Sciencemadness Discussion Board - Chromium Trioxide - Some questions

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.

woelen - 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.

j_sum1 - 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.

chornedsnorkack - 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?

S.C. Wack - 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.

How much of an issue is the water if I want to do the ethanol reaction?

Well, you could make acetic acid instead.

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.

chornedsnorkack - 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?

j_sum1 - 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]

diddi - 23-5-2015 at 00:17

11. Perfect for dissolving uncooperative students.

j_sum1 - 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.)

chornedsnorkack - 23-5-2015 at 22:22

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?

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?

j_sum1 - 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.

gdflp - 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)
This file has been downloaded 482 times

chornedsnorkack - 24-5-2015 at 10:50

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.

blargish - 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

. 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.

j_sum1 - 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.

chornedsnorkack - 25-5-2015 at 23:01

@ 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.

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?

j_sum1 - 26-5-2015 at 01:47

Quote:

Does chromic acid even exist?

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.

chornedsnorkack - 26-5-2015 at 23:48

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?

woelen - 26-5-2015 at 23:50

Even if it is wet and sticky, you still can do most experiments with it. E.g. the alcohol burning experiment also works when CrO3 is wet and sticky.

Solutions in water are deep orange and can be used as oxidizer and then are turned themselves in green or blue/grey solutions containing chromium(III).

For experiments with hexavalent chromium(VI) you need a lot of excess acid if you want smooth and clean reactions. Have a look at the half reaction:

CrO3 + 6H(+) + 3e --> Cr(3+) + 3H2O

A lot of additional acid is needed for the redox reaction in which CrO3 is reduced to a soluble chromium(III) salt. With dichromates even a little more acid is needed (now for each CrO3-unit, 7 H(+) ions are needed):

Cr2O7(2-) + 14H(+) + 6e --> Cr(3+) + 7H2O

-----------------------------------------------

CrO3 and K2Cr2O7 can oxidize HCl to Cl2, but not really cleanly and smoothly. At room temperature the reaction is very slow, but when the solution is heated to boiling, then indeed Cl2 can be produced. But you will have a hard time to get a purely green solution of trivalent chromium by heating a solution of CrO3 or K2Cr2O7 in aqueous HCl. You will end up with a brown solution, the brown color caused by a mix of green chromium(III) and remains of orange/red dichromate and red chlorochromate(VI). Only prolonged heating will make the solution purely green.
Oxidation of bromide and iodide proceeds more smoothly. Iodide is oxidized immediately (provided sufficient acid is present as well, see above), bromide can easily be oxidized completely by gently heating the solution. By stronger heating then the bromine can be boiled off and collected in another flask.

chornedsnorkack - 27-5-2015 at 00:57

Even if it is wet and sticky, you still can do most experiments with it. E.g. the alcohol burning experiment also works when CrO3 is wet and sticky.

Water dilutes reagents and dilutes reaction heat. Below which concentrations of chromic acid should C₂H₅OH be quietly oxidized to CH₃COOH and not burst in fire?

Also:
Acids are known to hydrolyze ethers.
How fast is hydrolysis of primary and secondary ethers in concentrated aqueous chromic acid, compared to oxidation of resulting primary or secondary alcohol? How well do such ethers, like diethyl ether, tetrahydrofurane, 1,4-dioxane et cetera tolerate aqueous chromic acid? And do they tolerate (and dissolve) dry chromium trioxide?

What are oxidized faster by chromic acid: primary alcohols, or aldehydes? When aqueous chromic acid is added to excess of ethanol, what is formed: ethanal (because it forms faster and chromic acid is consumed before acetic acid is formed) or acetic acid (because every mole of ethanal is oxidized to acetic acid before next amount of ethanol can react)?

