Sciencemadness Discussion Board

Acetic anhydride preparation

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madscientist - 20-5-2002 at 18:03

I'm theorizing that concentrated acetic acid could be dehydrated to acetic anhydride by mixing with concentrated sulfuric acid; then heating, and condensing the vapors, yielding acetic anhydride. Any comments or additional ideas?

Polverone - 20-5-2002 at 18:04

If the production were this easy, I think it would already be documented somewhere. But maybe not. Have you looked at Rhodium's materials on this topic? All of their methods require chlorides of phosphorus or sulfur, not so nice to work with and certainly not available as consumer products.

Polverone - 20-5-2002 at 18:04

It turns out that there's another synthesis route to acetic anhydride that requires more equipment but uses more accessible chemicals. Acetone is pyrolyzed with a catalyst in an electrically heated glass tube and the product is dissolved in glacial acetic acid to produce acetic anhydride. Search on the Hive for more info.

madscientist - 20-5-2002 at 18:04

Thanks for reminding me of that... I believe that is the ketene process.

Coen - 20-5-2002 at 18:05

I am sorry to dissapoint you, but it is not possible to dehydrate acetic acid to Ac2O.

Some time ago I was also working on the synthesis of this compound, but when I got from a friend about 1,5 L Ac2O, I stopped with that .

I strongly discourage you to do something with the route using ketene. Ketene is very dangerous. The setup is very difficult to realise. The procedure gives very low yields.
The whole process is only used in industry, and you can not always transform such syntheses to lab-scale.

I've did thought out another route using acetaldehyde which could be realised fairly easy. But you do have to make/buy acetaldehyde first for that.

Acetaldehyde can be made by leading acetylene through a solution of a Hg2+ salt and then condensing it.

Polverone - 20-5-2002 at 18:06

Not that I've actually done it, but I have good reason to believe that the ketene process CAN be performed on a lab scale. There was a thread on the Hive that included photographs of a home setup that had actually been used as part of drug manufacture in the early 1980s (now beyond statute of limitations). The person who used it was using thorium oxide as a catalyst, per instructions found in (I believe) Organic Syntheses (may have been a journal article; can't recall.) He claimed (IIRC) to be able to produce about a liter of acetic anhydride per day. There was an extended discussion on catalyst preparation and alternatives to the exotic and radioactive thorium catalyst. Certainly, I believe it may be hazardous, but not so hazardous as to disregard entirely. I believe the greater hazard is with potential legal trouble.

Coen - 20-5-2002 at 18:07

No, but only if there are no other alternatives. And there ARE in this case.
I will give some more details on the route via acetaldehyde some other time.

But if you're using it as a nitration catalyst there are several other possibilities to use, but also not very easy to get propably..

Further you must also not underestimate the power of simple mixed acid nitration. The concentration of the nitronium ion is such is mixtures is actually quite large.

Other acid-catalyzed nitrations can be done using mixtures of HNO3 with HF, HClO4, BF3, TFA, TFAA etc.

And of course nitronium salts are also a interesting group of nitration chemicals! NO2BF4, NO2SF6, NO2PF6.

And probably the most powerful nitration mixture is Magic Acid. This is a mixture of FSO3H, SbF5 and HNO3.

madscientist - 20-5-2002 at 18:08

es, I remember a thread that I started at The Forum a while back, "Holy Grail Oxidizer"? That included instructions on how to prepare [NF4+][ClO4-]. This was promptly dismissed by nbk2000, who then gave me a lesson on the basic fundamentals of chemistry. *sarcastic tone of voice* "Chemistry isn't just 1+1. Quit talking out of your ass". Perhaps sometime soon I shall repost that, here.

I believe it is rather easy for the home chemist to prepare CH3CHO. I haven't had the chance to test to see if it had actually worked, but recently I attempted to prepare CH3CHO from CH3CH2OH via the dehydrogenation route. Alcohols can be dehydrogenated by passing the hot vapor over hot copper metal, in the absence of reactive gasses such as oxygen. This would be the reaction for the formation of the CH3CHO:

(hot CH3CH2OH vapor passed over hot copper metal in oxygen-free environment)
CH3CH2OH --} CH3CHO + H2

If a small amount of oxygen gas contaminent is present, the results will not be disasterous, because the oxygen will first react with the hydrogen produced.

How I set up the experiment:
I took a 250mL flask, and filled it with 50mL of CH3CH2OH. I placed a rubber stopper in it, and connected a thin copper metal tube to the opening in the stopper (tight fit). I set the flask on a metal screen on top of a propane burner. I then bent the copper tube around so that it would come in close contact with the flames emmitted from the burner. The tube then continued on, bent down so that condensed vapors would drip into another flask (did not use a stopper at that end because then pressure might build inside of the two flasks and metal tubing, causing a stopper to shoot off). I of course had the metal tube, at the condensation end, packed with ice/salt filled baggies. I condensed a fair amount of liquid (don't remember just how much).

Coen - 20-5-2002 at 18:09

Hmm. This is another one of those industrial processes which have been transformed to home experiments.

But I must say that I find this rather stupid: heating with burners when working with the highly flammable and quickly vaporizing acetaldehyde. Even if you try to cover things up a little you can get quite dangerous situations. I have had some nasty experiences with a comparable experiment in the past so that's explains the nature of this reaction of me.

And I also doubt on the effectiveness of this experiment. How long was your copper tube? You'll need quite a long contact area.

BTW, I heard it is also possible to partial reduce ethanol to ethanal with K2Cr2O7. I find this also rather strange, since most simple alcohols get immediately reduced to the acid when you try to do this.

madscientist - 20-5-2002 at 18:09

I haven't worried about my experiment, considering that what acetaldehyde I have is either collected in a flask which is in an ice bath; or is being condensed in an environment free of oxygen. I have also not worried because I have been doing this outside (I live in a windy area), and because I am nowhere near the experiment while ethanol is still present in the flask that is being heated. I haven't even been able to ignite the vapors of boiling isopropanone that are jetting out of the neck of a flask (in the windy conditions outside). Please elaborate on your accident. :-)

The part of the copper tube that is being heated is about eight inches long. The length of the entire tube is a number that I am uncertain of, considering its now-curved nature. The acetaldehyde is immediately condensed after it passes through the heated part of the copper tube. The copper tube is around 2-3mm in diameter. I put eight kinks in the part of the copper tube that is being heated; at these kinks the width of the opening is about 1mm. I figured this would cause a lot of turbulence as well as slowing down the passage of the ethanol vapors, causing far more of it to be converted to acetaldehyde.

Even if this experiment failed to convert all of the ethanol to acetaldehyde, I could just run the ethanol / acetaldehyde mix through the experiment again, to dehydrogenate more of the ethanol into acetaldehyde.

As for oxidizing alcohols to aldehydes / ketones, I think it would be worth considering using hypochlorites to reduce an alcohol to its respective ketone / aldehyde. The problem with using permanganates or other oxidizing agents (this is just a theory) is probably that the permangante ion will react with just one alcohol molecule, reducing it fully, instead of reacting with multiple alcohols, reducing them partially, down to their respective ketone / aldehyde. A hypochlorite, however, would not have that problem.

CH3CH2OH + NaOCl --} CH3CHO + NaCl + H2O

The acetaldehyde could then be distilled. Water could be removed by adding the correct amount of MgSO4 (not the hydrate, MgSO4*7H2O!), which over an hour or two will absorb the water. Again, that mixture would then be distilled (but actively monitored to avoid heating the MgSO4*xH2O formed which would cause the water to be released again, defeating the purpose of that process).

Coen - 20-5-2002 at 18:10

Although I agree with you that it's quite certain there is atleast some CH3CHO in your product. But I think you have to realise more that often A LOT of side reactions are also possible that usually aren't listed in the standard literature. This story also goes for example on your theory of those carboxyl amines.

You would be maybe amazed how uncertain professional chemists often are about products obtained via (suspected) quite standard reaction mechanisms. Practically *every* compound that is synthesised is run through IR, NMR, and HPLC.

Our little 'research group' also walks to this problem. We only have IR and TLC capabilities and althoug that's already better then most hobbyists have, we really need for example at least a HPLC and a DSC.

For the possibility of using hypochlorites for that oxidation I also have my doubts. Hypochlorites are used in some other oxidation reactions, yes. But I'm really doubting what that OH group is going to do when ClO- approaches, several other possibilies also possible. And ClO- isn't a very powerful oxidator also. But I gues it would be quite easy to try.

But it's already proven that you can also do it using dichromates you why even bother further. Well, only if you can't get dichromates maybe.

BTW, how did the obtained product smelled?

PHILOU Zrealone - 20-5-2002 at 18:12

Hello Pyroweb euh Requiem II! Aren't you related to that guy on the weapon and explosive forum tsv or tvs + number?
I also think to remember you are part of that famous research group ;-) High energetic material & explosives.When I saw pichon, I thought it might be you, but when I saw that someone has provided you 1,5 L Ac2O (since it was me, there is no doubt :-)).
Why 3 different names?

