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Author: Subject: Copper(II) Aspirinate Synthesis
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[*] posted on 8-2-2008 at 12:48
Copper(II) Aspirinate Synthesis


I've been interested in this compound and analogues for a while now, and have prepared it by several procedures at least a few dozen times. I'll outline two procedures here, with the pros and cons of each listed--focusing on the easier and less time consuming one. I don't know if I have yield data in my notebook, so I'll post that when I prepare more.

Reagents:
acetylsalicylic acid - <em>abbrev. 'ASA' or asp., easiest to obtain is probably purified from uncoated OTC aspirin tablets (outlined below)</em>
copper(II) sulfate pentahydrate - <em>easier to dissolve than anhydrous? hardware, pottery, and gardening stores have this</em>
sodium bicarbonate - <em>baking soda (alternatives: sodium carbonate, potassium bicarbonate)</em>

Necessary Apparatus/Glassware/Equipment:
appropriate sized Erlenmeyer flasks and beakers
filter funnel (Büchner or Hirsch is preferred) and paper
(vacuum flask, pump, tubing, etc. for vacuum filtration)
a magnetic stirrer comes in handy and is highly recommended

Purification of Aspirin:
There are three ways to do this, and it's really up to you which one you like, I prefer the simple decantating method because it's cheap and I don't have HCl [sigh].
<ol type="1"><li>You can dump the <em>uncoated</em> tablets into a flask, add cold distilled water, and swirl until they break apart. Then wait for a short amount of time until the big chunks (the acetylsalicylic acid) settle, leaving the binder suspended--carefully decant and discard the liquid. Repeat as many times as necessary to obtain a reasonably clear solution. (I'll take some pictures of what the solution should look like before you decant and at the end of the procedure.) Then simply filter, rinse with distilled water, and dry. The purified product should be a coarse powder that pours easily and will probably smell of acetic acid due to decomp.--this is normal.</li><li>Dump tablets into flask; add methanol, ethanol, or isopropanol to dissolve the ASA; filter; evaporate filtrate; dispose of filtered solids. This product will form nice clear colorless crystals with little or no odor and should be powdered with mortar and pestle to speed later reaction.</li><li>Dump tablets into flask, add aqueous sodium bicarbonate or sodium carbonate to dissolve the ASA, filter to remove insoluble binders, then bring the pH back down with dil. HCl to precipitate the purified acetylsalicylic acid.</li></ol>Actually, a combination of the third preceded by either of the other two would probably give the best results.

Synthesis:
- Do your own calculations to figure out the quantities needed, but let me say that you need 2 moles of ASA per mole of Cu<sup>2+</sup>. <em>Be sure to use an excess of ASA, or your product will end up being a mixture of Cu<sub>2</sub>(asp.)<sub>4</sub> and copper carbonate!</em> If you need help with calculations, ask. I'll also put a simple 'recipe' at the very bottom of this post.
- Dissolve sodium bicarbonate in distilled water.
- Add weighed ammount of ASA to an Erlenmeyer flask 2 sizes larger than you think you'll need (now's the time to use your magnetic stirrer if you've got one) and dissolve by slowly adding the aqueous sodium bicarbonate solution. Be careful, it will foam a lot if you add too much, stirring and swirling can help to control this.
- Wait for gas evolution to stop (completely! be patient!), then vacuum filter for filtrate. Remember, you don't want <em>all</em> of the ASA to dissolve, but you don't want unreacted NaHCO<sub>3</sub> either.
- Prepare an aqueous solution of copper(II) sulfate while you're waiting (the more concentrated, the smaller the precipitate particle size seems to get, so concentration is really up to you).
- In a <em>beaker</em> (you'll have to scrape the sides to remove product), mix the copper and aspirin solutions slowly (magnetic stirrer here will help to increase crystal size). You should immediately see a beautiful bright blue precipitate form. Slower mixing of the two solutions and lower concentrations seem to generate larger crystals. Continue to stir/swirl for a few minutes. Scrape the precip. from the sides and bottom before filtering.
- Vacuum filter the product, rinse with distilled water (should be extremely easy to filter, aside from the tendancy of some of the product to form a floating film over the solution and inside of the beaker), and dry by allowing the vacuum to pull air through. The precip settles easily, so keep swirling and washing as you pour.
- Copper(II) aspirinate residues can be removed from glassware most conveniently by reaction with dil. HCl followed by alcohol washings.

