| Pages:
1
2 |
unionised
International Hazard
Posts: 5102
Registered: 1-11-2003
Location: UK
Member Is Offline
Mood: No Mood
|
|
Practically speaking, ammonium hydroxide doesn't exist.
If I were going to try making this stuff I would try the same method that is generally used for tetramine copper sulphate ie use a concentrated
solution of copper sulphate, add conc ammonia, then add alcohol to ppt the product.
|
|
|
The_Davster
A pnictogen
Posts: 2861
Registered: 18-11-2003
Member Is Offline
Mood: .
|
|
Unionised, the problem for me with your method is that the only ammonia that I have acess to(other than generating the gas) is the dilute solution
avaliable at the grocery store. Is there a way to concentrate it?
|
|
|
hodges
National Hazard
Posts: 525
Registered: 17-12-2003
Location: Midwest
Member Is Offline
|
|
I tried this again, and had an even less favorable result. Because of the humid weather I decided to dry it in an oven at around 80C. The resulting
product looked darker blue than copper nitrate, but it would not do a darn thing. I even heated a small amount of it with a blowtorch and all it did
was sizzle a bit (may have been just water escaping). Looks like it had decomposed in the 80C heat of the oven.
|
|
|
The_Davster
A pnictogen
Posts: 2861
Registered: 18-11-2003
Member Is Offline
Mood: .
|
|
I'm letting a sample of copper nitrate air dry. I have found that it seems to be slightly soluble in xylene, however the copper nitrate was not
dry at the time of adding to the xylene so iy could have just been the water in the copper nitrate allowing it to dissolve. Will bubble the ammonia
through tomorow.
|
|
|
hodges
National Hazard
Posts: 525
Registered: 17-12-2003
Location: Midwest
Member Is Offline
|
|
[Cu(NH3)4]SO4 Results
I made some tetraamminecopper sulfate. I started with 12.5 grams (0.05 moles) of copper sulfate pentahydrate. I dissolved this in water. I then
added an excess of ammonium hydroxide. This produced the characteristic deep blue color.
It took several weeks for all the water to evaporate from the solution at room temperature. I determined when it was dry by monitoring the weight
until it was no longer falling. The result was a blue salt. It is not as dark as the solution, but considerably darker than copper sulfate. It is
much more easily powdered than hydrated copper sulfate, but even when finely powdered it is hard to re-dissolve it in water. Heating caused a
definite evolution of NH3 gas.
The resulting product weighted 10.9 grams. Doing the math, there are 3.22 molecules of NH3 for every Cu molecule. Theoretical is 4. Of course, it
is also possible that there is less NH3 and more H2O, since the two have very similar masses.
|
|
|
t_Pyro
Hazard to Others
Posts: 120
Registered: 7-2-2004
Location: India
Member Is Offline
Mood: Volatile
|
|
Tetrammine copper compounds have got me confused now. Till now, I always used to regard it as [Cu(NH<sub>3</sub> <sub>4</sub>]<sup>++</sup>. Then, I came across a paragraph in
my inorganic chem book which referred to it as [Cu(NH<sub>3</sub><sub>4</sub>(H<sub>2</sub>O)<sub>2</sub>]<sup>++</sup>, which gives it a coordination number of 6.
So what is the coordination number of Cu? 4 or 6?
[Edited on 3-4-2004 by t_Pyro]
|
|
|
hodges
National Hazard
Posts: 525
Registered: 17-12-2003
Location: Midwest
Member Is Offline
|
|
I'm not having too good of luck drying my tetraamminecopper nitrate. The dark blue color fades from the solution after a couple days, leaving a
lighter blue color. I tried adding NH4OH again and the same thing happened after a couple of days. This problem did not occur when I was making
tetraammine copper sulfate. The crystals were lighter than the original solution, but the solution itself remained dark until it all evaporated. Why
should it matter whether the other ion is nitrate or sulfate as far as the rate of decomponsition?
I'm coming close to concluding that making tetraammine copper nitrate is impossible, at least in aqueus solution. I thought more about the
previous observation that tissue paper soaked in this solution burns leaving behind copper. Probably what is happening is that copper nitrate is
forming copper oxide as the paper burns. Then the remaining carbon from the paper reduces the copper oxide to copper. There probably isn't even
any TACN at all (I assumed there was due to Cu being formed when it burned).
|
|
|
The_Davster
A pnictogen
Posts: 2861
Registered: 18-11-2003
Member Is Offline
Mood: .
|
|
TACN success
Well as I did before, I made a concentrated solution of copper nitrate in a testube(15ml). Ammonia gas was bubbled through, first I got a tiny bit of
dark blue precipitate, then it all slowly turned light blue due to a precipitate. I continued bubling the ammonia through for about 10min. At this
time there was a mix of about 50/50 light blue and dark blue precipitates. some of this was filtered and the dark blue precipitate dissapeared for
some reason. the rest of the solution I dident want to bother with filtering so I dumped it onto an icecrean bucket lid, placed a filter paper on top
and absorbed all the water with paper towels. there was a mix of light and dark blue precipitate on the filter paper. this was allowed to dry and
today I lit the filter paper....success  I got little green puffs of fire and
some crackling as the crystalls burned.
|
|
|
chemoleo
Biochemicus Energeticus
Posts: 3005
Registered: 23-7-2003
Location: England Germany
Member Is Offline
Mood: crystalline
|
|
Hmm, I don't understand the problem of it so much.
