Sciencemadness Discussion Board - Copper(II) Aspirinate Synthesis

Copper(II) Aspirinate Synthesis

Pages: 1 2

bfesser - 8-2-2008 at 12:48

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

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 uncoated 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 Cu2+. Be sure to use an excess of ASA, or your product will end up being a mixture of Cu2(asp.)4 and copper carbonate! 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 all of the ASA to dissolve, but you don't want unreacted NaHCO3 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 beaker (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 CuCO3 by reaction of CuSO4 with Na2CO3, 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 Fig. 1

. 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">Fig. 1.</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. What the hell is copper(II) aspirinate anyways? What's it good for?
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 Recipe1:
-3.60 g aspirin
-2.01 g KHCO3 in 30 mL H2O
-2.50 g CuSO4 · 5 H2O in 20 mL H2O
-mix aspirin and potassium bicarbonate solution, slowly combine with copper(II) sulfate solution on magnetic stirrer

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

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

[Edited on 2/10/08 by bfesser]

[Edited on 7/26/13 by bfesser]

chemrox - 8-2-2008 at 17:52

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

bfesser - 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]

chemoleo - 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.

bfesser - 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.

microcosmicus - 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.

guy - 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.

Aqua_Fortis_100% - 9-2-2008 at 06:05

Very nice..

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%]

Sauron - 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.

bfesser - 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.

microcosmicus - 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]

pantone159 - 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.

Mr. Wizard - 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.

bfesser - 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 CuCl2. The ASA could probably even be reused a few times.

microcosmicus - 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.

bfesser - 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 might 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. If someone were to use it, however, I would recommend using a Gooch crucible for the filtration and drying, as any 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 Cu2(asp.)4 and filter with H2O rinse
-dry precip and redissolve Cu in cold dil. HCl, filter to remove ASA, and evaporate filtrate to recover CuCl2
-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]

Waffles - 10-2-2008 at 13:54

Sweet post. Refreshing synth! Keep it up..

bfesser - 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:

Figure 2
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). Acta Cryst. E59, m1171-m1173.
[ doi:<a href='http://scripts.iucr.org/cgi-bin/paper?S1600536803026126'>10.1107/S1600536803026126</a> ]

Jor - 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.

bfesser - 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.

Jor - 11-2-2008 at 08:38

We dont have aspirine here no

And when I buy it I'd rather buy it as a reagent.

bfesser - 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). Acta Cryst. 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]

16MillionEyes - 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)?

bfesser - 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 CO2 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 CuCO3, 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 CuCO3 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 1M acetic acid.

[Edited on 2/16/08 by bfesser]

16MillionEyes - 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!

)

bfesser - 16-2-2008 at 18:19

If you add baking soda, you can precipitate the copper out as the carbonate. But be careful, it will foam a lot. Then you can filter it, rinse it, dry it, and very carefully and slowly dissolve it in dil. H2SO4. Add the sulfuric acid dropwise with lots of swirling. If you can do this carefully enough to dissolve most of the solid, but just leave a little bit of solid remaining, you can filter it out and evaporate the filtered liquid down to copper(II) sulfate again.

As for the sodium acetate, you might as well just dump it. It would probably be more expensive to recover it than to just dump it. You can always make more with vinegar and baking soda.

It's also important to note that copper solutions are considered to be harmful to the environment, and you should avoid pouring such solutions down the drain at all costs. I always precipitate my copper out and throw it away if I end up having any in my filtrates from various syntheses. Sodium acetate, on the other hand, is pretty much harmless.

[Edited on 2/16/08 by bfesser]

16MillionEyes - 16-2-2008 at 19:05

I don't have any sulfuric acid (in reality I have but it's a yellow solution from drain cleaner, almost useless) but my main concern would be to get the copper out so I'm thinking that I can precipitate it as the hydroxide and filter it. I believe this way I could also get all my copper out and throw away the rest.
On the other hand, I dissolved some of this solution in 70% isopropanol and much to my surprise a clean green-blue precipitate forms. Although I believe it to be the acetate I'm not very confident of this, however, I think this could also be an alternative way to get my copper out of solution. Any ideas other ideas of what the precipitate may be? The original solution is only copper(II) sulfate and sodium acetate (a bit of excess acetic acid) in which the only green-bluish precipitate would the copper acetate as far as I know.

bfesser - 16-2-2008 at 20:21

If you read the revision of my last post, I mentioned that copper(II) sulfate will precipitate out of an aqueous solution of isopropanol, provided the isopropanol is in high enough concentration. Not sure if the same would be true of copper(II) acetate. I suppose you could always dissolve a little of the precip in water, then add a little of a barium salt, if you have the hydroxide or chloride, as a qualitative test for sulfate (forms a white precipitate). Anyone else have any ideas?

16MillionEyes - 16-2-2008 at 20:38

I did read it but it's exactly why I considered the green-blue precipitate to be the acetate and not the sulfate: the isopropanol solution remains blue (as I mentioned it's 70%, "significantly" wet to not be considered highly concentrated). I think I'll post some pictures on the precipitate tomorrow so that you may compare and see what I mean.

chemkid - 28-3-2008 at 08:38

At the bottom 'Basic Recipe' it says to use KHCO3 (pottasium bicarbonate) could NaCO3 (sodium carbonate) be substituted?

Nerro - 28-3-2008 at 11:02

Chelation therapy is looked at as a possible way to prevent Alzheimers desease. Could this mean that aspirin might help prevent Alzheimers?

bfesser - 28-3-2008 at 15:25

Sodium carbonate should work equally well. Please post your results. I know I've tried this before, but I can't remember if I had any issue with hydrolysis of the acetylsalicylic acid. I'll retry it when I get a chance.

Both copper and aluminum will form insoluble complexes with acetylsalicylic acid. I do recall seeing somewhere that someone's found a correlation between aluminum concentration and Alzheimers in the brain--I think. Other than that, I have no idea. Can anyone find any research on this subject? I haven't seen any yet.

