Sciencemadness Discussion Board - Making Copper Acetate?

Making Copper Acetate?

jgourlay - 8-9-2008 at 06:38

Gents, I've been casting about for some copper acetate to use in my continuing adventures of my son and I crystallizing everything under the sun. I want to hit copper acetate, but am having trouble.

Either I find reagent grade at $15+ per 100g, or it's one of a bajillion ingredients in cat food and vitamins. I'm looking for the ebay grade ziplock baggie full for $4.99.

20 years ago, I thought the chinese were going to take over the world so "speculated" in "commodities" by running out to a scrap metal dealer and buying several hundred pounds of copper scrap. I've sold most of it, and have about 50 lbs left.

Can I get what I want by (trying) to dissolve copper in glacial acetic acid? Would that be educational for the boy? Or would it be one of those "demonstrations" that starts with the "ooh"ing and "ah"ing and ends with the running and screaming?

panziandi - 8-9-2008 at 06:56

If you keep copper wire in acetic acid you will get a blue-green solution of copper (II) acetate. I have done this with vinegar and 33% acid, but it is SLOW. If you heat copper wire in a blue flame you can oxidise it to copper (II) oxide which will dissolve more easily in the acetic acid. Copper (II) oxide and copper (II) carbonate (basic) can be had cheaply and can be dissolved in the acetic acid quickly to yield the acetate. Copper (II) acetate are beautiful crystals, enjoy

Mr. Wizard - 8-9-2008 at 07:13

That's how I made it when I was a kid. I put some copper wire in a jar, open to the air, half full of vinegar. The vinegar does dry up, but just keep adding to it. If I recall, the CO2 from the air plays a part in the reaction. It took weeks, but I got some of the nice crystals.
As a caution, I'd recommend gloves and a dust mask. The stuff has a nasty irritating effect on your throat when the dry powder gets in the air. Has anyone else ever noticed this problem?

jgourlay - 8-9-2008 at 07:14

Thanks!

Magpie - 8-9-2008 at 07:38

There's some more copper acetate making experience here in this large thread:

http://sciencemadness.org/talk/viewthread.php?tid=5529&p...

[Edited on 8-9-2008 by Magpie]

Klute - 8-9-2008 at 07:54

If you have acces to nitric acid, it dissolves very easily (recycle the NOx fumes). You can then neutralize with alkali to form Cu(OH)2, heat until all the blue gel turns to a easily-decanted black solid (CuO), decant and wash with water a few times (leave it to decant, impossible to correctly filter), and then dissolve the CuO in 70% AcOH, and evaporate.

I've done it a few times at a 50-100g scale, because i can get nitric acid dirt cheap. Might not be the best option for everyone else.

gsd - 8-9-2008 at 09:03

You can make copper (II) Acetate by reaction of copper metal with acetic acid in presence of air (oxygen) at elevated temperature (boiling temperature of AA). I have successfully made it as under:

take dilute (about 50%) AA in a 3 neck round bottom flask kept in a heating mantle. Fit tube filled with copper metal pieces ( small stampings, cleaned wire scrap, turnings, tube cuttings etc. etc. - basically any small but CLEAN source of metal) on the centre neck. Attach a water cooled reflux condensor on top of the tube. In one side neck of the RBC, insert a tube for air entry. Use 3rd neck for thermometer pocket. Attach a gas trap to the top of condensor to scrub any escaping AA vapours.

Introduce a very small flow of air / oxygen using any suitable device ( I had used fish tank aspirator) in the flask and bring the liquid to boiling and then to very gentle reflux.

At the refulxing temperature of AA ( about 105 Dec C) the air oxidizes the copper and the copper oxide immediately reacts with AA to form CuAc2 which is washed down to the RBC by down coming liquid. It accumulates there and gets concentrated as the time passes. The experiment can run unattended for considerable duration say about 8-10 hrs. Stop the heating, dismantle the tubes and pour hot CuAc2 solution in a beaker.

Beautiful Dark Green crystals of CuAc2.H2O are obtained on cooling.

good luck

gsd

[Edited on 8-9-2008 by gsd]

Klute - 8-9-2008 at 11:49

Very clever! Looks much more practical that form the oxide and cleaning it.. I will try it out next time I need Cu(OAc)2!

kclo4 - 8-9-2008 at 16:30

A few years ago I produced Copper Acetate by electrolysis of copper wires in Vinegar IIRC.

Calcium Acetate could be reacted with Copper sulfate to form Calcium Sulfate which is basically insoluble and copper acetate.

The Calcium Acetate could be produced from chalk and Vinegar. - That is with chemicals you likley have around the house.

Also, by doing it that way you could then teach your son about reactivity series, naturalization, double displacement, etc.

chloric1 - 9-9-2008 at 15:21

Yeh but the calcium sulfate is a bear to filter! You need to boil with agitation for a few hours to get bigger crystals to filter.

kclo4 - 9-9-2008 at 16:53

How much of a difference does boiling make?
Does it take the ultra fine calcium sulfate to something the size of table salt crystals, or something?

ScienceSquirrel - 9-9-2008 at 17:00

Quote:

Originally posted by kclo4
How much of a difference does boiling make?
Does it take the ultra fine calcium sulfate to something the size of table salt crystals, or something?

Boiling a precipitate can make a big difference.

A precipitate of ferrous oxalate becomes a lot more granular on boiling for a few minutes.

watson.fawkes - 9-9-2008 at 18:20

Boiling a liquid with marginally soluble solids (assuming no vapor loss) is formally similar to reflux. In reflux, a phase change of the solvent causes part of it to go to gas, which then condenses. In boiling this mixture, a phase change of the solid causes part of it to go into solution, when then recrystallizes.

The ultimate reason that crystals get larger is that this configuration--fewer and larger crystals--is a lower energy configuration that the prior configuration--more and smaller crystals. The energy is, essentially, a surface tension. It doesn't seem to ordinarily act that way because the crystals doesn't deform in response to it.

Agitation, though not like reflux, adds a second means for reducing surface area: collision.

