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Author: Subject: thionyl chloride and potassium dichromate
woelen
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[*] posted on 3-11-2007 at 16:05
thionyl chloride and potassium dichromate


This is a funny experiment: I mixed some thionyl chloride and potassium dichromate. I expected either nothing, or a redox reaction between the potassium dichromate and the thionyl chloride. Dichromate ion easily oxidizes SO2, so I thought it also would do so with SOCl2 (both contain S in +4 oxidation state).

But something totally different happens. Potassium dichromate dissolves sparingly in thionyl chloride, giving a beautiful bright red solution, and a thin orange vapor above the liquid (this must be chromyl chloride, I cannot think of anything else).
When sulphuric acid is added, even then still no redox reaction occurs. The chromium remains in +6 oxidation state, but much more is converted to chromyl chloride. HCl gas (and SO2) escape from the liquid.

When this is added to water, then finally a redox reaction occurs. This reaction is quite spectacular. Have a look at this webpage, and if you happen to own the required chemicals, then try to repeat it. It is a fun experiment.

http://woelen.homescience.net/science/chem/exps/raw_material...

Be careful though, copious amounts of HCl and SO2 are produced, and also small amounts of chromyl chloride escape into the air. Use a fume hood, or do this outside.

EDIT(woelen): Made link working again.

[Edited on 25-10-16 by woelen]




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[*] posted on 3-11-2007 at 16:56


Two observations my friend:
1) From the pix it appears there's some reduction of the Cr+6 to Cr+3 as shown by the green color- does it last or not? why do you say the Cr+6 stays unconverted?

2) Thionyl chloride is presently a hard to come by commodity so I'm saving mine for the day when no other reagent will do.

Good pix!




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[*] posted on 3-11-2007 at 19:10


I suspect that you will find that half of the violent reaction was simply from putting conc. H2SO4 into water, although clearly other reaction happened too and this would have contributed. I have never seen chromyl chloride, but your mixture looked a lot like chromium trioxide. What I think you have to ask is what could have happened in a hypothetical reaction. I cant see how they could react without the presence of water. How much sulfuric acid did you put in?

Also if some reduction had occurred I don't think it would be visible in that deep red liquid.
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chemrox
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[*] posted on 3-11-2007 at 21:23


What, then, is the green if not chromium in solution? He's right he gets Cr2Cl2 by reduction with SO2



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[*] posted on 3-11-2007 at 21:33


Questions about your questions...

Chemrox: Which green color? In the flask at the end or are there colors in the little flask I'm not seeing (I'm colorblind dark green/red/brown, so I'm not much help I guess)? Cr3+ in the early stages would drop out as anhydrous chromium (III) sulfate which I imagine is conc. sulfuric acid insoluble, no?

Antwain: Have you poured a few drops worth of concentrated sulfuric acid into a large excess of water before? It isn't particularly dramatic. The thionyl chloride turning into two mole equivalents of hot gas as it hits the water is cause enough for violence I would think.

Edit: http://woelen.scheikunde.net/science/chem/exps/dichrom/index... The color of sulfite or SO2 reduced Cr(IV) is green.

[Edited on 11-4-07 by UnintentionalChaos]




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[*] posted on 3-11-2007 at 21:38


what the heck is Cr2Cl2?

Edit - Do you mean CrO2Cl2?

@ UnintentionalChaos- I have had a steam explosion by dumping, as opposed to pouring slowly, sulfuric acid into water. But I will concede that one, you get more heat from oxidising things with Cr(VI)

[Edited on 4-11-2007 by Antwain]
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[*] posted on 3-11-2007 at 21:43


Chemrox, what are you talking about? I think he meant CrO2Cl2 (think the acid chloride of chromic acid), but I don't know where the SO2 comes in. Chromyl chloride is hexavalent. Reaction with water affords chromic acid and HCl. Reduction of chromic acid by SO2 liberated by SOCl2 reacting with water yields some complexed Cr (III) salt.

Antwain:This is only 1 1/2 ml of sulfuric acid. I imagine the localized heating contributed to the rapid expansion of gases from hydrolyzing SOCl2 though, so yes, were both right I suppose.

[Edited on 11-4-07 by UnintentionalChaos]




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[*] posted on 4-11-2007 at 03:56


Answering your questions:

@chemrox: Chromium(VI) is reduced to chromium(III), but only when the red liquid is added to water. The green you see in the pictures indeed is chromium(III), but that only is in the water. The first large picture on the page shows the bottle without water, and there you only see a deep red liquid (and some orange vapor).

@UnintentionalChaos: Thionyl chloride reacts with water, but not that violently. That was the first experiment I did with thionyl chloride.

@Antwain: Chromyl chloride is a deep red liquid, with a somewhat purplish hue, and it is very volatile. Its vapor resembles the vapor of bromine very much, it is a little bit more red, less brown. Dilute vapor of chromyl chloride is orange. In the webpage, you can clearly see the vapor of chromyl chloride, formed in the reaction.