Ozonelabs - 28-4-2008 at 03:06

Hello everyone,

Apologies for digging up an old thread, but there is now a Youtube vid with pictures and method posted on this synthesis and it is on our website.

Youtube link: http://youtube.com/watch?v=PNJfWYKcjKY

Website link: http://www.freewebs.com/ozonelabsinc/syntheses.htm

If there are any questions/comments please get back to us!

Regards,

Ozonelabs

bfesser - 9-6-2008 at 12:48

Have you checked the identity/purity of your product? The color seems a bit off, even for a fine precipitate--although it could just be the camera. Also, it shouldn't be difficult to vacuum filter. What grade of filter paper did you use?

To test the identity of your compound, just add a little vinegar to a sample. If it dissolves, it's likely to be copper(II) carbonate. Adding a little HCl to another sample should leave a white precip of ASA and a blue CuCl2 solution.

[Edited on 6/9/08 by bfesser]

indigofuzzy - 1-3-2009 at 16:49

I tried synthesizing some Copper Acetylsalicylate two days ago. There seems to be little need to react the ASA with NaHCO3. I mixed ASA with Copper(II) Carbonate, and added a few drops of water. The slurry began to bubble slowly and over a few minutes changed from pale green to a saturated bluish-aqua color.

Upon drying (by evaporation at room temperature), the residual powder is a deep aqua color. The trace of green may be due to residual CuCO3, or residual insoluble component in the aspirin tablets. I'll look into this later tonight.

Cobalt, nickel, and other metal aspirinates

sargent1015 - 20-6-2013 at 20:13

First off, I want to say it feels good to be back on ScienceMadness, being a chem major and the research takes a serious toll on your freetime. But it is good to see you all, especially my old buddy Hexavalent.

Also, I think this belongs in beginnings?? But, it has been a while and maybe rules have changed.

Anyways, down to some inorganic complexes. I have seen a few threads and posts on copper aspirinate, which is good and dandy, but is acetylsalicylic acid a strong enough field ligand to bind any number of transition metals? Say, for example nickel or cobalt? Maybe even iron? Or is copper just the special oddball (Like usual)?

Well anyways, this is largely for discussion, since I'm going to peruse the literature to see what I come up with. I will update you on the progress of this.

Also, for reference, here is the wikipedia link to copper aspirinate to show the dimer structure. Wiki
I also like the synthetic approach the wiki gives to synthesizing the complex.

Lambda-Eyde - 21-6-2013 at 00:17

Interesting. I might give cobalt and nickel a shot with salicylic acid later this weekend. I have nickel sulfate, nickel chloride and cobalt chloride at my lab. Off the top of my head I can't think of any other colorful metal salts that I have lying around. I have no aspirin available. I'll report my findings.

Edit: Cobalt nitrate

[Edited on 21-6-2013 by Lambda-Eyde]

sargent1015 - 21-6-2013 at 05:12

I found a zinc complex that had been studied, so this sounds like a possibilty. However, they have no method for synthesis in this paper.

Anne Spasojević-de Biré, Nouzha Bouhmaida, Aleksandar Kremenović, Georges Morgant, and Nour Eddine Ghermani. Experimental Electron Density and Electrostatic Potential Analysis of Zinc(aspirinate)2(H2O)2 Complex: A 3d10 Metal Bonding to a Drug Ligand. The Journal of Physical Chemistry. November 22, 2002, pp 12170–12177.

I will do some aspirin extractions tonight and get enough material to study this over the weekend.

[Edited on 21-6-2013 by sargent1015]

sargent1015 - 21-6-2013 at 05:30

This is a pretty neat article too, salicylic acid complexes of copper are possible! Which can be a problem for the amateur chemist that overcooks his acetylsalicylic acid reaction and clams that is the complex he has obtained.

Orally active antioxidative copper(II) aspirinate: synthesis, structure characterization, superoxide scavenging activity, and in vitro and in vivo antioxidative evaluations. Fujimori, T. ; Yamada, S. ; Yasui, H. ; Sakurai, H. ; In, Y. ; Ishida, T.
JBIC Journal of Biological Inorganic Chemistry, 2005, Vol.10(8), pp.831-841

I guess this is no longer a beginnings post

[Edited on 21-6-2013 by sargent1015]

plante1999 - 21-6-2013 at 05:33

Welcome back. I have the feeling that iron won't make a stable complex, cobalt, copper, will and I don't know for nickel. For sure it is not a highly coordinating complexing agent. Iron generally does not complex well.

If one had uranyl acetate, I think it might make a pretty colorful complex.

sargent1015 - 21-6-2013 at 05:37

Hey Plante, thanks for welcoming me back

In regards to your iron statement, and of course not to be rude or undermining, I just found an article on its synthesis

Synthesis and magnetic properties of the iron(III) aspirinate complex, [Fe 3O(C 9H 7O 4) 6(H 2O) 3](C 9H 7O 4)·H2O. Lynch, S.K. ; Frazder, D.R. ; Carlisle, G.O.
Journal of Molecular Structure, 1982, Vol.82(1), pp.151-153.

Fairly old one at that, not the medicinal activity of copper aspirinate, not surprisingly.

7 moles of "aspirinate" to 1 mole iron with a water molecule complexed.

[Edited on 21-6-2013 by sargent1015]

plante1999 - 21-6-2013 at 06:29

Very interesting, but the question is it stable?

Hmm, I'll have to check more into this.

cyanureeves - 21-6-2013 at 07:01

yes this amateur chemist did over cooked the aspirin because my copper aspirinate looks like copper carbonate and isnt dark at all. i never gave up on wanting to do it again even though my Salicylic acid was very fluffy and crystally i never knew where i went wrong. i only have coffee warmers and not a regulated heat source like a heated stirer so i have to switch back and forth from heat to stirer.

bfesser - 21-6-2013 at 09:18

Quote: Originally posted by sargent1015

Also, for reference, here is the wikipedia link to copper aspirinate to show the dimer structure. Wiki
I also like the synthetic approach the wiki gives to synthesizing the complex.

