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Author: Subject: Antimony Pentachloride
Sauron
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[*] posted on 25-3-2009 at 01:53
Antimony Pentachloride


I have been thinking about further exploitation of this reagent beyond what we have already discussed.

How will it react with oxalic acid? Part of the hassle with the usual PCl5 method is that it is a solid/solid reaction. SbCl5 is a liquid. Ought to be better.

And with AcOH or NaOAc? Acetyl chloride or chloroacetic acid or chloroacetyl chloride?

Jf we bubble SO2 into SbCl5 might we obtain SOCl2?

These are not simple to preict. PCl55 fails to react with CS2 except under high temperature and pressure. SbCl5 reacts exothermically at room temperature. Go figure!

How about cyanuric acid? Will SbCl5 give TCT as PCl5 does?

The answers are likely all lurking in the lit. and I will ferret them out.

It is much cheaper and more practical to buy SbCl3 (solid) and prepare SbCl5 from it as needed than it is to try to store SbCl5, it is mega-corrosibe!

To whet the appetite here are a few pages deom Leo Paquette's book.

[Edited on 25-3-2009 by Sauron]

Attachment: SbCl5_Paquette.pdf (82kB)
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not_important
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[*] posted on 25-3-2009 at 03:19


Quote:

Antimony pentachloride is sometimes useful for chlorinating and is converted in the process into the trichloride. It is _ occasionally used in the presence of iodine for the \ exhaustive chlorination of aliphatic compounds of high molecular weight, such as palmitic acid.1 It is also capable of chlorinating aromatic compounds in the nucleus.

PREPARATION OF 3-4-DICHLORBENZOIC ACID.

Ten grammes of p-chlorbenzoic acid and 75 grm. antimony pentachloride are heated in a sealed tube to 200° for about eight hours. After cooling, the tube is opened and the contents treated with excess of dilute hydrochloric acid. The precipitated acid is collected, washed with cold water, and dissolved in dilute ammonia. After filtration the solution is evaporated to dryness and the ammonium salt decomposed with dilute hydrochloric acid. The acid is collected, washed, and dried in the usual way. It is recrystallised from dilute alcohol and forms colourless crystals melting at 201 °-202°.


The preparation of organic compounds By Edward de Barry Barnett

http://www.archive.org/details/preparationoforg00barnrich

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Nicodem
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[*] posted on 25-3-2009 at 03:26


SbCl5 (unlike PCl5) is a strong acid and therefore you can not use it to make oxalyl chloride from oxalic acid. It would only cause its decomposition to CO, CO2, H2O and perhaps a small amount of COCl2.
It also appears SbCl5 is similar to PCl5 in that it easily homolyticaly decomposes on heating, which could explain its use in catalysis of radical chlorinations (just like PCl5 is sometimes used). However, given the example not_important gave above, and the references in the SbCl5_Paquette.pdf attachment, it appears its redox potential is higher than PCl5 since I don't remember ever seeing PCl5 being used for oxidative chlorinations of so electron poor aromatic compounds.




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Sauron
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[*] posted on 25-3-2009 at 07:20


Well, in the case of oxalic acid there are only two reagents that succeed at all, PCl5 and TCT, and both produce the same crappy 50-52% yiels, I suspect that apart from the mechanical hassles of the solid phase PCk5 reaction, the chlorination proceeds stepwise

diacid-> half acid half acyl chloride -> diacyl chloride

And that the half acid half chloride undergoes a competing reaction of decomposition with a rate similar to the second chlorination. Hence 50%.

So yes perhaps SbCl5 might chew oxalic acid to gases.

But what about the others? AcOH, cyanuric acid, neither one likely to fall apart.

The instance of CS2 is interesting. Normally CS2 reacts rapidly with Cl2 even at room temperature so that Cl2 does not escape the liquid.. But is WP is dissolved in the CS2 chlorine reacts only with the P and not at all with the CS2, and PCl5 thus formed also does not react with CS2 under these conditions.

SbCl5 in contrast reacts in a few moments with CS2 at room temperature and if one does not take steps to moderate the reaction, the reaction mixture will be lost. Note that this is not the catalytic use of SbCl5 nor is UV employed. No heating is involved beyond the internal exotherm. Removing the evolved heat by cooling and using an excess of CS2 to moderate the reaction seem good ideas. Dropwise addn of the pentachloride also seems prudent. What is obtained is a slurry of CCl4, excess CS2, precipitated SbCl3 and sulfur along with SCl2/S2Cl2 as usual. More SbCl3 remains in solution.

