Precipitated copper powder

Self-made Copper powder (fine precipitation)

You can of course make copper powder yourself, by reducing it straight from solution. Not Zn powder, as the resulting Cu powder is at least the size of the Zn powder. Whereas precipitation straight from a copper II+ solution should yield nanosized Cu particles.

Here was my attempt at this, I boilded a CuSO4 solution (10g) and added 5 g of Vit C (ascorbic acid). The solution turned dark green (from blue), and soon after, copper powder was precipitated.



This is several grams worth. I tried to upscale this, but no success - I am not quite sure how the stoichiometry of the reaction would work out. Also, gas is devoloped during the course of the reaction.
The copper powder is so fine it coated a teflon-coated stirbar with it.

Anyway, this is something you may want to try.

PS what sort of reactions would you use the Cu powder for? Examples?


PS it is definitely Cu, as it is totally unreactive to HCl (Cu2O reacts with it), plus it is very heavy (hence easily washed and isolated).

[Edited on 10-10-2004 by chemoleo]

Hmm, I had a look at it now, and the 5M HCl which contained the copper powder has become very slightly green. I guess that is to be expected - i.e. oxygen aids in dissolving small amounts. Washing this, and adding 35% HCl didnt change things really, except that a slightly green colouration was noticed, which didn't become more.
Compared to this, when I dissolved commercial Cu2O, it would dissolve immediately to form white CuCl.

What is more peculiar that my copper powder, the batch that was left to dry, turned brown coloured - not all, only what was under the surface. So now I have a powder consisting of dark grains together with the copper coloured grains. I suspect the fineness of the Cu was enough to be oxidised by air, in the presence of trace acid and water? A bit strange nonetheless. But then, who has played with nanosized Cu before? When I get back, I will measure whether any current is conducted. I wish I had done that yesterday, as now it contains this brown/black stuff (it's very hard to say, as the particles are so small). Actually I will take half of that batch and try to redissolve it in HCl, I should remove the oxide that way presumably..

Anyway, that method is something I tried way back, when I had restricted amounts of chemicals. I was hoping to make Cu2O then, to turn this into CuCl. But as you can see I never got it to work then.

Reducing CuO with H2? Hmm... and the hassle of heat, hydrogen gas and so on? I am not quite so sure about that! Definitely more hassle than just boiling a solution. But then... H2 reduction is certainly more economical and quantitative!


Edit/Followup: Ah ok, no problem about the dark brown powder. It seems only coated at the surface. When adding 5M HCl, it turned immediately back to the rosy copper colour - and the solution above turned ever so slightly green - indicating a very thin oxide layer! I guess that's to be expected with finely divided copper like that; even copper sheets tarnish after a while.


[Edited on 10-10-2004 by chemoleo]

Yes, it definitely conducts, meaning it is truly copper.
Because the powder so rapidly oxidises in air, I cleaned it with HCl, and then measured the relative conductencies of the supernatant (the liquid (water) containing no copper powder, and the precipitate (copper powder and water). The supernatant had about 500 kOhm, while the precipitate in the same liquid went down to 10-40 kOhm.
Not ultimately a proof, but taken together with the chemical & observational data, I now have no doubt it truly is copper precipitated straight from a Cu2+ solution.

Apart from that... I rubbed the powder onto a surface, and it actually became shiney, as one would expect from a metal powder. But most convincing to me, anyway, is still how quickly the powder settles. It doesn't take 2 seconds. Only a high density compound would behave that way.

Good bye, Cu2O

PS to keep the powder reduced, I'd suggest to dry it in a vaccum dessicator, or over nitrogen/Ar/CO2 etc, and to keep it dry and away from air at all times. To me, the fact that it oxidises so happily actually shows it has a decent surface area (i.e. being very finely divided).

Yes it was still green, dark in my case (I guess I made sure to use an excess ). Copper ascorbate - I'm not sure. It would imply that ascorbic acid is stronger than H2SO4, which surely it is not. Yet the colour change seems to suggest that. It baffled me too.
As to the amount of copper - weigh it. It has a rel. high density, making you misjudge the amounts.
As to gas - it evolved slowly, while I continued to boil it. I can't really say whether this had an effect on the yield.
In fact, I was baffled by the yield alltogether, as I added more CuSO4/Vit C to the existing reaction, but no more Cu seemed to precipitate.
What might help is actually to construct a mechanism of the reaction.
Also, if you continue boiling it down, the colour will go very dark, and a distinct smell is produced - to my nose it smells like decomposing sugars.

I have a feeling that only a certain part of the reaction produces the copper powder, and thereafter, whatever you do, you won't get more. Maybe the product (copper, vitamin C's breakdown products) inhibiting the reaction. I just don't know. Looks like the subject for a PhD thesis


PS good to see someone else is doing some experimenting!

[Edited on 11-10-2004 by chemoleo]

Chemoleo, were you using lab grade ascorbic acid or OTC vit C? After disolving my vit C in water the aqueous layer was yellow so I am assuming some other impurities are in vit C. After filtering the hot dissolved VitC the residue left in the filter paper smelled strongly like citrus and unlike the vitC tablets prior to purification. Hmm, possibly because I was using citrus flavored vitC.