For experiments with hexavalent chromium(VI) you need a lot of excess acid if you want smooth and clean reactions. Have a look at the half reaction:

CrO3 + 6H(+) + 3e --> Cr(3+) + 3H2O

A lot of additional acid is needed for the redox reaction in which CrO3 is reduced to a soluble chromium(III) salt. With dichromates even a little more acid is needed (now for each CrO3-unit, 7 H(+) ions are needed):

Cr2O7(2-) + 14H(+) + 6e --> Cr(3+) + 7H2O

And CrO₃ is an acid you´ll need a lot.
Because the prevalent acid is H₂Cr₂O₇ below about pH of 6. Whereas Cr³⁺ is strongly hydrolyzed above pH of about 4. Meaning that if you reduce a chromic acid solution, your product should be Cr₂(Cr₂O₇)₃. 3/4 of your chromium (VI) remains unreacted, and at pH of below 4 is still fairly strong oxidant.

So - you´d need enough unreactive acid (like H₂SO₄) to keep all your resulting Cr³⁺ in solution.

blargish - 27-5-2015 at 14:45

I don't believe you'd necessarily have the species Cr₂(Cr₂O₇)₃, just a mixture of Cr3+_(aq) and Cr₂O₇-_(aq) in aqueous solution. Also, I've found that adding a reducing agent (sodium thiosulfate) directly to an aqueous solution of CrO₃ results in no apparent reduction whatsoever, suggesting that a solution of CrO₃ isn't sufficiently acidic. You need an external source of acid no matter what in order to drive the reduction to Cr3+. (As woelen showed in the half reaction for CrO₃ above)

[Edited on 27-5-2015 by blargish]

woelen - 28-5-2015 at 00:05

I have the same experience as blargish. Without acid, aqueous solutions of CrO3 and also Cr2O7(2-) only work like measly oxidizers. There is some redox reaction, but it is slow and has no nice stoichiometry.

E.g. if you mix a solution of excess Na2SO3 and CrO3 then initially the liquid remains orange, with a brown tinge. Slowly the solution turns brown, but the reaction is incomplete. As soon as some dilute H2SO4 or dilute HNO3 is added, the liquid turns green at once. With the acid, the reaction is really fast (much less than a second, even in the cold).

Quote:

Water dilutes reagents and dilutes reaction heat. Below which concentrations of chromic acid should C2H5OH be quietly oxidized to CH3COOH and not burst in fire?

As long as there still is solid CrO3, the reaction will be sufficient violent to ignite the alcohol. The practical performance of the experiment is quite simple. Take half a gram or so of the CrO3. Crunch the pieces somewhat (no need to make very small particles, just breaking up larger pieces somewhat is sufficient) and scrape them together to a little pile. Then drop some alcohol on the pile. Nearly immediately the alcohol inflames. This also works for wet and sticky material. If the CrO3 has liquefied completely, then I think it will not work anymore.

[Edited on 28-5-15 by woelen]

j_sum1 - 29-5-2015 at 03:52

An update.

I set about putting some of the CrO3 sludge in a desiccator. A couple of interesting things.

Firstly, I had some fun with the whole risk assessment process. We use a great online package to manage risk assessment at the school. The system flipped out at the mention of CrO3. It listed it as a prohibited substance. There is a bit of a work-around. We are a private school and in a different state and so we are not necessarily bound by the restrictions imposed by our software. However, I am filling out this paperwork with all kinds of metaphorical alarm bells ringing. It comes time to assess the risk without the control measures and it lists as extreme. I figure that is actually reasonable. I am setting up some equipment with an open container of concentrated sulfuric acid and an open container of chromium trioxide. There are obvious spill hazards, the whole corrosion issue with the sufuric acid, not to mention the CrO3 being simultaneously corrosive, acidic, desiccant, powerfully (or rather quickly) oxidising and carcinogenic. I figured "extreme" was ok. This is school-based risk assessment documentation that itemises the dangers of sword-fighting with burettes. There is no way that I am letting students near any of this. But without that control measure, yep, extreme is an appropriate word.
So, I went into detail on all the control measures that would be taken -- teacher only: no students present, limited quantity, full PPE, handling procedures, spill mitigation and containment practice, reduction of residues on spatulas etc to Cr(III), disposal of waste. I figured that with all of this in place the risks were reduced to low. But the use of a prohibited substance meant that the previously stated "extreme" was still the most prominent feature of the document. So, I have the school's workplace health and safety officer doing conniptions and my head of faculty asking deep and serious questions about why I was attempting extremely dangerous experiments. And then once I had reassured him there were all the necessary questions about documentation and liability and so forth. Fun moments.

