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Copper(II) acetate

Copper(II) acetate
Copper(II) acetate hydrate
IUPAC name
Other names Cupric acetate
Molecular formula Cu(CH3COO)2
Molar mass 182 g/mol
Appearance Dark green crystalline solid
Density 1.88 g/mL
Melting point

115 °C (388 K)

Boiling point

240 °C (513 K)

Solubility in other solvents 7.2 g/100 mL cold water
20 g / 100 mL hot water
Soluble in alcohol
Slightly soluble in ether and glycerol
Crystal structure Monoclinic
R-phrases 22-36/37/38-50/53
S-phrases 26-60-61
NFPA 704
Flash point Non-flammable
Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa)
Infobox references

Copper(II) acetate, also referred to as cupric acetate, is the chemical compound with the formula Cu(OAc)2 where AcO- is acetate (CH3CO2-). The hydrated derivative, which contains one molecule of water for each Cu atom, is available commercially. Cu(OAc)2 is a dark green crystalline solid, whereas Cu(OAc)2(H2O)2 is more bluish-green. Since ancient times, copper acetates of some form have been used as fungicides and green pigments. Today, Cu(OAc)2 is used as a source of copper(II) in inorganic synthesis and as a catalyst or an oxidizing agent in organic synthesis. Copper acetate, like all copper compounds, emits a blue-green glow in a flame.



The dinuclear structure of copper(II) acetate dihydrate

Copper acetate monohydrate is described by the formula Cu2(OAc)4(H2O)2. It adopts the "Chinese lantern" structure seen also for related Rh(II) and Cr(II) tetraacetates.[1][2] One oxygen atom on each acetate is bound to one copper at 1.97 Å (197 pm. Completing the coordination sphere are two water ligands, with Cu-O distances of 2.20 Å (220 pm). The two five-coordinate copper atoms are separated by only 2.65 Å (265 pm), which is close to the Cu--Cu separation in metallic copper. The two copper centers interact resulting in a diminishing of the magnetic moment such that near 90 K, Cu2(OAc)4(H2O)2 is essentially diamagnetic due to cancellation of the two opposing spins. Cu2(OAc)4(H2O)2 was a critical step in the development of modern theories for antiferromagnetic coupling.[3]


Copper(II) acetate has been synthesized for centuries by the method described in the history section. This method, however, leads to an impure copper(II) acetate. In a laboratory, a much purer form can be synthesized in a simple three-step procedure. The overall reaction is as follows:

2 CuSO4.5H2O + 4 NH3 + 4 CH3COOH → Cu2(OAc)4(H2O)2 + 2 [NH4]2[SO4] + 8 H2O

The hydrate form can be dehydrated by heating at 100 °C in a vacuum:[4]

Cu2(OAc)4(H2O)2 → Cu2(OAc)4 + 2 H2O

Heating a mixture of anhydrous Cu2(OAc)4 and copper metal affords colorless, volatile cuprous acetate:[5]

2 Cu + Cu2(OAc)4 → 4 CuOAc

Uses in chemical synthesis

The uses for copper(II) acetate are more plentiful as a catalyst or oxidizing agent in organic syntheses. For example, Cu2(OAc)4 is used to couple two terminal alkynes to make a 1,3-diyne:[6]

Cu2(OAc)4 + 2 RC≡CH → 2 CuOAc + RC≡C-C≡CR + 2 HOAc

The reaction proceeds via the intermediacy of copper(I) acetylides, which are then oxidized by the copper(II) acetate, releasing the acetylide radical. A related reaction involving copper acetylides is the synthesis of ynamines, terminal alkynes with amine groups using Cu2(OAc)4.


Copper(II) acetate is occasionally the primary component of verdigris, the blue-green substance that forms on copper during long exposures to atmosphere. It was historically prepared in vineyards, since acetic acid is a byproduct of fermentation. Copper sheets were alternately layered with fermented grape skins and dregs left over from wine production and exposed to air. This would leave a blue substance on the outside of the sheet. This was then scraped off and dissolved in water. The resulting solid was used as a pigment, or combined with arsenic trioxide to form copper acetoarsenite, a powerful insecticide and fungicide called Paris Green or Schweinfurt Green.

Copper (cupric) acetate tablets were believed to repel sharks in the mid twentieth century. Scuba divers strapped tablets of the compound to their belt and/or ankles to provide protection against sharks. It was used by Jacques-Yves Cousteau and his researchers with questionable results.[7]


  1. ^ van Niekerk, J. N. Schoening, F. R. L. (1953). "X-Ray Evidence for Metal-to-Metal Bonds in Cupric and Chromous Acetate". Nature 171: 36–37. doi:10.1038/171036a0. 
  2. ^ Wells, A.F. (1984). Structural Inorganic Chemistry, Oxford: Clarendon Press.
  3. ^ R. L. Carlin "Magnetochemistry" Springer: Berlin, 1986
  4. ^ S. J. Kirchner, Q. Fernando (1980). "Copper(I) Acetate". Inorg. Synth. 20: 53–55. doi:10.1002/9780470132517.ch16. 
  5. ^ Parish, E. J.; Kizito, S. A. "Copper(I) Acetate" Encyclopedia of Reagents for Organic Synthesis, 2001 John Wiley & Sons. DOI: 10.1002/047084289X.rc193.
  6. ^ P. Vogel, J. Srogl "Copper(II) Acetate" in "EROS Encyclopedia of Reagents for Organic Synthesis" Copper(II) Acetate, 2005 John Wiley & Sons.
  7. ^ Cousteau, Jacques-Yves; Frédéric Dumas (1953). The Silent World. Harper and Row, Publishers, Inc.. pp. 127–135. ISBN 0-7922-6796-6. 

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