Sciencemadness Discussion Board - Electrolytic tetraammine copper compounds (WITH PICTURES)

Electrolytic tetraammine copper compounds (WITH PICTURES)

AWLB - 26-7-2015 at 10:27

Tetraammine copper compounds are dark blue ammonia And copper (II) salt complexes. I have discovered a way to make this complex (chloride) and subsequently copper (I) chloride using only copper metal and electrolysis of ammonium chloride. the complex formed in this reaction is unstable and after a few days decomposes to ammonia and insoluble copper (I) chloride. At the anode Cl- ions are oxidized to form chlorine gas which instantly reacts with the copper electrode to form copper (II) chloride which subsequently reacts with ammonia in solution (which is produced at the cathode) to form the tetraammine copper complex in solution. You can replace the ammonium chloride with other ammonium salts which would result in the subsequent complex (eg. (NH4)2SO4--->tetraammine copper (II) sulphate). This experiment is not a particularly high yielding way to produce tetraammine complexes but is nevertheless very interesting.

Equipment

Copper wire (2 pieces 10cm long each).
53g Ammonium chloride
100ml Distilled water
DC power supply or 9v battery
Beaker

Method

1. Dissolve the NH4Cl in the water in the beaker (this is a 10 molar solution)
2. Set up power supply and attach the electrodes.
3. Make sure the cathode is beneath the anode in the liquid (also make sure they do not touch)
4. Turn on the power supply.
5. You will see the solution turn bright blue in around 3 minutes, indicating the presence of the tetraammine complex.
6. Stop electrolysis after around 5 minutes
7. The blue complex will decompose over a few days to copper (I) chloride which is a light green/white insoluble substance and ammonia gas.

Safety

Ammonia gas which is formed is toxic and irritant.
Dispose of liquid after experiment by soaking in paper and putting in a bin bag because ammonia and copper solutions are harmful to the environment.
This experiment can produce extremely poisonous and potentially explosive substances such as ammonium nitrite and chloramines(toxic).

Thank you for reading!

[Edited on 27-7-2015 by AWLB]

papaya - 26-7-2015 at 13:29

electrolyzing of ammonium chloride? do you understand the dangers ???

AWLB - 26-7-2015 at 14:28

Thank you for commenting. By "dangers" are you referring to ammonium perchlorate, because in this electrochemical reaction none is produced. As I said in the introduction to the post the copper anode reacts immediately with the chlorine to form copper chloride.

aga - 26-7-2015 at 14:34

Another 'i did this a while ago and maybe forgot' moments ?

At times the copper anode might not be looking, maybe having a beer or something, and forgets to react with the Cl<sup>-</sup> ion.

Of course some Cl₂ will escape.

Arguing otherwise would be sheer idiocy.

On that point, i shall retire bedwise, lest i forget to react correctly.

[Edited on 26-7-2015 by aga]

AJKOER - 26-7-2015 at 14:42

I agree with Papaya comment. One should not perform electrolysis of an increasing acidic ammonium solution. The summary explanation is the potential formation of aqueous NH4NO2 which is prone to detonation in lower pH solutions. Here is a relevant extract from one of my prior threads (link: http://www.sciencemadness.org/talk/viewthread.php?tid=18912 ):

Quote: Originally posted by AJKOER

OK back on topic, here is an interesting account on the electrolysis of aqueous ammonia in the presence of NaOH and, separately, Cu(OH)2 from an old (1905) report (see page 242 at http://books.google.com/books?pg=PA242&lpg=PA242&dq=... from Journal Chemical Society, London, Volume 88, Part 2), to quote:

"Electrolytic Oxidation of Ammonia to Nitrites. Erich Muller and Fritz Spitzer (Ber., 1905, 38, 778—782. Compare Traube and Biltz, Abstr., 1904, ii, 727).—In the presence of a small amount of sodium hydroxide, ammonia may be oxidised electrolytically to nitrite even in the absence of copper compounds.

