Verifiedfields = changed
| Watchedfields  = changed
| verifiedrevid  = 477003045
| Name           = Copper(II) sulfate
| ImageFile1     = Copper sulfate.jpg
| ImageCaption1  = Crystals of {{chem2|CuSO4*5H2O}}
| ImageSize1     = 80px
| ImageFile2     = Copper(II)-sulfate-pentahydrate-xtal-1985-Cu-coord-3D-bs-17.png
| ImageClass2    = bg-transparent
| ImageCaption2  = {{legend|rgb(256, 128, 0)|[[Copper]], Cu}}{{legend|yellow|[[Sulfur]], S}}{{legend|red|[[Oxygen]], O}}{{legend|white|[[Hydrogen]], H}}<br>Portion of the structure of the pentahydrate<br />(sulfate links {{chem2|Cu(H2O)4(2+)}} centers)
| ImageFile3     = Copper(II)-sulfate-pentahydrate-unit-cell-1985-3D-bs-17.png
| ImageClass3    = bg-transparent
| ImageCaption3  = [[Unit cell]] of the [[crystal structure]] of {{chem2|CuSO4*5H2O}}<br />with [[hydrogen bond]]s in black<ref name = "Varghese&Maslen">{{ cite journal | title = Electron density in non-ideal metal complexes. I. Copper sulphate pentahydrate | first1 = J. N. | last1 = Varghese | first2 = E. N. | last2 = Maslen | journal = [[Acta Crystallographica Section B]] | volume = 41 | year = 1985 | issue = 3 | pages = 184–190 | doi = 10.1107/S0108768185001914 }}</ref>
| IUPACName      = Copper(II) sulfate
| OtherNames     = {{ubl|Cupric sulphate|Blue vitriol (pentahydrate)|Bluestone (pentahydrate)|Bonattite (trihydrate mineral)|Boothite (heptahydrate mineral)|[[Chalcanthite]] (pentahydrate mineral)|Chalcocyanite (mineral)}}
Copper Sulphate pentahydrate
| SystematicName = 
| Section1       = {{Chembox Identifiers
| UNII_Ref = {{fdacite|correct|FDA}}
| UNII = KUW2Q3U1VV
| UNII_Comment = (anhydrous)
| UNII1_Ref = {{fdacite|correct|FDA}}
| UNII1 = LRX7AJ16DT
| UNII1_Comment = (pentahydrate)
| InChI = 1/Cu.H2O4S/c;1-5(2,3)4/h;(H2,1,2,3,4)/q+2;/p-2
| ChEBI_Ref = {{ebicite|correct|EBI}}
| ChEBI = 23414
| SMILES = [O-]S(=O)(=O)[O-].[Cu+2]
| InChIKey = ARUVKPQLZAKDPS-NUQVWONBAI
| ChEMBL_Ref = {{ebicite|correct|EBI}}
| ChEMBL = 604
| StdInChI_Ref = {{stdinchicite|correct|chemspider}}
| StdInChI = 1S/Cu.H2O4S/c;1-5(2,3)4/h;(H2,1,2,3,4)/q+2;/p-2
| StdInChIKey_Ref = {{stdinchicite|correct|chemspider}}
| StdInChIKey = ARUVKPQLZAKDPS-UHFFFAOYSA-L
| CASNo = 7758-98-7 
| CASNo_Ref = {{cascite|correct|CAS}}
| CASNo_Comment = (anhydrous)
| CASNo1_Ref = {{cascite|correct|??}}
| CASNo1 = 7758-99-8
| CASNo1_Comment = (pentahydrate)
| CASNo2_Ref = {{cascite|changed|??}}
| CASNo2 = 16448-28-5
| CASNo2_Comment = (trihydrate)
| CASNo3_Ref = {{cascite|changed|??}}
| CASNo3 = 19086-18-1
| CASNo3_Comment = (heptahydrate)
| Gmelin = 8294
| PubChem = 24462
| EINECS = 231-847-6
| RTECS = GL8800000 (anhydrous)<br/>GL8900000 (pentahydrate)
| ChemSpiderID_Ref = {{chemspidercite|correct|chemspider}}
| ChemSpiderID = 22870
| KEGG_Ref = {{keggcite|correct|kegg}}
| KEGG = C18713
 }}
| Section2       = {{Chembox Properties
| Formula = {{chem2|CuSO4}} (anhydrous)<br>{{chem2|CuSO4*5H2O}} (pentahydrate)
| MolarMass = 159.60 g/mol (anhydrous)<ref name=b92/><br/>249.685 g/mol (pentahydrate)<ref name=b92/>
| Appearance = gray-white (anhydrous)<br>blue (pentahydrate)
| Density = 3.60 g/cm<sup>3</sup> (anhydrous)<ref name=b92/><br/>2.286 g/cm<sup>3</sup> (pentahydrate)<ref name=b92/>
| Solubility = ''pentahydrate'' <br /> 316 g/L (0&nbsp;°C) <br /> 2033 g/L (100&nbsp;°C) <hr>''anhydrous'' <br /> 168 g/L (10 °C) <br> 201 g/L (20&nbsp;°C) <br /> 404 g/L (60&nbsp;°C) <br /> 770 g/L (100&nbsp;°C)<ref name="crc99">{{cite book |editor1-last=Rumble |editor1-first=John |title=CRC Handbook of Chemistry and Physics |date=2018 |publisher=CRC Press, Taylor & Francis Group |isbn=9781138561632 |pages=5–179 |edition=99th |language=en}}</ref>
