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Author: Subject: Dry distillation of Copper Acetate
k89
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[*] posted on 1-8-2008 at 03:11
Dry distillation of Copper Acetate


I read somewhere that some metal acetates like copper acetate,on dry distillation, give acetic acid.
I've dry distilled copper acetate (at 300 C) to get a greenish liquid.
It is probably an acid because it bubbles with potassium carbonate.
Iit smells nothing like acetic acid.
Why is it greenish?
I've done this experiment three times with similar results each time.

BTW,the copper acetate was made by precipitating copper hydroxide from copper sulphate using NaOH ,dissolving in excess of glacial acetic acid and boiling off the excess acetic acid and water to get (dark green)copper acetate crystals

[Edited on 1-8-2008 by k89]
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Mr. Wizard
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[*] posted on 1-8-2008 at 07:08


I'm assuming you washed the resulting sodium sulphate from the precipitate of copper hydroxide. Any 'left over' NaOH would need to be removed too. Pardon me if this is a stupid question. Any left over sodium sulfate or hydroxide reacting to give sodium acetate would leave a water rich salt behind, which may contain enough water to dissolve the copper acetate. Does a bit of the material color a clear flame yellow?

Have you weighed the materials you are using, and do the weights correspond to what you should get? How many moles of copper are you starting with, and how many do you have at the end?

Anyway, a green liquid that contains copper at 300 C and doesn't decompose sounds interesting.
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JohnWW
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[*] posted on 1-8-2008 at 14:22


Cu salts of aliphatic carboxylic acids - "soaps", including especially Cu naphthenate which is used as a wood preservative and is made by precipitation from naphthenic acid (made from the naphtha fraction of petroleum distillates), answer to that description, being an intense green color.
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kclo4
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[*] posted on 1-8-2008 at 21:17


The green is possibly, and sounds very likely that it is a contaminantion of some sort of copper salt coming over.
How did you distill it, would that be possible?
What color was the oxide left in the distillation flask? I wouldn't be to surprised if some acetic anhydride formed.
Could you perhaps tell us more about the physical properties of your distillate? I.e if it freezes easily or if it requires very low tempuratures, if it mixes with water or forms two layers?

Thanks! :D
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k89
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[*] posted on 2-8-2008 at 03:35


Quote:
Originally posted by Mr. Wizard
I'm assuming you washed the resulting sodium sulphate from the precipitate of copper hydroxide.


Yes,I washed the hydroxide precipitate thrice with tap water.Isn't this enough to remove NaOH as well?

Quote:

Have you weighed the materials you are using, and do the weights correspond to what you should get? How many moles of copper are you starting with, and how many do you have at the end?


No , I dont have any apparatus for taking weight measurements.Haven't done the flame test ,but I doubt that any sodium salts are present after the washings.


Quote:

The green is possibly, and sounds very likely that it is a contaminantion of some sort of copper salt coming over.


With an excess of potassium carbonate the green liquid gives a bluish-green precipitate and turns colourless. (sorry I forgot to mention this earlier) This indicates the presence of a copper salt.
I use a basic distillation setup - A conical flask sealed with a one hole cork.A long glass tube through the cork connected to the receiver.
There wont be any copper salt contamination unless a salt sublimes into the receiver.
But which copper salt is volatile enough to end up in the receiver?

I didnt use a condenser for this.The receiver was cooled by submerging it in a vessel containing cold water.This might explain why there were no salt deposits anywhere.


The residue in the flask was reddish (Cu2O?)
I added water to the green liquid.Acetic anhydride should be hydrolysed to acetic acid.But the smell remained the same.
Also , I did not see two layers , so it is probably miscible with water.
I was working with very a low quantity of it in a test tube(approximately 5 ml)
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blogfast25
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[*] posted on 2-8-2008 at 04:48


Quote:
Originally posted by k89

I use a basic distillation setup - A conical flask sealed with a one hole cork.A long glass tube through the cork connected to the receiver.
There wont be any copper salt contamination unless a salt sublimes into the receiver.
But which copper salt is volatile enough to end up in the receiver?


That's rather academic, because with that simple set-up some mechanical entrainment of tiny particles of the copper acetate by the evading distillate is not unlikely to happen. No sublimation of a copper salt is thus needed for traces of copper to be found in the distillate.

To actually prove some copper compound or other has effectively distilled over (through evaporation, not entrainment) you would have to use a short distillation column, like a Vigreux or something like that...

