Copper(II) nitrate

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Copper(II) nitrate
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Trihydrate
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alpha polymorph[1]
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beta polymorph[2]
Names
IUPAC name
Copper(II) nitrate
Other names
Cupric nitrate
Identifiers
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3D model (JSmol)
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  • 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
Page Module:Chem2/styles.css has no content.Cu(NO3)2
Molar mass Page Template:Plainlist/styles.css has no content.
  • 187.5558 g/mol (anhydrous)
  • 241.60 g/mol (trihydrate)
  • 232.591 g/mol (hemipentahydrate)
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  • blue-green crystals (anhydrous)
  • blue crystals (hydrate)
Density Page Template:Plainlist/styles.css has no content.
  • 3.05 g/cm3 (anhydrous)
  • 2.32 g/cm3 (trihydrate)
  • 2.07 g/cm3 (hexahydrate)
Melting point [citation needed] Page Template:Plainlist/styles.css has no content.
  • 114 °C (237 °F; 387 K) (anhydrous)
  • 80 °C (176 °F; 353 K) (trihydrate, decomposes)
  • 26.4 °C (79.5 °F; 299.5 K) (hexahydrate, decomposes)
Boiling point 170 °C (338 °F; 443 K)[citation needed] (trihydrate, decomposes)
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  • trihydrate:[3]
  • 381 g/100 mL (40 °C (104 °F; 313 K))
  • 666 g/100 mL (80 °C (176 °F; 353 K))
  • hexahydrate:
  • 243.7 g/100 mL (80 °C (176 °F; 353 K))
Solubility in ethanol hydrates very soluble
Solubility in ammonia hydrates very soluble
1570.0×10−6 cm3/mol (trihydrate)
Structure
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  • orthorhombic (anhydrous)
  • rhombohedral (hydrates)
Hazards
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[5]

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3
3
3
OX
NIOSH (US health exposure limits):[6]
PEL (Permissible)
 1 mg/m3 (TWA, as Cu)
REL (Recommended)
 1 mg/m3 (TWA, as Cu)
IDLH (Immediate danger)
 100 mg/m3 (TWA, as Cu)
Related compounds
Other anions
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Other cations
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Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).

Template:Chembox Footer/trackingTemplate:Short description

Copper(II) nitrate describes any member of the family of inorganic compounds with the formula Page Module:Chem2/styles.css has no content.Cu(NO3)2·x(H2O). The hydrates are hygroscopic blue solids. Anhydrous copper nitrate forms blue-green crystals and sublimes in a vacuum at 150–200 °C (302–392 °F).[7][8] Common hydrates are the hemipentahydrate and trihydrate.

Occurrence

No mineral of the ideal Page Module:Chem2/styles.css has no content.Cu(NO3) formula, or the hydrates, are known.

Likasite [it], Page Module:Chem2/styles.css has no content.Cu3(NO3)(OH)5·2H2O and buttgenbachite [it], Page Module:Chem2/styles.css has no content.Cu19(NO3)2(OH)32Cl4·2H2O are related minerals. Natural basic copper nitrates include the rare minerals gerhardtite and rouaite [it], both being polymorphs of Page Module:Chem2/styles.css has no content.Cu2(NO3)(OH)3. A much more complex, basic, hydrated and chloride-bearing natural salt is buttgenbachite.[9][10][11][12]

Structure

Anhydrous copper(II) nitrate

File:Copper(II)-nitrate-monomer-2D-dimensions.png
Structure of anhydrous copper(II) nitrate in the gas phase.[8]

Two polymorphs of anhydrous copper(II) nitrate, α and β, are known. 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, but the β form has two different copper centers, one with [4+1] and one that is square planar.[8][1][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.[13]

Heating solid anhydrous copper(II) nitrate under a vacuum to 150–200 °C (302–392 °F) leads to sublimation and cracking to give a vapour of monomeric copper(II) nitrate molecules. In the vapour phase, the molecule features two bidentate nitrate ligands.[8][14][15]

Hydrated copper(II) nitrate

Five hydrates have been reported: the monohydrate (Page Module:Chem2/styles.css has no content.Cu(NO3)2·2H2O), the sesquihydrate (Page Module:Chem2/styles.css has no content.2Cu(NO3)2·3H2O), the hemipentahydrate (Page Module:Chem2/styles.css has no content.Cu(NO3)2·2.5H2O), a trihydrate (Page Module:Chem2/styles.css has no content.Cu(NO3)2·3H2O),[citation needed] and a hexahydrate (Page Module:Chem2/styles.css has no content.[Cu(OH2)6](NO3)2.[2][16][17][18]

The crystal structure of the hexahydrate appeared to show six almost equal Page Module:Chem2/styles.css has no content.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 Page Module:Chem2/styles.css has no content.Cu−O bonds but it is probably due to nickel being misidentified as copper in the refinement.[citation needed]

Synthesis and reactions

Hydrated copper(II) nitrate

Hydrated copper nitrate is prepared by treating copper metal or its oxide with nitric acid:[19]

Page Module:Chem2/styles.css has no content.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.[citation needed]

In aqueous solution, the hydrates exist as the aqua complex Page Module:Chem2/styles.css has no content.[Cu(H2O)6]2+. Such complexes are highly labile and subject to rapid ligand exchange due to the d9 electronic configuration of copper(II).[citation needed]

Attempted dehydration of any of the hydrated copper(II) nitrates by heating affords the oxides, not Page Module:Chem2/styles.css has no content.Cu(NO3)2.[8] At 80 °C (176 °F) the hydrates convert to "basic copper nitrate", Page Module:Chem2/styles.css has no content.Cu2(NO3)(OH)3, which converts to Page Module:Chem2/styles.css has no content.CuO at 180 °C (356 °F).[19] 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:[citation needed]