| SolubleOther = ''anhydrous''<br/>insoluble in [[ethanol]]<ref name=b92/> <hr> ''pentahydrate''<br/>soluble in [[methanol]]<ref name=b92/><br/>10.4 g/L (18&nbsp;°C)<br/>insoluble in [[ethanol]] and [[acetone]] 
| MeltingPtC = 110
| MeltingPt_notes = ''decomposes'' 
<p>560&nbsp;°C ''decomposes''<ref name=b92>[[#CRC92|Haynes]], p. 4.62</ref>(pentahydrate)</p>
<p>Fully decomposes at 590 °C (anhydrous)</p>
| BoilingPt = decomposes to [[cupric oxide]] at 650 °C
| RefractIndex = 1.724–1.739 (anhydrous)<ref>{{cite book|editor1=Anthony, John W. |editor2=Bideaux, Richard A. |editor3=Bladh, Kenneth W. |editor4=Nichols, Monte C. |title= Handbook of Mineralogy|publisher= Mineralogical Society of America|place= Chantilly, VA, US|chapter-url=http://rruff.info/doclib/hom/chalcocyanite.pdf|chapter=Chalcocyanite|isbn=978-0962209741 |volume=V. Borates, Carbonates, Sulfates|year=2003}}</ref><br>1.514–1.544 (pentahydrate)<ref>[[#CRC92|Haynes]], p. 10.240</ref>
| MagSus = 1330·10<sup>−6</sup> cm<sup>3</sup>/mol
 }}
| Section3       = {{Chembox Structure
| Coordination = 
| CrystalStruct = [[Orthorhombic]] (anhydrous, chalcocyanite), [[space group]] Pnma, [[Pearson symbol|oP24]], a = 0.839 nm, b = 0.669 nm, c = 0.483 nm.<ref>{{cite journal|doi=10.1107/S0365110X58000955|title=The crystal structure of the anhydrous sulphates of copper and zinc|year=1958|last1=Kokkoros|first1=P. A.|last2=Rentzeperis|first2=P. J.|journal=Acta Crystallographica|volume=11|issue=5|pages=361–364}}</ref><br/>[[Triclinic]] (pentahydrate), [[space group]] P{{overline|1}}, [[Pearson symbol|aP22]], a = 0.5986 nm, b = 0.6141 nm, c = 1.0736 nm, α = 77.333°, β = 82.267°, γ = 72.567°<ref>{{cite journal |title=Neutron-diffraction studies of CuSO<sub>4</sub> · 5H<sub>2</sub>O and CuSO<sub>4</sub> · 5D<sub>2</sub>O |journal=Z. Kristallogr. |year=1975 |volume=141 |pages=330–341 |last1=Bacon |first1=G. E. |last2=Titterton |first2=D. H. |doi=10.1524/zkri.1975.141.5-6.330 |issue=5–6 |bibcode=1975ZK....141..330B }}</ref>
 }}
| Section4       = 
| Section5       = {{Chembox Thermochemistry
| DeltaHf = −769.98 kJ/mol
| Entropy = 5 J/(K·mol)
 }}
| Section6       = {{Chembox Pharmacology
| ATCCode_prefix = V03
| ATCCode_suffix = AB20
}}
| Section7       = {{Chembox Hazards
| ExternalSDS = [https://web.archive.org/web/20211208192904/https://hazard.com//msds/mf/baker/baker/files/c5920.htm anhydrous]<br/>[https://web.archive.org/web/20220316151633/http://hazard.com/msds/mf/baker/baker/files/c5918.htm pentahydrate]
| GHSPictograms = {{GHS05}}{{GHS07}}{{GHS09}}
| GHSSignalWord  = Danger
| HPhrases = {{H-phrases|302|315|318|319|410}}
| PPhrases = {{P-phrases|P264|P264+P265|P270|P273|P280|P301+P317|P302+P352|P305+P351+P338|P305+P354+P338|P317|P321|P330|P332+P317|P337+P317|P362+P364|P391|P501}}
| NFPA-H = 2
| NFPA-F = 0
| NFPA-R = 1
| FlashPt = Non-flammable
| LD50 = 300 mg/kg (oral, rat)<ref>[https://chem.nlm.nih.gov/chemidplus/rn/7758-98-7 Cupric sulfate]. US National Institutes of Health</ref>
87 mg/kg (oral, mouse)
| PEL = TWA 1 mg/m<sup>3</sup> (as Cu)<ref name=PGCH>{{PGCH|0150}}</ref>
| REL = TWA 1 mg/m<sup>3</sup> (as Cu)<ref name=PGCH/>
| IDLH = TWA 100 mg/m<sup>3</sup> (as Cu)<ref name=PGCH/>
 }}
| Section8       = {{Chembox Related
| OtherCations = {{ubl|[[Iron(II) sulfate]]|[[Manganese(II) sulfate]]|[[Nickel(II) sulfate]]|[[Zinc sulfate]]}}
 }}
}}

