Question

I have an almost-full quart of one of those blue solutions sold for producing a black finish on steel. I have had no use for it since I gained the capability of doing genuine black oxide, but with selenium selling for around $1/gm, it seems a shame to discard it without making an attempt at recovery. These products all contain selenious acid and cupric nitrate, plus a strong acid to prevent precipitation of cupric selenite. The MSDS for the product I have shows the presence of hydrochloric acid and a little phosphoric acid. That's all. Separation should be easy. (Famous Last Words)

The volatile acids could be evaporated, if need be, and the phosphate seems unlikely to cause trouble. When I looked up the reagents commonly used to precipitate selenium, though, I was dismayed to find that every one of them also precipitates copper. And vice versa. I hadn't expected so much difficulty separating an anion from a cation. That thought suggested perhaps a resort to electrolysis. Plate out the copper on the cathode while the selenite heads toward the anode. Alas, I found mention of this method in an old book on quantitative analysis. It is said to work, but imperfectly. Some of the selenium is deposited with the copper. There ought to be SOME reasonably straightforward way to pry selenium apart from Cu⁺⁺. What do you suggest I try?

Answer

A couple ideas spring to mind, but take them with a grain of salt as I have almost no experience with use of selenium or its compounds in my laboratory experience, so if you are going to try any of them, small scale is the way to go.  My first idea is to precipitate the selenium as selenium dioxide by adding concentrated sulfuric acid to the solution:

H₂SeO_{3 (aq)} + H₂SO₄ ---> SeO_{2 (s)} + H₂SO₄*H₂O

This is similar to dropping chromium trioxide out of solution  (Chromic acid is just another easily dehydrated acid).  One foreseeable problem is that as the concentration of sulfuric acid goes up, the solubility of your copper salts is going to drop conversely.  So you could end up with co-precipitation or even the copper salts dropping out before the selenous acid dehydrates from the sulfuric acid.  Regardless, at least in this form you have the option of subliming your SeO₂ away from any copper salts that drop out as well [Selenium dioxide sublimes at 340-350ºC, copper (II) nitrate decomposes at 170ºC, watch for NO₂]⁽¹⁾  but a sublimation attempt might be difficult at the temperatures involved.  From there dissolving the selenium dioxide back in water to give the selenous acid followed by reduction.  On Science Made Alive in the experiments section on chalcogen chemistry this is accomplished by metabisulfite or sulfite, but other reagents may work just as well.

H₂SeO_{3 (aq)} + Na₂CO_{3 (aq)} ---> Na2SeO_{3 (aq)} + CO_{2 (g)} + H₂O _(l)
Cu²⁺ _(aq) + Na₂CO_{3 (aq)} ---> CuCO_{3 (s)} + 2Na⁺ _(aq)

A second option would be to add in a sodium hydroxide or sodium carbonate solution.  Sodium selenate is listed as soluble in the Handbook of Chemistry and Physics⁽²⁾ although an exact number is not given whereas copper carbonate / hydroxide are relatively insoluble.  Thus precipitating out the copper and hopefully leaving the selenium salt in solution to be later reduced to the solid form.  Worry here would be that by basifying the solution you would give the copper selenite a chance to drop out as well.

Cu²⁺_(s) + 4NH_{3 (aq)} ---->  Cu(NH₃)₄²⁺_(aq)

Finally, in any of these cases, including the electrolytic methods you proposed, copper compounds increase their solubility greatly in ammonia solutions due to formation of tetramine copper complexes.  Even elemental copper is subject to this corrosion.  So, washing/soaking any obtained selenium with ammonia would help to cleanse it of copper contamination.

Evaporating the volatile acids first like you proposed may simplify matters, followed by re-dissolution.  In my experience electrolysis can be quite fickle, especially if you are trying to develop the method yourself, I would definitely investigate chemical separation first, but electrolysis could prove useful if all else fails.

(1)  Lewis, Richard J. Sr. Hawley's Condensed Chemical Dictionary.  13th ed.  New York, NY: John Wiley & Sons, Inc.  1997

(2) Hodgman, Charles D. Handbook of Chemistry and Physics. 39th ed. Cleveland, OH: Chemical Rubber Publishing Co. 1957