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woelen
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Interesting and colorful coordination chemistry with ethylene diamine
I did some experiments with nickel(II) ions and ethylene diamine. This combination affords remarkable colorful chemistry and as a bonus one can make a
nice energetic compound as well.
http://woelen.homescience.net/science/chem/exps/Ni_en_comple...
I hope to do more experiments, one of my goals is to isolate the perchlorate salt of the deep blue bis(en) complex. I can imagine that this is even
more energetic than the perchlorate salt of the tris(en) complex, because of better oxygen balance. The effect may be neutered, however, by the
presence of two water ligands in the bis(en) complex, but if I can isolate this complex, then I certainly will try this and compare it with the
tris(en) complex of which I still have almost 1 gram around and which can be stored safely.
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AndersHoveland
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I suspect the range of colors would be even more spectacular with 1,2-diaminobenzene, or anything with conjugated bonding between the vicinal carbon
atoms. I am thinking 1-amino-propyl-2-imine, but it might be rather difficult to synthesize, because from what I read, aminoacetone is not a stable
molecule (sorry do not mean to derail the topic here).
Woelen, you are aware that acetylacetone can also form a complex with copper?
Quote: Originally posted by woelen  | | one of my goals is to isolate the perchlorate salt of the deep blue bis(en) complex. I can imagine that this is even more energetic...
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It would be interesting to see what combined color the nitroformate salt would be, since the nitroformate ion has a bright yellow color. Might not be
a simple matter of mixing the colors blue and yellow, it depends exactly which wavelengths are absorbed, and how much overlap there is the visible
spectrum. It could be green, or purple, or reddish...
[Edited on 9-9-2013 by AndersHoveland]
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Bezaleel
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Woelen, you're just some time ahead of me. Great webpage you made, I can't say anything else. I intended to make this compound somewhere this winter.
In the energetic materials section I read this complex is reasonably stable, and can be well kept. Is the salt with sulphate also somewhat energetic
or not at all?
I was still in doubt whether I should synthesise my ethylene diamine or just purchase it. (Haven't really searched for how feasible making it is.)
First I intend to do two other syntheses I have slowly gathered the information for, so don't expect any work in the field of en-complexes in near
future from my side...
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AndersHoveland
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No, it would not be the slightest bit energetic at all. For most practical purposes in pyrotechnics, sulfate is not even considered an oxidizer.
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DraconicAcid
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I've been meaning to play with such complexes for quite some time. I'd like to try making the bis(ethylenediamine)nickel complexes with chloride,
bromide or iodide as the counterions; if made from anhydrous alcohol, they should coordinate, and give different colours.
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.
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woelen
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Yes, the bis(en) dichloro nickel(II) complex also is one of the things on my list. It is an example of a neutral complex. The chlorides are not free
ions, they are firmly coordinated to the nickel atom. Because it is a neutral (covalent) complex it also is soluble in many organic solvents. But
first I try the perchlorate salt of the aqueous bis(en) complex. If I have results, I'll extend the web page I have now and post the results here in
this thread.
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DraconicAcid
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| Quote: Originally posted by woelen | | Yes, the bis(en) dichloro nickel(II) complex also is one of the things on my list. It is an example of a neutral complex. The chlorides are not free
ions, they are firmly coordinated to the nickel atom. Because it is a neutral (covalent) complex it also is soluble in many organic solvents. But
first I try the perchlorate salt of the aqueous bis(en) complex. If I have results, I'll extend the web page I have now and post the results here in
this thread. |
And because they are coordinated, they should give different colours from the tris(en) complex. They should be labile, though, so I wouldn't expect
cis and trans isomers (not isolatable ones, at least)- you might get a grimy colour due to a mixture of the cis and trans. The analogous Co(III)
complexes are green and dark purple.
You could also try a bis(ethylenediammine)oxalato complex, by precipitating nickel oxalate, then dissolving it in ethanolic ethylenediamine. This
would have to be cis, so you wouldn't have to worry about two different, labile isomers.
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.
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atomicfire
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How might one obtain some ethylene diamine?
ban DHMO
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woelen
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Failed attempt to isolate the blue bis(ethylene diamine) bi aqua nickel(II) complex
In this experiment, an excess amount of nickel(II) carbonate was added to an approximately 15% by weight solution of perchloric acid. After the
initial fizzling, the solution was heated for quite some time to assure that all perchloric acid is consumed. The result is a turbid liquid, which has
a high concentration of nickel perchlorate, with finely suspended solid nickel carbonate in it.
The solution was poured in a thin test tube which was put aside for two days, with a rubber stopper on it to assure no dust could enter the test tube.
After two days there was a perfectly clear bright green solution on top of a thin layer of pale green precipitate (unreacted remains of the nickel
carbonate). Using a pasteur pipette with a long needle-like tip, the clear liquid is sucked away from the precipitate and transferred to a separate
test tube.