'''Copper(II) sulfate''' is an [[inorganic compound]] with the [[chemical formula]] {{chem2|[[Copper|Cu]][[Sulfate|SO4]]}}. It forms [[hydrate]]s {{chem2|CuSO4*''n''H2O}}, where ''n'' can range from 1 to 7. The pentahydrate (''n'' = 5), a bright blue crystal, is the most commonly encountered hydrate of copper(II) sulfate,<ref>{{Cite web |last=Connor |first=Nick |date=2023-07-24 |title=Copper (II) Sulfate {{!}} Formula, Properties & Application |url=https://material-properties.org/copper-ii-sulfate/ |access-date=2024-02-03 |website=Material Properties |language=en-US}}</ref> while its [[anhydrous]] form is white.<ref>{{Cite web |last=Foundation |first=In association with Nuffield |title=A reversible reaction of hydrated copper(II) sulfate |url=https://edu.rsc.org/experiments/a-reversible-reaction-of-hydrated-copperii-sulfate/437.article |access-date=2024-02-03 |website=RSC Education |language=en}}</ref> Older names for the pentahydrate include '''blue vitriol''', '''bluestone''',<ref>{{cite web|publisher = [[Oxford University]]|title = Copper (II) sulfate MSDS|url = http://ptcl.chem.ox.ac.uk/MSDS/CO/copper_II_sulfate.html|access-date = 2007-12-31|archive-url = https://web.archive.org/web/20071011161441/http://ptcl.chem.ox.ac.uk/MSDS/CO/copper_II_sulfate.html|archive-date = 2007-10-11|url-status = dead}}</ref> '''vitriol of copper''',<ref name=fou>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.</ref> and '''Roman vitriol'''.<ref name=oxdic>Oxford University Press, "[https://web.archive.org/web/20161113175134/https://en.oxforddictionaries.com/definition/roman_vitriol Roman vitriol]", Oxford Living Dictionaries. Accessed on 2016-11-13</ref> It [[Exothermic reaction|exothermically]] dissolves in water to give the [[metal aquo complex|aquo complex]] {{chem2|[Cu(H2O)6](2+)}}, which has [[octahedral molecular geometry]]. The structure of the solid pentahydrate reveals a polymeric structure wherein copper is again octahedral but bound to four water ligands. The {{chem2|Cu(II)(H2O)4}} centers are interconnected by sulfate anions to form chains.<ref>{{cite journal | last1 = Ting | first1 = V. P. | last2 = Henry | first2 = P. F. | last3 = Schmidtmann | first3 = M. | last4 = Wilson | first4 = C. C. | last5 = Weller | first5 = M. T. | year = 2009| title = In situ neutron powder diffraction and structure determination in controlled humidities | journal = Chem. Commun. | volume = 2009 | issue = 48| pages = 7527–7529 | doi = 10.1039/B918702B | pmid = 20024268 }}</ref>