I'm glad you like the synthetic approach, I wrote it.

I've done a fairly extensive literature search over the years on copper(II) aspirinate and related complexes. Perhaps, if you're interested, I'll dig out my file folder (yes, hard copy!) later and post what I have found. I also have performed experiments with several other cations, and have notes on this. I'll have to look those up, as well. They probably make up a third of my first chemistry laboratory notebook. Sorry, I don't recall precisely which complexes I managed to prepare (cobalt?).

[edit]
From what I remember, I tried copper, zinc, iron, cobalt, and magnesium.

[Edited on 7/8/13 by bfesser]

sargent1015 - 21-6-2013 at 09:33

Quote: Originally posted by bfesser

I'm glad you like the synthetic approach, I wrote it.

Well what do ya know!? Hahaha, thanks for looking all this up. I'd love to see your notes on the subject and hopefully repeat/expand on your studies. Hopefully I get some decent results.

The key procedure I'm seeing cited for these complexes is using a weakly basic buffer (K/Na HCO₃) and adding the aspirin and corresponding salt. Seems simple enough, though I only found iron and copper complexes in the literature. These should be wonderfully colored compounds (Minus iron, unless you like brown

bfesser - 21-6-2013 at 09:36

Quote: Originally posted by bfesser

(I'll take some pictures of what the solution should look like before you decant and at the end of the procedure.)

I just noticed that I promised photos five years ago and never delivered. Sorry about that. Maybe some day...

Regardless, repeat washing until the supernatant is clear.

[Edited on 7/8/13 by bfesser]

sargent1015 - 22-6-2013 at 10:43

Alright, so as far as forum etiquette as I work out some different complexes, should I form a new thread describing the syntheses or just continue here?

Also, I am currently making the iron and copper complexes, in case anyone was curious

bfesser - 22-6-2013 at 15:56

Personally, I would like to see them all in one 'aspirinate' thread.

sargent1015 - 22-6-2013 at 18:58

Well, for now, just updates on my progress:

Copper aspirinate made (easy)
Iron aspirinate made, but the solid particulate is too fine to collect with coffee filter paper. Going to try to get my hands on good filter paper.

To be done:
Cobalt aspirinate
Salicylic acid complexes
Copper mixed aspirinates and imidazole complexes (different ratios)

cyanureeves - 22-6-2013 at 19:19

i wonder if chromium/cobalt alloy will work as well because i have dental metal alloy that i believe is chromium/cobalt.it dissolved rather quickly in hot hydrochloric and turned light purple when i added sodium carbonate.btw this alloy is hard as shit to work with i mean it melts but it doesnt run and the glob will let the flame go through.the flame goes through the glowing mass like the exterminator's fist going through the mercury man's face. correction:the dissolved alloy was bluegreen and only after i added ammonium hydroxide did it turn indigo.i wanted to know if it made a blue complex as with copper

[Edited on 6-23-2013 by cyanureeves]

sargent1015 - 22-6-2013 at 22:45

That is interesting, I don't know anything about alloy complex chemistry, are there some papers on the subject? I would love to understand a little bit more about this.

Mixell - 23-6-2013 at 09:01

In those conditions the products of the dissolved alloy should act independently of each other. On some occasions different metal cations can form a an insoluble compound in which one of the cations is present as a charged coordinated complex (such as CuCl3 (-)/CuCl4 (2-)/CoCl4 (2-). Cations such as Cs+/Rb+ can be used to precipitate those complexes.

There are also mixed-valence compounds in which different cations (of the same or different element) are bridge by charge-transfer ligands, those compounds can sometimes have a low charge-transfer activation energy, resulting in absorption in the visible spectrum. A good example of that is the pigment"prussian blue".

There are also mixed oxides such as ilmenite and magnetite.

sargent1015 - 23-6-2013 at 14:09

More chemistry I would like to try:

Benzoate complexes
Imidazole complexes
Some amine complexes

I am an organic chemist in the lab, but at home I seem to have more of the inorganic mindset

I am also thinking of putting a pre-pub of these syntheses onto the site, heavily citing bfessor's previous work.

sargent1015 - 24-6-2013 at 11:50

Just took an NMR of the aspirin that I isolated over the weekend and it is beautifully pure. I will try to attach a picture of the spectra tonight.

Good to know my material is of high quality before going on to the next steps

DraconicAcid - 24-6-2013 at 13:52

One of my intended projects is making some anthranilates, and seeing if I can get crystals of them. I've got microcrystalline copper(II) anthranilate, and will try nickel, cobalt, iron, and possibly chromium.

sargent1015 - 24-6-2013 at 19:02

Oooo do that! If I may ask, where did you get the anthranilate?

DraconicAcid - 25-6-2013 at 07:30

We ordered anthranilic acid for the college shortly before it became listed as a possible drug precursor. Whether or not we'll ever get more remains to be seen.

[Edited on 25-6-2013 by DraconicAcid]

sargent1015 - 25-6-2013 at 08:21

Talk about lucky snag. That amine makes it an excellent ligand. And if it is the free acid version, you could even make it behave bidentate!

Update, iron aspirinate particulate is incredibly fine and difficult to filter

nezza - 26-6-2013 at 00:14

I haven't isolated any aspirinates, but came across an interesting colour reaction with cobalt salicylate complex. Cobalt complexes with various amines quite well and in alkaline solution cobalt III is readily formed by adding hydrogen peroxide. These complexes are usually blue to dark blue. On re-acidifying the EDTA complex it goes a dark magenta colour, much darker than the cobalt II ion. With salicylate and alkaline peroxide a dark green colour is produced which slowly (over several hours) goes yellow-brown. Green is unusual for cobalt so is this contamination ?.

