Sciencemadness Discussion Board - Interesting and colorful coordination chemistry with ethylene diamine

Interesting and colorful coordination chemistry with ethylene diamine

woelen - 9-9-2013 at 11:45

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.

AndersHoveland - 9-9-2013 at 12:11

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...

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]

Bezaleel - 10-9-2013 at 06:05

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...

AndersHoveland - 10-9-2013 at 06:25

Quote: Originally posted by Bezaleel

Is the salt with sulphate also somewhat energetic or not at all?

No, it would not be the slightest bit energetic at all. For most practical purposes in pyrotechnics, sulfate is not even considered an oxidizer.

DraconicAcid - 10-9-2013 at 08:08

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.

woelen - 10-9-2013 at 11:35

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.

DraconicAcid - 10-9-2013 at 13:49

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.

atomicfire - 11-9-2013 at 12:44

How might one obtain some ethylene diamine?

woelen - 15-9-2013 at 08:58

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.

Bezaleel - 16-9-2013 at 09:10

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.

woelen - 16-9-2013 at 23:08

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.

Bezaleel - 17-9-2013 at 03:29

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.

woelen - 17-9-2013 at 03:58

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'?

Bezaleel - 17-9-2013 at 05:45

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]

woelen - 17-9-2013 at 06:13

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(+).

Bezaleel - 17-9-2013 at 12:30

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.

PHILOU Zrealone - 18-9-2013 at 03:46

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?

woelen - 18-9-2013 at 03:58

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.

nezza - 18-9-2013 at 11:10

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.

woelen - 18-9-2013 at 12:13

@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]

nezza - 19-9-2013 at 01:15

As promised a picture and video (200fps) of the Nickel Ethylene Diamine perchlorate.

Attachment: Nickel EN Perchlorate2.tif (1.4MB)
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Attachment: My Movie.mp4 (1.8MB)
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woelen - 19-9-2013 at 02:12

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]

nezza - 19-9-2013 at 02:54

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]

woelen - 19-9-2013 at 10:20

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?

bfesser - 19-9-2013 at 10:32

<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?

blue_shifted Nickel EN Perchlorate2.png - 867kB

phlogiston - 19-9-2013 at 11:27

There are substantial differences between monitors in color rendering, so unless you guys are using color calibration on your displays, it is going to be difficult to compare subtle color differences like this.

bfesser - 19-9-2013 at 14:05

That is true, but I was just looking to get closer. Not an exact match. I just thought I'd give it a go while converting it to PNG.

nezza - 19-9-2013 at 23:45

The colour is better, but its lighter than that bfesser. As for the Nickel complex, has anyone tried to make an analogous cobalt or copper complex. I know that in alkaline conditions the cobalt would be trivalent, but it certainly complexes with en as does copper.

woelen - 19-9-2013 at 23:54

Bfesser, the color of your modified picture is very close to my sample. Maybe nezza's sample is lighter because of other particle properties (his particles look somewhat more 'rough' than mine, which are smaller, but have smoother surface). This kind of differences may make a compound look darker or less dark. Another issue is the lighting used when the picture was made. In my picture, the material looks somewhat darker, but this may be because the light was somewhat less intense. Hard to compare!

nezza - 21-9-2013 at 01:38

I have attempted the synthesis of a cobalt analogue of the Nickel ethylene diamine(en) perchlorate as below :-

1. Dissolve cobalt in Nitric acid (50:50 conc/water) - use excess cobalt.
2. Take an aliquot and add en until no further colour change is observed.
3. Add 100 vol H2O2 to the cold or warm(not hot) liquit to complete oxidation to Co(III). It will froth.
4. Heat and add concentrated hot Ammonium perchlorate.
5. Allow to cool.
6. Yellow-brown crystals will precipitate out of the solution.
7. Filter and wash sparingly with water.

Notes

The cobalt solutions are very dark so colour changes are difficult to observe.
I have no idea of the formula of the precipitate which looks impure (It has darker, almost black specks in)
I have added a picture of the damp material and will look at its reaction when heated once it has dried.
Next stop copper.

nezza - 22-9-2013 at 01:35

I have now had a look at the dry Cobalt salt. It deflagrates in a similar manner to the Nickel salt when heated. I have also prepared some dark blue crystals of a copper analogue ?? in low yield which I will show when they are dry. Video of the cobalt salt attached.

Attachment: Cobalt deflagrating.mp4 (1.9MB)
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nezza - 22-9-2013 at 02:40

Finally pictures and a video of the copper complex. Lovely blue flashes and a green coloured flame.

Attachment: CopperVGA.mp4 (1.3MB)
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woelen - 23-9-2013 at 23:17

I have done some research on cameras and I ordered the same type of camera as you have. I have been doubting between this one and an Exilim high speed camera, but the fact that the Panasonic has f/2.8 aperture even at full zoom level made me decide to take it. I could obtain it for EUR 399, which seems to be a rather good price for this beast. I have the camera already, I am still waiting for an SDHC card (16 GByte, SHI-I 600x).

I also intend to make the copper and cobalt complexes and I certainly want to make videos of the burning of these complex perchlorate salts. Did you make the copper salt in the same way as the nickel- and cobalt-salts. Just adding slight excess of (en) and then adding ammonium perchlorate?

I myself already tried what happens when vanadyl ion is added to (en), but apparently no complex is formed with that. When a little (en) is added to a solution of VOSO4, then a dark brown precipitate is formed. This is hydrous VO2. When more (en) is added, then the precipitate redissolves and a red/brown solution is obtained, which looks exactly the same as the red/brown solution, obtained when excess NaOH is added to a vanadyl(IV) solution. This red/brown solution contains the red/brown hypovanadate ion, V4O9(2-).

[Edited on 24-9-13 by woelen]

nezza - 25-9-2013 at 01:03

Hi Woelen. The method for copper is the same as for nickel. Concentrated solutions have to be used and the precipitate is quite soluble so the yield is not very high. I cooled the solution in an ice bath for an hour or so and got the crystals pictured.

woelen - 13-10-2013 at 05:17

I tried the experiment with copper. This is not as easy as the preparation of the nickel complex. I have the impression that I do not get any crystals from the solution, while it is quite concentrated.

I used CuO and 30% HClO4. I dissolved a slight excess amount of CuO in 30% HClO4. This requires quite some heating to get the last amount of HClO4 reacted. The final product is a turbid dark blue solution. When one or two drops of HClO4 are added to the turbid solution, then it becomes clear and deep blue, but then it also contains excess HClO4, which is not what I want.

To this turbid blue solution, I added some pure ethylene diamine. This reaction is more exothermic than I expected. The result is a deep blue liquid, which after standing for a while becomes turbid and on heating, it does not become clear again, but light blue solid material deposits on the bottom. This most likely is Cu(OH)2 or some basic copper perchlorate. So, I added quite a lot more ethylene diamine. When this is done, then the solution turns more purple/blue instead of deep royal blue. This purple solution remains clear, also on heating to boiling.

