Sciencemadness Discussion Board - Reactions Of Copper With Ferric Chloride

Reactions Of Copper With Ferric Chloride

Chemistry_Keegan - 1-4-2014 at 08:11

As most of you are probably aware, ferric chloride is often used as an enchant for printed circuit boards. I was particularly interested in the chemistry of this reaction, so I did a little research and found this. One of the reasons I researched this was because I knew it couldn't have been a plain single displacement reaction, because copper is less reactive than iron. So, as Woelen explained in that forum, the ferric ions actually oxidize the copper to form the tetrachlorocuprate coordination complex, which in turn oxidizes even more copper to form the dichlorocuprate complex after all the ferric ions have been transformed into ferrous ions. He also mentioned the chloride concentration needs to be really high for these reactions to proceed, which can easily be accomplished with sodium chloride.

My question is, once the solution has absorbed as much copper as it can, what would form upon evaporating all the water away? Would the coordination complexes stay as they are, and form ferric dichlorocuprate, or sodium dichlorocuprate if sodium chloride was used? Or would they both decompose into ferric chloride, sodium chloride, and cuprous chloride? Also, if the coordination complexes do indeed remain unchanged, could they still be used as a route to cuprous chloride? Sorry if these questions have obvious answers, I'm just beginning to learn about coordination complexes

[Edited on 1-4-2014 by Chemistry_Keegan]

blogfast25 - 1-4-2014 at 09:09

Evaporating anything 'cuprous' means exposing it to air and air oxygen oxidises cuprous salts very quickly. Even kept under water, CuCl is very prone to oxidising to Cu (II).

I'm not sure salts of the complex CuCl42- actually exist as solids. In the spent etching solution there are no cations other than Cu2+, Fe2+ and maybe Cu+.

I think CuCl42- on evaporation will manifest itself as:

Cu2+ + CuCl42- === > 2 CuCl2 (otherwise one would have to consider Cu (II) chloride as [Cu2+][CuCl42-], "copper (II) tetrachloro cuprate")

Similarly with the Fe (II) complex...

The stability constants of these two chloro complexes are quite low, if I recall well.

[Edited on 1-4-2014 by blogfast25]

aga - 1-4-2014 at 10:14

Quote: Originally posted by Chemistry_Keegan

easily be accomplished with sodium chloride.

Interesting, as i generally use Ferric Chloride to etch pcbs.

I knew that adding a little sulphuric acid speeds up the process (as does a reasonable temperature ~40 C) but adding salt ?

I'll try it and see if it does anything nasty to the photo-resist material.

Chemistry_Keegan - 1-4-2014 at 10:25

Quote: Originally posted by blogfast25

Evaporating anything 'cuprous' means exposing it to air and air oxygen oxidises cuprous salts very quickly. Even kept under water, CuCl is very prone to oxidising to Cu (II).

Oops, when I asked if this could be used as a route to cuprous chloride, I really meant to say cupric chloride, so if I understood correctly, this actually works out well for me. I am trying to make the cupric chloride for other experiments, and since the chemicals available to me are very limited at the moment, this appears to be the only plausible way for me to do so.

Correct me if I'm wrong, but this is my understanding of what I can do so far:

FeCl₃ _(s) + NaCl _(s) → Na⁺ _(aq) + FeCl₄^- _(aq)

2 FeCl₄^- _(aq) + Cu _(s) → CuCl₄^2- _(aq) + 2 Fe²⁺ _(aq) + 4 Cl^- _(aq)

2 Na⁺ _(aq) + CuCl₄^2- _(aq) + 2 Fe²⁺ _(aq) + 4 Cl^- _(aq) → 2 FeCl₂ _(s) + 2 NaCl _(s) + CuCl₂ _(s)

For an overall reaction of:

2 FeCl₃ _(s) + 2 NaCl _(s) + Cu _(s) → 2 FeCl₂ _(s) + 2 NaCl _(s) + CuCl₂ _(s)

I omitted this reaction:

CuCl₄^2- _(aq) + Cu _(s) → 2 CuCl₂^- _(aq)

Because, as you said, the oxygen just oxidizes the copper back to it's oxidation state of plus two. But I feel I misunderstood something. Does the oxygen simply aid in reversing the above reaction, like this:

2 CuCl₂^- _(aq) → CuCl₄^2- _(aq) + Cu _(s)

Or does it actually participate in yet another reaction?