The solids are filtered off and the liquids refluxed, CS2 reacts with SCl2 to give more CCl4 and sulfur. Finally, fractionation.

Anyway I bet there are more jewels like these in the older lit.




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watson.fawkes
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[*] posted on 25-3-2009 at 12:39


Might SbCl<sub>3</sub> act as a catalyst for chlorination under reflux by forming a reactive SbCl<sub>5</sub> species, or even SbCl<sub>4</sub><sup>-</sup>? If so, I've read here of plenty of reactions to speed up.
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[*] posted on 25-3-2009 at 14:19


No, SbCl3 is what is left behind after SBCl5 reacts and the three chlorines of the trichloride are much less reactive than the fourth and fifth chlorines of the pentachloride.

I do not know what basis you have for soeculating about a SbCl4 species, antimon comes in valences III and V never IV.

The exclusivity of reaction of P in CS2 with Cl2 is obviously a matter of kinetics, the formation of PCl5 and PCl5 being several orders of magnitude faster rate constant than formation of CSCl2, CSCl4, CCl4, SCl2 etc.




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[*] posted on 25-3-2009 at 17:03


@Sauron: watson suggested SbCl4(-) (you have to look closely to see the charge) which is an antimony (V) species.

@watson: the use of SbCl3 as a catalytic chlorinating agent has be discussed by Sauron in an earlier thread. The chlorine forms the SbCl5 in situ which is the actual chlorinating agent
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[*] posted on 26-3-2009 at 00:12


If SbCl5 reacts with CS2 and PCl5 does not, then certainly SbCl5 truly has a higher redox potential than PCl5. This is one thing more that makes it unsuitable for what you propose, the formation of acid chlorides from acids. But what makes it even less suitable is its acidity. Acid chlorides decompose in the presence of Lewis acids. They react with the Lewis acid in an equilibrium where R-CO<sup>+</sup> is one of the components and this carbocation fragmentates to R<sup>+</sup> and carbon monoxide. If things were so easy you could also use AlCl3 to form acid chlorides, but acyl chlorides react with AlCl3 to give alkyl chlorides, carbon monoxide and side products.

The unreactivity of the solvent in the chlorination of white phosphorous in CS2 is not necessarily due to kinetics. It is more likely that phosphorous reduces any of the first stage chlorination product of CS2 before this even rearranges or reacts further (for example, 3 S=C=SCl2 + 2 P => 3 CS2 + 2 PCl3). Otherwise you would certainly obtain a mixture of products with CS2 chlorination products present, but white phosphorous in solution is a strong reducent.

Quote: Originally posted by DJF90  
@Sauron: watson suggested SbCl4(-) (you have to look closely to see the charge) which is an antimony (V) species.

Well, actually SbCl4(-) is an antimony (III) species. :P




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[*] posted on 26-3-2009 at 00:45


Quote: Originally posted by DJF90  
@Sauron: watson suggested SbCl4(-) (you have to look closely to see the charge) which is an antimony (V) species.
Yes, the superscript typography is lacking. SbCl<sub>4</sub><sup>(-)</sup> certainly reads better on a browser, although still inadequately.
Quote:
@watson: the use of SbCl3 as a catalytic chlorinating agent has be discussed by Sauron in an earlier thread. The chlorine forms the SbCl5 in situ which is the actual chlorinating agent
I had only recalled its use as a reagent.

Incidentally, the thought arose by analogy with AlCl<sub>3</sub> in Friedel-Crafts reactions, which in that case acts as a chlorine acceptor, through a set of analogies in my head that I can't particularly reconstruct, except that I noted that the free energies would be of opposite sign for chlorine donors and acceptors. After some thought, it seems the critical reason that this works is that Sb can go from its (V) state to its (III) state, accepting a pair of electrons simultaneously. In the case of a C-H bond, it's accepting one from the (hypothetical) C-<sup>(1-)</sup> carbocation and another from a (hypothetical) short-lived intermediary Cl<sup>(1-)</sup> ion. And it does all this while donating chlorine simultaneously. No wonder it's aggressive.

With this in mind, it's possible that titanium or vanadium chlorides might also catalyze such reactions, since they exist in di-, tri-, and tetra- versions.
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[*] posted on 26-3-2009 at 03:39


And indeed Ti and V halides are useful at least as catalysts, although I am unaware of their being so reactive as SbCl5 in non-catalytic (stoichiometric) reactions.