Lastly, I would expect that putting the solution under a proper reflux column while boiling would increase crystal growth rate. An ion coming out of solution is driven locally by supersaturation at the crystal surface. Reflux, by keeping an ever-changing part of the solvent out of commission, increases the amount of volume with this local supersaturation property. Additionally, at the juncture where fresh solvent remixes, there's a zone of high solubility that will tend to dissolve fines. All these dynamics are still true at low solubilities; it's just that everything works more slowly.

kclo4 - 10-9-2008 at 16:48

Oh thats very interesting, thank you both for explaining that! I'll have to keep this in mind

crazyboy - 10-9-2008 at 20:55

Add a clean copper pipe to white vinegar or better yet concentrated acetic acid, add 35% H2O2 let sit over night or boil. When bubbling ceases filter and discard solids. Boil and collect crystals.

copper acetate

itchyfruit - 11-8-2009 at 11:41

Quote: Originally posted by kclo4

A few years ago I produced Copper Acetate by electrolysis of copper wires in Vinegar IIRC.

I have been trying this today without much(any if i'm honest)success
My procedure: Firstly i made a couple of electrodes by curling copper wire into a flat spiral then clipped them to a small beaker i filled the beaker to just above the spirals(40ml approx) with GAA and switched on the power(12vdc) after about an hour nothing had happened so i heated the beaker to 65oc still nothing so i added 5ml of 1mol cuso4 in the hope that this might make the solution a bit more conductive this just turned into a light blue precipitate and sank to the bottom of the beaker(i'm guessing the GAA stole all the h2o from the cuso4,perhaps a good way to make anhydrous cuso4 ?)I then added 10ml of 5% h2so4 the electrodes are now bubbling a bit but i don't have the deep blue solution i was hoping for!!!
Any suggestions?

Sorry i've put my post in the quote box

[Edited on 11-8-2009 by itchyfruit]

itchyfruit - 11-8-2009 at 11:50

Now i don't know what section this has ended up in.

itchyfruit - 11-8-2009 at 14:59

Now for an update, firstly my lab stinks of vinegar

The anode(+) is coated in a green/blue insoluble looking residue the cathode(-)is covered in what must be copper and most of the cuso4 precipitate has gone presumably onto the cathode after i added a bit more h2so4,not a good way of making copper acetate me thinks.

I also tried CU powder + GAA + 30% H2O2 this reacted quite violently and got pretty hot, but i now have a deep blue solution which i believe is copper acetate WAHOOO

woelen - 12-8-2009 at 00:27

I made copper acetate myself from copper sulfate, sodium hydroxide and dilute acetic acid.

The procedure is as follows:
- Dissolve copper sulfate in water
- Add sodium hydroxide in excess amount, such that a blue precipitate is formed
- Heat the liquid for quite some time, until all blue precipitate has become purely black. The copper hydoxide is converted to copper(II) oxide. This black precipitate is not so slimy as the blue copper hydroxide and can be filtered much more easily.
- Filter the black precipitate and rinse with distilled water several times to get rid of sodium and sulfate ions. Do not dry the precipitate, just rinse it.
- Add excess 30% acetic acid to the black material and let it stand in an open vessel for many many days until it has become perfectly dry. Everything will smell of vinegar when you do this inside, so it is best to put the vessel outside in a warm and dry place where animals and children cannot access it. When the acetic acid is added, the liquid becomes really dark green and black material seems not to dissolve at once, but while it is standing, slowly all of it is converted to copper(II)acetate. It is important though to have excess acetic acid.

Slowly, small crystals of copper acetate form. I obtained a crystalline meal of many glittering fine crystals. I did not get a nice single crystal, but you might obtain bigger crystal if you redissolve it in a lot of very dilute vinegar and let it evaporate very slowly.

The picture below shows the final result of this experiment. This is approximately 10 grams of copper(II)acetate.

[Edited on 12-8-09 by woelen]

itchyfruit - 12-8-2009 at 02:39

Thanks, I'll give it a try(i used a similar process to make cuprous chloride) i'll have to use koh as my naoh is not very pure.

kmno4 - 12-8-2009 at 04:13

In making Cu(OAc)2 it is important to to use excess of acetic acid - as woelen said. In another case you will get basic copper acetate, colored green.
Pure Cu(OAc)2 hydrate is colored deeply blue and when hydrolysis took place, colour changes to green-blue.
The best way to obtain pure crystals is evaporation of water from saturated solution at room temp. Unfortunately it is very slow and not effective (saturated sol. at r.t. is only 7g of acetate per 100g water). My Cu(OAc)2 is green-blue, because I had no time for playing in this way. I used CuO (contaminated with Fe and Zn, total ~5%) and acetic acid on boiling.
Prepared acetate was recrystallized from hot water and acetic acid. Analysed with x-ray spectrometer it has shown 99%+ of Cu.

itchyfruit - 12-8-2009 at 04:44

I'm just of to try this now,so by excess(try not to poke fun at my lack of knowledge'entropy51'i'm looking a you

) you mean for example 100ml of 1mol cuso4 + 110ml of 1mol koh this would be a slight excess of koh is that correct?

entropy51 - 12-8-2009 at 06:38

Surely even you can succeed with one of the recipes graciously posted for you.

But you'll probably drop it in the sink when you're finished.

Don't forget that KOH out of the bottle is usually quite a bit less than 100% pure. And NaOH is usually cheaper.

You might want to look up "peracetic acid". You're a strong contender for the Darwin Award.

[Edited on 12-8-2009 by entropy51]

[Edited on 12-8-2009 by entropy51]

bfesser - 12-8-2009 at 06:45

Quote: Originally posted by entropy51

But you'll probably drop it in the sink when you're finished.

I'm not exactly sure what entropy51 is talking about here, but please, don't dump copper waste down the drain. It's particularly harmful to marine environments. Just reduce it to the metal, or precipitate it as a salt, filter it, and keep it. Waste not, want not.

kmno4 - 12-8-2009 at 11:27

Quote: Originally posted by bfesser

... please, don't dump copper waste down the drain. It's particularly harmful to marine environments. Just reduce it to the metal, or precipitate it as a salt, filter it, and keep it. Waste not, want not.

Marine environments in the drain ?
From 10g or even 1kg of copper salts to environmental pollution is very long way. Especially the way leading via the drain ....

JohnWW - 12-8-2009 at 12:10

You may be able to use the stuff as a paint pigment, to tint existing paint, or by mixing it with clear polyurethane varnish or boiled linseed oil.

entropy51 - 12-8-2009 at 13:27

I try not to let heavy metals get into the sewer, BUT

The CuSO4 I buy at the hardware store is intended to be dumped into the sewer to kill tree roots invading the sewer line. The label has an EPA Registration number on it, but also states that it's harmful to fish and aquatic organisms.

