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]
 Crystals ofCuSO4·5H2O | |
-sulfate-pentahydrate-xtal-1985-Cu-coord-3D-bs-17.png) | |
-sulfate-pentahydrate-unit-cell-1985-3D-bs-17.png) | |
| Names | |
|---|---|
| IUPAC name Copper(II) sulfate | |
Other names
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| Identifiers | |
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3D model (JSmol) | |
| ChEBI | |
| ChEMBL | |
| ChemSpider |
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| ECHA InfoCard | 100.028.952 |
| EC Number |
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| 8294 | |
| KEGG |
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| RTECS number |
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| UNII |
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| Properties | |
| CuSO4 (anhydrous) CuSO4·5H2O (pentahydrate) | |
| Molar mass | 159.60 g/mol (anhydrous)[2] 249.685 g/mol (pentahydrate)[2] |
| Appearance | gray-white (anhydrous) blue (pentahydrate) |
| Density | 3.60 g/cm3 (anhydrous)[2] 2.286 g/cm3 (pentahydrate)[2] |
| Melting point | 110 °C (230 °F; 383 K)decomposes 560 °Cdecomposes[2](pentahydrate) Fully decomposes at 590 °C (anhydrous) |
| Boiling point | decomposes tocupric oxide at 650 °C |
| pentahydrate 316 g/L (0 °C) 2033 g/L (100 °C) anhydrous 168 g/L (10 °C) 201 g/L (20 °C) 404 g/L (60 °C) 770 g/L (100 °C)[3] | |
| Solubility | anhydrous insoluble inethanol[2] pentahydrate soluble inmethanol[2] 10.4 g/L (18 °C) insoluble inethanol andacetone |
| 1330·10−6 cm3/mol | |
Refractive index (nD) | 1.724–1.739 (anhydrous)[4] 1.514–1.544 (pentahydrate)[5] |
| Structure | |
| Orthorhombic (anhydrous, chalcocyanite),space group Pnma,oP24, a = 0.839 nm, b = 0.669 nm, c = 0.483 nm.[6] Triclinic (pentahydrate),space group P1,aP22, a = 0.5986 nm, b = 0.6141 nm, c = 1.0736 nm, α = 77.333°, β = 82.267°, γ = 72.567°[7] | |
| Thermochemistry | |
Std molar entropy(S⦵298) | 5 J/(K·mol) |
Std enthalpy of formation(ΔfH⦵298) | −769.98 kJ/mol |
| Pharmacology | |
| V03AB20 (WHO) | |
| Hazards | |
| GHS labelling: | |
   | |
| Danger | |
| H302,H315,H318,H319,H410 | |
| 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 704 (fire diamond) | |
| Flash point | Non-flammable |
| Lethal dose or concentration (LD, LC): | |
LD50 (median dose) | 300 mg/kg (oral, rat)[9] 87 mg/kg (oral, mouse) |
| NIOSH (US health exposure limits): | |
PEL (Permissible) | TWA 1 mg/m3 (as Cu)[8] |
REL (Recommended) | TWA 1 mg/m3 (as Cu)[8] |
IDLH (Immediate danger) | TWA 100 mg/m3 (as Cu)[8] |
| Safety data sheet (SDS) | anhydrous pentahydrate |
| Related compounds | |
Othercations | |
Except where otherwise noted, data are given for materials in theirstandard state (at 25 °C [77 °F], 100 kPa). | |
Preparation and occurrence
edit
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]
Chemical properties
editCopper(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.
Uses
editAs a fungicide and herbicide
editCopper sulfate has been used for control ofalgae in lakes and related fresh waters subject toeutrophication. It "remains the most effective algicidal treatment".[22][23]
Bordeaux mixture, a suspension of copper(II) sulfate (CuSO4) andcalcium hydroxide (Ca(OH)2), is used to control fungus ongrapes,melons, and otherberries.[24] It is produced by mixing a water solution of copper sulfate and a suspension ofslaked lime.
A dilute solution of copper sulfate is used to treataquarium fishes for parasitic infections,[25] and is also used to remove snails from aquariums and zebra 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.
Analytical reagent
editSeveral chemical tests utilize copper sulfate. It is used inFehling's solution andBenedict's solution to test forreducing sugars, which reduce the soluble blue copper(II) sulfate to insoluble redcopper(I) oxide. Copper(II) sulfate is also used in theBiuret reagent to test for proteins.
