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phlogiston
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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.
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"If a rocket goes up, who cares where it comes down, that's not my concern said Wernher von Braun" - Tom Lehrer |
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bfesser
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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.
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nezza
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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.
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woelen
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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!
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nezza
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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.
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nezza
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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)
This file has been downloaded 861 times
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nezza
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Finally pictures and a video of the copper complex. Lovely blue flashes and a green coloured flame.
Attachment: CopperVGA.mp4 (1.3MB)
This file has been downloaded 874 times
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woelen
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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]
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nezza
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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.
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woelen
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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]
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PHILOU Zrealone
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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...
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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bfesser
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Perhaps a stupid question, but do either of you plan to attempt syntheses of analogue Fe<sup>2+</sup> complexes?
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woelen
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@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.
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nezza
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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.
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bfesser
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woelen, I was thinking Schlenk techniques were necessary. I'm also looking forward to the pictures and video you promised of the
Cu<sup>2+</sup> complex.
Like I said, it was a stupid question. Thanks for the replies, though.
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woelen
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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?
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DraconicAcid
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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.
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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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.
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woelen
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As promised, a webpage about the copper(II) complex:
http://woelen.homescience.net/science/chem/exps/Ni_en_comple...
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bfesser
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Beautiful work, <strong>woelen</strong>. If I'm ever able to experiment again, this is the first synthesis I hope to do.
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PHILOU Zrealone
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| 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.
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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PHILOU Zrealone
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| Quote: Originally posted by PHILOU Zrealone | | 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 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]
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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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]
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PHILOU Zrealone
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| 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.
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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]
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
Super Administrator
Posts: 7977
Registered: 20-8-2005
Location: Netherlands
Member Is Offline
Mood: interested
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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]
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