RogueRose
International Hazard
Posts: 1592
Registered: 16-6-2014
Member Is Offline
|
|
Copper Acetate reducing (to CuO) while boiling? It just happened...
Something very odd happened when I was processing some copper acetate solution. I had about a gallon of copper acetate solution almost at saturation
at room temp. Solubility is 72g/L with "cold water and 200g/L @ boiling, so let's just assume that it was about 100g/L or so at room temp. I put 1/2
gallon in a ceramic coated aluminum pan (it's excellent for boiling down solutions as things don't seem to stick) and started boiling. After about 15
mins & about 500ml boiled off, I noticed the solution was black with a slight bit of blue. I extracted some solution and it was very dark and it
looked like it was filled with CuO which seemed impossible.
I've used this pot many times for this same process and many other salts. About a week before I did this exact same process and the result was very
small/fine beautiful crystals after reducing 4L of 200g/L solution. I did everything the same as the recent process where it reduced to CuO.
Now I poured out the liquid with the CuO in it and noticed the bottom was covered with black. Now from my calculations, the solution shouldn't have
even reached saturation by this point as it would have to boil off almost 2L before it reached a 200g/L concentration, at boiling, where the crystals
should start to fall out.
So I have to figure out how the temp was reached to decompose. The BP is determined to be 464F (anhydrous I'm guessing) but with the water present it
should have stayed in hydrate form. I just can't imagine how it could have reached this temp while in the water, even if boiling. What is even more
confusing is since I did this exact thing just a week ago and got the best results yet with copper acetate, not a bit of CuO and the most perfect
precip I've had of this acetate (it can be a bit of a PITA to work with from my experience).
Does anyone have any ideas of why this may have happened or similar experiences with this salt or others?
|
|
|
DraconicAcid
International Hazard
Posts: 4332
Registered: 1-2-2013
Location: The tiniest college campus ever....
Member Is Offline
Mood: Semi-victorious.
|
|
That's not reduction. If the solution does not have excess acetic acid in it, it will hydrolyze to give copper hydroxide, which will tunr to CuO when
heated.
Please remember: "Filtrate" is not a verb.
Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
|
|
|
RogueRose
International Hazard
Posts: 1592
Registered: 16-6-2014
Member Is Offline
|
|
Quote: Originally posted by DraconicAcid  | | That's not reduction. If the solution does not have excess acetic acid in it, it will hydrolyze to give copper hydroxide, which will tunr to CuO when
heated. |
That makes sense! The first attempt was after dissolving the copper in acid and there was excess acid, the second attempt was with the acetate that
was sub par or left over from the first processing even though I did filter well well.
Thank you.
Also, do you know if it is possible for copper acetate and or acetic acid & H2O2 can dissolve cotton? I had a strange result when making some
copper acetate with vinegar and H2O2 and it looked like white plastic when the blue liquid evaporated & was mixed in with the acetate crystals.
[Edited on 4-8-2018 by RogueRose]
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
My speculation a pin hole exposing aluminum to the Cu(ll) solution:
2 Al + 3 Cu(ll) --> 2 Al(lll) + 3 Cu
Cu + Cu(ll) --H+ --> Cu(l)
H2O = H+ + OH-
And/or, referencing the educational literature, see https://en.m.wikibooks.org/wiki/A-level_Chemistry/AQA/Module... which notes that "Aqueous solutions of transition metal ions are thus acidic'" as
a result of there tendency to form ligands which undergo deprotonation, like, for example:
[Cu(H2O)6]2+ (aq) + H2O (l) = [Cu(H2O)5(OH)]+ (aq) + H3O+ (aq)
Then, the following electrochemical reaction proceeds in the presence of air/oxygen by boiling in air:
Cu(l) + O2 + 2 H+ --> Cu(ll) + 2 OH-
Reference: In the case of iron in place of copper, see Equation (2) at https://pubs.acs.org/doi/abs/10.1021/ja01600a004 . Also my comments/sources on the commercial preparation of basic cupric chloride at http://www.sciencemadness.org/talk/viewthread.php?tid=80874#... .
You may have created a precipitate of basic copper acetate.