I now have the blue/purple solution standing and hope to see formation of crystals from this solution. This solution is a little viscous and it sticks to the glass.

---------------------------------------------------------------------

One day later: A large amount of crystalline solid has settled at the bottom, I am really amazed to see how much solid separates from the liquid. The liquid apparently can be hugely oversaturated.
I put the liquid in a freezer for some time, such that its temperature drops to well below 0 C, but not so cold that the dark blue liquid freezes. When this is done, then even more crystalline solid separates.

I decanted the ice cold purple/blue liquid from the crystal mass and put the material on a filter paper, which in turn is put on a pile of paper tissue. I firmly pressed the material in the filter paper between the paper tissue. By repeating this two times with fresh paper tissue, I obtained a fairly dry crystal mass (crystal size in the order of magnitude of 1 mm). The damp crystals were put in a warm dry place and after a few hours they were perfectly dry.

Properties of the material:
- dark blue/purple crystals
- dissolves in water with some difficulty (e.g. like KClO3)
- the solution is clear and has a very deep purple/blue color
- non-hygroscopic

On heating in a flame, the crystals pop and crackle with a beautiful blue flash. If the material is put on a small spatula, and the spatula is heated from below, then when the temperature reaches a sufficiently high value, the material explodes, giving only a weak blue flash, but an impressively loud and full report.

Compared to the nickel-complex I have the impression that this material is more explosive. It ignites more easily and the audible effect is stronger.

Pictures and video will follow soon.

[Edited on 14-10-13 by woelen]

PHILOU Zrealone - 14-10-2013 at 10:56

When playing with such complexes of ethylendiamine, you can without too much troubles induce the cristallization by playing with saturated aqua-ethanolic or aqua-methanolic solutions of the salt and of EDA...
If it doesn't cristallize on its own in the cold; then you can push a little by adding a few ml of diethyl-ether.

I usually do this by solvent migration (evapo-condensation induces a solvent gradient) in a closed vessel with ether in the bottom and the cup of interest in the middle (holding the aqua-alcool metalo-amino complex).Ether evaporates and condenses at the surface of the aqua-alcool recipient what evaporates much slowly than ether does.

Such complexes are usually unsoluble into ether...the ether phase remains uncolourised!
I keep some under ether for years into transparent film canishers...

bfesser - 14-10-2013 at 11:18

Perhaps a stupid question, but do either of you plan to attempt syntheses of analogue Fe2+ complexes?

woelen - 14-10-2013 at 22:53

@PHILOU Zrealone: What you say is interesting and I will certainly try that the next time I run into problems with crystallizing the materials. Using ether has the advantage that it leaves no residue, it is very volatile and once the crystals have formed, you can easily get rid of the ether by evaporation.

@bfesser: How do you envision the separation of the Fe(2+) complex in any state of purity? Fe(2+) in alkaline liquids is very sensitive to air, I think that it sucks oxygen from the air, producing iron(III). Another issue is that I do not expect any coordination of (en) with Fe(2+). I once tried with ammonia and iron(II) nor iron(III) form coordination complexes with ammonia. Because of the similarities of (en) compared to two NH3-ligands, I also expect the (en) to just cause formation of a precipitate of hydroxide.
Of course I can try. I'll prepare a solution of FeSO4 and add some (en) and see what happens. I'll come back on that later.

nezza - 14-10-2013 at 23:40

I'm glad to hear of your success with Copper Woelen. I look forward to the video. As for bfesser's question about Fe(II) complex, I too have my doubts about its formation as Iron in alkaline solution is very easily oxidised to Fe(III) and Iron does not form complexes as easily as the other metals we have used. Again I can try.

bfesser - 15-10-2013 at 03:55

woelen, I was thinking Schlenk techniques were necessary. I'm also looking forward to the pictures and video you promised of the Cu2+ complex.

Like I said, it was a stupid question. Thanks for the replies, though.

woelen - 15-10-2013 at 06:46

Reading this thread raised some concerns in me:

http://www.sciencemadness.org/talk/viewthread.php?tid=1778

It is about exotic primaries, based on transition metal ammine and ethylenediamine complexes with oxidizing anions as counterions. Most of the compounds, described in this thread, are peroxodisulfates. What concerns me is the instability on storage of some of the compounds. Some compounds decompose overnight (e.g. the ethylenediamine complex of copper with peroxodisulfate as counterion), others ignite after some time, especially when put in sunlight.

I always had the impression that perchlorates were quite stable (the ion is quite inert at room temperature), but after reading this thread I am concerned a little bit. I have a small vial of the nickel complex and a somewhat bigger vial of the copper complex. Could it be that the stuff ignites or explodes one bad day in the (near) future, simply by storing it at room temperature? I hardly can imagine that this can happen, but if someone over here with some authority can elaborate on that, then that would be very nice and hopefully can take away my concerns. I do not want unstable, potentially self-igniting or self-exploding stuff around.

Maybe PHILOU Zrealone can comment on this?

DraconicAcid - 15-10-2013 at 08:17

Quote: Originally posted by woelen

@bfesser: How do you envision the separation of the Fe(2+) complex in any state of purity? Fe(2+) in alkaline liquids is very sensitive to air, I think that it sucks oxygen from the air, producing iron(III). Another issue is that I do not expect any coordination of (en) with Fe(2+). I once tried with ammonia and iron(II) nor iron(III) form coordination complexes with ammonia. Because of the similarities of (en) compared to two NH3-ligands, I also expect the (en) to just cause formation of a precipitate of hydroxide.
Of course I can try. I'll prepare a solution of FeSO4 and add some (en) and see what happens. I'll come back on that later.

Iron doesn't complex with ammonia, but en is a bidentate ligand, so its complexes tend to be orders of magnitude more stable. It's worth a shot.

woelen - 15-10-2013 at 12:03

I tried the experiment with iron(II) and iron(III). None of these forms a nice complex, which can be isolated, but with iron(II) there is something interesting. Below follows what I did and what results I had.

Experiment 1:
-----------------
Add a solution of ferric chloride to an excess amount of a solution of (en): This leads to formation of a red/brown slimy and flocculent precipitate of iron hydroxide. Nothing interesting happens.

Experiment 2:
----------------
Add a solution of ferrous ammonium sulfate to an excess amount of a solution of (en): This leads to formation of a slimy and flocculent precipitate with a light yellow/brown color. This is different from other experiments with iron(II). Precipitates with that usually are light green. Apparently some complex with iron(II) is formed, but this complex most likely is basic and also has hydroxide ion incorporated, hence its low solubility.
The precipitate is very air-sensitive. In contact with air, it first darkens (becomes dark grey, probably due to formation of hydrous Fe3O4) and then it quickly becomes rust-colored, due to formation of hydrous Fe2O3 or Fe(OH)3).

