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[*] posted on 20-5-2015 at 23:02
Chromium Trioxide - Some questions


CrO3 -- what little I know of this compound is intriguing. I have just done a search on SM and found, well, not too much. In the top 10 hits was a comment of mine from last year that my school had a sealed container of the stuff that I wasn't going to open in a hurry. I have a couple of questions.

Question One. I understand it has applications in organic chemistry as an oxidant. OC is not my field and so I quickly get over my head reading. I have also seen a reaction of CrO3 with ethanol with an appropriately placed blast shield that was... let's just say exciting and dramatic. (Periodic table of videos for that one.) I recall reading last year about another reaction and thinking that I would add that to the list of possibilities but can't for the life of me remember what it was.
Anyway, question number one. I have located a cheap source of CrO3 and wondered if it was worth getting some for my home lab. What could I do with it?


Second question. I got around to peeling the tape off the container at my school and having a look at the contents. It had the crimson/rust colour that I expected. The consistency was not what I expected though. It looked wet. Could it have absorbed moisture from the atmosphere? I have no idea how many years the container has been sitting there. I thought CrO3 was incompatible with water and I wasn't expecting deliquescence. The question is what I should do with it now?
Is it actually water that I am seeing? (I didn't spend any time investigating thoroughly.)
Is there a way of cleaning it up?
If I dispose of it, what is the best route for rendering it safe and disposing?
Is there something I can do with it in its present state to use it up?


Any ideas, thoughts or advice muchly appreciated.
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[*] posted on 20-5-2015 at 23:18


CrO3 pulls water and makes chromic acid. it is Very corrosive and used as a pickling agent in powder coating preparation. the crystals are beautiful crimson translucent and well shaped. nice specimen. can be stored dry in a glass container for years without problem. your wet crystals may be dried in a desiccator. I have not used in a prep at this stage.



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[*] posted on 21-5-2015 at 00:08


CrO3 is extremely hygroscopic. In contact with air, it quickly becomes wet and liquefies to a dark brown/red liquid. It is very hard to dry the compound. Drying in a desiccator is nearly impossible, uness some concentrated sulphuric acid is put in the desiccator as well for absorbing water vapour.

I have some CrO3 and I store it in an extremely well sealed container (GL45 bottle with corresponding screw cap and soft silicone inlay to make the seal perfect). Any other container is not suitable and will lead to quick degradation (becoming sticky and wet).

The material is reactive. A nice experiment is to take a small pile of CrO3 (e.g. half a gram) and drop some ethanol or propanol on it. The alcohol inflames at once. Solutions in water are orange (like solutions of dichromates) or dark red/brown if the concentration is very high. Avoid contact with skin of the concentrated solutions, they are very corrosive and severely damage the skin in a matter of a few tens of seconds.

The reason why CrO3 is used in organic chemistry is that it dissolves quite well in many organic solvents, while dichromates, such as K2Cr2O7 do not. For aqueous chemistry experiments, it is better to use K2Cr2O7 and some acid instead of CrO3. K2Cr2O7 is much easier to keep around.




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[*] posted on 21-5-2015 at 02:08


So. Let's see if I have this right.

1. All things considered the CrO3 is in pretty good shape. It has been untouched in the same container (which did not really have a good seal) for at least 8 years and more likely 20.

2. I can use a bit of the CrO3 to make a nice little pyre with some alcohol. I could attempt this with the CrO3 as is or I can dry it in a dessicator with some H2SO4. (That will work quite well. I am teaching a unit on redox in a couple of months. I can demonstrate a change of oxidation states by a change in colour and show that redox is often accompanied by significant delta H. That works in the curriculum.) I will still have 499 grams left.

3. It can be used for some organic chem. That is probably why it was originally purchased. With the current curriculum and my level of inexpertise in the area, that is unlikely to happen.

4. It would be better if possible to dry the material out. That needs to be thought through. It will need a new container though. Double bagged with a dessicant wouldn't hurt.

5. Whatever I do, I need to play safe with hexavalent chromium. A combination of Cr(VI), highly acidic, hygroscopic and highly oxidating spells caution on several fronts.

Any more ideas/thoughts?

[Edit]
The container is exactly the same as in this video.
https://www.youtube.com/watch?v=ScNGJ7Axy1I 14 seconds in.


[Edited on 21-5-2015 by j_sum1]
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Leo Szilard
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[*] posted on 21-5-2015 at 12:47


Yes, chromium vi is not friendly stuff. Carcinogenic. The red/orange color is characteristic of chromium vi compounds. You might be able to dry out the chromium trioxide by carefully placing the bottle (with the cap ajar) into a desiccator with a tray of concentrated sulfuric acid acting as the desiccant.

