chempyre235
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Questions about Metal Chloride Hydrated Complexes
It's been awhile since I've worked on this project, but while I was in college, I decided to make various metallic salts during winter break. Keep in
mind, I was rather uninformed at the time, and took no measurements.
I used a pencil, a 9-volt battery and some alligator clips to electrolyze salt water with cathodes of different compositions. I used a copper wire to
obtain the hydrate of copper(II) chloride (I believe). The first time I tried this, I obtained some deep red precipitate, which I think is the
chloride complex of the chloride salt. However, I could not repeat this the two other times I attempted this. I wanted to obtain both the red chloride
and blue hydrate complexes. I also tried using a higher salt concentration, which still did nothing. What do you think went wrong?
On another round of electrolysis, I used a nickel (75% Cu, 25% Ni) to obtain mixed chlorides in a terra cotta colored solution, which yielded a
precipitate of the same color. I dried it for a couple of days on some filter paper, and to my surprise, it had turned to a bluish green color. Any
ideas as to why the color change took place?
The vials in the image (from right to left; the image won't cooperate  ) are
salts produced from various metals via electrolysis: zinc, from pennies after sanding the copper coating; lead, from shot; aluminum, from foil;
copper, from wiring; and lastly is the nickel/copper mixture.
P.S. If anyone knows how to embed the pictures correctly, I would also appreciate that!
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teodor
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On my knowledge the electrolisys of NaCl with copper electrode gives CuCl which could be hydrolized to red Cu2O. The rest would become CuCl2 by
oxidation in the electrolite solution. Try to dissolve the red precipitate in ammonia water solution.
As for other transition metals the change of colors could be connected with different oxidation states.
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bnull
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Try cropping the image. I took a bottom strip off:
It seems that it has to do with proportions. I would have to experiment to know what qualifies as an unacceptable proportion.
The deep red precipitate was copper (i) oxide. The color varies from red to orange (and yellow) depending on the size of the precipitate. There's no
way to have the complex salt with sodium hydroxide lurking around all along. You needed some hydrochloric acid to neutralise the hydroxide.
The same goes for the other salts. It was one of the things that frustrated me decades ago. Milky, bluish brine is not much interesting for a
12-year-old. It is one way to make oxide and hydroxides but I wasn't interested in them back then.
What you have are oxides and hydroxides and mixed stuff (from left to right in the image above): (1) zinc oxide and hydroxide; (2) lead (ii) oxide (it
is a little more complicated than that) and maybe some lead (ii) chloride (I don't think so); (3) aluminum oxide or hydroxide, or both; (4) copper
(ii) hydroxide and basic copper (ii) carbonate. (5) was a mixture of copper (ii) and nickel (ii) hydroxides, which became bluish-green because of the
formation of basic copper (ii) carbonate. Again, it was another thing that frustrated me at that time.
Test the solubilities: (1), (2), and (3) are insoluble in water and soluble in NaOH and other alkaline hydroxides; (4) and (5) are insoluble in water
and soluble in ammonia.
[Edited on 31-12-2024 by bnull]
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chempyre235
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I think you've answered all my questions. Thanks to you both!
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