DraconicAcid - 26-6-2013 at 10:02

Quote: Originally posted by nezza

I haven't isolated any aspirinates, but came across an interesting colour reaction with cobalt salicylate complex. Cobalt complexes with various amines quite well and in alkaline solution cobalt III is readily formed by adding hydrogen peroxide. These complexes are usually blue to dark blue. On re-acidifying the EDTA complex it goes a dark magenta colour, much darker than the cobalt II ion. With salicylate and alkaline peroxide a dark green colour is produced which slowly (over several hours) goes yellow-brown. Green is unusual for cobalt so is this contamination ?.

I think the tris(carbonato)cobaltate(III) complex is green, so I wouldn't be surprised if tris(salicylato)cobaltate(III) was also green. Then again, iron(III) turns salicylate purple, but decomposes, so I would also not be surprised if cobalt(III) will also oxidize and decompose a phenol such as salicylate.

This paper gives a synthesis of both the carbonato and salicylato complexes: http://pubs.acs.org/doi/abs/10.1021/ja01504a004
I don't have access to the full paper.

sargent1015 - 26-6-2013 at 10:06

Here's that paper.

Attachment: ja01504a004.pdf (287kB)
This file has been downloaded 923 times

I think I will attempt the Tris-carbonatocobaltate(III) Trihydrate tonight if time allows

[Edited on 26-6-2013 by sargent1015]

DraconicAcid - 26-6-2013 at 10:29

Quote: Originally posted by sargent1015

Here's that paper.

I think I will attempt the Tris-carbonatocobaltate(III) Trihydrate tonight if time allows

[Edited on 26-6-2013 by sargent1015]

Thanks for posting that. Now I'm wondering if cobalt(III) aspirinate and/or anthranilate would be stable

[Edited on 26-6-2013 by DraconicAcid]

nezza - 26-6-2013 at 23:58

Thanks for the post folks. I'll look into these complexes a bit more.

sargent1015 - 27-6-2013 at 05:05

Quote: Originally posted by DraconicAcid

Thanks for posting that. Now I'm wondering if cobalt(III) aspirinate and/or anthranilate would be stable
[Edited on 26-6-2013 by DraconicAcid]

Hmmm, that is an interesting idea. Let me know if you give it a shot. Might be a fun test tube reaction to run

Number 9 - 27-6-2013 at 08:25

Quote:

Quote: Originally posted by plante1999

I have the feeling that iron won't make a stable complex

Fe(ll)acetylsalicylate could be formed. You need a solution of FeCl3 and salicylic acid. Then reduce it with hydroxylamine. The slightly yellow colour is transformed in a orange solution. By spectroscopic determination you can estimate the Fe - acetylsalicylate molar ratio. It's proven, I have the documents.

[Edited on 27-6-2013 by Number 9]

Acidum - 27-6-2013 at 11:43

Very nice, interesting and inspiring articles! (referring to this one and Cu(II) related)

I like the methodical approach that You follow, puts science back in the title. Looking forward to see further results!

And yes, will try some of these as soon as I finish reorganizing my lab and purifying chemicals...

PS
Call me grammar/iupac/nerdy nazi, but I just can not look at titles and swallow "aspirinate" as proper name, especially from someone with master in chemistry who actually does research. And referring to ASA as "aspirin" as well. I would definitely be happier if You (or admin/mod) could rename threads at least to "acetylsalicylates (aspirinates)".
In Serbia, acetylsalicylic acid is more often sold as "Andol", "Anbol" and "Midol". If I put my results here, should I refer to them as "andolates, anbolates and midolates"? Or "polyandolate"...

PPS. Yes, I really have a grudge against pharmacy, and I think of their tradenames in my belovin' chemistry as of utter blasphemy! :mad:

sargent1015 - 27-6-2013 at 15:43

Quote: Originally posted by Acidum

This is in direct quote from the literature. Sooooo.... we will have to beg to differ in terms of naming (for some reason I am quite fond of the name aspirinates, probably since I have been messing with them for a week now

)

[edit] this is in reference to "copper aspirinates" almost exclusively, since most of the other metals are named as acetylsalicylates.

[Edited on 27-6-2013 by sargent1015]

sargent1015 - 27-6-2013 at 19:16

Synthetic updates!

The green cobalt complex:
Sodium Tris-carbonatocobaltate (III) Trihydrate
(The procedure is given in the above paper: Bauer & Drinkard, 1960).
This is cool since most cobalt III compounds are unstable

The brown stuff is the iron aspirinate, not great for gravity filtration setup due to the fine particle size. The procedure is given in Lynch, et. al. 1982.

[Edited on 28-6-2013 by sargent1015]

DraconicAcid - 28-6-2013 at 09:30

Very cool!

sargent1015 - 28-6-2013 at 09:45

Could one of the moderators move this to "Chemistry in General"? or should I make a thread for each class of compound (aspirinates, salicylates, cobaltates, etc.)?

Well, we proved one thing, cobalt has an amazing array of colored complexes

I'm excited to try the acetylsalicylate and salicylate complexes as well.

So there was an above post asking if one could make cobalt (III) aspirinate and if it was stable, that is my goal for tonight

DraconicAcid - 28-6-2013 at 10:16

Quote: Originally posted by sargent1015

Well, we proved one thing, cobalt has an amazing array of colored complexes. I'm excited to try the acetylsalicylate and salicylate complexes as well.

We knew that already. Cobalt(III) is great- its complexes (when they don't decompose due to internal oxidation) are quite stable, and colourful. If you've got ethylenediamine, there is a whole raft of cool complexes to be made (one favourite is [Co(en)2Cl2]Cl. The cis and trans isomers are dramatically different colours- bright green and dark purple. If the green looks cruddy, mix the solution with sodium nitrate to give insoluble salt, which is always pure and pretty (the other isomer's nitrate is soluble). The potassium hexanitrito complex is also simple to make if you have potassium nitrite- it's a lovely yellow, and its synthesis is often used as a test for cobalt in qualitative analysis.

