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Jor
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[*] posted on 28-1-2009 at 01:39
anhydrous CrCl3


Since I have seen some is this stuff in large crystal form, I really want to have some too. The compound forms beautiful intense purple/violet crystals, the most beautiful I have ever seen.

It is very expensive, more than 1 euro per gram.

So I have been looking for way of making it myself. These are the ones I have found:
-Pass hot CCl4 vapour over chromium oxide (Cr2O3) in a tube furnace, at 700C if I recall right. The CCl4 is not a real problem, the tube furnace is. So this is a no-go for me.
-Reflux the chromium chloride hexahydrate with thionyl chloride. I have thionyl chloride, about 20mL, but it is the last I will ever have probably, because of it's CWC status. There I'd like to avoid this route.

Are there any other suitable methods for the home chemist, not involving very hard to get chems or methods?

Would heating the hexahydrate be possible, or will you just decompose to oxychloride/oxide?
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[*] posted on 28-1-2009 at 02:43


The other listed method is passing chlorine over chromium metal, yet another tube furnace route.

CrCl3 heated in the absence of Cl2 breaks up into CrCl2 + Cl2, no temperature given for 'heated', and if heated in air reacts to give Cr2O3.

It is possible, but I make no promises, that anhydrous CrCl3 might be prepared by mixing the hexahydrate with at least a 12 mole ratio of NH4Cl, mix well and grind/mill if you can, and quite slowly heating the mixture to 300 C while applying mild suction (H2O, NH3, and HCl being released), and then finally increasing the suction to a decent vacuum while continuing heating in order to drive off excess NH4Cl.
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[*] posted on 28-1-2009 at 03:44


A tube furnace shouldn't be too hard to construct yourself, if you have Pyrex or Vycor (preferably the later as it can get a bit hotter than Pyrex) tube (easily purchased from eBay or a local glassblower), two ceramic heat proof mats with holes drilled in centre for the tube to go through. Then a couple of MAPP or Propane burners lined up to heat the centre of the tube is an easy set up to achieve and should work here.



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[*] posted on 28-1-2009 at 10:16


Is the by-product of the CCl4/Cr2O3 reaction phosgene?
If so, it might be good to know about that before you start.
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[*] posted on 28-1-2009 at 10:41


Yes, ofcourse.
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[*] posted on 28-1-2009 at 11:22


If you are having phosgene as a by product I would suggest dissolving it in toluene and keeping the solution for future use. It is a very usefull chemical and a solution of it in toluene is a very convenient storage method.



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[*] posted on 28-1-2009 at 20:56


You can buy quartz sleeves which housed UV bulbs used in water treatment quite cheaply. Replacement coils for 117 volt kilns are also cheap. The coil wraps around the tube and inexpensive kiln repair mix can be used to build up around the assembly creating a very cheap tube furnace. One simple way is to center this inside a larger metal tube and pour the refractory mix in to fill up the inner space which you then let dry and harden. Difficult to get back apart if you burn the coil open but simple to build a new furnace when this happens. Be sure the turns do not touch, steel wire can be twist tied at the ends to hold the coil in place. After your refractory sets the movement of the coils will not be an issue.

I took a light dimmer apart and paralleled two 10 amp triacs mounted on a heat sink for a cheap power control. Buy a dremel diamond cutting wheel for 15 bucks and you will find the quartz tube cuts like butter. You can also make fittings and seal them to the tube ends with the refractory mix and gas is easy to pass through the under 50 dollar tube furnace which works better than you might think for something so easy to build. I made a thermistor circuit work with the light dimmer pot to give me both adjustable heat range with temperature regulation. Mount the thermistor close but not too close to the heating coil and calibrate the heat control by taking measurements in the hot portion of the furnace.

I think you are shying away from a tube furnace when it is much easier and cheaper to build than you might think. On a side note a homemade bimetallic strip controller is also very easy to build which eliminates the electronics and can handle high current very well.
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[*] posted on 29-1-2009 at 03:23


IrC, that looks like a not too hard to perform method, however....