Outline of alternate synthesis procedures:
<ul type='circle'><li>Prepare CuCO<sub>3</sub> by reaction of CuSO<sub>4</sub> with Na<sub>2</sub>CO<sub>3</sub>, add acetic acid to dissolve some but not all of the precip, forming copper(II) acetate, filter for filtrate, dissolve purified ASA in isopropanol or ethanol, combine the solutions on a magnetic stirrer and allow the crystals to grow, scrape, filter, wash, dry.</li><li>The third procedure involves heating copper(II) salicylate with aspirin in 50% aqueous ethanol. It's in the literature, if you can't find it but want to try, let me know.</li></ul>When dried, the product from the first synthesis is very 'staticy', so be careful not to make a mess. The product from the first is a bright blue powder (see <strong>Fig. 1</strong>;). The product from the alternate synthesis I is small dark blue crystals with a slight greenish tint--I'll get a picture later. I haven't yet tried the second alternate synthesis, so if anyone tries it, post your results.

<table width="450" align="center"><tr><td colspan="2"><img src="http://upload.wikimedia.org/wikipedia/commons/d/d5/Copper_aspirinate.jpg" width="445" /></td></tr><tr><td width="40" valign="top"><strong>Fig. 1.</strong></td><td>Copper(II) aspirinate product under natural light conditions. Prepared by precipitation between sodium acetylsalicylate and copper(II) sulfate.</td></tr></table>

The aluminum salt of aspirin can be prepared by substituting aluminum sulfate for copper(II) sulfate in the first synthesis. Zinc nitrate and cobalt chloride failed to produce a precip. The zinc salt can be obtained by an evaporative method and a cadmium salt can be prepared (I don't know what method), according to lit.

Q. <em>What the hell is copper(II) aspirinate anyways? What's it good for?</em>
A. Well, it's been used as a treatment for rheumatoid arthritis, and it's been investigated for use as a blue pigment in PVC and other related polymers. It has an <a href="http://en.wikipedia.org/wiki/Copper_aspirinate" target="_blank">interesting structure</a> <img src="../scipics/_wiki.png" />, and forms polymer like chains (look into the lit). I put the melting point in the Wikipedia page, along with molecular formula, reaction equations, and all that. It's insoluble in everything I have, and according to the literature it's only sparingly soluble in DMSO--not much else--so recrystallization would be difficult.

Basic Recipe<sup>1</sup>:
-3.60 g aspirin
-2.01 g KHCO<sub>3</sub> in 30 mL H<sub>2</sub>O
-2.50 g CuSO<sub>4</sub> &middot; 5 H<sub>2</sub>O in 20 mL H<sub>2</sub>O
-mix aspirin and potassium bicarbonate solution, slowly combine with copper(II) sulfate solution on magnetic stirrer

<u>References:</u>
1. <em>J Biol Inorg Chem</em> (2005) 10: 831-841. DOI: 10.1007/s00775-005-0031-3

<u>See Also:</u>
Garcia, F., Mendez-Rojas, M. A., Gonzalez-Vergara, E., Bernes, S. & Quiroz, M. A. (2003). <em>Acta Cryst.</em> E59, m1171-m1173. DOI:<a href="http://dx.doi.org/S1600536803026126" target="_blank">10.1107/S1600536803026126</a> <img src="../scipics/_ext.png" />
<em><a href="http://scripturalphysics.org/etc/CopperAspirinate.html" target="_blank">Copper Aspirinate Synthesis</a></em> <img src="../scipics/_ext.png" /> by Brian Fraser, &copy;2003
<em><a href="http://www.youtube.com/watch?v=O0hxp_1w0MQ" target="_blank">Synthesis of Copper aspirinate</a></em> <img src="../scipics/_yt.png" /> by <strong>aonomus</strong>

[Edited on 2/10/08 by bfesser]

[Edited on 7/26/13 by bfesser]
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[*] posted on 8-2-2008 at 17:52


Interesting stuff! Thanks for the adventure. Does this come within the defination of chelate?