Isn't it likely that the Cu plates out on the Fe sewer pipes and mostly ferrous SO4 is left in the water heading out to sea? That's what happens in a beaker of CuSO4 when I add steel wool to it in order to make Cu powder.

itchyfruit - 12-8-2009 at 15:38

I don't go dropping anything down the drain unless i'm absolutely sure it's ok to do so,entropy51 was just having his customary dig about my general incompetence!!!
Funny you should mention peracetic acid,as i did do a little test before i tried the CU+GAA+H2O2 approach,1ml of each on a little AL cup and set about it with a mapp gas blow torch hoping to detonate it!! nothing just evaporated, USELESS :mad:

I opted for KOH as mine is from a large AR grade chemical supplier,whereas my NAOH is just sink unblocker(prills)from the local hardware store.
I don't think i'd be able to make much paint out of what i made,i suppose i could paint the mail box(on a dolls house)

itchyfruit - 16-8-2009 at 17:13

My copper acetate looks like woelens. wahooooo

woelen - 17-8-2009 at 08:49

And can you tell us more about the exact procedure, used to make it? You used H2O2 and copper metal? This is another (possibly new, at least for some of us) procedure. We learn the most of it, if you provide us with exact procedures, and numerical values for amounts of used chemicals.

UnintentionalChaos - 17-8-2009 at 14:28

Quote: Originally posted by itchyfruit

I don't go dropping anything down the drain unless i'm absolutely sure it's ok to do so,entropy51 was just having his customary dig about my general incompetence!!!
Funny you should mention peracetic acid,as i did do a little test before i tried the CU+GAA+H2O2 approach,1ml of each on a little AL cup and set about it with a mapp gas blow torch hoping to detonate it!! nothing just evaporated, USELESS :mad:

It's quite hard to make any appreciable amount of peracetic acid.

Look up industrial synthesis methods- high concentrations of H2O2 and acetic acid, sulfuric acid catalyst and up to 10 day reaction times.

A mixture of vinegar and dilute H2O2 agressively attacks lead and is commonly used to clean gun barrels.

Simply bubbling air through the acetic acid will also dissolve copper, albeit slowly.

[Edited on 8-17-09 by UnintentionalChaos]

itchyfruit - 17-8-2009 at 14:51

My H2O2 method didn't work !! I used your OH method, but i think i jumped the gun a little with my celebrations it did look like yours but now it's completely dried it has a bit of a green tint and some powdery white crystals on the surface.

By the sounds of it i was a bit impatient with the H2O2 method,and i didn't use any sulphuric acid in the peracetic acid or leave it long enough,but i don't think i'll temp fate by trying it again just yet

woelen - 17-8-2009 at 22:42

These powdery white crystals most likely are either sodium sulfate or sodium carbonate (from excess NaOH). You really need to rinse the black precipitate of CuO (which is formed after heating the suspension of Cu(OH)2 in water) in order to get rid of excess NaOH and/or sulfate.

itchyfruit - 18-8-2009 at 13:26

I'm going to start again from scratch, i must have got something wrong i'll double check all my measurements and make sure i get rid of any excess NAOH next time.

It's weird because i thought i had to much GAA as it really smelt vinegary...

chemoleo - 18-8-2009 at 13:57

You can also try the CaAc2 (from CaCO3 and white vinegar) and CuSO4 route, this way you get CuAc2 and CaSO4. CaSO4 can be filtered off reasonably well. I've done that a long time ago in a 5 liter quantity. This was then reacted with Pb metal to get PbAc2 (see the big PbO2 thread I think).
Anyway, it works becuase CaSO4 is badly soluble, so the amount of CaSO4 left in the CuAc2 solution is small.

By the way, woelen, I also did that a while ago (getting CuO from Cu[OH]2), but the resulting CuO was still a worse thermite than CuO from the pottery supply. Even though this CuO was roasted on red heat, and guaranteed free of Na2SO4 (it was washed for eons). It's an interesting powder actually, behaving/swirling almost like a liquid when hot!

[Edited on 18-8-2009 by chemoleo]

woelen - 19-8-2009 at 22:33

Quote: Originally posted by chemoleo

By the way, woelen, I also did that a while ago (getting CuO from Cu[OH]2), but the resulting CuO was still a worse thermite than CuO from the pottery supply. Even though this CuO was roasted on red heat, and guaranteed free of Na2SO4 (it was washed for eons). It's an interesting powder actually, behaving/swirling almost like a liquid when hot!

[Edited on 18-8-2009 by chemoleo]

I did not dry the CuO, because I wanted to make Cu(OAC)2 from it (and that can better be done with the wet stuff than with some inert dry powder). I can imagine that the roasting in red heat leads to loss of some oxygen and that part of it was decomposed to Cu2O. Your powder might have been some mixed oxide, somewhere between CuO and Cu2O. Such a mixed oxide definitely will perform worse in thermite reactions than pure CuO. On the other hand, I expect that such a mixed oxide still looks black and not red.

If you still have some of your roasted CuO left, then try dissolving some in concentrated HCl. If the solution becomes very dark brown, almost black, then it contains copper(I) and copper(II) and then my theory is right, if the solution becomes green/brown and not really dark, then it only contains copper(II) and then my theory is wrong and another explanation must be found for the different behavior in thermites.

chloric1 - 20-8-2009 at 13:41

I make other copper salts from store baught copper sulfate a different way. I make a solution of say 20% copper sulfate and it is moderately heated to 60 celsius. When solution is complete, i slowly add solid sodium bicarbonate until effervescense stops. The resulting basic copper carbonate is easier to filter and requires less washing. The copper sulfate should be somewhat dilute to keep most of the sodium sulfate that forms in solution. Copper is a weakly basic metal and the carbonate decomposes very easily on heating producing a pure cupric oxide. The basic carbonate can also be used for your copper acetate. I used this method to make cupric formate.

zed - 30-8-2009 at 14:46

Refluxing colloidal precipitates, like Calcium Sulfate, Calcium Phosphate, or the classic Silver Chloride.....In an effort to increase crystal size, and thereby filter-ability......Well, it's a pain in the ass.

A long time ago, in a galaxy, far, far away.......I donated this procedure, to my friend Carl Abrams, who was writing a text book.

Simply heat your reactant solutions near boiling. Then slowly, pour one solution, into the other.