Anyway, if the first four equations are indeed accurate, is their an easy way to go about separating the tree chloride salts? Thank you for all the info that you already provided! I apologize if I'm getting ahead of myself by asking too many questions at once

[Edited on 1-4-2014 by Chemistry_Keegan]

Chemistry_Keegan - 1-4-2014 at 10:37

Quote: Originally posted by aga

Quote: Originally posted by Chemistry_Keegan

easily be accomplished with sodium chloride.

This should normally just be the case for homemade ferric chloride. If you bought some specifically for that purpose, chances are the chloride concentration is already high enough for the reactions to proceed properly. However, it should never hurt to add some extra chloride if the solution isn't already saturated.

[Edited on 1-4-2014 by Chemistry_Keegan]

blogfast25 - 1-4-2014 at 11:54

Quote: Originally posted by Chemistry_Keegan

Or does it actually participate in yet another reaction?

Anyway, if the first four equations are indeed accurate, is their an easy way to go about separating the tree chloride salts? Thank you for all the info that you already provided! I apologize if I'm getting ahead of myself by asking too many questions at once

It is an actual oxidation: the oxygen actually absorbs the electron from:

Cu+ === > Cu2+ + e

1/2 O2 + 2 H+ + 2 e === > H2O

Separating the Cu(II) and Fe(II) chlorides may prove difficult. For one, the Fe(II) is oxidised back to Fe(III) fairly easily due to the second reaction (above).

I don't see it as a practical route to CuCl2. Much easier is to precipitate a known quantity of Cu(OH)2 and dissolve it in minimum needed amount of HCl.

Incidentally, I think your 'how to' link is wrong about generating Cu+: the etching solution would have become far too weak to still be useful, well before Cu(II) + Cu(0) === > Cu(I) would kick in. But prolonged exposure of a spent etching solution to copper would probably start that reaction.

[Edited on 2-4-2014 by blogfast25]

Chemistry_Keegan - 1-4-2014 at 13:28

Quote: Originally posted by blogfast25

It is an actual oxidation: the oxygen actually absorbs the electron from:

Cu+ === > Cu2+ + e

1/2 O2 + 2 H+ + 2 e === > H2O

But where do the positive hydrogen ions come from?

blogfast25 - 2-4-2014 at 04:14

Quote: Originally posted by Chemistry_Keegan

But where do the positive hydrogen ions come from?

Auto-dissociation of water:

2 H2O < === > H3O+ + OH-

In the case where ferric ions are present, hydrolysis of these:

[Fe(H2O6]3+ + H2O < === > [Fe(H2O5(OH)]2+ + H3O+

Further deprotonation of the ferric aqua ion leads to ferric hydroxide precipitation.

The oxidation of Cu(I) to Cu(II) also takes place in neutral/alkaline conditions (as does the oxidation of Fe(II) to Fe(III)):

1/2 O2 + H2O + 2 e === > 2 OH-

[Edited on 2-4-2014 by blogfast25]

Chemistry_Keegan - 2-4-2014 at 15:51

Sorry to bother you again, but what you're saying just isn't making sense to me. I can't seem to find a way for the oxidation of cuprous chloride to work with the reactions you described. You mentioned that it would oxidize in the presence of air, but then you implied water is required for it to happen, via its dissociation. Also, the equations you provided don't even seem to work with water anyway, due to fact that the charges and quantities of all the ions and chemicals can't balance.

I just don't understand how cuprous chloride can be oxidized to cupric chloride, without extra hydrochloric acid involved, like in this reaction:

4 CuCl + O₂ + 4 HCl → 4 CuCl₂ + 2 H₂O

Otherwise this just seems impossible. I hear many people refer to other salts being oxidized by air, but it seems as though they must always be in solution with an excess amount of their related acid in order for this to actually occur.

