This compound is rarely encountered because it is difficult to prepare[2] and readily reacts with water moisture from the air. The terms "copper carbonate", "copper(II) carbonate", and "cupric carbonate" almost always refer (even in chemistry texts) to a basic copper carbonate (or copper(II) carbonate hydroxide), such as Cu 2(OH)2CO 3 (which occurs naturally as the mineral malachite) or Cu 3(OH)2(CO 3)2 (azurite). For this reason, the qualifier neutral may be used instead of "basic" to refer specifically to CuCO 3.
Reactions that may be expected to yield CuCO 3, such as mixing solutions of copper(II) sulfateCuSO 4 and sodium carbonateinambient conditions, yield instead a basic carbonate and CO 2, due to the great affinity of the Cu2+ ion for the hydroxide anion HO− .[5]
Thermal decomposition of the basic carbonate at atmospheric pressure yields copper(II) oxide rather than the carbonate.
In 1960, C. W. F. T. Pistorius claimed synthesis by heating basic copper carbonate at 180 °C in an atmosphere of carbon dioxide (450 atm) and water (50 atm) for 36 hours. The bulk of the products was well-crystallized malachite Cu 2CO 3(OH)2, but a small yield of a rhombohedral substance was also obtained, claimed to be CuCO 3.[6] However, this synthesis was apparently not reproduced.[2]
Reliable synthesis of true copper(II) carbonate was reported for the first time in 1973 by Hartmut Ehrhardtet al. The compound was obtained as a gray powder, by heating basic copper carbonate in an atmosphere of carbon dioxide (produced by the decomposition of silver oxalateAg 2C 2O 4) at 500 °C and 2 GPa (20,000 atm). The compound was determined to have a monoclinic structure.[7]
The stability of dry CuCO 3 depends critically on the partial pressure of carbon dioxide (pCO2). It is stable for months in dry air, but decomposes slowly into CuO and CO 2 if pCO2 is less than 0.11 atm.[3]
In the presence of water or moist air at 25 °C, CuCO 3 is stable only for pCO2 above 4.57 atmospheres and pH between about 4 and 8.[8] Below that partial pressure, it reacts with water to form a basic carbonate (azurite, Cu 3(CO 3)2(OH)2).[3]
3CuCO 3 + H 2O → Cu 3(CO 3) 2(OH) 2 + CO 2
In highly basic solutions, the complex anion Cu(CO 3)2− 2 is formed instead.[3]
The solubility product of the true copper(II) carbonate was measured by Reiterer and others as pKso = 11.45 ± 0.10 at 25 °C.[2][3][4]
^ abcdH. Seidel, H. Ehrhardt, K. Viswanathan, W. Johannes (1974): "Darstellung, Struktur und Eigenschaften von Kupfer(II)-Carbonat". Z. anorg. allg. Chem., volume 410, pages 138-148. doi:10.1002/zaac.19744100207
^Ahmad, Zaki (2006). Principles of Corrosion Engineering and Corrosion Control. Oxford: Butterworth-Heinemann. pp. 120–270. ISBN9780750659246.
^C. W. F. T. Pistorius (1960): "Synthesis at High Pressure and Lattice Constants of Normal Cupric Carbonate". Experientia, volume XVI, page 447-448. doi:10.1007/BF02171142
^Hartmut Erhardt, Wilhelm Johannes, and Hinrich Seidel (1973): "Hochdrucksynthese von Kupfer(II)-Carbonat", Z. Naturforsch., volume 28b, issue 9-10, page 682. doi:10.1515/znb-1973-9-1021
^H. Gamsjäger and W. Preis (1999): "Copper Content in Synthetic Copper Carbonate." Letter to J. Chem. Educ., volume 76, issue 10, page 1339. doi:10.1021/ed076p1339.1