The light blue precipitate is Cu(OH)2 or copper hydroxide, which forms by adding the base ammonia. Continuous addition of NH3 eventually dissolves
this again, to form the complex. Ammonia solution is better in this case as it has all the ammonia dissolved already, unlike the NH3 gas where it has
to be absorbed first.
Why dont you try the following (similar to http://www.sciencemadness.org/talk/viewthread.php?tid=1778 ):
Make a saturated solution of Cu(NO3)2, and then add ethanol until copper nitrate starts to precipitate. I did this recently (by accident, with CuSO4),
a 50% addition of EtOH or so should be fine. Just check it how much EtOH this solution can hold.
Then take the alcoholic solution of copper nitrate, and add a mixture of ammonia in water/ethanol (the proportions of the mixture determined by the
result of the EtOH/Cu(NO3)2 mix, i.e. 50%NH3 solution/50% EtOH), and add this until you reach an excess of full saturation (calculate it).
At this point, you either get the deep blue precipitate of the nitrate complex, or everything remains in solution.
In case of precipitation, everything is good, you just purified the complex. In case of no precipitation, just add MORE ethanol until the complex
finally precipitates.
Filter it, dry it, and test!!
Honestly, I don't see at all why this shouldnt work! Good luck!
[Edited on 7-4-2004 by chemoleo]
Never Stop to Begin, and Never Begin to Stop...
Tolerance is good. But not with the intolerant! (Wilhelm Busch)
|
|
|
Theoretic
National Hazard
Posts: 776
Registered: 17-6-2003
Location: London, the Land of Sun, Summer and Snow
Member Is Offline
Mood: eating the souls of dust mites
|
|
I have myself prepared Cu(NH3)4SO4.
I have added copper sulfate to a 25%solution of ammonia, a deep blue colour appeared above the crystals, and the reaction proceeded quite quickly. For
some reason, it is now precipitating Cu(OH)2.
Hold on... shouldn't free copper ions destroy complexed copper ions:
Cu++ + Cu(NH3)4++ + 6H2O => 2Cu(OH)2 + 4NH4+
Otherwise there wouldn't be any precipitate when ammonia is added to copper salts. In my case there wasn't any, why are the textbooks
stating it? Or does the precipitate redissolve if it forms?
I have dissolved copper in a (NH4)2SO4 solution, with aid of atmospheric oxygen. Right now the copper is still dissoving, and the solution is
distinctly blue.
Dissolving copper in a NH4Cl solution should be fun, also an easy way to CuCl2 without HCl.., like so:
2Cu + O2 + 4NH4Cl => Cu(NH3)4CuCl4 + 2H2O.
Then heat to get rid of ammonia and you're done!
Edit: spelling, what else... 
[Edited on 14-6-2004 by Theoretic]
|
|
|
The_Davster
A pnictogen
Posts: 2861
Registered: 18-11-2003
Member Is Offline
Mood: .
|
|
I have been allowing some copper wire/strip/filings in an ammonium nitrate solution in the hopes of obtaining TACN. There were reports of success
with this method on the E&W I think. I am in the process of testing this out and so far the solution has turned a deep blue. I think the
reaction hapenning is 2Cu+4NH4NO3-->Cu(NH4)4(NO3)2 + Cu(NO3)2. Obviously this reaction produces two water soluble which would make extraction of
pure TACN quite difficult. So I got to thinking on how to obtain a purer product of TACN and came up with this
2Cu+2NH4NO3+2NH4OH-->Cu(NH4)4(NO3)2+Cu(OH)2. Here there is the advantage of an insoluble product obtained which leaves just TACN in solution after
filtering. However I am unsure if the second reaction will work. Also it would be advantageous I think to use a slight excess of copper to make sure
that all the ammonium nitrate would be reacted to keep it from contaminating the final product.
