Copper(II) sulfate: Difference between revisions
ClueBot NG (talk | contribs) m Reverting possible vandalism by 173.46.211.180 to version by Brichcja. False positive? Report it. Thanks, ClueBot NG. (1530541) (Bot) |
|||
| Line 94: | Line 94: | ||
Dehydration proceeds by decomposition of the tetraaquacopper(2+) moiety, two opposing aqua groups are lost to give a diaquacopper(2+) moiety. The second dehydration step occurs with the final two aqua groups are lost. Complete dehydration occurs when the only unbound water molecule is lost. |
Dehydration proceeds by decomposition of the tetraaquacopper(2+) moiety, two opposing aqua groups are lost to give a diaquacopper(2+) moiety. The second dehydration step occurs with the final two aqua groups are lost. Complete dehydration occurs when the only unbound water molecule is lost. |
||
We can induce water loss by heating the substance for extended amounts of time in a crucible. this created the dehydrated version of the cupric sulfate. we can create the fire the fire by burniung things under the crucible, such as methane, american flags, and propane. after heating it is very important that you go out and let the salt cool. the best way to do this is to let some cool air in, so open the door, get on the floor, and walk the dinosaur. |
|||
At {{convert|650|C}}, copper(II) sulfate decomposes into [[copper(II) oxide]] (CuO) and [[sulfur trioxide]] (SO<sub>3</sub>). |
At {{convert|650|C}}, copper(II) sulfate decomposes into [[copper(II) oxide]] (CuO) and [[sulfur trioxide]] (SO<sub>3</sub>). |
||
Revision as of 17:10, 27 February 2013
| |||
| |||
| Names | |||
|---|---|---|---|
| IUPAC name
Copper(II) sulfate
| |||
| Other names
Cupric sulfate
Blue vitriol (pentahydrate) Bluestone (pentahydrate) Bonattite (trihydrate mineral) Boothite (heptahydrate mineral) Chalcanthite (pentahydrate mineral) Chalcocyanite (mineral) | |||
| Identifiers | |||
3D model (JSmol)
|
|||
| ChEBI | |||
| ChEMBL | |||
| ChemSpider | |||
| ECHA InfoCard | 100.028.952 | ||
| EC Number |
| ||
| KEGG | |||
PubChem CID
|
|||
| RTECS number |
| ||
| UNII | |||
CompTox Dashboard (EPA)
|
|||
| |||
| |||
| Properties | |||
| CuSO4 | |||
| Molar mass | 159.62 g/mol (anhydrous) 249.70 g/mol (pentahydrate) | ||
| Appearance | blue (pentahydrate) gray-white (anhydrous) | ||
| Density | 3.603 g/cm3 (anhydrous) 2.284 g/cm3 (pentahydrate) | ||
| Melting point | 110 °C (·4H2O) 150 °C (423 K) (·5H2O) < 650 °C decomp. | ||
| pentahydrate 316 g/L (0 °C) 2033 g/L (100 °C) anhydrous 243 g/L (0 °C) 320 g/L (20 °C) 618 g/L (60 °C) 1140 g/L (100 °C) | |||
| Solubility | anhydrous insoluble in ethanol pentahydrate soluble in methanol 10.4 g/L (18 °C) insoluble in ethanol | ||
Refractive index (nD)
|
1.733 (anhydrous) 1.514 (pentahydrate) | ||
| Structure | |||
| Orthorhombic (chalcocyanite), space group Pnma, oP24, a = 0.839 nm, b = 0.669 nm, c = 0.483 nm[1] 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°[2] | |||
| Thermochemistry | |||
Std molar
entropy (S⦵298) |
109.05 J K−1 mol−1 | ||
| Hazards | |||
| NFPA 704 (fire diamond) | |||
| Flash point | Non-inflammable | ||
| Lethal dose or concentration (LD, LC): | |||
LD50 (median dose)
|
300 mg/kg (oral, rat) 87 mg/kg (oral, mouse) 470 mg/kg (oral, mammal) | ||
| Related compounds | |||
Other cations
|
Iron(II) sulfate Manganese(II) sulfate Nickel(II) sulfate Zinc sulfate | ||
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
| |||
-sulfate-unit-cell-3D-balls.png)
-sulfate-3D-vdW.png)
Copper(II) sulfate, also known as cupric sulfate or copper sulphate, is the chemical compound with the chemical formula CuSO4. This salt exists as a series of compounds that differ in their degree of hydration. The anhydrous form is a pale green or gray-white powder, whereas the pentahydrate (CuSO4·5H2O), the most commonly encountered salt, is bright blue. CuSO4·5H2O is in a shade of blue, and very toxic to the environment, irritating to the eyes and skin, and also can be harmful if swallowed. Copper(II) sulfate exothermically dissolves in water to give the aquo complex [Cu(H2O)6]2+, which has octahedral molecular geometry and is paramagnetic. Other names for copper(II) sulfate are "blue vitriol" and "bluestone".[3]
Preparation and occurrence
Copper sulfate is produced industrially by treating copper metal with hot concentrated sulfuric acid or its oxides with dilute sulfuric acid. For laboratory use, copper sulfate is usually purchased.
