fluorescence
Hazard to Others
Posts: 285
Registered: 11-11-2013
Member Is Offline
Mood: So cold outside
|
|
Discoloration of Copper Sulfate Pentahydrate by Lithium Chloride
So I was grinding some LiCl and CuSO4 both dry and perfectly clean and suddenly it turned orange. I tested it with NaCl, KCl,
BaCl2, CaCl2 and it would never get orange. Some where quite humid with Hydrates, some dry. The LiCl is fresh from Sigma Aldrich
and was never used before. Still a dry powder, no crystals formed and it doesn't say anything about Hydrates on the bottle.
I tested it with 2 different CuSO4 samples, one crystalline Pentahydrate with Food Quality and the other one is from a pigment supply,
still pentahydrate but a very dry fine powder, not sticky or anything. I will check it with anhydrous, too.
But does anyone have an idea what it could be ? Perhaps LiCl is sucking up water really well and it gets moist and forms some CuCl2, too ?
The blue one is a mixuture with NaCl, all the other Chlorides looked like this one, too and the orange one is with LiCl.
|
|
|
Zephyr
Hazard to Others
Posts: 341
Registered: 30-8-2013
Location: Seattle, WA
Member Is Offline
|
|
Wow! What a curious reaction...
Firstly, the anhydrous copper chloride I have seen is a much darker brown, so I have my money on something else having occured.
From sigma's immense inventory, it would appear copper it capable of forming some curious complexes with lithium, and although I don't see a
cupro-lithium one with both the sulfate and the chloride, there are several other possibly similar examples: http://www.sigmaaldrich.com/catalog/product/aldrich/224308?l...
Kudos on doing such diligent testing, I'm glad you were able to rule out contamination and moisture and whatnot.
From here I recommend two things, firstly, it may be enlightening to view the reaction under a microscope to see if the end crystals are a pure
substance of if you can distinguish between two compounds, as well as seeing if any moisture is produced from the pentahydrate (some of the crystals
looked a bit damp).
In addition I am interested in comparing if similar reactions take place with other salts of copper and the same or other salts of lithium...
Hopefully I will be able to try this out in a few months, good luck on solving this intriguing puzzle.
|
|
|
fluorescence
Hazard to Others
Posts: 285
Registered: 11-11-2013
Member Is Offline
Mood: So cold outside
|
|
Actually Amos has solved this yesterday already.
He tested the same thing with a Nickel salt (turning yellow) and a Cobalt salt (turning blue). It seems like
the Lithium Chloride is sucking up water so powerful that just grinding a dry hydrate salt with the dry LiCl
will cause it to become dehydrated. Interesting property. I only have Calcium Chloride as a Hydrate but one
could buy some water free and test it out with that. Since Calcium Chloride is commerically used it is way cheaper.
|
|
|
Zephyr
Hazard to Others
Posts: 341
Registered: 30-8-2013
Location: Seattle, WA
Member Is Offline
|
|
Interesting... But I thought anhydrous copper sulfate was white? And even if the copper and lithium switched anions would it not be a dark brown?
|
|
|
DraconicAcid
International Hazard
Posts: 4278
Registered: 1-2-2013
Location: The tiniest college campus ever....
Member Is Offline
Mood: Semi-victorious.
|
|
Quote: Originally posted by Pinkhippo11  | | Interesting... But I thought anhydrous copper sulfate was white? And even if the copper and lithium switched anions would it not be a dark brown?
|
The colour looks like CuCl4(2-) ions.
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.
|
|
|
Texium
|
Thread Moved
2-1-2016 at 12:37 |
blogfast25
International Hazard
Posts: 10562
Registered: 3-2-2008
Location: Neverland
Member Is Offline
Mood: No Mood
|
|
Mixed with Na2SO4 and some water that's what the left dish could look like.
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
The orange color may be result of mixture of a red salt like Cu2O (or, a Lithium-transition metal oxide, like LiCuO2 or Li2CuO2, see reference below)
and a white salt in the context of having constructed a Lithium-air battery with the associated products formed via a transition metal (Fenton
chemistry).
For example, Wikipedia on the Lithium-air battery (link: https://en.m.wikipedia.org/wiki/Lithium%E2%80%93air_battery ) describes the metal-air battery chemistry as follows:
"In a cell with an aprotic electrolyte lithium oxides are produced through reduction at the cathode:
Li+ + e− + O2 + * → LiO2*
Li+ + e− + LiO2* →Li2O2*
where "*" denotes a surface site on Li2O2 where growth proceeds, which is essentially a neutral Li vacancy in the Li2O2 surface.
