I've been trying to make potassium dichromate following the synthesis given by rhodium. I ude the Cr(III)salt - Cr2O3.
But i cant get the reaction with neither H2SO4 or HCl to go along. Have anyone tried to make potassium dichromate with this method using Cr2O3?
This file was written for informational purposes only. You should not attempt to make it and to use this controlled (in Europe it is)
No it isn't. I bought a kilo from an online pottery store.
Chromates usually oxidize HCl to Cl2 gas...chloro - 23-4-2004 at 11:16
Bought a kilo dichromate from a pottery store? Seriously doubt that... Perhaps you meant Cr2O3?
Good.. but.. ehm.. That didnt help. I already posses Cr2O3. Have you tried to make dichromate?
You think the synthesis is a fake?chemoleo - 23-4-2004 at 11:39
Hmm, I think photography's suppliers are a better approch (for k2Cr2O7). Last I checked you can get it in the UK and Germany.
For Cr2O3, pottery suppliers are indeed the best option.vulture - 23-4-2004 at 12:17
Quote:
Bought a kilo dichromate from a pottery store? Seriously doubt that... Perhaps you meant Cr2O3?
I bought a kilogram of sodiumdichromate from a dutch online pottery store. I could take a picture with my webcam if you don't believe me.Marvin - 23-4-2004 at 13:18
chloro, you dissolved the oxide in sodium hydroxide, then added the acid to that solution, and it didnt turn orange?
Was the solution strongly acid after you finished?
The colour change is not that marked I think. Just so long as its not green.Polverone - 23-4-2004 at 16:33
Ceramic-grade Cr2O3 will not dissolve in HCl or H2SO4. Fuse it with KNO3 to make K dichromate.chloro - 24-4-2004 at 06:55
No, Marvin. I tried reacting the oxide with H2SO4 and HCl, to make a soluble chromium[III]salt that sould be reacted with NaOH.
Polverone: Do you at wich temperature this occurs? Just melting temperature?Marvin - 24-4-2004 at 09:34
That might be a little extreme Polverone, though doubtless a mixture of hydroxide/nitrate would work for producing chromate.
Part of the process is devoted only to producing the hydrated oxide. I suggest using an equimolar amount of the ceramic oxide with the NaOH/Peroxide
solution and see if it dissolves. If it does it would be a lot easier. If it doesnt, or isnt fast enough, you may be stuck with molten salt methods.Polverone - 24-4-2004 at 11:19
I don't know exactly what temperature it occurs at. A while back I tried following an old method for producing chromates, heating a mixture of
K2CO3 and KNO3 with Cr2O3. It worked, but required too much heat for my tastes (the Pyrex test tube I was using melted) and chromate seemed to be
formed in only the very hottest parts of the tube.
More recently, I pleasantly discovered that fusing KNO3 or NaNO3 with Cr2O3 would permit the Cr2O3 to dissolve readily at alcohol-burner temperatures
(with the evolution of much NO2), and that the end product was the dichromate rather than chromate, so no separate acidification step
was needed to obtain dichromate.
Most of the Cr2O3 dissolved in the melt, but some did not, leaving me to wonder if there was some sort of physical difference that led some particles
to resist attack or if something else were at work.
I tried this several times on a test tube scale and once on a pyrex dish scale. The main limiting factor in scaling it up was the abundant production
of NO2, which the addition of hydroxide to the melt might mitigate.
I separated Cr2O3 particles by decantation (since I don't have any dichromate-resistant filters) and was able to obtain nice crystals of
potassium and sodium dichromates from their respective solutions by evaporation at ambient temperatures. K dichromate, being much less soluble, is
easier to separate as crystals.
A mixture of nitrate and hydroxide should fuse and work at a lower temperature, though I imagine you'll obtain chromate that way and not
dichromate directly.fritz - 25-4-2004 at 02:14
I know the reaction with KNO3/Na2CO3/Cr2O3 (Na2CO3/KNO3 1:1 twice the amount as Cr2O3) as a decomposition for insoluble Cr-compounds in analytical
chemistry. In this context I only used it with small batches (ca. 0,5-1g) I carried out this reaction in a porcelain-crucibles (which usually died
during the strong heating).
Polverone I think you could try to filter your slns. through normal papers. I used to do this and they used to survive this treating.
Maybe it is possible with greater batches to melt it in some kind of metal-container (e.g. iron-dishes) for several hours. If the material of the
container is attacked you may remove the contamination by adding NH4OH or KOH or H2S to precipiate unwanted heavy metals. The CrO4(2-) will stay
unaffected.chloro - 29-4-2004 at 08:43
Vulture: Okay, i've never seen it for sale anywhere. But i assume you live in Holland then? Export of a toxic material such as chromate or
dichromate, even within eu, would be impossible, or?