Since your Phosphate, Sulfate, Chloride, product....is slightly more soluable in very hot water, your initial crystals, form more slowly, and thereby attain a significantly larger size.

They simply fall out of solution.

Thus the simple expedient, of pre-heating your solutions, renders the otherwise maddening Phosphate-type of determination.....Easy.

The classic experiment, in my chem days, was the Chloride determination, via precipitation with Silver Nitrate. First, form an un-filterable Chloride precipitate. Then, boil endlessly in attempting to coagulate. Finally, filter, oven dry, and weigh.

At stake, was most of your chemistry grade.

Man, what a mess. The procedure made people hate chemistry.

[Edited on 30-8-2009 by zed]

Nathaniel - 26-12-2012 at 12:04

I have a problem with copper acetate preparation, but before I start describing everything, here's what troubles me:

1. CuCO3*Cu(OH)2 + acetic acid( large excess) --> uncomplete reaction
2. Heating the solution of copper acetate --> quite some blue-green percipitate forms (probably some basic copper salt)

I've tried to make copper acetate many times, usually with basic copper carbonate, made from sodium carbonate and a soluble copper salt (sulfate, chloride). Even though I used excess of acetic acid (usually 9% vinegar, but I tried pure acetic acid as well) the carbonate wouldn't react completely even after a few days on warm temperatures. I never actually determined the amount of undissolved carbonate, but it's usually approximately 25% or more. I plan to repeat the experiment again on Friday so I'd appriciate if someone could help me

I also had problems when I tried to boil down the filtrate of the above reaction. Even before boiling point a green-blue percipitate starts to appear and the amount of it even increases with boiling. It's probably some basic salt of copper which is surprising because there's still a lot of acetic acid present in solution. This also happens when I let the filtrate to partially evaporate and get some really nice crystals, but when I try to dissolve them in hot vinegar they decompose to the blue-green percipitate mentioned before. I couldn't find this mentioned anywhere so I'm turning to you guys for advice why this happens and how to prevent it (more the later as I have many theories why it happens but can't find the solution).

[Edited on 26-12-2012 by Nathaniel]

Rich_Insane - 29-12-2012 at 16:02

Will acetic anhydride substitute for GAA in the reaction of Cu2O to Cu(OAc)2, or will I have to add water?

Update: Heated 20 grams of CuSO4.5H2O with 3.5 grams of NaOH to produce Cu(OH)2. Boiled this down in an attempt to create CuO, but by the time all the water boiled away, there was still a lot of Cu(OH)2. So I decided to simply isolate the chunky green-black material and add 15 ml of Ac2O plus 20 ml of water. After an hour or two, there was a beautiful deep blue clear liquid and some blue-colored crystal precipitate on the bottom, in addition to some white precipitate (sodium sulfate?). I added 20 more ml of water to hydrolyze any remaining Ac2O. The solution quickly turned cloudy.

Now there is a slimy gray-blue precipitate on the bottom (probably sodium salts discolored by copper acetate) in addition to the deep blue transparent layer. Does anyone know how I might remove these impurities? It's not absolutely necessary: I plan to use the Cu(OAc)2 as a catalyst in a different reaction. However, the purer my reagents, the more accurate my calculations.

[Edited on 30-12-2012 by Rich_Insane]

[Edited on 30-12-2012 by Rich_Insane]

Metacelsus - 31-12-2012 at 19:14

I've never had problems with the basic copper carbonate + vinegar reaction. As long as the basic copper carbonate is pure (e.g., washed with copious water after synthesis), it should react completely with vinegar. I make mine by reaction of root killer (CuSO4*5 H2O) with washing soda (NaCO3)

Metacelsus - 31-12-2012 at 19:16

Also, as to the blue-green precipitate, it might be crystals of copper acetate. Copper acetate monohydrate is blue-green, but you probably know best.

Metacelsus - 1-1-2013 at 15:20

If you want to make copper acetate, use basic copper carbonate, not copper hydroxide. It's easier and cheaper to make.
2 CuSO4 + 2 Na2CO3 + H2O → Cu2(OH)2CO3 + 2 Na2SO4 + CO2
The basic copper carbonate precipitates.
Make sure to use an excess of acetic acid when making the acetate.
I have done this reaction several times with vinegar as my acetic acid source, and it works just fine. It's a pain to remove the water though.

Eddygp - 2-1-2013 at 07:03

Copper in a H2O2/AcOH mixture will work out slowly but OK.

Metacelsus - 2-1-2013 at 14:25

Have you ever heard of peracetic acid?

H2O2 + CH3CO2H -> CH3CO3H + H2O

H2O2 + AcOH is not something I'd like to try.

Sedit - 2-1-2013 at 19:10

I made some Copper acetate not to long ago using electrolysis of a EtOH solution and NaCl and Copper electrodes. Dunno what the yields where but I was just playing and was happy that I formed any Copper acetate by this method. Im sure it could be stream lined and improved using various electrolytes if the folks here put there minds to it.

Nathaniel - 3-1-2013 at 03:38

Thanks guys, I know the basic copper carbonate method is the cheapest and I've somehow managed to dissolve almost all the precipitate, but needed 4eq. of vinegar instead of 2...(until now I always added ca. 2,5eq.) The precipitate formed when the solution is heated is not copper acetate because it wouldn't dissolve in vinegar or distilled water...strange that only I'm having this problem

I'm planning to make a large amount of copper acetate for crystal growing (1 kg

) so I wanted to make the synthesis as efficient as possible.. I guess I'm stuck with slow evaporation..

DraconicAcid - 1-2-2013 at 10:18

My students make this material every semester, for a % yield lab. We start with copper shavings, dissolve in nitric acid, precipitate the basic carbonate (which gets washed well with water, and then acetone, so that it is less pasty to work with). The basic carbonate is then reacted with 50% acetic acid, and the solution is stuck in ice to get most of the product out. I haven't had much luck with growing crystals of it, but I never seem to get good results with that. One thing I have learned is that if hang a seed crystal in a solution of copper(II) acetate, do not suspend it with a copper wire.

If you're getting a precipitate when you heat the solution, then perhaps your acetic acid is too dilute, and you're getting a basic copper(II) acetate?

(First post!)