==Preparation and occurrence==
[[File:Synthesizing Copper Sulfate.jpg|left|thumb|210px|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 copper oxides with dilute sulfuric acid. For laboratory use, copper sulfate is usually purchased. Copper sulfate can also be produced by slowly [[Leaching (chemistry)|leaching]] low-grade [[copper ore]] in air; bacteria may be used to hasten the process.<ref name="copper.org">{{cite web|title=Uses of Copper Compounds: Copper Sulphate|url=http://www.copper.org/resources/properties/compounds/copper_sulfate01.html|website=copper.org|publisher=Copper Development Association Inc.|access-date=10 May 2015}}</ref>

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 of [[crystal size]] are provided based on its usage: large crystals (10–40&nbsp;mm), small crystals (2–10&nbsp;mm), snow crystals (less than 2&nbsp;mm), and windswept powder (less than 0.15&nbsp;mm).<ref name="copper.org"/>

==Chemical properties==
Copper(II) sulfate pentahydrate [[chemical decomposition|decomposes]] before melting. It loses two water molecules upon heating at {{cvt|63|C}}, followed by two more at {{cvt|109|C}} and the final water molecule at {{cvt|200|C}}.<ref>{{cite book |url=https://books.google.com/books?id=i9nyvTYBQtAC&pg=PA229 |pages=228–229 |title=Thermal decomposition of ionic solids |author=Andrew Knox Galwey |author2=Michael E. Green|publisher=Elsevier |year=1999 |isbn=978-0-444-82437-0}}</ref><ref>{{cite book|url=https://books.google.com/books?id=LxhQPdMRfVIC&pg=PA1263|page=1263|title=Inorganic chemistry|first=Egon |last=Wiberg |author2=Nils Wiberg |author3=Arnold Frederick Holleman |publisher=Academic Press|year=2001|isbn=978-0-12-352651-9}}</ref>

The chemistry of aqueous copper sulfate is simply that of copper [[aquo complex]], since the sulfate is not bound to copper in such solutions. Thus, such solutions react with concentrated [[hydrochloric acid]] to give [[tetrachlorocuprate]](II):
:{{chem2|Cu(2+) + 4 Cl− → [CuCl4](2-)}}
Similarly treatment of such solutions with zinc gives metallic copper, as described by this simplified equation:<ref>{{cite journal |doi=10.15227/orgsyn.014.0066|title=P-Nitrophenyl Ether |journal=Organic Syntheses |year=1934 |volume=14 |page=66|author=Ray Q. Brewster, Theodore Groening }}</ref>
:{{chem2|CuSO4 + Zn → Cu + ZnSO4}}