[Edited on 12-4-2018 by AJKOER]
|
|
|
DraconicAcid
International Hazard
Posts: 4332
Registered: 1-2-2013
Location: The tiniest college campus ever....
Member Is Offline
Mood: Semi-victorious.
|
|
The observed reaction can be explained without reference to redox reactions.
Please remember: "Filtrate" is not a verb.
Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
I, for one, would be interested in more details and sources on exactly how that could occur.
As adding oxygen to select metals whose salts are known to be reactive to oxygen, consuming H+, and are further known to produce basic salts
therefrom, that is a strong hypothesis.
If this reaction with water is your 'hydrolysis' argument:
Cu(H2O)6]2+ (aq) + H2O (l) = [Cu(H2O)5(OH)]+ (aq) + H3O+ (aq)
how does creating H+ lead to creating basic salts?
|
|
|
DraconicAcid
International Hazard
Posts: 4332
Registered: 1-2-2013
Location: The tiniest college campus ever....
Member Is Offline
Mood: Semi-victorious.
|
|
Because acetate ion is a significant base, and because he's boiling it, which means that acetic acid is capable of leaving solution by evaporation.
This is why anyone who grows crystals of copper(II) acetate will add excess acetic acid to prevent this from happening.
If you're so sure that there must be some redox pathway involved, why don't you whip up a solution of copper(II) acetate in a glass container, and
boil it for a while to see if you get any precipitate?
Please remember: "Filtrate" is not a verb.
Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
| Quote: Originally posted by DraconicAcid | Because acetate ion is a significant base, and because he's boiling it, which means that acetic acid is capable of leaving solution by evaporation.
This is why anyone who grows crystals of copper(II) acetate will add excess acetic acid to prevent this from happening.
.....
|
Yes, that would generally be a valid argument!
However, apparently acetate is an exception. See discussion "Is acetate a weak base due to its resonance structure?" at https://chemistry.stackexchange.com/questions/8087/is-acetat... .
|
|
|
DraconicAcid
International Hazard
Posts: 4332
Registered: 1-2-2013
Location: The tiniest college campus ever....
Member Is Offline
Mood: Semi-victorious.
|
|
That's a stupid discussion. The conjugate base of a weak acid is a weak base, regardless of resonance structure. The conjugate base of a strong acid
is a feeble base (and will have no basic properties in aqueous solution), and the conjugate base of a feeble acid will be a strong base.
[Edited on 12-4-2018 by DraconicAcid]
Please remember: "Filtrate" is not a verb.
Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
Here is a better discussion but stills I admit does not completely negate your claim:
Question: Why is an acetate ion a weak base? We know that the conjugate base of a weak acid is a strong base, because weak acids are less readily to
give off a proton but its conjugate base can accept a proton easily.
Answer by Barry Gehm, Asst. Prof. Of Chemistry/Biochemistry:
"The strength of a acid and its conjugate base (or vice versa) are inversely proportional: the stronger the acid the weaker the base, and vice versa.
But that does not mean that the conjugate base of a weak acid is strong in the absolute sense, only that it is stronger that the conjugate base of a
stronger acid. For instance, acetate is a stronger base than chloride (conjugate base of HCl) but a weaker base than bicarbonate (conjugate base of
carbonic acid). And both acetate and bicarbonate are weaker bases than hydroxide, because acetic acid and carbonic acid are stronger acids than
water."
Link: https://www.quora.com/Why-is-an-acetate-ion-a-weak-base-We-k...
---------------------------------------------------------------
Note, boiling ZnCl2 in air, where the chloride, being the conjugate base of the strong acid HCl, is a weak base, weaker than acetate, even can result
in a basic salt, see http://www.sciencemadness.org/talk/viewthread.php?tid=70303#... , to the surprise of Blogfast25, which is why I question this argument.
Interestingly, Blogfast25 may also have had an iron impurity.
Note also, zinc (not a transition metal but a post-transition metal) is a high on the anodic index above iron, which may have also been present. The
electrochemical redox, I cited, can occur in a metal/air battery. Apparently, iron, and to a lesser extent copper, are both viewed as anodic metals
(see http://www.zygology.com/cms/upload_area/pdf/Zyg-Anodic-Index... ).