Drops of a 25% solution of ethylene diamine are added to the bright green liquid, until the liquid has a deep blue color. After each drop the liquid
is shaked to assure good mixing. Addition of drops does not lead to formation of precipitates. First the liquid becomes cyan, then light blue like
dilute copper sulfate and then the liquid darkens after addition of each drop. As soon as a single drop did not cause any visible deepening of the
blue color, addition of ethylene diamine was stopped. The end result was a beautiful deep blue liquid, much like the liquid, shown in one of the
pictures above.
The deep blue solution was transferred to a petri dish and put in a warm dry place, free of dust. After one day, the volume of liquid was less than
when it was made and the color of the liquid still was deep blue. After two days, all liquid had evaporated, but no nice dry crystalline mass was
produced. A dirty-looking, very sticky, green/brown paste was sticking to the glass of the petri dish. After yet another half a day, the sticky
tar-like material still was tar-like (like a nearly dry syrup, which does not dry further). It could be scraped off, but it remained sticking to the
spatula.
The dark green/brown paste was not kept around, some water was added to it to see whether it can be reverted to the nice deep blue solution. This is
not the case. The paste does dissolve fairly easily, but it gives a turbid pale blue liquid. Probably the pale liquid contains some of the
bis-ethylene diamine complex, but also a lot of nickel hydroxide.
The isolation of the complex did not succeed. A small excess amount of ethylene diamine was added to the turbid pale blue liquid and this immediately
caused formation of a very fine purple precipitate and formation of the violet liquid. This liquid was heated and even gentle heating was sufficient
to make it completely clear and deep violet. On cooling down, many nice glittering violet crystals were produced of tris(ethylene diamine) nickel(II)
perchlorate. The liquid above these crystals was decanted and these crystals were dried in the same way as described above.
So, it can be concluded that the tris-complex can be isolated perfectly well, but the bis-complex cannot easily be isolated. Most likely this is due
to the more labile properties of the complex and partial loss of ethylene diamine from the liquid when it is allowed to evaporate to dryness.
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Bezaleel
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This is a funny result, woelen. In the liquid the complex seems to be stable, and I don't think it would change if you would let the solution with the
dark blue complex stand for a long time.
What I don't get is why the solution of the brownish paste does not dissolve clearly. Since en is a neutral ligand, the brownish paste should be
balanced in Ni2+ and ClO4- content. If en evaporates I can't see why that would change, unless a compound between en and perchlorate would form.
I am tempted to conclude that what you got was somewhere near an oxidised mono-en complex of nickel:
[Ni(en)2(H2O)2](ClO4)2 --> dehydration + oxidation --> [Ni(en)O(H2O)2] + en(ClO4)2
Presence of the oxygen in the complex would probably lead to the formation of a hydroxide on the addition of water:
[Ni(en)O(H2O)2] + H2O --> [Ni(en)(OH)2(H2O)2]
Thinking about it, probably both en ligands have evaporated from or moved out of the complex, pulling in 2 water from the environment
to replace the second en, yielding [NiO(H2O)4]. This, because I would expect [Ni(en)(OH)2(H2O)2] to be light blue as well.
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woelen
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The fact that the solution is not clear is not that surprising. Suppose the complex is labile and loses (en) in solution. Then you get Ni(2+) and
(en). The molecule (en) is quite basic and you easily get enH(+) in solution, together with OH(-). This OH(-) could replace a water ligand of Ni(2+)
and in this way an insoluble nickel hydroxide compound can be formed (possibly with another (en) still attached to it). No need to evoke the idea of
oxidation by oxygen from the air. Nickel(II) is not easily oxidized at all, it requires very strong oxidizers like peroxodisulfate. Even the strong
oxidizer oxone (peroxomonosulfate) is not capable of oxidizing nickel(II) as I have test recently.
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Bezaleel
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Okay, but based on your initial experimental results, it seems strange that the complex would be labile, as addition of a drop of (en) to your neutral
solution containing Ni2+ did not yield a precipitate, according to your attempted synthesis of the bis-complex.
I agree that "oxidation" in my previous post is a misnomer; I meant to say that oxygen reacted with Ni2+ in a replacement reaction, effectively
replacing 2 ClO4-. As you can tell by the formulas I gave, no change of oxidation state of the Ni is involved.
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woelen
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Your remark about adding drops of (en) sure makes sense. Even adding a tiny amount of (en) does not cause any formation of a precipitate, the liquid
simply becomes cyan instead of green in that case. This is in contrast with copper(II) and ammonia. When a drop of ammonia is added to a solution of
copper sulfate, then you get a light blue precipitate of copper hydroxide. Only if a lot more ammonia is added, the precipitate dissolves again and
the deep blue ammine-complex is formed. What I can do is make a royal blue solution of nickel(II) and (en) and keep this in a stoppered test tube for
a long time. Maybe it turns turbid over the course of days.