Was the Cu (II) acetate anhydrous or the monohydrate?
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[*] posted on 2-8-2008 at 21:01


Quote:
Originally posted by k89
Quote:
Originally posted by Mr. Wizard
I'm assuming you washed the resulting sodium sulphate from the precipitate of copper hydroxide.


Yes,I washed the hydroxide precipitate thrice with tap water.Isn't this enough to remove NaOH as well?

...


Not always, many hydroxides and hydrated oxides really retain alkali hydroxides strongly. Plus tap water can introduce contaminates of its own. Deionized or distilled water is your friend, at least until you've established a process's details after which you can try switching to tap water.


blogfast25 is correct that it is not unlikely that you had some mechanical carry-over of copper compounds. This is even more likely given you were doing dry distillations, during which crystals of the substance being heated can decrepitate or otherwise fragment into fine particles. As you said nothing about dehydrating the solid copper acetate, it most likely was the dihydrate, meaning that water would be boiling off during heating, most certainly disrupting the physival structure.
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[*] posted on 3-8-2008 at 04:47


One way of 'washing' a fresh precipitate of Cu(OH)<sub>2</sub>.n H<sub>2</sub>O is to boil it: it loses its 'crystal water' very quickly leaving you with a much smaller volume of black CuO (which also precipitates well, as it's quite dense). The blue Cu (II) 'hydroxide' is incredibly voluminous and can create the illusion that the few traces that were carried over (assuming that is what happened) represent quite a bit of copper.
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[*] posted on 3-8-2008 at 05:07


The acetate was a dihydrate.
The color was probably caused by mechanical carry over of copper salts.

Thank you all for the help.
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[*] posted on 3-8-2008 at 05:29


Copper Acetate (CuAc2) is a monohydrate salt (CH3COO)2Cu.H2O; MW: 199.5; Copper Content : 31.8 %.

CuAc2 has got relatively low solubility in water ( as compared to other metal acetates such as Manganese, Cobalt, Nickel etc.) IIRC it is about 6 % (w/w) at room temperature. And yes, CuAc2 is votile enough to get carried over to distillate. Even a trace of CuAc2 in water / Acetic Acid / Acetic Anhydride will impart distinct greenish colour to it

gsd
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[*] posted on 8-2-2011 at 23:40


The Wiki method to copper acetate and distillation


I found photographing this extremely difficult. The colour changes didn't seem to show up very well on the camera and I've had to take over 200 photos to find some that look right. Even then, they're not very good.

I don't know if this is the camera or some optical property of the salts. The camera is a Canon Z100fd. I absolutely hate the thing and am only using it because the other one has sand in it. This one has a remarkable tendency to blur, produce images covered in noise, not focus or do weird things with the colours anyway.

I decided to give the Wiki method a go as it says "Citation needed" over the method and someone is asking about it on the discussion page.

I've been up since 6am yesterday, and it's now 7.30am the next day. I've just finished doing this and thought I'd upload the photos.

Some of you may enjoy this thread I found whilst searching around.

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


I weigh out some dry copper sulphate


I add to this the appropriate amount of glacial acid (acetic). The sulphate remains unchanged but becomes a paste. I begin dripping in 10% ammonia and a vivid blue appears. This produces a lot of fuming and the solution warms up.




It's greener than that. I simply can't get the camera to show it. The mixture is now a slurry of aqua / mint green sand. It is thick, but easily stirred. There is a lightly sparkling suspension.


I vacuum filter the result until no more liquid is dripping. Then lightly rinse it with some cold, distilled water. You may be able to see there are whitish lumps in the cake still. Given how dilute the ammonia was, I'm not sure how these could be dry copper sulphate. This is the best picture of it in terms of it's colour. I would be interested to know why this method alone is up on the Wiki, as it seems complex and wasteful compared with the others.


Back with the weird colours again! You'll just have to trust me, it's green! I'm going to try putting it under vacuum distillation.


That is horrendously bad! The camera is either having a fit or the salt is like Ferric Chloride, changing colour depending on the type of illumination. That is a 100ml flask.


I begin heating and, as we might expect, water is first collected, at approximately 45C. The temperature sits still here, but then start gradually rising, over an hour as each drip comes through. As it rises, I note that the drips are forming swirls in the receiver as they splash into it - as you'd see adding brine to distilled water. I suspect the distillate is building in salt concentration. When it begins approaching 62C, I see a definite green tint in the drops leaving the tip of the condenser, and switch to another receiver.


