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Author: Subject: Selectively complexing copper at low pH
Amos
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[*] posted on 21-2-2019 at 08:34
Selectively complexing copper at low pH


I have a job at a hardware store right now that occasionally rewards me with small bounties of stainless steel items when these items can't be sold due to some condition. A recent interest of mine is maintaining a dedicated vessel to toss these in while they slowly digest in hydrochloric acid, as I can economically separate the chromium, nickel, and iron present to obtain useful compounds for my lab.

A recent issue, however, is that some grades of stainless steel contain small amounts of copper, and I've noticed to my dismay recently that all of my screws, nails, bolts, EVERYTHING are completely coated in a passivating layer of copper, as it tends to plate out of solution onto the more reactive metals. I can get the coating to redissolve by dumping a bunch of hydrogen peroxide in and covering the container to keep oxygen in, but it doesn't last long.

My question for you all is, do you know of a ligand that I can use to selectively copper to keep it in solution in a pH range of about 1-2 while allowing the acid to eat up the other metals? Most of the ligands I know that are easy to obtain aren't effective at low pH (tartrate, ammonia, alkylamines). Alternatively I could maybe just pump air through the solution but it wouldn't exactly be fast or cost-effective.
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Tsjerk
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[*] posted on 21-2-2019 at 08:55


Do you have a overhead stirrer? Connect it to a time-clock and you will have the best of both worlds; occasional stirring to get oxygen in, but not to much to lose cost-effectiveness by electricity bills.

Air pumps designed for aquaria use very little electricity and also don't have to run full-time though... A cheap time- clock could go for little money on aliexpress
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AJKOER
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[*] posted on 21-2-2019 at 13:48


Quote: Originally posted by Amos  
....

My question for you all is, do you know of a ligand that I can use to selectively copper to keep it in solution in a pH range of about 1-2 while allowing the acid to eat up the other metals? .....



Try adding much NaCl to create a chloride ligand.

Reference: See https://www.chemguide.co.uk/inorganic/complexions/ligandexch... . Also, a prior comment noting the value of pumping in air (hence the comment by Tsjerk above):

Quote: Originally posted by AJKOER  

There is a known electrochemical reaction in the presence of acid (H+) and oxygen (O2) with a metal in a lower valence state producing a basic salt. Some examples:

4 Cu(l)/Fe(ll)/Co(ll)...+ O2 + 2 H+ --> 4 Cu(ll)/Fe(lll)/Co(lll) + 2 OH- (see, for example, http://www.sciencemadness.org/talk/viewthread.php?tid=81800#... )

Add NaCl (good electrolyte as this is an electrochemical reaction) and also H2O2 (or pump in O2 or boil in air) to create the basic salt.


Note, the creation of the insoluble basic copper salt may occur at more neutral to basic conditions and effectively removes it from solution. You may produce more than one such basic salt depending on pH (which rises as the H+ is consumed in the above cited reaction).

[Edited on 21-2-2019 by AJKOER]
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fusso
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[*] posted on 21-2-2019 at 14:51


NO3-? It will oxidize any Cu to Cu2+ in acidic media. If NO3- is large excess than Cu2+ then Cu will not form.

[Edited on 190221 by fusso]




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[*] posted on 22-2-2019 at 00:24


I think no matter what you complex copper(II) with, iron and other metals will reduce it to copper metal.



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Amos
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[*] posted on 22-2-2019 at 06:27


AJKOER thinks adding a chloride salt to a vat full of dissolved chloride salts will help, sounds about par for the course on their contributions. I do however trust DraconicAcid's expertise when it comes to copper complexes. Nitrates are increasingly hard to find in the quantities I used to be able to but I may have to go that route. At least since I'm precipitating all of the metals with a base downstream I may be able to recycle some nitrate from the mother liquor if it hasn't all gone and been reduced.
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Tsjerk
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[*] posted on 22-2-2019 at 07:31


If too much nitrate is going to be uneconomic for you, you could think about adding small portions once in a while. This way it would not oxidize other metals than copper, as the other metals are plated with copper.

