N_I
Harmless
Posts: 4
Registered: 1-9-2018
Member Is Offline
|
|
Zinc complex synthesis
Hello everyone,
I know this forum since a long time and even if I never created an account until now, I have found tons of useful discussions and suggestions on
various topics. I am currently struggling to synthesize a molecule and I thought someone here might now better than I do:
I am a trained organic chemist working in a French university lab and, in the course of an industrial project, I was asked to provide a coordination
complex (guess why I need some help…  ) but it seems I can’t get the thing
to work.
This compound is the zinc diisocyanate diamine Zn(NCO) 2 (NH3) 2, (cas : 122012-52-6) with only 2 types of preparation described in the literature :
First one (Industrial & Engineering Chemistry Research (2013), 52(32), 10977, DOI : 10.1021/ie400745x) start with zinc oxide and urea, heated to
150°C for 45 min (that should give cyanic acid and ammonia, among a lot of other things) but after a couple of attempts, changing temperature and
reaction time, XRD shows a mixture of at least an unknown product (not the one I want, if I trust the diffraction diagram published within the paper)
and a lot of unreacted ZnO. MS dectects a small peak with the right mass and containing a Zn atom but surrounded with loads of other unidentified
things (probably some of them being organic material from the urea decomposition).
Second one (Journal of Molecular Structure (1992), 271(1-2), 27, DOI: 10.1016/0022-2860(92)80207-X) starts from a Zn2+ salt (chloride here, nitrate or
perchlorate when other metals are used to generate the intermediate M(NCO)2) to which is added potassium cyanate (2 eq.) and 35% conc. ammonia.
Evaporation should give crystals after 2 days but the only thing I got to crystallize by slowly evaporating the water at 40°C, since nothing
precipitated on its own, was some potassium chloride  .
I tried to redo this reaction in alcoholic solvent (MeOH or EtOH) to try precipitating the KNO3 or KCl before adding excess ammonia (6eq., 7N in
methanol) but there is a major problem. Contrary to KSCN, KOCN is scarcely soluble in alcohol, even when heating to 60°C for 30min (I did the
reaction with 3g of KOCN in 50ml 95% EtOH, but most of it stays as a solid). I didn’t try to reflux or heat even higher in a closed system for fear
of thermal decomposition of the cyanate.
I did get more solids to crash out after 1h of reaction (ZnNO3 solution in EtOH added dropwise on the KOCN suspension at 45°C) that I filtered before
adding ammonia at r.t. to the resulting solution, but so far, only a very small amount of a fine white powder appeared (maybe just some more
precipitated K salts?).
Also, there are not much info about the characterization of this compound and, as an organic chemist, I have a hard time figuring if it will be highly
soluble in water or polar organic solvents. I just assume it should be soluble in concentrated ammonia as it’s used for recrystallization.
Does anyone have some ideas or suggestion on what could be done? Both for improving the existing procedures or to try another synthesis. I’m sure
there are super-skilled inorganic chemists on this forum. I worked on this for like a week and I can’t seem to find a way to succeed. I can provide
more information and the papers if needed but they are easy to find using a certain website  .
Thanks in advance!
|
|
|
DraconicAcid
International Hazard
Posts: 4278
Registered: 1-2-2013
Location: The tiniest college campus ever....
Member Is Offline
Mood: Semi-victorious.
|
|
I would try dissolving zinc hydroxide or zinc oxide in concentrated ammonia, then adding two equivalents of ammonium cyanate.
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.
|
|
|
N_I
Harmless
Posts: 4
Registered: 1-9-2018
Member Is Offline
|
|
Thanks for your answer! I considered using ammonium cyanate, as analogous complexes M(NCS)2(NH3)2 have been made with NH4SCN in acetonitrile or
acetone (can't find the reference atm).
The only problem is that it's not commercially available and should be prepared freshly. It seems the only reasonnable method would be an anion
metathesis with a slight excess of AgOCN, as it would push the equilibrium by precipitating insoluble AgCl in water:
AgOCN + NH4Cl -> AgCl(s) + NH4OCN
I will see if the reagent is available in the lab tomorrow.
On a side note, don't you think that dissolving the zinc salt in ammonia before adding the cyanate will prevent the complex formation? If a Zn NH3
complex is formed, I am not sure it will be possible to displace the ammonia ligands. I was wondering if it might be easier to preform the Zn(NCO)2
(maybe with the same reaction as above, using ZnCl2 or ZnSO4?), filter the solids and add excess ammonia.
|
|
|
DraconicAcid
International Hazard
Posts: 4278
Registered: 1-2-2013
Location: The tiniest college campus ever....
Member Is Offline
Mood: Semi-victorious.
|
|
As the solution evaporates, it may lose ammonia to give the diammine complex.