In both cases of iron(II) and iron(III) one can safely conclude that isolating a pure crystalline complex is out of the question.

woelen - 19-10-2013 at 09:31

As promised, a webpage about the copper(II) complex:

http://woelen.homescience.net/science/chem/exps/Ni_en_comple...

bfesser - 19-10-2013 at 10:40

Beautiful work, woelen. If I'm ever able to experiment again, this is the first synthesis I hope to do.

PHILOU Zrealone - 24-10-2013 at 12:59

Quote: Originally posted by woelen

Reading this thread raised some concerns in me:

http://www.sciencemadness.org/talk/viewthread.php?tid=1778

It is about exotic primaries, based on transition metal ammine and ethylenediamine complexes with oxidizing anions as counterions. Most of the compounds, described in this thread, are peroxodisulfates. What concerns me is the instability on storage of some of the compounds. Some compounds decompose overnight (e.g. the ethylenediamine complex of copper with peroxodisulfate as counterion), others ignite after some time, especially when put in sunlight.

I always had the impression that perchlorates were quite stable (the ion is quite inert at room temperature), but after reading this thread I am concerned a little bit. I have a small vial of the nickel complex and a somewhat bigger vial of the copper complex. Could it be that the stuff ignites or explodes one bad day in the (near) future, simply by storing it at room temperature? I hardly can imagine that this can happen, but if someone over here with some authority can elaborate on that, then that would be very nice and hopefully can take away my concerns. I do not want unstable, potentially self-igniting or self-exploding stuff around.

Maybe PHILOU Zrealone can comment on this?

Of course I can comment on this

You have to take in account:
1°)the inherent stability or explosive properties of the couple amine-anion....
For example NH4NO3 is stabler than NH4ClO4 what are safer than N2H5NO3 itself safer than N2H5ClO4...

Primary amines are stabler than ammonia itself stabler than hydrazine...but where are placed NH2-OH, R-O-NH2, R-NHOH, R-NH-NH2, R2-N-NH2, R-NH-NH-R', ... to be placed?

Perchlorates are by definition more heat and shock sensitive than nitrates...and perchlorates are safer than chlorates... but what to say about persulfates, iodates, bromates, periodates, perbromates, nitroformates, tetranitroethandiates,...

2°)the heat of combustion and oxygen balance.

3°)I wouldn't trust peroxydes or peroxoanions especially when sticking to metallic core...what are known to be uncompatible.

4°)the oxydability of the amine or the power/sensitivity of the oxydant...it is an oxydoredox couple...as such it wants to be in its stablest form...decomposed

5°)the oxydoredox potential of the metalic cation

6°)the metallic core may induce a catalytic effect or a photosensitivity....

N2H5NO3 or N2H5ClO4 are relatively hard to detonate even in contact with the flame of a match, they simply burn... but Ni(N2H4)3(NO3)2 almost D2D transits while Ni(N2H4)3(ClO4)2 is reputed to detonate in water solution...

Long time ago I had an unexpected detonation of N2H5NO3 in a galvanized pipe due to contact with the metallic copper of the detonator.
Obviously the traces of HNO3 in the N2H5NO3 eather reacted with Cu to generate Cu(2+) and NxOy or dissolved traces of CuO and Cu(OH)2 generating Cu(2+); then Cu(2+) and NxOy were in contact with the N2H4.HNO3 leading to decomposition of the hydrazine, formation of the unstable Cu(N2H4)2(NO3)2 and overheating!
By chance no person was armed, only material damages!

I worked with Ruthernium (III) with tris (diamino) complexes for my end study work and they where light sensitive.

PHILOU Zrealone - 25-10-2013 at 15:26

Quote: Originally posted by PHILOU Zrealone

Quote: Originally posted by woelen

Of course I can comment on this

You have to take in account:
1°)the inherent stability or explosive properties of the couple amine-anion....
For example NH4NO3 is stabler than NH4ClO4 what are safer than N2H5NO3 itself safer than N2H5ClO4...

Primary amines are stabler than ammonia itself stabler than hydrazine...but where are NH2-OH, R-O-NH2, R-NHOH, R-NH-NH2, R2-N-NH2, R-NH-NH-R', ... to be placed?

Perchlorates are by definition more heat and shock sensitive than nitrates...and perchlorates are safer than chlorates... but what to say about persulfates, iodates, bromates, periodates, perbromates, nitroformates, tetranitroethandiates,...

2°)the heat of combustion and oxygen balance.

3°)I wouldn't trust peroxydes or peroxoanions especially when sticking to metallic core...what are known to be uncompatible.

4°)the oxydability of the amine or the power/sensitivity of the oxydant...it is an oxydoredox couple...as such it wants to be in its stablest form...decomposed

5°)the oxydoredox potential of the metalic cation

6°)the metallic core may induce a catalytic effect or a photosensitivity....

N2H5NO3 or N2H5ClO4 are relatively hard to detonate even in contact with the flame of a match, they simply burn... but Ni(N2H4)3(NO3)2 almost D2D transits while Ni(N2H4)3(ClO4)2 is reputed to detonate in water solution...

Long time ago I had an unexpected detonation of N2H5NO3 in a galvanized pipe due to contact with the metallic copper of the detonator.
Obviously the traces of HNO3 in the N2H5NO3 eather reacted with Cu to generate Cu(2+) and NxOy or dissolved traces of CuO and Cu(OH)2 generating Cu(2+); then Cu(2+) and NxOy were in contact with the N2H4.HNO3 leading to decomposition of the hydrazine, formation of the unstable Cu(N2H4)2(NO3)2 and overheating!
By chance no person was armed, only material damages!

I worked with Ruthernium (III) with tris (diamino) complexes for my end study work and they where light sensitive.

So in conclusion risk factor can be evaluated by:
1°)The delta-H of reaction of the amino-anion couple is important (entropic reactants store more energy (like hydrazine))...

The higher it is, the higher the energy output and the driving force for decomposition.

2°)The delta-G linked to oxydoredox couples of the amino-metallic-anion triangle.
Each couple has to be evaluated as part of an electric cell in shortcut...

The delta-G allows one to see if a reaction is favourable or not, if it is on the edge of happening or if there is a safer activation energy barrier with relation to the temperature and the entropy.

3°)The activation energy of the amino-anion couple and the lowering of this activation energy barrier induced by the metallic core!

In a system with a negative delta-H, the catalytic effect of certain metallic core may reduce the activation energy barrier by modifying the chemical pathway and its critical step; so at a same temperature the average number of molecule with a sufficient energy to pass from one side of the barrier to the other exothermically (thus heating up the rest of the molecules) is much higher. If the system is not strong enough to dissipate the energy (dillution, venting, radiation, convection, vibrational modes...) then the system will go to exothermic runaway (take fire, burst, explode or detonate).