There is a very small quantity of white powder in the base. At the time, I suspect this is the last of the water leaving, and am waiting for the temperature to sit still at a specific point. But it doesn't. It floats up towards 70C, and then back down towards 45C - all in a continuous band.


I see that the column and condenser are clouding up with a white, dust like, solid coating. I see 'snow flakes' of the powder sputtering into position around the thermometer and worry the column is on it's way to clogging. But it never does. You can see the problems I am having with the camera here, as I know a number of you have seen this room in other photos, and the walls are bright white. Yet they appear yellow here. All of the colours in the flask are also tinted yellow. The whitish powder on the side of the flask is A4 paper white.



There are numerous colours all over the insides, ranging from green to yellow to white to grey.



Inside the receiver, I also see white, with a hint of bluey / green. It is quite obviously spraying out as a mist or vapour as opposed to dripping. You can see where drips are cutting through the layer.


Looks like it's clogged, but it's not.


Traces of brown appeared at the edges of the flask early on. As the still head has now cooled back to 45C, despite continuous heating, this is the end of the process. This is clearly a decomposition product.


This is green in person. Giving it a sniff, it smells like acetic to me.



This also smells acetic like. But neither is as intense as the glacial acid used at the start. The one above is close however. I wonder if those swirls I saw were acetic?


VIEWERS! PREPARE YOUR SPECTROSCOPIC TECHNICOLOUR PERL AND DEAN GOOGLES FOR SOME EXCITING CHANGES!

That IS blue. As soon as the glass was opened to the atmosphere, it all began rehydrating again, quickly!


Lined, but not clogged. Interesting the way it goes from green to blue.



There is a free running, fine chocolate powder in the flask.


I had to get a photo of this fast. On emptying the powder out, it immediately began to glow and smoulder.




Conclusion

Use the glacial and peroxide method, or glacial onto the hydroxide.

If you're thinking of distilling it, you may want to look into sublimation instead.

I couldn't find the triple point of the acetate.

[Edited on 9-2-2011 by peach]




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[*] posted on 9-2-2011 at 01:12


Nice experiment. The pyroforic copper powder is the most exciting feature though.
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[*] posted on 9-2-2011 at 04:41


the copper acetate decompose to acetone an other organic product but arrond 20-60% is acetone
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[*] posted on 9-2-2011 at 06:50


Woelen has given a nice preparation for copper acetate which does not require glacial acetic acid. It was so easy that even I could do it.
Quote: Originally posted by woelen  
I made copper acetate myself from copper sulfate, sodium hydroxide and dilute acetic acid.
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[*] posted on 9-2-2011 at 07:40


Quote: Originally posted by plante1999  
the copper acetate decompose to acetone an other organic product but arrond 20-60% is acetone

What is the reference for that claim?
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[*] posted on 9-2-2011 at 10:17


The white sublimate seems likely to be cuprous acetate. Yellow colors could be CuOH or something close to it. To the extent that there is copper in the receiver, it seems likely to have passed over as cuprous acetate; on contact with water and air, the following happens (according to Gmelin, referencing Berzelius; I paraphrase):

Hydrolysis: CuCH3COO + H2O -> CuOH + CH3COOH (CuOH is apparently yellow, though I think I have also seen claims that the pure substance is white)
Oxidation: 2CuOH + O + 2COOHCH3 -> 2Cu(OH)(CH3COO) + H2O

Following this, if you have acetic acid in the solution, you also get the conversion of the monobasic cupric acetate to the regular Cu(CH3COO)2 that you started with. Rather clever of the cupric acetate to pass over and re-form itself like that! However the last reaction is really an equilibrium; at higher temperatures the monobasic cupric acetate and acetic acid tend to predominate, which is why you can't just boil a solution of cupric acetate to dryness to obtain the material.

Quote:
I suspect the distillate is building in salt concentration.


What colorless salt solution would you have here? I would suspect instead that you are getting acetic acid, following on the initial run of water (or maybe even backwards...). Cu(CH3COO2) + H2O -> CuOH(CH3COO2) + CH3COOH proceeds easily. I think the difference in index of refraction is enough to give you the little visible swirls.

Quote:
What is the reference for that claim?

In trying to answer this question myself, I came upon the following:
http://article.pubs.nrc-cnrc.gc.ca/ppv/RPViewDoc?issn=1480-3...

