In CuSO4 solution, why doesn't Cu(OH)2 precipitate out?

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The discussion centers on the solubility behavior of Cu(OH)2 in an aqueous CuSO4 solution with a pH of around 4. The presence of excess H+ ions leads to a lower concentration of OH- ions, preventing the precipitation of Cu(OH)2, which is typically insoluble. Participants noted that copper may form soluble complexes such as CuOH+, Cu2(OH)22+, and Cu(OH)42-, with the latter being relevant only at high pH levels due to its low stability constant. The conversation highlights the importance of stability constants and solubility products in understanding these equilibria.

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wywong
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An aqueous solution of CuSO4 has a pH around 4. That means there are more H+ ions than OH-. Since the H+ ions come from dissociation of water, there must be an equal amount of OH-. Where have all the OH- ions gone? I suppose they are locked up as Cu(OH)2. However, the latter is highly insoluble in water. So why doesn't Cu(OH)2 precipitate out? I guess chelation of Cu(OH)2 by H3O+ may make it soluble. However, I have been unable to find any information about such chelation. Can anyone help?

CuCl2, AlSO4, CaCl2 etc. all pose a similar problem.

Many TIAs.

Wai Wong
 
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Solution may contain some amount of copper (or other cation) complexed by OH- groups. There usually whole families of soluble complexes - like CuOH+, Cu2(OH)22+, Cu(OH)42- - involved.

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Thanks Borek!

BTW, I think Cu(OH)42- shouldn't be there, or otherwise Cu(OH)2 would be soluble in water or alkaline solutions.

Wai Wong
 
It depends on the relative values of the stability constant and solubility product. These are taken from the equilibria database, that means it is possible to trace the source of the information.

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I am not good at finding those constants, but according to several sources, Cu(OH)42- forms only at very high pH, suggesting a low stability constant.

Never mind. The root of my question has already been well answered.

Thanks

Wai Wong
 

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