DraconicAcid - 2-4-2014 at 16:00

If there's no acid, then 2 H₂O + 4 CuCl + O₂ --> 2 Cu(OH)₂ + 2 CuCl₂

Chemistry_Keegan - 2-4-2014 at 16:42

Quote: Originally posted by DraconicAcid

If there's no acid, then 2 H₂O + 4 CuCl + O₂ --> 2 Cu(OH)₂ + 2 CuCl₂

Just as I thought, I misunderstood what Blogfast was trying to say. When he said that the copper would be oxidized again, I made the assumption that it would simply be reverted back to cupric chloride, without considering the fact that the copper could also form a salt with the hydroxide anion. Silly me

A big thanks to both of you for putting up with my temporary obliviousness

[Edited on 3-4-2014 by Chemistry_Keegan]

eidolonicaurum - 9-4-2014 at 02:22

A while back I dissolved some copper up in concentrated hydrochloric acid (with a very little sodium chlorate to help oxidise the copper and dissolve it. This may be total rubbish) and evaporated the solution. I obtained a dark green solid. Cupric chloride is blue, the tetrachlorocuprate(II) ion is dark green. So why, if this complex is unstable in solid form, does the solid resemble it? On dilution in water, a white precipitate forms, with a blue solution.

blogfast25 - 9-4-2014 at 04:46

Quote: Originally posted by eidolonicaurum

I obtained a dark green solid.

For starters, are you sure your product corresponded to the stoichiometry of CuCl2? You had excess HCl in there and KClO3 which yields KCl. Now you boil all that down, but what have you got?

That proper CuCl2 is blue kind of proves that Cu II tetrachlorocuprate doesn't exist as a solid, or it would be green.

[Edited on 9-4-2014 by blogfast25]

aga - 9-4-2014 at 13:32

Quote: Originally posted by blogfast25

kind of proves that Cu II tetrachlorocuprate doesn't exist as a solid, or it would be green.

Note to self : never wear tinted eye-protection, and always use sunlight + the Good eye for basic spectrometry.

eidolonicaurum - 11-4-2014 at 04:51

Quote: Originally posted by blogfast25

are you sure your product corresponded to the stoichiometry of CuCl2?

No, but I know that cuprous compounds are almost always colourless (excepting cuprus oxide), so it should be a cupric compound. The colour is the same as tetrachlorocuprate(II) but it clearly isn't. What do you think it might be?

The crystals are also somewhat hygroscopic...

[Edited on 11-4-2014 by eidolonicaurum]

blogfast25 - 11-4-2014 at 05:32

Odd, that green really is very reminiscent of tetrachlorocuprate. Try drying them: in solution CuCl2 really is green, because those pesky tetrachlorocuprate anions then form. My guess is that if you dry your product properly it will end up blue.

Quote: Originally posted by aga

Note to self : never wear tinted eye-protection, and always use sunlight + the Good eye for basic spectrometry.

Totally true: observing colour objectively is hard. Rose tinted specs can be fun but not in science.

[Edited on 11-4-2014 by blogfast25]

eidolonicaurum - 11-4-2014 at 05:49

The crusting around the edge is more of the same substance, and there are some crystals in the pot, but I'll separate them out and dry them. The browny colour is just a reflection of something, I don't know what.

DraconicAcid - 11-4-2014 at 07:45

Hydrated copper(II) chloride is green, possibly greenish-blue. Anhydrous copper(II) chloride is a chocolate brown.

blogfast25 - 11-4-2014 at 10:38

Quote: Originally posted by DraconicAcid

Hydrated copper(II) chloride is green, possibly greenish-blue. Anhydrous copper(II) chloride is a chocolate brown.

Hmmm... see for instance the Wiki photo: looks more blue than green to me....

eidolonicaurum - 11-4-2014 at 10:40

It looks like the crystals must have had a little liquid adhering to them, or a similar problem has occured that mrhomescientist encountered here

Texium - 22-6-2014 at 14:10

I'm having that same issue with the weird green copper(II) chloride, but I take it from this last picture that yours did dry out properly.
Just for clarification, was it necessary to use a desiccator, or were you able to dry them solely by evaporation? I left mine outside for a few days and it was still quite green.