First of, yes CH3-CH=O is easy to do from acidic KMnO4/ K2Cr2O7 by dropping (in a closed vessel with a cold trap) drop by drop ethanol on the heated previously mentionned oxydiser/acid solution; of course don't do this with 96-98% H2SO4 and dry oxydiser, otherwise flash kabooom from Cr2O6 or Mn2O7! No here we are talking about dilluted solutions of acid and mediumly concentrated solution of oxydiser.
True that primary alcools are oxydised to aldehyde first and then to carboxy acid. Here is the reason why this has to be performed at high T (actually 10°C over the boiling point of the aldehyde and lower than the bp of the alcool).
For the low molecular aldehydes, it is often the case that bp is lower than the parent alcool and always than the parent acid; over its boiling point, most of the aldehyde is then volatilised and gets out of reach of oxydiser!
A tiny portion of it remains and react further reason why you mustn't be too concentrate in oxydiser (would you trust KMnO4/acetic acid- it has been involved in many lab accidents). The cold trap is there to collect the volatile aldehyde! This is one of the safest way to get pure aldehydes!

Now ketene process is absolutely hard and waytoo dangerous to do in a lab.

There are many ways to make Ac2O, but all of them involve toxic halides!
The general idea is to get CH3-CO-Cl; now use your brains how to get it from acid halides!
Most of acid halides are easily done, but I'm a bit tired today, maybe another day!
(Tips: S2Cl2, SCl2, SOCl2, SO2Cl2, COCl2, PCl3, PCl5, SCl4, SCl6, POCl3, ....)

Also to madscientist:
CH3-CH2OH + NaOCl --) no acetaldehyde, but a complex mixture of crotonisation of CH3-CH=O (see aldehydes in basic media) into CH3-CH=CH-CH=O and the like; and... chloroform; characteristic reaction of CH3-CO- groups (aceton, acetaldehyde, acetic acid, acetophenone, ...see haloform reaction in basic media); resulting in the splitting of the CH3 and the formation of an acid.
CH3-CH=O + NaOCl -OH(-)-) CCl3(-) + HO-CH=O
CCl3(-) + H2O --) CHCl3 + OH(-)

P.S.: NaOCl is always basic media!


madscientist - 26-5-2002 at 21:18

I'm not sure, but I think that it's possible that I prepared acetic anhydride by accident today. I mixed 164g CH3COONa (contaminated with a small amount of NaHCO3, around 1-2g) with 75g 94% H2SO4. I then poured that into a flask; began heating it, and condensing the vapors (typical distillation). I noticed a very strange, sickly-sweet odor; very difficult to describe. I got a whiff of a very small amount of it, causing me to choke for a few moments. The condensed liquid (which I got 22mL of, if I remember correctly) was still liquid, showing no signs of imminent freezing, at -10C. Now, if that had been acetic acid, it would have frozen at a far higher temperature than that... acetic anhydride, on the other hand, would not freeze until the temperature was FAR lower. I'm postulating that the following reactions were occuring, forming at least a fair-sized quantity of acetic anhydride (there is probably a significant amount of acetic acid remaining). Keep in mind that there was a slight stoichemical excess of sulfuric acid.

2CH3COONa + H2SO4 --> 2CH3COOH + Na2SO4
7CH3COOH + H2SO4 --> 7CH2CO + H2SO4*7H2O

And of course, the following occurs:


Tomorrow I'll try droppering a small amount of the distilled liquid onto an aluminum plate; if there is no visible reaction, then it is definitely high-purity acetic anhydride. Otherwise, it contains at least a medium amount of acetic acid.

madscientist - 27-5-2002 at 08:26

I placed a few drops of the distilled liquid on a piece of aluminum foil. No visible bubbling, or audible bubbling, resulted. It has a pH of 1, though.


Polverone - 27-5-2002 at 12:15

I wouldn't be so quick to jump to the conclusion that this is acetic anhydride. I do not observe a rapid reaction with my own glacial acetic acid and aluminum foil. Many favorable reactions of aluminum can be difficult to start because of the oxide coating. Try mixing some of the liquid with water and stirring. Does it form two phases? I think acetic anhydride should be immiscible with water; it will of course be hydrating, but the reaction is slow enough that you should have time to identify a separate phase if you really have the AA.

vulture - 8-7-2002 at 01:23

Aluminium foil usually has a resistant oxide layer, better try zinc powder.

vulture - 9-8-2002 at 08:21

Okay, maybe a brute force way to produce anhydride but it's maybe worth a try.

How about mixing glacial acetic acid with sulfuric acid and magnesium sulfate (anhydrous!). Magnesium sulfate is a powerful dehydration agent.

Here's a hypothetical reaction:

2CH3COOH + H2SO4 + MgSO4
-> CH3(CO)O(CO)CH3 + H2SO4 + MgSO4.H2O

This should be done in a distilling setup and the temperature shouldn't go above 60C to prevent the MgSO4.H2O from dehydrating again.

I doubt if it will work, but who knows?


PrimoPyro - 9-8-2002 at 08:26

Ac2O is more hydrophilic than salt hydrates, so the reverse reaciton would predominate. Has anyone ruled out P2O5 dehydration? I dont know too many things that cant be dehydrated by P2O5....

CH3COOH can be dehydrated to acetic anhydride, with ketene, CH2=C=O.

Ketene reacts with acetic acid to dehydrate it straight to Ac2O. Ketene is quite noxious though.

If someone wants to try aluminum, try an aluminum amalgam with HgCl2. This is very commonly used to strip the oxide coating off the aluminum, usually to use the aluminum metal as a gentle reducing agent.


vulture - 9-8-2002 at 11:53

Oh well. I think P2O5 should work, I think I read that somewhere. Too bad it is also a controlled substance.

Hang on a sec. See the thread for preparation of elemental phosphorus. I mentioned something about calciumphosphide. If this is burnt, it will produce P2O5. Now I just have to find a way to collect it.

vulture - 9-8-2002 at 11:54

I can't edit...

Well, if it is more hygroscopic than salt hydrates, that's why I added the H2SO4, this might fix the water. Just trying to combine forces, he.....;)


Polverone - 9-8-2002 at 11:54

Nobody has ruled out dehydration with P2O5, but you should keep in mind that P2O5 isn't exactly growing on trees around here. Hmm, that reminds me: I've always heard P2O5 referred to as a "powerful" dehydrating agent. Is there anything *more* powerful than P2O5?


PrimoPyro - 9-8-2002 at 12:28

Extremely waterscavenging species do exist, but they are never used for this purpose.

Take alkyllithiums for example, or grignard reagents. They are hard and cumbersome to prepare, and are very useful indeed, and they are extremely reactive toward water and similar loose protons, but no one would ever make these for this purpose.

P2O5 can be bought from chem supply, but it isn't very cheap. I think P2O5 is the strongest "common" dehydrating agent.


magnesium perchlorate for dehydration?

Hoffmann-LaRoche - 10-9-2002 at 05:19

what about using magnesium perchlorate?
its used in place of P2O5( but much easier to prepare...

(ref.: Gmelin, Syst.-Nr. 27, Mg, Tl. B, 1937, S. 154–158 ï Hager 5, 643 ï Hommel Nr. 284.)

vulture - 23-9-2002 at 02:23

Good idea. Only drawback is that perchlorate ions are maybe considered too precious for this purpose?

Also, I've been thinking. Would it be possible to hook up two ethanol molecules by the O by splitting H2O and then oxidizing it to anhydride? Hmm just realize ethers are too unreactive for that...
Mkay, how about hooking up acetaldehyde?

Hoffmann-LaRoche - 23-9-2002 at 10:07

Anodic oxidation of MgCl2 at low current density.
Then carefully heating the resulting
Mg(ClO4)2*6H2O to give Mg(ClO4)2.
Just dont know wether magnesium perchlorate would decompose, or even explode...

Another the acetaldehyde pathway for Ac2O, how do they manage in industry to get Ac2O and not CH3COOH when oxidizing acetaldehyde???


vulture - 23-9-2002 at 11:08

I can tell from personal experience that converting KCl to KClO4 SUCKS VERY BADLY. The yield is shit and it takes ages. Considering that MgCl2 needs double oxidation, it would take forever.:mad:

Btw, the acetaldehyde pathway is used in the industry by catalytical oxidation of acetaldehyde.. odd, cause this should just yield acetic acid

Hoffmann-LaRoche - 23-9-2002 at 11:52

Maybe you didnt have a good procedure?-I will look up one.....

NaClO3 is cheap.

1)4NaClO3>NaCl + 3NaClO4

2)NaClO4 + 1/2K2CO3>KClO4(insoluble in cold water) + 1/2Na2CO3

How is that?:)

Just dont know how we come from ClO4- to Mg(ClO4)2......mixing with a magnesium salt would yield some Mg(ClO4)2, sure...but how to seperate?(magnesium perchlorate is easily soluble in water, and the NaClO4 needed would yield a soluble salt at the equilibrium)



vulture - 24-9-2002 at 10:26

1)4NaClO3>NaCl + 3NaClO4

This equation is incorrect. When heated to its melting point, chlorates will decompose according to the followin equation:
2NaClO3 -> NaCl + NaClO4 + O2

The trouble is that one can't do this with metal containers since they will catalyticly decompose the chlorate into mere chloride and oxygen.

