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Author: Subject: Chromium Trioxide - Some questions
woelen
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[*] posted on 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.




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[*] posted on 27-5-2015 at 00:57


Quote: Originally posted by woelen  
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)?
Quote: Originally posted by woelen  


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.

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[*] posted on 27-5-2015 at 14:45


Quote: Originally posted by chornedsnorkack  

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.



I don't believe you'd necessarily have the species Cr2(Cr2O7)3, just a mixture of Cr<sup>3+</sup>(aq) and Cr2O7<sup>-</sup>(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 Cr<sup>3+</sup>. (As woelen showed in the half reaction for CrO3 above)


[Edited on 27-5-2015 by blargish]




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[*] posted on 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]




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[*] posted on 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.
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[*] posted on 15-6-2015 at 18:23


Quote: Originally posted by woelen  
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.




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[*] posted on 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.




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[*] posted on 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?
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[*] posted on 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.
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[*] posted on 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.




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[*] posted on 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.
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[*] posted on 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)
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[*] posted on 1-8-2015 at 07:58


Head cold. Can't sleep.
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[*] posted on 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.




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[*] posted on 2-8-2015 at 02:04


How does wet chromium trioxide behave on heating?
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[*] posted on 2-8-2015 at 02:35


Quote: Originally posted by woelen  
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.


Quote: Originally posted by chornedsnorkack  
How does wet chromium trioxide behave on heating?

I don't know. I bet it sticks to whatever it is sitting on. I'll call that approach plan C.
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[*] posted on 2-8-2015 at 05:17


In my country, CrO3 is often used as a glassware cleaning agent. Two kinds of acid mixtures are prepared from this: chromosulfuric acid (with H2SO4) and chromonitric acid (with HNO3). Chromosulfuric acid, as far as I know, is known worldwide, but chromonitric appears to be a secret weapon of Soviet chemists. The piranha solution was virtually unknown in the Soviet bloc, and chromonitric acid, a mixture with comparable or even superior cleaning powers, was used instead.
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