In the presence of copper hydroxide and sufficient alkali, the oxidation of ammonia to nitrite does not cease suddenly when the nitrite concentration has reached a certain value, but appears to proceed quite independently of the nitrite concentration. In these experiments, the oxidation was allowed to proceed for a comparatively short time only, so that the amount of alkali present was not greatly reduced. The formation of nitrite is intimately connected with the amount of alkali present, and when no sodium 'hydroxide is present, but only ammonia, nitrite, and copper hydroxide, it is found that the nitrite is transformed into nitrate more rapidly than the ammonia into nitrite, and thus the concentration of the nitrite tends to decrease.

Nitrogen is also formed during the oxidation. J. J. S."

The source also notes, to quote:

"In continuation of the previous experiments, the influence of changing the concentration of the free alkali or ammonia on the rate of the electrolytic oxidation of ammonia has been investigated. In presence of much ammonia, the amount of nitrite can be increased to about 11 per cent, before oxidation to nitrate begins, whilst from an 11 per cent, nitrite solution to which ammonia, sodium hydroxide, and copper hydroxide had been added a solution containing as much as 17 per cent, nitrite was obtained on hydrolysis."

Apparently replacing NaOH with Cu(OH)2 favors the formation of nitrate over nitrites, and increasing the ammonia concentration raises the yield. Caution: product could include some copper ammonium nitrate, see discussion at http://www.pyrosociety.org.uk/forum/topic/3303-electrolysis-... . This experiment may be inherently dangerous as the author states "In these experiments, the oxidation was allowed to proceed for a comparatively short time only, so that the amount of alkali present was not greatly reduced" together with the observed formation of N2. From this I suspect the presence of NH4NO2 (decomposing to form nitrogen), which is inherently unstable (explosive) as the pH is lowered, which could be particular problematic in the presence also of any copper ammonium nitrate.

On the surface IMHO, this appears to be a simple, educational and safe experiment, but upon adding NaOH and/or Cu(OH)2 to the aqueous ammonia, things apparently could go very wrong, especially if one attempts to recover the dry salts.
[Edited on 30-3-2013 by AJKOER]

Now as to how such products can be formed, I provide this link http://chemistry.stackexchange.com/questions/12315/electroly... to the underlying processes involved in electrolysis of water itself which involves the formation of several radical species including the active hydroxyl radical. The latter acts on NH3 in the presence of O2 to form NH4NO2 and NH4NO3. Here is an abstract outlining the process to quote from "Removal of ammonia by OH radical in aqueous phase" by Huang L1, Li L, Dong W, Liu Y, Hou H.:

"Many advanced oxidation technologies have been developed to remove ammonia in wastewater. All these technologies have one common characteristic, that is, the removal processes involve OH radical (*OH). In this research work, H2O2 was selected as *OH precursor. The removal of ammonia under 253.7 nm irradiation from low-pressure mercury lamp in the presence of H2O2 was studied to investigate the ammonia removal efficiency by *OH. Results show that the *OH, generated by H2O2 photolysis, could oxidize NH3 to NO2- and further to NO3-. Removal efficiencies of ammonia were low and were affected by initial pH value and ammonia concentration. Laser flash photolysis technique with transient absorption spectra of nanosecond was used to investigate the oxidation pathway and kinetics of ammonia oxidation by *OH. Results illustrate that *OH could oxidize NH3 to form *NH2 with a second-order rate constant of (1.0 +/- 0.1) x 10(8) M(-1) s(-1) (20 degrees C). *NH2, the main product of *OH with NH3, would further react with H2O2 to yield *NHOH. Since *NHOH could not stay stable in solution, it would rapidly convert to NH2O2- and consequently NO2- and NO3-. The rate constants for these elementary reactions were also given. The low removal efficiency of ammonia by *OH was mainly due to the slow reaction rate constant"

Link: http://www.ncbi.nlm.nih.gov/pubmed/19031904

[Edited on 26-7-2015 by AJKOER]

papaya - 26-7-2015 at 14:55

Amazingly noone mentioned possible formation of chloramines (toxic) AND NCl3 ! This last might explode violently and that's why electrolyzing NH4Cl is bad idea. Don't say copper reacts instead of that because you never know - maybe under some conditions it reacts completely, but in other cases some Cl2 will escape and dissolve in solution further reacting with ammonia part. Didn't you smell any strange pungent smell during your experiment? I bet you did!

