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woelen
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[*] posted on 21-2-2006 at 10:58
Strange copper/iron reaction


Recently I have done quite some experimenting with copper ions and also with iron ions. Now I decided to combine the two metals and a whole bunch of new and special phenomena pops up :).

Many of you probably know the FeCl3.6H2O in the form of small pea-sized globules, used for etching copper from printed circuit boards. Especially when combined with some HCl it works quite well.

If you add some copper to a solution of FeCl3.6H2O in conc. HCl, then the copper dissolves and the solution darkens. Most people think that the following reaction occurs:

Cu + 2Fe(3+) --> Cu(2+) + 2Fe(2+)

Well, I found that the real reaction is MUCH more complex!

What do you expect to happen, when CuCl2 (containing Cu(2+) ion) is added to a solution of FeCl2 in HCl? If you look at the above equation, then you would not expect a reaction at all. In reality the solution turns black or very dark brown.

If someone has access to FeCl2 (or FeSO4) and also has access to CuCl2 (or CuSO4), then dissolve some of the iron (II) salt in conc. HCl (appr. 30%) and also dissolve some of the copper (II) salt in conc. HCl. Then add the two solutions to each other. You'll be surprised that the liquid turns almost black. So, there definitely is a reaction. So, apparently Cu(2+) oxidizes Fe(2+) to Fe(3+), itself being reduced to Cu(+).

Another experiment: Add a small amount of copper wire to a concentrated solution of FeCl3.6H2O (from an electronics parts store) in conc. HCl. The liquid becomes amazingly dark, almost black, the copper dissolves. Even with a large excess amount of FeCl3 the liquid still becomes black. So even a large excess of FeCl3 does not oxidize all copper to Cu(2+), but part of the copper remains in solution as Cu(+).

Finally, add a small amount of CuCl2 (or CuSO4) to a large excess of a solution of FeCl2 (or FeSO4) in conc. HCl. Again, the solution becomes almost black.

Conclusion:
Excess Fe(3+) ---> Not all copper (I) is oxidized to copper (II)

Excess Fe(2+) ---> Not all copper (II) is reduced to copper (I).

This seems a contradiction, so there must be a copper (A) species in solution with 1 < A < 2, in other words a fractional oxidation state of copper or a mixed valency complex of copper. This seemingly very simple reaction raises a lot of questions and a lot of interesting things to research. If anybody has ideas or is willing to repeat the experiment and play around with it, you're welcome :).

EDIT: I just posted this, but I wonder if this is the correct forum. I severely doubt this is a beginners question.

[Edited on 21-2-06 by woelen]




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[*] posted on 21-2-2006 at 16:01


There was once this heated discussion of JohnWW with me on the subject of PCB etching, using FeCl3.
https://sciencemadness.org/talk/viewthread.php?tid=2519


Essentially, it is possible that complex ions form, such as [FeCl4]- or [CuCl4]2- form.
However, you mentioned that the black colour only occurs when the CuCl2/FeCl2 are mixed, and are in the presence of HCl.
That is, CuCl2/HCl and FeCl3/2/HCl in isolation are *not* black? That'd mean that the respective chloro-complexes are not responsible for the colour.

One thing that made me think are the redox potentials.

FeCl2 and CuCl2 could possibly react, to form FeCl3 and CuCl. The latter is insoluble and white, but it is also forming a chloro complex, [CuCl2]- (just checked), which is soluble (possibly explaining why this reaction might not occur in the absence of HCl). I am not so familiar with electrochemistry as I used to be, so correct me if I am wrong.

The reduction potentials are

(Cu2+) + e- --> Cu+ |+0.16 V

(Fe2+) ---> (Fe3+) + e- | -0.77 V


This reaction should proceed, so Cu2+ act as an oxidiser.

Conversely, the reaction between Cu and FeCl3 is

(Fe3+) + e- ---> (Fe2+) | +0.77 V (the reverse)

Cu ---> Cu2+ + 2 e- | - 0.34 V


THe reduction potential is greater in this reaction, so this reaction proceeds readily, and so should it proceed for the first reaction.


Anyway - if I am not mistaken you could conclude that the mixture after Cu etching contains, when in the presence of FeCl3/HCl, a mixture of [CuCl2]-, [CuCl4]2- and the various iron complexes.

It would be interesting to see what colour clean, precipitated CuCl, which is redissolved in strong HCl/NaCl (to provide the counter ion), has, or conversely, what the colour of [FeCl4]- is.


What do you think?




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[*] posted on 21-2-2006 at 16:33


Quote:

(Cu2+) + e- --> Cu+ |+0.16 V

(Fe2+) ---> (Fe3+) + e- | -0.77 V


This reaction should proceed, so Cu2+ act as an oxidiser.