I would want to build a tube furnace, but I dont have the time, room and money for it. Besides, I have never needed it before, and I'm not wanting to waste time, room and money, for the synthesis of a single compound.

But I appreciate the advice on how to build one greatly ! :P
When I want to build one sometime, I know how!

Panziandi, I don't have toluene. Besides, phosgene may be useful, but it is simply not fun to play with, too toxic. I tend to avoid the gasses AsH3, H2Se and COCl2 at all cost unless milligram amounts are evolved.
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[*] posted on 29-1-2009 at 03:38


Then I think your best bet is to dehydrate the CrCl3.6H2O with your precious thionyl chloride... Unless heating the hexahydrate in a brisk current of dry HCl would work?



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[*] posted on 29-1-2009 at 08:33


Heating with NH4Cl is similar to heating in a stream of dry HCl, but a little easier to do because there's no HCl generator to make. It works with many metal halides, for some the carbonate or oxide can be used instead of the hydrated halide.
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[*] posted on 29-1-2009 at 21:41


Here is pontentially useful method of producing anhydrous metal chlorides from metal oxides. Dichloro Disulfane is extremly easy to make and the procedure could possibly be adapted to hydrated metal chlorides. Petrolium ether may work as a replacement for the carbon disulfide extraction of the final product.
Procedure for producing vanadium (III) chloride from vanadium pentoxide. Copied from the Hand Book of preparative Inorganic Chemistry; available in the Sciencemadness library.
2V2O5 + 6S2C12 = 4VCL, + 5SO2 + 7S
Fine, pure V2O5 powder (18 g.) and 40 ml. of S2Cl2 are refluxed
under anhydrous conditions for 8 hours (constant stirring). The
excess S2C12, containing dissolved S, is decanted and the VC13
formed is washed with dry CS2 . Adhering volatiles are removed by
heating the material at 120-150°C under vacuum or by extracting
it for several hours with CS2 in a Soxhlet apparatus. After thorough
purification, the residual sulfur content of the resulting fine crystals
of VC13 is about 0.2%. The yield is about 30 g.
Coarse (and hence less hygroscopic) crystals of VC13 are
obtained by heating the fine crystalline product with a small amount
of fresh S2C12 in a sealed tube at 240°C. Since no gas is evolved
in this operation, large amounts can be treated at one time.
The same reaction can also be carried out in a sealed tube at
300°C; however, smaller quantities must be used in this case (6-7 g.
of V205 and 20 ml. of S2C12).


:D

[Edited on 29-1-2009 by benzylchloride1]
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[*] posted on 29-1-2009 at 22:20


That works for metals with fairly reactive oxides, which isn't the case with the Cr2O3 you normally find. I believe it does work at higher temperatures, similar to the CCl4 based method.

Note that the example reduces V5+) to V(3+)
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[*] posted on 29-1-2009 at 22:32


This reaction may work for the hydrated chlorides; the S2Cl2 reacts with the water to produce SO2 and HCl. The reaction with the V2O5 is a redox reaction and sulfur is produced as a by product and would not work with unreactive oxides. Another method that I have read about is heating the metal oxide mixed with carbon under a stream of chlorine gas in a tube furnace. This procedure starts with the unreactive TiO2 and produces TiCl4. Anhydrous chromium (III) Sulfate can be made by heating the hydrated salt in a 110 celsius oven for about 4 days, I recently used produced Cr2(SO4)3 by this procedure for an inorganic lab at school. The sulfate was then reacted with anhydous ethylenediamine to produce tris-(ethylenediamine) Chromium (III) Sulfate. I have conducted the procedure with the S2Cl2 for producing VCl3, the product is extremely moisture sensitive and soon turned gray due to atmospheric moisture.
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[*] posted on 4-4-2009 at 07:35


A few questions. First I will explain what I am going to do.