Regarding your remark about oxygen replacing something, then it goes from oxidation state 0 to oxidation state -2. Then there must also be something
which is oxidized. Which is that 'something'?
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Bezaleel
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The theory I suggested was that (en) got oxidised:
| Quote: Originally posted by Bezaleel | I am tempted to conclude that what you got was somewhere near an oxidised mono-en complex of nickel:
[Ni(en)2(H2O)2](ClO4)2 --> dehydration + oxidation --> [Ni(en)O(H2O)2] + en(ClO4)2 |
Now that I think about it, this won't happen, as en(ClO4)2 does not exist. What may happen instead is the following, where water is the "oxidiser":
[Ni(en)2(H2O)2]2+(ClO4-)2 + H2O -->
[Ni(en)O(H2O)2]0 + en(H+ClO4-)2
Edit: Of course, (en)(ClO4)2 might not be stable and not be formed at all. Instead Cl2O7, (en) and H2O might escape....
[Edited on 17-9-2013 by Bezaleel]
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woelen
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Yes, what you write here could occur, but this is not called "oxidizing", but hydrolysis. I personally am inclined to think that not a single
O(2-)-unit connects to the nickel ion, but two hydroxide-units, so I propose a somewhat modified reaction equation:
Ni(en)2(H2O)2(ClO4)2 + 2 H2O ---> Ni(en)(OH)2(H2O)2 + (en)(HClO4)2
The latter can better be written as (enH2)(ClO4)2, the H(+) units are coordinated to the nitrogen atoms of (en) and actually you get
(+)H3NCH2CH2NH3(+).
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Bezaleel
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So, what you think is that the brownish complex consists of Ni(en)(OH)2(H2O)2. Earlier you said (bold face added):
| Quote: Originally posted by woelen | (...)
The dark green/brown paste was not kept around, some water was added to it to see whether it can be reverted to the nice deep blue solution. This is
not the case. The paste does dissolve fairly easily, but it gives a turbid pale blue liquid. Probably the pale liquid contains some of the
bis-ethylene diamine complex, but also a lot of nickel hydroxide.
(...) |
Did you also add a bit of (en) to the hydroxide that remained in the tube, to see what would happen? I would expect that from hydrated (wet)
"Ni(OH)2", you could make Ni(en)(OH)2(H2O)2, by the addition of (en).
My guess is, however, that this will give a light blue complex, same colour as [Ni(en)(H2O)4]2+, and not a dark green/brown one.
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PHILOU Zrealone
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| Quote: Originally posted by woelen | I did some experiments with nickel(II) ions and ethylene diamine. This combination affords remarkable colorful chemistry and as a bonus one can make a
nice energetic compound as well.
http://woelen.homescience.net/science/chem/exps/Ni_en_comple...
I hope to do more experiments, one of my goals is to isolate the perchlorate salt of the deep blue bis(en) complex. I can imagine that this is even
more energetic than the perchlorate salt of the tris(en) complex, because of better oxygen balance. The effect may be neutered, however, by the
presence of two water ligands in the bis(en) complex, but if I can isolate this complex, then I certainly will try this and compare it with the
tris(en) complex of which I still have almost 1 gram around and which can be stored safely. |
As usual  Nice work woelen!
About the bis(en) complex you may reduce the putative "neuterization" of the two water ligands by passing some dry NH3 gas through the dry complex to
expell the H2O and replace those by NH3...
You would then end up with Ni(en)2(NH3)2(ClO4)2.
Alternatively you could make a tiny test with hydroxylamine...
But I don't think hydroxylamine as a ligand is wel defined (may complexate via the O atom or via the N atom or both) and it could be potentially
dangerous owing to its inherent reductive power...to mind comes the infamous related compound with hydrazine as a ligand wich is uncompatible with Ni
perchlorate and that is said to explose even when dilluted into water solution by the shock of a glass rod while swirling into the mix onto the walls
of the beacker.
I did some testing long time ago with with Cu(NO3)2 and Ni(NO3)2 aside with ethylenediamine...I got deep dark blue-violet cristals of parallelipipedal
fashion. But I use very concentrated ethanolic solutions and I added ether for precipitation and drying of the complexes...
What is surprising is that despite their apparent brotherhood in their complexing abilities Cu, Ni and Co complexes displays very different properties
(color, stability, number of ligation sites and geometry).
Cu(NO3)2 is turquoise blue:
-tetracoordinable
-and turns into a deep blue amino cristaline complex soluble into water
-its hydrazino complex is very unstable and unsoluble most of the reactants turns black brown with N2 evolution (probably Cu and CuO aside with some
Cu2O) but some forms an unstable turquoise blue precipitate that bruns energetically with a green blue flame flash.Thanks to you I now imagine that
there can be 2 complexes like Cu(N2H4)(H2O)2(NO3)2 and Cu(N2H4)2(NO3)2... and one of the two is hell unstable (immediate decomposition) and the other
is unstable (spontaneous ignition upon drying).