A correct characterization of my argument is pay attention to the anodic nature of the metal (and also possible impurities).
[Edited on 13-4-2018 by AJKOER]
|
|
|
TheMrbunGee
Hazard to Others
Posts: 364
Registered: 13-7-2016
Location: EU
Member Is Offline
Mood: Phosphorising
|
|
https://www.youtube.com/watch?v=UrA3zmI2pZM
This is how I destroyed my batch of 100+g Copper acetate.. just put it in nearly boiling water..
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
I bet you used tap water as you are intending to destroy the Copper acetate. Normal tap water has some ferrous, other transition metals like Mn and a
little dissolved oxygen. My take on a possible path:
First, the action of water on an aqua cupric complex:
[Cu(H2O)6]2+ (aq) + H2O (l) = [Cu(H2O)5(OH)]+ (aq) + H3O+ (aq)
Second, a redox couple equilibrium reaction leading to a presence of cuprous:
Cu(ll) + Fe(ll) = Cu(l) + Fe(lll)
which is acted on by oxygen bubbles in near boiling water consuming H+ based on the net reaction derived from a 2013 radical reaction supplement,
"Impacts of aerosols on the chemistry of atmospheric trace gases: a case study of peroxides radicals"', by H. Liang1, Z. M. Chen1, D.
Huang1, Y. Zhao1 and Z. Y. Li, link: https://www.google.com/url?sa=t&source=web&rct=j&... :
R24 O2(aq) + Cu+ → Cu2+ + O2− ( k = 4.6xE05 )
R27 O2− + Cu+ + 2 H+ → Cu2+ + H2O2 ( k = 9.4xE09 )
R25 H2O2 + Cu+ → Cu2+ + .OH + OH− ( k= 7.0 xE03 )
R23 .OH + Cu+ → Cu2+ + OH− ( k = 3.0×E09 )
Net reaction: O2 + 4 Cu+ + 2 H+ → 4 Cu2+ + 2 OH-
Electrolysis reference: See p. 7 at https://www.utc.edu/faculty/tom-rybolt/pdfs/electrochem2014.... for the reverse reaction with 2 H+ adding to each side. Alternate source of the
above reaction, per my records, but access to the full article is no longer free, see: https://www.researchgate.net/publication/262451840_Review_of... .
[Edited on 14-4-2018 by AJKOER]
|
|
|
TheMrbunGee
Hazard to Others
Posts: 364
Registered: 13-7-2016
Location: EU
Member Is Offline
Mood: Phosphorising
|
|
| Quote: Originally posted by AJKOER |
I bet you used tap water as you are intending to destroy the Copper acetate. Normal tap water has some ferrous, other transition metals like Mn and a
little dissolved oxygen. My take on a possible path:
First, the action of water on an aqua cupric complex:
[Cu(H2O)6]2+ (aq) + H2O (l) = [Cu(H2O)5(OH)]+ (aq) + H3O+ (aq)
Second, a redox couple equilibrium reaction leading to a presence of cuprous:
Cu(ll) + Fe(ll) = Cu(l) + Fe(lll)
which is acted on by oxygen bubbles in near boiling water consuming H+ based on the net reaction derived from a 2013 radical reaction supplement,
"Impacts of aerosols on the chemistry of atmospheric trace gases: a case study of peroxides radicals"', by H. Liang1, Z. M. Chen1, D.
Huang1, Y. Zhao1 and Z. Y. Li, link: https://www.google.com/url?sa=t&source=web&rct=j&... :
R24 O2(aq) + Cu+ → Cu2+ + O2− ( k = 4.6xE05 )
R27 O2− + Cu+ + 2 H+ → Cu2+ + H2O2 ( k = 9.4xE09 )
R25 H2O2 + Cu+ → Cu2+ + .OH + OH− ( k= 7.0 xE03 )
R23 .OH + Cu+ → Cu2+ + OH− ( k = 3.0×E09 )
Net reaction: O2 + 4 Cu+ + 2 H+ → 4 Cu2+ + 2 OH-
Electrolysis reference: See p. 7 at https://www.utc.edu/faculty/tom-rybolt/pdfs/electrochem2014.... for the reverse reaction with 2 H+ adding to each side. Alternate source of the
above reaction, per my records, but access to the full article is no longer free, see: https://www.researchgate.net/publication/262451840_Review_of... .