I attempted this reaction again moments ago and I appear to have gotten much more copper powder than before. Last time as I was filtering the citrus impurities out of the Vit C I just had the stem of the funnel leading into the copper sulfate solution. I assume this lead to some cooling of both reactants. However this time I allowed the filtrate from purifing the vit C to go into a 250 mL beaker. The solutions of copper sulfate and vit C were separatly heated to nearly boiling then mixed quickly and boiled again. I recieved much more copper powder than before.

Also, Chemoleo, in your pic of the copper powder, did it look like that upon precipitation and filtering? Or is that pic after you rinsed it with HCl? Mine is much darker after precipitation and filtering.

A more exact way for chemoleo's excellent method

Wow, that's actually quite recent! Should have published it back in 92 or so when I first tried it - teenage production of Cu powder

Rogue - I used Vit C powder, 100%, straight from a drug store. It might indeed help to get rid of citrus smell
The copper powder you see on the picture is the water rinsed powder straight after the precipitation reaction- no HCl involved. I only used HCl to remove the brown coating after I allowed the Cu powder to dry on air. If yours is darker - I wouldnt know why. As long as it does look like copper, and sinks very rapidly to the bottom of the tube/beaker, you should be fine.

Possibly it would. It would be interesting to follow the pH during the course of the reaction.
Possibly the pH will stop the reaction lateron. I think I mentioned above that I had the feeling that not much more precipitation happened once I added more ascorbic acid/CuSO4. Didn't make much sense to me at the time. This may actually explain it. I guess someone (me?) ought to measure the pH during the course of the reaction

Well, since you and I are the only ones who have done this reaction(so far), and I currently only possess broad scale pH paper...

Anyway, I weighed my yields from the two tries at this experiment. The first trial yielded 0.2g of a dark red precipitate, the second trial yielded 0.6g of identically colored precipitate.

The precipitated powder was added to dilute HCl to attempt to remove any oxide coating(and I wanted it shiny). This was unsuccesful, the precipitate retained its origional dark color.

I believe I can attribute such low yields to the hippie brand of vitamin C.
Quote from the bottle
"Each tablet contains;
Vitamin c.........500mg
(ascorbic acid, sodium ascorbate)
Prepared in a specially formulated nutrient base of natural fruit pulp extracts, acerola, rosehips and the complete vitamin C complex of rutin, hesperidin and citrus bioflavonoids. Sweetened with Crystallose(tm)( a complex of phenylalanine, glucose, and cellulose)."
So as I said before, a hippie brand. The tablets do taste good, in fact, having the bottle in front of me while I quoted caused me to eat a couple. I believe the fact that this vit C is ascorbic acid and sodium ascorbate may have reduced my yields due to not all 500mg being ascorbic acid which is necesary for this reaction.
Tomorow I will try to obtain a bottle of (relatively)pure ascorbic acid and try that instead of this brand.

PS. I currently am evaporating some of the green filtrate at room temperature in the hopes that I obtain some sort of green crystalls. I realize it will still be contaminated with the other crud in the vit C tablets, unfortunatly.

[Edited on 14-10-2004 by rogue chemist]

And precipitated copper is somewhat brighter ... at least as long as one uses pure reagents! (i.e. no citrus flavours etc )

As I had never made copper powder in any way I thought I'd give it a try.

I took ~80ml of strong CuCl2 solution (I don't know how strong it was as I was evaporating it down to get the CuCl2) and added few pieces of Al-foil. It boiled vigorously and some dark red powder witch was mostly very fluffy formed on the bottom. I then took a piece of pure Al-bar ( ~40cm x 5cm x 0.4cm) and supported it so that it would sink into the solution as it dissolves.
This method as I noted does not make very fine powder. But it's ok until I get some ascorbic acid as I think I’ll clean this batch with some HCl and close it into a glass vial so it will remain nice and shiny in my shelf.

Would formic acid perhaps prove sufficient for reduction? just for the reason that would entail no extra expense to purchase ascorbid acid, as I have a bottle ofr 40% formic lying around much untouched

Well, I went out and bought the purest vit C I could find OTC. The contaminants are methylcellulose, microcrystalline cellulose, talc, magnesium stearate.

The yield is much higher, and shinier, than with the aforementioned hippie brand I was previously using but....

The copper was bound by one or more of the contaminants, so all the copper appears to be plasticised. It looks, and feels like chewing gum.

Sorry, for the crappy pics, my digital camera is very bad. In the pictures the copper is actually brighter than it appears.

EDIT: I am just going to order some ascorbic acid instead of dealing with these contaminants

Sorry about the double post, just wanted to show another picture.

Rogue - why don't you use lemon juice? It might be cleaner than the stuff you are working with!!