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[*] posted on 8-2-2008 at 18:37


Yes, it does. Also, here's the pic of the product prepared by the acetate method, having larger crystal size:

<img src="http://www.medievalsiege.net/bfesser/images/chem/cu2(asp)4_crystals_small.jpg" />

[Edited on 2/8/08 by bfesser]
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[*] posted on 8-2-2008 at 18:59


Very nice, I'd have never thought something so simple forms an interesting complex structure, with additional biological benefits compared to plain aspirin! I guess this theme could be used to start exploring plain salicylic acid, resorcinol (here there is no COOH), terephtalic acid, phtalic acid, and quite a large number of other organic acids..... copper benzoate is of course known, and does to my knowledge not form a complex. It is also used as a blue colouring agent in pyrotechnics, so I imagine the salicylate should be equally suitable.



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[*] posted on 8-2-2008 at 19:24


On a related note, I'm hoping to experiment with polymerizing aspirin and aspirin derivatives when I can scrounge up money to get the reagents. There have been quite a few news stories on chem sites about aspirin-based polymers in the last few years.
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[*] posted on 8-2-2008 at 23:22


I just wanted to let you know that I tried out the experiment and it worked for me.
To extract the HC9H7O4, I first washed out the water-soluble binders from a
ground-up pill of asprin, then added 20 ml EtOH (95%) to the filtrate and evaporated
the supernatant to get 100 mg of flakes. I added 25 mg NaHCO3 to the ground flakes.
Upon adding a few drops of the result to dilute CuSO4, there formed a flocculent
gel of a precipitate. I didn't have any luck with NiCl2, though. It's late now, so
further experimentation will have to wait for another day.
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[*] posted on 9-2-2008 at 01:07


Quote:
Originally posted by chemoleo
Very nice, I'd have never thought something so simple forms an interesting complex structure, with additional biological benefits compared to plain aspirin! I guess this theme could be used to start exploring plain salicylic acid, resorcinol (here there is no COOH), terephtalic acid, phtalic acid, and quite a large number of other organic acids..... copper benzoate is of course known, and does to my knowledge not form a complex. It is also used as a blue colouring agent in pyrotechnics, so I imagine the salicylate should be equally suitable.

The structure is what copper(II) salts usually form with carboxylate anions. Copper (II) acetate has the same structure.
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[*] posted on 9-2-2008 at 06:05


Very nice.. :D

Unfortunatelly I'm away from ASA since I recently wasted mine in TNP production .

I've produced months ago some copper terephthalate.. Simply mixing sodium terephthalate solution with copper sulfate soln. It will form instantaneously a clear blue ppt ..

Unfortunately I think I dont have made a pure material since probably some Na2CO3 left unreacted in starting terephthalate soln..

The powder didnt appeared to work well for flame coloured pyro compositions.. I wish to repeat experiments latter..

I'm wondering how good would be this copper aspirinate for pyro purposes..

[Edited on 9-2-2008 by Aqua_Fortis_100%]




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[*] posted on 9-2-2008 at 06:40


I can't say I approve of referring to salts of acetylsalicylic acid as aspirinates.

Aspirin is a chemically meaningless trade name (Bayer) gone generic.

There is no aspirinic acid.

Call it what it is: Copper (II) acetylsalicylate. I can't help it if that does not fall trippingly off the tongue.




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[*] posted on 9-2-2008 at 11:03


Quote:
Upon adding a few drops of the result to dilute CuSO4, there formed a flocculent gel of a precipitate.

