Sciencemadness Discussion Board

Cu (III) and other transition metals in high oxidation states

UnintentionalChaos - 14-4-2008 at 15:18

In a recent thread on Periodic acid, the topic of copper (III) came up and I have some rough experience with what I think may have been Cu2O3. I am particularly interested in copper chemistry. I want to see if anyone can replicate my experimental results. In a test tube with a little water, dissolve a small spatula of copper sulfate. Add a few mls of H2O2 solution (I was using 15%). A pale blue gelatinous looking precipitate forms. Add a few drops of ammonia solution (I was using 5% I believe). In my case the solution darkened intensely and reddish brown flakes were visible floating around. If I'm not mistaken, this material is copper (III) oxide. I observed a similar precipitate of gelatinous blue material from direct addition of bleach to copper sulfate solution but did not get the dark compound to form. I wonder if this material is Cu(OH)3 If so, then some very interesting experiments with Cu (III)...a generally unheard of oxidation state (Other than in superconductors)...are well within our grasp. I do not have any persulfate and am currently at school and away from my lab.

Copper (III) by way of persulfate oxidation in alkaline medium either with or without periodate as a complexing agent would make a very nice addition to woelen's experiment page here on Ag (III) and Ni (III/IV): http://81.207.88.128/science/chem/exps/Ni+persulfate/index.html

I also stumbled upon this patent while looking for info on periodate complexes. [The link I had posted doesn't seem to work, so here is an extract of the important material]

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Said trivalent silver complexes were subsequently evaluated as to their efficacy in killing gram positive and gram negative bacteria in algae in accordance with the EPA protocols for swimming pools, which require 100% kills of bacteria within ten minutes. The compounds far exceeded the bacteria requirements at concentrations of one PPM or less of silver. They were evaluated with and without persulfate salts at 10 PPM and were effective without persulfates as bactericides.

The complexes were then evaluated with salt concentrations as high as 10% without precipitating halide.

The complexes, which were colored from deep orange to brown and maroon, were left exposed in clear glass bottles for three months with constant exposure to daylight. The complexes were stable and did not decompose to silver.

Ag(III) complexes were applied to human skin in concentrated form containing as much as 5,000 PPM silver without any silver staining of the skin whatsoever.

Of all the Ag(III) complexes prepared, the easiest to prepare were periodate complexes. Accordingly, a particular Ag(III) periodate complex was selected for evaluation against algae and it proved effective.

The particular Ag(III) periodate selected was prepared by the action of potassium hydroxide on tetrasilver tetroxide (Ag 4 O 4 ) and is depicted by the following reaction: Ag 4 O 4 +6KOH+4KIO 4 =2K 5 H 2 [Ag(IO 6 ) 2 ]+Ag 2 O+H 2 O

An Ag(III) periodate was prepared by calculating the amounts of reactant necessary to form the periodate complex according to the aforementioned reaction equation involving KOH, 0.15 grams; tetrasilver tetroxide 1.0 gram and potassium periodate 2.0 grams. The KOH was first dissolved in 20 ml. of distilled water and the subsequent ingredients were added, and the entire mixture was heated and kept at 65 degrees C. for two hours. At the end of that time, the supernatant liquid, which had a rich orange-maroon color, was separated and was filtered and submitted for analysis to an independent laboratory, after it was diluted to a total volume of 100 ml. in a volumetric flask. The resulting solution assayed 3290 mg./L Ag and represented a 76% yield of Ag(III) periodate based on the aforementioned reaction stoichiometry.

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I wonder if analogous stable complexes of Ni(IV) and perhaps Fe(VI) as well as other transition metals in high oxidation states exist.

I added this same paper in the other thread, but I'll attach it here for reference.

[Edited on 4-14-08 by UnintentionalChaos]

Attachment: Copper (III).pdf (157kB)
This file has been downloaded 2270 times

JohnWW - 14-4-2008 at 16:40

Quote:
Originally posted by UnintentionalChaos
I also stumbled upon this patent while looking for info on periodate complexes.
http://www.freepatentsonline.com/5223149.html.

Page not found.

woelen - 15-4-2008 at 11:53

I downloaded the paper and tried to reproduce what is described in the paper, but no success at all.

I used reagent grade chemicals (except the bleach, which is ordinary household bleach, but without additives).