I've tried several methods of electrolysis and the processes are very unsactisfactory due to certain restrictions to the home chemist. Low voltage and high current requires special power supplies and evolves alot of poisonous chlorine gas.

rikkitikkitavi - 24-9-2002 at 10:58

when making chlorates i can only recomend wouters webpage. i dont have the link here but it is the best web info about making them.

anyway, i seriously doubt that mg(clo4)2 is a strong enough dehydratant. p2o5 relies on chemical reaction to pull water.
mg(clo4)2 doesent.


andreas - 25-9-2002 at 07:09
Here's the link:)

Interesting process

vulture - 25-9-2002 at 08:27

I found out a rather interesting process for acetic anhydride today. It consists of heating a mixture of sodiumacetate and chloroacetic acid:

CH3COONa + CH2ClCOOH -> (CH3CO)2O + NaCl

Hmm somethings wrong with this equation...

madscientist - 25-9-2002 at 08:51

I think that reaction chloroacetic acid with sodium acetate would yield CH3COOCH2COOH, not acetic anhydride. Obviously acetic anhydride can be prepared from sodium acetate and acetyl chloride (CH3COCl).

vulture - 25-9-2002 at 10:20

Obviously acetic anhydride can be prepared from sodium acetate and acetyl chloride

That was what I meant, what a brain fart :(

How hard to prepare is acetyl chloride?

Very Hard

PrimoPyro - 25-9-2002 at 18:29

Without PCl3 or PCl5, very hard.

Next question...

vulture - 26-9-2002 at 12:49

..How hard to prepare is PCl3 or PCl5?:D
Considering phosphates and maybe phosphides as a starting material..

Damn, does acetic acid like phosphorus compounds or what? :D

Acetaldehyde with K2Cr2O7

chemoleo - 27-7-2003 at 15:35

There are definitley easy ways to do this (i.e. in the lab at home). the point is to control the reaction carefully (in terms of time and temperature) to avoid generation of acetic acid. Did it once myself, but need to look up details again if anyone is interested. Also, acetaldehyde is rather interesting for other syntheses, like pentaerythritol :)

maybe hard, but not that hard

BASF - 20-8-2003 at 08:18


Very Hard

Without PCl3 or PCl5, very hard.

In the middle of the text is a procedure starting with S2Cl2 and Na-acetate.
The method IS extensive and involves several steps, but S2Cl2 can be made quite easily...

(IF there is a lot of interest, i can translate it, at the moment the formulas should suffice.)

[Edited on 22-8-2003 by BASF]

BASF - 22-8-2003 at 07:01

BTW, what about zeolithes("molecular sieves";) for dehydrating HAc to Ac2O?

Molecular sieves are one of the most powerful dehydrating agents used...

BASF - 22-8-2003 at 07:06

Anorganisch-chemisches Praktikum II

Darstellung von Schwefeldichlorid SCl2 aus Schwefelmonochlorid S2Cl2

Theoretischer Teil

Das Schwefeldichlorid ist eine dunkelrote, toxische, an der Luft rauchende Flüssigkeit mit einem Schmelzpunkt von –122°C und einem Siedepunkt von +59-60°C. Die Substanz besitzt einen chlorähnlichem Geruch und ist sehr hydrolyseempfindlich. Mit einer Bildungsenthalpie von DHB = -49.4 kJ/mol ist SCl2 nur eine mässig exotherme Verbindung. Es ist bei Raumtemperatur nicht stabil, da es nach folgendem Gleichgewicht (Gleichung 1) in Schwefelmonochlorid und Chlor zerfällt.

So kann auch der Siedepunkt nicht ohne Zersetzung erreicht werden, wenn nicht das Gleich-gewicht durch Stabilisatoren auf der rechten Seite fixiert wird. Als solche dienen PCl3 oder PCl5.

Die Darstellung erfolgt gemäß obiger Gleichung (Gleichung 1). Die Reaktion läuft bei Raum-temperatur nur langsam, kann aber durch Zugabe von Katalysatoren wie Eisendi- oder trichlorid, elementares Eisen oder Iod erheblich beschleunigt werden.

Das SCl2 besitzt C2v-Symmetrie. Das Schwefelatom ist von den Chloratomen und den zwei freien Elektronenpaaren tetraedrisch umgeben. Der ClSCl-Bindungswinkel beträgt etwa 103°, der SCl-Bindungsabstand wird mit 2.01Å angegeben.

Für das Schwefeldichlorid finden sich in etwa die gleichen Anwendungsmöglichkeiten wie für die als Edukt verwendete Substanz, das Schwefelmonochlorid. Letztere Verbindung ist um einiges stabiler als das Dichlorid und wird in der Technik bei Vulkanisationsprozessen verwendet, da sich darin der Schwefel hervorragend unter Bildung von Polyschwefelchloriden löst. Weiterhin wird es für die Härtung von Öl- und Lackstrichen verwendet, sowie als Sulfidierungs- und Chlorierungsmittel in der Synthese.

Das Schwefelmonochlorid, S2Cl2, ist eine orangegelbe, toxische, hydrolyseempfindliche Flüssig-keit von widerlichem, erstickendem Geruch, deren Siedepunkt bei 137°C liegt. Die Verbindung wird dargestellt durch Überleiten eines trockenen Chlorstroms über geschmolzenen und wieder erstarrten Schwefel bei Raumtemperatur oder mittels Durchleiten eines trockenen Chlorstroms durch geschmolzenen Schwefel bei 240°C.

Das Schwefelmonochlorid liegt nicht – wie der Name es vielleicht impliziert – als SCl vor, sondern existiert als Dimer. Es besitzt eine kettenförmige, gewinkelte, nichtebene Struktur (Abbildung 2) und gehört der Punktgruppe C2 an. Im Unterschied zum S2F2 liegt das S2Cl2 unter Normalbedingungen nur in dieser gauche-Struktur vor, nicht jedoch als das Isomer mit einer S-S-Doppelbindung (S=SCl2). Letzteres kann jedoch mit Hilfe der Blitzlichphotolyse aus einer Tieftemperaturmatrix isoliert werden.

Die Bildungsenthalpie für das S2Cl2 ist mit DHB = -58.2 kJ/mol nur wenig mehr exotherm als jene für das SCl2, dennoch ist das S2Cl2 bei Raumtemperatur um einiges stabiler als das Dichlorid.

Für die Darstellung dieser Schwefelhalogenide wird eine Folge von a-Addition und a-Eliminierung vorgeschlagen (Schema 1).

S-S-Bindungen werden schon bei Raumtemperatur von Chlor merklich angegriffen, so dass die Darstellung der Chloride – wie es hier geschieht – bei eben dieser Temperatur erfolgen kann.

Sowohl von S2Cl2 als auch von SCl2 gibt es Verbindungen der leichteren (Fluor) wie auch der schwereren (Brom, Iod) Homologen des Chlors. Dabei sind S2F2, S2Br2 und SF2 (dimerisiert zum FSSF3) bei Raumtemperatur stabile Verbindungen, wohingegen das SBr2 instabil ist und nur als Intermediat bei der Darstellung von S2Br2 nachgewiesen werden kann. Auch S2I2 ist eine Substanz, die sich bereits unterhalb von –31°C zersetzt. Das SI2 lässt sich schließlich gar nicht mehr darstellen.

Darstellung von S2Cl2

Die Innenwand eines 500-mL-Dreihalskolbens wurde mit einer Schmelze von Schwefel (8.55 g, 266.7 mmol) überzogen. Man ließ die Schmelze erkalten. Dieser über Kopf stehende Kolben wurde mit einem zweiten Zweihalskolben über einen Doppelkern so verbunden, dass ent-stehendes Produkt in letzterem aufgefangen werden konnte. Durch diese Anordnung wurde ein trockener Chlorstrom geleitet. Der Auslass des Auffangkolbens wurde gegen eindringende Feuchtigkeit mittels einer Gaswaschflasche, die konz. Schwefelsäure enthielt, gesichert. Das überschüssige Chlorgas wurde zur Entsorgung in Natriumthiosulfat-Lösung eingeleitet. Nach einiger Zeit und gelindem Erwärmen des Schwefels setzte die Reaktion mit der Entwicklung einer orangegelben Flüssigkeit ein. Der Chlorstrom wurde solange duch die Apparatur geleitet, bis aller Schwefel aufgebraucht war. Das Rohprodukt wurde mit einer Spatelspitze Schwefel versetzt und unter Stickstoff destilliert. Es wurde eine orangegelbe Flüssigkeit aufgefangen, deren Siedepunkt bei 135-136°C lag. Ausbeute: 10.62 g (103 mmol, 39%).

IR (neat), n (cm-1) 426, 511, 540, 880, 990.