They claim acetone as a main product for copper acetate decomposition, with acetic acid and methane as minor constituents, and CuO as the end result for the copper. Frankly however, this is at odds with the older authors (not to mention the experiment peach conducted), and I'm not sure I believe the result. It is possible that starting with anhydrous cupric acetate would lead to quite different results than what peach saw, since the initial conversion to monobasic acetate along with release of acetic acid requires some water. I suppose differences in the rate of temperature increase could also account for discrepancies.



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[*] posted on 9-2-2011 at 10:28


Thermal decomposition of Cu(CH3COO)2 x H2O was investigated experimentally in many publications (interested can start their search in Thermochimica Acta and Journal of Thermal Analysis and Calorimetry).
The only interesing product of pyrolysis is Cu(I) acetate as white sublimate, in low amounts (<5%).
The rest are: acetic acid, water, CO2/CO/C and Cu/CuO.



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[*] posted on 9-2-2011 at 11:39


I recommend the old paper 'Observations on the Phenomena and Products of Decomposition When Normal Cupric Acetate is Heated' (Journal of the Chemical Society, Volume 81, Part 2), from 1902. Lots of detailed observations on colors, exact temperatures at which things started to distill and so on. The products described there match what kmno4 indicates, except that traces of acetone are also mentioned (0.01% or so).
However... their approach consisted of slow heating, and it seems that they generally held the acetate at whatever minimum temperature was required for some reaction to proceed, even if it took hours to complete. They mention running one experiment for a week (at which point decomposition was still incomplete) and another for nine and a half hours. It's entirely reasonable to suppose that faster heating (16C per minute is mentioned in the earlier paper I linked) would reach temperatures sufficient to produce other products, like acetone, long before decomposition to lower-temperature products was complete.
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[*] posted on 11-2-2011 at 13:30


This is the glassware from the distillation attempt, the day after. There's a solution of copper sulphate in the top left as a reference for the colours.



Yep, I also agree that kmno4 has got it; having had a chance to sleep on it, that seems right.

I don't know about the acetone. I couldn't smell it but, if it's produced in such small amounts, I doubt it'd be possible over the acetic anyway.

It took me about 3 hours to do the 30 or so grams in there.

Anyway, I decided to give electrolysis a go as well. The level of conductivity is stupidly low so, again, it takes a long time. It's not like sulphuric. I thought I'd try electrolysis as I'd not only just run out of peroxide but also thought it was nice in terms of variables. The only things in the process are copper from mains cables (which are produced to a high standard) and the acetic, so not a lot of options for contaminants.

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


Negative on the top, positive on the bottom. Bit of tubing along the positive to shield it on the way down. This is so the salt forms at the bottom and stays there. If it floats back up to the negative at the top, 'snot' ensues from that electrode.

I tried various different power supplies, ranging from 4 to 19Vdc and fractions of an amp to a couple of amps.

I think I ruined this attempt by playing with the electrodes whilst it was going, encouraging the salt to start mixing in the beaker. The result, jellyfish like tails of copper began forming and floating down. They're mesmerising to watch but, once they start forming, it gets a lot worse.

I ran this for about 12h. I expect the current flow was around 0.1A, before the copper snot appeared. Afterwards, I believe the copper bridges were partially shorting the supply and heating the solution, as it got hot enough to steam. The copper shorts tend not to touch the coil at the bottom. As they get closer, gas evolution on them increases, making them float upwards and appear to be swimming.

The sludge at the base is from me swirling the top electrode to knock it off. You can also knock sludge off electrodes by momentarily reversing the polarity (diode with a resistor across it). Using a pulsed power supply can also help avoid bonding. But the stuff should be avoided in the first place by keeping the electrodes further apart.








Do not attempt to microbake the bulk of the solution off, even small amounts of it will go ape shit! :D John has conducted this experiment for you, and does not recommend a repeat.







Half a gram, sir?! For real!? For keeps!?



Conclusions

Don't touch the beaker or electrodes at any point once it's started.

Expect it to take absolutely ages.

I think you'd probably need something on the order of 100 to 200V to get a decent current through this; decent as in, a single amp.

Don't connect your experiment directly to the rectified mains. You need an isolation transformer involved to float the voltage. A yellow box site transformer could be a good candidate. Do you use them in the US? I just realised, I've not consciously thought about it before.



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


A 3-15Vdc, 25 amp supply turned up in the post this morning. I thought I'd give it another go. This time, exercising more restraint with regards to poking at it and I'll probably leave it going a lot longer. I'll see how it goes.

In the Mark 2 version, I have modified the electrodes a little.

