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Copper(II) nitrate

(Redirected from Copper(II) nitrate - trihydrate)
Not to be confused with Copper(I) nitrate.

Copper(II) nitrate describes any member of the family of inorganic compounds with the formula Cu(NO3)2(H2O)x. The hydrates are blue solids. Anhydrous copper nitrate forms blue-green crystals and sublimes in a vacuum at 150-200 °C.[5][6] Common hydrates are the hemipentahydrate and trihydrate.

Copper(II) nitrate
alpha polymorph[1]
beta polymorph[2]
Names
IUPAC name
Copper(II) nitrate
Other names
Cupric nitrate
Identifiers
3D model (JSmol)
ChEBI
ChemSpider
ECHA InfoCard 100.019.853 Edit this at Wikidata
RTECS number
  • GL7875000
UNII
  • InChI=1S/Cu.2NO3/c;2*2-1(3)4/q+2;2*-1 checkY
    Key: XTVVROIMIGLXTD-UHFFFAOYSA-N checkY
  • InChI=1/Cu.2NO3/c;2*2-1(3)4/q+2;2*-1
    Key: XTVVROIMIGLXTD-UHFFFAOYAG
  • [Cu+2].[O-][N+]([O-])=O.[O-][N+]([O-])=O
Properties
Cu(NO3)2
Molar mass 187.5558 g/mol (anhydrous)
241.60 g/mol (trihydrate)
232.591 g/mol (hemipentahydrate)
Appearance blue crystals
hygroscopic
Density 3.05 g/cm3 (anhydrous)
2.32 g/cm3 (trihydrate)
2.07 g/cm3 (hexahydrate)
Melting point 114 °C (237 °F; 387 K) (anhydrous, decomposes)
114.5 °C (trihydrate)
26.4 °C (hexahydrate, decomposes)
Boiling point 170 °C (338 °F; 443 K) (trihydrate, decomposes)
trihydrate:[3]
381 g/100 mL (40 °C)
666 g/100 mL (80 °C)
hexahydrate:[3]
243.7 g/100 mL (80 °C)
Solubility hydrates very soluble in ethanol, ammonia, water; insoluble in ethyl acetate
+1570.0·10−6 cm3/mol (~3H2O)
Structure
orthorhombic (anhydrous)
rhombohedral (hydrates)
Hazards
Occupational safety and health (OHS/OSH):
Main hazards
 Irritant, Oxidizer
NFPA 704 (fire diamond)
NFPA 704 four-colored diamondHealth 1: Exposure would cause irritation but only minor residual injury. E.g. turpentineFlammability 0: Will not burn. E.g. waterInstability 3: Capable of detonation or explosive decomposition but requires a strong initiating source, must be heated under confinement before initiation, reacts explosively with water, or will detonate if severely shocked. E.g. hydrogen peroxideSpecial hazard OX: Oxidizer. E.g. potassium perchlorate
1
0
3
OX
NIOSH (US health exposure limits):
PEL (Permissible)
 TWA 1 mg/m3 (as Cu)[4]
REL (Recommended)
 TWA 1 mg/m3 (as Cu)[4]
IDLH (Immediate danger)
 TWA 100 mg/m3 (as Cu)[4]
Safety data sheet (SDS) Cu(NO3)2·3H2O
Related compounds
Other anions
 Copper(II) sulfate
Copper(II) chloride
Other cations
 Silver nitrate
Gold(III) nitrate
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
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Synthesis and reactions

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Hydrated copper(II) nitrate

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Hydrated copper nitrate is prepared by treating copper metal or its oxide with nitric acid:[7]

Cu + 4 HNO3 → Cu(NO3)2 + 2 H2O + 2 NO2

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 [Cu(H2O)6]2+. 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 Cu(NO3)2.[6] At 80 °C the hydrates convert to "basic copper nitrate", Cu2(NO3)(OH)3, which converts to CuO at 180 °C.[7] 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:

2 Cu(NO3)2 → 2 CuO + 4 NO2 + O2
3 NO2 + H2O → 2 HNO3 + NO

Treatment of copper(II) nitrate solutions with triphenylphosphine, triphenylarsine, and triphenylstibine gives the corresponding copper(I) complexes [Cu(EPh3)3]NO3 (E = P, As, Sb; Ph = C6H5). The group V ligand is oxidized to the oxide.[8]

Anhydrous copper(II) nitrate

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Anhydrous Cu(NO3)2 is one of the few anhydrous transition metal nitrates.[9] It cannot be prepared by reactions containing or producing water. Instead, anhydrous Cu(NO3)2 forms when copper metal is treated with dinitrogen tetroxide:[6]

Cu + 2 N2O4 → Cu(NO3)2 + 2 NO

Structure

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Anhydrous copper(II) nitrate

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Structure of anhydrous copper(II) nitrate in the gas phase.[6]

Two polymorphs of anhydrous copper(II) nitrate, α and β, are known.[6] 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,[1] but the β form has two different copper centers, one with [4+1] and one that is square planar.[2]

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.[10]

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.[6][11] In the vapour phase, the molecule features two bidentate nitrate ligands.[12]

Hydrated copper(II) nitrate

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Five hydrates have been reported: the monohydrate (Cu(NO3)2·2H2O),[2] the sesquihydrate (Cu(NO3)2·1.5H2O),[13] the hemipentahydrate (Cu(NO3)2·2.5H2O),[14] a trihydrate (Cu(NO3)2·3H2O),[15] and a hexahydrate ([Cu(OH2)6](NO3)2.[16] 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.