A further illustration of such [[Single displacement reaction|single metal replacement reactions]] occurs when a piece of iron is submerged in a solution of copper sulfate:
:{{chem2|Fe + CuSO4 → FeSO4 + Cu}}
In high school and general chemistry education, copper sulfate is used as an electrolyte for [[galvanic cell]]s, 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.<ref>{{cite book|last1=Zumdahl|first1=Steven|last2=DeCoste|first2=Donald|title=Chemical Principles|date=2013|publisher=Cengage Learning|isbn=978-1-285-13370-6|pages=506–507}}</ref>
:{{chem2|Cu(2+) + 2e− → Cu (cathode)}}, E<sup>°</sup><sub>cell</sub> = 0.34 V

Copper sulfate is commonly included in teenage [[chemistry set]]s and undergraduate experiments.<ref>{{cite journal |doi=10.1021/ed079p486|title=A Copper-Sulfate-Based Inorganic Chemistry Laboratory for First-Year University Students That Teaches Basic Operations and Concepts |year=2002 |last1=Rodríguez |first1=Emilio |last2=Vicente |first2=Miguel Angel |journal=Journal of Chemical Education |volume=79 |issue=4 |page=486 |bibcode=2002JChEd..79..486R }}</ref> It is often used to grow crystals in [[school]]s and in [[Copper electroplating]] 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 {{chem2|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.{{cn|date=January 2025}} Copper(II) sulfate pentahydrate can easily be produced by crystallization from solution as copper(II) sulfate, which is [[hygroscopic]].

==Uses==

===As a fungicide and herbicide===
Copper sulfate has been used for control of [[algae]] in lakes and related fresh waters subject to [[eutrophication]]. It "remains the most effective algicidal treatment".<ref>{{cite journal |doi=10.1051/jp4:20030547|title=Fate and forms of Cu in a reservoir ecosystem following copper sulfate treatment (Saint Germain les Belles, France) |year=2003 |last1=Van Hullebusch |first1=E. |last2=Chatenet |first2=P. |last3=Deluchat |first3=V. |last4=Chazal |first4=P. M. |last5=Froissard |first5=D. |last6=Lens |first6=P. N.L. |last7=Baudu |first7=M. |journal=Journal de Physique IV (Proceedings) |volume=107 |pages=1333–1336 }}</ref><ref>{{cite journal |doi=10.1016/S0043-1354(00)00054-3|title=Forms and fate of Cu in a source drinking water reservoir following CuSO4 treatment |year=2000 |last1=Haughey |first1=M. |journal=Water Research |volume=34 |issue=13 |pages=3440–3452 }}</ref>

[[Bordeaux mixture]], a suspension of copper(II) sulfate ({{chem2|CuSO4}}) and [[calcium hydroxide]] ({{chem2|Ca(OH)2}}), is used to control fungus on [[grape]]s, [[melon]]s, and other [[berries]].<ref>{{cite journal|title = Uses of Copper Compounds: Copper Sulfate's Role in Agriculture|journal = Annals of Applied Biology|volume = 20|issue = 2|pages = 342–363|doi=10.1111/j.1744-7348.1933.tb07770.x|year = 1933|last1 = Martin|first1 = Hubert}}</ref> It is produced by mixing a water solution of copper sulfate and a suspension of [[slaked lime]].

<!-- too dated: ''Cheshunt compound'' (discontinued), a commercial mixture of copper sulfate and [[ammonium carbonate]], was used in [[horticulture]] to prevent [[damping off]] in seedlings.<ref>Coutts, J, Edwards, A, Osborn, A, & Preston, GH, ''The Complete Book of Gardening'', p. 533, Ward Lock, London (1954)</ref> As a non-agricultural [[herbicide]], copper sulfate is used to control invasive [[aquatic plants]] and the roots of plants situated near water pipes. It is used in [[swimming pool]]s as an algicide. -->
A dilute solution of copper sulfate is used to treat [[aquarium]] fishes for parasitic infections,<ref>{{cite web|title = All About Copper Sulfate|publisher = National Fish Pharmaceuticals|url = http://www.nationalfishpharm.com/Q&A/all_about_copper.html|access-date = 2007-12-31}}</ref> and is also used to remove snails from aquariums and [[Zebra mussel|zebra mussels]] from water pipes.<ref>{{Cite web|date=2020-10-26|title=With Zebra mussels here to stay, Austin has a plan to avoid stinky drinking water|url=https://www.kxan.com/news/local/austin/with-zebra-mussels-here-to-stay-austin-has-a-plan-to-avoid-stinky-drinking-water/|access-date=2020-10-28|website=KXAN Austin|language=en-US}}</ref> Copper ions are highly toxic to fish. Most species of algae can be controlled with very low concentrations of copper sulfate.