Quote:

So there was an above post asking if one could make cobalt (III) aspirinate and if it was stable, that is my goal for tonight.

Good luck!

sargent1015 - 28-6-2013 at 12:09

Quote: Originally posted by DraconicAcid

If you've got ethylenediamine, there is a whole raft of cool complexes to be made (one favourite is [Co(en)2Cl2]Cl.

Is there a convenient source of this? I have not checked ebay at the time of this posting, but I have always wanted to get my hands on that compound. Fun to work with, but holy crap smelly lol

DraconicAcid - 28-6-2013 at 12:24

None that I know of. Ammonia forms the same kinds of complexes, but it's harder to control the number of amine ligands that attach.

[<a href="u2u.php?action=send&username=bfesser">bfesser</a>: removed unnecessary quoting]

[Edited on 7/8/13 by bfesser]

sargent1015 - 28-6-2013 at 12:58

Oh good ol' ammonia. I will let you know if i find some (en)

Also, as promised, NMR's of the aspirin!

Attachment: Aspirin 1.pdf (98kB)
This file has been downloaded 640 times

Attachment: Aspirin 2.pdf (83kB)
This file has been downloaded 596 times

[<a href="u2u.php?action=send&username=bfesser">bfesser</a>: removed unnecessary quoting]

[Edited on 7/8/13 by bfesser]

sargent1015 - 28-6-2013 at 21:01

Well, I actually found some ethylenediamine suppliers:
1) HMS-Beagle
http://stores.intuitwebsites.com/beaglewebstore/-strse-1411/...
Just under a different name, 1,2-ethanediamine

2) Ebay as the dihydrochloride salt.
http://www.ebay.com/itm/Ethylenediamine-dihydrochloride-reag...

nezza - 29-6-2013 at 01:57

Here's a picture of various cobalt (III) complexes I have prepared in solution.

sargent1015 - 29-6-2013 at 07:22

And that is beautiful! Where did you obtain your ethylenediamine?

bfesser - 30-6-2013 at 07:24

Quote: Originally posted by Acidum

Call me grammar/iupac/nerdy nazi, but I just can not look at titles and swallow "aspirinate" as proper name, especially from someone with master in chemistry who actually does research. And referring to ASA as "aspirin" as well. I would definitely be happier if You (or admin/mod) could rename threads at least to "acetylsalicylates (aspirinates)".
In Serbia, acetylsalicylic acid is more often sold as "Andol", "Anbol" and "Midol". If I put my results here, should I refer to them as "andolates, anbolates and midolates"? Or "polyandolate"...

We covered this in the first thread on the topic:

Quote: Originally posted by Sauron

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.

Quote: Originally posted by bfesser

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.

sargent1015 - 30-6-2013 at 08:23

Hehe someone always has to question the naming

sargent1015 - 30-6-2013 at 20:30

Well, more updates, I attempted the Cobalt (II) aspirinate and the Cobalt (III) aspirinate syntheses and both behaved poorly.

The cobalt (II) never crashed out of solution upon cooling, as per the paper given above. I will try boiling off the solvent, but I fear that may decompose the product, but it can't hurt on the small scale I ran.

The cobalt (III) was actually my carelessness that caused it to fail, forgetting to fully dissolve the the acetylsalicylic acid and the sodium bicarb BEFORE throwing it into the cobalt solution

Also, This is the final product of the cobalt (III) carbonato (yeah, yeah, poor naming, but I don't care) synthesis. I did not record the yield since my cobalt chloride was already wet before the synthesis.

Very cool olive colored powder

Also, I don't know if I have shown off my copper (II) aspirinate, which is a beautiful aqua color:

[Edited on 1-7-2013 by sargent1015]

sargent1015 - 6-7-2013 at 06:28

Well, I am on vacation now, so my project here is on hold.

But I will say this, the iron aspirinate synthesis did yield a red-brown solid after the LONG gravity filtration (this is why I am looking for an aspirator/pump currently

). After I do get something for filtration, I will attempt this synthesis again, but It does seem to work and it also seems stable.

I will post a picture of the product sometime next week.

cyanureeves - 6-7-2013 at 08:21

darn it i threw away my copper aspirinate because it was aqua color and thought it was supposed to be dark navy blue like the one in the youtube video.before i threw it away i added hydrogen peroxide to it and it turned yucky brown.

[Edited on 7-6-2013 by cyanureeves]

bfesser - 6-7-2013 at 08:39

Adding hydrogen peroxide to the complex will yield muddy brown <a href="http://en.wikipedia.org/wiki/Copper(II)_oxide" target="_blank">CuO</a> <img src="../scipics/_wiki.png" /> (similarly to <a href="viewthread.php?tid=17397" target="_blank">copper(II) carbonate</a>

.

[Edited on 7/26/13 by bfesser]

DrSchnufflez - 17-8-2013 at 04:54

I attempted making it tonight by following the procedure for making copper (II) acetylsalicylate:

Approx 4g of acetylsalicylic acid was dissolved in 50ml of distilled water on an oil bath.
Nickel (II) carbonate was added bit by bit until no more would dissolve.
A distinct lack of gas evolution was noted during the addition of the nickel (II) carbonate.
This resulted in a very deep red bordering on black solution with a dark purple blue precipitate.
The precipitate was vacuum filtered off and the filtrate was set aside.
The precipitate was then washed on the frit with 99% propan-2-ol which resulted in a crimson coloured solution.

Upon looking at the first aqueous filtrate a layer of apparently red, opaque crystals had deposited.

Hash anyone attempted the synthesis of nickel(II) acetylsalicylate complexes?
Does anyone know of nickel exhibiting a deep red colour in solution?

Ill take some photos when I can gain access to a camera.

prof_genius - 18-8-2013 at 05:58

I will try making it when I get new chemicals.