AJKOER - 26-7-2015 at 15:19

Quote: Originally posted by papaya

Amazingly noone mentioned possible formation of chloramines (toxic) AND NCl3 ! This last might explode violently and that's why electrolyzing NH4Cl is bad idea. Don't say copper reacts instead of that because you never know - maybe under some conditions it reacts completely, but in other cases some Cl2 will escape and dissolve in solution further reacting with ammonia part. Didn't you smell any strange pungent smell during your experiment? I bet you did!

Usually, the formation of NCl3 requires an excess of chlorine relative to the ammonia and we are starting with a balance situation from the NH4Cl.

However, to the extent that ammonia is removed from the mix per reactions involving problematic nitrites formation and such, I guess NCl3 is still possible, but perhaps not overly likely (but perhaps more so if the vessel was a long tube and the formed Cl2 gas flowed up through the NH4Cl solution).

[Edited on 26-7-2015 by AJKOER]

AWLB - 26-7-2015 at 15:42

Thank you very much papaya, aga and AJKOER for highlighting the possible dangers of this experiment to both me and others reading this post, I will add extra safety notices to my method.

blogfast25 - 26-7-2015 at 15:55

Quote: Originally posted by AWLB

Thank you very much papaya, aga and AJKOER for highlighting the possible dangers of this experiment to both me and others reading this post, I will add extra safety notices to my method.

It's likely that the amounts of copper(II)tetrammine chloride you produced are really small. These complexes are very, very intensely coloured and and even low concentrations will show up as colour. It's not a practical way of preparing such compounds.

AWLB - 26-7-2015 at 21:23

Do you agree with what I observed, i.e. that this particular tetraammine chloride complex was unstable(decomposing after several days), resulting in the formation of copper (I) chloride (CuCl) and ammonia ? Or was copper oxychloride formed due to an excess of Cu2+ and alkaline conditions ? or maybe could excess Cl- anion make the complex unstable. There appears to be very little information online about this particular tetraammine copper compound. I also must agree with you on the fact that this method is not practical to produce large quantities of tetraammine copper (II) chloride, it was simply a interesting reaction I observed and wished to share.

Thanks.

[Edited on 27-7-2015 by AWLB]

papaya - 27-7-2015 at 03:53

From you pictures - Cu dissolves at anode and plates out at cathode, you will never get much amount of copper to stay in solution that way. Better you try to prepare it as follows: Copper fillings + NH4Cl + NH4OH (excess) + oxygen (bubbling air), hopefully after many hours you'll dissolve all the copper.

AWLB - 27-7-2015 at 04:03

Thank you very much for commenting papaya, from your advice I have added more safety notes to my post.

The method you have proposed looks very interesting and I will certainly attempt it. I have one question about it though; does the reaction [you proposed] require elevated temperatures or can it be performed efficiently at room temperature.
I do agree with you, my method does not produce a large quantity of Cu2+ ions in solution, which does make it very inefficient.

Thank you.

AJKOER - 27-7-2015 at 04:35

Quote: Originally posted by papaya

From you pictures - Cu dissolves at anode and plates out at cathode, you will never get much amount of copper to stay in solution that way. Better you try to prepare it as follows: Copper fillings + NH4Cl + NH4OH (excess) + oxygen (bubbling air), hopefully after many hours you'll dissolve all the copper.

Better per my experience is Copper fillings + NH4Cl + NH4OH (excess) + H2O2 (even dilute works) + a touch of sea salt. Microwave (best as it also creates some hydroxyl radicals and other associated reactive oxygen species) for a few seconds to jump start the reaction.

