Copper
Harmless
Posts: 17
Registered: 24-10-2015
Member Is Offline
Mood: No Mood
|
|
Schweizer's reagent and ammonia
Ammonia is difficulty to find in the concentrated form needed to make Schweizer's reagent (30% ammonia is suggested). Would making ammonia from
ammonium carbonate work?
(NH3)2CO3 + 2NaOH -> 2NH4OH + Na2CO3
NH4OH <-->>> NH3 + H2O
Would the resulting mixture of ammonium hydroxide, ammonia and sodium carbonate still be suitable for making Schweizer's reagent? Thanks
|
|
|
MolecularWorld
Hazard to Others
Posts: 110
Registered: 30-10-2015
Member Is Offline
Mood: No Mood
|
|
The carbonate might interfere with the formation of the Schweizer's reagent, by precipitating the copper. If the solution is basic enough, quite some
copper could stay in solution as cuprates and various complexes, though I don't know if such a solution is suitable for dissolving cellulose. Try it
and see!
Do you have any other ammonium compounds? If you make a concentrated solution of an ammonium salt and sodium hydroxide, you could use a setup like the
one in this video, possibly combined with the salting out technique in this thread, to produce pure ammonia solution from it (30% might be a stretch).
[Edited on 29-11-2015 by MolecularWorld]
|
|
|
Amos
International Hazard
Posts: 1406
Registered: 25-3-2014
Location: Yes
Member Is Offline
Mood: No
|
|
I would recommend reacting the ammonium carbonate with an acid first and then crystallizing the ammonium salt. Ammonium chloride, nitrate, or sulfate
are good ones. These can be heated with sodium hydroxide to generate ammonia that can then be bubbled through ice-cold water to make strong aqueous
ammonia.
The problem with trying to use ammonium carbonate for this is that its degradation products will include carbon dioxide, which will just react with
ammonia and water again to form ammonium carbonate.
|
|
|
MolecularWorld
Hazard to Others
Posts: 110
Registered: 30-10-2015
Member Is Offline
Mood: No Mood
|
|
@Amos: Good point. If the ammonium carbonate / sodium hydroxide reaction mixture gets too hot, carbon dioxide is liable to evolve with the ammonia
instead of reacting with the hydroxide. I imagine this could be mitigated by first dissolving the sodium hydroxide in water, allowing the solution to
cool, then slowly adding the ammonium compound (the dissolution of the ammonium compound should even be endothermic). Other ammonium compounds would
be better, which is why I asked Copper if he had any.
Similarly, heating a mixture of a less-volatile ammonium compound and sodium hydroxide and bubbling the ammonia through water would also be better, if
acids, a ground glass distillation setup, and gas absorption bottles are available (or can be improvised). I suggested the Tupperware method on the
basis that it doesn't require any additional chemicals or specialized equipment, but the yield with such a method will definitely be lower, as will be
the concentration of the purified ammonia solution.
[Edited on 29-11-2015 by MolecularWorld]
|
|
|
deltaH
Dangerous source of unreferenced speculation
Posts: 1663
Registered: 30-9-2013
Location: South Africa
Member Is Offline
Mood: Heavily protonated
|
|
Another idea is to mix the required amount of ammonium carbonate and crushed ice [to ultimately yield a 30% solution] with an excess of slaked lime.
This will yield highly insoluble calcium carbonate and the slurry can be filtered to yield a concentrated ammonia solution.
(NH4)2CO3(aq) + Ca(OH)2(s) => CaCO3(s) + 2NH3(aq) + 2H2O(l)
Don't forget to factor in the water evolved by the reaction and subtract it from the ice you are putting in initially.
You can test that your ammonia reagent is carbonate free enough by adding a few drops of filtered lime water (at room temperature) to it in a test
tube. If it doesn't go milky, you got rid of the carbonate. If it does, then repeat the lime washing step and filtering.
Handy hint: Store 30% ammonia in a freezer before working with it, else it releases a $#@$-load of ammonia when opening and pouring.
[Edited on 29-11-2015 by deltaH]
|
|
|
ave369
Eastern European Lady of Mad Science
Posts: 596
Registered: 8-7-2015
Location: No Location
Member Is Offline
Mood: No Mood
|
|
You can just take any household ammonia, put it in a big flask connected to a funnel-and-beaker trap and boil. All escaping ammonia will be trapped in
the beaker, where you wil have a much more concentrated solution.