Im going to prepare some hydrated CrCl3 soon.
This will be done by heating a solution of potassium (or sodium) dichromate in concentrated hydrochloric acid (chlorine evolves). My first idea was using ethanol or methanol as the reducer, but I'm afraid acetate/formate will form a complex with chromium(III). Next I will evaporate ALMOST to dryness (to prevent hydrolysis) and let crystallise.
After filtering, I should have a 1:1 mixture of CrCl3.6H2O and KCl.

Next, I will add ethanol to solvate the chromium(III)chloride. NaCl and KCl are both 'slightly' soluble in ethanol. Wich one is the least soluble? I find data like 'slightly soluble' , but I have not find yet exact data. This will determine if I use Na of K dichromate.

After the extraction of CrCl3.6H2O with ethanol, and filtering, I will evaporate the ethanol. Problem is, will i just get the hydrated chromium chloride, or will I get chromium chloride with one or more ethanol ligands? If yes, would, after evaporating ethanol, adding some water and evaporating again replace all ethanol ligands with water ligands?

I have also been thinking about CrO3, but I only have 25g of this (donation from woelen :) ), and it is also nasty stuff to handle.

I will then try heating the hydrated salt with ammonium chloride, and after sufficient CrCl3 anhydrous has formed, i will add water to solvate all other materials (CrCl3 is insoluble in water, in the abscence of Cr(II) ).

Otherwise, maybe heating with acetyl chloride will help instead of heating with thionyl chloride. Or can't the water ligands be removed by acetyl chlordie on reflux?
I guess the acetic acid can't form a complex with the chromium? i think you need acetate for that, wich will not form in conditions where there is no water, just acetyl chloride.
And even if some will form, it can later be extracted from the purple anhydrous CrCl3, leaving CrCl3 behind (unless my flask is contamined with a very small piece of zinc :D ;) ).

EDIT: I just found out dichromate oxidises Cl(-) in very acidic environment only very slowly. I added 1mL of sodium dichromate solution to 1ml of 37% HCl, and after boiling for 1 minute, the solution is dirty brown (a mixture of dichromate and Cr(III) ). So i have to find another reductor now, wich will still give me CrCl3.6H2O as product. I could also start from chromium aluim, but I rather don't as I only have 30 grams or so, and I would have to do it in 2 steps:
-precitipate hydrated Cr2O3 with dilute ammonia
-dissolve in concentrated HCl
Any ideas?

[Edited on 4-4-2009 by Jor]
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[*] posted on 4-4-2009 at 21:07


Once again, download Seidall's solubilities of inorganic and organic compounds from the Internet Archive or Google books. There's several pages for each of solubilities in various alcohol-water mixtures as well as several different alcohols.

Make a strong solution of the dichromate, heat it on a water bath. Make a mixture of hydrochloric acid and alcohol, drip this into the hot dichromate. The alcohol is oxidised to acetaldehyde and vapourises off; some get oxidised further to acetic acid.

Don't count on the CrCl3 in the heated CrCl3-NH4Cl mixture not to have any Cr(II) in it. Really, the best way to clean it up is to start with a goodly excess of NH4Cl, and finish by heating under reduced pressure to sublime off the excess NH4Cl.

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[*] posted on 8-4-2009 at 03:48


The best way to make CrCl3 is by dissolving pure chromium metal. Right now, you can buy 200 grams of 99.9% chromium on eBay (Jarmondbrinkley) for a very good price. Add this metal to excess 25% HCl and let it dissolve. It takes a while, but in a day or so all of it has dissolved and you get a remarkably dark green and somewhat viscous liquid. This can be evaporated on a luke-warm place and you end up with a very viscous green syrup of CrCl3.xH2O with some traces of HCl in it.