Note, within the first 10 minutes you may witness a significant gas release from the solution. This will cause a tightly sealed vessel to burst and even in a large mouth open vessel, a possible overflow. This is caused by the formation of an unstable HNO2/NH4NO2 that decomposes abruptly when the pH drops releasing N2. As a reference source, please see http://www.academia.edu/292096/Kinetics_and_Mechanism_of_Cop... ) which identifies this side reactions among the other primary electrochemical reactions, also presented.

[Edited on 27-7-2015 by AJKOER]

AWLB - 27-7-2015 at 04:42

Thanks AJKOER, that also looks like a good way to produce the complex, are there other ways of introducing hydroxyl radicals to the reaction mixture without using a microwave (I do not have access to one), does UV cause •HO to be formed (as with chlorine it produces the Cl· radical) ?
If no microwave is used is external heating required?

papaya - 27-7-2015 at 04:43

Depends where you live, or what is your room temperature

That method was proposed long ago here in this forum for nitrate salt(and it works), but I thing chloride will do also, if not even better. It's important for copper to have high surface area (foil,filings, not a thick wire), and ammonia to be concentrated enough (25% is good) and in excess (lot's of it is lost while blowing air through it).

AWLB - 27-7-2015 at 04:48

Thanks papaya, I will give it a try! My room will undoubtedly smell strongly of ammonia after performing this experiment (maybe I should do it outside

). As for my room temperature, I live in England, despite it being the middle of summer it is only around 15 degrees celsius

, but it should be sufficient.

Thanks again for the details on this experiment.

[Edited on 27-7-2015 by AWLB]

AJKOER - 27-7-2015 at 05:00

A microwave is not required (I believe it speeds things up) , and I do not recommend even the use of concentrated aqueous ammonia or H2O2, as concentrated side reaction products like NH4NO2 are less stable when concentrated and even more problematic.

I have added a reference source that you may find interesting.

Note, the product first created is tetraamminecopper(II) hydroxide, a base, which acts on the NH4Cl forming aqueous ammonia (which, with continuing addition/presence of H2O2, is consumed) and then the tetraamminecopper(II) chloride.

[Edited on 27-7-2015 by AJKOER]

AWLB - 27-7-2015 at 05:10

Thanks for the extra information, I appreciate the help you have given me, I will attempt the experiment under your recommended instructions and use all reagents in a dilute form at first(I will try 1g copper powder,100 ml 10% ammonia, 3g NH4Cl, 20 ml 6% H2O2 and use an air pump). In your method you included using NaCl, what is the purpose of this, is it to ensure there is an excess of Cl-? I also did not know that tetraamminecopper(II) hydroxide is initially formed and then reacts with more NH4Cl, thank you for informing me. I believe this sort of electrochemistry is fascinating, but quite complex

Thanks.

[Edited on 27-7-2015 by AWLB]

blogfast25 - 27-7-2015 at 05:45

Going by the second photo, the concentration of copper(II) tetrammine complex obtained was probably less than 0.001 M.

papaya - 27-7-2015 at 05:50

Who said anything about H2O2 ? Ajoker, honestly, did YOU ever try anything close to what is the discussion here? Speculating ?
Here are facts
1. Copper ammino complexes are STRONG catalysts for H2O2 decomposition, I'll not google for "references" instead of you. this means that all peroxide WILL decompose before it does anything to copper metal (homogeneously) sputtering solution all around and form god knows what product

2. from practical experience: all solutions must be as STRON as possible - excess 25% ammonia, dissolve NH4Cl directly into it (some excess also here) + oxygen. Weak solutions work VERY slow, weeks...

Don't advice things you've never seen!

AJKOER - 27-7-2015 at 06:28

Papaya:

Believe me when I say I have experienced many of the points I have stated including gas pressure bursting and overflows, performance enhancement via the microwave (see my research on a recent thread at http://www.sciencemadness.org/talk/viewthread.php?tid=62978#... ) and even the use of sea salt over NaCl (which I personally suspected until I saw it explicitly mentioned in a journal article a while back and recently, explicitly incorporated into Patent US 6740220 available at https://www.google.com/patents/US6740220 to quote:
"The aluminum anode was an ALUPOWER™ alloy designated EB50V. The electrolyte contained 3.0M NaOH, 0.5M H2O2 and 40 gram/Liter (g/L) of sea salt. Its temperature was maintained at 55° C. and the flow rate was 100 cm3/min."