The size of the system is of major importance because the ratio surface/volume drops down very fast with increasing size... this explains the main risk of scaling up of an apparently masterized chemical process ... without taking in account thermodynamic properties...even a factor 2 can lead to dramatic consequences.

4°)Also last but not least...the use of dilluted solutions is a must in the case of complex coordination chemistry.
As the one that have performed the experiment have certainly noticed: the complexation reaction heats a lot... this is due to the loss of freedom levels of the molecules.
In the case of Ni(NO3)2 and N2H4: adding anhydrous hydrazine onto dry nickel (II) nitrate will more than certainly lead to a fire or an explosion...when using saturated water solution of Ni(NO3)2 and N2H5OH (N2H4 at 80%) the reaction media goes well over the 50°C.
The vibrationnal energy of 3 molecules of hydrazine is trapped into the Ni(N2H4)3(2+) complex and provided to the surrounding molecules...the effect is enhanced by the precipitation process of Ni(N2H4)3(NO3)2 then the vibrationnal energies of Ni (2+), 3 N2H4 and 2 NO3(-) (so 6 molecules, anions and cations) are transfered to the media.

[Edited on 26-10-2013 by PHILOU Zrealone]

woelen - 26-10-2013 at 07:35

Thanks for your comments. If I read your comments, then I have the impression that the hydrazine-based complexes are much more risky than the (en)-based complexes. Hydrazine is much more endothermic than (en). I did quite some mishandling of the perchlorate salts of the (en)-complexes and the only thing which can have them explode is heating them in a flame and having them react all at once. Strongly crunching the solid material with a rough metal spatula does not cause ignition.

I now did the experiment with zinc as well. This gives a white crystalline compound, which is approximately as energetic as the nickel-complex. The procedure is as follows:

Prepare a concentrated solution of ZnCl2 in distilled water. Use appr. 1 part of ZnCl2 with 2 parts of water. When the ZnCl2 is dissolved in that amount of water, then the solution becomes quite warm.
Prepare a solution of appr. 1 part of (en) in 4 parts of water. Make around 10 ml of this solution. Slowly add this solution to the solution of ZnCl2. You get a white precipitate. When more of the solution of (en) is added, then the white precipitate redissolves again. Add so much solution of (en) that all white precipitate redissolves again. This requires quite some excess of (en), the liquid has a clearly noticeable smell of (en) when it is totally clear again.
In a separate test tube dissolve roughly 2 times as much of NH4ClO4 as ZnCl2 (not measured, just estimated on eye) in a small amount of water. Use heat to get such a concentrated solution of NH4ClO4.
Pour the boiling hot solution of NH4ClO4 and the (en)/ZnCl2 solution in a small erlenmeyer. When this is done, a white powdery precipitate is formed. Heat the liquid, until it becomes entirely clear again and boil for a while.
Finally, put the hot liquid aside in a dust-free and quiet place. Let stand overnight. After one night there are large needle-like white crystals. Place the liquid in a fridge and allow to cool to around 0 C to have a little more crystals.
Pour the ice cold liquid from the crystals and then put the wet solid on a piece of filter paper, which in turn is put on a pile of paper tissues. Fold the filter and the paper tissue and press firmly to have most adhering water absorbed by the paper tissue. Repeat this procedure two times with fresh dry paper tissue.
Finally, unfold the filter paper and allow the crystalline (damp) material to dry for a day in a warm place. Crunch the crystals somewhat and stir them up every few hours, to get all of them perfectly dry.

[Edited on 26-10-13 by woelen]

PHILOU Zrealone - 29-10-2013 at 10:23

Quote: Originally posted by woelen

Thanks for your comments. If I read your comments, then I have the impression that the hydrazine-based complexes are much more risky than the (en)-based complexes. Hydrazine is much more endothermic than (en). I did quite some mishandling of the perchlorate salts of the (en)-complexes and the only thing which can have them explode is heating them in a flame and having them react all at once. Strongly crunching the solid material with a rough metal spatula does not cause ignition.

I now did the experiment with zinc as well. This gives a white crystalline compound, which is approximately as energetic as the nickel-complex.

Wel concluded

Stil ethylendiamine diperchlorate on its own is quite a powerfull HE, but its sensitivity is not critical vs hydrazine perchlorate and a little more shock sensitive than the dinitrate (what is in the sensitivity range of TNT).
The main difference comes from the fact EDA is not as strong a reducer as hydrazine.

Nice to see Zn also makes a complex.
I would start from Zn(ClO4)2 to avoid chloride and ammonium.

Note that in principe all those complexes should be reachable via interaction of the oxydes or hydroxydes of the metals and the salt of the amine:

CuO + 2 H2N-CH2-CH2-NH2.2HNO3 --> Cu(EN)2(NO3)2 + H2O + 2 HNO3

2 CuO + 2 H2N-CH2-CH2-NH2.2HNO3 --> Cu(EN)2(NO3)2 + Cu(NO3)2 + 2 H2O

Ni(OH)2 + 3 H2N-CH2-CH2-NH2.2HClO4 --> Ni(EN)3(ClO4)2 + 2 H2O + 4 HClO4

3 Ni(OH)2 + 3 H2N-CH2-CH2-NH2.2HClO4 --> Ni(EN)3(ClO4)2 + 2 Ni(ClO4)2 +6 H2O

This was one of the plausible cause of the unexpected detonation of N2H4.HNO3 in contact with Cu(OH)2 and CuO at the surface of metallic copper detonator!

[Edited on 29-10-2013 by PHILOU Zrealone]

woelen - 29-10-2013 at 12:51

The reason why I prefer to use NH4ClO4 in many cases is that I do not want to use up my HClO4 quickly. NH4ClO4 is much easier for me to come by than HClO4. If I were in a lab with supply of chemical without restrictions, then I would use the metal oxide or carbonate in all cases and would dissolve that in aqueous HClO4.

With copper, however, I used HClO4 in which I dissolved CuO. This was because it was written that the copper complex is very soluble and less easily crystallized and in such cases it is best not to have foreign ions in solution as well.

My next one will be the cobalt(III) complex and maybe the silver(I) complex, although I am somewhat reluctant to make the latter. Silver(I) with ammonia is a dangerous combination on standing for a longer time and I can imagine that silver(I) with (en) also is dangerous, regardless of the anion. So, if I try the silver complex, I will try it on a very small scale.

I already tried chromium(III), iron(II), iron(III) and vanadium(IV) and all of these form a slimy/flocculent precipitate with (en), which does not lead to formation of clear solutions and easily crystallizable compounds.

[Edited on 29-10-13 by woelen]

DraconicAcid - 29-10-2013 at 13:36

When you try the iron with en, perhaps try it under anhydrous conditions? That would prevent formation of slimy hydroxides, and en should form a complex with Fe(II) (based on the fact that Lang's handbook gives a formation constant for it).

