Author Topic: CuCl2 from CuO/HCL (Pottery reference)  (Read 771 times)

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CuCl2 from CuO/HCL (Pottery reference)
« on: September 02, 2003, 06:02:00 PM »
I consulted TFSE for methods on synthesizing CuCl2 and I found a several methods at Rhodiums site at

, one of which was written by Hellman. He mentions using CuCO3 with HCL(30%) to produce CuCl2, but he also mentions that CuO can be used in place of CuCO3, but I did not find any references to this with the exception of the following link (pretty cool)...

It's a pottery reference that details the synth of CuCl2 + H2O using black copper oxide (CuO) and HCL(37%).

I'm assuming the water can be removed from the CuCl2 after the synth by boiling, but it would require that the xtals be placed in an air tight container immediately afterwards, otherwise the CuCl2 will absorb water from the surrounding air.

If CuO is black copper oxide, what the hell is CuCO3? Copper carbonate I'm assuming, but if this is also sold at pottery stores, does someone ask for it by name? I mean, does it have a brand name?


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« Reply #1 on: September 04, 2003, 02:39:00 AM »
From Ullmann`s:

Copper(II) chloride [7447-39-4] , CuCl2, Mr 134.45, mp (extrapolated) 630 °C, d425 3.39, begins to decompose to copper(I) chloride and chlorine at about 300 °C. The often reported melting point of 498 °C is actually a melt of a mixture of copper(I) chloride and copper(II) chloride. Decomposition to copper(I) chloride and chlorine is complete at 993 °C. The deliquescent monoclinic crystals are yellow to brown when pure; their thermodynamic data are as follows: cp (298 K) – 579.2 J kg–1 K–1, cp (288 –473 K) – 621.7 J kg–1 K–1, cp (288 – 773 K) – 661.9 J kg–1 K–1, DH° (25 °C) – 247.2 kJ/ mol. In moist air, the dihydrate is formed. Figure (4) shows the solubility of copper(II) chloride in water and hydrochloric acid at two temperatures [89]. At higher concentrations of hydrogen chloride, [CuCl3]– and [CuCl4]2– complexes are formed. Copper(II) chloride is easily soluble in methanol and ethanol and moderately soluble in acetone.
The more common commercial form of copper(II) chloride is the dihydrate [10125-13-0] , CuCl2 · 2 H2O, Mr 170.45, mp around 100 °C (with decomposition to the anhydrous form). This occurs in nature as blue-green orthorhombic, bipyramidal crystals of eriochalcite, d425 2.51. Its solubility characteristics are proportionally similar to those of the anhydrous form. In moist air the dihydrate deliquesces, and in dry air it effloresces.
Production. Because of the relative stabilities of copper(I) chloride and copper(II) chloride at high temperature, it is improbable that a pure anhydrous copper(II) chloride can be prepared by excessive chlorination of copper in a melt, even though such methods have been reported. The most common method for the production of anhydrous copper(II) chloride is by dehydration of the dihydrate at 120 °C. The product must be packaged in air-tight or desiccated containers.
The dihydrate can be prepared by the reaction of copper(II) oxide, copper(II) carbonate, or copper(II) hydroxide with hydrochloric acid and subsequent crystallization. Commercial production of copper(II) chloride dihydrate uses a tower packed with copper. An aqueous solution is circulated through the tower. Sufficient chlorine is passed into the bottom of the tower to oxidize the copper completely [72] , [73]; to prevent hydrolysis [precipitation of copper(II) oxychloride] of concentrated copper(II) chloride solutions, they are kept acidic with hydrochloric acid. The tower can be operated batchwise or continuously; Figure (5) shows the continuous operation. A hot, concentrated liquor is circulated continuously through the tower, and the overflow from the tower is passed through a crystallizer where the liquor is cooled; the product is then centrifuged, dried, and packaged. The addition of hydrogen chloride is pH controlled; the addition of water is controlled by specific gravity. Copper is added daily or twice daily, as needed.

[89]  as reported in A. Siedell (ed.): Solubilities of Inorganic and Metal Organic Compounds, vol. I, D. Van Nostrand, New York 1940, p. 478.
[72]  Schering,

Patent DE1080088

1958 (H. Niemann, K. Herrmann). =

Patent GB927313

[73]  Harshaw,

Patent US2367153

  1945 (C. Swinehart).