Lithium oxides are insoluble in aprotic electrolytes, which leads to cathode clogging.[26]
In a cell with an aqueous electrolyte the reduction at the cathode can also produce lithium hydroxide:
Acidic electrolyte
2Li + 1⁄2O2 + 2H+ → 2Li+ + H2O
A conjugate base is involved in the reaction. The theoretical maximal Li-air cell specific energy and Li-air cell energy density is 1400 W·h/kg and
1680 W·h/l, respectively.[9]
Alkaline aqueous electrolyte
2Li + 1⁄2O2 + H2O → 2LiOH
Water molecules are involved in the redox reactions at the air cathode. "
----------------------
Or, the surface chemistry reaction may more parallel the classic Lithium-ion battery, described per one source (see http://www.nexeon.co.uk/about-li-ion-batteries/ ), to quote:
"Overall reaction on a Li-ion cell: C + LiCoO2 ↔ LiC6 + Li0.5CoO2
At the cathode: LiCoO2 – Li+ – e- ↔ Li0.5CoO2 ⇒ 143 mAh/g
At the anode: 6C + Li+ + e- ↔ LiC6 ⇒ 372 mAh/g
Materials other than graphite have been investigated, with silicon offering the highest gravimetric capacity (mAh/g)."
In the current context employing the transition metal copper in place of cobalt, for the related Lithium salt, please see http://www.sciencedirect.com/science/article/pii/S0167273897... , "Electrochemical and structural study of Li2CuO2, LiCuO2 and NaCuO2", by Hajime
Arai, et al.
Also, per another source (see http://www.jmbatterysystems.com/technology/cells/how-cells-w... ), to quote:
"Lithium Battery cells consist of three main components :
"The anode : on discharge gives up electrons to the external circuit and is oxidised during the electrochemical reaction. Most commercial cells
currently employ a carbon/graphite based electrode; however metal or an alloy can also be used....
The cathode : on discharge accepts electrons from the external circuit and is reduced during the electrochemical reaction. It is usually a transition
metallic oxide or phosphate....
The electrolyte (an ionic conductor but electronic insulator) separates the two electrodes and provides the medium for charge transfer inside the cell
between the anode and cathode. The electrolyte is typically a non-aqueous inorganic solvent containing a dissolved Lithium salt, e.g. LiPF6 in
propylene carbonate."
Or, the answer to what is occurring could be more simple, but I really doubt it.
[Edit] Note, the reported formation with a Copper salt (orange), a Nickel salt (turning yellow) and a Cobalt salt (turning blue), but not NaCl, KCl,
BaCl2, CaCl2, suggests to me the required employment of a transition metal, and associated Fenton products derived therefrom, for the targeted complex
formation. Otherwise, the argument of a simple complex formation and ignoring the apparent selectivity of the transition metal cation, appears weak to
me (especially based on references approaching a half century old).
[Edited on 4-1-2016 by AJKOER]
|
|
|
unionised
International Hazard
Posts: 5102
Registered: 1-11-2003
Location: UK
Member Is Offline
Mood: No Mood
|
|
| Quote: Originally posted by AJKOER | The orange color may be result of mixture of a red salt like Cu2O (or, a Lithium-transition metal oxide, like LiCuO2 or Li2CuO2, see reference below)
and a white salt in the context of having constructed a Lithium-air battery.
...
...
Or, the answer to what is occurring could be more simple, but I really doubt it. |
Not really, no.
There's no strong oxidant or reductant so the copper is almost certainly not going to change oxidation state.
Why do you doubt something simple, but propose something practically impossible?
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
Unionised:
I want to agree with you, but the apparent chemistry around the Lithium-ion battery is daunting alone!
Here is yet another source "Electrochemical Energy: Advanced Materials and Technologies", edited by Pei Kang Shen, et al., page 184 at https://books.google.com/books?id=ujk0CwAAQBAJ&pg=PA184&... with cited references to formation of 'superoxide' and 'peroxide'.
Some of my recent postings discuss copper based Fenton-like reactions (see, for example and links that I have provided therein at: https://www.sciencemadness.org/whisper/viewthread.php?tid=64... ) involving O2/H2O2 with movements between the Cupric and Cuprous states, in the
presence of just sunlight (and more so, I suspect, with radical based reactions involving chlorides and uv). Even more advanced is the reference to
surface chemistry based complexes and so called (hydr)oxides.
If someone wants accurate simple chemistry, my advice is to avoid Lithium, all transition metals (especially in the presence of uv, reductants..) and
surface chemistry based reactions.