Experiment 1:
------------------
Dissolve some copper sulfate in water.
Dissolve some NaIO4 in 5% bleach
Mix the two solutions. A lime green precipitate is formed. Initially it is really green, but in ten seconds or so, it shifts to lime green, a nice bright color. Not any of the red/brown complexes, mentioned in the paper. There also is a faint odour of free chlorine, which is not present in the bleach solution.


Experiment 2:
------------------
Dissolve some copper sulfate in water.
Dissolve some NaIO4 in water.
Mix the two solutions (excess NaIO4). Again, a lime-green precipitate is formed.
Add some bleach to the precipitate. The precipitate does not change color.
Heat the liquid with the precipitate. No visible changes in color, the precipitate becomes a little bit more compact and more granule-like and less flocculent.




Experiment 3:
------------------
Dissolve some copper sulfate in water.
Dissolve some NaOH in water and let this solution cool down.
Mix the two solutions. A blue precipitate of Cu(OH)2 is formed.
Dissolve some Na2S2O8 in water.
Add this solution to the blue precipitate. The blue precipitate slowly turns black. After a somewhat longer time, tiny bubbles are produced (most likely oxygen).
Dissolve some NaIO4 in water.
Add this solution to the black precipitate. The color of the black precipitate shifts to dark green (like dark green olives). I think this dark green color is due to a mix of the lime green periodate compound, as shown above in the picture, and the black precipitate.


Experiment 4
-----------------
Dissolve some copper sulfate in water.
Add some bleach to this solution. This results in formation of a green/blue precipitate. I think this is nothing special, just copper(II) hydroxide, probably contaminated with chloride ions. A similar precipitate can be obtained when a solution of a mix of NaOH and NaCl is added to a solution of copper(II) sulfate. You get a copper(II) hydroxide, contaminated with copper(II) chloride/hydroxide.
Heating results in blackening of the precipitate. This black stuff probably is CuO. A similar blackening also occurs when simple Cu(OH)2 is heated in water.


The only thing I can confirm is the reddish (rather dark, I would call it brown) material when H2O2 is added to an alkaline copper(II) compound. I have read about this, but according to what I have read, this is not copper(III) oxide, but a mix of copper(II) oxide and copper(II) peroxide. I have a procedure for making CuO2, which is copper(II) peroxide. I never cared for making this compound, because it is unstable and in a few weeks looses all its excess oxygen, leaving plain CuO behind.


Altogether, no interesting results were obtained.

I'm really wondering how they made those red/brown complexes. I could not make them.

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[Edited on 7.1.14 by bfesser]

Axt - 15-4-2008 at 12:30

This is the example from US5336416.

A Cu(III) periodate complex was prepared as follows 2.0g of cupric sulfate pentahydrate were dissolved in 150 cc. of distilled water, 9.2g. of potassium periodate (reagent grade) were added to the solution and 24 cc. of 45% KOH was then added. This was followed by the addition of 8.4g. of sodium persulfate to the solution which was heated to 90°-95° C. The solution was subsequently cooled to room temperature and filtered. After removal of insoluble solid matter, the remaining solution was diluted to 200ml. volume in a volumetric flask and assayed via AA spectroscopy. The yield of Cu(III) periodate was 30%. Aliquots of this solution were taken for subsequent pathogenic evaluations.

Klute - 16-4-2008 at 06:01

Complexation of a Ni II salt with dithiolene ligands, such as dmit 2- ( 2-thione-1,3-dithiole-4,5-dithiolate) and oxydation with I2 in acetone produces Ni IV complexes.

The dmit2- ligand can be produced by reaction of CS2 and qsodium metal in presence of DMF, under inert atmospher, and complexed with a tetraalkylammonium salt to yield a stable precursor. Benzoylation offers a product that liberates the ligand upon reaction with sodium methoxide, addition of a Ni II salt produces the II complexe. Oxydation with I2/I- in acetone offers the IV complexe.
All the reactions except the oxidation must be done under strict inert conditions, granted this isn't something easy to do in amateur settings.

References:

G.Steimecke et al., Phosphorus, Sulfur 7, 49 (1979)
N.Svenstrup, J.Becher, Synthesis, 3, 215 (1995)
CS.Wang et al. , Synthesis, 11, 1615 (1998)

[Edited on 16-4-2008 by Klute]

Taoiseach - 17-4-2008 at 07:54

"Dissolve some copper sulfate in water.
Dissolve some NaOH in water and let this solution cool down.
Mix the two solutions. A blue precipitate of Cu(OH)2 is formed."