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.
|
|
|
N_I
Harmless
Posts: 4
Registered: 1-9-2018
Member Is Offline
|
|
Sadly, I didn't found any AgOCN in our stockroom so I can't try to synthesize NH4OCN. I will see if we can buy some later.
I tried to use Zn(ClO4)2 with KOCN to generate the zinc cyanate but it seems the KClO4 that is formed is still soluble enough in water and after
filtration of the precipitate and evaporation of the solution, I got much more product than expected.
As for the ZnO method, I wasn't able to dissolve it in ammonia (500mg in 15ml, it only forms a thick white suspension. I still added 2eq KOCN but
nothing changed after a couple of hours) and I don't have any zinc hydroxide at hand.
|
|
|
DraconicAcid
International Hazard
Posts: 4278
Registered: 1-2-2013
Location: The tiniest college campus ever....
Member Is Offline
Mood: Semi-victorious.
|
|
Making zinc hydroxide should be easy- mix any soluble zinc salt with a stoichiometric amount of dilute hydroxide, and allow it to settle (or
centrifuge it). Wash it a couple of times to get rid of the other ions, but don't bother trying to filter it- it will be too gelatinous.
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.
|
|
|
WGTR
National Hazard
Posts: 971
Registered: 29-9-2013
Location: Online
Member Is Offline
Mood: Outline
|
|
The first reference is a bit short on experimental details. Have you tried contacting the author for advice?
Anyway, the reaction is conducted under reflux, but the condenser temperature isn't specified. The assumption is that it is room temperature, but
maybe it is much colder.
ZnO, like other metal oxides, can respond sluggishly if they were previously calcined at high temperatures. The source of ZnO is not specified from
what I can tell. However, figure 1 mentions calcined ZnO. Also, I have dissolved 0.9g of calcined ZnO into 12.01g of NaOH in 50ml under heavy
stirring at room temperature over a period of hours, and everything dissolved. This was performed inside of a closed bottle to keep CO2
away.
Is it possible that the reaction is somewhat moisture-sensitive at these high temperatures, and the reactants need to be rather dry? Urea hydrolyses
readily in warm water.
For the second reference, ammonia has a high vapor pressure in such strong solutions. Perhaps the evaporation of ammonia from the solution was too
fast. Also, the reference implies that evaporation was commenced immediately after mixing the reactants, but perhaps leaving it in a sealed container
for a few days beforehand might make a difference.
|
|
|
N_I
Harmless
Posts: 4
Registered: 1-9-2018
Member Is Offline
|
|
| Quote: | | Making zinc hydroxide should be easy |
Yeah, and I think that is what I got in my last attempt with the second method (ZnCl2+KOCN+NH3) : I got some crystals but most of what precipitated
after 3 days was a white goo-like mass. I wonder if I should let the salts react for a much longer time before adding ammonia...
| Quote: | | Have you tried contacting the author for advice? |
I thought about it. I might try to send an email next week, but the paper is 5 years old and I wouldn't be surprised if they didn't even bother to
give me details about the procedure (which it is obviously lacking, as you noticed)
| Quote: | | The assumption is that it is room temperature, but maybe it is much colder. |
I don't really see what I should condense here appart from ammonia itself (but I guess it should react pretty fast at this temperature). I have
observed a fine white solid, soluble in water, sublimating on the air condenser and, from what I've found it might be some ammonium carbamate that is
also a decomposition product of urea with carbon dioxide. I ran the reaction with an without an argon atmosphere but it still doesn't work.
| Quote: | | ZnO, like other metal oxides, can respond sluggishly if they were previously calcined at high temperatures |
I guess it has something to do with the oxide but I only have one source at hand and there's no specification on it, appart from particule size.
Technical datasheets on Sigma don't seem to have much more information about the thermal treatment applied to what they sell either...
One thing that bugs me is that there are videos on youtube for the synthesis of calcium cyanate (for the formation of CaCN2, used for aminoguanidine
and tetrazoles synthesis) and absolutely no precaution is taken (usually done in a metal pot in air, without stirring and heated to high temp with a
gas burner), so making metal cyanates seems to be a quite easy thing and I don't think the reaction is very sensitive to moisture.
The only thing I have not tried is to mix and crush the reactant in a mortar like they do in the videos. I guess it is useful without stirring but for
my attempts, I used a stir bar and since the urea melts at 135°C, it forms a thick white liquid that should mix everything easily.
I will also try using zinc carbonate instead of the oxide as it was reported in an old patent from Ig farben for the synthesis of Zn(NCO)2.
| Quote: | | For the second reference, ammonia has a high vapor pressure in such strong solutions. Perhaps the evaporation of ammonia from the solution was too
fast |
Yes, that is what I was wondering, due to my observation of most product being possibly zinc hydroxide. If only the authors could write decent
experimental sections without half of the information missing...
I will keep you updated if I have more luck in the future.