Page Module:Chem2/styles.css has no content.2 Cu(NO3)2 → 2 CuO + 4 NO2 + O2
Page Module:Chem2/styles.css has no content.3 NO2 + H2O → 2 HNO3 + NO

Treatment of copper(II) nitrate solutions with triphenylphosphine, triphenylarsine, and triphenylstibine gives the corresponding copper(I) complexes Page Module:Chem2/styles.css has no content.[Cu(RPh3)3]NO3 (E = P, As, Sb). The group V ligand is oxidized to the oxide.[20]

Anhydrous copper(II) nitrate

Anhydrous Page Module:Chem2/styles.css has no content.Cu(NO3)2 is one of the few anhydrous transition metal nitrates. It cannot be prepared by reactions containing or producing water. Instead, anhydrous Page Module:Chem2/styles.css has no content.Cu(NO3)2 forms when copper metal is treated with dinitrogen tetroxide:[21][8]

Page Module:Chem2/styles.css has no content.Cu + 2 N2O4 → Cu(NO3)2 + 2 NO

Applications

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.[19] It is often used in school laboratories to demonstrate chemical voltaic cell reactions.[citation needed] It is a component in some ceramic glazes and metal patinas.[citation needed]

Organic synthesis

Copper nitrate, in combination with acetic anhydride, is an effective reagent for nitration of aromatic compounds, known as the Menke nitration.[22]

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.[23] A related reagent based on montmorillonite has proven useful for the nitration of aromatic compounds.[24]

Electrowinning

Copper(II) nitrate may also be used for copper electrowinning on small scale with a ammonia (Page Module:Chem2/styles.css has no content.NH3) as a byproduct.[25]

References

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  1. ^ a b Page Module:Citation/CS1/styles.css has no content.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 Page Module:Citation/CS1/styles.css has no content.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. ^ Page Module:Citation/CS1/styles.css has no content.Perry's chemical engineers' handbook (7th ed.). New York: McGraw-Hill. 1997. ISBN 0070498415.[page needed]
  4. ^ Sigma-Aldrich Co., Copper(II) nitrate.
  5. ^ Page Module:Citation/CS1/styles.css has no content."SDS - Copper(II) nitrate". fishersci.com. ThermoFisher Scientific. 24 December 2021. Retrieved 6 November 2025.
  6. ^ Page Module:Citation/CS1/styles.css has no content."NIOSH Pocket Guide to Chemical Hazards".
  7. ^ Page Module:Citation/CS1/styles.css has no content.Pass and Sutcliffe (1968). Practical Inorganic Chemistry. London: Chapman and Hall.
  8. ^ a b c d e f Page Module:Citation/CS1/styles.css has no content.Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. p. 1190. doi:10.1016/C2009-0-30414-6. ISBN 978-0-08-037941-8.
  9. ^ Page Module:Citation/CS1/styles.css has no content."Buttgenbachite". www.mindat.org.
  10. ^ Page Module:Citation/CS1/styles.css has no content."Gerhardtite". www.mindat.org.
  11. ^ Page Module:Citation/CS1/styles.css has no content."Rouaite". www.mindat.org.
  12. ^ Page Module:Citation/CS1/styles.css has no content."Likasite". www.mindat.org.
  13. ^ Page Module:Citation/CS1/styles.css has no content.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.
  14. ^ Page Module:Citation/CS1/styles.css has no content.Addison, C. C.; Hathaway, B. J. (1958). "628. The vapour pressure of anhydrous copper nitrate, and its molecular weight in the vapour state". Journal of the Chemical Society: 3099–3106. doi:10.1039/JR9580003099.
  15. ^ Page Module:Citation/CS1/styles.css has no content.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.
  16. ^ Page Module:Citation/CS1/styles.css has no content.Dornberger-Schiff, K.; Leciejewicz, J. (1958). "Zur Struktur des Kupfernitrates Cu(NO3)2.1.5H2O". Acta Crystallographica. 11 (11): 825–826. doi:10.1107/S0365110X58002322.
  17. ^ Page Module:Citation/CS1/styles.css has no content.Morosin, B. (1970). "The crystal structure of Cu(NO3)2.2.5H2O". Acta Crystallographica Section B. 26 (9): 1203–1208. doi:10.1107/S0567740870003898.
  18. ^ Page Module:Citation/CS1/styles.css has no content.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.
  19. ^ a b c Script error: No such module "Template wrapper"..
  20. ^ Page Module:Citation/CS1/styles.css has no content.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.
  21. ^ Page Module:Citation/CS1/styles.css has no content.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.
  22. ^ Page Module:Citation/CS1/styles.css has no content.Menke J.B. (1925). "Nitration with nitrates". Recueil des Travaux Chimiques des Pays-Bas. 44: 141. doi:10.1002/recl.19250440209.
  23. ^ Page Module:Citation/CS1/styles.css has no content.Balogh, M. (15 April 2001). "Copper(II) Nitrate–K10 Bentonite Clay". In Paquette, L. (ed.). Encyclopedia of Reagents for Organic Synthesis. New York: J. Wiley & Sons. doi:10.1002/047084289X.
  24. ^ Page Module:Citation/CS1/styles.css has no content.Collet, Christine (1990). "Clays Direct Aromatic Nitration". Angewandte Chemie International Edition in English. 29 (5): 535–536. doi:10.1002/anie.199005351.
  25. ^ Page Module:Citation/CS1/styles.css has no content.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.

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