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Author: Subject: Sodium Chromate from stainless stell 18/10
tetrahedron
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[*] posted on 5-11-2012 at 17:26


not sure, but colorimetrically i would say 3. > 1. > 2.

edit. meanwhile, i found some interesting patents:

electrochemical process using nitric acid and air

chromic acid from chromium hydroxide

chromic acid from ferrochrome

[Edited on 6-11-2012 by tetrahedron]
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blogfast25
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[*] posted on 6-11-2012 at 12:42


Tetrahedron:

Hmmm…without neutralising you cannot be sure what you’ve precipitated.

Remember what I wrote above:

Quote: Originally posted by blogfast25  
This solubility products table:

http://www.csudh.edu/oliver/chemdata/data-ksp.htm

… has Ks = 3.6 x 10<sup>-6</sup> and 2.2 x 10<sup>-20</sup> for the respective solubility products of CuCrO4 and Cu(OH)2.



Since as Cu(OH)2 is far more insoluble than CuCrO4, so in alkaline conditions the former precipitates rather than the latter.

The green in test 1 is due to CuCl<sub>4</sub><sup>2-</sup> (chlorocuprate anions) which the excess Cu<sup>2+</sup> forms with the free chloride ions. Perhaps also in 3.

Remember also that the second part of Weiming’s scheme calls for adding alkali to the washed CuCrO4, to convert it to insoluble Cu(OH)2 and liberate the chromate anions.

Your tests should have been conducted in near-neutral conditions.



[Edited on 6-11-2012 by blogfast25]




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tetrahedron
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[*] posted on 6-11-2012 at 13:41


Quote: Originally posted by blogfast25  
The green in test 1 is due to CuCl<sub>4</sub><sup>2-</sup> (chlorocuprate anions) which the excess Cu<sup>2+</sup> forms with the free chloride ions.

i seriously doubt that. the concentration of Cl- is simply too low, yes the complex could be contributing to the color of the supernatant, but it wouldn't explain the precipitate. moreover, the Cu(OH)2 would have decomposed to CuO by now, which it didn't.

OTOH in 3. no Cl- is present. after standing overnight the precipitate is now completely brown (decomposed Cu(OH)2).

but you're right, i should have at least neutralized the samples first.
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[*] posted on 6-11-2012 at 14:32


Quote: Originally posted by tetrahedron  
i seriously doubt that. the concentration of Cl- is simply too low, yes the complex could be contributing to the color of the supernatant, but it wouldn't explain the precipitate. moreover, the Cu(OH)2 would have decomposed to CuO by now, which it didn't.

OTOH in 3. no Cl- is present. after standing overnight the precipitate is now completely brown (decomposed Cu(OH)2).

but you're right, i should have at least neutralized the samples first.


Yes, I meant the chlorocuprate in the supernatant solution, of course, not as a precipitate. Its green is quite characteristic, very much like what I'm seeing there.

As regards decomposition of Cu(OH)2 to CuO, that would only happen on mild heating. I didn’t see any mention of heat in your report?




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[*] posted on 6-11-2012 at 14:51


Quote: Originally posted by blogfast25  
As regards decomposition of Cu(OH)2 to CuO, that would only happen on mild heating. I didn’t see any mention of heat in your report?

that's right, no heating. from wikipedia (referenced):
Quote:
Moist samples of copper(II) hydroxide slowly turn black due to the formation of copper(II) oxide

this is indeed what happened to sample 3.

i am now in the process of reclaiming the chromate from the precipitate in 1. as per weiming's method, second part. i'll post the results asap.
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[*] posted on 6-11-2012 at 15:03


Tetrahedron, the electrolisis process is not the main process I developed, the main one is the drano fusion of iron oxide/chromium oxide and blogfast can attest that. The electrolysis process was a side process and you cant precipitate copper chromate without neutralizing the solution. Secondly your claim that it don't work goes with my original statement: "It take a large amount of time to get a product".

I can prove it work using experimental evidence, as well theoretical evidence. I originally precipitated the salt by neutralizing with HCl and precipitated using Zinc or lead salts. I did not make any claim that the product was enough pure to be used as it.

Hopefully you don't take this comment as offensive.



[Edited on 6-11-2012 by plante1999]




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[*] posted on 6-11-2012 at 15:36


not at all, i really admire your experimentalism and your 'industrial processes', but you need to provide more detail beyond 'abstract style' if you expect others to successfully follow your procedures. also it would make it so much easier to gauge the yields.

i tried to replicate your electrochemical procedure with the details you provided but the result was unatisfactory (this was clear despite the flawed test). i'm not saying it doesn't work, i just wasn't able to confirm it given the data, in the spirit of 'peer review'.

as to the 'drano' method, as i mentioned upthread you told me to 'use stoichiometry' but i doubt your drano has the right stoichiometry for endogenous oxidation of the chromium, so i wonder whether you apply an external stream of air or an excess of drano to make it work or just a looot of time. anyway the molten approach is well documented and certainly works.

i have another 'one pot' electrochemical method in mind that might salvage your idea and make it more efficient, stay tuned.
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[*] posted on 6-11-2012 at 16:26


I agree that the drano method doesn't use stoichiometry but at that moment (and at the moment also) I didn't had nitrates, but my drano contained up to 40% Nitrate by weight. It worked very well in the sense that at that time I was not (and I'm still not now, but its better then at that time) very well equipped.