The anhydrous form occurs as a rare mineral known as chalcocyanite. The hydrated copper sulfate occurs in nature as chalcanthite (pentahydrate), and two more rare ones: bonattite (trihydrate) and boothite (heptahydrate).
Chemical properties
Copper(II) sulfate pentahydrate decomposes before melting at 150 °C (302 °F), losing two water molecules at 63 °C (145 °F), followed by two more at 109 °C (228 °F) and the final water molecule at 200 °C (392 °F).[4][5]
Dehydration proceeds by decomposition of the tetraaquacopper(2+) moiety, two opposing aqua groups are lost to give a diaquacopper(2+) moiety. The second dehydration step occurs with the final two aqua groups are lost. Complete dehydration occurs when the only unbound water molecule is lost. We can induce water loss by heating the substance for extended amounts of time in a crucible. this created the dehydrated version of the cupric sulfate. we can create the fire the fire by burniung things under the crucible, such as methane, american flags, and propane. after heating it is very important that you go out and let the salt cool. the best way to do this is to let some cool air in, so open the door, get on the floor, and walk the dinosaur.
At 650 °C (1,202 °F), copper(II) sulfate decomposes into copper(II) oxide (CuO) and sulfur trioxide (SO3).
Its blue color is due to water of hydration. When heated in an open flame the crystals are dehydrated and turn grayish-white.[6]
Copper sulfate reacts with concentrated hydrochloric acid very strongly. In the reaction the blue solution of copper(II) turns green, due to the formation of tetrachlorocuprate(II):
- Cu2+ + 4 Cl– → CuCl2−
4
It also reacts with more reactive metals than copper (e.g. Mg, Fe, Zn, Al, Sn, Pb, etc.):
- CuSO4 + Zn → ZnSO4 +Cu
- CuSO4 + Fe → FeSO4 + Cu
- CuSO4 + Mg → MgSO4 + Cu
- CuSO4 + Sn → SnSO4 + Cu
- 3 CuSO4 + 2 Al → Al2(SO4)3 + 3 Cu
Some metals more reactive than others like magnesium and aluminium will cause a secondary reaction. They will form hydroxides with the water while releasing some hydrogen gas.
The copper formed is deposited on the surface of the other metal. The reaction stops when no free surface of the metal is present anymore.
Uses
As a herbicide, fungicide and pesticide
Copper sulfate pentahydrate is a fungicide.[7] However, some fungi are capable of adapting to elevated levels of copper ions.[8] Mixed with lime it is called Bordeaux mixture and used to control fungus on grapes, melons, and other berries.[9] Another application is Cheshunt compound, a mixture of copper sulfate and ammonium carbonate used in horticulture to prevent damping off in seedlings. Its use as a herbicide is not agricultural, but instead for control of invasive aquatic plants and the roots of plants near pipes containing water. It is used in swimming pools as an algaecide. A dilute solution of copper sulfate is used to treat aquarium fish for parasitic infections,[10] and is also used to remove snails from aquariums. Copper ions are highly toxic to fish, so care must be taken with the dosage. Most species of algae can be controlled with very low concentrations of copper sulfate. Copper sulfate inhibits growth of bacteria such as Escherichia coli.