I know this compound quite well and have done more experiment with it:

http://woelen.scheikunde.net/science/chem/exps/cro2cl2/index...


Finally, any theoretical reaction with dichromate and thionylchloride would be some reaction, in which the sulphur goes from oxidation state +4 to oxidation state +6. I would expect transfer of oxygen atoms from the dichromate (or CrO3, or CrO2Cl2), leving behind a chromium(III) compound and the thionyl chloride being converted to sulphuryl chloride, or to sulfate.
Many oxidizers can work without water (e.g. manganese(VII) does so very violently), but chromium really needs the water (at least in this case).




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[*] posted on 4-11-2007 at 04:40


Quote:

Many oxidizers can work without water (e.g. manganese(VII) does so very violently), but chromium really needs the water (at least in this case).


This is true, but I would tend to expect that chromium(VI) does need it. certainly the oxidation of alcohols by CrO3 actually goes through chromic acid (at least that is my understanding, but it was a complicated process and I don't think our lecturer understood it completely either). The acid then forms an ester which breaks with the bridging oxygen forming a second bond to the carbon and through electron transfer one of the other oxygens on the chromium extracting the hydrogen from the other side. But that is besides the point.

I guess in some cases hydrogen ions are all that is needed.

This is interesting, you definitely had some chromyl chloride but you think not SO2? (before the water). How about this...

SO2 is "very soluble in water" (79 volumes @ 0*C and 40vol @ 20*C) how soluble is it in sulfuric acid? Whilst it may not be soluble in a mixture of acid and water, I wouldn't be surprised if it has some solubility in the anhydrous acid. You could try heating your mixture gently and condensing/ bubbling the gas through alkali. If the temp is low enough, SOCl2 SO2Cl2 and CrO2Cl2 will not come across. you can then test for sulfur using any of several methods.

Or the reaction without water may be slow. Did it change immediately in colour, or was there a progression?

Or there may be an equilibrium. consider SO2 oxidised to SO3 in sulfuric acid. This is an oxidising mixture. And since chromium does funny stuff in conc. H2SO4 this may change the potential required to oxidise Cr(III) to Cr(VI).

Just some thoughts.
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[*] posted on 4-11-2007 at 05:28


Sorry if things were not formulated sufficiently clear. I did not mention the SO2 together with the chromyl chloride, but I did not mean to tell there was no SO2.

In my reaction equation on the webpage, I do mention the SO2. I indeed expect simulataneously existing SO2 and CrO2Cl2 in the experiment, before water is added. After adding water, these chemicals cannot coexist anymore, the SO2 immediately reduces the CrO2Cl2.




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[*] posted on 4-1-2008 at 17:25
Thionyl Chloride synth?


After reading through the forums and seeing the chromyl chloride synth, I was wandering if it would be possible to make thionyl chloride in a similar way, but using NaS2O7 instead of using dichromate? I'm guessing the equation would look something like this:

Na2S2O7+6H2SO4+4NaCl => 2SOCl2+6NaHSO4+3H2O+O2

Not sure if the O2 would be produced or if SO2Cl2 would be made instead of SOCl2. If this is the case then are the properties of the two sulphur dichlorides similar enough to use SO2Cl2 in place of thionyl chloride? And I'm assuming that hydrolysis of the SO2Cl2 would look something like this:

SO2Cl2+H2O => SO3+ 2HCl

The pyrosulphate could be made from bisulphate:

2NaHSO4 → Na2S2O7 + H2O @ 315°C (wikipedia)
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[*] posted on 4-1-2008 at 19:37


I'm afraid that is not how things go.

See Merck Index for reaction of SO2Cl2 with water.

These really are not ionic compounds.

SOCl2 is prepared by oxidation of SCl2/S2Cl2 system with SO3, or, by reaction of phthaloyl chloride with dry SO2.

SO2Cl2 is prepared by direct union of SO2 and Cl2 over a catalyst, usually GAC or camphor.

It is possible to disproportionate a mixture of PCl3 and SO2Cl2 to POCl3 and SOCl2 again over a GAC fixed bed. Tis is subject of a US patent.

But that's about all there is for thionyl chloride preparation.




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[*] posted on 4-1-2008 at 20:46


It is indeed interesting that SO2 and S(IV) compounds are so resilient against oxidation in anhydrous systems, very contrary to its behavior in aqueous solution.
I have myself once tried to oxidise SOCl2 to SO2Cl2 using KClO3, but without success. The analogous oxidation of PCl3 to POCl3 does work with KClO3, I have read.

SO2Cl2 is a unique compound, very different from SOCl2. You will see this if you drop some into water, the droplets are still there the next morning!

BTW, woelen, to your question about dimethyl sulfite, here is a preparation of this compound:
http://www.lambdasyn.org/synfiles/dimethylsulfit.htm
You see that alcohols can give dialkyl sulfites with SOCl2, it requires adding a slight excess of the alcohol to the SOCl2 and refluxing afterwards. It is distilled in vacuum.