Chromium vi species are very useful in oxidizing alcohols to ketones, aldehydes, and carboxylic acids. When you conduct the experiment of pouring ethanol over CrO3, you're essentially oxidizing the alcohol into acetaldehyde. This reaction is very exothermic, so the acetaldehyde and any unreacted ethanol "flash boil" and burst into flames.

There are lots of variations of chromium vi oxidizing reagents out there. Two common recipes are a) Jones reagent and b) PCC in dichloromethane. These two recipes, in principle, do the same thing. They differ, however, when you look at the oxidation of primary alcohols like ethanol.

Jones reagent is essentially an aqueous solution of chromic acid. This is prepared by combining water, sulfuric acid, and some chromium vi species like CrO3 or K2Cr2O7. With this reagent, primary alcohols like ethanol will tend to oxidize twice: once to the aldehyde and again to the carboxylic acid.

PCC (pyridinium chlorochromate) can be prepared from CrO3 and is very similar to Jones reagent. The difference is that PCC can be used in an organic solvent such as dichloromethane to the exclusion of water. It turns out that the absence of water makes a big difference in what product you get. With PCC you can selectively oxidize primary alcohols like ethanol once: yielding the aldehyde.

CrO3 is a good oxidizer in general. It would probably oxidize halides into halogen. It might also react in a spectacular fashion with metals but I'm just speculating. With CrO3, you could also prepare chromyl chloride: CrO2Cl2. Now I wouldn't advise this because chromyl chloride is very dangerous (and it's a volatile liquid). But chromyl chloride is a very aggressive oxidizing agent and you might be able to find some pretty spectacular demonstrations of its oxidizing potential.

CrO3 is definitely a cool compound. Have fun experimenting! Just remember your ppe and deal with this compound and its waste responsibly.


[Edited on 21-5-2015 by Leo Szilard]

[Edited on 21-5-2015 by Leo Szilard]
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[*] posted on 21-5-2015 at 20:09


I normally use it with boiling diluted H3PO4 to strip anodizing from aluminum. It's used to indirectly measure the thickness of the anodized layer.

I guess this much is obvious, but since the solution oxidizes alcohol so easily, this also is an easy means of rendering the waste solution less toxic. I slowly add alcohol to the warm, stirred, solution until it all turns dark green. Gradually, the green powder of Cr(II) precipitates out.

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[*] posted on 21-5-2015 at 21:16


Quote: Originally posted by WGTR  
I normally use it with boiling diluted H3PO4 to strip anodizing from aluminum. It's used to indirectly measure the thickness of the anodized layer.

I guess this much is obvious, but since the solution oxidizes alcohol so easily, this also is an easy means of rendering the waste solution less toxic. I slowly add alcohol to the warm, stirred, solution until it all turns dark green. Gradually, the green powder of Cr(II) precipitates out.


I thought it was Cr(III) that was generally produced.
And yes: I had thought that excess ethanol would be the go-to for safe disposal. It's a powerful oxidant and hazardous in the VI state. Reducing it back solves both issues -- at least in part.

How much of an issue is the water if I want to do the ethanol reaction? I noticed that the periodicvideos demo was also quite wet. But mine might be a fraction more so.

What are some safe / interesting alternatives to using an alcohol?

In terms of drying it out, dessicating with H2SO4 has been mentioned. Is it possible or desirable to have the sulfuric acid in contact with the CrO3 in the same manner that one dries / stores iodine?
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[*] posted on 21-5-2015 at 22:35


Quote: Originally posted by j_sum1  

I thought it was Cr(III) that was generally produced.


That's because...you're correct!!! Now if you'll excuse me, I need to get back to smoking my socks.
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[*] posted on 21-5-2015 at 23:54


Drying CrO3 with H2SO4 is NOT meant to be done by adding the CrO3 to the H2SO4. It dissolves quite well in H2SO4 and separating it is nearly impossible.

What you have to do is put a dish with CrO3 in a dessicator and also put a dish with conc. H2SO4 in the same desiccator. Water vapor escaping the CrO3 will then be absorbed by the H2SO4, simply because H2SO4 is even more hygroscopic and water, bonded to the H2SO4 does not escape again from the H2SO4.