The actual CrO3 had the appearance of tar or treacle. The container originally held 500g. I have no idea how much had been used in the past. The container plus contents now weighs 590g. I estimate 50g for the container and probably no more than 100g used (although that is a guess). That would suggest 400g of material that had absorbed a further 140g of water.
Whatever it was, there was no lovely flaring reaction with alcohol. A pity. I concur with what blargish and woreen stated -- that a bit of acid is needed to assist the redox reaction with sodium sulfite. Really sluggish without. For some of my utensils I switched to zinc powder for the clean up. It was a bit simpler.
So now I need to leave it for a few weeks to see how much the stuff dries out. I have weighed both the chromium trioxide and the sulfuric acid and both containers so I can track the mass transfer. Let's see how it works.

S.C. Wack - 15-6-2015 at 18:23

But you will have a hard time to get a purely green solution of trivalent chromium by heating a solution of CrO3 or K2Cr2O7 in aqueous HCl. You will end up with a brown solution, the brown color caused by a mix of green chromium(III) and remains of orange/red dichromate and red chlorochromate(VI). Only prolonged heating will make the solution purely green.

In my hands dichromate turns green quickly; the whole thing closely resembles MnO2/HCl except for the big mess at the end with MnO2. I wonder if maybe you're not using the right amount of HCl. I've seen different suggestions such as 2 parts dichromate to 17 parts 32% HCl, but I think something closer to the theoretical amount is better such as 180 g. per liter.

K2Cr2O7 + 14 HCl = 2KCl + 2CrCl3 + 3Cl2 + 7H2O

If the generator set up so that any reflux on the glassware is heated, to actually make the chlorine leave the flask, the dichromate/HCl is perfectly fine. It's supposed to give a high-purity product, after the usual water and sulfuric acid traps at least.

woelen - 17-6-2015 at 23:17

If it closely resembles MnO2/HCl, then you only have partial reduction of hexavalent chromium to trivalent chromium. Pure chromium(III) has a beautiful bright deep green color and no brownish/olive colored hue at all. MnO2/HCl gives dark green/brown solutions, which are not like chromium(III) in HCl at all.

If you want to see the difference do the following:
- In one test tube put some dichromate and add HCl
- In the other do the same and add a pinch of sulfite and watch the difference.

Tdep - 1-8-2015 at 06:45

Hey j_sum1, is there another update on this? I made some CrO3 today from a boiling solution of K dichromate and sulfuric acid. It honestly was rather beautiful, it's such a lovely red.

Currently it's a red sludge as I can't pull much a of vacuum to dry it. It's in a desiccator with CaCl2, which is likely to just dry out the CaCl2 more but hey, worth a shot? How is the acid helping the drying?

j_sum1 - 1-8-2015 at 07:18

Not much to report.
Two months with 40 grams of the wet CrO3 in a desiccator with an open container of H2SO4. About 3 grams of water made its way from the CrO3 to the acid. Strangely, the H2SO4 gained more mass than the oxide lost -- by nearly a gram. Not too sure where that came from.

It reached a point where, as far as I could tell, it wasn't losing more water. I wasn't weighing it continuously but it was no longer changing appearance and it looked to be dry on the surface. It turns out that it had the consistency of badly made toffee. Black and treacly in appearance, tacky at the surface, pliable with a lot of effort, but not at all brittle, It behaved as a very viscous liquid -- flowing flat again if left for a while. I could not prise any from the glass petri dish without breaking it. And it stained anything that touched it with a tarry-looking smear. Those smears were visibly pulling moisture from the atmosphere. It started to dribble on my gloves after a short time.