[Edited on 22-2-2019 by Tsjerk]
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Bezaleel
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[*] posted on 22-2-2019 at 08:32


A thought: do you think it could be removed electrolytically by adjusting the potential so that copper will plate out, but iron, chromium and nickel will not?
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[*] posted on 22-2-2019 at 09:03


This is a problem because the copper coat the metal and passivates right? You can oxidize the copper back into solution using air, so just get and air stone and an aquarium pump.
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Tsjerk
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[*] posted on 22-2-2019 at 09:17


Wait a minute... how is it possible the copper dissolves in the first place? There must be oxygen, and when the oxygen gets consumed faster than it can dissolve it starts to plate the metals... If there is no oxygen dissolving in the mixture the copper won't dissolve in the first place and will be left as a spongy material.
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RedDwarf
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[*] posted on 22-2-2019 at 09:28


Out of interest how do you "economically separate the chromium, nickel, and iron present to obtain useful compounds" and what sort of useful compounds do you end up with?
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AJKOER
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[*] posted on 7-11-2019 at 11:58


OK, on further thought, a concept for a possible successful approach.

Try keeping the chemistry in the sphere of electrochemistry, like an electrochemical cell, for example.

Why? Because the metal/alloy will be singularly consumed, in order, per their respective anodic index. This is a central concept in the application of a sacrificial electrode (see, for example, discussion and references at https://chemistry.stackexchange.com/questions/7850/what-can-... ).

As a possible embodiment, use a bleach battery cell with the more active metals being attacked in order of their anodic index. So, adding a mix of, say, metal Al, Fe and Cu to NaOCl in the presence of high surface area carbon, the expected successive products are Al(OH)3, FeO and CuO. I would perform this in a microwave-assisted procedure. Upon taken periodic samples, I would expect pure Al(OH)3, followed by a mixed sample collection of Al(OH)3 and particles of FeO, then just FeO, FeO with CuO and then only CuO.

However, in the presence of alloys, it is the anodic index of the alloy that decides when it is attacked, producing mixed product oxide that could result in protective layers. Adding crushed glass together with the agitation of the solution may be able to mechanical remove the protective coating (I have read electrolysis where mechanical means were employed to clean the electrode to increase product yield).
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VSEPR_VOID
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[*] posted on 7-11-2019 at 12:48


Try boiling a solution of your dissolved metal chlorides with copper metal. The copper will reduce Cu+2 to Cu+1, which is very insoluble. You can filter it out after words.



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unionised
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[*] posted on 7-11-2019 at 13:39


Iodide.

Drop the copper out of solution as CuI

There are 2 ways to look at it.
Either just add iodide and hope that keeps the Cu content down far enough that copper doesn't plate out (and, with a side benefit that iodine/ iodide/ I3- attacks most metals rather well)
Or, from time to time, add enough peroxide to get all the Cu into solution, then pour the liquid off the remaining scrap metal and carefully add just enough iodide to ppt the CuI. (adding something like bisulphite to reduce iodine to iodide)
Decant the acid off the CuI
then pour the (copper free) acid back onto the scrap and let it get on with the job again.
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[*] posted on 7-11-2019 at 14:06


Not sure how applicable this is, but it came to mind:

https://www.youtube.com/watch?v=jRiBYMv6Tz4




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AJKOER
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[*] posted on 7-11-2019 at 15:15


Quote: Originally posted by VSEPR_VOID  
Try boiling a solution of your dissolved metal chlorides with copper metal. The copper will reduce Cu+2 to Cu+1, which is very insoluble. You can filter it out after words.


Per Wikipedia (https://en.wikipedia.org/wiki/Dicopper_chloride_trihydroxide ), to quote:

" CuCl2 + Cu + 2 NaCl → 2 NaCuCl2 (eq.6) "

See also https://books.google.com/books?id=0uwDTrxyaB8C&pg=PA512&... and https://books.google.com/books?id=iZd-vg7RopYC&pg=PA839&... .
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[*] posted on 7-11-2019 at 15:29


I think unionized is on the right track here wrt precipitation. But Cu+ forms CuCl2- in a strong chloride solution, so reducing Cu2+ alone is not going to give you a precipitate.

PABA precipitates Cu2+, but also precipitates Fe2+:

https://www.ncbi.nlm.nih.gov/m/pubmed/18960272/

This may be good or bad. Fe2+ should mostly be oxidized to Fe3+ under suitably strong oxidizing conditions.




[Edited on 04-20-1969 by clearly_not_atara]
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Amos
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[*] posted on 8-11-2019 at 06:51


Some of the responses seem to neglect that this is a dynamic system involving not only a mixed solution of metal ions but also the metals themselves; precipitating the copper can't be done unless all of the copper is dissolved first, since it's currently plating itself onto all of my stainless steel. I'd like to form a complex of copper that is stable at low pH and will not be reduced to metallic form by less reactive metals like iron and nickel.
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