Nathaniel - 22-2-2013 at 13:23

I just restarted making a large batch of copper acetate so I noticed your post a bit late, but yes, it's possible it's the basic salt that's the problem... I've just recently found out that "basic copper sulphate" exists as well - it's formula is something like Cu₄(OH)₆SO₄ or something similar so I got scared a bit that there's some sulphate in my copper acetate as well (just btw

)...

I was thinking about using stronger acetic acid (I have some) but copper acetate is quite poorly soluble so I would end up adding extra water anyway...
If I manage to make the crystals anyway, I'll post the pictures

blogfast25 - 22-2-2013 at 14:08

7.2 g/100 ml in cold water isn't that poorly soluble (Wiki). And 20 g/100 ml in hot water (Wiki) is someting that could be exploited: from a hot, saturated Cu(OAc)2 with excess acetic acid to suppress hydrolysis, crystals of the salt should be formed on cooling.

AJKOER - 22-2-2013 at 17:22

Here is a new approach. Use the apparent fact that Citric acid can effectively leach copper oxide. The following extract (see http://link.springer.com/article/10.1007%2Fs12613-012-0628-9 ) provides some detail:

"Leaching of an oxidized copper ore containing malachite, as a new approach, was investigated by an organic reagent, citric acid. Sulfuric acid is the most common reagent in the leaching of oxide copper ores, but it has several side effects such as severe adverse impact on the environment. In this investigation, the effects of particle size, acid concentration, leaching time, solid/liquid ratio, temperature, and stirring speed were optimized. According to the experimental results, malachite leaching by citric acid was technically feasible. Optimum leaching conditions were found as follows: the range of particle size, 105–150 µm; acid concentration, 0.2 M; leaching time, 30 min; solid/liquid ratio, 1:20 g/mL; temperature, 40°C; and stirring speed, 200 r/min. Under the optimum conditions, 91.61% of copper was extracted."

So dissolve an excess of powdered Cu in Citric acid/dilute H2O2 solution at 40 C with stirring. Replenish the H2O2 as needed. Add Na2CO3 to form a precipitate of Copper carbonate. Treat with acetic acid to create the Copper acetate.

Alternative Synthesis: As oxidation of Cu maybe a key and difficult step, alternately, one could heat the powdered Copper (or, a large number of copper pennies as we will be using only the surface oxide) in air and/or treat with NaOCl. Rinse and treat with Citric acid (or Citric acid/H2O2 as before) and contnue with the prior synthesis. Namely, add Na2CO3 to form a precipitate of Copper carbonate. Finally, add acetic acid to the CuCO3 to create the Copper acetate.

[EDIT] As Wikipedia cites the following reaction:

2 CuSO4 + 2 Na2CO3 + H2O → Cu2(OH)2CO3 + 2 Na2SO4 + CO2

most likely we will be treating Basic copper(II) carbonate with Acetic acid and not CuCO3.

[EDIT EDIT] However, the reaction with Baking Soda, NaHCO3, could be more favorable forming CuCO3(?) in conc solutions.

CuSO4 (aq) + NaHCO3 --> CuCO3 (s) + NaHSO4 (aq)

so I would substitute NaHCO3 for Na2CO3 in the synthesis.

[Edited on 23-2-2013 by AJKOER]

[Edited on 23-2-2013 by AJKOER]

KonkreteRocketry - 22-2-2013 at 23:32

So i just keep to put CuO in Acetic acid and heat it a bit ?

Nathaniel - 23-2-2013 at 04:52

Quote: Originally posted by blogfast25

7.2 g/100 ml in cold water isn't that poorly soluble (Wiki). And 20 g/100 ml in hot water (Wiki) is someting that could be exploited: from a hot, saturated Cu(OAc)2 with excess acetic acid to suppress hydrolysis, crystals of the salt should be formed on cooling.

Yeah, I already looked that up, but the problem is that I don't have that much acetic acid. I'm making 1mol copper acetate, which dissolves in almost 1l of hot water. So if I would want the solution to be let's say 10% acetic acid AFTER reaction is complete, that's almost 100ml pure acetic acid excess (210ml total) and I don't have that much...

I guess I'll try to dissolve the basic carbonate in multiple portions of 9% vinegar...with a large excess it should work

blogfast25 - 23-2-2013 at 07:07

Quote: Originally posted by KonkreteRocketry

So i just keep to put CuO in Acetic acid and heat it a bit ?

Nope. Commercial CuO is likely to have been calcined hard. That stuff doesn't usually dissolve into anything but the strongest and most concentrated acids.

[Edited on 23-2-2013 by blogfast25]

blogfast25 - 23-2-2013 at 07:10

Quote: Originally posted by Nathaniel

I guess I'll try to dissolve the basic carbonate in multiple portions of 9% vinegar...with a large excess it should work

Or you could try and concentrate the acetic acid a bit. React the vinegar with stoichiometric amounts of slaked lime (Ca(OH)2) to obtain calcium acetate, by boiling in. Treat the solid with conc. H2SO4 and distill over the acetic acid (ethanoic acid).

KonkreteRocketry - 23-2-2013 at 11:45

Quote: Originally posted by blogfast25

Quote: Originally posted by KonkreteRocketry

So i just keep to put CuO in Acetic acid and heat it a bit ?

Nope. Commercial CuO is likely to have been calcined hard. That stuff doesn't usually dissolve into anything but the strongest and most concentrated acids.

[Edited on 23-2-2013 by blogfast25]

I got my CuO by decomp from Copper IInitrate, so that shall work ?

blogfast25 - 23-2-2013 at 12:45

Quote: Originally posted by KonkreteRocketry

I got my CuO by decomp from Copper IInitrate, so that shall work ?

If you didn't calcine it to death then probably, yes.

KonkreteRocketry - 24-2-2013 at 05:45

Quote: Originally posted by blogfast25

Quote: Originally posted by KonkreteRocketry

I got my CuO by decomp from Copper IInitrate, so that shall work ?

If you didn't calcine it to death then probably, yes.

what do u mean by calcine it to death ? would heating CuO do something to it ? I heat it in a 200 degree alcohol flame.

woelen - 24-2-2013 at 06:29

Calcining to death of an oxide is heating very strongly (1000 C or so) for an extended period of time. Many oxides become very inert when this is done. E.g. they cannot be dissolved in acids anymore when they are calcined.

