Sciencemadness Discussion Board

Wet chromium trioxide & troublesome sodium chromate

The Mad Plater - 12-3-2017 at 12:44

Let's try a different approach for a change. First, the "short" version - 3 most pressing questions:

1. WTF is that superfine green precipitate which occurs as a side effect of (fully) neutralizing CrO3 with NaOH to sodium chromate? (FWIW, it begins to appear well before the stoich amount of NaOH is added).

2. Any ideas for removing it from the solution? Filtration seems ineffective - filters get clogged fast, yet most of the green gunk stays in solution.

3. Any good ideas for a quantitative analysis of the (sodium chromate) solution concentration, other than a SG measurement?


With that out of the way, it's time for some backstory:

A few weeks ago, my order of CrO3 had finally arrived, all 250g of it. In a flimsy little "flip-top" plastic container.
Also, shaking the container made no noticeable sound - which should have set off alarm bells at that point already, but it didn't, since I was far too preoccupied with other, far more pressing problems at that time.

Because I had no time to do anything useful with it, I just set it aside on the shelf.
About a week or two later, I noticed something most peculiar: a strange black blob a few mm in size had grown on top of the white container lid.
Wiping the blob off revealed underneath it a crack in the plastic lid, which certainly wasn't there before.

At this point, I was quite alarmed by this revelation, and decided to open the (factory sealed) container immediately.
Inside was a sticky, wet black mess - which was a surefire indicator that something had gone horribly wrong.
Whatever that was, it was certainly not chromic acid anhydride. Not anymore, at least.

Well, since what I really needed was not the anhydride, but the aqueous chromic acid (for sealing the phosphate coatings), I set out to make a stock solution of chromic acid (500g/L).
Dumped most of the black goop into a flask, washed out the remainder into the flask with DI water, added water to ~80% of final volume.

And then, more trouble. Roughly 0.5% of that black gorp turned out to be insoluble.
Given that the CrO3 was of analytical purity when it left the factory, the insoluble crap was likely the product of the reaction between the CrO3 and the container walls.
The inside of the CrO3 container was full of small cracks, in addition to the one larger (penetrating) crack in the lid.

I brought the issue up with the supplier (a local brick-and-mortar shop), but didn't press it too hard, given the circumstances - and nothing further became of that.

A few days later, my sweep of the literature turned up an even better process of sealing phosphate coatings than the long-proven chromic acid rinse: a sodium chromate rinse.
Better results, less hassle. If it's good enough for the military, it's good enough for me.

And that's where the REAL trouble started.

To make a long story short, my first attempt at making a stock solution of 350g/L sodium chromate, from stock solutions of chromic acid and NaOH, could best be described as "a total disaster".

Because of a stupid mistake (a simple mathematical error), I had only measured out half of the correct amount of NaOH, and when the reaction (predictably) failed to produce the yellow sodium chromate, I failed to realize what was going on, and kept adding more (unmeasured) NaOH in order to get the pH into the right range. (VERY STUPID MISTAKE! Don't do that!)

Result: flask full of yellow liquid, itself full of gelatinous green precipitate (Cr(III) hydroxide?).
After filtering (most of) that crap out, which took HOURS and at least 8 coffee filter papers, I was left with only 0.74L of solution - far from the intended 1L - and a quick SG measurement put the concentration at 260g/L, instead of the intended 350g/L.

However, I do have to say, it was quite clear and transparent after filtering it once - only a tiny amount of the green precipitate remained, despite using cheap'n'crappy coffee filter papers.
NB: tried qualitative filter papers - clogged up 10x faster than the coffee filters. Gave up after the first one.

Fixing this mess would have required considerably more chromic acid than I had remaining, so I ordered another 250g container, advising the shopkeeper to watch out for the quality issues.

About a week ago, I picked it up from the shop. At that time, the contents did rattle somewhat when shaken, although it wasn't quite a "dry" rattling sound.