Polverone: Sounds interessting. How large a batch did you use with the xNO3/Cr2O3 - method? And at with wich ratios? Also, you describe adding NaOH to
the mix to absorb - have you tried this? Couldn one lead the NOx with kind of a tube into water, making HNO3 ? Or...
Expensive but it's there.Polverone - 29-4-2004 at 14:59
I ran several test-tube size batches and one batch in a pyrex dish with probably 15 grams or so of Cr2O3. As I said, I was discouraged from scaling up
any further due to the fumes and my lack of open spaces to do this sort of thing. I wasn't particularly careful with the ratios, using nitrates
in excess since I have a lot of them and the dichromate isn't hard to separate.
My idea about adding the NaOH is merely an idea, not tested. But I see no reason why it would not work. I don't recall seeing any NO2 when using
the carbonate/nitrate mixture, and I would thing hydroxide would similarly suppress production of NO2. The NO2 might be used for something or other,
but from such a high temperature reaction I wouldn't want to try taming it for other purposes, at least not without better equipment than I have
currently.
What a joke!
Polverone - 29-4-2004 at 15:03
The Photographer's Formulary says they require a DEA form to order potassium dichromate or any of a number of other chemicals (including dilute
sulfuric acid!)
Is it even possible to get a DEA form that says "this person is authorized to purchase 48% H2SO4"? It'd be like trying
to get a permit to buy distilled water, I would think! And I can't imagine how the company sells those materials at all if they require this
mysterious form that no photographer has any reason to possess in the first place. It's a mystery why they are even listed on the site.frogfot - 23-5-2004 at 03:24
I've successfully tested the KNO3/Cr2O3/KOH procidure, where dichromate were precipitated by acetic acid.
But on testing K2Cr2O7 in simple isopropanol oxidation, I realised that stuff is not that soluble, which made me use twice as much water
So, now I wanna produce Na2Cr2O7 by NaNO3/Cr2O7/NaOH, but the problem is how to separate and purify the product.. The only thing I've came up
with is to use the good solubility of Na2Cr2O7..
Basically, if we ubtain (from experiment above) a mix of Na2CrO4, NaNO2 and some traces of NaOH and NaNO3.. On addition of H2SO4 (with a tiny excess)
while heating, one would get Na2Cr2O7, Na2SO4 and lots of NOx.
The NOx could be easily collected, but how would one destroy them? Would bubbling through NaOH soln. work?
Then on partial evaporation and cooling, Na2SO4 would precipitate. Leaving pretty pure Na2Cr2O7 soln. This could be evaporated again and
redissolving/cooling is repeated... does this sound practical? or have I forgotten something obvious?..
Only thing is, finding source of NaNO3 is hard. I don't think they use it today as a fertiliser.
Oh another unanswered question that have bugged me a long time, I've noticed that when I add acid to precipitate dichromate to abovementioned
reaction mix, there evolves a colorless gas that have no smell.. what could this be??
Only thing that could evolve is NOx from decomposition of nitrite.. but I sure would've noticed this!!
This happens also when some acetic acid is poured on homemade KNO2 from Pb/KNO3 method.. no color/smell..
Btw, this gas evolves until solution becomes acidic..
[Edited on 23-5-2004 by frogfot]t_Pyro - 24-5-2004 at 08:15
Quote:
Originally posted by frogfot
I've noticed that when I add acid to precipitate dichromate to abovementioned reaction mix, there evolves a colorless gas that have no smell..
what could this be??
[Edited on 23-5-2004 by frogfot]
It could be oxygen gas produced by the reduction of Cr<sup>6+</sup> to Cr<sup>3+</sup>.
The NOx should get absorbed by NaOH I think, to give the corresponding NOx salts which have N in the same oxidation state.S.C. Wack - 29-4-2005 at 11:35
I've had some success getting the calcined pottery Cr2O3 to react with H2SO4 and HCl in test-tube experiments. The secret weapon was MnO2, also
from the pottery supply.
This might not be worthwhile as a preparative method, but I thought that I would mention it anyways.
Cr2O3, MnO2, and 90% H2SO4 were heated to around 70C for 8 hours. There was a green liquid and a black solid. The whole was neutralized with excess
NaOH, H2O2 was added, and a lovely yellow chromate layer was filtered from a typical brown precipitate of MnO2.
Cr2O3, 30% HCl, and MnO2 were allowed to sit for 8 hrs. Then cold, conc. NaOH was added. 4 hours later the solids were separated by decantation and
washed, then NaOH soln and H2O2 was added, again giving chromate and MnO2.
This was only a qualitative test, but there seemed to be very little unreacted Cr2O3. The amounts of the other materials used, while in excess, were
nowhere near excessive.