Darstellung von SCl2

S2Cl2 (10.62 g, 103 mmol) wurde in einem mit Zweihalsaufsatz versehenen Rundkolben, dessen eine Öffnung mit einem Gaseinleitungsrohr versehen war und dessen andere Öffnung gegen eindringende Feuchtigkeit an eine Gaswaschflasche mit konz. Schwefelsäure angeschlossen war, vorgelegt und mit einer Spatelspitze Eisenpulver versetzt. Durch die Flüssigkeit wurde unter Eis-kühlung für 45 Minuten ein trockener Chlorstrom geleitet. (Überschüssiges Chlor wurde durch Einleiten in Thiosulfat-Lsg. entsorgt. Während dieser Zeit färbte sich die Flüssigkeit dunkel. Nach Beendigung der Chlordurchleitung wurde für eine weitere Stunde gerührt. Anschliessend wurde die Flüssigkeit mit einigen Tropfen PCl3 versetzt und unter Stick-stoff destilliert. Es konnte eine rote Flüssigkeit aufgefangen werden, die bei einer Temperatur von 58°C überging. Ausbeute: 3.26 g (24.1 mmol, 23%).

IR (neat), n (cm-1) 436, 510.

Diskussion der Ergebnisse
Die dargestellten Substanzen lassen sich aufgrund von Farbe, Geruch und Siedepunkt als S2Cl2 und SCl2 identifizieren.

Auch die IR-Daten bestätigen die Identität der Verbindungen, wie im folgenden gezeigt wird.

S2Cl2 gehört, wie bereits erwähnt, der Symmetriegruppe C2 an, besitzt also nur eine zweizählige Drehachse. Für das vieratomige Molekül werden 3n-6, also 6 Fundamentalschwingungen erwartet, von denen vier der Schwingungsklasse A und zwei der Klasse B zugeordnet werden (Abbildung 3).

ns ns ds

nas das g

Abb. 3: Fundamentalschwingungen des S2Cl2

Die erwarteten und gemessenen Schwingung sind in Tabelle 1 aufgeführt.

Tabelle 1
Art der Schwingungen
Gemessene Werte

n in cm-1

n in cm-1

n1: ns(S-S)

n2: ns(S-Cl)
Unscharf, siehe unten

n3: ds(SSCl)
n. g.

n4: g(SCl)
n. g.

n5: nas(SCl)

n6: das(SSCl)
n. g.

n. g.: nicht gemessen

Das S2Cl2 wurde als Film zwischen zwei KBr-Plättchen gemessen. Beobachtet werden konnten eine Bande bei 540 cm-1 sowie eine im Bereich von 436 cm-1, wobei letztere nicht als scharfer Peak auftaucht, sondern verschwommen ist, da in ihr vermutlich die erwarteten Schwingungen bei 448 cm-1 und 434-436 cm-1 zusammenfallen. Jegliche Schwingung unterhalb von 400 cm-1 konnte nicht mehr gemessen werden, da KBr ab etwa 350 cm-1 undurchlässig für IR-Strahlung ist.

Neben den Fundamentalschwingungen lassen sich auch einige der erwarteten Ober- und Kombinationsschwingungen beobachten (Tabelle 2).

Tabelle 2
Gemessene Werte

n in cm-1

n in cm-1

n7 = 2n1

n8 = 3n2
n. a.

n9 = n1+n6

n10 = n2+n5

n11 = n1+n2

n12 = n1+n2+n8
n. a.

n. a.: nicht ausgewertet. Bei ungefähr den Literaturwerten tauchen zwar Banden

auf, denen aber aufgrund der geringen Intensität keine genauen Werte zu-

geordnet werden konnten.

Die im Bereich von 3448 cm-1 sowie um 1654 cm-1 erhaltenen Banden rühren vom Wasser her, dessen Erfassung aufgrund der Hygroskopizität von KBr nicht vermieden werden kann.

Das SCl2 besitzt C2V-Symmetrie, hat also eine zweizählige Achse sowie zwei Spiegelebenen. Erwartet werden drei Fundamentalschwingungen (Abbildung 4).

ns ds nas

Abb. 4: Fundamentalschwingungen von SCl2

Dabei gehören die symmetrische Valenz- und die symmetrische Deformationsschwingung zur Schwingungsklasse A1, die asymmetrische Valenzschwingung zur Klasse B2.

Erwartete und gemessene Werte sind in Tabelle 3 aufgeführt.

Tabelle 3
Art der Schwingung
Gemessene Werte

n in cm-1

n in cm-1

n1: ns

n2: ds
n. g.

n3: nas
n. b.

n. g.: nicht gemessen; n. b.: nicht beobachtet

Die zusätzlich auftretende Bande bei 436 cm-1 ist auf Verunreinigungen durch S2Cl2, die sich wegen des im Theorieteil erwähnten Gleichgewichts nicht vermeiden lassen, zurückzuführen.

Da auch hier die Messung von SCl2 als Film zwischen KBr-Plättchen durchgeführt wird, werden ebenfalls Wasserspuren detektiert.

Ober- und Kombinationsschwingungen (Tabelle 4) werden nur andeutungsweise im gekenntzeichneten Bereich (Siehe Spektrum.) beobachtet.

Tabelle 4


n in cm-1

n4 = 2n1
1027 und 1237

n5 = n1+n2

n6 = n1+n3

n7 = 2n1+2n2

Die trotz des vollständig umgesetzten Schwefels niedrige Ausbeute bei der Darstellung von S2Cl2 lässt sich vermutlich darauf zurückführen, dass ein Teil des Schwefels durch darübergelaufendes Produkt zu Polyschwefeldichloriden gelöst wurde, womit sich sowohl ein Teil des Produktes als auch weiterer Schwefel der Produktgewinnung entzog. Hinzu kommen Verluste bei der Destillation, z. T. auch hier dadurch verursacht, dass sich der für die Destillation zugegebene Schwefel in Form von Polyschwefelhalogeniden gelöst hat.

Für SCl2 liegt die Ausbeute um 48% niedriger als die angegebene Literaturausbeute von 70%. Hier hat sich wahrscheinlich ein erheblicher Anteil der Produktes beim Destillieren zersetzt, da nicht ausreichend PCl3 für die Stabilisierung zugegeben wurde. Es konnte neben dem Produkt auch einiges an Edukt aufgefangen werden.


[1] Gmelin, Handbuch der Anorganischen Chemie, Ergänzungsband 2: Schwefelhalogenide,

Springer Verlag, Berlin, Heidelberg, NY 1978, S.220ff.

[2] A. F. Holleman, E. Wiberg, Lehrbuch der Anorg. Chemie, 101. Auflage, W. de Gruyter,

Berlin, NY 1995, S. 561ff.


AA - 10-9-2003 at 23:42

Acetic anhydride can be made by the dehydration of acetic acid by PO5 but in response to mad scientists original question it can also be made by bubbling SO3 throught acetic acid. Or probably just adding sufficient fuming sulphuric (~30% SO3) to dehydrate and then stilling off the anhydride pushing the equalibrium towards the right.

BASF - 16-9-2003 at 06:25


Acetic anhydride can be made by the dehydration of acetic acid by PO5 but in response to mad scientists original question it can also be made by bubbling SO3 throught acetic acid.

Source is?-Just curious.

Acetic anhydride........ as easy as a surrey girl....... and as cheap as gold tassles at a shriners convention.

Hermes_Trismegistus - 4-3-2004 at 10:11

Sodium Acetate reacts with Carbon Dioxide in aqeous solution to produce acetic anhydride and sodium bicarbonate, under suitable conditions the sodium bicarbonate precipitates and can be removed by centrifugal seperation.

However, presumably the cold water solution can be extracted with an organic solvent such as chloroform or ethyl acetate, to furnish acetic anhydride.

However, this synth would need to be carried out with no small amount of haste.

The half-life of acetic anhydride in aqueous sol'n at 19 degrees celcius is published as being ~60 min.

Some other data suggests a half-life of approx six minutes at 20 degrees Celcius.

In either case, frequent and hasty extractions coupled with lower temperatures would be condusive to good yields.

Half life in moist chloroform may be as much as 120 minutes.

For those skeptical scientists among the assembly, I will be attempting this reaction withing the next couple weeks and I'll publish a full report, including references, at that time.

Organikum - 4-3-2004 at 12:33

At what temperature is the reaction between sodium acetate and CO2 thought to be carried out, or better what temperature does this reaction require to proceed in a reasonable speed?

And I guess the reaction has to be carried out in a running centrifuge anyways - for to get the AA to separate as a single blob at the bottom an so reducing the surface area exposed to the water. The surface area might explain the different times mentioned - if a blob sits at the bottom the AA takes lots of time to decompose - if the shit is mixed it will go much faster.

Awaiting the results of the experiment - interesting approach!

[Edited on 4-3-2004 by Organikum]

Organikum - 5-3-2004 at 02:29

Are you sure you dont confuse this with the methylacetate + CO to AA way of production? (rhodium-catalyst)


yes, more details to follow.