Some noteworthy examples are:

Cr2O3
Al2O3
Fe2O3
Co3O4
TiO2

These oxides do not dissolve in strong acids, not even boiling hot 35% hydrochloric acid, 65% nitric acid or pure H2SO4 are capable of dissolving these oxides after calcining. This change of property is due to formation of a more compact crystalline form, which slowly occurs at very high temperatures, even well below the real melting point of the oxides. An inert calcined oxide usually requires workup with molten NaOH or molten NaHSO4 in order to dissolve it. This requires working with melts at several hundreds of degrees centigrade.

I think that CuO, on the other hand, remains fairly reactive, even after strong calcining.

KonkreteRocketry - 24-2-2013 at 09:00

Quote: Originally posted by woelen

Calcining to death of an oxide is heating very strongly (1000 C or so) for an extended period of time. Many oxides become very inert when this is done. E.g. they cannot be dissolved in acids anymore when they are calcined.

Some noteworthy examples are:

Cr2O3
Al2O3
Fe2O3
Co3O4
TiO2

These oxides do not dissolve in strong acids, not even boiling hot 35% hydrochloric acid, 65% nitric acid or pure H2SO4 are capable of dissolving these oxides after calcining. This change of property is due to formation of a more compact crystalline form, which slowly occurs at very high temperatures, even well below the real melting point of the oxides. An inert calcined oxide usually requires workup with molten NaOH or molten NaHSO4 in order to dissolve it. This requires working with melts at several hundreds of degrees centigrade.

I think that CuO, on the other hand, remains fairly reactive, even after strong calcining.

So yeah mine was not even over 200 degree, so my CuO shall work with household vinegar right ? My vinegar is 6% Acetic acid. Shall i heat it while i put the CuO insdie or its ok if i just put CuO.

And in the end i shall get Copper acetate right ?

Cool.

I have some K2S and idk what to do with it,

Do u know what K2S thermal decmpose into ?

and when i put K2S with water, this will happen right ?

K2S + H2O = KOH + KSH

blogfast25 - 24-2-2013 at 09:55

Quote: Originally posted by KonkreteRocketry

Cool.

Maybe. Vinegar is a weak solution (5 - 6%) of a weak acid, i.e. acetic acid (ethanoic acid). Only a small part of the ethanoic acid is actually dissociated via:

HOAc(aq) + H2O(l) < === > H3O+(aq) + OAC-(aq)... (1)

And it's the oxonium ions (H3O+) that are supposed to do the work:

CuO(s) +2 H3O+(aq) === > Cu2+(aq) + 3 H2O(l)...(2)

Since as the concentration of oxonium ions in a commercial vinegar is really small, the reaction rate for the second reaction is small, compared to when you use a strong acid like HCl or H2SO4 where the dissociation step (1) is almost 100 %.

What's more, even if your CuO dissolved completely in your vinegar you still only have a very weak solution of copper (II) acetate, from (2).

The proof is in the trying, of course...

[Edited on 24-2-2013 by blogfast25]

kingkey24 - 27-2-2013 at 16:11

why not just react CuCO3 with acetic acid, then boil down the solution and place in a desiccator for crystallization?

AJKOER - 29-5-2013 at 06:10

Quote: Originally posted by panziandi

If you keep copper wire in acetic acid you will get a blue-green solution of copper (II) acetate. I have done this with vinegar and 33% acid, but it is SLOW. If you heat copper wire in a blue flame you can oxidise it to copper (II) oxide which will dissolve more easily in the acetic acid. Copper (II) oxide and copper (II) carbonate (basic) can be had cheaply and can be dissolved in the acetic acid quickly to yield the acetate. Copper (II) acetate are beautiful crystals, enjoy

I tried the following reaction recently and was surprised. I started by adding dilute aqueous household ammonia to some copper coated pennies (US currency). I was not expecting much, and 'SLOW' is too rapid an adverb for what does occur.

However, per this not too dated 1962 paper "Kinetics and Mechanism of Copper Dissolution In Aqueous Ammonia" fully available after signing on to ones Facebook account at http://www.academia.edu/292096/Kinetics_and_Mechanism_of_Cop... the author cites a rate for Cu dissolution as a function of available O2 and NH3.

Some of the underlying reactions include:

2 Cu + 4 NH3 + 1/2 O2 + H2O --> 2 [Cu(NH3)2]OH

2 [Cu(NH3)2]OH + 4 NH3 (aq) + 1/2 O2 + H2O --> 2 [Cu(NH3)4](OH)2

Cu + [Cu(NH3)4](OH)2 <---> 2 [Cu(NH3)2]OH

Side reaction:

2 NH3 (aq) + 3 O2 + [Cu(NH3)4](OH)2 --> [Cu(NH3)4](NO2)2 + 4 H2O

and I suspect further oxidation to Copper ammonium nitrate as well.

As a result, I added some dilute H2O2 to add oxygen to speed up the reaction.

To my surprise, the pennies became readily covered with O2 in agreement with a cathodic reduction reaction of oxygen at the copper's surface per the electrochemical dissolution model. The reaction is also apparently exothermic as the solutions became warmer. Within an hour, a dark blue was apparent. In 8 hours, a different lighter shade of blue was apparent that is characteristic of the usual cupric salts. Expected products could include tetraamminediaquacopper(II) dihydroxide, [Cu(NH3)4(H2O)2](OH)2, as well a monohydroxide, tetraamminecopper(II) nitrite and also the nitrate. Note, the reaction produced more gas than I suspected (do not used sealed vessels) perhaps due to the formation of both O2 and N2 (via some nitrite formation and decomposition reaction).

Conclusion, if you wish to form a cupric salt, this appears to be a fairly quick and inexpensive procedure using dilute ammonia, 3% H2O2 and copper.

