Sciencemadness Discussion Board

Copper Acetate Fiasco

CHRIS25 - 16-7-2012 at 00:59

I added:

19 grams copper
56 mL 40% conc. Acetic acid
19.5 mL H2O2

Left for a few hours and the solution turned an expected deep Blue

Observations: unexpected Mottled turquoise/black colour on 20% of the copper surface. Reaction stopped. no bubbles detected.

Added a further 25 Ml H2O2.

Reaction kick started and the black on surface of copper turned turquoise then slowly back to black again after a few hours. Reaction stopped after 10 hours.

Added a further 30 mL acetic acid, the black disappeared after four hours. Began heating solution for three hours, a brown heavier than water precipitate formed.

Gave up on whole solution and poured into a jar to use as is anyway. After a couple of hours there was a turquoise/blue residue on the inside of the glass, the solution is still deep blue as expected but the residue that has formed is difficult to scrape off although it can be removed.

Conclusion: Have not got a clue what is happening here, could find absolutely NO references at all about the making of acetic acid on the web (only the home - making blue solutions and a few kitchen videos - nothing professional)

The whole reaction consumed 4.2 grams of copper - the problem I have here is what is the point of 19 grams of copper.

Help needed please.





[Edited on 16-7-2012 by CHRIS25]

watson.fawkes - 16-7-2012 at 04:26

What concentration is the H2O2? It's acting as an oxidizer for the copper, so it matters how much of it there is.

m1tanker78 - 16-7-2012 at 05:34

What's your source of copper? Part of the bubbling you see could be catalytic decomposition of hydrogen peroxide which obviously spells premature death to your oxidizer.

Tank

CHRIS25 - 16-7-2012 at 08:57

Quote: Originally posted by watson.fawkes  
What concentration is the H2O2? It's acting as an oxidizer for the copper, so it matters how much of it there is.
Hallo, H2O2 is 6% but I calculated this in in my stoichemetry from an original:
Cu + 2CHCOOH + H2O2 = Cu(CH3COO)2 + 2H2O
63.5 + 120 + 34 = 182 + 36

The above values for peroxide and acetic are for concentrated so my maths adjusted for my own percentages.

[Edited on 16-7-2012 by CHRIS25]

CHRIS25 - 16-7-2012 at 08:59

Quote: Originally posted by m1tanker78  
What's your source of copper? Part of the bubbling you see could be catalytic decomposition of hydrogen peroxide which obviously spells premature death to your oxidizer.

Tank


Source is scrap dealer. But my cleaning procedures are thorough, they receice three separate cleans: Scrub scrub with Sodium Carbonate, then in Sulphuric (for the de- oxidising) then citric for the final rinse, and washed in de-ionized water and then another final treatment in sulphuric just for fun!.

[Edited on 16-7-2012 by CHRIS25]

zoombafu - 16-7-2012 at 10:35

I'm pretty sure your stoichiometry is off. 6% hydrogen peroxide is 6gH202/100gH20. If you are using 19g copper, you will need .30 moles of H202, that's 10g. 44.5mL of 6% H202 is not .30moles.

Also Is the copper in big chunks? If it is, breaking it up will increase the surface area for the reaction.

CHRIS25 - 16-7-2012 at 11:05

Quote: Originally posted by zoombafu  
I'm pretty sure your stoichiometry is off. 6% hydrogen peroxide is 6gH202/100gH20. If you are using 19g copper, you will need .30 moles of H202, that's 10g. 44.5mL of 6% H202 is not .30moles.

Also Is the copper in big chunks? If it is, breaking it up will increase the surface area for the reaction.

Hi well 10g equates to approx 10mL because it's basically water and the density of H2O2 is 1.135 which is 11.35ml. Since I added more than 19mL to the reaction where am I wrong here?

and sorry for my lack of mathematical know how (not being sarcastic) but how on earth did you arrive at 0.3moles by seeing the 19 grams of copper? I see that concentrated H2O2 is 34moles, and that copper is 63.5moles. 6% equals 2.04moles. But I really don't know how to calculate what you just calculated?