-A gel? Sounds like something went wrong.

Do you guys try to burn everything?

Quote:
I can't say I approve of referring to salts of acetylsalicylic acid as aspirinates.

-I don't really care for the name either, but that 's what it's referred to in most of the literature. Personally, I'd prefer dicopper 2-acetyloxybenzoate (IUPAC) or just copper(II) acetylsalicylate as you've said.
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[*] posted on 9-2-2008 at 11:57


Quote:

-A gel? Sounds like something went wrong.


Given that my solution of sodium acetylsalicilate was dense and denser
solutions give smaller precipitate, that might have been what happened.
It is also possible that I did not wait long enough for all the bicarbonate to
react (It was getting quite late) or that some other component of the asprin
remained and got in the way. To figure out what is going on, I will do the
experiment more carefully, starting with reprecipitating the HC9H7O4 using
HCl. I will get back to you all later when I do so in a few days.

Update: I had a look at the vial which had the flocculent precipitate --- it
went away of its own accord and I am now left with solid precipitate which
looks a lot like what you show in your picture. I also bought a bottle of
asprin from the dollar store for use in future experiments. Time to start
extracting and cleaning it up.

[Edited on 9-2-2008 by microcosmicus]
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[*] posted on 9-2-2008 at 19:32


I got good results doing this, on a small scale, I am not trying to save the products, at least not this time. It did take a little patience to react all the HCO3-, my ASA was in fluffy xtals so there were a lot of little bubbles. After filtering, and adding Cu(II)SO4, I got a nice bluish ppt.

I tried also with Ni(II)SO4 and PdCl2/NaCl, and got no ppt or apparent reaction with either, other than seemingly some ASA ppt, a few crystals seem to have appeared in the otherwise clear solution.
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[*] posted on 10-2-2008 at 01:40


Very interesting posts. I noted that there was no precipitate with Nickel Sulfate, which may allow separation of the two metals with easily available ingredients.
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[*] posted on 10-2-2008 at 09:26


Quote:
Originally posted by Mr. Wizard
Very interesting posts. I noted that there was no precipitate with Nickel Sulfate, which may allow separation of the two metals with easily available ingredients.


That is true. Simply adding dil. HCl will free the Cu from the chelate forming a white acetylsalicylic acid solid and aqueous CuCl<sub>2</sub>. The ASA could probably even be reused a few times.
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[*] posted on 10-2-2008 at 11:11


This also sounds like it might be useful for analysis. Since the precipitate seems to be so
insoluble in most solvents, one should be able to get nearly all the Cu out of a solution and wash the precipitate well. (By the way, does someone happen to have a numerical value
the solubility of copper acetylsalicylate?) As a precipitating agent, ASA has the advantage
over H2S that it is much better for your health and does not make your experiment stink
like a rotten egg. Since the molecule is so darned big (m.w. 844) and the copper accounts
for only 15% of the total weight, this should be a good form for weighing copper.
Alternatively, maybe titrate the precipitate away with HCl.
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[*] posted on 10-2-2008 at 12:18


I haven't yet been able to find any numerical value for the solubility. It's not in the Merck or any other reference material I have access to, and I haven't been able to find any mention of it in any journal articles.

It <em>might</em> have some value for us amateurs as a readily available and relatively non-toxic gravimetric means for copper determination. But there are definitely better means for analytical labs. <em>If</em> someone were to use it, however, I would recommend using a Gooch crucible for the filtration and drying, as <em>any</em> transfers would result in a significant loss of the precip.--as I said, it sticks to everything! It might also be a little difficult to transfer all the precip. from the reaction glassware into the filtration glassware. Perhaps a <a href="http://kimble-kontes.com/html/pg-292900-Filtration.html">one-piece unit</a> would be in order? Again, forming the precip. slowly and in larger crystals would help.

(My personal favorite means for copper determination is spectrophotometric titration with EDTA and a Spec 20.)