Thoughts?
|
|
|
hodges
National Hazard
Posts: 525
Registered: 17-12-2003
Location: Midwest
Member Is Offline
|
|
If you add excess NH4OH it will convert any Cu(NO3)2 to [Cu(NH4)3](NO3)2. So yes that should work. I never had much luck with TACN though, even
starting with pure Cu(NO3)2. At best it would deflagrate, did have a pretty green flame though. Unfortunately it is also a powerful explosive if
detonated, which to me makes it the worst of both worlds. It is hard to get to work and gives you a false sense of security becuase usually it does
not do much. This might make it tempting to use larger amounts, which might actually detonate, and remember it is very powerful if detonated.
|
|
|
The_Davster
A pnictogen
Posts: 2861
Registered: 18-11-2003
Member Is Offline
Mood: .
|
|
This summer I plan do some in depth work with the various methods of production of TACN and test the products relative to each other. Does anyone
have the VoD of TACN?
|
|
|
Magius
Harmless
Posts: 20
Registered: 9-6-2004
Location: Green Bay
Member Is Offline
Mood: Constrianed
|
|
Hey guys, I to have prepared some tetraaminocopper(II), I think. I set up an electrolysis experinemt with ammonia as the electrolyte, and copper
annodes and cathodes. After running it for a few hours, I had a dark blue solution that appears to be tetraaminocopper(II). But I'm not sure what
the anion of the solution is, anyone have any suggestions? No percipitate has formed, so OH is out of the question.
Wait for it...
|
|
|
Geomancer
Hazard to Others
Posts: 228
Registered: 21-12-2003
Member Is Offline
Mood: No Mood
|
|
Isn't tetraamine copper (II) strong enough to be soluble even with hydroxide counter-ion?
|
|
|
Nitro-esteban
Harmless
Posts: 39
Registered: 10-4-2013
Location: Fifth dimension
Member Is Offline
Mood: inert
|
|
I made tetraamminecopper (II) perchlorate by adding a saturated solution of copper perchlorate to an excess of ammonia (30% concentrated). Adding the
ammonia to the copper perchlorate only presipitates copper hydroxide.
|
|
|
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 Nitro-esteban  | | I made tetraamminecopper (II) perchlorate by adding a saturated solution of copper perchlorate to an excess of ammonia (30% concentrated). Adding the
ammonia to the copper perchlorate only presipitates copper hydroxide. |
You needed to add more ammonia. Copper(II) hydroxide should be plenty soluble in enough xs ammonia. Either that, or add some ammonium perchlorate as
a buffer.
[Edited on 4-5-2014 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.
|
|
|
nezza
Hazard to Others
Posts: 324
Registered: 17-4-2011
Location: UK
Member Is Offline
Mood: phosphorescent
|
|
The ethylenediamine perchlorate complex is an interesting dark purple compound which explodes erratically when heated giving a beautiful blue flame.
|
|
|
Nitro-esteban
Harmless
Posts: 39
Registered: 10-4-2013
Location: Fifth dimension
Member Is Offline
Mood: inert
|
|
| Quote: Originally posted by DraconicAcid | | Quote: Originally posted by Nitro-esteban | | I made tetraamminecopper (II) perchlorate by adding a saturated solution of copper perchlorate to an excess of ammonia (30% concentrated). Adding the
ammonia to the copper perchlorate only presipitates copper hydroxide. |
You needed to add more ammonia. Copper(II) hydroxide should be plenty soluble in enough xs ammonia. Either that, or add some ammonium perchlorate as
a buffer.
[Edited on 4-5-2014 by DraconicAcid] |
Copper (II) hydroxide forms when the ammonia is added to the copper perchlorate but if the reactants are mixed the other way around (copper
perchlorate is added to the ammonia) an exothermic reaction occurs and tetraamminecopper (II) perchlorate precipitates as the temperature drops. In
both attempts I used the same amount of ammonia and copper perchlorate, I just mixed them in different ways.
Here is a link to a thread about TACN: http://www.sciencemadness.org/talk/viewthread.php?tid=16220#...
[Edited on 4-5-2014 by Nitro-esteban]
|
|
|
The Volatile Chemist
International Hazard
Posts: 1981
Registered: 22-3-2014
Location: 'Stil' in the lab...
Member Is Offline
Mood: Copious
|
|
| Quote: Originally posted by chloric1 | | Quote: |
2) Is there a way to convert the tetrachlorocopper ions back to copper(II) ions?
Thanks
[Edited on 23-3-2004 by rogue chemist] |
Yes there is, add to excess H2O! I have played around alot with CuSO4 and other common copper salts. When in excess chloride, especially acidic, you
get bright green tetrachlorocuprate. Now you add water until you get a aquamarine color. If you want to really play, you should take acidic CuSO4
and add large portions of concentrated NaBr solution. Bromo complexes are Purple! My favorite color. If you boil this it is unstable and it will
deposit cuprous bromide on dilution!! FUN FUN!!
OMG! Post 100!! I am hazardous to others now!! MAybe I have always been! LOL!
[Edited on 3/23/2004 by chloric1] |
When you make the bromide complex, this person said "acidic CuSO4. Did they mean in-acid, or were they considering the copper sulfate
'acidic'?
Also, someone mentioned an acetate complex. What? I hadn't heard of this. How would one go about making it (w/o just making CuAc?
Thanks,
Nathan
|
|
|
| Pages:
1
2 |
![[*]](images/xpblue/default_icon.gif){kind=icon}