Boffis - 20-8-2013 at 13:24

I can't reproduce your reaction> Tried two different approaches. First I repeat the experiment above with 0.53g of acetylsalicylic acid (M & B lab grade) in 20ml of boiling water (about the minimum it will dissolve in) and added basic nickel carbonate until no more would dissolve even after boiling for 10 minutes. The solution was slightly brownish red rather than green and was filtered quickly through a hirsh funnel and lefts to cool/evaporate in a shallow basin. After 8hr a little brown liquid remains but no crystals yet. The insoluble excess carbonate was stained orange red in places by tiny amorphous looking patches of another compound.

Next I dissolve 1.03g of acetylsalicylic acid in 40ml of boiling water and added 0.71g of nickel acetate tetrahydrate (2:1 M ration). The latter dissolved to give a pale green solution but no crystal formed on cooling and the solution has remained pale green for about 7 hours. This is curious since normally the acetylsalicylic acid crystallizes swiftly on cooling. The lack of crystals suggests that a more soluble complex has formed but it appears to be green.

I report back if any crystals form.

[Edited on 21-8-2013 by Boffis]

bfesser - 20-8-2013 at 15:06

Boiling <a href="http://en.wikipedia.org/wiki/Aspirin" target="_blank">acetylsalicylic acid</a> <img src="../scipics/_wiki.png" /> (ASA) in water will <a href="http://en.wikipedia.org/wiki/Hydrolysis#Esters_and_amides" target="_blank">hydrolyze the ester</a> <img src="../scipics/_wiki.png" /> to <a href="http://en.wikipedia.org/wiki/Salicylic_acid" target="_blank">salicylic acid</a> <img src="../scipics/_wiki.png" /> and acetic acid—you are aware of this, correct? The excess water and heat shift the equilibrium in favor of the products.

<img src="http://2012books.lardbucket.org/books/general-chemistry-principles-patterns-and-applications-v1.0/section_18/f574f030712b50f72c75ab7b0ebcd468.jpg " width="500" />

[edit] If you want to prepare complexes of ASA from aqueous solution, first react the acid with a gentle base (e.g. NaHCO3

to avoid hydrolyzing the ester. Then prepare a salt of your transition metal with a weak acid (acetic is suitable). Mixing of the two solutions will precipitate an insoluble <a href="http://en.wikipedia.org/wiki/Chelate" target="_blank">chelate</a> <img src="../scipics/_wiki.png" />, if you pick the right metal and conditions.

[Edited on 4.1.14 by bfesser]

sargent1015 - 26-8-2013 at 04:57

I agree with Bfessor, you need the weak base (bicarb) to react with the acetylsalicylic acid completely in solution before adding the metal salt, no heat required. It will take a sizeable amount of stirring since sodium bicarb is not all that soluble either. I used 1:1.1 (bicarb:ASA) equivalence, since you don't want any sodium bicarb left to react with the metal. After this, add a metal salt solution.

My guess is that the complex is insoluble in water, but then again doesn't nickel like to form octahedral complexes? In which cases it would be a salt.

Does anyone have a reference for the Nickel complex we are trying to make?

Boffis - 8-9-2013 at 15:49

@DrSchnufflez I have continued to monitor the experiments I did with nickel and aspirin and described above. The experiment with the nickel carbonate produced no crystals even after a couple of days so I evaporated down on a hotplate on the minimum setting and ended up with a glassy semi-solid reddish brown residue that still would not crystallise. I tried triturated it with alcohol but this made no obvious difference apart from the fact I ended up with a cruddy mess. The fact that the aspirin itself wouldn't crystallise out from the residue does suggest that it has formed a complex with the nickel.

With the nickel acetate-aspirin experiment I got a pale green glassy residue that also would not crystallise under any condition.

I think that the difference in colour is due to the partial hydrolysis of the aspirin to salicylic acid around the basic nickel carbonate and that it is the salicylic acid that forms the red complex. In the nickel acetate experiment the nickel compound was added only after complete solution of the aspirin so hydrolysis was minimal.

Contrary to the comments about aspirin above it can be recrystallized from boiling water, in the absence of an acid or alkali hydrolysis is slow and in my experiment with nickel carbonate the boiling was not long enough to cause much hydrolysis except possibly in the immediate vicinity of the basic nickel carbonate.

If you doubt the difference between the complexes of aspirin and salicylic acid are that great consider the difference between the respective ferric complexes. That of aspirin is brown and insoluble while that of salicylic acid is deep purple and freely soluble. I have used this reaction many time as a colorimetric method to analyse iron but I have never made it concentrate enough to know if you can crystallise the complex. The colour of the ferric complex raises an interesting possibility; how pure is your nickel carbonate? Could it contain a little iron?

It should be easy enough to test this by repeating the experiment with salicylic acid, I am not at home at present so I can't try it but someone else might.

bfesser - 8-9-2013 at 18:43

Quote: Originally posted by Boffis

Contrary to the comments about aspirin above it can be recrystallized from boiling water, in the absence of an acid or alkali hydrolysis is slow and in my experiment with nickel carbonate the boiling was not long enough to cause much hydrolysis except possibly in the immediate vicinity of the basic nickel carbonate.

This contradicts my past observations. Do you have empirical evidence to back up this claim? I'd conduct some hydrolysis experiments and follow-up with melting point determinations, but my entire lab is stowed away in storage.

Regardless, I still think you're going about this all wrong... you shouldn't have to force anything to crystallize out. I'd expect the complexes to be highly insoluble. Try as I suggested above, by using a weak base to deprotonate the acid, etc. I'd dig myself out of apartment hiatus and do it myself, but I don't have any nickel salts.