I tried making anhydrous CrCl3 from dark green CrCl3.6H2O by refluxing the solid suspended in SOCl2 for half an hour or so and then leaving the solid standing in the warm SOCl2 for several hours. No result at ALL!! Some of the SOCl2 reacts giving SO2 and HCl (you can see bubbles of gas coming from the solid CrCl3.6H2O) but this reaction soon stops. I think that the only water which is decomposed is the water which is not in the CrCl3.6H2O, but extra water, due to this compound's highly hygroscopic nature. So, Im could make dry CrCl3.6H2O, but no further dehydration occurs. I conclude that SOCl2 is not suitable for making anhydrous CrCl3.

I also did the experiment with PCl5 and this does give pink/purple CrCl3 but only with great difficulty. I obtained spots of pink/purple in the green solid. the main problem is that PCl5 is a solid and the solid-solid reaction does not proceed smoothly.

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

Another option is to take Na2Cr2O7 (or the potassium salt), use dilute sulphuric acid, and use sulfite as a reductor. Add a slight excess of sulfite. Then add excess ammonia to precipitate the chromium(III) and boil to make the precipitate somewhat more compact. Filter, rinse with distilled water and dissolve in dilute HCl. Reprecipitate with slight excess of ammonia and repeat the action. In this way you will have a product free of coprecipitated sulfate ion. It's more of a hassle than dissolving Cr in conc. HCl, but if you have plenty of dichromate and plenty of sulfite, then it is another option.




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[*] posted on 8-4-2009 at 06:43


I am now indeed seriously considering buying some Cr-metal. I can also use it to make some Cr(II) in solution and who knows what more.

Thank you very much for performing the experiment woelen! :)
Strange, I have seen a reference about the route, although I cannot remember where I have seen it. The 2 routes in Vogel are Cr reaction with Cl2 and Cr2O3 rxn with CCl4. Both use very high temperatures.
I have thought about reacting chlorine with chromium at lower temperatures I can reach with my burner, and although I think it will react albeit very slowly, it will be impossible AFAIK to remove the unreacted chromium metal from the anydrous CrCl3. For example, adding hydrochloric acid will produce Cr(II) and CrCl3 will quickly dissolve.

Very interesting about the reaction with PCl5 though. Maybe it will be possible to heat to the melting point of PCl5 to liquify it, and then, all POCl3 is instantly boiled away. Then, PCl5 is either evaporated as well, or one could destroy it by slowly adding water, and washing out the phosphoric acid by several washes with water, followed by drying at high temperatures (200C). I think none Cr(II) will be present, I don't see where it should come from, so no anhydrous CrCl3 should dissolve.
After all Brauer washes the CrCl3 with boiling HCl is both procedures.

Or does PCl5 sublime before it melts? If so, maybe a solution of the stuff in CCl4 can be refluxed with Cr2O3 (active Cr2O3, this can be ontained by decomposing ammonium dichromate), followed by distillation of the CCl4, and evaporating/subliming the PCl5, followed by washes with water. This way you don't have a solid-solid reaction.

I haven't made any PCl5 yet. Can you maybe perform the first or 2nd experiment on a test-tube scale woelen, look if it is possible to get pure CrCl3 from these methods? :P I'm not sure if they work, I just made them up myself.

I'm not sure, if you get those very nice crystals by these methods. Maybe the stuff should be sublimed.

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

I indeed have plenty of sulfite and dichromate, but it seems like a lot of hassle for CrCl3.6H2O. I rather leave from Cr-metal, I will buy it soon I think.

[Edited on 8-4-2009 by Jor]
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[*] posted on 9-4-2009 at 17:34


What could you use the CrCl3 for? An obvious use would be as chromium plating, by electrolytic reduction..
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[*] posted on 12-4-2009 at 09:14


Jor, PCl5 hardly liqufies, it sublimes, giving some kind of rime at cooler parts of the test tube. But I tried the experiment with CrCl3.6H2O + PCl5 again. I took a large excess amount of PCl5, added some CrCl3.6H2O and carefully heated. The PCl5 sublimes, the CrCl3.6H2O starts bubbling (loosing water due to heating) and seems to melt (no real melting, just dissolving in own water of crystallization). The globules of "molten" chromium chloride are covered by a purple layer, but inside they remain green. When they are added to a large excess of water, then they quickly dissolve, but the purple parts dissolve much more slowly, albeit that in a minute or so, these also dissolve.

