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Poppy
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[*] posted on 10-3-2012 at 13:08
Dissolution of Chromium by conc. HNO3


Hello,

Does you people have some advice on the dissolution strenghts of nitric acid against chromium metal?

I tried dissolving chromium based stainsless steel with dilute (53%) nitric acid and it didn't break the passivation layer, neither by heating the acid. Addition of sulphuric acid yields the same result, although working with sulphuric acid only in solution did attack the steel vigorously when heated.
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[*] posted on 10-3-2012 at 14:42


May I have your pardon.
Still, I am not sure if all this is correct relating to stainless steel.

From WiKi:

Passivation
Although chromium (Cr), iron (Fe) and aluminium (Al) readily dissolve in dilute nitric acid, the concentrated acid forms a metal oxide layer that protects the metal from further oxidation, which is called passivation. Typical passivation concentrations range from 20-50% by volume. (See ASTM A967-05) The metals which are passivated by concentrated nitric acid are Iron, Cobalt, Chromium, Nickel, and Aluminium.[4]

Still it does not counts for the very acid rich FNA, any disposals?


[Edited on 3-10-2012 by Poppy]

[Edited on 3-10-2012 by Poppy]
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[*] posted on 10-3-2012 at 15:49


i never could dissolve stainless steel with my home made nitric either.always resorted to electricity.
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[*] posted on 10-3-2012 at 16:20


I did once set corrosion to overrun by mixing the 3 common inorganic acids together, i.e, hydrochloric, sulphuric and nitric. The pickled sheet piece showed corrosion pits and none of those black substrates sticking to the metal. It's appealing I know, but fast and operates at room temperature. I check and publify further details.

C,

Poppy
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[*] posted on 10-3-2012 at 17:09


I've succesfully dissolved SS in 37 % HCl, quite fast and w/o problems.



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[*] posted on 14-3-2012 at 08:38


Blogfast25:
Guess how even higher concentrations of hydrochloric acid would behave on the purpose? Any means on trying that?
------------------------
A method dealing with sulfuric acid only has been tried again. The passivation layer absolutely is the major problem here. Suppose a stainless stell plate, as for sulphuric acid, I could break the layer at ~120°C IF the whole piece of ss could hold this temperature, i.e., applying the solution locally will not corrode the plate. Once broken the corrosion starts and keep until 90°C. Temperature however should not cool below 110°C to prevent the passivation from forming again. The sulfuric acid solution was no less than 10 molar (10 mol/ L), yes, 850g/ L H2SO4 (50% by volume).
Dissolving then turns even more problematic with this acid because as acid molarities goes down the reaction stops:mad:
But a cool and well-known effect was observed after cooling the part: the common ion effect. Salts of ironII and chromium III could be retrieved in anhydrous form from the slurry, thus being that the supernatant can be reused to dissolve the plate and so on. But this as presented requires expendure of energy sources and a heat resistent vessel to be put on oil baths or directly upon flame, the later being so a 1L borosilicate glass to ensure cracking will not occur.


[Edited on 3-14-2012 by Poppy]
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[*] posted on 14-3-2012 at 10:46


Quote: Originally posted by Poppy  
Blogfast25:
Guess how even higher concentrations of hydrochloric acid would behave on the purpose? Any means on trying that?


[Edited on 3-14-2012 by Poppy]


37 % HCl is about the theoretical maximum concentration in STP conditions.




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[*] posted on 14-3-2012 at 11:17


Strange I think I've seem HCl concentrations up to 90%, you just had to keep the lid on the jar to stop HCl from scaping the solution, hows that if I manage to create a 48% solution from aqueous CaCl2 and slightly dilute H2SO4 wont it remain in there??

Nvm, just figured CaCl2 doesn't gives HCl that easy, it need a pressure container, nothing Imma going to build right now, not right not...

[Edited on 3-14-2012 by Poppy]
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[*] posted on 14-3-2012 at 12:55


Quote: Originally posted by Poppy  
Strange I think I've seem HCl concentrations up to 90%, [...]

[Edited on 3-14-2012 by Poppy]


No, sorry poppy but you're 'seeing things'. Concentrated HCl is basically 37 w% HCl.

Careful breathing in those chemfumes! ;)

[Edited on 14-3-2012 by blogfast25]




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[*] posted on 24-3-2012 at 17:14


For curiosity, first make some HOCl by adding vinegar to Bleach (NaClO/NaCl/NaOH). This produces dilute HOCl in a solution of Sodium acetate. Or, add a very dilute solution of NaHSO4/H2O to Bleach. Cool to extract the Na2SO4, leaving HOCl in a salt solution. For those seeking higher strength and/or purity, extract the HOCl via acetone.

A similar product in terms of reaction characteristics, I would argue, is obtained by combining HCl and dilute H2O2.

Next, add the chromium based stainless steel to HOCl mixture and see if there are any bubbles. The reaction with Iron is surprising even with dilute HOCl, and with chromium based stainless steel, I wouldn't be too surprised if you reported something interesting.