CrO3 is quite a strong oxidizer, but what it makes so useful is how easily it works as oxidizer. Oxidizers can be classified as 'strong' in two ways. One way being the redox potential, which is a measure of how well thermodynamically speaking the compound should be capable of oxidizing something. The other way of 'strong' is that a compound has easy/many mechanistic pathways for working as oxidizer. This is not easily quantified. A compound which is strong in both ways is a dangerous compound, which oxidizes nearly everything very easily. CrO3 is quite strong thermodynamically speaking and it is very strong in the sense that if its redox potential is sufficient to oxidize something, then it almost certainly also does so, and usually very quickly.

Comparing it with other oxidizers, there are much more potent oxidizers, which are considered less strong. E.g. peroxodisulfate, S2O8(2-) is one of the oxidizers with the highest redox potential (more than 2 V), but still it is not considered very strong. This is because mechanistically, the breakdown of the S2O8(2-) ion is not that easy and reactions with peroxidisulfate tend to be slow or even non-existent, despite its very high redox potential (only ozone, fluorine and a few highly fluorinated compounds have higher redox potential).

CrO3 is not capable of oxidizing chloride ion to chlorine, at normal temperature. It can react with chlorides, however, to make the very reactive chlorochromate ion CrO3Cl(-) and it can also be used to make the even more reactive CrO2Cl2. These reactions are not redox reactions, the chromium remains in oxidation state +6.

Making the salt KCrO3Cl is interesting. Dissolve some KCl in water and dissolve some CrO3 in ice cold conc. HCl. Mix the two solutions and allow to cool in a fridge. Crystals of KCrO3Cl separate. This salt is deep orange/red and very reactive. It decomposes in contact with water, then it gives K2Cr2O7 and HCl.

Making CrO2Cl2 is also very interesting. Crunch some CrO3 and mix this with NaCl. Add conc. H2SO4 to this mix. Big dark red drops of CrO2Cl2 are formed. This liquid looks very much like bromine, it is slightly more red instead of brown. Its vapor is somewhat more orange instead of brown. Be very careful with CrO2Cl2, it is even more corrosive than CrO3. Use it immediately, do not keep the liquid around. Avoid inhaling the vapor. The vapor is carcinogenic.




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[*] posted on 22-5-2015 at 04:39


Thanks woelen. I have read that CrO2Cl2 is a conclusive test for chloride ions. No analogues are produced with other halides. This makes it superior to a silver nitrate test that is really just a test for halides even though Cl is its most common application.
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[*] posted on 22-5-2015 at 14:15


Quote: Originally posted by Leo Szilard  


Jones reagent is essentially an aqueous solution of chromic acid. This is prepared by combining water, sulfuric acid, and some chromium vi species like CrO3 or K2Cr2O7. With this reagent, primary alcohols like ethanol will tend to oxidize twice: once to the aldehyde and again to the carboxylic acid.

What is the function of the added sulphuric acid?
Quote: Originally posted by Leo Szilard  

CrO3 is a good oxidizer in general. It would probably oxidize halides into halogen.

Which halogen?
Obviously not fluorine.
It´s said that chlorine tends to give CrO2Cl2, not Cl2.
Bromides, yes.
How would you test for chloride if what you get is a brown liquid? Are Br2 and CrO2Cl2 mutually soluble? And how would you spot an impurity of CrO2Cl2 in Br2 vapour or liquid?
Now, iodides, sure, it´s easy to get I2. But is it what CrO3 gives? Can CrO3 be used to oxidize I2 to HIO3 or I2O5?
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[*] posted on 22-5-2015 at 15:28


A saturated solution of CrO3 definitely precipitates after addition of sulfuric acid. The CrO3 is then washed with nitric acid in the funnel and dry air might be sucked through it there, or it goes on a heated tile if nitric vapors are OK.

The more used AFAIK and best form of CrO3 to have IMHO is sodium chromate/H2SO4.

Quote: Originally posted by j_sum1  
How much of an issue is the water if I want to do the ethanol reaction?


Well, you could make acetic acid instead.

Quote: Originally posted by chornedsnorkack  

It´s said that chlorine tends to give CrO2Cl2, not Cl2.


Dichromate/HCl instead does give chlorine, and nicely. Chromyl chloride happens when sulfuric acid gets involved.




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[*] posted on 22-5-2015 at 22:10


Quote: Originally posted by S.C. Wack  
A saturated solution of CrO3 definitely precipitates after addition of sulfuric acid. The CrO3 is then washed with nitric acid in the funnel and dry air might be sucked through it there, or it goes on a heated tile if nitric vapors are OK.