===Analytical reagent===
Several chemical tests utilize copper sulfate. It is used in [[Fehling's solution]] and [[Benedict's solution]] to test for [[reducing sugar]]s, 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 dropped into a solution of copper sulfate of known [[specific gravity]]—blood with sufficient [[hemoglobin]] sinks rapidly due to its density, whereas blood which sinks slowly or not at all has an insufficient amount of hemoglobin.<ref>{{cite book|title = Basic Medical Laboratory Techniques|first= Barbara H. |last=Estridge |author2=Anna P. Reynolds |author3=Norma J. Walters |page = 166|publisher = Thomson Delmar Learning |year = 2000 |isbn = 978-0-7668-1206-2}}</ref>  Clinically relevant, however, modern laboratories utilize automated blood analyzers for accurate quantitative hemoglobin determinations, as opposed to older qualitative means.{{citation needed|date=January 2023}}

In a [[flame test]], the copper [[ion]]s of copper sulfate 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]].<ref>{{cite book|last=Hoffman |first=R. V. |title= Copper(II) Sulfate, in Encyclopedia of Reagents for Organic Synthesis|year= 2001|publisher= John Wiley & Sons|doi=10.1002/047084289X.rc247|chapter=Copper(II) Sulfate |isbn=978-0471936237 }}</ref> The anhydrous salt is used as a dehydrating agent for forming and manipulating [[acetal]] groups.<ref name="Kocienski2005">{{cite book|author=Philip J. Kocienski| title=Protecting Groups|url=https://books.google.com/books?id=eI8p5B1uTJMC&pg=PA58| year=2005| publisher=Thieme| isbn=978-1-58890-376-1|page=58}}</ref> The hydrated salt can be intimately mingled with [[potassium permanganate]] to give an oxidant for the conversion of primary alcohols.<ref>{{OrgSynth|last = Jefford |first=C. W. |last2=Li |first2=Y. |last3=Wang |first3=Y.|prep = cv9p0462|title = A Selective, Heterogeneous Oxidation using a Mixture of Potassium Permanganate and Cupric Sulfate: (3aS,7aR)-Hexahydro-(3S,6R)-Dimethyl-2(3H)-Benzofuranone|collvol = 9|collvolpages = 462}}</ref>

=== Rayon production===
Reaction with [[ammonium hydroxide]] yields [[tetraamminecopper(II) sulfate]] or [[Schweizer's reagent]] which was used to dissolve [[cellulose]] in the industrial production of [[Rayon]].

===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 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.<ref name="ReferenceA">{{cite web|last1=Copper Development Association|title=Uses of Copper Compounds: Table A - Uses of Copper Sulphate|url=http://www.copper.org/resources/properties/compounds/table_a.html|website=copper|publisher=Copper Development Association Inc.|access-date=12 May 2015}}</ref> 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.<ref>{{cite web|last1=Partin |first1=Lee |title=The Blues: Part 2 |url=http://www.skylighter.com/fireworks/how-to-make/blue-copper-fireworks-stars.asp |website=skylighter |publisher=Skylighter.Inc |access-date=12 May 2015 |url-status=dead |archive-url=https://web.archive.org/web/20101221162804/http://www.skylighter.com/fireworks/how-to-make/blue-copper-fireworks-stars.asp |archive-date=21 December 2010 }}</ref>

[[File:Copper sulfate etching bath.JPG|thumb|Lowering a copper etching plate into the copper sulfate solution]]