[Edited on 25-3-2012 by AJKOER]
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[*] posted on 24-3-2012 at 17:26


According to this info(I know Wikipedia's overused, but that's the first thing that comes up) http://en.wikipedia.org/wiki/Chromium
Chromium forms a passivating layer of oxide against corrosion, so when you use nitric acid, which is oxidizing, it won't go through. Same thing with HOCl. The way to go is probably to dip your steel in oxalic acid, then, when the oxide layer comes off, immediately place in a non-oxidizing acid (like dilute sulfuric, NaHSO4(aq),HCl, etc.)
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[*] posted on 24-3-2012 at 18:27


Is there a way to separate chromium from the mix? I have a LOT of scrap stainless, and a bit of HCl (or oxalic, if needed).



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[*] posted on 26-3-2012 at 13:23


That is exactly the point.
The passivation layer regenarates rapdly once the dissolving bath goes trought saturation. Twice excess acid is generally required to dissolve a given ammount of the ss, ant that with heating. The trick here we might guess is to create a bath wherein a catalyst constantly breaks passivation, allowing the acid to keep doing its job longer.
Chromium ore seems more fruitful than dissolving the own ss IMHO, because Cr2O3 is expected to react easier.
Once dissolved chromium can be pushed into Na2Cr2O7, decanted and put to react with (NH3)2SO4, where ammonium dichromate shall coprecipitate with sodium sulfate. Then both are placed in a heated bowl and thermal decomposition of the ammonium dichromate takes place, leaving Cr2O3 together with soluble impurities.
The Cr oxide is then washed and, as wiki says, its hygroscopic, which means it probably turns into Cr(OH)3 once it contacts water. A mass verification will probe the reversibility of the hygroscopicity upon heating.

Might oxalic acid be this accelerator?
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[*] posted on 26-3-2012 at 19:46


Excellent! Thanks so much, but still a few questions:
1) How do I get the sodium dichromate made? (I know it's a double salt, but I also don't know anything about double salts.)
2) Can the ammonium sulfate be bought pure as fertilizer, or made from a sulfate salt (epsom?) and an ammonium compound?
3) Are you sure Cr(OH)3 forms? From what I know, the hydroxides of transition metals tend to turn into some oxide over time and Cr(OH)3 isn't well documented.
Also, I do have quite a bit of oxalic on hand. Time to experiment!:D




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[*] posted on 27-3-2012 at 08:04


I could find this out of google:
http://www.sciencemadness.org/talk/viewthread.php?action=pri...
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[*] posted on 27-3-2012 at 13:47


I don't have sulfuric acid or methylated spirits.
Also, when trying the SS dissolution into concentrated HCl, I get amounts of two products that are not very soluble: Lime green, floating stuff, which I assume is a hydrate of chromium chloride; and a darker green stuff on the bottom (oxide?) Does anyone know what this is?




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[*] posted on 2-4-2012 at 13:45


Quote: Originally posted by Poppy  
That is exactly the point.
The passivation layer regenarates rapdly once the dissolving bath goes trought saturation. Twice excess acid is generally required to dissolve a given ammount of the ss, ant that with heating. The trick here we might guess is to create a bath wherein a catalyst constantly breaks passivation, allowing the acid to keep doing its job longer.
Chromium ore seems more fruitful than dissolving the own ss IMHO, because Cr2O3 is expected to react easier.
Once dissolved chromium can be pushed into Na2Cr2O7, decanted and put to react with (NH3)2SO4, where ammonium dichromate shall coprecipitate with sodium sulfate. Then both are placed in a heated bowl and thermal decomposition of the ammonium dichromate takes place, leaving Cr2O3 together with soluble impurities.
The Cr oxide is then washed and, as wiki says, its hygroscopic, which means it probably turns into Cr(OH)3 once it contacts water. A mass verification will probe the reversibility of the hygroscopicity upon heating.

Might oxalic acid be this accelerator?


Sorry folks,
A little reconsideration on the making of Cr2O3 by this method. I dont think it works perfectly. Last time I checked on this reaction a dark green powder was obtained, which is similar to the so told pure Cr2O3, but this time now I obtained a black powder probably because I fried the mixture of Na2Cr2O7+(NH)3SO4 with a aluminium pan, which almost surely released Al3+ ions in product. Darn it!
Gotta use glass inside my oven, but I am a bit afraid of the gases going off and increasing the pressure in the oven making the expected worst to come by.
I will try with a smal ammount. I am just uncertain as maybe chromium III sulfate may be depositing after the reaction

About its color, Wiki says:
" It is not readily attacked by acids or bases, although molten alkali gives chromites (salts with the Cr
2O2−
4 anion, not to be confused with the related mineral chromite). It turns brown when heated, but reverts to its dark green color when cooled. It is also hygroscopic."
Between this is the picture Wiki throw for it:



Also Wiki says:
"The conversion of chromite to chromia proceeds via Na2Cr2O7, which is reduced with sulfur at high temperatures:[5]
Na2Cr2O7 + S → Na2SO4 + Cr2O3"
So we could at least expect better results with just a sulfate than with sulfur itself.
As long as we talking about ion trapping and reactiojn yield.