[Edited on 16-7-2012 by CHRIS25]

zoombafu - 16-7-2012 at 12:16

copper is 63.5g/mol
H202 (pure) is 34.0g/mol

you used 19g of copper.
19g/1mol x 1mol/63.5g = .30moles

You need one mole of H202 to react with one mole of copper
so if you are using .30 moles of copper you need .30 moles of H202
.30mol x 34.0g H202/1mol = 10. g H202

You need 10 grams of pure H202.
Your solution is a 6% solution, not pure.

Ill let you find out how much solution you need to have a total of 10g H202.

I recommend getting a first level college chemistry book, because learning stoichiometry is a must. You can get one at a used book store, or off of ebay pretty cheaply. You can also find lots of pdfs online in torrent form


CHRIS25 - 16-7-2012 at 13:44

Quote: Originally posted by zoombafu  

Ill let you find out how much solution you need to have a total of 10g H202.

>>>Hi, I am learning stoichemetry and have a couple of books already, my problem is working out maths, I try but am very very weak at thinking in a mathematical problem, I can not work out how to solve things when it comes to thinking in a mathematical way and believe me I do my best. This problem is not so much understanding stoichemetry as it is working out HOW you work something out, I spent an agonising 10 minutes on this and simply can not do it.

The best I came up with was: I have 0.6g in a 1000mL solution of H2O2, I juggled with figures and well - I'm lost here, I really do not know how to work it out, the books I have are good, but they don't help you with maths unfortunately.<<<<





Ok, just had a brain wave!! I need 56 litres of 6% H2O2 to equal 1 litre of 100%. Therefore if I need 10mL of 100% and only have 6% I need ....yep there you go lost it!!!!!!!!!!!

[Edited on 16-7-2012 by CHRIS25]

zoombafu - 16-7-2012 at 13:51

If there are 6g H202 in 100g water, then there are 12g in 200g water (only 10g are needed). When dealing with H202 in an inorganic synthesis I usually put in a slight excess, as it has the tendency to decay in storage. So 200mL of the solution should be more than enough to carry out the reaction.

CHRIS25 - 16-7-2012 at 16:04

Trouble here is that I never saw 6% as 6 grams , that's confusing to be honest, I did not think that 6% was 6 grams in 100 mL of water - this constant shifting between grams and percent is very confusing when each separate calculation is being measured with totally different parameters, ie, moles grams and percent and then these have to be standardised within the equations in order to achieve the blasted answer. Your example above I know is very good, but for me it would be more understandable and manageable if parameters were consistent. I was trying to work out what 6% would be in grams a few posts ago, can't damn well believe it Aagh - - 6% HCl is not 6 grams in 100 mL of water, it's 2 moles or 170mL for example, in 1 litre of water. or 17 mL in 100 mL water which then makes it 17% but as we know that does not make it 17% concentration or 17 grams, it's still 6% right? Can you see where I am with all this? So when one thinks with a one tracked mind I am at a disadvantage as regards Maths. I appreciate your help to be honest here and hope you forgive my utter simplicity, maths just leaves me screaming - brain is not configured for this.


So now we are at the end of this, Do I boil the solution until there is no more copper left?(chemistry books don't teach How you know when to freeze boil or just let be, they are full of calculations and atomic numbers and very interesting things I need to learn, but never give example of how you know WHAT exactly one needs to do withthe reactants in order to ensure a successful reaction...
[Edited on 17-7-2012 by CHRIS25]

[Edited on 17-7-2012 by CHRIS25]

[Edited on 17-7-2012 by CHRIS25]

Sedit - 16-7-2012 at 16:28

You can not boil a solution of Copper acetate as it will decompose as you do it.

zoombafu - 16-7-2012 at 17:11

Do some research on concentrations. Some are weight/weight, some are volume/volume. You have to look on the container, it usually says which (w/w, v/v). All the situations I have come across it is w/w. It is stupid and confusing, which is why I don't use percents. I prefer molarities.