The chromium is probably only giving a slimy precipitate because it's non-labile. The complex [Cr(en)₃]³⁺ should be very stable, just slow to form.

[Edited on 29-10-2013 by DraconicAcid]

PHILOU Zrealone - 29-10-2013 at 14:26

Quote: Originally posted by woelen

The reason why I prefer to use NH4ClO4 in many cases is that I do not want to use up my HClO4 quickly. NH4ClO4 is much easier for me to come by than HClO4. If I were in a lab with supply of chemical without restrictions, then I would use the metal oxide or carbonate in all cases and would dissolve that in aqueous HClO4.

With copper, however, I used HClO4 in which I dissolved CuO. This was because it was written that the copper complex is very soluble and less easily crystallized and in such cases it is best not to have foreign ions in solution as well.

My next one will be the cobalt(III) complex and maybe the silver(I) complex, although I am somewhat reluctant to make the latter. Silver(I) with ammonia is a dangerous combination on standing for a longer time and I can imagine that silver(I) with (en) also is dangerous, regardless of the anion. So, if I try the silver complex, I will try it on a very small scale.

I already tried chromium(III), iron(II), iron(III) and vanadium(IV) and all of these form a slimy/flocculent precipitate with (en), which does not lead to formation of clear solutions and easily crystallizable compounds.

[Edited on 29-10-13 by woelen]

I use HClO4 because I have quite a lot of it at hand (70%).
But I don't have NH4ClO4.
Everybody has its own troubles (cross to bare)

.

By metathesis NH4ClO4 is as potent as HClO4...:
1°)It can turn many metal hydroxyde or oxyde into the desired metal perchlorate or amino metal perchlorate (what is relatily unstable and in open system free amonia by exchange towards water ligands).
2°)It can turn EDA into EDA diperchlorate and free NH3 as a gas.

There are many complexing metal specific of amines, even calcium (what I was unaware of) or mercury that has a special affinity for it.

Silver is particularly uncompatible with ammonia because of formation of fulminating silver Ag-NH2, Ag2NH and Ag3N (nightmare of the mirror makers); but other amines will be much less of a trouble... with exception of hydrazine what is immediately decomposed into metallic silver and nitrogen.

It is always a good idea to test in tiny quantity because the potency of the catalytic effect of the metallic core (of not wel documented complexes or unmade to date complexes) is not clearly known until done and tested towards time decay, temperature sensitivity, shock sensitivity and light sensitivity.

[Edited on 29-10-2013 by PHILOU Zrealone]

Chemstudent - 3-11-2013 at 07:34

I've been working on Nickel complexes in my Inorganic lab this semester. It really has been a fun experience. Working with Ni(II) complexes has really helped me learn the spectrochemical series and colour theory. Although some of these complexes have been super hard to produce (namely the water sensitive varieties)

I have one this week that I'll be tackling for a second time - [Ni(NO2)2(en)2] which is a red linkage isomer. I begin with (K4[Ni(NO2)6].H20) <-- Hexanitronickelate (II) which is an orange-clay compound. It is very hygroscopic so all the glassware & reagents must be as moisture-free as possible. Any moisture and you end up getting a purple aqua/nitro complex.

Last thing though, I am still unsure what commercial value or medical/scientific use these linkage isomers have. Is there any use for these compounds?

woelen - 3-11-2013 at 11:47

Your complexes, are they nitrito complexes or nitro complexes? If you can give more details on how to prepare them (and if I have the reagents and needed apparatus), then I certainly would like to try them myself.

---------------------------------------------------------

In the meantime, I now tried the cobalt(III) (en) complex. This is not hard to make, I took a fairly concentrated solution of cobaltous nitrate, but if you don't have that, you can use cobaltous carbonate and nitric acid to make such a solution.
To this solution, I added quite some excess amount of (en) and swirled. This leads to formation of a reddish/brown solution. In contact with air, this solution becomes MUCH darker. You can see the dark color near the surface and on the glass.
Next, I added 10% H2O2 while swirling and continued adding this, until there is weak effervescence. Initially, the H2O2 immediately reacts with the cobalt(II) complex and oxidizes this to cobalt(III). But at a certain point, the H2O2 is not used up anymore and then it slowly decomposes.
I then boiled the liquid for a while to be sure that all excess H2O2 is destroyed.
In a separate test tube I dissolved a slight excess amount of NH4ClO4 in a small amount of water (I needed to heat the water to near boiling). This solution of NH4ClO4 is added to the cobalt(III)/(en) solution and then the liquid is boiled again for some time, just to be sure that no solid material is present in the liquid.
On cooling down, a lot of crystalline material is formed. This is dried as described above in the thread. Pictures and a more extensive write-up will follow one of these days.

--------------------------------------------------------------------

I also tried the silver(I) complex, but this failed.
I prepared a solution of AgNO3 in water and added drops of (en) to this solution. Initially, this leads to formation of a yellow/brown precipitate but if enough (en) is added, then the precipitate redissolves again and a clear solution is obtained.
To this solution I added a solution of NH4ClO4. When this is done, then the liquid becomes turbid again, it becomes dirty white. On heating, the liquid becomes more turbid and color of the precipitate turns to light brown and part of the small particles coalesce to somewhat larger particles. No crystals are formed, a firtly looking water-insoluble precipitate is formed. I think it is a mix of Ag2O, AgOH, some Ag(en) complex and maybe ammonia complexes. Not something which I wanted to keep around for a long time (knowing the danger of ammonical silver solutions), so I destroyed all of it by adding excess dilute HNO3, which dissolves all of it to a colorless solution.

[Edited on 3-11-13 by woelen]

DraconicAcid - 4-11-2013 at 09:06

Quote: Originally posted by woelen

No crystals are formed, a firtly looking water-insoluble precipitate is formed. I think it is a mix of Ag2O, AgOH, some Ag(en) complex and maybe ammonia complexes. Not something which I wanted to keep around for a long time (knowing the danger of ammonical silver solutions), so I destroyed all of it by adding excess dilute HNO3, which dissolves all of it to a colorless solution.

I wouldn't expect the en complex to be particularly stable, since silver(I) pretty much always forms linear complexes, and ethylenediamine isn't long enough to chelate in a linear fashion.

woelen - 4-11-2013 at 10:22

Ah, that is interesting info you give about the linearly shaped complexes. This also explains why silver(I) usually has only two ligands coordinated to it (or one bidentate ligand).

I dried the cobalt complex. I think that the cobalt complex also is less stable than other complexes. I obtain a mustard-colored solid, while the liquid from which it is obtained is much more red (nearly as red wine, albeit with a brownish tinge). The yield is very low, when compared to the yield of the other complexes.
The solid material is very easy to ignite, it is much more sensitive to flame than any of the other complexes I made. It is not the most energetic though. The copper complex is more energetic.

