Copper(II) nitrate
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Copper(II) nitrate describes any member of the family of inorganic compounds with the formula Cu(NO3)2(H2O)x. The hydrates are hygroscopic blue solids. Anhydrous copper nitrate forms blue-green crystals and sublimes in a vacuum at 150-200 °C.<ref>Template:Cite book</ref><ref name="G&E">Template:Greenwood&Earnshaw2nd</ref> Common hydrates are the hemipentahydrate and trihydrate.
Synthesis and reactionsEdit
Hydrated copper(II) nitrateEdit
Hydrated copper nitrate is prepared by treating copper metal or its oxide with nitric acid:<ref name=Ullmann/>
The same salts can be prepared treating copper metal with an aqueous solution of silver nitrate. That reaction illustrates the ability of copper metal to reduce silver ions.
In aqueous solution, the hydrates exist as the aqua complex Template:Chem2. Such complexes are highly labile and subject to rapid ligand exchange due to the d9 electronic configuration of copper(II).
Attempted dehydration of any of the hydrated copper(II) nitrates by heating affords the oxides, not Template:Chem2.<ref name="G&E" /> At 80 °C the hydrates convert to "basic copper nitrate", Template:Chem2, which converts to Template:Chem2 at 180 °C.<ref name=Ullmann/> Exploiting this reactivity, copper nitrate can be used to generate nitric acid by heating it until decomposition and passing the fumes directly into water. This method is similar to the last step in the Ostwald process. The equations are as follows:
Treatment of copper(II) nitrate solutions with triphenylphosphine, triphenylarsine, and triphenylstibine gives the corresponding copper(I) complexes Template:Chem2 (E = P, As, Sb; Ph = Template:Chem2). The group V ligand is oxidized to the oxide.<ref>Template:Cite book</ref>
Anhydrous copper(II) nitrateEdit
Anhydrous Template:Chem2 is one of the few anhydrous transition metal nitrates.<ref>Template:Cite journal</ref> It cannot be prepared by reactions containing or producing water. Instead, anhydrous Template:Chem2 forms when copper metal is treated with dinitrogen tetroxide:<ref name="G&E" />
StructureEdit
Anhydrous copper(II) nitrateEdit
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Two polymorphs of anhydrous copper(II) nitrate, α and β, are known.<ref name="G&E" /> Both polymorphs are three-dimensional coordination polymer networks with infinite chains of copper(II) centers and nitrate groups. The α form has only one Cu environment, with [4+1] coordination,<ref name="Wallwork&Addison">Template:Cite journal</ref> but the β form has two different copper centers, one with [4+1] and one that is square planar.<ref name=Troyanov>Template:Cite journal</ref>
The nitromethane solvate also features "[4+1] coordination", with four short Cu-O bonds of approximately 200 pm and one longer bond at 240 pm.<ref>Template:Cite journal</ref>
Heating solid anhydrous copper(II) nitrate under a vacuum to 150-200 °C leads to sublimation and "cracking" to give a vapour of monomeric copper(II) nitrate molecules.<ref name="G&E" /><ref>Template:Cite journal</ref> In the vapour phase, the molecule features two bidentate nitrate ligands.<ref>Template:Cite journal</ref>
Hydrated copper(II) nitrateEdit
Five hydrates have been reported: the monohydrate (Template:Chem2),<ref name=Troyanov/> the sesquihydrate (Template:Chem2),<ref>Template:Cite journal</ref> the hemipentahydrate (Template:Chem2),<ref>Template:Cite journal</ref> a trihydrate (Template:Chem2),<ref>J. Garaj, Sbornik Prac. Chem.-Technol. Fak. Svst., Cskosl. 1966, pp. 35–39.</ref> and a hexahydrate (Template:Chem2.<ref>Template:Cite journal</ref> The crystal structure of the hexahydrate appeared to show six almost equal Cu–O distances, not revealing the usual effect of a Jahn-Teller distortion that is otherwise characteristic of octahedral Cu(II) complexes. This non-effect was attributed to the strong hydrogen bonding that limits the elasticity of the Cu-O bonds but it is probably due to nickel being misidentified as copper in the refinement.
ApplicationsEdit
Copper(II) nitrate finds a variety of applications, the main one being its conversion to copper(II) oxide, which is used as catalyst for a variety of processes in organic chemistry. Its solutions are used in textiles and polishing agents for other metals. Copper nitrates are found in some pyrotechnics.<ref name=Ullmann>H.Wayne Richardson "Copper Compounds" Ullmann's Encyclopedia of Industrial Chemistry 2005, Wiley-VCH, Weinheim. {{#invoke:doi|main}}.</ref> It is often used in school laboratories to demonstrate chemical voltaic cell reactions. It is a component in some ceramic glazes and metal patinas.
Organic synthesisEdit
Copper nitrate, in combination with acetic anhydride, is an effective reagent for nitration of aromatic compounds, known as the Menke nitration.<ref>Template:Cite journal</ref> Hydrated copper nitrate adsorbed onto clay affords a reagent called "Claycop". The resulting blue-colored clay is used as a slurry, for example for the oxidation of thiols to disulfides. Claycop is also used to convert dithioacetals to carbonyls.<ref>Balogh, M. "Copper(II) Nitrate–K10 Bentonite Clay" in Encyclopedia of Reagents for Organic Synthesis (Ed: L. Paquette) 2004, J. Wiley & Sons, New York. {{#invoke:doi|main}}.</ref> A related reagent based on montmorillonite has proven useful for the nitration of aromatic compounds.<ref>Template:Cite journal</ref>
ElectrowinningEdit
Copper(II) nitrate may also be used for copper electrowinning on small scale with a ammonia (NH3) as a byproduct.<ref>Template:Cite journal</ref>
Naturally occurring copper nitratesEdit
No mineral of the ideal Template:Chem2 formula, or the hydrates, are known. Likasite, Template:Chem2 and buttgenbachite, Template:Chem2 are related minerals.<ref name=min2399>{{#invoke:citation/CS1|citation |CitationClass=web }}</ref><ref name=min811/>
Template:Anchor Natural basic copper nitrates include the rare minerals gerhardtite and rouaite, both being polymorphs of Template:Chem2.<ref>{{#invoke:citation/CS1|citation |CitationClass=web }}</ref><ref>{{#invoke:citation/CS1|citation |CitationClass=web }}</ref><ref name=IMA>{{#invoke:citation/CS1|citation |CitationClass=web }}</ref> A much more complex, basic, hydrated and chloride-bearing natural salt is buttgenbachite.<ref name=min811>{{#invoke:citation/CS1|citation |CitationClass=web }}</ref><ref name=IMA/>