Oh, and I almost forgot to mention; don't store sodium acetylsalicylate solutions. They tend to turn pink and decompose.

[edit]
Process to separate aqueous Co/Ni from Cu and recover as chlorides:
-prepare sodium acetylsalicylate (excess acetylsalicylic acid) and filter
-add an excess of the acetylsalicylate solution to the metals solution to precipitate copper as Cu<sub>2</sub>(asp.)<sub>4</sub> and filter with H<sub>2</sub>O rinse
-dry precip and redissolve Cu in cold dil. HCl, filter to remove ASA, and evaporate filtrate to recover CuCl<sub>2</sub>
-add aqueous sodium carbonate to Co/Ni filtrate and cool to precipitate
-filter for solid, rinse, dry, redissolve in dil. HCl, evaporate to recover

How does that procedure sound? Should result in relatively pure chlorides with low sodium contamination depending on the quality of your HCl.

[Edited on 2/10/08 by bfesser]
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[*] posted on 10-2-2008 at 13:54


Sweet post. Refreshing synth! Keep it up..



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[*] posted on 10-2-2008 at 18:11


Here's a cool image illustrating the polymer chains formed by the compound. Sorry for not including it before, it took me a while to find the source I originally got it from.

<img src="http://www.medievalsiege.net/bfesser/images/chem/cu_asp_poly.gif" />
Quote:
<strong>Figure 2</strong>
The crystal packing of (I). One polymeric chain is shown, with Cu atoms represented by spheres of arbitrary dimensions. H atoms have been omitted for clarity.

Taken from:
Garcia, F., Mendez-Rojas, M. A., Gonzalez-Vergara, E., Bernes, S. & Quiroz, M. A. (2003). <em>Acta Cryst.</em> E59, m1171-m1173.
[ doi:<a href='http://scripts.iucr.org/cgi-bin/paper?S1600536803026126'>10.1107/S1600536803026126</a> ]
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[*] posted on 11-2-2008 at 05:11


I retried the experiment at school.
Cobalt and nickel indeed do not form a precipitate (well I had some , but that was because the sodium acetylsalicylate solution was slightly contaminated with carbonate.).
I also tried iron(II) (as Mohr salt), iron(III) (as nitrate), chromium(III) (as nitrate), lead (as nitrate), manganese(II) (as sulphate). They all gave copious ammount of precipitate, so those do not interfere (unless any of the used metals has a carbonate wich has a huge volume, then one should retry the experiment). Only chromium(III), didnt have much precipitate, and after adding the sodium acetylsalicylate solution, the precipitate formed after like 10 seconds (like cobalt and nickel), while the other metals form instantanouesly. So maybe it was only chromiumcarbonate?
Can someone please retry it? I dont have acetylsalicylic acid at home.
I forgot to test zinc.
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[*] posted on 11-2-2008 at 06:17


Quote:
I dont have acetylsalicylic acid at home.


You must have Aspirin tablets at home, right? The procedure for purifying these is outlined in the first post of this thread. Granted it's not analytical grade, but it's more than pure enough for testing precipitate formation with transition metal salts.
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[*] posted on 11-2-2008 at 08:38


We dont have aspirine here no :P
And when I buy it I'd rather buy it as a reagent.
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[*] posted on 15-2-2008 at 11:00


I just realized that I neglected to include something in my earlier posts . . .

There is one more way to synthesize copper(II) acetylsalicylate found in the literature, and it claims to afford the largest and purest crystals. It's an electrochemical method from the same article as the polymer image above. It'd be an interesting experiment to conduct for someone who can afford the reagents and equipment [hint hint].