DrSchnufflez - 8-9-2013 at 20:26

That's very interesting Boffis,
My nickel carbonate was made by electrolytically dissolving a cupronickel spoon in iron free HCl. The resultant green liquor was filtered from the copper sponge that had deposited on the lead cathode.
The nickel chloride was then precicpated as the carbonate by addition of a stoichiometric amount of sodium bicarbonate.

I will buy some more asprin, hydrolyse and try to make a nickel salicylic acid complex.
Also
bfesser, I shall try your method of making the Acetylsalicylic acid complexes using a weak base to deprotonate the acid.

sargent1015 - 23-9-2013 at 21:08

Feels good to have some free time to do amateur chemistry again. It has been way too long since I have been in my home lab!

Well, time for another go at aspirinate complexes. This time, it is bis(aspirinato) bis(imidazole) Cu (II).

The procedure, modified from using 2-methylimidazole, is found in:
Abuhijleh, A. L. Inorganic Chemistry Communications, 14, 2011, 759-762.

Basically, imidazole is dissolved in ethanol and copper aspirinate is added directly to the reaction mixture. I hope copper aspirinate has enough solubility to get into solution and react with the imidazole, since the solution is quite cloudy. I will let the solution stand overnight after filtering and crash the product out (Hopefully) with ether tomorrow evening.

[Edited on 24-9-2013 by sargent1015]

violet sin - 23-9-2013 at 22:18

I bought some aspirin the other day to undertake some of these reactions. started by dissolving 31 x 325mg aspirin in MeOH, going to distill off the methanol tonight, hopefully be left with ~10g ASA to play with. a lot I know but didn't wanna make some for each rxn, and with the talk of ASA vs. SA versions should have plenty for test tube size tests and perhaps more colors. also I was wondering if using an acetate salt of the desired metals instead of the sulfate ( like in Cu version ) would have any ill consequences? still have some Ni(CH3COO)2 laying around from electro-chem fun and thought it might make it gentle on the ASA or even not need the sodium bicarbonate portion. PURELY speculation on my behalf, but I thought I would ask the opinions of members if it was a waste of time/materials. I still planned on doing things in the conventional manner but why not have a lill fun. also what about heating the ASA with borax as the base? could lead to boron aspirinate perhaps. seems unlikely but again why not have a lill fun n try.

I have Li, Na, Mg, Al, Ca, Ti, Cr, Fe, Co, Ni, Cu, Zn, Ga, Zr, Nb, Pd, Ag, In, Sn, Ta, W, Pt, Ag, Hg, Pb, Bi metals and a few other salts also. Deff not going to try them all but I have some to work with for sure.
any thoughts?
-Violet Sin-

edit: realized I don't have a means to get some of those into sol. to participate though. I just love metals

[Edited on 24-9-2013 by violet sin]

bfesser - 24-9-2013 at 07:29

sargent1015, I very much look forward to reading (hopefully seeing) the results of your preparation! I've got a large jar (500 g) of high purity <a href="https://en.wikipedia.org/wiki/Imidazole" target="_blank">imidazole</a> <img src="../scipics/_wiki.png" /> that's been sitting unopened in my storage bins for many years. If I can find a safe place (i.e. not in my apartment) I will try to conduct a parallel experiment to yours. Do you have a copy of that paper to share in References?

Violet sin, an acetate should work even better, as it is the salt of a weak acid, CH3COOH, rather than the strong acid H2SO4. I suggest (see first post) washing out all the binder materials from aspirin tablets with water before doing any recystallization. As for the boron idea; I'm no expert (speculating), but I don't think boron forms coordination complexes with carboxylic acids. I'd be interested to see photos of the result of sucn an experiment, however—perhaps something interesting would form.

DraconicAcid - 24-9-2013 at 08:16

Quote: Originally posted by violet sin

also what about heating the ASA with borax as the base? could lead to boron aspirinate perhaps. seems unlikely but again why not have a lill fun n try.

I have Li, Na, Mg, Al, Ca, Ti, Cr, Fe, Co, Ni, Cu, Zn, Ga, Zr, Nb, Pd, Ag, In, Sn, Ta, W, Pt, Ag, Hg, Pb, Bi metals and a few other salts also.

Don't use borax, or your metals may precipitate as the borate instead of the aspirinate.

Don't bother with the alkali metals- they won't precipitate with aspirinate. Don't bother with the alkaline earth metals or aluminum. If they do precipitate, it will be a boring white material. I similarly wouldn't bother with titanium, zirconium, niobium or tantalum- these tend to stick with their highest oxidation state, with no d-electrons to give lovely colours. The Cr, Fe, Co, and Ni have potential to give nice precipitates; the Pd and Pt may work, but I wouldn't want to waste precious platinum on making an aspirinate. Things like mercury and lead I prefer to just leave on the shelf.

violet sin - 24-9-2013 at 12:11

@bfesser: the aspirin I got were from a discount store, only aspirin and starch if I am to believe the packaging. so they dissolved in the MeOH rather nice leaving an absolutely clear sol. with a little dust in the bottom. was going to filter it and go from there. like 500 count bottle of 325mg pills for <3$ so if there is a prob I have more. also don't have a stirrer so would be all by hand to get the starch out. and cool beans I was hoping the acetates would be easier/gentler.