How would the reaction proceed if Eo is negative?

And by the way, is the liquid a CLEAR black or like colloidial solution?

[Edited on 2/22/2006 by guy]




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[*] posted on 21-2-2006 at 17:07


Ah, I forgot that Eo has to be negative, for some reason I thought only a potential difference is required which doesnt really make sense on second thought :o.
True I did think that Fe3+ is a stronger oxidiser than Cu2+ so how could Cu2+ oxidise Fe2+ to Fe3+? Silly me.

Still... I do think the solution lies in those chloro complexes.

[Edited on 22-2-2006 by chemoleo]




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[*] posted on 21-2-2006 at 18:00


I have an idea. Reduction potentials are a result of thermodyanamic stability, so if the the product is more stable, then the reaction can proceed. I haven't heard of the complex [FeCl4]2- before this post, but I think it will be unstable (K is less than 1). The complex [CuCl2]- is a lot more stable than either [FeCl4]2- or [CuCl4]2-. Therefore it is more favorable and stable to form the [CuCl2]-.

[FeCl4]- + [CuCl4]2- --> [CuCl2]- + FeCl3 + 2Cl-

Just a guess.




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[*] posted on 21-2-2006 at 18:32


Well absolutely, it's always a matter of equilibria. Always, regardless how strong or favourable the reaction is, regardless the potentials.

What makes you think that CuCl2- is more stable, than the tetrachloro version [CuCl4]2-, or the ferro-versions?




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[*] posted on 21-2-2006 at 19:38


Well I tried to replicate the experiement. I don't have HCl so I used NaCl. Somehow it didnt form the [CuCl4] 2- complex when I added NaCl to CuSO4. I also tried to add NaCl to FeSO4 and it didnt form a complex. Do these complexes only form in acidic solutions?

[Edited on 2/22/2006 by guy]




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[*] posted on 21-2-2006 at 20:01


Yes they do.
Dilution of [CuCl2]- into water results in precipitation of CuCl, and [CuCl4]2- can only be formed in strongly acidic solutions.

Also, I'd rather stick to the same counterions, i.e. only H2SO4, or only HCl. No mixture of FeSO4 and Cl-. It will just complicate things.




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[*] posted on 21-2-2006 at 22:14


Quote:
Also, I'd rather stick to the same counterions, i.e. only H2SO4, or only HCl. No mixture of FeSO4 and Cl-. It will just complicate things.


But then copper and iron dont form sulfate complexes.

Also WHY do you need to have acidic conditions in order to form the complex? It doesn't appear anywhere in the equation. [Cu(H2O)4]2+ + 4Cl- <---> [CuCl4]2- + 4H2O




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[*] posted on 22-2-2006 at 00:15


Some additional observations:

FeCl4(-) (iron (III)) is deep yellow. At very high concentration it looks brown.
CuCl4(2-) (copper (II)) is green/yellow and exists at very high chloride concentration.

CuCl is white like snow. I made this several times. It is hard to keep it white though.
CuCl2(-) is colorless. I made this as well, by adding Cu metal to a solution of CuCl2 in conc. HCl and stoppering the test tube. The liquid becomes (almost) colorless after a day or so.

Mixing FeCl3 with CuCl2 in conc. HCl does not result in a black or dark brown liquid.
Mixing FeCl3 with Cu in conc. HCl gives a black liquid.
Mixing FeCl2 with CuCl2 in conc. HCl gives a black liquid.
Dissolving FeCl2 in HCl gives a very pale green liquid, so here no complexes.

Tonight I'll try mixing FeCl2 and FeCl3 in conc. HCl. That combo I did not try yet. I expect, however, this will simply give the color of FeCl3 in HCl (the waek color of FeCl2 being masked).

I also did other experiments with copper alone as I mentioned, and there I found that a mix of copper (I)/copper (II) is dark brown. Have a look at my page on copper (I) / copper (II) complexes.
In the iron/copper experiment I also think that I have such complexes, but what surprises me is that I get these. I always though that Fe(3+) would oxidize copper to Cu(2+), itself being reduced to Fe(2+). But this reaction apparently does not occur or only partially. I now think that there is an in-between situation, with the dark brown/black complexes being a stable species.

Indeed, the reactions described only work good in very strong chloride solutions. That is why I use conc. HCl. I, however, do not think the acid is essential, but I cannot achieve the high concentrations of chloride with other chemicals.

You can perfectly use CuSO4 and FeSO4 for the experiments, if you do not have the chlorides. If I dissolve a spatula of CuCl2 or a spatula of CuSO4 in a few ml of conc. HCl, then I see no difference. The SO4(2-) ion just is a spectator ion.