P.S. - I am sure that I got a good yield of Na chromate (in solution at least). It was isolated as a light yellow heavy Ba salt by adding CaCl2
sol'n, filtering, and BaCl2 sol'n was then added to the filtrate.
P.P.S. - When I made the BaCl2, I opened a new batch of BaCO3 from a major US pottery supplier, rather than using "recycled" BaCO3 on hand.
I had never used this suppliers chems before but assumed that it was what they said it was. Well, the fact that it contains BaS was not mentioned
anywhere and I did not discover this until I added the HCl to it. Another good reason to always be ready for unlikely dangers when in the lab.
[Edited on 29-4-2005 by S.C. Wack]Polverone - 17-4-2006 at 00:07
I recently had a more pleasant experience with chromate production from pottery-grade Cr2O3. I used a mixture of Na2CO3 and NaNO3, prepared by fusing
Na2CO3 and ammonium nitrate. Chromium oxide dissolved in it with much bubbling and no production of nitrogen oxides, which was a big
plus over using nitrate alone. My earliest experiments used nitrate:carbonate ratios taken from some old industrial text (maybe Muspratt), which was
optimized for economy rather than small-scale convenience. Using more nitrate, it is possible to achieve an easily-fused mixture that is more
convenient to work with, but some carbonate is still useful to prevent the evolution of nitrogen dioxide.
I'll have to try S.C. Wack's methods one of these days, but if you want to go for the high-temperature methods, this is very convenient. I also found
it more convenient this time around to use a small stainless steel dish in place of a test tube, since it holds up better to the heat and will never
annoy you with frothing as test tubes are so prone to do.S.C. Wack - 17-4-2006 at 09:17
I found something quite easy around the time that I wrote the last post, but I had a series of experiments planned. But this led to another...and...so
on. So despite quite a few exacting experiments using various amounts of this and that I never found the best method with KMnO4, Cr2O3, and H2SO4, but
it is quite fast, hot, vigorous, and led to the isolation of large amounts of potassium dichromate. It is preferable to MnO2. The yields can be high
and isolation easy. I did these tests on a large scale rather than use my very many dust gathering test tubes. I suppose I should order some more
Cr2O3 get serious with it.
Now I know that it might not make sense to make a weaker oxidizer from a stronger oxidizer, but oh well. But if anyone else wants to pick this up
before I try it again...
The reaction can proceed instantly to a hot purple cloud with H2SO4 >50%. There is also not enough liquid and/or time to entirely wet the solids.
The yield is lower and the reaction much slower with less concentrated acid and external heating. It was found best to make the reaction happen fast
to take advantage of the Mn while it is in some +7 form, yet with enough liquid to mix everything together, yet maintain a very exothermic reaction.
The Cr2O3 and KMnO4 are mixed quite well and the H2SO4 is dumped in with stirring, and a reaction soon ensues. When the reaction stops, it is allowed
to cool down and it forms a cement. This all takes just 2-3 minutes. The brown/black mass was then heated with water until the water had evaporated.
The mass is then leached with boiling hot water 4 times. KOH is added to make the pH 3.6-4. There is some contamination but all of this can be worked
out due to the low solubility of the dichromate in the cold.guy - 17-4-2006 at 11:46
[Edited on 4/17/2006 by guy]neutrino - 17-4-2006 at 14:27
Of course, there is a difference between the chromic oxide they use and the stuff sold in pottery stores. Theirs was freshly precipitated
Cr(OH)<sub>3</sub> with a very high surface area which made it relatively reactive. The material pottery stores sell is dried and solid
Cr<sub>2</sub>O<sub>3</sub> with a very low surface area. Pottery grade oxide is also probably calcined to an inert crystal
structure.
Chromate Synthesis
12AX7 - 21-4-2007 at 22:39
Performed my own. Second time with this method, this time more controlled and, most importantly, photographed.
Did anyone know potassium chromate is thermochromic?
Timgarage chemist - 22-4-2007 at 04:26
Why did you add KCl? Thats a very bad idea, as you will be unable to obtain dichromate from this. Upon acidification, the dichromate will oxidise the
chloride to chlorine.
Thats the reason that Frogfots site uses acetic acid for acidification, which is the best acid to use as acetates are very soluble and there is no
risk of crystallizing potassium acetate along with the dichromate.
Potassium chromate is best obtained by combining K-dichromate with KOH and evaporating the solution, as the chromate is too soluble to make
crystallization directly from the leached melt economic.
[Edited on 22-4-2007 by garage chemist]12AX7 - 22-4-2007 at 07:28
If you read my article in detail you would notice I mentioned I don't have KOH.