Hermes_Trismegistus - 5-3-2004 at 17:09

BASF - 25-3-2004 at 12:46

That´s really interesting...i have found patents on the MeOAc + CO2-method a year ago, but they weren´t quite detailed:

GB480953 Process for the conversion of alkali metal chlorides into carbonates

GB353318 Improvements in and relating to the manufacture and production of ethylidene diacetate

GB368835 Process of producing acetic anhydride from ethylidene diacetate

GB441956 Improvements in the manufacture of acetic anhydride

GB441956 Improvements in the manufacture of acetic anhydride

GB404333 Process for the treatment of the reaction mixtures formed in the preparation of acetic anhydride by the thermal decomposition of acetic acid

GB238825 Process of manufacture of acetic anhydride and aldehyde

GB643700 Purification of acetic anhydride

(notice the english origin(?)...strange, eh?-I did a worldwide search, in spite of this, the most promising ones were these old british patents...)


unionised - 15-4-2004 at 14:00

Hermes, you will probably have to be even faster than you thought. In the presence of a base like bicarbonate or acetate the half life will be even shorter.:(


Hjalmar_Poelzig - 28-4-2004 at 10:11

Nobody mentioned the patent which uses sodium pyrosulfate and sodium acetate? (can't remember it off-hand, have it around somewhere)

Basically it goes like this:

2 NaHSO4 --heat--> Na2S2O7 + H2O

Na2S2O7 + 2CH3COONa -----> (CH3CO)2O + 2Na2SO4

Sodium bisulfate is heated above 300°C, where it loses a molecule of water and becomes sodium pyrosulfate.

An intimate mixture of this and 2 molar equivalents anhydrous sodium acetate is wetted with some glacial acetic acid and refluxed for a couple of hours, then subjected to (fractional) distillation. Acetic acid comes first, then the anhydride.


BASF - 29-4-2004 at 08:12

Not that i am not able to search myself...but which patent is that exactly?


BASF - 29-4-2004 at 23:46

Ok. I can now spend a little time for searching:)

A more economical method for producing Na2S2O7(although a bit more complicated):

Description of US3885024


At 25 DEGC, 8 m.moles of KOH dissolved in isopropyl alcohol were added to a solution containing 15 mM of SO 2 and 0.1 m.mole of CuCl 2 in 20 ml of wet isopropyl alcohol (0.3 % H 2 O ), the solution being kept under an oxygen atmosphere at 760 Torr.

The reaction consumed 4 m.moles of O 2 and, after filtration, 3.75 m.moles of K 2 S 2 O 7 were isolated. The SO 2 conversion was 50 %, while the yield in pyrosulphate was 95 %.


94 m.moles of NH 3 dissolved in 40 ml of alcohol were slowly added to 100 ml of isopropyl alcohol containing 115 m.moles of SO2, 31 m.moles of H 2 O and 0.1 m.moles of CuCl 2 Pu 2 under an oxygen atmosphere. The reaction totally consumed 31 m.moles of oxygen and, after filtration, 5.8 g of a white solid were isolated.

We ascertained that it was constitued by ammonium pyrosulphate on the basis of the melting point (236 DEGC), I.R. spectrum (1,260, 1,095, 1,045 cm@@- 1) and the characteristic reactivity thereof. The SO 2 conversion was 50 %, and the yield in pyrosulphate was 89 %.

Prep. of anhydrides using pyrosulfates:

Abstract of GB424573

In the production of aliphatic anhydrides such as acetic anhydride by reacting a salt of the organic acid with a pyrosulphate as described in Specification 136,574, [Class 2 (iii)], or a pyrophosphate as described in Specification 299,342, [Class 2 (iii)], one of the reactants is a salt of ammonium or other volatile base which may be a primary, secondary or tertiary acyclic or cyclic amine, for instance the mono-, di-, and tri-methyl, ethyl, propyl and isoamyl amines, and pyridine. Ammonium acetate or acid ammonium acetate may be heated with sodium sulphate, yielding acetic anhydride and sodium and ammonium sulphates, which on further heating yield sodium pyrosulphate and ammonia which is available for the production of more ammonium acetate. Preferably the ammonium salt should be anhydrous; the dehydration of the salt may be effected in a vacuum or the water may be removed as an azeotropic mixture with benzene, toluene, xylene, dichlorethylene, ethylene dichloride or the like. Preferably the reaction between the acetate and pyrosulphate or pyrophosphate is effected in the presence of a diluent such as a hydrocarbon or chlorinated hydrocarbon or the acid whose anhydride is to be produced. The temperature of the reaction may be 100--140 DEG C., and may be the boiling-point of the diluent. Alternatively the organic salt may be a salt of a non-volatile base such as soda or potash, which is reacted with a pyrosulphate or pyrophosphate of the volatile base such as ammonia. In an example, anhydrous neutral ammonium acetate is heated with anhydrous sodium pyrosulphate and glacial acetic acid under reflux, and the anhydride is then distilled off. The residue is then heated to liberate ammonia and convert the sodium salt into pyrosulphate.

Another new pathway:

Abstract of GB352176

Acetic and other anhydrides, dichlor- and diaceto diethyl ether.-Aliphatic anhydrides are produced by heating diaceto-diethyl ether, CH3CH (O,CO,CH3)-O-CH (O,CO,CH3),CH3, or analogous bodies with a catalyst. Suitable catalysts are substances of acid or salt but not basic nature, such as hydrochloric and sulphuric acids, sulphur dioxide, and zinc and stannous chlorides. These ethers are produced by reacting with the aliphatic acid or its salt upon the corresponding dichlor ether, which is obtained by the reaction of hydrochloric acid on acetaldehyde or paraldehyde. In the examples, (1) the dichlor diethyl ether is allowed to stand at ordinary temperature with anhydrous alkali acetate; decahydronaphthalene and a small quantity of zinc chloride in acetone are added, and the acetaldehyde formed is distilled off; the acetic anhydride forms a layer separate from the decahydronaphthalene; (2) the dichlor ether is heated to 80 DEG C. with benzene and glacial acetic acid until the hydrochloric acid is expelled; the benzene is then distilled off and zinc chloride in acetone is added, when acetaldehyde distils off followed by acetic anhydride; (3) the diaceto ether is heated with zinc chloride in glacial acetic acid at 110-115 DEG C., acetaldehyde and acetic anhydride distilling off together; they are separated by fractional distillation; (4) the diaceto ether is heated with decahydronaphthalene and zinc chloride in alcohol; (5) the diaceto-ether is heated at 100-105 mm. pressure with zinc chloride in acetone; (6) the diaceto-ether is heated to 120 DEG C. and sulphur dioxide is passed in, acetaldehyde, paraldehyde, and acetic anhydride being distilled over; (7) the dipropio-ether is heated with zinc chloride in alcohol at 110 DEG C., propionic anhydride and acetaldehyde being distilled off.

[Edited on 30-4-2004 by BASF]

Marvin - 30-4-2004 at 14:06

Umm more economical pyrosulphate methods!

What about,

Na2SO4 + H2SO4 => 2NaHSO4

Then at 300C or so as mentioned,

2NaHSO4 => Na2S2O7 + H2O.

Nice tip btw, Hjalmar, method for AA is completely new to me.

BASF - 3-5-2004 at 08:28

"What about,

Na2SO4 + H2SO4 => 2NaHSO4

Then at 300C or so as mentioned,

2NaHSO4 => Na2S2O7 + H2O. "

Rather 400°C than 300°C, no?

Heating also costs money. I like the method for Na2S2O7 using SO2-gas because it does not involve such excessive heating.

...although I have nothing at all against simplicity :P

[Edited on 3-5-2004 by BASF]

the patent

Hjalmar_Poelzig - 3-5-2004 at 08:55

It is in US patent 1430304

Also, I seem to remember that it is not possible to get a quantitative yield of the pyrosulfate, as at some point SO3 will start to form.

But considering the cheapness and easy availability of the reagents this should not be much a problem.

Your method using KOH, CuCl2 and SO2 in IPA seems *very* nice to me, BASF, thanks! :D

[Edited on 3-5-2004 by Hjalmar_Poelzig]

frogfot - 11-5-2004 at 13:38

Problem with NaHSO4 decomp is that Na2S2O7 decomposes to SO3 and Na2SO4 at 460*C. Although temp diffeence is about 100 degrees, it may be hard to control with just a gas burner..
But still, not like I have any cheap SO2 to go with the other route, this is worth to test.

Acetate + CO2 method sounds to good to be true.. :o If this really gives good results, I know where to apply CO2 from next (theoretical) wash batch :P (ofcaurse a big gas drying column should be made..)

frogfot - 16-7-2004 at 07:09

KHSO4 dehydration seems to work very nicely. I've heated some, it melted and started to boil vigorously. After, boiling slowed down to almost zero, there appeared white fumes. So everything seems to go as planned. (Btw, OT, how about this one as an SO3 source)

I then broke solid pyrosulfate to small pieces, mixed with about 1,6 parts dry sodium acetate and 100 ml ethyl acetate. Stuff has refluxed now for about 0,5h (and counting..) with a CaCl2 drying tube and pyrosulfate chunks seem to dissapear/react.. While reaction mix becomes a bit more mobile.