SOCl2 can not only make acyl chlorides from organic acids, but it can also serve as a dehydration agent.
I have synthesized benzonitrile (C6H5-CN) from benzamide by refluxing with an excess of SOCl2. The benzonitrile is a liquid with a sweet almond smell, it can be obtained pure by steam distillation.
Benzamide in turn can be made by dripping benzoyl chloride into ammonia solution.




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[*] posted on 4-1-2008 at 21:38


SOCl2 made by the standard industrial method (SO3 oxidation of mixed sulfur chlorides) is usually sold in rather impure form. Hence the classical purification method in the literature, using boiled linseed oil and quinoline. A rather tedious and lossy procedure.

SO2Cl2 loves to slowly return to SO2 and Cl2. In its preparation from those gases, the challenge is to keep the proportions close to stoichiometry. In the US patent I mentioned, this is accomplished using a UV-Vis spectrophotometer set to monitor the peak absorption of Cl2 in a flow cell, the system being a loop reactor or stirred tank reactor. Peristaltic pumps are used for the circulation through a GAC bed and reservoir, Cl2 and SO2 are fed in through mass flow controllers.

When combined with the interaction of PCl3, to produce a mix od SOCl2 and POCl3, the result is a readily fractionated mixture.

MFCs being a little pricey, it is just as effective, although not as elegant, to let chlorine be in excess, and then turn off the chlorine when sufficient SO2Cl2 has accumulated, and add only SO2 until the yellow green color disappears. Again, a good UV-Vis spec is helpful. I acquired the same model named in the patent, a Sequoia-Turner Model 340. This is a dual voltage device so I have no trouble operating it here on 220. The damned flow cell cost as much new as the used spectrophotometer did, damn it, and then was lost in Hurricane Katrina, so I have to order another one.




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[*] posted on 5-1-2008 at 12:44


Quote:
It is indeed interesting that SO2 and S(IV) compounds are so resilient against oxidation in anhydrous systems, very contrary to its behavior in aqueous solution.
I have done another very nice experiment. I made some SO2 from Na2S2O5 and conc. H2SO4 and heating this mix. This slowly produces some SO2 (only slowly, if water is added, then much more SO2 is produced, but I wanted it really dry).

I lead the SO2 into a dry erlenmeyer, with some finely ground I2 in it. There was absolutely no reaction. I2 granules, sticking to the bottom and side walls were not affected at all, not even after an hour with the erlenmeyer carefully stoppered to prevent the SO2 from escaping.

I heated the glass of the erlenmeyer, and this gives purple vapor of iodine. On cooling down, feather-like crystals of iodine are deposited on the glass, but no reaction with the SO2.

Then I opened the erlenmeyer, put in a single drop of water and quickly closed it again. Only seconds later, the feather-like crystals of iodine liquefied and turned into a brown puddle/droplets. The water vapor now allows reaction of the iodine, giving HI and H2SO4 and this in turn picks up more water and causes the excess iodine to dissolve and form the red/brown droplets.

This is very instructive and really nicely demonstrates that dry SO2 is not oxidized by iodine. Almost certainly, the same is true for oxidation with chromyl chloride.

@DJF90: Your reactions certainly do not occur. Na2S2O7, H2SO4 and NaCl will give rise to formation of HCl, nothing more. On VERY strong heating, some SO3 will be produced as well.

The other reaction also does not occur. With many acid chlorides, the acid chloride is much more reactive than the oxide (e.g. CrO2Cl2 is more reactive than CrO3, NbOCl3 is MUCH more reactive than Nb2O5, COCl2 is much more reactive than CO2, SOCl2 is much more reactive than SO2, etc.), but sulphuryl chloride is an exception. I now have a small amount of SO2Cl2 and I am surprised to see how sluggish its reactions with water are. This compound slowly reacts with water, not giving SO3 and HCl, but only if there isenough water, it will react to form H2SO4 and HCl. But again, this reaction is remarkably slow. A few tenths of ml of SO2Cl2, added to 10 ml of water, takes several days before it has reacted. In alkaline solution it reacts faster. Unfortunately, SO2Cl2 does not form SO2(OH)Cl with a small amount of water. With a small amount of water, only part of the SO2Cl2 is converted to H2SO4 and HCl and the rest remains unreacted.

Making chromyl chloride is not difficult at all:

http://woelen.homescience.net/science/chem/compounds/chromyl...

I made a small fairly pure sample, by adding concentrated ice cold H2SO4 to a very finely ground mix of NaCl and K2Cr2O7. With a very finely and intimately mixed powder, with a slight excess of NaCl, droplets of liquid Cr2O2Cl2 are separating from the acid and these can be pipetted away with a pasteur pipette. I did not distill this compound, it is not very stable and probably needs distilling at reduced pressure. It is really nasty stuff, nice to have a sample of 0.1 ml or so, but nothing more should be stored for a longer time.

Edit(woelen): Made link to my webpage work again.

[Edited on 30-7-16 by woelen]




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