@DraconicAcid: I wasn't planning on using borax as the base to get Na(ASA), but rather, was hoping the basic nature of the borax itself with sodium roaming around would allow the borate to complex. kinda opening its own door to a precipitate as it were. who knows something may come of it. I read a bit on boron and zinc aspirinates(or at least salicylate), but it was on some hippy looking site about healing stuff, tonics and all. and I couldn't find any scientific papers to back any of it up.

http://cancernaturaltherapyfoundation.org/nl/200909.pdf
"In addition to copper salicylate complexes also a salicylate complex with zinc and boron is a good healing remedy. This has been called the Schweitzer Formula, and is formed from zinc (oxide or carbonate), boron (boric acid) and salicylic acid. It has been used as an antibiotic, disinfectant, fungicide, and anti-inflammatory agent
The Schweitzer Formula supposedly was developed 1915 in Germany and sold worldwide since 1920. In addition to any kind of infection or
inflammation, it has been used in cancer treatment, to improve the immune response and blood oxygenation. Applied externally it is
claimed to heal injuries and skin diseases, including acne, scarring varicose veins and varicose ulcers. To make the Schweitzer Formula
dissolve 9.2g of salicylic acid, 2.1g of boric acid and 2.7g of zinc oxide or 4g of zinc carbonate in 2 litres of hot water. You may get these
ingredients from a pharmacist or supplier of fine chemicals and have these quantities weight out. However, it is sufficient to use approximate
amounts. You may use 2level teaspoons of salicylic acid and half a teaspoon each of boric acid and zinc oxide or one level teaspoon of zinc
carbonate"

when they use words like supposedly and give a by volume measurement for recreating the mix, I have serious doubts.

also I know a bunch of the metals I listed are impractical for this. but I been collecting for a while so they got added on the list.

unfortunately I didn't get enough done last night and am going out of town again today here in a couple hours so any progress is going to have to wait a week or so. our baby shower is this sunday n I'm trying to get some work in before I come back so I may be gone a few weeks.

-Violet Sin-

bfesser - 24-9-2013 at 13:55

Quote: Originally posted by violet sin

cancernaturaltherapyfoundation.org
. . .
when they use words like supposedly and give a by volume measurement for recreating the mix, I have serious doubts.

You should have had doubts as soon as you saw the domain name. Also, any published media using <a href="https://en.wikipedia.org/wiki/Papyrus_(typeface)" target="_blank">Papyrus</a> <img src="../scipics/_wiki.png" /> immediately loses any and all potential credibility.
<img src="http://www.cancernaturaltherapyfoundation.org/images/natures_way.gif" />
<a href="http://www.cancernaturaltherapyfoundation.org/health-articles-view.php?id=2" target="_blank">Copper Salicylate: A Potent Inflammation Fighter and Rejuvenator</a> — Walter Last

The references are unimpressive at best and—as far as I see—the author has no qualifications to speak of.

[Edited on 24.9.13 by bfesser]

sargent1015 - 24-9-2013 at 15:28

Well looks like I am out of ether for the time being, probably whip up a quick batch soon. For now, evaporation will be the key to finding if I made the correct complex.

Bfesser, here is the reference, check the cited works for the synthesis, #14, and realize that this is for 2-methylimidazole, my hopes is that I do not need that methyl group to get the imidazole to complex.

http://www.sciencemadness.org/talk/viewthread.php?tid=26316#...

Violet Sin, I have an article on several different metal aspirinates, apparently they all follow the same synthetic procedure, however my cobalt (II) aspirinate didn't crash like it was supposed to. Check the thread above for the paper.

And just because I like pretty pictures, here is the current state of my product.

I love the beautiful dark blue solution, which makes me quite hopeful. The 2-methylimidazole version claims purple crystals, so I wonder if these will be similar.

bfesser - 24-9-2013 at 18:01

Hmm. I thought I had some <a href="http://monographs.iarc.fr/ENG/Monographs/vol101/mono101-014.pdf" target="_blank">2-methylimidazole</a> <img src="../scipics/_pdf.png" />, but after searching through my bins I was only able to find an unopened 500 g bottle of Fisher <a href="http://www.fishersci.com/ecomm/servlet/fsproductdetail_10652_632433__-1_0" target="_blank">certified imidazole</a> <img src="../scipics/_ext.png" /> and an unopened 5 g bottle of <a href="http://www.sigmaaldrich.com/catalog/product/aldrich/425842?lang=en&region=US" target="_blank">4-imidazolecarboxylic acid</a> <img src="../scipics/_ext.png" />, 98% from Aldrich. I guess my memory was flawed. Until now, I had no idea they were so valuable—I doubt I payed more than $20 each!

<hr width="80%" />
I've noticed that there's been <a href="viewthread.php?tid=26095">recent interest</a> in the coordination compounds of <a href="https://en.wikipedia.org/wiki/Ethylenediamine" target="_blank">1,2-diaminoethane</a> <img src="../scipics/_wiki.png" />. Perhaps one of the participants [cough]woelen[cough] in that thread could attempt the synthesis of 2-methylimidazole as outlined in this paper:

<table><tr><td width="20"></td><td>The staged synthesis of 2-methylimidazole from ethylenediamine and acetic acid in the presence of a bifunctional aluminoplatinum catalyst K.M. Gitis, N.I. Raevskaya, G.V. Isagulyants Bulletin of the Russian Academy of Sciences, Division of chemical science.September 1992, Volume 41, Issue 9, pp 1551-1554.
DOI: <a href="http://dx.doi.org/10.1007/BF00863570" target="_blank">10.1007/BF00863570</a> <img src="../scipics/_ext.png" /></td></tr></table>
The preparation of the γ-Al2O3/Pt catalyst would be interesting and challenging in and of itself. The whole project from 1,2-diaminoethane to the final complex would make an excellent addition to <a href="forumdisplay.php?fid=20">Prepublication</a>.

<a href="viewthread.php?tid=25035&page=7#pid300833">[reference requested]</a>

P.S. woelen, I know that you <a href="viewthread.php?tid=15513">'have a thing for copper'</a>—how can you resist?

Oscilllator - 27-9-2013 at 02:08

This isn't stricly speaking an aspirinate synthesis, but I thought I'd post it here anyway.
Recently I successflly made a copper/salicylic acid complex from copper sulfate using the procudure that bfesser outlined, although I didn't bother to measure out the reagents.
It is interesting to note that although the aspirinate complex is insoluble, the salycilic acid complex is very soluble, having a deep green colour similar to a concentrated solution of nickel nitrate.
Here is a picture of the compound after vacuum filtration:

I also attempted to prepare a nickel analogue, however is seemed that only nickel carbonate formed

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