Is ethyl acetate a good choice of a solvent? It shouldn't consume any acetic anhydride.. or?
In the end of reaction I wanna filter off solid K2SO4, evap solvent to hopefully recover AA anhydride..

Duh, ethyl acetate didn't do a squat.. however I've added some dry acetic acid to the pyrosulfate/acetate, and it heated alot! Nice to know that something is happening :)

[Edited on 16-7-2004 by frogfot]

[Edited on 16-7-2004 by frogfot]

Theoretic - 1-10-2004 at 10:55

What about Ca(CH3COO)2 + NH4HCO3? Partial hydrolysis of HCO3- should give CO2 in a dissolved state, ammonium carbonate shouldn't be alkaline enough to cause rapid hydrolysis of Ac2O, while CaCO3 should precipitate and thus facilitate the reaction. I think this is better than the NaCH3COO + CO2 method.

A better method for quickly forcing the pyrosulfate route is dropping all your acetate into the molten pyrosulfate and condensing the fumes, the anhydryde might be affected by hot pyrosulfate or heat generally.

neutrino - 1-10-2004 at 13:30

What I don't understand is why only salts of acetic acid work in the pyrosulfate method and not the acid itself. The acid is, after all, the hydrogen salt. Pyrosulfates will dehydrate sulfuric acid to its anhydride, so why not acetic?

Acetic anhydride- it's really that easy!

garage chemist - 16-2-2005 at 04:05

Acetic anhydride from sodium acetate, using S2Cl2

Translated from russian by Antoncho

Prepare 100g freshly fused NaOAc and 65g S2Cl2. A small quantity of NaOAc is placed in a thin-walled glass cooled in an ice bath. To this is added some sulfur chloride, the mixtr is vigorously stirred w/a wooden spatula, not allowing the temp to rise. Then some NaOAc is added again, and the process is repeated several times until all is mixed in. The semi-liquid mass is transferred into a 1 liter RBF. The previous operation is repeated 4 times, so that 400g NaOAc and 260 g S2Cl2 total are taken into work. The RBF is then equipped w/a reflux condenser and gently heated on a water bath to ~80-85°C. As soon as the rxn starts, the heating is removed, and in case the rxn gets too vigorous it's cooled w/cold water. After 20-30 min SO2 evolution ceases and the mixture is heated for 10 more min's on a boiling water bath. The rxn product is then distilled off under vacuum, then fractionally re-distilled at ordinary pressure, collecting the fraction boiling between 132-142°C.

For further purification it's distilled with 2-3% KMnO4 or K2Cr2O7 for breakage of sulfurous contaminants (test for their presence: 1 ml of the distillate upon neutralization w/pure NH3 mustn't give a dark precipitate on treatment with Pb(AcO)2) The yield is ~90% based on S2Cl2.

(From Rhodium's site)

I am surprised that nobody here works with sulfur chlorides. They are easy to make!
As I wanted to tell everyone in the "disulfur dichloride" thread, you can produce S2Cl2 by the liter when you have a ground- glass still.
Sulfur and chlorine are most easy to get.
Just insert a gas inlet tube into the thermometer adapter down into the boiling flask, exactly as you would do with a boiling capillary for vacuum distillation.
The chlorine doesn't even have to be dry!

I have 20ml of S2Cl2 sitting around, I only didn't make more because I didn't know of its usefulness to produce acetic anhydride until now.

Esplosivo - 16-2-2005 at 05:46

On this thread somewhere were stated the half-lives of acetic anhydride at different temperatures. At rtp what is the half life of acetic anhydride? Could it be possibly as low as stated previously (a few minutes IIRC!!!)? At what temperature should it be preferably stored then to slow down this decomposition, and at such temp. what would be the half life?

I have searched around the net but found not references to the half life ot ethanoic anhydride. Any information about it's half life would be awesome! Thanks.

chemoleo - 16-2-2005 at 06:38

We use litres of acetic anhydride for peptide synthesis, and the stuff is kept in a normal glass bottle in the dark, for months/years. I don't think it decomposes, because else it'd 1) say on the bottle, and 2) be useless for the synthesis.

Hangon - halflife - to what is it supposed to decay to?

Esplosivo - 16-2-2005 at 08:51

Right now I noticed that I had a mistake, Hermes was refering to the 'half-life' of acetic anhydride in water. I still don't like the term half-life though, but it makes more sense like that.

Thank you for your help, since at least now I know that acetic anhydride 'does not' decompose on storage.

trilobite - 16-2-2005 at 12:48

I've read that the downside of using sulfur chlorides as reagents is that the reaction products often have sulfur compounds as contamination. This is the case with preparation of ethylene chlorohydrin from ethylene glycol with S2Cl2, for example. This is not to say that the method is useless, but in some cases it is going to give trouble. One notorious example is hydrogenation catalyst poisoning.

Esplosivo - 23-2-2005 at 12:14

Would passing acetone (propanone) vapour through a thin (approx 5mm in diameter) copper tube heated with 4 bunsen burners of the LP type (fueled with butane/natural gas) be enough for the pyrolysis of acetone to ketene to occur? I am very interested in such a method of production, and will probably try it out next summer. Thanks for the help.

alchemie - 23-2-2005 at 13:24

Originally posted by Esplosivo
Right now I noticed that I had a mistake, Hermes was refering to the 'half-life' of acetic anhydride in water. I still don't like the term half-life though, but it makes more sense like that.

I'm sure what is meant is: "the half-life of the reaction of acetic anhydride with water".

The half-life of a reaction is just the amount of time that it takes for the consumption of half of the limiting reactant.

The confusion here is with the decai of radiactive nuclei, where "half-life" is the time necessary for the half of the original nuclei to decay.


aldol - 10-3-2005 at 19:51

hi hoffman i persume you want the anhydride for p2p production
i have used it before its not fun
as far as for making it tou will need a distillation of your anhydride
I WOULD NOT LIKE TO BE AROUND FOR THAT there is betta way to use phenylacetic acid after the making of p2p you have to disstill down the anhydride not fun

S.C. Wack - 10-3-2005 at 20:57

Hermes never reported back. The best patent wasn't mentioned, GB486964, an addendum to Consalvo's other patent, the aforementioned GB480953. (There are also FR versions of these two, FR47873E and the earlier FR809731)

Acetic anhydride from acetate

SAM4CH - 24-3-2005 at 10:43

Can I get acetic anhydride from molten sodium acetate by reaction with Cl2 gas?!

Phel - 4-5-2005 at 16:55

Sorry to bring this topic up again, but an alternative route is described, where an alkali salt of acetic acid is not used, just plain ol' KNO3:

Acetyl Chloride (40gms) is mixed with KNO3 (7gms) in a flask attached to an upright condenser, which is closed by a calcium chloride tube.
A vigorous reaction is accompanied by the evolution of chlorine and nitrosyl chloride, and after standing for half an hour, the mixture is heated on a waterbath and gradually raised to boiling temperature, at which it is maintained for two hours. The colourless liquid mass is the extracted with ether to remove the KCl, and the ethereal extract fractionally distilled.
Yield 15-20gms; bp., 138°C

Ref. The synthetic use of metals in organic chemistry by Arthur J. Hale p. 125

PainKilla - 4-5-2005 at 17:04

Too bad acetyl chloride is just as hard/harder to get than the acetic anhydride :(. most methods I have seen involve the use of phosphorous or obscure reageants.

Very interesting if you do have access to the compund though.

neutrino - 4-5-2005 at 17:07

We're trying to make this stuff from readily available chemicals, not exotic ones like acetyl chloride.

Phel - 4-5-2005 at 17:18

Indeed you are right Painkilla (edit: and neutrino), but luckily, not all countries has banned phosphorous from the surface of earth, thus the route via PCl<sub>3</sub> seems a little "easier" to come by.

[Edited on 5-5-2005 by Phel]

Mahlzahn - 5-5-2005 at 06:34

Is it easy prepare acetic anhydride from
CH3COOH and the method correct ?
First buy some dilute vingar in a food store and than the vingar is concetrated by distillation (<118 gegree celsius) to concetrated CH3COOH.
(first the water is boiled out).

The oxidation (dehydration) of CH3COOH
with a homemade desillation equipment
are difficultly.
(Desillation bridge with a electrolytic catalyst-grid)

Examble: ;)
4CH3COOH + O2 [catalyst] -->
2(CH3-CO)2O + 2H2O

Other methods walking fine too.

[Edited on 5-5-2005 by Mahlzahn]

S.C. Wack - 5-5-2005 at 23:57

I don't remember seeing these anywhere before, so here are some interesting variations on known routes.

US1926087 is unusually detailed for a Dreyfus patent. Sodium metaphosphate and acetate are used here.

DE283163 uses chlorine and acetate as usual, but more conveniently uses thiosulfate instead of S or SO2.