Copper sulfate was once used to kill [[bromeliads]], which serve as mosquito breeding sites.<ref>{{cite book|last1=Despommier|last2=Gwadz|last3=Hotez|last4=Knirsch|title=Parasitic Disease|date=June 2005|publisher=Apple Tree Production L.L.C|location=NY|isbn=978-0970002778|pages=Section 4.2|edition=5|url=http://www.medicalecology.org/diseases/malaria/malaria.htm|access-date=12 May 2015}}</ref> Copper sulfate is used as a molluscicide to treat [[bilharzia]] in tropical countries.<ref name="ReferenceA"/>

====Art====
In 2008, the artist [[Roger Hiorns]] filled an abandoned waterproofed [[council 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 titled ''Seizure''.<ref>{{cite web|url=https://www.artangel.org.uk/project/seizure/ |title=Seizure |publisher=Artangel.org.uk |access-date=2021-10-05}}</ref> Since 2011, it has been on exhibition at the [[Yorkshire Sculpture Park]].<ref>{{cite web|url=http://www.ysp.co.uk/exhibitions/roger-hiorns-seizure|title=Roger Hiorns: Seizure|publisher=Yorkshire Sculpture Park|access-date=2015-02-22|archive-url=https://web.archive.org/web/20150222005608/http://www.ysp.co.uk/exhibitions/roger-hiorns-seizure|archive-date=2015-02-22|url-status=dead}}</ref>

====Etching====
Copper(II) sulfate is used to etch zinc, aluminium, or copper plates for [[Intaglio (printmaking)|intaglio printmaking]].<ref>[http://www.greenart.info/galvetch/bordeaux.htm greenart.info], Bordeau etch, 2009-01-18, retrieved 2011-06-02.</ref><ref>[http://www.ndiprintmaking.ca/?p=74 ndiprintmaking.ca], The Chemistry of using Copper Sulfate Mordant, 2009-04-12, retrieved 2011-06-02.</ref>
It is also used to etch designs into copper for jewelry, such as for [[Champlevé]].<ref>[http://mordent.com/etch-howto/ http://mordent.com/etch-howto/], How to Electrolytically etch in copper, brass, steel, nickel silver or silver, retrieved 2015-05-2015.</ref>

====Dyeing====
Copper(II) sulfate can be used as a [[mordant]] in vegetable [[dyeing]]. It often highlights the green tints of the specific dyes.{{citation needed|date=December 2022}}

====Electronics====
An aqueous solution of copper(II) sulfate is often used as the resistive element in [[liquid resistor]]s.{{citation needed|date=December 2022}}

In electronic and microelectronic industry a bath of {{chem2|CuSO4*5H2O}} and [[sulfuric acid]] ({{chem2|H2SO4}}) is often used for [[Electroplating|electrodeposition]] of copper.<ref>{{Cite book |url=https://www.worldcat.org/oclc/868688018 |title=Copper Electrodeposition for Nanofabrication of Electronics Devices |date=2014 |author1=K. Kondo |author2=Rohan N. Akolkar |author3=Dale P. Barkey |author4=Masayuki Yokoi |isbn=978-1-4614-9176-7 |location=New York |oclc=868688018}}</ref>