Here's the excerpt of the synthesis (is it legal to post this much of an article? if not, please edit this out):
Quote:
Electrosynthesis experiments were performed in a conventional
three-electrode cell with a Pt (99.999%, Aldrich) wire as auxiliary
electrode (cathode), a saturated calomel electrode (SCE) as reference
and a Cu (97%, Merck) foil as working electrode (sacri®cial
anode). The Cu foil was cleaned with dilute hydrochloric acid,
washed thoroughly with threefold-distilled water and dried at 383 K
before use. Acetonitrile (HPLC grade, Caledon), the reaction
medium, was dried over molecular sieves and the supporting electrolyte
was generated in situ. A small amount (ca 0.04 mmol) of the
Na salt of the carboxylate ligand was generated by dissolving the
respective acid (1.0 mmol) in dry acetonitrile (30 ml) and adding
fresh metallic Na (ca 1.0 mg, 0.04 mmol, cleaned with dry ether).
After 5 min of stirring at 298 K, the solid residues were separated by
®ltration. The resulting solution, containing the deprotonated ligand
(ca 0.04 mmol) and the precursor acid (ca 0.96 mmol), was then used
to ®ll the electrochemical cell. The electrolysis experiments were
carried out at a constant current of 10 mA, which was imposed for a
period of 180 min by using an EG&G PARC potentiostat-Galvanostat,
model 362, affording (I) in 89±90% yield.


Source:
Garcia, F., Mendez-Rojas, M. A., Gonzalez-Vergara, E., Bernes, S. & Quiroz, M. A. (2003). <em>Acta Cryst.</em> E59, m1171-m1173.
[ doi:<a href='http://scripts.iucr.org/cgi-bin/paper?S1600536803026126'>10.1107/S1600536803026126</a> ]

[Edited on 2/15/08 by bfesser]




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[*] posted on 16-2-2008 at 17:47


In the first post for alternative synthesis number 1 you say that a precipitate of copper acetate forms as acetic acid neutralizes the carbonate. I have to admit I didn't do it this way but instead I used a solution of sodium acetate and added it to a solution of Copper(II)Sulfate expecting the precipitate, this, unfortunately, never came to be.
I'm considering that perhaps my solution of copper sulfate was too concentrated, do you know if this affects solubility (or have you experienced it at least)?
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[*] posted on 16-2-2008 at 17:53


Copper(II) acetate precipitate should never form, it's highly soluble in water. As the acetic acid reacts with the carbonate it dissolves the copper ions to form copper(II) acetate, water, and CO<sub>2</sub> gas.

Likewise, a solution of sodium acetate and copper(II) sulfate should not form a precipitate. The concentrations of the solutions won't affect the solubility unless they're super saturated or something, in which case they'd just crystallize out when disturbed.

[edit 1]

I also should have added something to the first alternate synthesis procedure. If a solution of acetylsalicylic acid in isopropanol is used, and the isopropanol is too concentrated, it can actually drive the copper sulfate right out of solution forming a light blue crystalline solid. To test for this precipitate and make sure you have authentic copper(II) acetylsalicylate and not copper(II) sulfate, just put a small amount of the solid in a test tube and add distilled water. If the solid dissolves in water, it's copper(II) sulfate and your alcohol solution was too concentrated.

[edit 2]
Sorry to edit this again, but I forgot to post about something I thought of last night. If you think the copper(II) acetylsalicylate product is contaminated with copper(II) carbonate, maybe a wash with fairly dilute acetic acid would be sufficient to dissolve any CuCO<sub>3</sub>, but the acid shouldn't be strong enough to dissolve the copper(II) acetylsalicylate. Anyone have a second opinion on this?

I've got a bottle of CuCO<sub>3</sub> that I prepared a year or two ago, and it seems like the longer I let it sit, the less of it wants to dissolve when I put some in 1<u>M</u> acetic acid.

[Edited on 2/16/08 by bfesser]




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[*] posted on 16-2-2008 at 18:12


Ah I see, it's mostly a disconsideration on my side my apologies. I was meaning to get some copper acetate but I'm thinking I better go the acid-base way since it's a fairly easy reaction and affords relatively pure product. As for the copper sulfate and sodium acetate solution, I hope I can make any use out of it now (perhaps crystallization but, damn, I can't hold of any solubility curves! :()
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