Esplosivo - 6-5-2005 at 00:17

The first one is surely a very interesting patent, the second I can't read because its in german. Can anyone at least give a quick translation of the most relevant part - shouldn't take more than a couple of minutes for a german speaking guy right? Please.... Thanks.

CD-ROM-LAUFWERK - 7-5-2005 at 04:09

i slowly mixed 35g P2O5 to 100g of acetic acid (under a strange reaktion, much heat)
i think that over 95% of the acetic acid reacted to the anhydrid, after destillation the stuff dont frezze at temperaturs around -20°C!
per 2 mol. of P2O5 u can use 16 mol. of acetic acid, that will give u 8 mol. acetic anhydrid (per 30g P2O5 use 100g acetic acid)

garage chemist - 7-5-2005 at 04:37

Relevant part of the second patent:

Crystalline sodium thiosulfate is made anhydrous by heating and then pulverized. 185 parts of this are mixed with 820 parts pulverised, anhydrous sodium acetate.
This mixture is reacted with chlorine at 40- 50°C. The chlorine addition is stopped when no more is absorbed. Approx. 285 parts are needed. Too much chlorine decreases the yield and the purity of the final product by chlorination of the carbon chain.
A high vacuum is applied to the reaction chamber and it is heated with pressurized steam of 7 bar pressure from the outside (oil bath should also work).
The acetic anhydride distills off.
With correct procedure, about 460 parts acetic anhydride are obtained (a 90% yield). A good yield is only obtained when water in any form is excluded from the reaction.

[Edited on 7-5-2005 by garage chemist]

Mahlzahn - 7-5-2005 at 06:45

Of interest are for instance the incursion:

4CH3COOH + 4NaBr ;) -->
4CH3COOBr + 4Na + 2H + 2e + H2

4CH3COOBr + O2 catalyst -->
2(CH3-CO)2O + 2Br2O


kmno4 - 1-6-2005 at 14:18

Originally posted by CD-ROM-LAUFWERK
i slowly mixed 35g P2O5 to 100g of acetic acid (under a strange reaktion, much heat)
i think that over 95% of the acetic acid reacted to the anhydrid, after destillation the stuff dont frezze at temperaturs around -20°C!
per 2 mol. of P2O5 u can use 16 mol. of acetic acid, that will give u 8 mol. acetic anhydrid (per 30g P2O5 use 100g acetic acid)

I tried to do the same , but effects was strange...
First, I prepared 218g AcOH(wet) and 76g P2O5. I placed P2O5 in distillation flask eqiped with vertical condenser and added about 20cm3 AcOH through it. In ten seconds it started to get brown, then blac and I saw a little white fumes inside flask. I added the rest of acid in 1/2 hour, temperature inside reached about 80C.
The mixture was deep brown, almost black, but I tried to distill this. I collected about 160g liquid between 118-121C. I was afraid to colect more, because the residue in flask was looking like boiling tar. So I recovered 160g of pure acetic acid:( I took this acid, placed it in a bottle, and slowly added 46g P2O5 in portions. First portion added and... nothing happens !! Another one - still nothing.
After adding this 46g nothing special has happened. No heat, no traces of any reaction*. Then I tried to calculate heat of such reaction:
Enthalpy(dH) this reaction is:+18kJ - not exothemic!
dG is:+38kJ - not possible:P:P!
Conclusions ?
At first time, wet(97%) AA was dried by P2O5, in small volume there was great evolving of heat and this caused exothermic charring of AA...
Propably H3PO4 is not formating in this reaction, maybe HPO3 and other phosphorus acids are. There must be used much more than 35gP2O5 per 100gAA to conver it (quantitiv)in Ac2O, at least I think so.
*in a bottle P2O5 started to get yellow(why?) and slowly disappears. I am wondering if it all is worth my job(P2O5 is cheap, about 10 dollars per 1kg:D)

CD-ROM-LAUFWERK - 2-6-2005 at 05:43

I tried to do the same , but effects was strange...
First, I prepared 218g AcOH(wet) and 76g P2O5. I placed P2O5 in distillation flask eqiped with vertical condenser and added about 20cm3 AcOH through it. In ten seconds it started to get brown, then blac and I saw a little white fumes inside flask. I added the rest of acid in 1/2 hour, temperature inside reached about 80C.
The mixture was deep brown, almost black, but I tried to distill this. I collected about 160g liquid between 118-121C. I was afraid to colect more, because the residue in flask was looking like boiling tar. So I recovered 160g of pure acetic acid I took this acid, placed it in a bottle, and slowly added 46g P2O5 in portions. First portion added and... nothing happens !! Another one - still nothing.
After adding this 46g nothing special has happened. No heat, no traces of any reaction*.

u added the acetic acid to the P2O5????
ther must be keten formed, the anhydrid from acetic anhydrid (its a gas)
i dont think that u used clean P2O5 :P

garage chemist - 2-6-2005 at 09:21

My preferred production method for Ac2O would be the production of acetyl chloride from benzalchloride or benzotrichloride (from toluene + chlorine) and acetic acid (ZnCl2 catalyst is required), then reaction of the acetyl chloride with sodium acetate.

Or you could directly use the acetyl chloride in place of acetic anhydride, they are both powerful acetylating agents.

IPN - 2-6-2005 at 10:38

Has anyone tested the method in US1926087?

2NaPO3 + 4CH3COOH -> 2NaH2PO4 + 2(CH3CO)2O

That seems just too easy. Distill off the anhydride (or filter, then distill), possibly with a column, and heat the orthophosphate to +300C (?) for few hours to convert it back to the metaphosphate.
This really should be tested IMO.
Phosphates and acetic acid are both OTC in their pure forms.

Ok, this seems to be similar to the experiments kmno4 and CD-ROM-LAUFWERK did. Maybe using a phosphate salt doesn't dehydrate it so strongly as P2O5 does and results in a purer product?
Sodium orthophosphate decomposes to pyrophosphate at 240C according to Patnaik.
Haven't found any other info on pyrophosphate than the MSDS, which just said that it decomposes to sodium and phosphorus oxides on heating.
I'll try to make some sodium metaphosphate from phosphoric acid and sodium carbonate tomorrow.

[Edited on 2.6.2005 by IPN]

[Edited on 2.6.2005 by IPN]

ketene lamp

blazter - 4-6-2005 at 09:06

I still don't understand why the ketene lamp method is so unpopular. Vogel has a rather descriptive diagram of the setup and explaination on page 373. If it could be done in 1948, it shouldn't be that hard to repeat.

Obviously there is the toxicity to deal with, but this is not so bad since ketene is supposed to have a characteristic odor (anyone know what it smells like?) The other problem is just how reactive/corrosive ketene is to various types of tubing. Vogel's device is made entirely from pyrex glass, but why couldn't say copper tubing be used? The electrical connections for the NiChrome grid he uses is supported by platinum, nickel, and brass terminals which would be directly exposed to the newly formed ketene.

Due to the reactivity of ketene towards organic materials, I could see that most plastics are probably unsuitable, but what about PTFE plastic? Anyways, I also dont see the need for two condensors as Vogel's setup has either. All one would really need is a very effecient condenser in a reflux configuration to recondense the unreacted acetone, while the newly formed ketene and other gas byproducts would flow out of the reactor.

I do not feel the need to so quickly discount this method. After all, once the equipment is dealt with, acetone is very cheap, and even if 30% molar yeilds are realized (conservative estimate from Orgsyn) one could obtain 15g of AA from 58g acetone without that much work (aside fromt the equipment setup again).

I'd really like to know more about what materials ketene is compatable with and what it will degrade (specifically plastics).

docberto - 4-6-2005 at 10:24

On the topic of ketene lamps, what would happen if you used MEK in a ketene lamp instead of acetone. Would you get methyl-ketene?

IPN - 6-6-2005 at 09:41

I neutralized 50ml of phosphoric acid with sodium carbonate and heated the viscous paste in a steel crucible over a propane burner for few hours. I should now have at least sodium pyrophosphate.
According to the Handbook of Preparative Inorganic Chemistry, metaphosphate requires temperatures over 1000C. Can this be correct? If so then this doesn't seem that good method after all..
Maybe it would work with just the pyrophosphate?

Acetic Anhydride

armo - 26-8-2005 at 05:17

Hi, I just like to know why all that interest in making AA. I think organic chemistry has a lot more interesting subjects than this one...


neutrino - 26-8-2005 at 15:57

This answer applies to just about every synthesis on this board: it's a useful chemical that can't be easily bought.

The_Davster - 26-8-2005 at 21:40

Also it is laughing in the face of the powers that be, they do not want us to have it as it is an 'evil chemical' with potentially illegal uses, when we just want it for regular chemisty.

Ketene lamp

armo - 29-8-2005 at 09:13

Ok, if you want to make AA, why not? I think the easy way is using a ketene lamp. I´m going to build one here since I have a research project were I will use ketene. I´m building the Vogel´s equipment and I´m going to use polypropilene tubes. I think they should resist very well to ketene. Anyway, as long as I start the ketene production I will put a photo and tell if PP tubes are resistant enough.

garage chemist - 4-9-2005 at 14:35

Preparation of Acetic Anhydride from AcOH and my new favourite water scavenger (HPO3)n has been unsuccessful.