Today I opened that container... and guess what? Same. Exact. Problem. Similar cracks in the container, same wet black mess, and even more of insoluble contaminants.

OK, made more of the chromic acid stock solution.
Calculated and measured out the amount needed, dumped it into the previous sodium chromate "screwup".
Then I added the correct amount of (solid) NaOH, slowly, with lots of stirring.

Strangely, the results seem to defy any and all reason:
1. Although only a fairly minor amount of the green precipitate resulted, it started to appear well before the stoich amount of NaOH was added. It's possible it started to appear right away, but the sodium chromate + chromic acid mix was far too opaque to determine that.
2. Filter papers appear to be completely ineffective, which I find most vexing. While they clog really fast, it doesn't seem that the clarity of the solution is improving, even after 2 consecutive filtrations.
3. WTF is happening with the specific gravity? Although I didn't have the time to measure it accurately today, a quick'n'dirty measurement seems to indicate a similar value as before - but clearly the concentration of sodium chromate has (or at least should have?) increased???

What a mess.

I'm half tempted to write this off as a total loss, and just buy some sodium chromate and make a stock solution with that.
I already wasted far too much time on this nonsense.

chornedsnorkack - 12-3-2017 at 13:28

Stoichiometric amount as what?
The pKa of Cr(III) is quoted as around 4.

JJay - 12-3-2017 at 13:37

If you reduced some of the chromium trioxide with plastic, you might have some trivalent chromium oxide in the mix. It's green and insoluble in water.

[Edited on 13-3-2017 by JJay]

The Mad Plater - 12-3-2017 at 14:16

Quote: Originally posted by chornedsnorkack  
Stoichiometric amount as what?
The pKa of Cr(III) is quoted as around 4.


Uh, I'm quite confused now. Most likely because of the very late hour.

If you're referring to the amount of NaOH, then by stoich I meant 2 moles NaOH for every mole of chromic acid - as is proper for sodium chromate.
1:1 ratio would yield the dichromate instead.

I'm probably missing something really obvious here, but how is the pKa of Cr(III) supposed to be of use in this case?

Regarding the insoluble contaminants in the CrO3 - as far as I could tell, there were 2 major types, based on how they appeared as observed on the walls and bottom of the flask:
- small (0.2-0.5mm) granules, possibly white or colorless, and transparent/translucent (possibly bits of destroyed plastic?),
- similarly small dark-colored (black? dark green?) flakes.

I didn't investigate said debris more closely than that - I was running low on time, and I also have no desire to be closely examining unknown substances still wet with a solution of hexavalent chromium ions; even though I very much value Cr(VI) compounds for their corrosion protection properties.

As far as I'm concerned, Cr(VI) compounds are, for practical purposes, analogous to the HF-based stainless pickling paste: the benefits are fantastic, but both of these substances are highly deadly, if only in somewhat different ways.

chornedsnorkack - 13-3-2017 at 03:34

Quote: Originally posted by The Mad Plater  
Quote: Originally posted by chornedsnorkack  
Stoichiometric amount as what?
The pKa of Cr(III) is quoted as around 4.


Uh, I'm quite confused now. Most likely because of the very late hour.

If you're referring to the amount of NaOH, then by stoich I meant 2 moles NaOH for every mole of chromic acid - as is proper for sodium chromate.
1:1 ratio would yield the dichromate instead.

I'm probably missing something really obvious here, but how is the pKa of Cr(III) supposed to be of use in this case?

The pKa of dichromate is about 6.
Meaning Cr(III) should be insoluble in chromate solutions - or indeed in chromate/dichromate buffer.

Bobymartinez - 14-3-2017 at 07:34

The green precipitate is certainly Cr (III) oxide as told previously. So you might have some reducing agent in your solution. To get rid of it, try filtration over Celite (it might help sometimes with fine precipitates). Or you can also try to filtrate your chromic acid before the addition of the sodium hydroxide.