[Edit] It may be easy to test if there is any O2 based pathway by simply grinding the salts in a larger vessel filled with CO2, that is, under a
blanket of CO2.
[Edit] [Edit] Found a reference on the CuO2- anion. "Two isomers of CuO2: The Cu(O2) complex and the copper dioxide" by Hongbin Wu, Sunil R. Desai,
and Lai-Sheng Wanga), to quote from the extract:
"The CuO2 anion is also observed to undergo photodissociation to Cu- + O2 at both 532 nm and 355 nm detachment wavelengths."
and not, Cu + O2-.
Link: https://www.google.com/url?sa=t&source=web&rct=j&...
So, under photodissociation, a breakdown of LiCuO2.
[Edited on 4-1-2016 by AJKOER]
|
|
|
woelen
Super Administrator
Posts: 7976
Registered: 20-8-2005
Location: Netherlands
Member Is Offline
Mood: interested
|
|
Very interesting observation. I think it can be compared to the effect I observed with CsCl in aqueous solution, where I obtained nice red and yellow
solid complexes:
http://woelen.homescience.net/science/chem/exps/CsCuCl3/inde...
I think that you get some Li-salt of a similar complex.
|
|
|
blogfast25
International Hazard
Posts: 10562
Registered: 3-2-2008
Location: Neverland
Member Is Offline
Mood: No Mood
|
|
Your page actually mentions a paper:
A Redetermination of the Crystal Structure of CsCuCl3", written by Albert W. Schlueter, Robert A . Jacobson, and Robert E . Rundle.
That paper (abstract) can be found here:
http://pubs.acs.org/doi/abs/10.1021/ic50036a025
It actually mentions LiCuCl3.2H2O, and chain-like Cu2Cl4<sup>2-</sup> anions. And long
'unsymmetric' (?) Cu-Cl chains.
The latter would almost certainly be coloured and I'll see if the bonding structure can reveal anything.
But 'mystery solved', I think. Well done.
[Edited on 4-1-2016 by blogfast25]
|
|
|
kmno4
International Hazard
Posts: 1495
Registered: 1-6-2005
Location: Silly, stupid country
Member Is Offline
Mood: No Mood
|
|
http://lib.dr.iastate.edu/cgi/viewcontent.cgi?article=1207&a...
also:
http://scripts.iucr.org/cgi-bin/paper?a03978
Слава Україні !
Героям слава !
|
|
|
blogfast25
International Hazard
Posts: 10562
Registered: 3-2-2008
Location: Neverland
Member Is Offline
Mood: No Mood
|
|
Interesting, especially the first one.
So the structure would be [Li(H2O)2]2Cu2Cl6.
On redissolving OP's orange mixture it should probably revert to green/blue.
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
More precisely, with respect to Blogfast's comment to quote from page 9:
"The crystals were quite unstable to water. In a moist atmosphere they turned green and then, after prolonged standing, became a
green solution. It is believed the compound breaks down into its component salts, LiCl and CuC12 • 2H2O"
---------------
My opinion based on what was observed is still Cuprous oxide or related Lithium salt. For example, to quote from U.S. Patent 2,474,533 (see http://www.google.com/patents/US2474533 ):
"The cuprous oxide produced by this method is a brick red to orange color, depending on the fineness of the particles. Most of the-tests were made
using oxide that had-been dried and ground immediately after filtering, and so retained its high cuprous oxide contents. Samples of freshly
precipitated-wet filter cake oxide were also used; It was found that with freshly precipitated oxide the cupric oxide dissolved out more rapidly and
gave a product of slightly higher cuprous oxide content; however, it was a darker red color, presumably because of the larger particle size. On
grinding this material it was found that the color could be changed from brick red to lighter red, or to orange by finer grinding."
A rough delineation of a possible path is the action of O2 on LiCl2 in the presence of CuSO4 forming Li2O2. The action of water vapor forms LiOH (and
H2O2), which acting on CuSO4 forms Cu(OH)2 that may react or not further with LiOH or LiCl2. The resulting salt is dehydrated by anhydrous LiCl2 and
possibly further reduced by light to a cuprous salt.
The last property of light to transform higher valence states to lower ones is observed for Iron and Copper . Source, see, for example, https://www.researchgate.net/publication/222308414_Destructi... , and, I suspect, possible for other transition metals like Cobalt and Nickel (and
certainly Mn) in presence of H2O2, which would account for an apparent transition metal selectivity.