==Other forms of copper sulfate==
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 as [[chalcocyanite]].<ref>{{Cite web|url=https://www.mindat.org/min-963.html|title=Chalcocyanite|website=www.mindat.org}}</ref> The pentahydrate also occurs in nature as [[chalcanthite]]. Other rare copper sulfate minerals include [[bonattite]] (trihydrate),<ref>{{Cite web|url=https://www.mindat.org/min-718.html|title=Bonattite|website=www.mindat.org}}</ref> [[boothite]] (heptahydrate),<ref>{{Cite web|url=https://www.mindat.org/min-720.html|title=Boothite|website=www.mindat.org}}</ref> and the monohydrate compound poitevinite.<ref>{{Cite web|url=https://www.mindat.org/min-3249.html|title=Poitevinite|website=www.mindat.org}}</ref><ref name="IMA">{{Cite web|url=https://www.ima-mineralogy.org/Minlist.htm|title=List of Minerals|date=March 21, 2011|website=www.ima-mineralogy.org}}</ref> There are numerous other, more complex, copper(II) sulfate minerals known, with environmentally important basic copper(II) sulfates like langite and posnjakite.<ref name="IMA"/><ref>{{Cite web|url=https://www.mindat.org/min-2322.html|title=Langite|website=www.mindat.org}}</ref><ref>{{Cite web|url=https://www.mindat.org/min-3265.html|title=Posnjakite|website=www.mindat.org}}</ref>
{{Gallery
| title        = Forms of copper(II) sulfate
| align        =
| footer       =
| style        =
| state        =
| height       =
| width        =
| captionstyle =
| File:Copper sulfate anhydrous.jpg
 | Anhydrous {{chem2|CuSO4}}
 | alt1=
| File:Сульфат_меди_одноводный.jpg
 | Copper(II) sulfate monohydrate
 | alt2=
| File:Síran měďnatý.jpg
 | Copper(II) sulfate pentahydrate
 | alt3=
| File:Arseniosiderite-Boothite-sea69a.jpg
 | The rare mineral ''boothite'' ({{chem2|CuSO4*7H2O}})
 | alt4=
}}

== Toxicological effects ==
{{See also|Copper toxicity}}
Copper(II) salts have an [[LD50]] of 100&nbsp;mg/kg.<ref>Windholz, M., ed. 1983. ''The Merck Index''. Tenth edition. Rahway, NJ: Merck and Company.</ref><ref name="EPA-Guidance">{{citation |publisher=U.S. Environmental Protection Agency, Office of Pesticide Programs |year=1986 |title=Guidance for reregistration of pesticide products containing copper sulfate. Fact sheet no. 100. |location=Washington, DC}}</ref>

Copper(II) sulfate was used in the past as an [[emetic]].<ref name="pmid4385403">{{cite journal |last1=Holtzmann |first1=N. A. |last2=Haslam |first2=R. H. |title=Elevation of serum copper following copper sulfate as an emetic |journal=Pediatrics |volume=42 |issue=1 |pages=189–93 |date=July 1968 |doi=10.1542/peds.42.1.189 |pmid=4385403 |s2cid=32740524 }}</ref> It is now considered too toxic for this use.<ref name="isbn0-8385-8172-2">{{cite book |last=Olson |first=Kent C. |title=Poisoning & drug overdose |publisher=Lange Medical Mooks/McGraw-Hill |location=New York |year=2004 |page=[https://archive.org/details/poisoningdrugove00olso/page/175 175] |isbn=978-0-8385-8172-8 |url=https://archive.org/details/poisoningdrugove00olso/page/175 }}</ref> It is still listed as an [[antidote]] in the [[World Health Organization]]'s [[Anatomical Therapeutic Chemical Classification System]].<ref>{{ATC|V03|AB20}}</ref>
{{clear}}

==See also==
*[[Chalcanthum]]
*[[Vitriol]]

==References==
{{reflist|30em}}

==Bibliography==
*{{cite book|ref =CRC92|editor=Haynes, William M.|year=2011|title=CRC Handbook of Chemistry and Physics (92nd ed.). |place=Boca Raton, FL |publisher=CRC Press|isbn=978-1439855119}}

==External links==
*{{Commons category-inline}}
*{{ICSC|0751|07}}
*{{ICSC|1416|14}}
*[https://web.archive.org/web/20080302034606/http://www.npi.gov.au/database/substance-info/profiles/27.html National Pollutant Inventory – Copper and compounds fact sheet]
{{Copper compounds}}
{{Sulfates}}
{{Antidotes}}
{{Sulfur compounds}}
{{Authority control}}

{{DEFAULTSORT:Copper(Ii) Sulfate}}
[[Category:Copper(II) compounds]]
[[Category:Sulfates]]
[[Category:Desiccants]]
[[Category:Herbicides]]