I haven't had a problem so far as far as chloride. Actually, the solution I didn't acidify smelled like bleach for the longest time (its color seems
to be weak and turbid now), whereas the dichromate solution is currently forming crystals. Both produced (yellow-stained) salt first.
My solubility table says potassium chromate is 60-75g/100ml soluble (from 0-100C respectively), no especially big problem.
Timgarage chemist - 22-4-2007 at 09:13
You could have heated your KCl with H2SO4 to make K2SO4, you could have used K2CO3 for this, etc... The potassium ion can even be added after the
oxidation melt.
You need to avoid chloride, otherwise your product can be contaminated by Cr2O3 or toher Cr(III) compounds.
[Edited on 22-4-2007 by garage chemist]12AX7 - 22-4-2007 at 13:35
- Fixed a typo on the page (I had said that "potassium chromate is less soluble").
I wonder what happens if I fuse KClO3 and Cr2O3. Is a strong oxidizing environment strong enough to kick out Cl2? (HCl would be preferred, but I
suppose with no hydration available, that can't happen.)
Tim12AX7 - 26-4-2007 at 15:24
So I did put Cr2O3 and KClO3 together and heated the mixture. Lo and behold, it drove me out of the room! Should've done this outdoors...
The main reaction should be something like:
2 KClO3 + Cr2O3 > K2Cr2O7 + Cl2 + O2 (I don't know that O2 is present, but it balances)
But there will be others, some releasing HCl from hydration and some leaving KCl, which appears to be stable in a CrO4 melt (I don't know about
Cr2O7).
Goes off cleanly, Cl2 notwithstanding. No boiling over, because it fuses to a permeable mush as the chlorate decomposes. That chlorine is pretty
harsh though.
Tim12AX7 - 13-9-2007 at 22:49
I made more K2Cr2O7 today, in two experiments: scaled up, and done outdoors with
fire for heat.
First I tried melting a pound of KClO3 in a rusty steel crucible, which gave me a bubbling salt melt. A lot of oxygen seems to be given off, for a
process that's supposed to make perchlorate in reasonable yield. I added a few scoops of Cr2O3, which resulted in the exothermic decomposition
(reaching up to about red heat, 700C or so) and probably quantitative reaction of what little chrome I added. The melt turned yellow to orange on the
surface, while the rest frothed up with prodigious vigor, giving off choking fumes and, presumably, a lot of oxygen. When it finished, there was a
dull red glow (around 700C in daylight), but all material was solid, so it was all decomposed to KCl.
So, this reaction must be done as a combination. That said, I then blended the remaining 250g Cr2O3 with 450g KClO3 and began heating it slowly.
Careful heat control is required to keep the reaction from running away, especially on this scale (I went cautiously; I didn't intend to find out what
happens when a whole pound of this stuff cooks off all its chlorine and oxygen at once!). At one point, the entire mass became a beet red,
crystalline semisolid sort of consistency, which I suspect was total conversion. I proceeded to melt the material and pour it out. The resulting
slab remained a drab color on cooling. I don't know if I simply couldn't see the unreacted Cr2O3 at that beet-red stage, or if more had formed (due
to oxidation of the crucible? decomposition without bubbling? I don't know).
I am presently hot filtering and recrystallizing the material, which appears to have left a pretty good bit of Cr2O3 on the filter. It's going
slowly, so I'll be at it some more, reheating the liquor, dissolving crude K2Cr2O7 and filtering it. So far, I already have more consistently
orange-red product sitting in the filtrate than I have collected to date (e.g. see above), so in an absolute sense, yield was good.
Timgarage chemist - 14-9-2007 at 04:37
How do you separate potassium dichromate from potassium perchlorate? KClO4 is far less soluble than dichromate.
Can you test your product for presence of perchlorate?12AX7 - 14-9-2007 at 08:30
Chlorate and perchlorate both decompose at the final temperature I brought it to, but I could seperately try fusing KClO3 to produce KClO4.
Timgarage chemist - 14-9-2007 at 11:40
Yes, and preferably not in a metal container, since metals catalyse evolution of oxygen. A porcelain or quartz crucible would be good, a small beaker
would do the job as well. You also need patience.
I havent tried the decomposition of KClO3 since I found NaClO3 to be a better starting material due to its lower melting point, at which it doesnt
already start to decompose.
Whats wrong with Cr2O3 + NaOH + KNO3? I found that to be a very good method of (di)chromate synthesis, completely converting all Cr2O3 if heated and
stirred enough, with only water vapor as byproduct.12AX7 - 14-9-2007 at 14:35
...And N2?
I have neither KNO3, nor NaOH at the moment. KClO3 gives off chlorine, but it's just two things and some heat. Besides which, KClO3 can be made
indefinitely in one's own lab.