For most of the twentieth century, chromated copper arsenate (CCA) was the dominant type of wood preservation for uses other than deep driven piles, utility poles, and railroad ties. To make pressure-treated wood, a large cylinder is filled with an aqueous chemical bath. Copper sulfate pentahydrate is dissolved in the water along with other additives prior to the lumber being placed inside the cylinder. When the cylinder is pressurized, the chemicals are absorbed by the wood, giving the wood fungicidal, insecticidal, and UV-light-reflecting properties that help preserve it.
Niche uses
Being a relatively benign and cheap reagent, copper(II) sulfate has attracted many niche applications over the centuries.
Analytical reagent
Several chemical tests utilize copper sulfate. It is used in Fehling's solution and Benedict's solution to test for reducing sugars, 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 also used to test blood for anemia. The blood is tested by dropping it into a solution of copper sulfate of known specific gravity – blood which contains sufficient hemoglobin sinks rapidly due to its density, whereas blood which does not floats or sinks less rapidly.[11]
In a flame test, its copper ions emit a deep green light, a much deeper green than the flame test for barium.
In the presence of chlorine, copper ions emit a deep blue light.
Organic synthesis
Copper sulfate is employed in organic synthesis.[12] The anhydrous salt catalyses the transacetylation in organic synthesis.[13] The hydrated salt can be intimately mingled with potassium permanganate to give an oxidant for the conversion of primary alcohols.[14]
Chemistry education
Copper sulfate is a commonly included chemical in children's chemistry sets and is often used to grow crystals in schools and in copper plating experiments. Because of its toxicity, it is not recommended for small children. Copper sulfate is often used to demonstrate an exothermic reaction, in which steel wool or magnesium ribbon is placed in an aqueous solution of CuSO4. It is used in school chemistry courses 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, and is known as blue copperas.[15] Copper(II) sulfate pentahydrate can easily be produced by crystallization from solution as copper(II) sulfate is quite hygroscopic.
In an illustration of a "single metal replacement reaction", iron is submerged in a solution of copper sulfate. Upon standing, iron reacts, producing iron(II) sulfate, and copper precipitates.
- Fe + CuSO4 → FeSO4 + Cu
Medical
Copper sulfate was also used in the past as an emetic.[16] It is now considered too toxic for this use.[17] It is still listed as an antidote in the World Health Organization's Anatomical Therapeutic Chemical Classification System.[18]
Art
In 2008, the artist Roger Hiorns filled an abandoned waterproofed council flat in London with 75,000 liters of copper sulfate 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.[19]
Etching
Copper sulfate is also used to etch zinc plates for intaglio printmaking.[20][21]
Dyeing
Copper sulfate can also be used as a mordant in vegetable dyeing. It often highlights the green tints of the specific dyes.
Toxicological effects
Copper sulfate is an irritant.[22] The usual routes by which humans can receive toxic exposure to copper sulfate are through eye or skin contact, as well as by inhaling powders and dusts.[23] Skin contact may result in itching or eczema.[24] Eye contact with copper sulfate can cause conjunctivitis, inflammation of the eyelid lining, ulceration, and clouding of the cornea.[25]
Upon acute oral exposure, copper sulfate is only moderately toxic.[23] According to studies, the lowest dose of copper sulfate that had a toxic impact on humans is 11 mg/kg.[26] Because of its irritating effect on the gastrointestinal tract, vomiting is automatically triggered in case of the ingestion of copper sulfate. However, if copper sulfate is retained in the stomach, the symptoms can be severe. After 1–12 grams of copper sulfate are swallowed, such poisoning signs may occur as a metallic taste in the mouth, burning pain in the chest, nausea, diarrhea, vomiting, headache, discontinued urination, which leads to yellowing of the skin. In cases of copper sulfate poisoning, injury to the brain, stomach, liver, or kidneys may also occur.[25]
References
- ^ 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. doi:10.1524/zkri.1975.141.5-6.330.