Caution: The presence of Copper Ammonium nitrite and/or Ammonium nitrite may present a potential spontaneous nitrogen gas decomposition issue, which are more likely in slightly acidic or concentrated solutions. I would also be concerned on heating an acidified form of the solution just prepared due to known stability issues with hot aqueous NH4NO3 in the presence of metallic impurities (including Copper, Tin and Nickel see http://www.google.com/url?sa=t&rct=j&q=ammonium%20ni... ).
--------------------------------------------------

Now, having forming one or more cupric salts, to prepare Cupric acetate, just add NaHCO3 and filter out the CuCO3. Add Acetic acid to the washed precipitate to form Cupric acetate.
--------------------------------------------------

[EDIT] Here is a less authoritative 2011 study ("Copper-Mediated Non-Enzymatic Formation of Nitrite from Ammonia and Hydrogen peroxide at Alkaline pH" ) that is pertinent relating to nitrite formation noted above (please see http://www.google.com/url?sa=t&rct=j&q=reaction%20of%20nh3%2Ch2o2%20and%20cu&source=web&cd=4&ved=0CDwQFjAD&url=http%3A%2F%2Fsp hinxsai.com%2Fvol3.no2%2Fchem%2Fchempdf%2FCT%3D23(646-656)AJ11.pdf&ei=iS-mUfCNN4nr0gGYw4D4BA&usg=AFQjCNFaObAi5_3NNOdt8e1DiRoiHzg9bg&bvm=bv .47008514,d.dmQ ). To quote:

"Hydrogen peroxide with lowest recorded redox
potential of - 0.68 V compared to that of Cu++ / Cu+, +
0.15 V15 acts as a strong reducing agent particularly in
presence of hydroxide ions [13], [18] to donate electrons to
copper (II) forming copper (I) oxide,

H2O2 + 2 OH- → 2 H2O + O2 + 2 e- (1)
2 Cu++ + 2 e- + H2O2 → Cu2 O + H2O (2)

Reddish-yellow cuprous oxide is rendered colorless in
presence of sufficient ammonia to form
diamminecopper (I) [15],

Cu2 O + 2 NH4OH → 2 [Cu (NH3)2] OH + H2O (3)

[ not balanced, corrected per ajkoer:
Cu2O + 4 NH3 + H2O → 2 [Cu(NH3)2]OH (3)]

Diamminecopper (I), generated from reduction of
copper (II) or added exogenously facilitates oxidation
of ammonia, a reducing agent [14], by hydrogen
peroxide,

...[Catalyst].....Cu (NH3)2]OH.........................
NH3 + 3 H2O2 -----------------> HNO2 + 5 H2O (4)

[ not balanced, corrected by ajkoer:
NH3 + 3 H2O2 -----------------> HNO2 + 4 H2O (4)]

Further studies are required to elucidate the actual role
of diamminecopper (I) in the reaction; whether it is
converted to tetramminecopper (II), or undergoes a
reversible changes during the process."

With additional ammonia, the reaction with nitrous acid proceeds as follows:

HNO2 + NH3•H2O --> NH4NO2 + H2O

Interesting observations by the author includes "The reaction is mediated by copper (II) as it fails to occur in absence of copper", and that the best order of addition of reactants is Cu then aqueous NH3 and finally H2O2. The author also notes the need for excess ammonia, to quote: "as it is needed to maintain: (i) solubility of copper; (ii) optimal alkalinity for expression of reducing potential of hydrogen
peroxide; (iii) adequate concentration of free ammonia; and (iv) conversion of nitrous acid to ammonium nitrite."

[Edited on 30-5-2013 by AJKOER]

subsecret - 6-8-2013 at 10:12

You can damage the fauna in your septic tank if you throw large quantities of copper salts down the drain. Same goes for sodium hypochlorite.

Copper Acetate Synthesis

thebean - 13-11-2013 at 16:00

I've never seen this synthesis before and I attempted the other day. I believe it worked and thought I'd share my findings. I'm going to transcribe the notes.

31.2 grams of CuSO4·5H2O are weighed and dissolved in minimal water with ice. 9.99 grams of NaOH are added over the course of a few minutes. The solution was then filtered. Cu(OH)2 is an is a blue insoluble compound. After filtration the Cu(OH)2 was added to a beaker and acetic acid (6% vinegar) was added until no reaction occurred. The solution goes a dark green and then after sitting for a while turns blue which indicates Cu2+ in solution. This is then boiled down to a solid.

You could use higher concentration acetic but I don't have a lot of glacial acetic acid and didn't want to use it on this synthesis so I used a lot of vinegar instead of a lot less glacial. The acetic acid strips the hydroxide off of the Cu(OH)2 and gives you Cu2+ in solution and water. Once you boil the solution down the Cu2+ ion attaches itself to the C2H3O2- ion and forms copper (II) acetate. I've seen a lot on vinegar, pennies and hydrogen peroxide being boiled but this gives you impure copper acetate and a mediocre yield.

violet sin - 13-11-2013 at 18:12

I just made a nice pile of copper acetate by taking all the sludge from copper hydroxide electrolysis experiments, and leaving it in vinegar. After reducing water content a few times, then adding more vinegar to get all the sludge dissolved, I filtered and crystallized it. Apparently I still had some sulfate in there from the Epsom salt electrolyte. This was evident from crust on top of the sol. When viewed with a back light. ~+80% teal crystals with a few much more blue rhombus patterns every so often. But I ended up with a lill more than two pill bottles worth of nice sized crystals. Rock salt-ish size, maybe a bit smaller. Fun and at least useful in the future as opposed to mixed oxides and hydroxide. I used a coffee carafe on low heat with a computer fan blowing on the surface took a while, but I didn't end up with more hydroxide/carbonate by decomp. The last lill bit had some of the magnesium salts in it, but the first 4 or so crystallizations never went to dry, so they were much cleaner.

I was hoping to leave some to crystallize to much larger size like in the 3rd page of pretty pictures 2 thread.

cyanureeves - 13-11-2013 at 18:37

nice!i will give this a try for sure thebean because it took me weeks to get just a few tiny crystals of copper acetate using peroxide and vinegar and natural summer heat.

thebean - 13-11-2013 at 20:15

Violet Sin, you could make larger crystals by putting the solution of copper acetate in a more moist environment because this slows down the crystallization process and gives you larger crystals. I would also make sure that the solution is as saturated as possible and then add a seed.

PHILOU Zrealone - 14-11-2013 at 05:08

Simply leave metallic Cu wire or pieces into a closed jam glass jar with white vinegar up to 1/4 height of the recipient.

The metallic Cu wire have to come from the liquid phase to the aerian phase!
The air is important above the liquid!

Swirl the closed jar from time to time and open from time to time to allow fresh air to come inside.
The oxygen is consumed in the process so the jar will be in slight depression.

A nice blue colour will appear in a day, then green colour will darken from day to day until you get concentrated Cu acetate.

Cu is oxydisable by air and forms a layer of CuO and Cu(OH)2 what is dissolved by the vapours of water and acetic acid. The naked metal is then further oxydised and the cycle can continue...theorically until all Cu has dissolved (if enough reactants are present).