If you look at this method, then it seems possible to make anhydrous CrCl3, but isolating this from unreacted hydrated CrCl3 and PCl5 is not easy at all. Too strong heating makes a total mess of the mix of chemicals. All PCl5 sublimes, the chromium trichloride looses water and HCl and a mix of purple anhydrous CrCl3 and dark green, almost black chromium oxochloride remains. This mix does not look nice at all.

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

I did a similar experiment with hydrated chromium(III)formiate and PCl5. Here, no reaction occurs at all. All PCl5 simply sublimes and the dry chromium(III)formiate remains behind.




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[*] posted on 12-4-2009 at 21:14


I refluxed thionyl chloride with chromium (III) chloride hexahydrate for about 2 days and obtained a purple solid. The mixture does not start to turn purple for several hours. The product is still soluble in water due to CrCl2 being present. Dissolving chromium metal in concentrated hydrochloric acid is the best way to go to make the hydrated salt. After concentrating the solution and allowing it to sit in a large beaker, crystals started to form. It is best to let the solution slowly evaporate at room after concentration. Metal salts often form highly concentrated solutions from which they are unwilling to crystallize from. I have prepared titanium (III) chloride hexahydrate by dissolving titanium metal in concentrated hydrochloric acid, evaporating and salting the product out by passing anhydrous HCl gas into the solution until the solution fumed in air. upon cooling, purple crystals formed which where filtered off using a fritted funnel. At least chromium (III) chloride hexahydrate does not require these drastic measures to isolate it in a solid form.



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[*] posted on 13-4-2009 at 11:04


Wow, you refluxed with thionylchloride for 2 days :o ? That's a very long time. Now I understand why I did not see anything interesting happen. I stopped after half an hour or so. How do assure safety with such long-lasting refluxes? I never let any active system (reflux, distillation, heating bath, heating mantle, etc.) run unattended, so for me, such refluxes for such a long time are no option at all.

The CrCl3.6H2O I used is somewhat wet, it is made from Cr-metal and conc. HCl, which works quite well, at least if you allow the liquid with the metal pieces to stand for a long time (having such tame liquids standing and brewing for a long time is no problem for me). I usually do the drying on a heat radiator, with a paper tissue loosely covering the beaker, such that no dust gets in the crystals. In this way, I made a lot of metal salts, like NiSO4.6H2O, Co(NO3)2.xH2O, Cu(HCOO)2.xH2O, CrCl3.6H2O, PrCl3.xH2O, and many more.

The wetness of the CrCl3.6H2O quickly reacts with the thionyl chloride, but the water molecules which really are part of the hydrate do not react, at least not quickly. I also find it quite unusual that an inorganic chemistry reaction is going so slowly. Probably this is unique for chromium(III), because the water-ligands are attached to chromium(III) very firmly. Complexes of chromium(III) and cobalt(III) are rather inert and ligand-exchange or ligand-removal is very slow.




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[*] posted on 13-4-2009 at 21:13


The 2 day reflux was conducted at the university that I attend. The reaction was carried out in a fume hood and the water supply to the condenser was wired on. We routinely conduct reactions that run over night. I have not conducted an over night reflux in my home laboratory because I have not run water into my fume hood. For added safety, there are apparatus made by I2R that automatically shut the power off to a reaction if the temperature goes above a certain point or the water supply fails. I have aquired an I2R Thermowatch and a water sensor, these will be useful in the future. It is always advisible to heat reaction flasks with a heating mantle controlled with a variac transformer. No sparks are produced by these unlike normal laboratory hot plates. I love heating mantles, and magnetic stirring can be accomplished by placing a stirring hot plate below the heating mantle. For many inorganic and organometallic reactions, a hexammine comple of the metal is used because there is no coordinated water. An example is hexamine nickel (III) which I used in my synthesis of nickelocene described in another thread. Too bad hexammine chromium (III) chloride has to be prepared from anhydrous chromium (III) chloride in liquid ammonia as this compound has some interesting chemistry including this complex [Cr(NH3)6][CuCl5] that can be prepared by mixing Cr(NH3)6Cl3 with CuCl2 in HCl solution.