[Edited on 22-2-06 by woelen]




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[*] posted on 22-2-2006 at 17:00


Looking at the CuCl2 + CuCl reaction on your site, I think maybe the reaction is [Cu(H2O)4]2(aq)+ + 2Cl-(aq) + CuCl(s) <--> Cu(CuCl2)Cl(s) + 4H2O(l) or something along those lines. It looks like it makes sense because when adding excess water, it can revert back to the original products. The reaction of [CuCl2]- and Cu2+ is definetly the key to the black substance. Try adding NaCuCl2 to CuSO4 and see if that makes a black substance.

I finally managed to produce [CuCl4]2- by adding a TON of NaCl to CuSO4. I halved the solution. One solution, I put a piece of copper in. No reaction. In another solution, I added FeCl2. Still no reaction. The concentrations must be very concentrated to work? or only in acidic conditions?

[Edited on 2/23/2006 by guy]




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[*] posted on 8-10-2016 at 11:39


To quote Woelen, in part, from above:

"Conclusion:

Excess Fe(3+) ---> Not all copper (I) is oxidized to copper (II)
Excess Fe(2+) ---> Not all copper (II) is reduced to copper (I).

This seems a contradiction, so there must be a copper (A) species in solution with 1 < A < 2, in other words a fractional oxidation state of copper..."

From a recent thread, a link to an educational reference with a somewhat succinct comment as to a "coupled redox reaction" (see https://www.sciencemadness.org/whisper/post.php?action=reply... ):
Quote: Originally posted by AJKOER  
Per my Wikipedia reference on what I contend is a similar reaction scheme occurring with acetate, cited equations to quote:

" CuCl2 + Cu + 2 NaCl → 2 NaCuCl2 (eq.6)
6 NaCuCl2 + 3/2 O2 + H2O → 2 Cu2(OH)3Cl + 2 CuCl2 + 6 NaCl (eq.7) "

where Equation (7) indicates redox chemistry.

So, assuming we can move the cupric into cuprous, a redox reaction could proceed. In the above cited system, the presence of copper metal assisted in forming cuprous from cupric. Alternately, to quote a source ( https://www.researchgate.net/publication/11374766_Generation...):

"The process is enhanced by contaminating Fe3+ and Cu2+;"
"The addition of Fe2+ and Cu+ (0-20 microM) to KH resulted in a concentration-dependent increase in *OH formation, as measured by the salicylate method."

where an iron contamination could arise from using tap water (containing some ferrous bicarbonate, for example, and noting in the opening thread, to quote MrbunGee, "I was not using distilled water, but there just can’t be that much CO3 ions in my water. :?" ).

[Edit] Yet another reference:

Fe2+ + Cu2+ ↔ Fe3+ + Cu+ (coupled redox reaction)

See: https://www.google.com/url?sa=t&source=web&rct=j&...

[Edited on 23-9-2016 by AJKOER]


Note, the comment above "The process is enhanced by contaminating Fe3+ and Cu2+", which is repeated in the research literature, citing the apparent beneficial impact of a mixed transition metal system which can introduce a coupled equilibria. In the same 2000 paper "Generation of .OH initiated by interaction of Fe2+ and Cu+ with dioxygen; comparison with the Fenton chemistry" by Norbert K. Urbañski and Andrzej Berêsewicz, available at https://www.google.com/url?q=http://www.actabp.pl/pdf/4_2000... , the authors noted to quote "The Fe2+-mediated .OH yield was enhanced not only by Fe3+ but also by Cu2+ (Fig.3)".

Also, supporting material from more recent work on hetergeneous transition metal catalysts, "Review on the application of modified iron oxides as heterogeneous catalysts in Fenton reactions", by Shima Rahim Pouran, et al, 2011 available at http://www.researchgate.net/publication/257353836_Review_on_... , where the authors note in the case of transition metal substituted iron oxides (TMSIOs), to quote:

"Two mechanisms were suggested for enhanced activity of TMSIOs: (i) the participation of the thermodynamically favourable redox pairs, Fe3+/Fe2+ and Mn+/Mn+1, in H2O2 oxidation cycle, to produce OH radicals and (ii) generation of oxygen vacancies on the surface of catalyst, resulted from adjustments of unequal charge replacements".

In another work, "Impact of MnO2 on the efficiency of metallic iron for the removal of dissolved CrVI, CuII, MoVI, SbV, UVI and ZnII", by C. Noubactep K.B.D., et al. , 2011 (link: http://www.sciencedirect.com/science/article/pii/S1385894711... ), the authors confirms the observation that manganese oxide (MnO2) sustains the reactivity of metallic iron (Fe0) in a multi-elemental aqueous system containing Cr(VI), Cu(II), Mo(VI), Sb(V), U(VI), and Zn(II).