Applications

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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.[7] 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 synthesis

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Copper nitrate, in combination with acetic anhydride, is an effective reagent for nitration of aromatic compounds, known as the Menke nitration.[17] 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.[18] A related reagent based on montmorillonite has proven useful for the nitration of aromatic compounds.[19]

Electrowinning

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Copper(II) nitrate may also be used for copper electrowinning on small scale with a ammonia (NH3) as a byproduct.[20]

Naturally occurring copper nitrates

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No mineral of the ideal Cu(NO3) formula, or the hydrates, are known. Likasite, Cu3(NO3)(OH)5·2H2O and buttgenbachite, Cu19(NO3)2(OH)32Cl4·2H2O are related minerals.[21][22]

Natural basic copper nitrates include the rare minerals gerhardtite and rouaite, both being polymorphs of Cu2(NO3)(OH)3.[23][24][25] A much more complex, basic, hydrated and chloride-bearing natural salt is buttgenbachite.[22][25]

References

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  1. ^ a b Wallwork, S. C.; Addison, W. E. (1965). "526. The crystal structures of anhydrous nitrates and their complexes. Part I. The α form of copper(II) nitrate". J. Chem. Soc. 1965: 2925–2933. doi:10.1039/JR9650002925.
  2. ^ a b c Troyanov, S. I.; Morozov, I. V.; Znamenkov, K. O.; Yu; Korenev, M. (1995). "Synthesis and X-Ray Structure of New Copper(II) Nitrates: Cu(NO3)2·H2O and β-modification of Cu(NO3)2". Z. Anorg. Allg. Chem. 621 (7): 1261–1265. doi:10.1002/zaac.19956210727.
  3. ^ a b Perrys' Chem Eng Handbook, 7th Ed
  4. ^ a b c NIOSH Pocket Guide to Chemical Hazards. "#0150". National Institute for Occupational Safety and Health (NIOSH).
  5. ^ Pass and Sutcliffe (1968). Practical Inorganic Chemistry. London: Chapman and Hall.
  6. ^ a b c d e f Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. p. 1190. ISBN 978-0-08-037941-8.
  7. ^ a b c H.Wayne Richardson "Copper Compounds" Ullmann's Encyclopedia of Industrial Chemistry 2005, Wiley-VCH, Weinheim. doi:10.1002/14356007.a07_567.
  8. ^ Gysling, Henry J. (1979). "Coordination Complexes of Copper(I) Nitrate". Inorganic Syntheses. Inorganic Syntheses. Vol. 19. pp. 92–97. doi:10.1002/9780470132500.ch19. ISBN 9780470132500.
  9. ^ Addison, C. C.; Logan, N.; Wallwork, S. C.; Garner, C. D. (1971). "Structural Aspects of Co-ordinated Nitrate Groups". Quarterly Reviews, Chemical Society. 25 (2): 289. doi:10.1039/qr9712500289.
  10. ^ Duffin, B.; Wallwork, S. C. (1966). "The crystal structure of anhydrous nitrates and their complexes. II. The 1:1 copper(II) nitrate-nitromethane complex". Acta Crystallographica. 20 (2): 210–213. doi:10.1107/S0365110X66000434.
  11. ^ Addison, C. C.; Hathaway, B. J. (1958). "628. The vapour pressure of anhydrous copper nitrate, and its molecular weight in the vapour state". J. Chem. Soc.: 3099–3106. doi:10.1039/JR9580003099.
  12. ^ LaVilla, R. E.; Bauer, S. H. (1963). "The Structure of Gaseous Copper(II) Nitrate as Determined by Electron Diffraction". J. Am. Chem. Soc. 85 (22): 3597–3600. doi:10.1021/ja00905a015.
  13. ^ Dornberger-Schiff, K.; Leciejewicz, J. (1958). "Zur Struktur des Kupfernitrates Cu(NO3)2.1.5H2O". Acta Crystallogr. 11 (11): 825–826. doi:10.1107/S0365110X58002322.
  14. ^ Morosin, B. (1970). "The crystal structure of Cu(NO3)2.2.5H2O". Acta Crystallographica Section B. 26 (9): 1203–1208. doi:10.1107/S0567740870003898.
  15. ^ J. Garaj, Sbornik Prac. Chem.-Technol. Fak. Svst., Cskosl. 1966, pp. 35–39.
  16. ^ Zibaseresht, R.; Hartshorn, R. M. (2006). "Hexaaquacopper(II) dinitrate: absence of Jahn-Teller distortion". Acta Crystallographica Section E. 62: i19–i22. doi:10.1107/S1600536805041851.
  17. ^ Menke J.B. (1925). "Nitration with nitrates". Recueil des Travaux Chimiques des Pays-Bas. 44: 141. doi:10.1002/recl.19250440209.
  18. ^ Balogh, M. "Copper(II) Nitrate–K10 Bentonite Clay" in Encyclopedia of Reagents for Organic Synthesis (Ed: L. Paquette) 2004, J. Wiley & Sons, New York. doi:10.1002/047084289X.
  19. ^ Collet, Christine (1990). "Clays Direct Aromatic Nitration". Angewandte Chemie International Edition in English. 29 (5): 535–536. doi:10.1002/anie.199005351.
  20. ^ Oishi, Tetsuo; Koyama, Kazuya; Konishi, Hirokazu; Tanaka, Mikiya; Lee, Jae-Chun (November 2007). "Influence of ammonium salt on electrowinning of copper from ammoniacal alkaline solutions". Electrochimica Acta. 53 (1): 127–132. doi:10.1016/j.electacta.2007.06.024.
  21. ^ "Likasite". www.mindat.org.
  22. ^ a b "Buttgenbachite". www.mindat.org.
  23. ^ "Gerhardtite". www.mindat.org.
  24. ^ "Rouaite". www.mindat.org.
  25. ^ a b International Mineralogical Association (21 March 2011). "List of Minerals". www.ima-mineralogy.org.
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