-its ethylendiamine complex is deep dark violet and soluble into water
Ni(NO3)2 is emerald green:
-hexacoordinable
-and turns into a deep blue amino cristaline complex soluble into water
-its hydrazino complex is stable and unsoluble like a pale pink lillac precipitate with transitory blue and violet blue precipitate (high
concentration of Ni and low concentration of N2H4 at contact). The precipitate bruns energetically and detonates upon mild confinement and heating.
-its ethylendiamine complex is deep dark violet and soluble into water
Co(NO3)2 is ruby red:
-hexacoordinable
-and turns into a deep ruby red amino cristaline complex soluble into water. I would have tought it to provide also a blue complex...but chemistry is
full of wonders 
-its hydrazino complex is stable and unsoluble like a pale orange-brown precipitate. The precipitate bruns energetically and detonates upon mild
confinement and heating.
-i have not performed the ethylendiamino complex.
So all a world of colors and properties full of surprises.
Did you tried ethylendiamine ligand with Ni(III) perchlorate?
It could also be a way to improve oxygen balance...but I don't know the number of coordination sites for Ni(III).
Did you tried ethylendiamine ligand with Co(II) and (III) perchlorate/nitrate?
PH Z (PHILOU Zrealone)
"Physic is all what never works; Chemistry is all what stinks and explodes!"-"Life that deadly disease, sexually transmitted."(W.Allen)
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woelen
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I added (en) to the turbid liquid and when this is done, it immediately becomes clear and purple, due to formation of the tris(en) complex. I added
too much (en) for formation of intermediate complexes.
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nezza
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Thanks Woelen. Another interesting preparation. I have prepared a few grams of the solid which is a beautiful lilac colour. I will post pictures and
videos shortly.
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woelen
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@PHILOU Zrealone: Thanks for your interesting information. I certainly will continue my research on this kind of complexes. I now have plenty of
ethylene diamine and can continue with many more experiments. I already did the experiment with cobalt(III), but I did not yet isolate the complex.
@nezza: I am looking forward to your pictures and videos. Always good to see other's work and comparing my results with yours.
@Bezaleel: At the moment I have a test tube with the dark blue complex, dissolved in water. The test tube is stoppered. I just let it stand for
several days to see whether the solution remains clear, or decomposition occurs with formation of a pale precipitate.
[Edited on 18-9-13 by woelen]
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nezza
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As promised a picture and video (200fps) of the Nickel Ethylene Diamine perchlorate.
Attachment: Nickel EN Perchlorate2.tif (1.4MB)
This file has been downloaded 1158 times
Attachment: My Movie.mp4 (1.8MB)
This file has been downloaded 980 times
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woelen
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Thanks for this video and picture. Your sample looks very much like mine, so that is a good sign. My sample looks slightly more blue, but the picture
was made in daylight. Under TL-light, the material has exactly the same color as shown by your picture.
The video is particularly interesting. Which model camera did you use for making that video? I consider buying a new digital camera and I am looking
for one with high speed video capabilities.
[Edited on 19-9-13 by woelen]
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nezza
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The camera is a Panasonic Lumix FZ 200. The PAL version does 100fps @ 1280x720 (HD) and 200fps @ 640x480 (VGA) resolution. The American (NTSC) version
does 120fps & 240fps at the equivalent resolutions. Its expensive but a good camera for still and video. Further technical information on the
video :-
Focusing can be an issue with small objects off centre in the frame so I use a white sheet of print with a light shining on it in the same focal plane
as the subject to focus and set exposure.
Exposure compensation -3 stops (Any pyrotechnics is going to be very bright compared to incident illumination)
Start the video and remove the paper.
Any post editing is done in Windows Movie Maker
The VGA videos are good enough to be upscaled to 1024x768 for viewing on a larger screen.
As for the colour of the salt, the uploaded image does look too magenta compared with reality.
Hope this info is of use.
[Edited on 19-9-2013 by nezza]
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woelen
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Thanks for this information. This camera is a good candidate for what I buy.
Also nice to see that the color of the picture is somewhat too maganta, compared with reality. So, your compound was more bluish than the picture
suggests?
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bfesser
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<a href="http://en.wikipedia.org/wiki/Tagged_Image_File_Format" target="_blank">TIFF</a> <img src="../scipics/_wiki.png" /> isn't
particularly well-suited for the web, so here are <a href="http://en.wikipedia.org/wiki/Portable_Network_Graphics" target="_blank">PNG</a>
<img src="../scipics/_wiki.png" /> versions of your image. I've slightly blue shifted the one on the right; is it closer to the observed color?
 
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