It is important to note that the coupled redox equilibrium can be dependent on the solubility of the lower valent state salt. So, for example, Cu(l) is more soluble in presence of chloride (complexation), which could explain the beneficial effect of say NaCl in a Fenton-type reaction occurring in a mixed transition metal system.

For additional references particular to coupled redox systems see, for example: https://www.google.com/url?sa=t&source=web&rct=j&... which restates the Fe/Cu system and for a more complex case example, see http://pubs.acs.org/doi/abs/10.1021/om960749b .

Per another source (link: http://chem.libretexts.org/Textbook_Maps/Inorganic_Chemistry_Textbook_Maps/Map%3A_Inorganic_Chemistry_(Wikibook)/Chapter_04%3A_Redox_Stability_and_Red ox_Reactions/4.4%3A_Redox_Reactions_with_Coupled_Equilibria ) to quote:

"Coupled equilibria (solubility, complexation, acid-base, and other reactions) change the value of E°, effectively by changing the concentrations of free metal ions."

where, for example, a pH change could arise from metal autoxidation via dioxgen as was noted above.

Twenty years ago the importance of coupled redox reactions was bearly addressed. This was unfortunate given its subsequent benefit it apparently provided in some adanced oxidation processes/technologies (where, for the most part, AOP/AOT focus is on drinking water purification and environmental remediation of polluted soil and waste water). As an example, a quote from the last page of a published 1996 article, "Catalytic Metals, Ascorbate and Free Radicals: Combinations to Avoid", by Garry R. Buettner and Beth Anne Jurkiewicz, (link: https://www.google.com/url?sa=t&source=web&rct=j&... ), to quote:

"One area that has only been examined briefly is the potential synergy of metals in oxidations. Our colleagues in environmental research have noted that iron and copper are co-conspirators in the oxidation of organics in atmospheric waters (79). It was determined that reduced copper transferred an electron to iron, which in turn participates in the oxidation process.

k= 3 x E07 M–1s–1 (79)

Cu1+ + Fe3+ → Cu2+ + Fe2+

These same processes may be of significance in a biological setting."
-----------------------------------------------------------------------
It is important to understand the mechanics of Fenton and Fenton-type reactions as they move lower state transition metals to higher valent states. As an example, see "Fenton chemistry in biology and medicine*" by Josef Prousek, to quote reaction (15) on page 2330:

"For Fe(II) and Cu(I), this situation can be generally depicted as follows [20,39],

Fe2+/Cu+ + HOX → Fe3+/Cu2+ + .OH + X- (15)

where X = Cl, ONO, and SCN. "

[Edited on 9-10-2016 by AJKOER]
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[*] posted on 8-10-2016 at 17:59
Interesting Cu / Fe species


Woelen, just for fun, is there any way you could let your old solutions evaporate? I wonder if anything interesting crystallises out! That might give some insight as to what is present.
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[*] posted on 14-10-2016 at 01:37


I think you will find that the coloured species is a polynuclear chloro-complex of either Cu and Fe or Cu alone. I haven't seen information on complexes of this type for these metals but there is plenty of information for some other metals such as Co and Cr.

Both Fe and Cu are notorious for forming highly labile and highly coloured chloro-complexes that transform one into another with only a small change of chloride ion or free HCl concentration, so in this case the very high HCl/Cl- concentration is required but that does not mean that Cu+ is a premanent species (though presumably the copper must pass through this valency enroute to Cu2+). Anhydrous ferric chloride exists in the vapour phase (and possibly in some organic solvents) as just such a binuclear Fe2Cl6, so maybe you are generating say a CuFeCl5 neutral complex.

Colour is not a good indicator of valency. The dark colour could be due to complementary absorption spectra of seperate complex ions of any valency and a polynuclear ion could have similar complementary absorption spectra associated with each centre.

It might be possible to abtain a solid crystalline complex or a mixture by placing a drop of the solution on a microscope slide and adding a few grains of caesium chloride. If you can't crystallise a dark caesium salt then I suspect a neutral complex is formed. I have used CsCl for analysing mineral under the microscope and as far as I can recall I have never made such a dark crystalline compound, instead separate crystalline compounds are formed; copper forms both deep red and deep yellow crystals while Fe3+ forms amber coloured monoclinic blades.
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[*] posted on 14-10-2016 at 13:13


Maybe you can form FeClx[CuCl2](4-x)-, i.e. one Fe3+ coordinated to four of either Cl- or CuCl2-. Iron needs four electrons to fill a d-block orbital and CuCl2- has one extra electron. Apparently woelen has been down this rabbit hole before:

http://www.scienceforums.net/topic/15728-nails-in-copper-ii-...

EDIT: found this:

http://pubs.acs.org/doi/abs/10.1021/ie51394a037

[Edited on 14-10-2016 by clearly_not_atara]
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