Real Cu(OH)2 is only formed if excess NaOH is used. One would expect Cu(OH)2 to look like basic copper carbonate (turqouise to pale blue) but it has quite a brilliant blue color. Without excess NaOH only the turqouise basic sulphate is produced:

2[Cu(H2O)6](2+) + 2OH- + SO4(2-) ---> Cu(OH)2 * CuSO4 + 12H2O

There is a "nickel peroxide" NiO2 that can be made from NiCl2 and NaOCl. I have a prep laying around if anyone is interesting. I never cared making it because yield is low and a HUGE excess of NaOCl is required.

Also there is a prep for pyridine silver(III) persulphate in Brauer: Inorganic Preparations. Its blood-red and non-explosive. The ethylenediamine silver(III) persulphate might be worth a try.

Klute - 19-4-2008 at 09:13

Oh, BTW, found there an org syn procedure to form the dmit(COPh)2 i mentionned above:

http://www.orgsyn.org/Content/pdfs/procedures/CV9P0203.pdf

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[Edited on 7.1.14 by bfesser]

woelen - 20-4-2008 at 06:44

I tried the method of Axt, and this gives striking results!

* I took 40 mg of solid CuSO4.5H2O and dissolved this in 3 ml of water.
* I added 180 mg of solid KIO4. This results in formation of a small amount of the lime green precipitate, which I have shown in the picture above. Most of the KIO4 remains undissolved.
* In a separate test tube, I dissolved 250 mg of solid KOH in a small amount of water, and this solution I added to the first test tube. The result is that all KIO4 dissolves, and the solution becomes deep blue. Apparently some copper(II)/periodate complex is formed. The solution also becomes clear.
* Finally, I added 170 mg of solid Na2S2O8. This partly dissolves. The solution becomes a little darker blue and somewhat turbid. Nothing special.

Now, I started heating the liquid. The result of this is remarkable. Slowly, the color shifts (through greens and browns of all kinds of shades) to a beautiful bright red, somewhat like a saturated solution of bromine water, but the color is a little more intense. This deep red color is obtained when the solution is near boiling.

After this, I left the solution standing and it does not change any further. More experiments and a wrtie-up on this will follow. I've never seen a copper-complex which is deep red. That color really is remarkable for copper. I'm quite sure that this is the copper(III) compound, which is mentioned above!

16MillionEyes - 21-4-2008 at 06:12

You're great at building suspense aren't you Woelen? Where is your usual nice picture to illustrate the exotic compound! :D

woelen - 25-4-2008 at 11:09

Here follows the write-up with pictures. This ends the suspense :D

http://woelen.homescience.net/science/chem/exps/CuIII/index....

[Edited on 25-4-11 by woelen]

JohnWW - 25-4-2008 at 15:38

Now that you have mastered Cu(III), has anyone obtained higher oxidation states, like Cu(IV) and (V)? Au(V) has been made as AuF5 and AuF6-, so there is a possibility. Also, I wonder if anyone has managed to obtain Zn(III), Cd(III), or Hg(III), utilizing a d electron as well as the two s electrons.

UnintentionalChaos - 25-4-2008 at 17:20

The problem with the high oxidation state flourides is that synthesis usually (maybe always) requires you to expose, if not heat the lower oxidation state compound in fluorine gas.

Beautiful experiment woelen. I am very impressed as always with both the result, pictures, and writeup....it's a shame that the silver (III) complex couldn't be produced by an analagous procedure though.

woelen - 11-6-2008 at 12:59

I now isolated the copper(III) compound as a solid. It is a dark red solid, which slowly dissolves in water, giving yellow/brown aqueous solutions when dilute, and red/brown when more concentrated:



This compound most likely is K3[Cu(H3IO6)2(OH)2], potassium bis-orthoperiodato dihydroxo cuprate(III).

The sample, shown above, is approximately 200 mg.


[Edited on 11-6-08 by woelen]

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Jor - 11-6-2008 at 13:37

:o

Time to isolate it as well!
How exactly did you do it? Want to do it right in one go, not wasting expensive periodate.

EDIT:

Nevermind, you edited your webpage :P
Don't you have to rinse or recrystallise the product? It might still contain quite some contamination from the yellow/white compound, although maybe not visible.