Anyway - on the degree of fineness- tacho/esplosivo you may have noticed that there is very fine stuff, and the more grainy variant. The more grainy variant can, when dry, be crushed, to be very fine. This is when I got a piece of paper shiny with copper dust
Also - I think the overall surface area of the precipitated powder has to be greater by definition than Zn-reduced Cu, simply because reduction is not localised in space to the reductant particles (i.e. Zn). Instead, reduction occurs everwhere at once, i.e. in solution.
What I rather suspect is that the Cu conglommerates, aggregates or whatever, in a loose kind of fashion. This, as it seems from the Chemical Abstracts, above, can be prevented with various agents, here a polymer. Might be a nice idea to do this whole thing with polyethyleneglycol, or even glycerol, to see whether the coarsity of the Cu changes. However, here the hydroxyls might interfere.
Hell, I might give it a go tomorrow

Edit: Rogue, the picture is above 100 kb, so it might be deleted. But I don't know whether that's true for pictures in the upload directory only.

[Edited on 18-10-2004 by chemoleo]

Tacho, I should think that we are dealing with amorpheus aggregates of Cu here, rather than 'crysallised' Cu. I.e. one Cu atom binds randomly to another, in an unordered fashion. This explains why the powder can be crushed easily. While, when it is derived from Zn powder, the particles will at least have the size of the Zn powder, whereby the reaction allows for better alignment of the Cu atoms - so if anything - Zn-derived Cu is more crystalline than Cu precipitated straight from solution.
Anyway, I fully agree with vulture on this.

Nice idea about the dextrin btw. Shame I haven't got any here

Mick, of course one can make Cu that way too - but this is bound to be not quite as simple as with Zn/Vit C. But ultimately that should yield extremely fine metal powder.

Condutive paint powder...

Low yield

A solution of 43.2 g (≈ 245 mmol) ascorbic acid in 140 ml 50 °C warm, distilled water was added to a second solution of 25 g (≈ 100 mmol) copper(II)-sulphate pentahydrate in 80 ml 50 °C warm, distilled water.


↑ Before mixing.

↑ After 30 minutes reaction time (the mix was stirred in this time).

The formed copper powder was filtered off and washed two times with distilled water and once with spirit to give 1.8 g (dry) copper powder. This is 28% of the theoretically possible yield. (Note: scale inaccuracy ±0.2 g)

Maybe higher concentrated solutions of CuSO4 and ascorbic acid were yielding more copper powder. Since I need only some grams Cu-powder to catalyze a reaction, the yield isn't very important for me.


I found easily a Vit C product that consists of pure ascorbic acid, without any contaminants. 100 g costs 1,49 € but it could be somewhere cheaper, since I bought it in the first drugstore, which was on my way.

BTW: German version of this experiment on www.LambdaSyn.org

[Edited on 29-10-2004 by Mephisto]

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[Edited on 7/22/13 by bfesser]

If a am not mistaken aluminum could be isolated from an electrolyte based on DMSO as a solvent. I am not sure what aluminum compound though. Besides, by the time you work it out, the immense investment in reagents would far outweigh what puny yeilds you could possibly obtain.

"Just a thought about the conductivity tests: Cu2O is a semiconductor with photoelectric properties."
Maybe a Cu2O mirror can be created on a suitable surface with slow depositing (low concentration), then playing around with the photosensitive element thus made? Integrate it into an electric circuit with an alarm in it, so the first rays of the sun at sunrise wake you up and you start your day nice and early?

I have been doing some more experiments with the reaction between ascorbic acid and CuSO4 recently. I have pure acsorbic acid now, so no more problems with purity.
Using 5g vitC and 10g CuSO4, a dark precipitate of copper powder is produced(much more than in my previous experiments). However it takes a long time to settle. The green solution left over is what I have mainly been experimenting with.
1. Green solution +NaOH results in an orange/yellow precipitate.
2. Green solution + HNO3 results in a solution the normal Cu2+ colour(light blue)
3. Solution produced in 2 +NaOH results in the same precipitate as in 1
4. Green solution + Zn results in a black precipitate and bubbling (the most confusing result so far)
5. Green solution +Na2CO3 results in a green precipitate

[Edited on 4-9-2005 by rogue chemist]

Bright metallic concentrate

You'll get metallic crystals from cementation (e.g. powdered zinc). Good old fashioned sintering of the colloidial powder (under hydrogen at maybe >600°C) will of course do something, although it may just stick together as a solid lump.

I've got fine copper before, but I can't quite remember why. Maybe it was Cu2O I dissolved in H2SO4 (which disproportionates to CuSO4(aq) and Cu(s)).

Tim

That's pretty cool. Nanozized Cu.. Nice. Never even thought of reducing copper(II) salts with ascorbic acid..
(Seems like one often overlooks those conceptually simple but useful reactions for some reason..)
This could be quite useful and seems v. easy if done right and pure ascorbic acid is used..
Luckily I can get it undiluted it at the grocery-store for a few dimes here.

Though for those that can't...
Shouldn't Oxalic acid work too? I think it's also more widely available as a pure chemical than ascorbic acid, as it's sold as a rust remover.