Same happens with metallic Cu and concentrated NH4OH. You can form concentrated Cu(NH3)4(OH)2 solution and crystallization.
The Eurocents (1, 2 and 5) Cu layer (Cu plating) can be completely dissolved that way leaving the silvery steel core of the coins.

Metacelsus - 14-11-2013 at 05:50

I like to use basic copper carbonate (from copper sulfate and sodium carbonate) to react with acetic acid. Sodium carbonate is easier to obtain for me than sodium hydroxide.

cjevancich - 16-11-2013 at 02:35

I just wanted to add, since I didn't notice it already mentioned, that Copper acetate has the (annoying) property of decomposing under heat. Therefore, if you synthesize it, in a large, aqueous solution, you will have to slowly evaporate it. You cannot get away with boiling off the water. At best, you can keep the water warm with something like a crock pot on a low setting. If you try to boil it, the Copper acetate will decompose into Copper hydroxide and you'll lose acetic acid.

Arthur Dent - 18-11-2013 at 05:46

I like the Copper Carbonate method, even if it's a bit labour intensive.

After dissolving pure copper wire in HCl (which at first, looks like a brownish, turbid solution until it oxidizes into a bright emerald green solution), I neutralize it by adding sodium bicarbonate, which produces an insoluble pale greenish blue Copper Carbonate precipitate and a Sodium Chloride solution. It's a bitch to filter so my suggestion is to let the precipitate settle, decant as much liquid as you can, add more distilled water and repeat the procedure a few more times, finally filter off the precipitate with ample quantities of water. The resulting copper carbonate mud can then be poured on a plate of glass to dry, then scraped off. The other annoyance is that if you don't do several water/decant cycles, then a thorough filtering, there will still be a lot of Sodium Chloride/Bicarbonate impurities in the Copper Carbonate.

The resulting light greenish blue dust readily dissolves in Acetic Acid and turns into a clear blue solution. Gently heat the solution to evaporate the water and voilà! The resulting crystals will look like dark blue sand. I tried this experiment with glacial acetic acid, never tried with plain vinegar (but it should work equally, although there would be a lot of water to evaporate).

Robert

Robert

DraconicAcid - 18-11-2013 at 09:45

Copper(II) carbonate (actually a basic carbonate) is a goopy mess, but it becomes much more tractable if you wash it with acetone. Same with copper(II) oxide.

BobD1001 - 18-11-2013 at 11:39

I made some relatively pure Copper Acetate recently. Simply dissolved some pure copper wire in dilute glacial acetic acid, with added H2O2 to force the reaction. Let it dissolve for several days until no more reaction was noted. Then I simply vacuum dessicated it,and got some nice dark copper acetate crystals.

Ascaridole - 19-11-2013 at 02:14

Many years ago I remember making copper (II) acetate from copper (II) oxide formed via thermal decomposition of copper (II) carbonate in an oven. The still hot copper oxide was then added to concentrated acetic acid and crystals began to form quite rapidly. The acetic acid was not glacial as I distilled it from sodium acetate and sodium bisulfate that was not perfectly dry. Most likely I formed the monohydrate.

The stupid part was I attempted to purify the already beautiful crystals in a dilute acetic acid solution and over heated the solution and began forming basic copper salts. Never did recover my beautiful crystals. Lesson learned, copper likes to form basic salts.

Perhaps a repeat of the experiment is in order...

Sciencevamos - 25-1-2014 at 19:27

To make pure Cu(CH3COO)2 you need copper metal vinegar and acetic acid, you could also mix copper carbonate and vinegar. A third alternative would be to pass an electric current threw an aqueous solution of calcium acetate with 2 copper electrodes.

Here's a link to youtube on how to make it: http://www.youtube.com/watch?v=mOpxpoOemN0

mr.crow - 25-1-2014 at 21:14

Just a note that vinegar is not suitable for chemistry use

I have a beaker of copper acetate solution that smells like burnt fish and chips with lots of weird fine precipitate. There's probably lots of proteins left from the vinegar fermentation process.

As for the video putting beakers on those stove burner hotplates makes me sad

Sciencevamos - 30-1-2014 at 17:49

To make pure Cu(CH3COO)2 you need copper metal vinegar and acetic acid, you could also mix copper carbonate and vinegar. A third alternative would be to pass an electric current threw an aqueous solution of calcium acetate with 2 copper electrodes.

Here's a link to youtube on how to make it:
http://www.youtube.com/watch?v=6-4xhVWxLnE

DraconicAcid - 30-1-2014 at 17:51

Quote: Originally posted by mr.crow

Just a note that vinegar is not suitable for chemistry use

I have a beaker of copper acetate solution that smells like burnt fish and chips with lots of weird fine precipitate. There's probably lots of proteins left from the vinegar fermentation process.

As for the video putting beakers on those stove burner hotplates makes me sad

White vinegar's probably fine- malt vinegar is right out.

Sciencevamos - 2-2-2014 at 10:19

http://www.youtube.com/watch?v=vc5oR2KlggQ

Sciencevamos - 8-2-2014 at 12:37

https://www.youtube.com/watch?v=ppWguhJCMGw
how to quikly make it

BOBFOOSER - 21-2-2014 at 08:15

Hi, guys!

I've mixed the vinegar with the hydrogen peroxide, but instead of turning into a blue solution, it gets kinda orange-red. Do you have any idea why this happens?

BOBFOOSER - 21-2-2014 at 08:16

Oh, I forgot to say, that the liquid got pretty thick too.

macckone - 21-2-2014 at 10:02

hydrogen peroxide + food grade vinegar = a lot of nasty crap

food grade vinegar needs to be purified before use.
distillation will get rid of most proteins and sugars but not
some other stuff.
purification through freezing is documented in a number of
threads but may not get rid of sugars and some other stuff.

Also instead of mixing the vinegar and peroxide
you may want to drip the peroxide into your bath
with copper and vinegar.

BOBFOOSER - 21-2-2014 at 11:42

Oh, I see! Thanks for the tip, man!

Also, I've found that brazilian coins are made of bronze! Does that have anything to do with the final color of the liquid?

testimento - 21-2-2014 at 12:55

Vinegar is very suitable for chemistry if "white" or "distilled" is used. The flavoured crap is not very good, of course, depending on the use. Because it's availability I'm using it for most purposes and products can be usually quite easily purified by fractional crystallization, saturation or evaporation.

For high purity acetic acid vinegar though, distillation is preferred.