Let's quickly review some known reactions

SbCl5 catalyzed vapor phase chlorinations:

CHCl3 -> CCl4
CH2=CH2 -> ClCH2-CH2Cl
CBr2Cl2 ->Cl3CBr
Cl3CBr -> CCl4
Acetylene -> tetrachloroethylene

By analogy CH2Cl2 -> CHCl3

Liquid phase non-catalytic or catalytic

CS2 -> CCl4

This is useful and versatile

Nicodem, I accept that oxalic acid is likely yo fall apart in SbCl5, but the de Barry Barnett reference above statesw that this reagent has been used for exhaustive chlorination of higher MW aliphatic carboxylic acid, e.g. palmitic acid. Therefore it seems that the -COOH group can survive. My question re AcOH was will the product be the alpha chloroacid, the acyl chloride or the alpha chloroacyl chloride? I have no particular preference for one over another as all would be useful. The following is entirely hypothetical AFAIK

AcOH + SbCl5 -> ClCH2-C(=O)-OH +SbCl3 + HCl

I2 catalyzed, suggestive of ICl or a mixed halogen Sb (V) comound like SbCl3I2 maybe. If it exists. I am familiar with the SbCl3F2.

By manipulating the amount of reactants this might be driven to dichloroacetic or trichloroacetic acid and all without UV or Cl2 gas. Admittedly more expensive but the SbCl3 is not lost and can be rechlorinated to SbCl5.


[Edited on 26-3-2009 by Sauron]




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[*] posted on 26-3-2009 at 06:05


Quote: Originally posted by Sauron  

Nicodem, I accept that oxalic acid is likely yo fall apart in SbCl5, but the de Barry Barnett reference above statesw that this reagent has been used for exhaustive chlorination of higher MW aliphatic carboxylic acid, e.g. palmitic acid. Therefore it seems that the -COOH group can survive. My question re AcOH was will the product be the alpha chloroacid, the acyl chloride or the alpha chloroacyl chloride? I have no particular preference for one over another as all would be useful.

I did not express myself properly. I should say that I do not know what the product of SbCl5 + RCOOH at any condition is, but I do know that it can not be RCOCl for the reasons I explained in my previous post (that is, as far as I know, strong Lewis acids do not form acid chlorides with carboxylic acids and if they would, the acid chlorides would decompose in their presence).
Alpha chlorination seems more than a reasonable outcome of such a reaction given the high redox potential and the capability of Lewis acids to enolize the carboxy group (SbCl5 is both an acid and an oxidant). However, the reference about palmitic (Berichte, 24, 1011-1026) is unavailable to me. Besides, "exhaustive chlorination" to me sounds like the COOH group is gone. Maybe you should check/request that paper, it could have some interesting information for this topic.




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Sauron
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[*] posted on 26-3-2009 at 10:12


That band of Ber. ought to be in Gallica, I will have a look.

My best guess is that exhaustive chlorination replaces all H with Cl, there are two standard reagents, SbCl5/I2 is one, the other is SO2/Cl2-AlCl3/S2Cl2 the so called BMC reagent. This is an old tool in analytical chemistry for elucidating structure in the pre NMR days gone by. In contemporary chemistry the same reagents are more often used to prepare perchloro derivatives of aromatics for use as analytical standards (GC ets.)

[Edited on 26-3-2009 by Sauron]




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[*] posted on 27-3-2009 at 19:08


Antimony trichloride according to Brauer reacts slowly with water to form SbOCl crystals, and 2 mols HCl.

The stoichiometry is SbCl3 + H20 -> SbOCl + 2 HCl but the procedure calls for 100 g SbCl3 and 70 ml water to be mixed at room temperature and left standing 2 days. It states that water hydrolyzes SbOCl to Sb2O3 but one must suppose this is even slower at RT since 70 ml is >11x excess water for a 1/3 molar reaction.

With anhydrous ethanol at 150 C in a sealed tube, same compound is formed in larger crystals.

The text is shtun as to other products but I suppose

EtOH + SbCl3 -> EtCl + SbOCl + HCl

SbCl3 is a much weaker Lewis acid than SbCl5

So why not

AcOH + SbCl3 -> AcCl + SbOCl + HCl ??

[Edited on 28-3-2009 by Sauron]




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