DraconicAcid - 4-11-2013 at 10:27

Quote: Originally posted by woelen

I dried the cobalt complex. I think that the cobalt complex also is less stable than other complexes. I obtain a mustard-colored solid, while the liquid from which it is obtained is much more red (nearly as red wine, albeit with a brownish tinge). The yield is very low, when compared to the yield of the other complexes.
The solid material is very easy to ignite, it is much more sensitive to flame than any of the other complexes I made. It is not the most energetic though. The copper complex is more energetic.

I was about to post that this information is very surprising, as it's an extremely stable cation... but then I remembered you had perchlorate counterions. When I was a grad student, I assisted with a course that made a set of cobalt(III) en complexes every year, and the tris(en) complex was an impurity in most of the products.

woelen - 5-11-2013 at 07:16

Could you provide information about the other Co-(en) complexes you prepared in that course and do you have procedures for that? I might be able to modify these procedures, such that I can make similar complexes with the reagents which I have.

DraconicAcid - 5-11-2013 at 09:29

Alas! this was more than fifteen years ago, so I don't remember the exact procedures. I didn't keep the lab manuals after graduating.

Basically, we made [Co(en)₂(NO₂)₂]Cl by reacting CoCl₂, en, and sodium nitrate in a solution through which air was pulled through overnight (using a vacuum aspirator). Once this was isolated, we added hydrochloric acid to get the green trans-[CoCl₂(en)₂]Cl salt (this always looked like crap, because it was contaminated with the tris(en) complex, among other things). Heating the green complex in water with a steam bath converted it to the purple cis complex. We could also dissolve the green complex in water and add a concentrated solution of potassium nitrate, which formed the much-less-soluble trans-[CoCl₂(en)₂]NO₃, which were very nice green xtals (since the tris(en) complex and all the other crap that was in the original green stuff did not form a less-soluble nitrate).

In basic solution, these would hydrolyze to a pinkish [Co(H₂O)₂(en)₂]³⁺ ion.

I heard that they replaced the "overnight aspiration" prep with a hydrogen peroxide one, which took much less time. I seem to recall that activated charcoal was added as a catalyst.

At one point I tried replacing the the chlorides with oxalate (got a red material which might have been the aquo complex) or iodide (Co(III) plus a reducing agent- you can guess how well that worked).

Some references:
http://www.uiowa.edu/~c004153a/Co%28en%293-2004.pdf
http://chemlab.truman.edu/CHEM131Labs/CoordChem.asp

An attempt at tris(ethylene diamine) nickel(II) hydroxide

Bezaleel - 6-11-2013 at 08:17

Yesterday, I made an ethylene diamine nickel complex in the following manner. Half a spatula full of wet nickel(II)hydroxide was put in a test tube and about 2cm of demi water were added. To this approximately 0.7ml of ethylene diamine were added.

The hydroxide attached to the glass of the tube immediately received a pink-purplish layer on the outside. On the inside it remained green, and in between was a dark blue region. The hydroxide in the water turned a bit whitish. After shaking and standing for a minute, everything had become pink-purple, just as the tris(ethylene diamine) complexes produced by Woelen at the beginning of this thread. It surprised me that the liquid on top of the pink-purple precipitate was also pink-purple. Apparently the newly formed substance is somewhat soluble in water.

The picture was taken under incandescent light, and white balance corrected to represent the colour correctly. Under the light of mercury vapour lamps, the colour of the pictures became far too blue.

Note that the nickel(II)hydroxide was not entirely pure. It was created by adding an NaOH solution to a solution of black nickel oxide in hydrochloric acid solution. The hydroxide was used without further purification.

DraconicAcid - 6-11-2013 at 09:10

Nice nickel.

I've seen a sample of [Ni(en)₃]Cl₂, which was made by accident. The students were supposed to be making NEt₄NiCl₄ by dissolving nickel(II) chloride in ethanol, adding triethylformate, and then adding tetraethylammonium chloride. One student accidentally added ethylenediamine instead of triethylformate (they both said "ethyl" in the name somewhere, you can't blame her for not reading the entire label....honest!), and got an instant precipitate of a very similar purple material.

I'm not surprised that it's water soluble- many hydroxides of complex ions are perfectly water-soluble, in stark contrast to the simple hydroxides.

woelen - 6-11-2013 at 23:27

Interesting to see that you can get a hydroxide complex as well. This can be a good starting point for many other complexes. It would be interesting to see how it behaves if dilute sulphuric acid is added very carefully, such that the hydroxide is just neutralized. Does the hydroxide react first, or is the (en) protonated, resulting in destruction of the complex?

I have done a similar small experiment, by adding some CuO to a solution of (en) in water. If this is done, then the CuO partially dissolves, giving a blue/purple solution.

DraconicAcid - 15-11-2013 at 16:47

Quote: Originally posted by woelen

I already tried chromium(III), iron(II), iron(III) and vanadium(IV) and all of these form a slimy/flocculent precipitate with (en), which does not lead to formation of clear solutions and easily crystallizable compounds.

I was just looking up a prep for Cr(en)3 complexes- if you add a bit of granular zinc, it can catalyze the replacement of the ligands by reducing the Cr to the labile Cr(II) state. Try that.

Chemstudent - 22-11-2013 at 08:14

Quote: Originally posted by woelen

Your complexes, are they nitrito complexes or nitro complexes? If you can give more details on how to prepare them (and if I have the reagents and needed apparatus), then I certainly would like to try them myself.
[Edited on 3-11-13 by woelen]

Sorry for not having returned an answer sooner! I have since been able to synthesize a small quantity of the troublesome Nickel(II) complex which I had mentioned earlier. And no, it wasn't a nitrito complex. Below I will include any relevant information from my lab manual:

"The intent of this experiment is to extend the student's ability to synthesize, purify, and characterize coordination compounds. In addition, the ambidentate nature of the NO₂ ligand is investigated...."

Three linkage isomers of Nickel(II) were synthesized from the precursor: hexanitronickelate(II)monohydrate, K₄[Ni(NO₂)₆].H₂O

The precursor was prepared in ~40 minutes time, however was placed in a desiccator over an extended period for drying. See attached word document for detailed procedure. Il a été très facile!

Now...

Two of the linkage isomers found in the experiment were easily synthesized in good yield. One however took me 3 attempts to produce; even then the yields were quite poor. I have not included the procedure for the other two compounds (pm me if you want the others). Included is the procedure for producing the compound I found very troublesome.

The complex is: [Ni(NO₂)₂(en)₂] it was found to be ruby red in colour.

The procedure transcribed from my lab manual is attached in a second word document. For this procedure I had not offered additional input, rather I merely transcribed the directions. It is however important to note the reaction conditions should be moisture free as possible (wet glassware/reagents was a big issue).