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[*] posted on 6-11-2012 at 17:09


tetrahedron, you wiki fact about Cu(OH)2 turning black over time when moist... I have had a solution of exactly that sitting for near 8months in an airtight container and not a speck of black. it was my wash up after a few weeks of Cu(OH)2 production leftovers for the black oxide. filled a gallon to let sit and was gonna decant it off but got bussy with work. settled to 4inches of light blue ppt under like ~8" of water. no change in 8 months. but with heat it changes no prob.
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[*] posted on 7-11-2012 at 00:52


Quote: Originally posted by tetrahedron  
i am now in the process of reclaiming the chromate from the precipitate in 1. as per weiming's method, second part. i'll post the results asap.

here goes. the washed CuCrO4 precipitate from sample 1. suspended in some water:
1c.jpg - 84kB

this was heated to boiling and two drops NaOH 3.8M were added: it immediately turned black. after cooling and standing overnight, the filtrate was recovered as a pale yellow liquid. upon further boiling a dark yellow solution remained:
1d.jpg - 79kB

[Edited on 7-11-2012 by tetrahedron]
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[*] posted on 7-11-2012 at 06:29


Quote: Originally posted by violet sin  
tetrahedron, you wiki fact about Cu(OH)2 turning black over time when moist... I have had a solution of exactly that sitting for near 8months in an airtight container and not a speck of black. it was my wash up after a few weeks of Cu(OH)2 production leftovers for the black oxide. filled a gallon to let sit and was gonna decant it off but got bussy with work. settled to 4inches of light blue ppt under like ~8" of water. no change in 8 months. but with heat it changes no prob.


That's also my experience: no change over time w/o heating. It's probably dependent on other conditions.

@Tetrahedron:

That shows neatly and conclusively that the method works, as far as I'm concerned. Case closed.

[Edited on 7-11-2012 by blogfast25]




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[*] posted on 7-11-2012 at 08:52


Excellent! So it was my reagents at fault, most likely not pH-balanced. Good to know, in the meantime I'll have to make some more copper sulfate.



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[*] posted on 7-11-2012 at 10:04


actually you should be getting a precipitate no matter the pH, so either no copper or no chromate is present. i don't recall your procedure, but maybe your chromium only reached oxidation state +3, and what makes the solution yellow is actually the iron.
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[*] posted on 7-11-2012 at 11:13


Well, I suppose that's possible, but it was reduced to a green color with sulfuric acid and isopropyl, which is almost entirely Cr (III). Also, color matches exactly with the test tubes posted above. So, I would say either there is heavy dilution of my chromate, in which case I need to boil it down, or the pH is indeed messing with my precipitate somehow. (There is probably plenty leftover NaOH / Na2CO3 in my chromate, and there was almost definitely leftover sulfuric acid in my copper sulfate.)



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[*] posted on 7-11-2012 at 11:33


if theory serves me right, as soon as you have copper ions and chromate or hydroxide or carbonate in solution, the corresponding insoluble copper compounds should precipitate immediately, quantitatively consuming the Cu2+ or the anions (whichever is less). someone better informed than me could maybe point out the effect of extreme pH on the solubility of CuCrO4?
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[*] posted on 7-11-2012 at 13:48


Quote: Originally posted by tetrahedron  
if theory serves me right, as soon as you have copper ions and chromate or hydroxide or carbonate in solution, the corresponding insoluble copper compounds should precipitate immediately, quantitatively consuming the Cu2+ or the anions (whichever is less). someone better informed than me could maybe point out the effect of extreme pH on the solubility of CuCrO4?


In the case of 'extreme pH' (i.e. very high pH) Cu2+ could manifest its amphoterism and cobalt blue cuprate (Cu(OH)<sub>4</sub><sup>2-</sup>;) anions could form. Not very likely, judging EC1's conditions superficially, IMHO...

[Edited on 7-11-2012 by blogfast25]




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[*] posted on 7-11-2012 at 15:47


I'm going to separate nichrome soon with bleach and NaOH. Bleach is going to be 6% while NaOH will be made from Ca(OH)2 and NaHCO3.