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.
Organic synthesis
editCopper 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]
Rayon production
editReaction withammonium hydroxide yieldstetraamminecopper(II) sulfate orSchweizer's reagent which was used to dissolvecellulose in the industrial production ofRayon.
Niche uses
editCopper(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 used in firework manufacture as a blue coloring agent, but it is not safe to mix copper sulfate with chlorates when mixing firework powders.[32]
Copper sulfate was once used to killbromeliads, which serve as mosquito breeding sites.[33] Copper sulfate is used as a molluscicide to treatbilharzia in tropical countries.[31]
Art
editIn 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, leavingcrystal-covered walls, floors and ceilings. The work is titledSeizure.[34] Since 2011, it has been on exhibition at theYorkshire Sculpture Park.[35]
Etching
editCopper(II) sulfate is used to etch zinc, aluminium, or copper plates forintaglio printmaking.[36][37]It is also used to etch designs into copper for jewelry, such as forChamplevé.[38]
Dyeing
editCopper(II) sulfate can be used as amordant in vegetabledyeing. It often highlights the green tints of the specific dyes.[citation needed]
Electronics
editAn aqueous solution of copper(II) sulfate is often used as the resistive element inliquid resistors.[citation needed]
In electronic and microelectronic industry a bath ofCuSO4·5H2O andsulfuric acid (H2SO4) is often used forelectrodeposition of copper.[39]
Other forms of copper sulfate
editAnhydrous 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.[40] The pentahydrate also occurs in nature aschalcanthite. Other rare copper sulfate minerals includebonattite (trihydrate),[41]boothite (heptahydrate),[42] and the monohydrate compound poitevinite.[43][44] There are numerous other, more complex, copper(II) sulfate minerals known, with environmentally important basic copper(II) sulfates like langite and posnjakite.[44][45][46]
AnhydrousCuSO4
Copper(II) sulfate monohydrate
Copper(II) sulfate pentahydrate
The rare mineralboothite (CuSO4·7H2O)
Toxicological effects
editCopper(II) salts have anLD50 of 100 mg/kg.[47][48]
Copper(II) sulfate was used in the past as anemetic.[49] It is now considered too toxic for this use.[50] It is still listed as anantidote in theWorld Health Organization'sAnatomical Therapeutic Chemical Classification System.[51]
See also
editReferences
edit- ^Varghese, J. N.; Maslen, E. N. (1985). "Electron density in non-ideal metal complexes. I. Copper sulphate pentahydrate".Acta Crystallographica Section B.41 (3):184–190.doi:10.1107/S0108768185001914.
- ^abcdefgHaynes, p. 4.62
- ^Rumble, John, ed. (2018).CRC Handbook of Chemistry and Physics (99th ed.). CRC Press, Taylor & Francis Group. pp. 5–179.ISBN 9781138561632.
- ^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.ISBN 978-0962209741.
- ^Haynes, p. 10.240
- ^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.
- ^Bacon, G. E.; Titterton, D. H. (1975). "Neutron-diffraction studies of CuSO4 · 5H2O and CuSO4 · 5D2O".Z. Kristallogr.141 (5–6):330–341.Bibcode:1975ZK....141..330B.doi:10.1524/zkri.1975.141.5-6.330.
- ^abcNIOSH Pocket Guide to Chemical Hazards."#0150".National Institute for Occupational Safety and Health (NIOSH).
- ^Cupric sulfate. US National Institutes of Health
- ^Connor, Nick (2023-07-24)."Copper (II) Sulfate | Formula, Properties & Application".Material Properties. Retrieved2024-02-03.
- ^Foundation, In association with Nuffield."A reversible reaction of hydrated copper(II) sulfate".RSC Education. Retrieved2024-02-03.
- ^"Copper (II) sulfate MSDS".Oxford University. Archived fromthe original on 2007-10-11. Retrieved2007-12-31.
- ^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.PMID 20024268.
- ^ab"Uses of Copper Compounds: Copper Sulphate".copper.org. Copper Development Association Inc. Retrieved10 May 2015.
- ^Andrew Knox Galwey; Michael E. Green (1999).Thermal decomposition of ionic solids. Elsevier. pp. 228–229.ISBN 978-0-444-82437-0.