Smells like ammonia....
|
|
|
Copper
Harmless
Posts: 17
Registered: 24-10-2015
Member Is Offline
Mood: No Mood
|
|
Thanks for all the suggestions. I can get ammonium chloride. Would a conical flask for the reaction stoppered with one hole stopper (on hot plate)
connected using glass tubing to another conical flask with water in ice bath work? Also if it does should I use one hole stopper, no stopper, or two
hole stopper? Also, how much ammonia gas would dissolve (I'm worried the gas will just bubble through and not much will dissolve).
Would titration suffice to determine the concentration of the ammonia?
Thanks
|
|
|
NedsHead
Hazard to Others
Posts: 409
Registered: 9-12-2014
Location: South Australia
Member Is Offline
Mood: No Mood
|
|
This is my jerry rig Ammonia generator, https://www.youtube.com/watch?v=xFVPx4L-DG4 it works well and produces some very strong ammonia solution from Urea and sodium hydroxide. nothing
fancy or technical needed, the air stone wasn't even necessary and I removed it in future runs
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
Schweizer's reagent specifically refer to the chemical complex with ammonia tetraamminediaquacopper dihydroxide, [Cu(NH3)4(H2O)2](OH)2. However, in
the presence of copper metal, dilute household ammonia further diluted by 3% H2O2 and a touch of sea salt (for the galvanic electrochemical reaction
afoot with your standard chemical reactions and ligand formations, and dare I speculate, even more per this path), with a burst of energy from a
microwave to jump start the reaction (reducing the inception period), one may be able to rapidly form Schweizer's reagent. I would also expect other
routes based on standard chemical reaction of Cu(OH)2 and very strong ammonia, for example, would more likely be a mix of the lower ammines
corresponding to [Cu(NH3)k(H2O)(6-k)](OH)2 where k is at most 4 (with the blue color intensifying as k increases).
Now, in the galvanic cell path I recommend, one can still separately heated dilute aqueous ammonia and drive the NH3/air fumes into the cell with the
copper metal present, hoping to increase the presence of [Cu(NH3)4(H2O)2](OH)2. However, I am usually satisfied with the intensity of the developing
blue coloration in time without this step, per my cited path based on the direct action of NH3 on Cu in the presence of H2O2 (or O2) and an
electrolyte. Normally, depending solely on the ammonia concentration in solution, one could have anywhere from [Cu(NH3)(H2O)5]2+ (the greenish-blue
thin layer) to [Cu(NH3)5(H2O)]2+ (a royal blue complex), with the latter occurring in very concentrated ammonia solutions (see, for example, https://docs.google.com/viewer?a=v&q=cache:IjHK0vuBZhcJ:... ).
However, in the recommended variation at hand, the major source of the OH- ion, I argue is not be from the ammonia, but the cathodic reduction of H2O2
or O2, per the cited half-reaction:
1/2 O2 + H2O + 2 e- ---> 2 OH- (cathodic reduction of O2 at surface of the Copper)
And, at the Copper anode, the formation of the complex:
Cu + 4 NH3 + 2 H2O --) [Cu(NH3)4(H2O)2]2+ + 2 e- (anodic dissolution of Cu by a complexing agent)
With an overall reaction:
Cu + 4 NH3 + 1/2 O2 + 3 H2O ---> [Cu(NH3)4(H2O)2]2+ + 2 OH-
Interestingly, having performed this reaction many times, I recall once actually having seem layers of color in my galvanic cell with dilute ammonia,
including a royal blue around the copper at the base of the flask.
To be clear, what I am saying in this thread is that in a zone of high OH- ions, dilute ammonia may be able to fulfill its sole remaining role of a
complexing agent leading to the formation of [Cu(NH3)4(H2O)2](OH)2 and other products.
Also, some interesting cited standard chemical reactions, where copper oscillates between cuprous and cupric states as noted by the source below:
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
Reference: "Kinetics and Mechanism of Copper Dissolution In Aqueous Ammonia", at https://www.google.com/url?sa=t&source=web&rct=j&...