The same as with P2O5 happens: there is a reaction and heat is evolved, but you just distill off pure acetic acid and the residue is boiling tar.

Preparation of Acetic Anhydride from AcOH and a water scavenging agent is NOT POSSIBLE.
Formation of a carboxylic anhydride is a VERY DIFFERENT reaction than formation of the anhydride of a mineral acid.
It cannot be effected by water scavengers. Different methods must be employed to form the CO-O-CO (carboxylic anhydride) group.

trilobite - 8-9-2005 at 15:44

I think SO3 has been used for the purpose. Funny that it doesn't work, but interesting nevertheless.

BASF - 14-10-2005 at 12:48

An addendum to the translation of garage chemist, at the top of the page:

water vapor at 7atmospheres is approx. 165°C hot.

I think after all, the thiosulfate-patent should be simple enough to try. (the ketene-lamp seems simple, but can i really operate a ketene lamp in a small room with only a small window for ventilation?:( )

[Edited on 15-10-2005 by BASF]

Illegal Parkinson - 18-10-2005 at 12:15

Hmm I dont really trust that can S2Cl2 can be used to make acetyl chloride. All the textbook procedures make use of PCl3.

[Edited on 20-10-2005 by Illegal Parkinson]

Just a further undiscovered method..

BASF - 22-10-2005 at 12:09

The preparation of acetic anhydride by the action of nitrogen peroxide on sodium acetate. Rodionov, V. M.; Oblitseva, T. A. Trudy VI Vsesoyuz. Mendeleev. S'ezda Teoret. i Priklad. Chim. 2(No. 1), 1002-3. From: Chem. Zentr. 1938, II, 4054. Journal language unavailable. CAN 34:43152 AN 1940:43152 CAPLUS


R. and Smarin found that when N2O4 was passed over dry NaOAc, Ac2O was formed. The reaction is 2AcONa + 2N2O4 = Ac2O + 2NaNO3 + N2O3. The N2O3 does not act on the NaOAc. Formation of Ac2O takes place also in mixts. with air but the yield is smaller than with the pure N2O4. The Ac2O yield is 85%. The N2O4 was prepd. either by strongly heating Pb (NO3)2 or from nitrosylsulfuric acid and NaNO3 according to the equation: HOSO2ONO + NaNO3 = NaOSO2OH + N2O4. The action of N2O4 on the salts of other fatty acids was investigated and good yields of propionic and butyric anhydrides were obtained. Somewhat poorer yields were obtained in the prepn. of the anhydrides of isovaleric and stearic acids.

What the?

chloric1 - 22-10-2005 at 13:24

So, I wonder what keeps the N2O4 or byproduct N2O3 from deteriating the Ac2O? Is this done at room temperature?

[Edited on 10/22/2005 by chloric1]

The_Davster - 22-10-2005 at 13:59

Cool method:) Literally...
In order for the N2O4<->2NO2 equilibrium to lie to the left, the N2O4 side, this reaction would likely have to take place at a sugnificantly lower than room temp.

[Edited on 22-10-2005 by rogue chemist]

acetic anhydride

ramarao - 27-11-2005 at 19:17

You can try distilling a mixture of Anhydrous sodium acetate and acetyl chloride.Yields is fairly good.You have to add acetyl chl;oride drop wise to sodium acetate .The distilled crude iquid has to be redistilled to get pure product.

stygian - 27-11-2005 at 19:44

I'm going to assume that this was overlooked, otherwise I don't see why there'd be a continuing discussion on this.

Acetic Anhydride and Propionic Anhydride
Drone #342:

Here's what everyone's looking for. Some things are a little wierd about it, like the fact in the acetic anhydride synth they use a small quantity of acetic anhydride as a solvent. However, as one sees in propionic anhydride, such a solvent may not be necessary.

From "Chemistry & Industry", 1945, p. 382; "LABORATORY METHOD FOR THE PREPARATION OF ORGANIC ACID ANHYDRIDES" by Jehuda Orshansky and Eliahu Bograchov.

"...(1) Acetic anhydride. To 50 g. acetic anhydride in a round-bottomed flask of 1500 cc. capacity, placed in cold water, 440 g. of powdered sodium acetate (dried by fusion at 320 C) and at the same time a solution of 22. g of sulfur in 320 g. bromine is added while stirring. The operation takes about 30 minutes.

"The mixture is then stirred for a further 5 minutes, after which period the initially dark brownish-red colour has changed into pale yellow, and the anhydride is distilled off from a water bath under reduced pressure. The crude anhydrie (310 g) is redistilled under normal pressure, and the fraction boiling between 134-138 C is collected. Yield, 295 g. of 98% purity = 87.5%. The so purified anhydride contains neither bromine nor sulphur compounds and leaves no residue on evaporation..."

"(2) Propionic anhydride. To 40 g. fused and powdered sodium propionate in a flask of 250 cc. capacity a solution of 2 g. sulphur in 22 g. bromine was added while stirring. The temperature was kept at about 50 C. When the operation was completed, the anhydride was distilled off in vacuo. The crude product (25g.) was fractioned under normal pressure, and the fraction 155-156 C was collected. Yield, 23 g. propionic anhdride of 90% purity = 85%..."

The paper mentions that other alkali metals and alkali earth metals work just fine. Calcium propionate is a food preservative added to cheap white bread to keep it from molding. With these nuggets of information, the most closely watched reagent on the DEA's watched list, propionic anhydride, just turned OTC. I can almost see the fentanyl analogs clogging the opioid market already.

The one reason, and justafiably so, they poo-poo using chlorine (which does indeed work nicely) is that its a hassle to work with, especially considering the fact that they'd be adding a gas to a solid to make a liquid. I propose that perhaps with the use of chloroform as a solvent, chlorine could be bubbled in readily, and the reaction would go as previously stated. I assume they tried to avoid using extra solvents in hopes of staying away from azeotropes messing up their products' purity, so this may or may not work, depending on what you're trying to do.


To 66.5g fused AcONa (made from AcOH + NaOH) was added a soln. of 3.3g S in 48.4g Br2 over 5mins, under manual stirring. Brown colour disappears rapidly on stirring to give slightly off-white mixture. Stirring continued for another 15mins. Mixture has v. strange consistancy; becomes almost liquid while stirring but as soon as one stops stirring it becomes solid again. Thus formed sludge was scooped up and dumped in a distillation flask (spilling much in the process - do this rxn in the same flask you're gonna distil from), and distilled.

Yield of Ac2O = 29.2g as a clear pungent liquid.

Note that basically all you need is NaOAc, S, and Br2. And some distilling equipment.

BASF - 28-11-2005 at 11:36


I don´t want to discourage you after finding this "unknown" method....
1) the first method uses Ac2O as a solvent. It´s always nice if you need the target compound when you attempt to make the target compound...
2)Bromine is not exactly OTC, nor is it cheap or anything else. -Agreed?

Please no more acetylchloride/bromine - methods!

[Edited on 28-11-2005 by BASF]

garage chemist - 28-11-2005 at 14:09

As soon as you get a kilogram of a bromide salt (CaBr2*H2O in my case, it is sold for making high-density drilling fluids, a saturated CaBr2 solution has got a density of 1,64 g/ml) and adapted the process of bromine production according to your ressources (like what oxidiser you use: I use NaClO3, but H2O2, MnO2, Cl2 and others just work as well) you will find bromine to be a cheap chemical.
Also, the bromine can be fully regenerated from the residues of Ac2O production, it is present in the form of bromide there which just needs to be treated with acid and an oxidiser.

But the N2O4/NaAc process is most interesting, if it works it has got enormous potential.
Let's discuss methods of NO2/N2O4 production. Heating of lead nitrate might be the method of choice for research labs, but for us lead nitrate is too expensive and difficult to make.
Heating calcium nitrate might be a better idea, as this is easier to get (fertilizer).
But this is the tetrahydrate, dehydrating the salt using heat is really messy, I've tried it.

Any more ideas for convenient and controllable production of dry NO2/N2O4?

[Edited on 28-11-2005 by garage chemist]

stygian - 28-11-2005 at 14:20

Originally posted by BASF

I don´t want to discourage you after finding this "unknown" method....
1) the first method uses Ac2O as a solvent. It´s always nice if you need the target compound when you attempt to make the target compound...
2)Bromine is not exactly OTC, nor is it cheap or anything else. -Agreed?

Please no more acetylchloride/bromine - methods!

[Edited on 28-11-2005 by BASF]

1) I can't say for sure but it does say that the anhydride as a solvent "may NOT be needed"

2) Actually, bromine IS "exactly OTC". Go to a pool supply store, buy sodium bromide, sodium bisulfate, and potassium persulfate. All your ingredients right there in one shot. Or you could get a cheap gallon of HCl at a hardware store for $5, and use an oxidizer that may be more convenient.

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