- ^Wiberg, Egon; Nils Wiberg; Arnold Frederick Holleman (2001).Inorganic chemistry. Academic Press. p. 1263.ISBN 978-0-12-352651-9.
- ^Ray Q. Brewster, Theodore Groening (1934). "P-Nitrophenyl Ether".Organic Syntheses.14: 66.doi:10.15227/orgsyn.014.0066.
- ^Zumdahl, Steven; DeCoste, Donald (2013).Chemical Principles. Cengage Learning. pp. 506–507.ISBN 978-1-285-13370-6.
- ^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.
- ^"All About Copper Sulfate". National Fish Pharmaceuticals. Retrieved2007-12-31.
- ^"With Zebra mussels here to stay, Austin has a plan to avoid stinky drinking water".KXAN Austin. 2020-10-26. Retrieved2020-10-28.
- ^Estridge, Barbara H.; Anna P. Reynolds; Norma J. Walters (2000).Basic Medical Laboratory Techniques. Thomson Delmar Learning. p. 166.ISBN 978-0-7668-1206-2.
- ^Hoffman, R. V. (2001). "Copper(II) Sulfate".Copper(II) Sulfate, in Encyclopedia of Reagents for Organic Synthesis. John Wiley & Sons.doi:10.1002/047084289X.rc247.ISBN 978-0471936237.
- ^Philip J. Kocienski (2005).Protecting Groups. Thieme. p. 58.ISBN 978-1-58890-376-1.
- ^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".Organic Syntheses;Collected Volumes, vol. 9, p. 462.
- ^abCopper Development Association."Uses of Copper Compounds: Table A - Uses of Copper Sulphate".copper. Copper Development Association Inc. Retrieved12 May 2015.
- ^Partin, Lee."The Blues: Part 2".skylighter. Skylighter.Inc. Archived fromthe original on 21 December 2010. Retrieved12 May 2015.
- ^Despommier; Gwadz; Hotez; Knirsch (June 2005).Parasitic Disease (5 ed.). NY: Apple Tree Production L.L.C. pp. Section 4.2.ISBN 978-0970002778. Retrieved12 May 2015.
- ^"Seizure". Artangel.org.uk. Retrieved2021-10-05.
- ^"Roger Hiorns: Seizure". Yorkshire Sculpture Park. Archived fromthe original on 2015-02-22. Retrieved2015-02-22.
- ^greenart.info, Bordeau etch, 2009-01-18, retrieved 2011-06-02.
- ^ndiprintmaking.ca, The Chemistry of using Copper Sulfate Mordant, 2009-04-12, retrieved 2011-06-02.
- ^http://mordent.com/etch-howto/, How to Electrolytically etch in copper, brass, steel, nickel silver or silver, retrieved 2015-05-2015.
- ^K. Kondo; Rohan N. Akolkar; Dale P. Barkey; Masayuki Yokoi (2014).Copper Electrodeposition for Nanofabrication of Electronics Devices. New York.ISBN 978-1-4614-9176-7.OCLC 868688018.
{{cite book}}: CS1 maint: location missing publisher (link) - ^"Chalcocyanite".www.mindat.org.
- ^"Bonattite".www.mindat.org.
- ^"Boothite".www.mindat.org.
- ^"Poitevinite".www.mindat.org.
- ^ab"List of Minerals".www.ima-mineralogy.org. March 21, 2011.
- ^"Langite".www.mindat.org.
- ^"Posnjakite".www.mindat.org.
- ^Windholz, M., ed. 1983.The Merck Index. Tenth edition. Rahway, NJ: Merck and Company.
- ^Guidance for reregistration of pesticide products containing copper sulfate. Fact sheet no. 100., Washington, DC: U.S. Environmental Protection Agency, Office of Pesticide Programs, 1986
- ^Holtzmann, N. A.; Haslam, R. H. (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.
- ^Olson, Kent C. (2004).Poisoning & drug overdose. New York: Lange Medical Mooks/McGraw-Hill. p. 175.ISBN 978-0-8385-8172-8.
- ^V03AB20 (WHO)
Bibliography
edit- Haynes, William M., ed. (2011).CRC Handbook of Chemistry and Physics (92nd ed.). Boca Raton, FL: CRC Press.ISBN 978-1439855119.