[Edited on 14-4-2018 by AJKOER] |
I love when such simple action has so much under it. Yes, tap water it was!
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
Good guess on my part, but I would not be surprised if some disagree with my suggested path.
Interestingly, I was looking at "A Dictionary of Chemistry and the Allied Branches of Other Sciences", Volume 1, p.14 by Henry Watts, at
https://books.google.com/books?id=A-dYAAAAYAAJ&pg=PA14&a... . I noticed the interesting property of chromium acetate, to quote:
" It forms red transparent crystals, which when moist absorb oxygen very rapidly from the air, undergoing a true combustion."
Now, as Chromium has many valence states (see https://en.wikipedia.org/wiki/Chromium ), together with being a somewhat anodic metal (less than zinc and more than copper) along with also some
transition metal properties, it would seem to be a good candidate for my argued oxygen reaction and apparently is.
[Edited on 14-4-2018 by AJKOER]
|
|
|
DraconicAcid
International Hazard
Posts: 4332
Registered: 1-2-2013
Location: The tiniest college campus ever....
Member Is Offline
Mood: Semi-victorious.
|
|
Because your suggested path is still nonsense.
Please remember: "Filtrate" is not a verb.
Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
Draconic Acid:
Possibly, but how could Chromium acetate have such an appetite for O2 if not an electrochemically promoted reaction?
It reminds me of a report of how sewer workers were asphyxiated when a iron pipe was drained exposing a hungry iron surface!
[Edited on 15-4-2018 by AJKOER]
|
|
|
DraconicAcid
International Hazard
Posts: 4332
Registered: 1-2-2013
Location: The tiniest college campus ever....
Member Is Offline
Mood: Semi-victorious.
|
|
| Quote: Originally posted by AJKOER | | Possibly, but why could Chromium acetate have such an appetite for O2 if not an electrochemically promoted reaction? |
Chromium(II) acetate is very air-sensitive, and that is an electrochemical reaction. But it's not relevant here.
Please remember: "Filtrate" is not a verb.
Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
To me your claimed 'not relevant' net reaction:
O2 + 4 Cu+ + 2 H+ → 4 Cu2+ + 2 OH-
screams basic salt formation!
and is actually part of the cited reaction path for the commercial preparation of, for example, basic copper chloride (see Eq 7 at https://en.wikipedia.org/wiki/Dicopper_chloride_trihydroxide ).
|
|
|
DraconicAcid
International Hazard
Posts: 4332
Registered: 1-2-2013
Location: The tiniest college campus ever....
Member Is Offline
Mood: Semi-victorious.
|
|
But that's not the net reaction, because we're starting with copper(II) ions, which you are assuming get reduced to copper(I) by iron.
Please remember: "Filtrate" is not a verb.
Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
| Quote: Originally posted by DraconicAcid |
....
But that's not the net reaction, because we're starting with copper(II) ions, which you are assuming get reduced to copper(I) by iron.
|
Or, the cuprous is already present (assuming as a soluble Cu(l) acetate complex formed with acetic acid) as apparently if you have actually prepared
cupric (and not cuprous) in the presence of copper metal with acetic acid per:
Cu + Cu(ll) ---> 2 Cu(l)-complex
some cuprous can be formed from copper acting on Cu(ll). Here is a quote from the Watts reference summarizing the two step reaction:
"The same compound [referring to a basic copper acetate] is obtained by exposing copper plates to damp air in contact with normal acetate of copper
made into a paste with water."
This would be a parallel path to the cited formation of basic cupric chloride per Eq. 6 at https://en.wikipedia.org/wiki/Dicopper_chloride_trihydroxide with chloride serving as the complexing agent for Cu(l) created from the action of
copper metal and cupric, followed by Eq 7.
Note, the actual amount of soluble Cu(l) present (from whatever path) would limit the amount of basic copper acetate created as a visible precipitate.