It looks darker green in person

[Edited on 6-22-2014 by zts16]

Brain&Force - 22-6-2014 at 14:46

Isn't solid copper(II) chloride a neutral complex with copper(II) ions coordinating to two water molecules and two chloride ions?

I would think anhydrous CuCl₂ would be copper(II) tetrachlorocuprate.

Texium - 22-6-2014 at 14:51

I don't know, I was under the impression that it was a normal ionic compound, usually seen as the blue-green (much more blue than what I have) dihydrate, and that it only forms complexes when there are excess chloride ions in solution.

Texium - 22-6-2014 at 18:24

I left it in my desiccator for a few hours, and it's all brown now, as is to be expected with anhydrous copper(II) chloride. I think I will add a stoichiometric amount of water to make it into the dihydrate without letting it get overly saturated and see if it assumes the proper color.

Brain&Force - 22-6-2014 at 18:58

Maybe there's a slight excess of chloride from some other source, like sodium chloride, that's turning the product green. But the light color, that's what gets me. The copper chloride source I've been using for the past few months is really dark.

Texium - 22-6-2014 at 19:27

There aren't any other chloride sources in there. The calcium chloride I used is very high quality. The copper sulfate was store bought, but I recrystallized it a couple times, and it seemed very pure from the start. If there are any contaminants, it would just be calcium sulfate, but I recrystallized the final product to get rid of all that was visible. I'll see what happens tomorrow when I rehydrate it.

AJKOER - 22-6-2014 at 19:45

OK, this point may already be clear, but when one hears the words metal, O2, NaCl (which acts as a good electrolyte) in an acidic or basic medium, think galvanic cell ( metal air battery). That is, electrochemistry.

This system is further complicated with some standard chemical reactions involving, for one, the formation of complexes.

Many of the reactions and half reactions have been presented, but I think it is important to understand the nature of/sources of the reactions.

[Edited on 23-6-2014 by AJKOER]

Texium - 30-6-2014 at 09:02

Ok, so I strongly heated my acidic copper(II) chloride until it was the familiar brown anhydrous color, and then added a bit of water, but it turned olive green, a bit different shade than what it was before. I thought that all of the HCl would have boiled off. It certainly looked (and smelled) like it did. Is there something else that could have happened?

blogfast25 - 30-6-2014 at 12:42

zts16:

I've seen this with cupper chloride hydrate before: I own two batches of CuCl2.2H2O and the hue of the green is distinctly different one from the other. Not sure what that is all about.

MrHomeScientist - 30-6-2014 at 13:03

I've also grown separate batches of copper(II) chloride crystals that turned out different colors - one green, one blue. The wiki page on this chemical has a photo of how its solutions change color based on chloride concentration, which looks pretty similar to what I saw. So I imagine the colors in the crystals are due to chloride concentration as well, and the conditions with which they were grown. How about this: try isolating and drying some green solid crystals of CuCl₂. Redissolve in a minimum of distilled water, then allow to slowly evaporate and crystallize. See if these new crystals are more blue in color. If not, repeat recrystallization once more. If it's still green by then, it would seem extra chloride isn't the issue.

Texium - 30-6-2014 at 13:44

Quote: Originally posted by MrHomeScientist

How about this: try isolating and drying some green solid crystals of CuCl₂. Redissolve in a minimum of distilled water, then allow to slowly evaporate and crystallize. See if these new crystals are more blue in color. If not, repeat recrystallization once more. If it's still green by then, it would seem extra chloride isn't the issue.

That's exactly what I did. First I allowed it to dry in a desiccator, then hydrated and recrystallized it, but it was still green. Then I took it and heated it until it was definitely all anhydrous and rehydrated it again, but it is still green. I was completely expecting it to turn blue this time, but no luck with it.
I'll continue to recrystallize it though, just in case.