The Mad Plater - 14-3-2017 at 11:45

Quote: Originally posted by chornedsnorkack  
The pKa of dichromate is about 6.
Meaning Cr(III) should be insoluble in chromate solutions - or indeed in chromate/dichromate buffer.


Oh, OK.

Actually, is Cr(III) oxide even soluble at all in any "commonly encountered" substances?

Because that would seem like a fairly easy route to Cr(VI) compounds... dissolve the cheap'n'plentiful Cr(III) oxide in {something}, then add some of the cheap'n'plentiful 30% hydrogen peroxide, and Bob's your uncle.


Regarding the filtration of the "dirty" chromic acid: on paper, it's a great idea, and I did in fact try it.

Unfortunately, because all I have are regular filter papers - the ones made out of actual paper, as opposed to, say, asbestos (the kind of filter papers normally called out for in the literature for this sort of use) - they really don't get along with the concentrated chromic acid.
As in, they rapidly become black on contact with even a single drop.
I'm no expert, but I'm pretty sure that whatever this reaction is, I don't want it to be occuring on any larger scale than that.

Who'd have known that mixing strong oxidizers with reducible substances (commonly known as "fuels") isn't a terribly great idea?


BTW: this is just a hunch, but I'm suspecting that the reducing agent contamination most likely came from the crappy plastic container the CrO3 arrived in.
Seeing as the plastic was cracking, but only inside the container, the first thing that comes to mind is plasticizer leaching.
And I'd imagine that most of the common plasticizers would also have reducing properties (they do burn just fine...).

Another thing in support of that theory is the fact that the previous batch of CrO3 resulted in a relatively huge amount of Cr(III) on neutralization (but it was sitting in that crappy plastic container for a good few weeks on my shelf), whereas the newer batch yielded only minor amounts of Cr(III) - but it was also removed from the container with far less delay.

Bobymartinez - 14-3-2017 at 12:02

To transform your cheap Cr (III), you actually need to fuse together your oxide (or even directly the chromite ore) with sodium carbonate under oxidative conditions. Few years ago I transformed some chromite I had using this procedure and used sodium nitrate as oxidizer (I put everything in a stainless steel container mixed evenly and heated on direct flame the container) After reaction, you end up with sodium carbonate, chromate (yellow), iron oxide and ureacted chromite. Separation is quite easy since only sodium carbonate and chromite will be soluble in water. So you filter the reaction and then transform you chromate in dichromate by acidification. Then finally isolate your sodium dichromate.

So for your problem, if you only have paper filter, you indeed are quite in trouble to get rid of your Cr (III) oxide...

chornedsnorkack - 14-3-2017 at 13:35

Quote: Originally posted by The Mad Plater  
Quote: Originally posted by chornedsnorkack  
The pKa of dichromate is about 6.
Meaning Cr(III) should be insoluble in chromate solutions - or indeed in chromate/dichromate buffer.


Oh, OK.

Actually, is Cr(III) oxide even soluble at all in any "commonly encountered" substances?

Because that would seem like a fairly easy route to Cr(VI) compounds... dissolve the cheap'n'plentiful Cr(III) oxide in {something}, then add some of the cheap'n'plentiful 30% hydrogen peroxide, and Bob's your uncle.

Chromium(III) is supposed to be amphoteric. Meaning that, unlike iron(III), but like aluminum(III), it may dissolve in excess of aqueous alkali.

The Mad Plater - 14-3-2017 at 13:44

Sounds like way too much effort for something I can just go and buy relatively inexpensively...

An even crazier idea I've run across was generating Cr(VI) compounds by some process involving electrolysis with scrap stainless steel electrodes.
Which seems rather inefficient w.r.t. time and resources, since common stainless steels are less than 20% Cr.

Back to my Cr(III) problem - what about recrystallization?
If I understand it correctly, recrystallizing the sodium chromate should cause the Cr(III) precipitate to become concentrated in the remaining liquid phase, since there's no proper place for it in the crystal structure.