EdIt: april 03
Checking on the mass of the product it can be reported the double salt decomposition method is worthy although care must be taken to avoid contamination, as the highly oxydising dichromates are likely to inpart stains from metallic surfaces.


[Edited on 4-3-2012 by Poppy]

Here the doublesalt can be seen just prior to thermal decomposition and the resulting blackpowder thus formed.


[Edited on 4-3-2012 by Poppy]

mixture.JPG - 134kB fstpdr.JPG - 143kB

Washing the black powder yields the expected ammount of Cr2O3, but because of contamination the best I could get until now is a very dark green powder.

[Edited on 4-3-2012 by Poppy]
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[*] posted on 5-4-2012 at 20:18


AJKOER:

A test was carried with the devilish chloric acid from NaClO3 owed from a phosphating solution motif. It was abut 20 -30% in think and then sulfuric acid was added to release the beast, nothing out the sketches happened unfortunatly though.

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[*] posted on 5-4-2012 at 22:29


Even decomposition of pure (NH4)2Cr2O7 does not yield the nice green powder, such as shown in the picture above.

The decomposition of ammonium dichromate mostly goes according to the reaction:

(NH4)2Cr2O7 --> Cr2O3 + N2 + 4H2O

There are, however, side reactions in which the chromium is not fully reduced. I have read that also some CrO2 is formed, which is dark brown, nearly black. The presence of this compound makes the resulting Cr2O3 dark green and not like the nice green of pure Cr2O3.

Also, part of the ammonium dichromate does not decompose at all.

In yet another reaction, also some ammonia is released (you can smell it). Release of ammonia means that chromium remains hexavalent. Maybe this hexavalent chromium converts to CrO2 at the high temperatures, with release of oxygen, which can oxidize any further ammonia.




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[*] posted on 6-4-2012 at 05:38


I need trivalent chromium for the purpose of assisted ADP crystal growth, probably purity is an important factor here. Do you think the reduction with aldehyde could bring it anything near purity in the reaction of dichromate, even if it takes days?
Better try it myself, and then bring some results.
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[*] posted on 9-4-2012 at 18:45


Couldn't you reconvert the chromium oxides to chlorides and then precipitate out Cr(OH)3, heating mildly to decompose to Cr2O3?

Alternatively, couldn't you precipitate out both iron and chromium hydroxides, convert to the (III) oxides by heating, and magnetize the iron out?

[Edited on 10-4-2012 by elementcollector1]




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[*] posted on 11-4-2012 at 07:14


That could be a simple, but since the hydroxides co-precipitates at first the then formed oxides would stick together rendering separation impossible. Besides, carrying the separation in water is actually done with gold, but requires considerable density differences so that magnetite Fe3O4 floats while gold sinks. Somehow ferromagnetism of the oxides does not show off very efficiently in water, probably because of hydration.

Also, the conversion must be carried in nitric acid because Cr is bad and wont release its complexes that easy.
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[*] posted on 11-4-2012 at 12:16


Quote: Originally posted by AJKOER  
For curiosity, first make some HOCl by adding vinegar to Bleach (NaClO/NaCl/NaOH). This produces dilute HOCl in a solution of Sodium acetate.


Here is a reference that briefly outlines the chemistry of Chloride and HOCl on Stainless Steel per an article on Stainless Steel at NACE Resource Center, which apparently occurs naturally as a consequence of Chlorine in the drinking water:

http://events.nace.org/library/corrosion/MatSelect/rouging.a...

To quote:
"Class II Rouge forms in a two stage reaction, the first is the dissolution of the chromium oxide passive layer, the second the oxidation of the iron in the substrate:

Cr2O3 + 6Cl- + 6H2O --> 2CrCl3(aq) + 6OH-

2Fe + 4H2O --> 2FeO(OH) + 3H2

This reaction is self-perpetuating by the chloride reacting with the chromium to form hypochlorous acid as a byproduct, and the hypochlorous acid oxidizing the iron and forming more chloride."

Also, the author notes:

2FeO(OH) --> Fe2O3.H2O

I would recommend Vinegar and Bleach (NaOCl/NaCl) as the Sodium acetate may catalyze the reaction of HOCl and NaCl on Stainless Steel. Wait a few hours or a day to judge visually the efficiency of this reaction.

EDIT: the author's first equation has a typo, should be:

Cr2O3 + 6Cl- + 3H2O --> 2CrCl3(aq) + 6OH-


[Edited on 11-4-2012 by AJKOER]
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[*] posted on 22-4-2012 at 11:17


Reporting in to say that my sodium dichromate is a beautiful sunset orange, and should I convert to the ammonium salt before decomposing?



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[*] posted on 22-4-2012 at 23:16


High-silicon cast irons also offer excellent resistance to nitric acid at all temperatures and concentrations, with the exception of dilute hot acid. Aluminum is actually resistant to nitric acid if the acid concentration is over 95%, but if the concentration is below 80%, or if the nitric acid is heated above 40degC, the corrosion rate is much faster. Handbook of corrosion data. Bruce D. Craig, David S. Anderson, ASM International
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