The more used AFAIK and best form of CrO3 to have IMHO is sodium chromate/H2SO4.

Presumably dichromate is better - less metal for the acid...
The solubility of CrO3 in water is quoted as 165 g/100 ml at 0 degrees and 169g/100 ml at 25 degrees. What is the air humidity over that saturated solution?
And what is the saturated solubility of CrO3 in H2SO4? In HNO3?
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[*] posted on 22-5-2015 at 22:55


Ok. So I am building up a picture of what I need to do.

For the wet reagent at school
1. Dry some in a dessicator with conc H2SO4. Do a small batch at first. Write up a procedure and risk assessment and change the storage procedure so this issue does not arise again.
2. Write up a list of what can be done with CrO3 so that the stuff actually gets used. And if it gets used then it can be replaced in due course.

Uses at school
1. Redox demonstrations using ethanol
2. Testing for chlorides
3. Potentially some OC oxidations at some stage in the future.

For my home lab
1. Double check the shipping costs -- dangerous goods shipping may push a cheap reagent into the expensive zone.
2. Dramatic oxidation of ethanol because that's what you do.
3. Production of chromic acid, dichromate and chromate because I only have gram quantities of dichromate
4. Production of chromochloride and interesting chlorochromate salts.
5. Having the ability to test for chlorides when I want to
6. Make glacial acetic acid
7. Experiment with other reducers and CrO3
8. Experiment with Cr(III) salts produced. There are a few interesting reactions that I have seen but not done.
9. Electrolysis and chrome plating
10. Use what I learn to expand the repertiore in the school environment

I think there is enough on that list to justify going ahead.
Anything I have missed?


[Edited on 23-5-2015 by j_sum1]
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[*] posted on 23-5-2015 at 00:17


11. Perfect for dissolving uncooperative students.



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[*] posted on 23-5-2015 at 00:30


Haha.
Actually, I don't have any of those at the moment.

(And I don't mean that I have already dissolved them.)
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[*] posted on 23-5-2015 at 22:22


Quote: Originally posted by j_sum1  

3. Production of chromic acid, dichromate and chromate because I only have gram quantities of dichromate

Does chromic acid even exist?
At room temperature, saturated aqueous solution is in equilibrium with dry CrO3. Can any hydrates be formed on freezing, or does dry CrO3 precipitate at all temperatures to eutectic?
Quote: Originally posted by j_sum1  

7. Experiment with other reducers and CrO3
8. Experiment with Cr(III) salts produced. There are a few interesting reactions that I have seen but not done.

Can CrO3 be reduced to Cr(II) salts? With which reducers? Besides Jones reductor?
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[*] posted on 23-5-2015 at 22:34


Chromic acid? refer to diddi's comment near the top of this thread.

Cr(II) I haven't even investigated. I did see a series of reactions in a YT clip using the Cr(III) product of the ammonium dichromate volcano. If nothing else I can thermite back to metallic Cr.
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[*] posted on 24-5-2015 at 06:22


If you think that you may be interested in doing some organic reactions in the future, an interesting reaction may be to make some pyridinium chlorochromate. It's a very useful compound in organic chemistry that can be prepared using inorganic chemistry. You do need pyridine though. Attached is a paper with the synth(which I stole from the Pyridine thread).

Attachment: 2647-2650.pdf (205kB)
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[*] posted on 24-5-2015 at 10:50


Quote: Originally posted by j_sum1  
Chromic acid? refer to diddi's comment near the top of this thread.

Saturated solubility of CrO3 is under 200 g/100 ml water at 100 degrees. In other words, 2 moles in 5,5 moles of water.
H2CrO4 would mean 555 g in 100 ml water. Which is not a miscible composition under 100 degrees.
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[*] posted on 24-5-2015 at 15:53


I got some CrO3 a couple years or so back. It was a local source for me, so I lucked out in not having to pay any hazmat fees :). It came as large flakes in an HDPE container, which has kept it safely so far with minimal deliquescence. The ethanol reaction is really awesome :D. Once I saw the periodic table of videos clip I knew I had to see it for myself.

I have also used it to produce a few peroxochromate complexes such as potassium tetraperoxochromate(V) and Diperoxotriamminechromium(IV) as per the syntheses found in Brauer's Handbook. Chromates and dichromates can be used for the production of those as well I believe.

Another interesting application that I found in Brauer was the use of CrO3 in producing polychromate ions, such as trichromate and tetrachromate. I knew that potassium trichromate could be produced by dissolving the dichromate in concentrated nitric acid; however, apparently both the trichromate and tetrachromate variants can be produced via careful evaporation of solutions of potassium dichromate in increasing excesses of CrO3. I have yet to try this for myself though.