Then wash with alcohol (is that even safe? ignition/explosion hazard?), dry and weigh the crystals - this would also answer the tricky question of "what was the actual yield?".
(EDIT: Oops. Woudn't it just reduce the alcohol???)

Haven't actually purified anything by recrystallization yet (not on purpose, anyway), so that might be just wishful thinking on my part, though.
(also, don't have a crystallizing dish - but I do have an evaporating dish; close enough)

[Edited on 14-3-2017 by The Mad Plater]

Bobymartinez - 14-3-2017 at 14:04

Indeed it's easier to just buy it but the preparation is quite interesting to be done once.
I thought you had your Cr (III) oxide precipitation during the neutralization of your trioxide with NaOH. If you try to recrystalize this solution, you'll inevitably end up with Cr(III) precipitate into your product so it won't solve you problem. Then wasing up chromate or dichromate with alcohol is really a bad idea, you'll reduce some of your Cr(VI) in Cr (III) once again.
I stick on the fact that the best idea is filtering over something very fine. Just try to find some fine sand or crush it into a very fine powder to put it over you filter paper if you can't have access to celite (which surprises me, should be quite easy to find...)

The Mad Plater - 14-3-2017 at 14:59

IIRC, Celite = diatomaceous earth?
If so, then yes, it's not hard to find, nor very expensive (not sure of the purity though).
Although I don't know where I can buy it locally - though I'm sure it should be available somewhere - but it's easily available online.

Don't have any of it on hand, though.
I do have a whole bunch of hollow glass microspheres, but I don't think that would work, since they float in water - not surprising, with their stupidly low density.

For now, I think I'll just try filtering the offending solution a few more times with the qualitative filter papers - after the second pass, it is significantly clearer than it was to begin with, and the filter papers don't clog up nearly as fast, either.

Now I just had an idea (either really great, or really stupid) - if the (hypothetical) sodium chromate crystals were bone dry, would washing with nonpolar solvents do any good? Or would that just result in a big fireball?

I do have a lot of brake cleaner - it lists the ingredients as "n-hexane, cyclohexane, hydrocarbons, yadda yadda yadda". Quite volatile, evaporates very fast with essentially no residue.
Sure does smell very nice, which is a pleasant change from the normally rather malodorous technical-grade light hydrocarbon fractions.

Then there's that old can of an unknown ether (diethyl?), but it's contaminated with lubricating oil and other unknown additives, and I'm not sure if I like the idea of messing around with trying to purify that.

chornedsnorkack - 14-3-2017 at 23:32

Quote: Originally posted by The Mad Plater  

Back to my Cr(III) problem - what about recrystallization?
If I understand it correctly, recrystallizing the sodium chromate should cause the Cr(III) precipitate to become concentrated in the remaining liquid phase, since there's no proper place for it in the crystal structure.

Incorrect.
Precipitates can easily be trapped in growing polycrystalline precipitates!
Quote: Originally posted by The Mad Plater  

Then wash with alcohol (is that even safe? ignition/explosion hazard?), dry and weigh the crystals - this would also answer the tricky question of "what was the actual yield?".
(EDIT: Oops. Woudn't it just reduce the alcohol???)

Ethanol could fairly easily be oxidized.
Then again, chromates are weaker oxidants in basic conditions.
Quote: Originally posted by The Mad Plater  

Haven't actually purified anything by recrystallization yet (not on purpose, anyway), so that might be just wishful thinking on my part, though.
(also, don't have a crystallizing dish - but I do have an evaporating dish; close enough)


Your problem is chromium(III) impurity. Additional problem is that it is a fine colloidal precipitate which is not easy to settle/filter out.

You might dissolve it, in either acid conditions or in excess alkali. But in solution, it would still be chromium(III).