[Edited on 15-4-2018 by AJKOER]
|
|
|
DraconicAcid
International Hazard
Posts: 4332
Registered: 1-2-2013
Location: The tiniest college campus ever....
Member Is Offline
Mood: Semi-victorious.
|
|
| Quote: Originally posted by AJKOER |
Or, the cuprous is already present (assuming as a soluble Cu(l) acetate complex formed with acetic acid) as apparently if you have actually prepared
cupric (and not cuprous) in the presence of copper metal with acetic acid per:
Cu + Cu(ll) ---> 2 Cu(l)-complex
some cuprous can be formed from copper acting on Cu(ll). |
Copper(I) acetate is unstable and water-sensitive. It's not going to spontaneously form in aqueous solution.
Also, the OP started with copper(II) acetate, with no elemental copper present.
Please remember: "Filtrate" is not a verb.
Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
|
|
|
Tsjerk
International Hazard
Posts: 3032
Registered: 20-4-2005
Location: Netherlands
Member Is Offline
Mood: Mood
|
|
@AJOEKER; this reaction also happily happens with the cleanest reagents possible.
There is hydroxide in water, this precipitates copper hydroxide, which decomposes to the oxide. Very simple but complete explanation of the process.
So the first comment was correct and complete:
| Quote: Originally posted by DraconicAcid | | That's not reduction. If the solution does not have excess acetic acid in it, it will hydrolyze to give copper hydroxide, which will tunr to CuO when
heated. |
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
| Quote: Originally posted by RogueRose | Something very odd happened when I was processing some copper acetate solution. I had about a gallon of copper acetate solution almost at saturation
at room temp. Solubility is 72g/L with "cold water and 200g/L @ boiling, so let's just assume that it was about 100g/L or so at room temp. I put 1/2
gallon in a ceramic coated aluminum pan (it's excellent for boiling down solutions as things don't seem to stick) and started boiling. After about 15
mins & about 500ml boiled off, I noticed the solution was black with a slight bit of blue. I extracted some solution and it was very dark and it
looked like it was filled with CuO which seemed impossible.
I've used this pot many times for this same process and many other salts. About a week before I did this exact same process and the result was very
small/fine beautiful crystals after reducing 4L of 200g/L solution. I did everything the same as the recent process where it reduced to CuO.
....................
|
The argument that the precipitate occurred due to a base always presence in water is not consistent with the claim that this was an exceptional run as
previously no such reaction occurred.
If there was some exposure to the underlying aluminum pan, this could imply elemental copper formation. Copper metal plus cupric plus acetate (to
serve as a complexing agent) could then form some cuprous. To repeat the quote from Watts:
"The same compound [referring to a basic copper acetate] is obtained by exposing copper plates to damp air in contact with normal acetate of copper
made into a paste with water."
Watts states on cuprous acetate:
"Water decomposes it into normal cupric acetate and yellow cuprous hydrate."
to which I concur if the water contains any dissolved oxygen.
Note that Lead acetate per Watts is claimed to be highly sensitive to the presence of CO2 in air!
--------------------------------------------
If I have correctly guessed the source of the cuprous formation (aluminum exposure when boiling in an open pan), it should repeat with the next run,
unless it was unusual contact with the shiny aluminum alloy lip of the ceramic coated aluminum pan which could have been overly filled, or from an
unusual violent boiling/frothing event.
[Edited on 16-4-2018 by AJKOER]
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
If no contact with the shiny pan's top, perhaps the metal fasteners to hold the pans handle (see, for example, https://www.jcpenney.com/p/cooks-ceramic-12i-skillet/ppr5007... )?
|
|
|
RogueRose
International Hazard
Posts: 1592
Registered: 16-6-2014
Member Is Offline
|
|
I think this is an older style pot, maybe 60's to 70's cornell ware. The entire pot is one piece and dipped into ceramic. The only exposed part is
a circle that touches the element, the rest is one solid, thick coating of ceramic. I picked this pot because the way it was made and it is easy to
clean.
|
|
|
![[*]](images/xpblue/default_icon.gif){kind=icon}