- ^ "Copper(II) sulfate MSDS". Oxford University. Retrieved 2007-12-31.
- ^ Andrew Knox Galwey; Michael E. Brown (1999). Thermal decomposition of ionic solids. Elsevier. pp. 228–229. ISBN 0-444-82437-5.
- ^ Wiberg, Egon (2001). Inorganic chemistry. Academic Press. p. 1263. ISBN 0-12-352651-5.
{{cite book}}: Unknown parameter|coauthors=ignored (|author=suggested) (help) - ^ Holleman, A. F.; Wiberg, E. (2001). Inorganic Chemistry. San Diego: Academic Press. ISBN 0-12-352651-5.
- ^ Johnson, George Fiske (1935). "The Early History of Copper Fungicides". Agricultural History. 9 (2): 67–79. JSTOR 3739659.
- ^ Parry, K. E.; Wood, R. K. S. (1958). "The adaption of fungi to fungicides: Adaption to copper and mercury salts". Annals of Applied Biology. 46 (3): 446. doi:10.1111/j.1744-7348.1958.tb02225.x.
- ^ "Uses of Copper Compounds: Copper Sulfate's Role in Agriculture". Copper.org. Retrieved 2007-12-31.
- ^ "All About Copper Sulfate". National Fish Pharmaceuticals. Retrieved 2007-12-31.
- ^ Estridge, Barbara H. (2000). Basic Medical Laboratory Techniques. Thomson Delmar Learning. p. 166. ISBN 0-7668-1206-5.
{{cite book}}: Unknown parameter|coauthors=ignored (|author=suggested) (help) - ^ Hoffman, R. V. (2001). Copper(II) Sulfate, in Encyclopedia of Reagents for Organic Synthesis. John Wiley & Sons. doi:10.1002/047084289X.rc247.
- ^ Hulce, M.; Mallomo, J. P.; Frye, L. L.; Kogan, T. P.; Posner, G. H. (1990). "(S)-( + )-2-(p-Toluenesulfinyl)-2-Cyclopentanone: Precursor for Enantioselective Synthesis of 3-Substituted Cyclopentanones". Organic Syntheses
{{cite journal}}: CS1 maint: multiple names: authors list (link); Collected Volumes, vol. 7, p. 495. - ^ 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.
- ^ "Process for the preparation of stable copper (II) sulfate monohydrate applicable as trace element additive in animal fodders". Retrieved 2009-07-07.
- ^ Holtzmann, N. A.; Haslam, R. H. (1968). "Elevation of serum copper following copper sulfate as an emetic". Pediatrics. 42 (1): 189–93. PMID 4385403.
{{cite journal}}: Unknown parameter|month=ignored (help) - ^ Olson, Kent C. (2004). Poisoning & drug overdose. New York: Lange Medical Mooks/McGraw-Hill. p. 175. ISBN 0-8385-8172-2.
- ^ V03AB20 (WHO)
- ^ "Seizure homepage". Artangel.org.uk. Retrieved 2009-09-21.
- ^ Bordeau etch. Greenart.info (2009-01-18). Retrieved 2011-06-02.
- ^ The Chemistry of using Copper Sulfate Mordant. Ndiprintmaking.ca (2009-04-12). Retrieved 2011-06-02.
- ^ Windholz, M., ed. 1983. The Merck Index. Tenth edition. Rahway, NJ: Merck and Company.
- ^ a b 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
- ^ 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.
- ^ a b Clayton, G. D. and F. E. Clayton, eds. 1981. Patty's industrial hygiene and toxicology. Third edition. Vol. 2, Part 6 Toxicology. NY: John Wiley and Sons. ISBN 0-471-01280-7
- ^ National Institute for Occupational Safety and Health (NIOSH). 1981–1986. Registry of toxic effects of chemical substances (RTECS). Cincinnati, OH: NIOSH.
External links