You might try to convert a fine colloidal precipitate into coarser grains of chromium(III) hydroxide/oxide, easier to remove. A common means for that is digestion - prolonged heating of precipitate with mother liquor.

Alternatively, you might oxidize chromium(III) back to chromium(VI). That would be easier in basic conditions. But what could be a suitable oxidant to oxidize a chromium(III) impurity in excess of chromium(VI)?

The Mad Plater - 15-3-2017 at 07:50

Quote: Originally posted by chornedsnorkack  
Precipitates can easily be trapped in growing polycrystalline precipitates!

Good to know.
So if I understand this right, only the dissolved impurities get excluded from the growing crystal lattice?
Quote: Originally posted by chornedsnorkack  

Ethanol could fairly easily be oxidized.
Then again, chromates are weaker oxidants in basic conditions.

I was thinking more along the lines of isopropanol, actually, since I do have quite a bit of it, and in reasonably high purity, too.
Quote: Originally posted by chornedsnorkack  
Alternatively, you might oxidize chromium(III) back to chromium(VI). That would be easier in basic conditions. But what could be a suitable oxidant to oxidize a chromium(III) impurity in excess of chromium(VI)?

Ah, and we arrive at the crux of the problem.

According to this page, hydrogen peroxide would do just fine.
Which would be very nice, since I have it on hand, and I normally use rather small amounts, so the remainder of the 1L bottle would eventually just decompose on the shelf anyway.

Now, what I can't quite grasp is: if I need to use excess NaOH for this to work, then how would I go about getting rid of the excess afterwards?
Or would it just get used up by the newly (re-)created Cr(VI) ions, if the "excess" is not actually too excessive?

Well, I guess I'll try a test-tube-scale experiment to check that later. At least the peroxide can't ignite, nor oxidize any further...

chornedsnorkack - 15-3-2017 at 14:22

Quote: Originally posted by The Mad Plater  
Quote: Originally posted by chornedsnorkack  
Precipitates can easily be trapped in growing polycrystalline precipitates!

Good to know.
So if I understand this right, only the dissolved impurities get excluded from the growing crystal lattice?

Yes.
Quote: Originally posted by The Mad Plater  

Quote: Originally posted by chornedsnorkack  

Ethanol could fairly easily be oxidized.
Then again, chromates are weaker oxidants in basic conditions.

I was thinking more along the lines of isopropanol, actually, since I do have quite a bit of it, and in reasonably high purity, too.

A secondary alcohol. Easily oxidized to acetone.
Acetone, however, is fairly resistant to oxidation, and polar:
https://en.wikipedia.org/wiki/Jones_oxidation

The Mad Plater - 19-3-2017 at 10:15

Chromic anhydride and acetone. I don't quite think I'm cut out for that kind of excitement - as far as I'm concerned, there are other, far more fun things to do than mixing highly toxic strong oxidizers with highly flammable fuels.

But back to the subject at hand - my precipitate problem has now cleared up significantly.
It was so simple, too! All that was needed was ~15mL of 30% H2O2, to get rid of the precipitate from nearly 1L of sodium chromate solution.
(of course I first did a small scale test on ~1mL!)

It was quite strange, too - at first, the solution turned a dark amber color, more typical of dichromates - but after several minutes of gentle heating to decompose the excess H2O2, it gradually returned to the proper, intensely yellow color.

I currently have some difficulties measuring pH in this range ("universal" test papers are IMO the most worthless!), but it does appear to be in the correct range of roughly 9-10, as far as I can tell.

A very small amount of some unidentified precipitate still remains, but it's very different from before - not at all green, and more "flaky"/fuzzy rather than gelatinous.
Given its current form, and how little of it remains, it should now filter out easily.

The density was still short of the target, which still perplexes me (how did that happen?), so I evaporated ~200mL of water from that solution, to verify if the solution density is still a (nearly) linear function of concentration in this range - and it is, to within the margin of measurement errors.