Attachment: Precursor - Hexanitronickelate(II) monohydrate.docx (32kB)
This file has been downloaded 570 times

Attachment: red complex - [Ni(NO2)2(en)2].docx (96kB)
This file has been downloaded 543 times

DraconicAcid - 22-11-2013 at 11:54

Looks cool, Chemstudent. For the precursor, however, it says to make it in de-ionized water, but it also calls for 100 mL methanol (but you only wash with 15 mL methanol). Is there an error there?

Chemstudent - 22-11-2013 at 14:21

Quote: Originally posted by DraconicAcid

Looks cool, Chemstudent. For the precursor, however, it says to make it in de-ionized water, but it also calls for 100 mL methanol (but you only wash with 15 mL methanol). Is there an error there?

H₂O in the production of the precursor is used to help catalyze the reaction. The final product is a mono-hydrate complex, however it will have more H₂O in its coordination sphere which is why the final product must be dried in a desiccator. I was told by my laboratory instructor that the compound is hygroscopic, thus it must remain in a desiccator for extended storage. Further, it must be in the mono-hydrate state in order to successfully form the Ni(NO₂)₂(en)₂ complex. As stated, the second synthesis is very sensitive to moisture, and a wee drop can ruin the yield. H₂O in that capacity would act as a catalyst to push out ALL the NO₂, and the en would fully react with the Nickel to form a purple Ni(en)₄ complex. What is desired is to only displace 1/2 of the NO2 ligands, and replace them with the (en) ligand. Controlling not only moisture, but the rate of ethylenediamine addition was equally important. This synthesis really challenged my understanding of kinetic/thermodynamically favored products.

Lastly, I always state DI water because I never want there to be any confusion as to if other sources of H₂O are acceptable. I never use anything other than DI in lab procedures.

Oh, and the 100mL Methanol was just me saying have it on hand in a wash bottle in case you want to clean the product above and beyond what is necessary. I ended up using more to wash my product, but only because I am one who likes to go overboard. I wanted to clean off as much impurity as possible. It only hurt my final yield slightly and that I was not concerned with.

[Edited on 22-11-2013 by Chemstudent]

Ethylene diamine with MoO[sub]3[/sub]

Bezaleel - 29-11-2013 at 02:29

I tried to create a molybdenum complex with (en). As anticipated, this was not a success.

Half a spatula of MoO3 (ceramic grade, medium grey powder) was put into a test tube with some 3 cm of demiwater. This gave a whitish suspension on shaking. On heating the oxide did not dissolve, neither after addition of a few drops of concentrated H2SO4. To the still hot solution, 1 ml of (en) was added. Immediately the suspension became clear, leaving only the larger particles of MoO3 on the bottom of the test tube. Some of these were bluish, indicating that the starting material used did not have all molybdenum in the +6 oxidation state.

After addition of another ml of (en) and shaking the mixture near its boiling point for a few minutes, all of the oxide had dissolved into a completely clear solution, almost colourless, with only a slight hint of yellow.

I do not think that (en) did any complexation to the oxide. Instead, a (poly)molybdate has been formed, with (en) acting as the counter-cation.

MoO3 (s) + H2O + (en) (aq) --> MoO4(2-) (aq) + H3NCH2CH2NH3(2+) (aq)

MoO4(2-) may in turn have complexated with more MoO3 to form a polymolybdate anion, although I think that there was too much (en) in solution for this to happen.

woelen - 29-11-2013 at 05:51

In this case, (en) simply acted as a base, needed to dissolve the MoO3. Nothing really interesting. Apparently (en) is a strong enough base to dissolve MoO3.

DraconicAcid - 29-11-2013 at 15:56

Quote: Originally posted by Chemstudent

I tried this reaction today- that was interesting. Adding the nitrite gave a deep green solution, and slowly gave an orange-red precipitate. The filtrate was olive-green, until I washed the precipitate with alcohol and more compound precipitated from solution.

Bezaleel - 2-12-2013 at 07:50

Quote: Originally posted by woelen

Thanks [nezza] 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.
(...)
[Edited on 19-9-13 by woelen]

Funny that these compounds look different under filament/discharge lamp light. Here's the nickel(en)hydroxide, that I crystallised from its solution under filament lamp light. Though not visible in thepicture, the crystals are glistening a bit in the lamp light.

Chemstudent - 2-12-2013 at 19:49

Quote: Originally posted by DraconicAcid

That's great! So I assume you obtained the Orange-Red Hexanitronickelate complex? It looks like the clay found on a baseball field to be honest. Now leave it in a desiccator to dry and then you can go from there to the red complex. I will be very impressed when you do, but the final product is worth it. You get crimson red crystals. They are nice to look at.

[Edited on 3-12-2013 by Chemstudent]

DraconicAcid - 2-12-2013 at 19:56

Quote: Originally posted by Chemstudent

I don't actually have a desiccator, so I tried it briefly in the oven at low temp. I'll let you know when I try the final product. Couldn't compare it to the stuff on a baseball field, though- I'm not a baseball kind of guy.

Chemstudent - 2-12-2013 at 20:13

Quote: Originally posted by DraconicAcid

Quote: Originally posted by Chemstudent

Well, there may be the issue of pulling atmospheric moisture when cooling down the complex to room temp as it is somewhat hygroscopic. You would do well to quickly store the dry complex in a container immediately after pulling it from the oven so as to minimize adding moisture. Still shouldn't be an issue. And by low temp, I hope you mean no greater than 60C

[Edited on 3-12-2013 by Chemstudent]

DraconicAcid - 2-12-2013 at 20:35

It was about 50, and I sealed it up immediately after. It still looks nice.

Bisethylenediamminenickel(II) bistri-iodoplumbate(II)

Bezaleel - 12-1-2020 at 16:56

===A description with follow when I edit this post===

1.00g of NiCl2.6H2O in 25.5ml of water with ethylenediammnie (1:1 molar)

Addition of Ni(en) solution to a mixture of KI solution with Pb(AcO)2 (4+:1 molar)

Bisethylenediamminenickel(II) bistri-iodoplumbate(II)

Bezaleel - 17-11-2020 at 09:58

From this thread (I think this is a more logical thread to post it):

Quote: Originally posted by MidLifeChemist

Got it - ok I'm assuming the idea is that the iodide ions from the KI will replace the chloride ions. So why do you need to do this second step - to get the Ni(en)3I2 the first time, did you simply boil away or evaporate the solution, is that why you need to recrystallize? And how do you know this is the iodide and not the chloride? I couldn't find any references to Tris(ethylenediamine) nickel (II) iodide or to its solubility. It looks like an interesting compound that I may want to try to make one day. Thanks in advance for the info!

Quote: Originally posted by Bezaleel

Quote: Originally posted by MidLifeChemist

Quote: Originally posted by Bezaleel

Made some Ni(en)3I2. Dissolved this in dilute KI to recrystallise.

Did you do that to reduce the solubility via the common ion effect?

No, it's just a final purification step, to get rid of assumed co-crystallised chloride.