Since I'm using strands of conductor, if they get eroded through near the plastic tubing, the rest of it becomes useless, and you shouldn't be poking around with it once it's going as it'll mix the salt with the rest of the solution. To this end, I've made the electrodes into a coil, with no sections bridging.

I made the coils like flat pancakes, to keep them as far apart as possible. I then bent the centre of each up a little (towards each other). I was thinking that this would encourage them to erode from the free end towards their supply, rather than the supply being cut off with lots of wire left dangling.

I formed two little arms on the top coil to help it stay in place at the top.






Computer sez no



Jar sez yes



[Edited on 12-2-2011 by peach]




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[*] posted on 13-2-2011 at 10:19


I let that go for 24h in total.

I kept checking back every now and again to note the voltage, current and temperature. For all intents, the voltage stayed between 14.8 and 15.1Vdc throughout the entire thing. And the current, the needle was barely off the first mark. At it's peak it reached what I would estimate to be 0.3A.

This means you could meet the peak current capacity of the process at this voltage with a wall adaptor. A laptop power supply is about 19Vdc @ 3A.

The temperature rose to 38C. This was a lot less than the first try due to the lack of direct copper bridges between the two.

I was planning to leave this going for 24h and, as luck would have it, I noticed the electrodes stopped fizzing and the temperature beginning to fall around this point. Emptying it all out, I discovered the electrode had done what I thought it would - eroded through at the supply end; but with a fair bit of the rest gone.

By not touching the electrodes or jar at all, and letting it run longer, it has yielded a usable amount of the acetate. Not a lot (7g once vacuum dried), as the amount of copper falling back was still significant. But it's clean, and if you don't have much to use at home, all it takes is some acetic in some form or another.

Also, check out the last two photos. That is exactly the same pile of salt in two different places. One is teal greeny blue, the other is copper sulphate blue.











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[*] posted on 17-2-2011 at 23:15


awesome pictures peach :)
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[*] posted on 16-7-2011 at 07:03


US2073686
Quote:

Proocess for manufacture of acetic anhydrid, which comprises heating to a temperature between 200-450c a mixture of cupric acetate with copper salt of a strong of hydrochloric acid and sulfuric acid




Attachment: US2073686.pdf (113kB)
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[*] posted on 16-7-2011 at 18:56


Interesting patent. Looks like it would be easy enough to try... but reading it, I suspect it is worthless. Among other things:
- the single example provided is not written as if it was actually executed (quantities that really should be given as a single concrete value are listed as a range)
- the author suggests drying cupric acetate and cupric sulfate separately and then combining them before performing pyrolysis. No mention is made of any process or step for achieving intimate admixture of the two solids (which I have to think would be essential).
- no yields are given

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[*] posted on 17-7-2011 at 00:22



Quote:

Band 61 B5, Seiten 125-6.

2.2.7.2 Thermische Zersetzung

Die Zersetzung von AgCH3COO, die bereits von Chevenix [1] beim Erhitzen über einer Kerzenflamme (starker Geruch nach Essigsäure) beobachtet wurde, setzt bei 210°C ein (Dunkelfärbung), erfolgt hauptsächlich zwischen 220 und 240°C und ist vollständig bei etwa 300°C [2]. Beim Erhitzen im offenen Rohr oder im bedeckten Porzellantiegel werden als Zersetzungsprodukte im wesentlichen Essigsäure und Ag neben wenig C02 und C beobachtet entsprechend 4AgCH3COO ->4Ag + 3CH3COOH + CO2 + C [3 bis 5]. Von Kachler [2] werden als Reaktionsprodukte gefunden (in Gew.-%, nach vorstehender Gleichung berechnete Werte in Klammern): 64.61 Ag (64.67), etwa 26 bis 27 Essigsäure (26.95), 4.09 bzw. 8.42 CO2 (6.59) und 1.21 C (1.79). Auch bei thermogravimetrischer Untersuchung wird als Zersetzungsprodukt (bei 280°C) metallisches Ag und kein Ag2O erhalten [6]. Erfolgt die Zersetzung jedoch unter streng wasserfreien Bedingungen, so werden als Zersetzungsprodukte nur Ag2O und Essigsäureanhydrid erhalten nach 2AgCH3COO -> (CH3CO)2O + Ag2O. Bei der Zersetzung von 2.1077 g trocknem AgCH3COO bei 300 bis 400°C in einem Quarzgefäß unter Argon (1 atm) werden nahezu quantitativ Ag2O (96% der Theorie) und Essigsäureanhydrid (93%) gebildet neben wenig Ag (0.012 g) und CO2 (0.016 g) [7]. Die primäre Essigsäureanhydridbildung ist bereits von Kanewskaja, Schemiakin [8] bei Untersuchung der thermischen Zersetzung von AgCH3COO (im Gemisch mittrocknem Sand) unter einem CO2-Strom von 20 bis 25 Torr angenommen worden. Zwar fanden diese Autoren nur wenig Essigsäureanhydrid neben viel Essigsäure (entsprechend den früheren Angaben von Kachler [2] und Iwig, Hecht [4]), doch ist dies auf die leichte Hydratisierung des Essigsäureanhydrids zurückzuführen, die sowohl durch H2O-Spuren in dem schwer zu trocknenden Silberacetat als auch durch H2O erfolgen kann, das beim teilweisen Zerfall der Essigsäure gebildet wird.
Für die vollständige thermische Zersetzung von AgCH3COO in Gegenwart von H20 muß auf etwa 230 bis 340°C erhitzt werden. Abweichend von der trocknen Destillation wird neben den Hauptprodukten der Zersetzung Ag, Essigsäure und CO2, kein Kohlenstoff gefunden [2].
Über die Darstellung von Ag-Schwamm durch Erhitzen einer getrockneten Paste von AgCH3COO im Tiegel auf 400 bis 500°C s. [10].