Well, now that I have this mess figured out, next weekend it's time to wrap it up, and (hopefully) never have to do this again.

chornedsnorkack - 20-3-2017 at 03:31

Quote: Originally posted by The Mad Plater  
Chromic anhydride and acetone. I don't quite think I'm cut out for that kind of excitement - as far as I'm concerned, there are other, far more fun things to do than mixing highly toxic strong oxidizers with highly flammable fuels.

Yes, but ethanol also is highly flammable.
It's interesting that hydrogen peroxide acted so nicely. There were at least 2 possible side reactions I worried about :
formation of CrO5;
catalysis of hydrogen peroxide decay.

The Mad Plater - 20-3-2017 at 06:56

Quote: Originally posted by chornedsnorkack  
Yes, but ethanol also is highly flammable.

Sure, but I also don't "use" it by mixing it with hazardous oxidizers, either.

Energy wise, it also makes a rather poor fuel, compared to say, diesel - the energy density is woefully low.
Quote: Originally posted by chornedsnorkack  
formation of CrO5;

I guess that's what could have caused that transient color change?
And then it decomposed on its own, and everything went back to normal.

Although strangely, it didn't cause more Cr(III) precipitation. Maybe it decomposed faster than the peroxide did?
Quote: Originally posted by chornedsnorkack  
catalysis of hydrogen peroxide decay.

And how would that be a bad thing exactly, seeing that the excess peroxide needs to be decomposed anyway?

Well, other than the possibility of a runaway decomposition, but with a concentration that small, it would just fizz a bunch and then do nothing.

chornedsnorkack - 20-3-2017 at 10:56

Quote: Originally posted by The Mad Plater  
Quote: Originally posted by chornedsnorkack  
Yes, but ethanol also is highly flammable.

Sure, but I also don't "use" it by mixing it with hazardous oxidizers, either.

You did propose washing chromate with "alcohol".
So: primary and secondary alcohols (like isopropyl alcohol you proposed) would be prone to oxidation. Tertiary alcohols would resist oxidation at low temperature (but are no less flammable if ignited).
Of course, chromates are weaker oxidants than dichromates which are weaker oxidants than chromium trioxide. Basic conditions make chromium(VI) less oxidizing.
What you can have is either bursting on fire, or quiet oxidation, or no reaction. Since your goal is pure chromate, you don´t want any reaction at all, even a quiet reaction. And chromate might be basic enough to not be reduced?
Quote: Originally posted by The Mad Plater  

Quote: Originally posted by chornedsnorkack  
formation of CrO5;

I guess that's what could have caused that transient color change?
And then it decomposed on its own, and everything went back to normal.

Although strangely, it didn't cause more Cr(III) precipitation. Maybe it decomposed faster than the peroxide did?
Quote: Originally posted by chornedsnorkack  
catalysis of hydrogen peroxide decay.

And how would that be a bad thing exactly, seeing that the excess peroxide needs to be decomposed anyway?

Well, other than the possibility of a runaway decomposition, but with a concentration that small, it would just fizz a bunch and then do nothing.


What you want:
1) 2 Cr(OH)3+3H2O2+4NaOH->2Na2CrO4+8H2O
What you don´t want:
2) 2H2O2->O2+2H2O (and Cr(OH)3 unchanged)
3) 2Na2CrO4+3H2O2+2H2O->2Cr(OH)3+4NaOH

Possibly the basic conditions favoured reaction 1) over 2) and 3)

The Mad Plater - 20-3-2017 at 12:59

Quote: Originally posted by chornedsnorkack  
You did propose washing chromate with "alcohol".

Sure, but not without asking around first.

And it's a good thing that I didn't actually try that, because at this point I'm quite certain that whatever the resulting reaction would be, I would most likely not be at all happy with the result.

Also, that's one of those suspicious things which IMO absolutely merit a sub-gram scale test first, if even trying it at all, especially seeing how that was just throwing crap at a wall to see if it sticks.