[Edited on 5-1-2016 by AJKOER]
|
|
|
blogfast25
International Hazard
Posts: 10562
Registered: 3-2-2008
Location: Neverland
Member Is Offline
Mood: No Mood
|
|
| Quote: Originally posted by AJKOER |
A rough delineation of a possible path is the action of O2 on LiCl2 in the presence of CuSO4 forming Li2O2. The action of water vapor forms LiOH (and
H2O2), which acting on CuSO4 forms Cu(OH)2 that may react or not further with LiOH or LiCl2. The resulting salt is dehydrated by anhydrous LiCl2 and
possibly further reduced by light to a cuprous salt.
|
Nope. With high concentrations of chloride, CuCl4<sup>2-</sup> forms. In even higher concentrations of chloride and specific
conditions Cu2Cl6<sup>2-</sup> forms.
Both have relatively low values of Kf (complex formation constant), hence the need for high Cl<sup>-</sup>
concentrations/activities.
As with many D-block complexes, Ligand Field Theory explains the colour, at least qualitatively. Similar salts of Cs and Rb also exist.
[Edited on 5-1-2016 by blogfast25]
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
Blogfast:
My comments relate to the opening context of this thread, namely the action of "Copper Sulfate Pentahydrate by Lithium Chloride".
Your point relating to high chloride ion concentration would likely be not applicable.
|
|
|
unionised
International Hazard
Posts: 5102
Registered: 1-11-2003
Location: UK
Member Is Offline
Mood: No Mood
|
|
| Quote: Originally posted by AJKOER | Blogfast:
My comments relate to the opening context of this thread, namely the action of "Copper Sulfate Pentahydrate by Lithium Chloride".
|
No they don't: wrong oxidation state.
|
|
|
blogfast25
International Hazard
Posts: 10562
Registered: 3-2-2008
Location: Neverland
Member Is Offline
Mood: No Mood
|
|
| Quote: Originally posted by AJKOER | Blogfast:
My comments relate to the opening context of this thread, namely the action of "Copper Sulfate Pentahydrate by Lithium Chloride".
Your point relating to high chloride ion concentration would likely be not applicable. |
You grind together dry LiCl and dry CuSO4.5H2O and you don't have a high chloride concentration? Who knew? 
Have you read the method of preparation in kmno4's first link?
As I wrote above, let the OP dissolve his orange mixture containing [Li(H2O)2]2Cu2Cl6 and he'll
get a greenish/blue solution.
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
Unionised:
I have added a source link to Photo-Fenton and Photo-Fenton-like reactions, which I will repeat https://www.researchgate.net/publication/222308414_Destructi... .
The latter reactions account for my needed explanation, for example, in the creation of a cuprous state. The precursors required (like H2O2) may have
been supplied, as I detailed above, via the chemistry surrounding the Lithium-air battery with the formation of Lithium peroxide.
Remove oxygen, and the reaction proceeds, my roughly suggested path is definitely wrong. Remove light, and the reaction proceeds, my roughly suggested
path is likely wrong.
Or, more easily, as Blogfast just suggested, try dissolving the orange mix in hot water.
[Edited on 5-1-2016 by AJKOER]
|
|
|
MrHomeScientist
International Hazard
Posts: 1806
Registered: 24-10-2010
Location: Flerovium
Member Is Offline
Mood: No Mood
|
|
Very interesting! I'm curious to see what happens when dissolved in water, and then what happens when that solution is evaporated.
|
|
|
DraconicAcid
International Hazard
Posts: 4278
Registered: 1-2-2013
Location: The tiniest college campus ever....
Member Is Offline
Mood: Semi-victorious.
|
|
| Quote: Originally posted by blogfast25 | | Nope. With high concentrations of chloride, CuCl4<sup>2-</sup> forms. In even higher concentrations of chloride and specific
conditions Cu2Cl6<sup>2-</sup> forms. |
Surely, higher concentrations of chloride would give the tetrachloro complex, since it has a higher ratio of chlorine to copper?
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.
|
|
|
blogfast25
International Hazard
Posts: 10562
Registered: 3-2-2008
Location: Neverland
Member Is Offline
Mood: No Mood
|
|
| Quote: Originally posted by DraconicAcid |
Surely, higher concentrations of chloride would give the tetrachloro complex, since it has a higher ratio of chlorine to copper?
|
Good point, actually. Well, not sure how it works, but clearly that is what happens: at very high chloride concentrations the dimeric complex forms.
That's what woelen's experiment showed, as well as the preparation procedure for the Li compound in kmno4's *.pdf.
[Edited on 6-1-2016 by blogfast25]
|
|
|
![[*]](images/xpblue/default_icon.gif){kind=icon}