Anyone have experience with more dillute solutions? SO I don't waste my time if it wont dissolve fully. Maybe even Ca(OH)2 could be used with bleach I dunno.
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sad.gif posted on 7-11-2012 at 16:02


Quote: Originally posted by tetrahedron  
yesterday i took some of the sludge that i got when "electrodissolving" stainless steel in a Cl- containing electrolyte and poured it into a fresh solution of NaOH and NaCl (about 1 spoonful each in 400ml tap water), which i then electrolyzed for a couple hours with a 5mm diameter graphite "gouging rod" anode, initially @ ~7V, soon reduced to 4.4V & 3A in order to reduce the loss of chlorine. i noticed a faint chlorine smell during the procedure (it was done outside), as well as after unplugging the current. even after standing overnight i still noticed some bubbling, which i can't explain.

i'm running this procedure anew, with a few modifications: i dried (calcined) the mixed oxide slurry to a powder, i switched to a copper cathode instead of stainless steel, and diluted the electrolyte with distilled water instead of tap water. the electrolysis has been running for hours now without evolution of Cl2 and without apparent change in the electrolyte (other than a lot of bubbling, which i presume consists of only H2 and O2).
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[*] posted on 8-11-2012 at 16:25


more food for thought, freshly baked:

1. ions of Cr (III) are oxidized to hexavalent state with (NH4)2 S2O8 in the presence of AgNO3 (catalyst) in acid solution
ammonium persulfate is prepared by electrolysis of ammonium sulfate in sulfuric acid; this could be a viable electrolyte for the electrochemical oxidation of Cr3+.

2. electrodialysis to regenerate hexavalent chromium baths

3. Aspects of Electrochemical Regeneration of Chromium Containing Solutions from Metal Finishing Industry
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[*] posted on 13-11-2012 at 12:13


Looking around for more insoluble chromates, I spied calcium's as having low solubility (4.5g/100mL at 0 C, 2.25g/100mL at 20 C). I have plenty of calcium chloride, so should I try this?




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[*] posted on 13-11-2012 at 12:46


Quote: Originally posted by elementcollector1  
Looking around for more insoluble chromates, I spied calcium's as having low solubility (4.5g/100mL at 0 C, 2.25g/100mL at 20 C). I have plenty of calcium chloride, so should I try this?


4.5 g/100 ml isn't that low. But you could 'release' the chromate again by treatment of the CaCrO4 with the right amount of Na2CO3 (Or K2CO3) because CaCO3 is so insoluble...

4.5 g/100 ml corresponds to a Ks of about 8 x 10<sup>-2</sup> (ignoring any hydrolysis), much much higher than for CuCrO4.

[Edited on 13-11-2012 by blogfast25]




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[*] posted on 9-1-2013 at 11:45


I electrolyzed some stainless steel and filtered the solution.

I treated the precipitate with NaOCl and Na2CO3, filtered again.

It seems I can treat HCl to get rid of the carbonate, but I will still have NaCl in the solution with the dichromate.

I want to prepare some Cr+3 salts. My understanding is that will oxidize alcohols to produce ketones. Can I use Isoproply alcohol or is ethanol or methanol preferred.
Then I could add NaOH to precipitate the Cr hydroxide and then treat it with an acid to create the salt.

Does this seem workable or am I missing something?
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[*] posted on 9-1-2013 at 12:48


Quote: Originally posted by michael971  

I want to prepare some Cr+3 salts. My understanding is that will oxidize alcohols to produce ketones. Can I use Isoproply alcohol or is ethanol or methanol preferred.
Then I could add NaOH to precipitate the Cr hydroxide and then treat it with an acid to create the salt.

Does this seem workable or am I missing something?


Primary alcohols like methanol and ethanol are easier oxidised than secondary ones like IPA. Use 'methylated spirits' and you can't go wrong. Heat the mixture dichromate + alcohol on a steam bath for quick conversion: it'll go from amber/yellow to green/blue.

Then neutralise most of the excess acid with strong NaOH, then add Na2CO3 to avoid the Cr3+ to go into solution as chromite (CrO2(-)). Cr(OH)3.nH2O should precipitate. Filter and wash. Pronto.




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cool.gif posted on 6-3-2018 at 13:13
Potassium dichromate from stainless steel


These methods are ardous . They are not as simple as they seem to . I have brought a efficient and effective method to make potassium dichromate from stainless steel .
Step 1: take stainless steel wool used for washing utensils and dissolve it in 50 ml of 6.3 molar HCl. Wait for completion of reaction and take out the residue stainless steel out.
Step 2 : dissolve 16.8 g Na2CO3(anhydrous) in 100 ml H2O and pour it into that solution.
Step 3 : separate the ppt by any method you like (sedimentation - decantation or filtration) .now take the ppt and heat it in a steel bowl to convert it into oxides
Step 4: take the mix with 9.7g KNO3 and 2.6g Na2CO3
Powder it ,mix it and heat it in the bowl at high temp .
Step 5 :.dissolve the mass in hot H2O and filter.
Step 6 : add 5.2g NH4Cl in the filtrate and boil it to decompose all the potassium nitrite formed
Step 7: concentrate the solution to less than 10 ml by boiling .
Step 8: acidify the solution with 3.8ml 6.3M HCl
Step 9: cool the solution to 0℃
Step 10 : filter the orange crystals and wash them with ice cold H2O
Maximum yield of K2Cr2O7 will not be more than 3.6g.

Soon going to upload a video on this at my channel must watch .




Subscribe to my youtube channel named akhil the chemist. search it and you will get it this channel is unique .
https://www.youtube.com/channel/UC9GD00yhAoKajgjRWvqyH-w
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