On radical formation from H2O2 via microwave and uv please see http://www.hindawi.com/journals/ijp/2013/854857/

Here is my reference again at: https://www.academia.edu/292096/Kinetics_and_Mechanism_of_Co...

As to my actually preparing the salt in question and its nitrate cousin with dilute ingredients, warning: link to the Energetic material section of SM, see http://www.sciencemadness.org/talk/viewthread.php?tid=26053#... .

[Edited on 27-7-2015 by AJKOER]

papaya - 27-7-2015 at 07:03

I don't get the point - use peroxide(most of which will decompose catalytically) instead of cheap air? What? Nitrite formation side reactions? OK, have you measured nitrite levels that can form (and that has nothing to do with the salt we talk about here) ? No? Then how can you state that nitrite is a problem here, etc, etc. Though, I don't understand anything from what you're writing, every sentence leads to some distant "reference", that demands days to follow, I don't care. What I proposed WORKS (at least for nitarate salt) and that's it.

AJKOER - 27-7-2015 at 07:18

Papaya:

Yes, I agree that what you say works, but my route leads also to a side reaction involving HNO2/NH4NO2 formation and a significance rapid nitrogen gas release. Oddly, but an apparent producer of lab explosions, is the sensitivity of aqueous NH4NO2.
-----------------------

I once took days to make prepare Copper acetate from copper metal, vinegar and dilute H2O2.

I now can prepare the same amount in way less than an hour by adding sea salt and heating in a microwave. Here is an example of reaction speed (picture near bottom of page per cited link):

Quote: Originally posted by AJKOER

What surprised me recently was when I placed an old nail in salt sea and 3% H2O2. No reaction, zip. Then, in a microwave for 30 seconds, and wow! Bubbles up like super strong acid from sea salt and dilute H2O2? Actually, constructed an Iron-Oxygen galvanic cell with an NaCl electrolyte.

Cool way to Fe2O3.xH2O (see pictures at http://www.sciencemadness.org/talk/viewthread.php?tid=153&am... ).

[Edited on 15-8-2014 by AJKOER]

What is even more unexpected, repeat the above with sterling silver (Ag/Cu), critic acid and dilule H2O2 (instead of iron and acetic acid) and the reaction rate is similar!

[Edited on 27-7-2015 by AJKOER]

papaya - 27-7-2015 at 07:31

So what? Copper + acid + H2O2 is used in large scale for copper etching (circuit boards ), and that's not a surprise - under acidic conditions H2O2 is stable. HOWEVER, try to mix some CuSO4 + NH4OH until everything is dissolved and add a drop of H2O2. A very intense "explosive" decomposition of H2O2 is the result, just in few seconds! Using H2O2 is a waste in this situation, it WILL totally decompose.
Anyway, I'll not distract you guys from your business anymore, do it, at least this time you didn't mention bleach, but I suspect you'll soon arrive there

AJKOER - 27-7-2015 at 07:49

My last words, the paper I gave you may mention the preferred order in which the H2O2 is added is last. In other words, the copper oxide coating on the copper metal is first attacked by ammonia and then the dilute H2O2 is added.

In the microwave, O2, reactive oxygen species and hydroxyl radicals are then formed attacking the ammonia along with the copper.

[Edited on 27-7-2015 by AJKOER]

MeshPL - 20-8-2015 at 07:27

A few days ago I did something simmilar but with 25% aqueous ammonia. I was able to obtain quite intensly colored solution, too dark to see anything through it unless put directly against a source of light even after diluting it with 2-3 times more water. I used a jar with diameter around 8cm as a container.

Turns out you can synthesise tetraaminacopper hydroxide electrolytically. With some help of HCl it can be turned into chloride. Obviously due to strong coloration of tetraamina copper (II) I didn't obtain very concentrated solution.