Then of course there are the things that should NEVER be attempted at all, such as combining anything containing acetone with H2O2 of any concentration.
That's certainly one way to push the limits of otherwise well-explored "household chemistry" into a new, poorly explored dimension - but of more immediate importance, it will also invariably push shards of your labware into poorly explored areas of the far wall.

*****

Now, for a change of pace, let's try tackling a related problem from a different angle:
Would it be safe and/or effective to dispose of highly dilute Cr(VI)-containing waste (to the tune of a few grams per liter) by mixing it with contaminated isopropyl or ethyl alcohol, and then acidifying the mixture?

The idea here is that these 3 substances are wastes from the processing of parts, so it incurs no further costs to use them for reducing the Cr(VI), which is otherwise the most troublesome from a waste disposal standpoint.

chornedsnorkack - 21-3-2017 at 02:00

It would certainly be safe. Fairly dilute aqueous alcohol cannot be ignited by strong oxidants (because water absorbs the heat). Oxidation of ethanol gives ethanal and then acetic acid (both miscible with water); oxidation of isopropanol, acetone (also soluble).

What I'm not sure of is how effective it is, in rapidly reducing traces of chromium(VI) in dilute aqueous solutions.

Also: acetone peroxide IS a nuisance in dilute aqueous solutions, because it is poorly soluble in water - it can precipitate and then explode.
Basic conditions are said to diminish formation of acetone peroxide.

If you have hydrogen peroxide present then you might oxidize isopropanol to acetone... and then to acetone peroxide.

The Mad Plater - 26-3-2017 at 11:03

Quote: Originally posted by chornedsnorkack  
If you have hydrogen peroxide present then you might oxidize isopropanol to acetone... and then to acetone peroxide.

Which is exactly what I was worried about, since IPA (and/or acetone) is my preferred choice for cleaning parts.

Plan B, I guess, is to use denatured alcohol instead. It's fairly cheap, as far as solvents go, and I have it in large quantities for deicing purposes - so it's OK to sacrifice a (relatively) small amount.

Then there's also plan C: since most of my acid pickles contain citric acid in varying proportions, then - if I understand it correctly - using waste citric acid should do the job by itself, since it both acidifies the mixture, and is also itself oxidizable.

I tried it on a sub-gram scale earlier today - ~5mL of 0.35% sodium chromate solution added to ~0.5g citric acid and ~0.5mL 30% H2O2 (a very useful combination!).

This rapidly (instantly?) resulted in the Cr(VI) being reduced to Cr(III), as evidenced by the characteristic color of the [Cr(H2O)6]3+ ion.
Further addition of ~1mL of an unknown alcohol (certainly not IPA; most likely wood alcohol?) resulted in no obvious reaction.

Heating the resulting solution resulted in no further transformation, and no new interesting properties - other than a curious tendency to bump and erupt with extreme prejudice when heated close to boiling point, despite very slow and careful heating.
I haven't tried this without adding the H2O2 (yet), however I suspect that the result would be essentially the same.


BTW, I went back and decided to try that Cr(III) to Cr(VI) oxidation-by-H2O2 trick again, to figure out what's going on.

As expected, if there's any dichromate present at all, this fails miserably - the H2O2 just decomposes immediately upon addition, with no other noticeable effect; the Cr(III) precipitate remains unchanged.
However, as long as the solution is kept sufficiently basic, it seems to work flawlessly.

chornedsnorkack - 27-3-2017 at 13:40

Quote: Originally posted by The Mad Plater  
citric acid should do the job by itself, since it both acidifies the mixture, and is also itself oxidizable.

I tried it on a sub-gram scale earlier today - ~5mL of 0.35% sodium chromate solution added to ~0.5g citric acid and ~0.5mL 30% H2O2 (a very useful combination!).


Was citric acid/H2O2 mixture itself stable to reaction?
Citric acid as tertiary alcohol should be hard to oxidize.