The synthesis was from nickel sulphate which contained a few % nickelchloride as impurity. Ni(en)3SO4 and chloride hava a much higher solubility at low temperatures. Hence adding KI solution will make the Ni(en)3I2 crystallise out on cooling. This is both concentration and pH critical (do add the full 3 moles of (en)).

I used 7.38 g of NiSO4/NiCl2 (assumed to be NiSO4) in 50ml water. Then added 5.7ml (en) to 50ml water, and mixed both solutions. Let stand for a while and filter off Ni(en)3(OH)2. Add at least 9.3g of KI and let crystallise in the cold (I used 5 C). I obtained a second batch by evaporating about half the solution and cooling again.
As said, recrystallisation was an assumed purification step. The iodide is able to form beautiful multi-faceted crystals, see pictures.
[Modified from Gmelins Handbuch, 58 C pp 152-158]

20170427_124507_detail_adj_den_small.jpg - 74kB

woelen - 17-11-2020 at 11:07

Very nice crystals! How large are they?

Bezaleel - 17-11-2020 at 15:34

These are close-ups. The long side is 4 mm. I'm making more of this compound to continue growing.

woelen - 18-11-2020 at 02:09

That's quite big already. My crystals of nickel complexes are at most 1 mm, but most of them are smaller, like table salt.
I really like the deep blue with a hint of purple of your crystals.

I also should try other anions than perchlorate. Perchlorate is nice and convenient, because it crystallizes so easily and for nearly all complex and bigger cations, it is not hygroscopic, but the disadvantage is that with this you cannot make large crystals. Due to its lower solubility, it tends to form many small crystals of sub-mm size.

I replenished my stock of perchloric acid two weeks ago (a new bottle of 500 ml of 60% acid), so that I can do more experiments with metal complexes, but I'll also try with HBr (I also have 300 ml of 48% HBr). With chloride I never had really success. With many metal complexes I had issues of concurrent coordination. Besides the ammine or (en) ligands, I also had chloride ligands, attached to the metal ions (especially with copper this is easily the case) and strong hygroscopic properties, making crystallizations very difficult. Perchlorate itself is nearly inactive as a ligand, and combined with its ease of crystallization, it almost is the perfect anion for this purpose. Apparently this also can be done very well with BF4(-), but this ion is prone to hydrolysis to some extent (formation of HF and borate)., which is highly undesirable.

Bedlasky - 18-11-2020 at 05:38

Woelen: Bromides are also strongly coordinating. Try sulfate or nitrate, they are weakly coordinating. Nitrate salt is energetic, so sulfate is better option.

Bezaleel - 18-11-2020 at 15:38

Quote: Originally posted by woelen

That's quite big already. My crystals of nickel complexes are at most 1 mm, but most of them are smaller, like table salt.
I really like the deep blue with a hint of purple of your crystals.

I also should try other anions than perchlorate. Perchlorate is nice and convenient, because it crystallizes so easily and for nearly all complex and bigger cations, it is not hygroscopic, but the disadvantage is that with this you cannot make large crystals. Due to its lower solubility, it tends to form many small crystals of sub-mm size.

I replenished my stock of perchloric acid two weeks ago (a new bottle of 500 ml of 60% acid), so that I can do more experiments with metal complexes, but I'll also try with HBr (I also have 300 ml of 48% HBr). With chloride I never had really success. With many metal complexes I had issues of concurrent coordination. Besides the ammine or (en) ligands, I also had chloride ligands, attached to the metal ions (especially with copper this is easily the case) and strong hygroscopic properties, making crystallizations very difficult. Perchlorate itself is nearly inactive as a ligand, and combined with its ease of crystallization, it almost is the perfect anion for this purpose. Apparently this also can be done very well with BF4(-), but this ion is prone to hydrolysis to some extent (formation of HF and borate)., which is highly undesirable.

Perchloric acid would be a really nice to have. Have you bought it or made it? Since it's practically non-coordinating, it would be a very good starter for making complexes. My guess is that Ni(en)3(ClO4)2 will have a high solubility, so addition of KI or KBr should provide the iodide or bromide variety on cooling. I' m guessing here, though, since Gmelin has some complicated entries on the Ni(en)x perchlorates.

You mention other ligands, like BF4-. I guess that's a small ligand? Maybe IO4- could also work? It's larger than the perchlorate, bond lengths are 144 and 178pm respectively.

Bedlasky - 18-11-2020 at 16:15

Periodate is strongly oxidizing. Ammonium periodate is shock sensitive, very unstable and explosive. [Ni(en)3](IO4)2 would be even worse - more reducing agent and Ni as a metal catalyst.

DraconicAcid - 18-11-2020 at 16:48

Tetrafluoborateis not a ligand- it is very non-coordinating. Unlike perchlorates, it's also non-oxidizing.

Bezaleel - 19-11-2020 at 02:02

Quote: Originally posted by Bedlasky

Periodate is strongly oxidizing. Ammonium periodate is shock sensitive, very unstable and explosive. [Ni(en)3](IO4)2 would be even worse - more reducing agent and Ni as a metal catalyst.

Thanks, I never worked with it, so I wasn't aware.

Quote: Originally posted by DraconicAcid

Tetrafluoborateis not a ligand- it is very non-coordinating. Unlike perchlorates, it's also non-oxidizing.

Correct, sorry about the wrong wording. It's just an ion, non-coordinating indeed.

(I must have been more tired than I thought wen posting my previous message...)

[Edited on 19-11-2020 by Bezaleel]

woelen - 19-11-2020 at 02:26

I bought my perchloric acid, together with another member I bought one liter and we shared the cost, so that we both have 500 ml.

BF4(-) indeed can be used as a non-coordinating anion. I have NaBF4, but I noticed, that in the long run, hot near boiling solutions of this attack glass (very slowly, but noticeably). The BF4(-) ion does hydrolyse a little, especially when heated. This is not fun at all, I value my glassware. Maybe even some HF escapes from the solutions, but I doubt whether these quantities are significant.

Ni(en)3(ClO4)2 is not very soluble in cold water, while being more soluble in hot water (not exceptionally, but clearly more than in cold water). This allows me to recrystallize this complex and get very pure and nice dry crystalline samples. This actually is the case for many perchlorates with bigger anions. I already made numerous perchlorates from different metal complexes, but also from protonated amines (e.g. NH2CH3, NH(CH3)2, (en), N(CH3)4(+)), and also from Cs(+) and Rb(+). The starting material is HClO4, or in some limited cases, NH4ClO4. Easiest and cheapest would be the freely soluble NaClO4, but in the EU you cannot buy that anymore legally, so for me no NaClO4

. I noticed that in the cold, the perchlorates are remarkably stable. When heated in a flame, however, they burn quickly or even deflagrate, if the cationic species contains organic parts. So, perchlorate is a nice alley in the quest for easily separated and purified complexes, but its use requires some care.

[Edited on 19-11-20 by woelen]