-------------------------------------------------
Now, I would like to comment on a relevant side reaction, which I suspect is an attempt on part of the authors to explain the net the formation of a
nitrite ion:
2 NH3 (aq) + 3 O2 + 2 OH- --Cu,Cu(I),Cu(Il)--> 2 NO2- + 4 H2O
Upon drawing on more recent knowledge of transition metal Fenton-like reactions, I will postulate a possible reaction pathways. In particular, per the
source cited below, the following reactions, numbered (30) and (31), occurring between cuprous and cupric ions and hydrogen peroxide as Fenton-type
reactions:
Cu(II) + H2O2 → Cu(I) + HO2· + OH− (30)
Cu(I) + H2O2 → Cu(II) + OH· + OH− (31)
Reference: "Review of iron-free Fenton-like systems for activating H2O2 in advanced oxidation processes", by Alok D. Bokare and Wonyong Choi, fully
available at http://www.researchgate.net/publication/262451840_Review_of_... .
Also:
HO2· + HO2· → H2O2 + O2
So, we should expect a continuing generation of the OH- ion and the powerful hydroxyl radical, OH·, which can attack ammonia, which is otherwise
generally resistance to aqueous decomposition. As a reference, I came across an interesting article, where hydroxyl radicals were generated not via a
Fenton process, but by photolysis of H2O2 at pH 9.3, resulting in the eventual formation of some nitrite and nitrate. Per the source, "Removal of
Ammonia by OH Radical in Aqueous Phase" by Lihuang, Liangle,.., Department of Environmental Science and Engineering, Fudan University, Shanghai,
China, 2008, to quote a possible reaction chain:
"H2O2 --hv->2•OH (2)
NH3 + •OH --> •NH2 + H2O (3)
•NH2 + H2O2 --> •NHOH + H2O (4)
•NH2 + •OH --> NH2OH (5)"
where, in the present case, the source of the hydroxyl radicals, as mentioned previously, is a Copper based Fenton-like reaction discussed above, and
not photolysis. Further comments include:
"When attacked by •OH, ammonia would be oxidized to •NH2. Then, •NH2 would be rapidly oxidized to •NHOH and further to NH2O2-. Afterward, the
unstable NH2O2- splits to NO2-, which could be oxidized to NO3-. These reaction processes were fairly comprehensible since they were in good
accordance with the results in Section 3.1. The concentration of NO2- first increased and then decayed with irradiation time while the concentration
of NO3- ascended monotonously with irradiation time."
Full paper available at http://www.google.com/url?sa=t&rct=j&q=nh3%20%2B%20h...
Unfortunately, the formation of unstable HNO2/NH4NO2 usually leads to a significant and relatively rapid release of N2. This can cause part of the
reaction vessel to overflow, and I would recommend placing the reaction vessel in a sink or tray to reduce the spillage issue. Close vessels will leak
or worst, burst.
[Edited on 1-12-2015 by AJKOER]
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
Here is the results from a recent run where I used 100 cc of 3% H2O2 added to several pieces of copper tubing (99.9% pure Cu, see http://www.copper.org/applications/plumbing/techref/tpf_stds... ) in a solution of 131 cc of household ammonia plus a teaspoon of Morton's Kosher
Salt (got it on sale, currently low on sea salt).
No jump starting in a microwave was needed. In the first 10 minutes, a significant column of foam developed and by the use of a large vessel with a
wide mouth which permitted stirring of the foam, I narrowly avoided an overflow. Within the first hour, two layers visible consisting of an intense
royal blue surrounding the Cu extending in height just over 1 cm, and a less intense coloration above. However, after 12 hours, the entire solution
was royal blue with a small amount of an insoluble blue solid on some of the copper (see picture) and at the bottom of the vessel. Likely, a quantity
of diammine tetra-aqua copper chloride (over did the salt electrolyte).
My goal was to obtain a more visible clear zone of [Cu(NH3)2]OH by employing a large excess of copper. I just, apparently, produced mostly
[Cu(NH3)4(H2O)2]2+ .
[Edited on 6-12-2015 by AJKOER]
|
|
|
Texium
|
Thread Moved
22-12-2015 at 08:06 |
|
![[*]](images/xpblue/default_icon.gif){kind=icon}