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[*] posted on 24-5-2015 at 16:11


@ choenedsnorkack
I am not sure I understand the problem that you describe. CrO3 + H2O --> H2CrO4. That seems pretty logical to me (now that I have seen it. I have learned a bit over the past few days.) It is pretty much analogous to SO3 in that regard.
The bottom line is that my CrO3 has absorbed enough moisture to become a viscous tarry-looking substance. I fully expect it to be acidic and the acid to be H2CrO4.


@Blargish
Fascinating stuff. I have not even heard of trichromate and tetrachromate.

I think Cr is slipping up the ranks of my favourite transition metals.
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[*] posted on 25-5-2015 at 23:01


Quote: Originally posted by j_sum1  
@ choenedsnorkack
I am not sure I understand the problem that you describe. CrO3 + H2O --> H2CrO4. That seems pretty logical to me (now that I have seen it. I have learned a bit over the past few days.) It is pretty much analogous to SO3 in that regard.

Mixtures of SO3and water are liquid at +20 degrees for all compositions between water and approximately 30 % oleum. (30 % oleum is saturated with solid disulphuric acid, which melts at +35 degrees). The exact composition of H2SO4 is a single phase liquid for all temperatures from +10,4 degrees to +280 degrees. And crystallizes as a single phase solid below +10,4 degrees.
The gross composition of H2CrO4 is not available between 20 and 100 degrees as a single phase, either liquid or solid.
Quote: Originally posted by j_sum1  

The bottom line is that my CrO3 has absorbed enough moisture to become a viscous tarry-looking substance. I fully expect it to be acidic and the acid to be H2CrO4.

I also expect it to be acidic, but still I expect it to be mostly water.
Also, how much of the solutes is H2CrO4, how much is H2Cr2O7 and how much is other acids?
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[*] posted on 26-5-2015 at 01:47


Quote:
Does chromic acid even exist?

Quote: Originally posted by chornedsnorkack  
The gross composition of H2CrO4 is not available between 20 and 100 degrees as a single phase, either liquid or solid.

Thanks for clarifying what you mean. In the context of this discussion, that is perhaps a narrow definition of acid. Dilute solutions of H2SO4, HCl, or whatever are commonly called acids and rightly so. If a quantity of CrO3 absorbs sufficient moisture from the atmosphere to fully dissolve then I would think it is entirely appropriate to label it chromic acid.
What the exact composition of my container is, I don't rightly know. It appears as an extremely thick plasticky substance like a melted toffee or maybe window putty. I have not weighed it yet, but the contents purportedly started out at 500 grams. (I don't know how much has been used in its history, but I would bet very little.) It is certainly homogenous at a macro scale. It might be good to speculate on its composition: H2O, H2CrO4, H2Cr2O7, H2Cr3O10, H2Cr4O13, CrO3 are all contenders. However, all of that is academic rather than practical. What matters to me is
(a) Can I dry it out?
(b) Do I need to dry it out?
(c) What fun can be had and what interesting and educational things done?

Thanks to everyone's help, I am slowly building up some answers to these questions. Desiccator is being set up tomorrow.
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[*] posted on 26-5-2015 at 23:48


Quote: Originally posted by j_sum1  
In the context of this discussion, that is perhaps a narrow definition of acid. Dilute solutions of H2SO4, HCl, or whatever are commonly called acids and rightly so. If a quantity of CrO3 absorbs sufficient moisture from the atmosphere to fully dissolve then I would think it is entirely appropriate to label it chromic acid.
(b) Do I need to dry it out?
(c) What fun can be had and what interesting and educational things done?


These two questions are closely related.
What you now have is
a) a concentrated aqueous solution,
b) of not exactly known concentration.

If you want to perform experiments with a dilute aqueos solution of chromic acid (I assume mostly H2Cr2O7 below pH of about 6) then you can produce a dilute solution by diluting concentrated solution just as well as by dissolving dry chromium trioxide.
If you want a dilute aqueous solution of exactly known concentration, like for redox titration, then you might dry the concentrated solution of unknown concentration and then weigh the dry chromium trioxide - or else you might produce a dilute aqueous solution of not exactly known concentration and then measure its exact concentration against a reagent whose exact concentration you do know.

Are there any interesting reactions which require using chromium trioxide in dry media? Like production of pyridinium chlorocromate?
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