I attempted the classic reaction with alcohol. Nothing happened at all. I guess it was probably oxidising the alcohol, but not the vigorous reaction that causes the alcohol to flare up.

My thoughts are twofold.
1. I doubt a desiccator is an effective means for drying this stuff out. A bit of heat might be more effective.
2. I think that this particular container may have pulled other substances than water fro the atmosphere. It could easily be 20 years old. It has been in a chemical store that also holds HCl, organic liquids and iodine. Sure it was lidded but the lid obviously wasn't effective enough to stop it absorbing something. Quite what it is now is anyone's guess. It is an ugly black and not a nice red.

What to do with it now is a problem. There are few options.
1. Attempt to dry a sample using heat.
2. Use it to make some sodium or potassium dichromate -- and recognise that it is going to have some impurities in it.
3. Reduce it back to Cr(III) and dispose of it.

I am open to suggestions.

deltaH - 1-8-2015 at 07:25

Did you use enough H2SO4? Maybe it worked somewhat but hit an equilibrium before completion.

I'd remove the wet H2SO4, add much more fresh H2SO4 and leave it to dry again. If necessary, repeat again until it's done.

j_sum1 - 1-8-2015 at 07:35

It's possible. I used 170g of acid to 40 grams of CrO3. I would have thought that would be enough. But ultimately I don't see this as a practical route at the scale we have. My school is not even supposed to have CrO3. Leaving large volumes of it in containers with open vessels of concentrated sulfuric acid is a bit of a risk assessment headache. I could do a few grams. But with the very limited success so far I don't think it is worthwhile to pursue this route. Besides, I am quickly coming to the conclusion that it is not just water contaminating.

Tdep - 1-8-2015 at 07:44

Ah, that's disappointing. I'm not going to have much luck with CaCl2 then, especially not in the impatient timescale I like to work in. Might try some heat or a few other things, when it stops raining.

(I like your commitment to science at all hours, what is sleep)

j_sum1 - 1-8-2015 at 07:58

Head cold. Can't sleep.

woelen - 1-8-2015 at 10:48

If I read this, then I would go for making K2Cr2O7. Dissolve all of the CrO3 in water and add a stoichiometric amount of a solution of KOH or K2CO3 (careful: mix very slowly, the acid/base reaction is exothermic and if the carbonate is used you will have a lot of foaming and droplets with hexavalent chromium are released into the air). The solution will be bright orange. Making pure K2Cr2O7 is not hard. The salt crystallizes very well and is not hygroscopic. Even if your CrO3 is impure, you still can make nice bright orange K2Cr2O7 fairly easily.

Try with a small amount whether you can get a nice bright solution with KOH.

chornedsnorkack - 2-8-2015 at 02:04

How does wet chromium trioxide behave on heating?

j_sum1 - 2-8-2015 at 02:35

If I read this, then I would go for making K2Cr2O7. Dissolve all of the CrO3 in water and add a stoichiometric amount of a solution of KOH or K2CO3 (careful: mix very slowly, the acid/base reaction is exothermic and if the carbonate is used you will have a lot of foaming and droplets with hexavalent chromium are released into the air). The solution will be bright orange. Making pure K2Cr2O7 is not hard. The salt crystallizes very well and is not hygroscopic. Even if your CrO3 is impure, you still can make nice bright orange K2Cr2O7 fairly easily.

Try with a small amount whether you can get a nice bright solution with KOH.

That is my current plan. I will put the broken petri dish in some water, filter out the glass and take it from there. The problem is that, since I don't know how much water is absorbed, working out the mass ratio of reagents will be impossible. I guess I could do a titration with some iodide and starch. But it might be easier simply to stir thoroughly and watch the pH.