CHRIS25 - 16-7-2012 at 23:37

Quote: Originally posted by zoombafu  
Do some research on concentrations. Some are weight/weight, some are volume/volume. You have to look on the container, it usually says which (w/w, v/v). All the situations I have come across it is w/w. It is stupid and confusing, which is why I don't use percents. I prefer molarities.
I did that months ago, no problem in isolation understanding moles and percent by volume oor weight - it's just maths and thinking in that kind of unique problem solving context that seems to have no working horsepower in my head.


My acetic acid calculations were also wrong! If I need 35.03 mL of 100% and I only have 40% then I need twice that amount +1/10 more, ie 75.66mL of acetic acid. Right?

[Edited on 17-7-2012 by CHRIS25]

CHRIS25 - 16-7-2012 at 23:43

Quote: Originally posted by Sedit  
You can not boil a solution of Copper acetate as it will decompose as you do it.
So how does the copper 2 acetate precipitate out of solution without having to wait for weeks and weeks for it to evaporate?

My assumption was that you heat it up to about boiling temperature of water (which is what meant - apologies for not being scientifically accurate, I know that boil point of Cu acetate is somewhere around 200c) and then keep it at this temp range?

But since I have been wrong about so many things lately my confidence is waning here.

watson.fawkes - 17-7-2012 at 08:01

Quote: Originally posted by CHRIS25  
So how does the copper 2 acetate precipitate out of solution without having to wait for weeks and weeks for it to evaporate?
Perhaps it cannot. This wouldn't be the first chemical whose manipulations were limited by heat sensitivity.

In this particular case, you can improve evaporation rate by using more surface area. To get large surface area, use wide flat trays. Metallic trays are likely a bad idea, given the copper ions. In addition, blowing a fan over the trays will speed the evaporation by reducing vapor saturation at the liquid surface.

You would also use vacuum evaporation, but that's a lot of gear for small batches of product, particularly when there are cheaper alternatives.

CHRIS25 - 17-7-2012 at 09:39

Quote: Originally posted by watson.fawkes  
Quote: Originally posted by CHRIS25  
So how does the copper 2 acetate precipitate out of solution without having to wait for weeks and weeks for it to evaporate?
Perhaps it cannot. This wouldn't be the first chemical whose manipulations were limited by heat sensitivity.

In this particular case, you can improve evaporation rate by using more surface area. To get large surface area, use wide flat trays. Metallic trays are likely a bad idea, given the copper ions. In addition, blowing a fan over the trays will speed the evaporation by reducing vapor saturation at the liquid surface.

You would also use vacuum evaporation, but that's a lot of gear for small batches of product, particularly when there are cheaper alternatives.
Thankyou, well I have only plastic trays, and I could squeeze one into my home-made container for dangerous and smelly chemicals, I picked a broken down commercial Coffee urn from a skip a while back, modified it, and that comes in very handy for outdoor work in the rain and wind. (which never stops here even now). But I thought if you heated the solution say keeping it around 60c after the reaction was completed you could drive off a lot of the excess water? But it appears that this would still decompose some of the acetate?

[Edited on 17-7-2012 by CHRIS25]

triplepoint - 24-8-2012 at 07:54

I have made and dried some copper acetate. It forms very satisfying crystals, but there is also some decomposition no matter how I do it. The bright side is that the decomp is a relatively small % and it is excluded by the crystals, which can be rinsed with cold water to remove the impurities.

SM2 - 24-8-2012 at 09:47

If my brain remembers, I just found the non-zinc pennies, couldn't dissolve in glacial. Dissolved in some other heavy duty acid easily. Than just swapped, and relied on solubility or lack of get my copper acetate (or was it cupric). All the copper salts I have made are all different shades of green. But those big sulphate x-tals at Ace are beautiful.