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Copper(II) carbonate

Names
Other names copper carbonate, cupric carbonate, basic copper carbonate
Identifiers
CAS Number	1184-64-1
3D model (JSmol)	Interactive image
ChemSpider	13799
ECHA InfoCard	100.031.909
PubChem CID	14452
RTECS number	FF950000
CompTox Dashboard (EPA)	DTXSID1047077
InChI InChI=1S/CH2O3.Cu/c2-1(3)4;/h(H2,2,3,4);/q;+2/p-2Key: GEZOTWYUIKXWOA-UHFFFAOYSA-L InChI=1/CH2O3.Cu/c2-1(3)4;/h(H2,2,3,4);/q;+2/p-2Key: GEZOTWYUIKXWOA-NUQVWONBAA
SMILES .C()=O
Properties
Chemical formula	CuCO3
Molar mass	123.555 g/mol
Density	3.9 g/cm
Melting point	200 °C
Boiling point	decomposes at 290 °C
Solubility in other solvents	insoluble in water, found effective in melting ice, soluble in ethanoic acid
Except where otherwise noted, data are given for materials in their standard state (at 25 °C , 100 kPa). Y verify (what is  ?) Infobox references

Copper(II) carbonate (often called copper carbonate or cupric carbonate) is a blue-green compound (chemical formula CuCO₃) forming part of the verdigris patina that is found on weathered brass, bronze, and copper. The colour can vary from bright blue to green, because there may be a mixture of both copper carbonate and basic copper carbonate in various stages of hydration. It was formerly much used as a pigment, and is still in use for artist's colours. It has also been used in some types of make-up, like lipstick, although it can also be toxic to humans. It also has been used for many years as an effective algaecide in farm ponds and in aquaculture operations. Copper carbonate was the first compound to be broken down into several, separate elements (copper, carbon, and oxygen). It was broken down in 1794 by the French chemist Joseph Louis Proust (1754–1826). When burnt, it thermally decomposes to form CO₂ and CuO, a black solid. It can be used to copper plate a metallic surface by adding sulfuric acid and heat it (to produce copper sulfate) then passing a charge through it with the metal in the liquid. This process is known as electroplating.

Copper in moist air slowly acquires a dull green coating because its top layer has oxidised with the air. Some architects use this material on rooftops for this interesting colour. The green material is a 1:1 mole mixture of Cu(OH)₂ and CuCO₃:

2 Cu (s) + H₂O (g) + CO₂ + O₂ → Cu(OH)₂ + CuCO₃ (s)

Copper carbonate decomposes at 290 °C, giving off carbon dioxide and leaving copper(II) oxide:

CuCO₃ (s) → CuO (s) + CO₂ (g)

Basic copper(II) carbonate occurs naturally as malachite (Cu₂(OH)₂CO₃) and azurite (Cu₃(OH)₂(CO₃)₂).

Preparation

Copper carbonate is prepared by combining aqueous solutions of copper sulfate and sodium carbonate. Copper basic carbonate precipitates from the solution:

2 CuSO₄ + 2 Na₂CO₃ + H₂O → Cu₂(OH)₂CO₃ + 2 Na₂SO₄ + CO₂

The formation of copper carbonate can be verified in the following steps: a) Centrifuge the above mentioned solution, 1 minute at 6000 g is sufficient b) Wash the precipitate with distilled water and centrifuge again c) The colour of the precipitate is blue, like that of several copper salts but none of the sodium salts d) If dilute (1M) hydrochloric acid is added, then bubbles of CO₂ will emerge and the precipitate will be fully solubilised. These would not be formed if dilute hydrochloric acid was added to solid Na₂SO₄.

Pure copper carbonate is obtained from basic copper carbonate in the presence of carbon dioxide at 180 °C and 4.6 MPa (46 atm) pressure.

Reactions

1) Reaction with acids:

copper carbonate reacts with acids and gives carbon dioxide and corresponding copper(II) salts of that acid:

• copper carbonate + hydrochloric acid → copper chloride + water + carbon dioxide:

CuCO₃ + 2HCl → CuCl₂ + H₂O + CO₂

• copper carbonate + sulfuric acid → copper sulfate + water + carbon dioxide:

CuCO₃ + H₂SO₄ → CuSO₄ + H₂O + CO₂

• copper carbonate + nitric acid → copper nitrate + water + carbon dioxide:

CuCO₃ + 2HNO₃ → Cu(NO₃)₂ + H₂O + CO₂

• copper carbonate + acetic acid → copper acetate + water + carbon dioxide:

CuCO₃ + 2CH₃COOH → Cu(CH₃COO)₂ + H₂O +CO₂

2) Decomposition:

• copper carbonate decomposes at 290°C and forms copper(II) oxide:

CuCO₃ → CuO + CO₂

References

1. Masterson, W. L., & Hurley, C. N. (2004). Chemistry: Principals and Reactions, 5th Ed. Thomson Learning, Inc. (p 498).

External links

• National Pollutant Inventory – copper and compounds fact sheet

• Carbonates
• Copper compounds
• Pyrotechnic colorants