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[*] posted on 4-6-2009 at 14:32


Ok, I finally made some CrCl3, although it is not yet pure.
I first dissolved some chromium metal in concentrated hydrochloric acid, blew some air through it, and evaporated to dryness. The crystals were not that hygroscopic it seems.

A test-tube was placed almost horizontally, and 0,5mL of CCl4 was placed in the bottom of the test tube. Next in the middle CrCl3.6H2O crystals were placed. The crystals were heated with a small propane flame, and with the same burner I heated the CCl4 for short intervals such that it boiled slowly. The volume expanded, and I could see the purple color, but the conversion was far from complete. I decided change my plan somewhat. I heated the entire test tube more strongly (still pretty small flame, the reaction does't require high temperatures). Next I pippetted some CCl4 and brought the pipette as far in the test tube as possible and spewed it , such that it would reach as far to the bottom of the test tube as possible. The CCl4 instantly evaporated. After repeating a few times I collected my solid. It was a purple powder, but still contaminated with a small amount of green material. I attempted to dissolve theis with boiling 2M HCl, but it did not work. The CrCl3 did also not dissolve, wich is nice, as it shows no CrCl2 is present. Adding zinc dissolved the purple materials in less than a minute.

So I did the same procedure described above once more, and tried something else: putting the test tube vertically and heating the CrCl3 at the bottom, and next dropping CCl4 (I used up 1mL CCl4 in this step) on the solid.

After these treatments the solid was less contaminated, but still green residues remained. It is a waste of CCl4 to remove the last traces (I used up like 5mL today in trying many of these things). I think the solid is at least 90-95% CrCl3.
I will make some more batches (I don't have to glasware to work on a larger scale, as for larger scale you need a long tube, where I could lead in CCl4 vapour on one side and absorb the phosgene on the other side (ground glass joints).

When I have enough crude material, I will attempt to make the beautiful purple metallisch crystals, by subliming in a stream of Cl2.

Slowly heating the hydrated chloride alone gave a purple solid, but it decomposed on further heating, so it was probably [Cr(H2O)6]Cl3. The green hydrate is [Cr(H2O)4Cl2]Cl.2H2O .

Is it possible to extract the CrCl3 with a solvent? Maybe ether? CrCl3 is a pretty strong Lewis acid IIRC, and it should form an adduct with ether right? If yes, would it be possible to decompose the adduct by heating?

I am happy I will get 250mL CCl4 in a week, as I only have a small amount left (maybe 20mL). :)

[Edited on 4-6-2009 by Jor]
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[*] posted on 4-6-2009 at 15:09


Cr(III) forms 'inert' complexes, right? Once H2O is bound and the Cr reduced to Cr(III) then it is stuck. What you might need is a way of getting the chlorines on first before reduction (or oxidation) to Cr(III).

Your method of using Cr metal in conc HCl should work, another way might be to persist with your reaction of dichromate and chloride - you should be able to get the chlorines on this way before reduction to Cr(III). Would a dry reaction be advisable I wonder (K2Cr2O7/NaCl)?




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[*] posted on 4-6-2009 at 19:35


Anhydrous chromium (III) chloride forms a complex with tetrahydrofuran; CrCl3(THF)3. The CrCl3 can be extracted with THF in a Soxlet extractor and the complex obtained by evaporating the solvent. The complex is claimed to be very water sensitive. A procedure for the preparation of this complex can be found in Synthesis and Technique in Inorganic Chemistry, 2nd edition.



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