[Edited on 11-6-2008 by Jor]

woelen - 11-6-2008 at 13:52

Rinsing is done, very briefly, while the material still was in the test tube. The crystalline matter only dissolves VERY slowly.

A recrystallization could be an option, but I only have 200 mg, and I think that will waste a lot of the material. Keep in mind that the formation of these crystals took several days (my second crop even took more than 1 week) and such slow formation of crystals usually results in very pure material. The material is good enough for further experimenting and research, and it definitely is good enough for demonstrating its properties.

Right now, I have a test tube with a silver complex standing, and it forms ruby-red crystals, not nearly as dark as the stuff I have from the copper. The silver(III) complex does form as well, but its formation takes weeks :o Slowly, very slowly, my silver/periodate solution turns darker and darker, it started off light yellow, now it is red like wine.

Engager - 14-7-2008 at 17:18

Try to add some solid sodium peroxide to blue cuprite solution, this results in red precipitate of copper (III) oxide Cu2O3. Copper (IV) can be made in complexes, for example with biguanide.

Xenomorph - 25-4-2011 at 02:41

Copper (III) periodate is easy one, but with Ag(III) I got extremely small yields.

For Copper(III) I first dissolved 2g of CuSO4*5H2O in 150ml of water and added 10g of KIO4 (green precipitate formed, if I remember corectly).
Then I slowly with stirring added 10g of KOH solution in 25ml of water and precipitate dissolved forming intense blue solution.
Finally I added 9g of Na2S2O8 and heated solution until it turned black and stained walls of flask intense yellowish-brown.
Then I poured solution in beaker for crystalization (there was none of undissolved solid). After about one day I poured solution in another beaker and discarded the light yellow crystals of impurities (when I previously made Cu(III) periodate, I learned that impurities crystalizes first). Repeated this step until black spheres of Cu(III) periodate crystals started to form. After one more day solution became transparent and thick layer of dark needle-like crystals were visible. Poured off solution and let the crystals dray on filter paper. Then I rubed them through fine mesh to fet rid of few remaining crystals of impurities (they are much larger). Got large amount of really pure Cu(III) periodate, no need for recrystalization. I guess they are K3H4[Cu(IO6)2]*4H2O. Am I right?

Image: Diperiodatocuprate(III) on the left, Diperiodatoargentate(III) on the right.

P1010468cr (Small).jpg - 42kB


[Edited on 25-4-2011 by Xenomorph]

Xenomorph - 25-4-2011 at 03:09

Diperiodatoargentate(III) was tricky one.
I tried to follow this protocol:
(Hieremath DC, Hieremath CV, Nandibewoor ST. 2006. Oxidation of paracetamol drug by a new oxidant diperiodatoargentate(III) in aqueous alkaline medium. E-Journal of chemistry, 3(10):13-24.)
Quote:

Preparation of DPA:DPA was prepared by oxidizing Ag(I) in presence of KIO4 as described elsewhere: the mixture of 28g of KOH and 23g of KIO4 in 100cm3 of water along with 8.5g of AgNO3 was heated just to boiling and 20g of K2S2O8 was added in several lots with stirring then allowed to cool. It was filtrated through a medium porosity fritted glass filter and 40g of NaOH was added slowly to the filtrate, whereupon a voluminous orange precipitate agglomerates. The precipitate is filtered as above and washed three to four times with cold water. The pure crystals were dissolved in 50cm3 water and warmed to 80C with constant stirring therebay some solid was dissolved to give a red solution. The resulting solution was filtered when it was hot and on cooling at room temperature, the orange crystals seperated out and were recrystalized from water.


I tried 2 times with 10x smaller amounts of reagents and sodium persulfate instead of potassium one. Upon reaction with persulfate, solution turned intense dark orange but large amount of broun precipitate remained undissolved.
Following the protocol I got extremely low yields of orange-red crystals. They seems to be very stable at prolongated exposure of light.
Can the problem be with persulfate?
(K,Na,H)7[Ag(IO6)2]*nH2O /// Any idea of exact chemical formula?

vano - 1-6-2022 at 07:29

Copper(III). K3[Cu(H3IO6)2(OH)2]

283707564_418979096744707_6160515379421631002_n.jpg - 294kB

clearly_not_atara - 1-6-2022 at 15:01

Very interesting! Looks like dark brown sugar. How was it made?

vano - 2-6-2022 at 05:41

You can read here:
https://woelen.homescience.net/science/chem/exps/CuIII/index...