...

[1] R. Chevenix (Ann. Chim. [Paris] 69 [1809] 5/58, 19, 22; Ann. Physik 32 [1809] 156/201, 167, 179). — [2] J. Kachler (Monatsh. Chem. 12 [1891] 338/49, 340). — [3] K. Birnbaum (Ann. Chem. 152 [1869] 111/21, 119). — [4] F. Iwig, 0. Hecht (Ber. Deut. Chem. Ges. 19 [1886] 238/42). — [5] J. Redtenbacher, J. Liebig (Liebigs Ann. Chem. 38 [1841] 113/40, 131).
[6] D. A. Edwards, R. N. Hayward (Can. J. Chem. 46 [1968] 3443/6). — [7] A. D. Kirschenbaum, A. G. Streng, M. Hauptschein (J. Am. Chem. Soc. 75 [1953] 3141/5, 3143).— [8] S. J. Kanewskaja, M. M. Schemiakin (Ber. Deut Chem. Ges. 69 [1936] 2152/7, 2154). — [9] V. I. Yakerson (Izv. Akad. Nauk SSSR Otd. Khim. Nauk 1963 1003/11, 1007; Bull. Acad. Sei. USSR Div. Chem. Sci. 1963 914/21, 916). — [10] T. Yamanaka, H. Nidorikawa (Japan.P. 6720 [1956] nach C.A. 1958 10462).

http://www.sciencemadness.org/talk/viewthread.php?tid=9&...



Quote:

The JACS article uses inert gas sweeping, which is often necessary to get yields of desired pyrolysis products from salts. Yield 600 mg on heating 2.1 g at 300-400C 2-1/2 hrs
The other refs above:
DE556247
US2073686
Proc. Chem. Soc. 20 (1957)
Zhur. Neorg. Khim. 5, 558 (1960)
Izv. Akad. Nauk SSSR, Otd. Khim. Nauk 1003-1011 (1963)
Ber. 69B, 2152 (1936)
http://www.sciencemadness.org/talk/viewthread.php?tid=8024&a...


I think same result will achive with copper acetate too.


Quote:


The thermal decomposition of calcium, sodium, silver and copper(II) acetates
Journal of Thermal Analysis and Calorimetry
M. D. Judd, B. A. Plunkett and M. I. Pope
Volume 6, Number 5, 555-563,
DOI: 10.1007/BF01911560
http://www.sciencemadness.org/talk/viewthread.php?tid=15455&...

Intresting patent ,but there is no AC2O in decomposition components
With thanks to@ solo


[Edited on 17-7-2011 by Waffles SS]
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[*] posted on 17-7-2011 at 05:50


Quote:
(J. Am. Chem. Soc. 75 [1953] 3141/5, 3143).

It is about Ag salts, but if it would go like this :
2 Ag(CH3COO) + 3/2 S -> Ag2S + (CH3CO)2O + 1/2 SO2
, then what about Cu acetate (anhydrous) ? It is interesting in itself, not beacuse of stupid Ac2O.
Does anybody know something about this reaction ?
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