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
I once reported on my successful experience with the reaction of the base [Cu(NH3)4(H2O)2](OH)2 with a magnesium salt:
[Cu(NH3)4(H2O)2](OH)2 + MgX2 --> [Cu(NH3)4(H2O)2]X2 + Mg(OH)2 (s)
I would suggest trying the same with corresponding zinc salt.
--------------------------------------------------------------------
I created the copper salt by the action of copper metal with ammonia followed by adding dilute H2O2, in that respective order, along with a touch of
sea salt (this is an electrolyte for the electrochemical reaction which I jump start in a microwave). The kinetics of Cu/NH3/H2O2 or O2
electrochemical cell is discussed here: https://onlinelibrary.wiley.com/doi/pdf/10.1002/bbpc.1963067... .
Caution: The Cu/NH3/O2 battery cell produces, in my opinion, electrons that can act on the NH3 in the presence of an oxygen source to form also a
nitrite side product, NH4NO2, as has been reported (see sources cited at http://www.sciencemadness.org/talk/viewthread.php?tid=81755#... ). The latter with a pH shift can rapidly decompose releasing nitrogen, which may
rupture closed vessels (I can provide an outline of this NO2- formation reaction for anyone interested). Assuming zinc follows a similar path as
copper, and given that Zn has a higher anodic index than copper, I would not be too surprised if there was a significant side product formation,
however, if the cell does not act to form any NH4NO2, the rarity of the formation of Zn(l) (unlike the cuprous complex, which is likely key to a
recycling reaction) may be a factor (as to more precisely how, please see my discussion in this SM thread http://www.sciencemadness.org/talk/viewthread.php?tid=87781#... ).
-------------------------------------------------------------------------
Here is a source citing various corrosive conditions for zinc metal (like temperature above 60 C, aqueous chloride, the presence of cathodic metal
like Pb or Sn, O2, NH4+ and in particular (NH4)2SO4 or (NH4)2SO3 or NH4NO3 or NH4NO2, pH acidic to neutral,...) suggesting to me possible
electrosynthesis (https://en.wikipedia.org/wiki/Electrosynthesis ) path to zinc salts (https://books.google.com/books?id=8C7pXhnqje4C&pg=PA526&... ).
See also my comments on attempting an electrosynthesis of ZnO/Zn(OH)2 based on a so called Bleach Battery replacing the aluminum metal with zinc
(which perhaps, as given to me as a free sample, may have been a zinc alloy) at http://www.sciencemadness.org/talk/viewthread.php?tid=75990#... .
[Edited on 10-9-2018 by AJKOER]
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
To quote N_l:
"compound is the zinc diisocyanate diamine Zn(NCO)2(NH3)2"
Actually, what if the compound is [Zn(NH3)2(H2O)4](NCO)2 (which I would label as zinc diamine tetraaqua diisocyanate) ?
One problem is that per my new thread (see http://www.sciencemadness.org/talk/viewthread.php?tid=91632 ) to quote:
"More particularly, research in physical chemistry journals report the formation of the Zn(NH3)3(H2O)3 complex (see, for example, ‘Zinc complexes of
water, hydroxide, and ammonia, by Douglas B. Kitchen and L. C. Allen, in J. Phys. Chem., 1989, 93 (20), pp 7265–7269, DOI: 10.1021/j100357a046 at https://pubs.acs.org/doi/10.1021/j100357a046 ). This specie, in more recent (2007) research, is apparently more stable than the
[Zn(NH3)2(H2O)4]2+ complex (see ‘Stability of different zinc(II)-diamine complexes in aqueous solution with respect to structure and dynamics: a
QM/MM MD study’ by Fatmi MQ, Hofer TS, Randolf BR and Rode BM, in J Phys Chem B. 2007 Jan 11;111(1):151-8, link: https://www.ncbi.nlm.nih.gov/pubmed/17201439 ). There also exist a stable [Zn(NH3)(H2O)5]2+ complex (see https://www.ncbi.nlm.nih.gov/pubmed/16633651 )."
So, technically, one problem is that [Zn(NH3)2(H2O)4]2+ complex is not particularly stable, and the actually compound (if exists) is either
[Zn(NH3)3(H2O)3]2+ or [Zn(NH3)(H2O)5]2+ complex in the presence of the NCO- anion.
|
|
|
WGTR
National Hazard
Posts: 971
Registered: 29-9-2013
Location: Online
Member Is Offline
Mood: Outline
|
|
Zinc oxide as the precursor of homogenous catalyst for synthesis of dialkyl carbonate from urea and alcohols
There is more experimental detail in this reference here:
Attachment: zhao2009.pdf (344kB)
This file has been downloaded 385 times
The catalyst is synthesized under reflux, but still inside of an autoclave, not open to atmosphere.
|
|
|
|
![[*]](images/xpblue/default_icon.gif){kind=icon}
