Sciencemadness Discussion Board
Not logged in [Login ]
Go To Bottom

Printable Version  
 Pages:  1  
Author: Subject: Oxalate complexes
sasan
Hazard to Self
**




Posts: 92
Registered: 22-2-2014
Location: TEHRAN / IRAN
Member Is Offline

Mood: Radiative

cool.gif posted on 15-6-2014 at 01:14
Oxalate complexes



Hi guys,I know that oxalate complexes had discussed in madness so much,but they are not complete and discussed in different topics with different dates,so I decided to open new topic to discuss about oxalate complexes

potassium trisoxalatoferrate(III):Fe2(SO4)3 + 3 BaC2O4 + 3 K2C2O4 → 2 K3[Fe(C2O4)3] + 3 BaSO4 (mixing in hot water and stiring and filtering the barium sulfate and drying the complex)

second method:2FeC2O4.2H2O + H2C2O4 + H2O2(aq) + 3K2C2O4 ==) 2K3[Fe(C2O4)3]+...... avaporating and drying the
product

potassium tirsoxalatochromate(III):K2Cr2O7 + 7H2C2O4 + K2C2O4==) K3[Cr(C2O4)3] + 6CO2 +H2O precipitating by alcohol and cooling and drying

potassium bisoxalatochromate(III):K2Cr2O7 + 7H2C2O4 + ??H2O ==)2K[Cr(C2O4)2]·4H2O+ ??CO2 drying the bluish violet solution

potassium bis(oxalato)cuprate(II)dihydrate:2K2C2O4 + CuSO4.5H2O ==)K[Cu(C2O4)2].2H2O + ......mixing stochiometric amounts of copper sulfate and potassium oxalate solutions at 60 C.precipitated by cooling,it has easiest pocedure in the oxalate stuff.
this is the picture of copper complex,very beautiful blue color

There is aluminium complex too,but it is boring and white to make that

Now my questions:is there other complex of oxalate with transition metals?and have you guys ever work on oxalate complexes?

I know there is titanium(III) complex too,but nothing on googling.

DSC_0244.jpg - 393kB DSC_0244.jpg - 393kB

[Edited on 15-6-2014 by sasan]
View user's profile View All Posts By User
Brain&Force
Hazard to Lanthanides
*****




Posts: 1299
Registered: 13-11-2013
Location: UW-Madison
Member Is Offline

Mood: Incommensurately modulated

[*] posted on 15-6-2014 at 11:10


Nice work sasan! Have you checked the tris(oxalato)ferrate complex for fluorescence?

Are there vanadium, manganese, nickel, and cobalt complexes as well?




At the end of the day, simulating atoms doesn't beat working with the real things...
View user's profile View All Posts By User
unionised
International Hazard
*****




Posts: 4939
Registered: 1-11-2003
Location: UK
Member Is Offline

Mood: No Mood

[*] posted on 15-6-2014 at 11:13


Does this help?
http://scripts.iucr.org/cgi-bin/paper?S0567740877009819
View user's profile View All Posts By User
blogfast25
Thought-provoking Teacher
*****




Posts: 10346
Registered: 3-2-2008
Location: Old Blighty
Member Is Offline

Mood: No Mood

[*] posted on 15-6-2014 at 13:11


Quote: Originally posted by Brain&Force  
Nice work sasan! Have you checked the tris(oxalato)ferrate complex for fluorescence?



I made that complex and it fluoresces mildly in visible light. Didn't check in UV.

It's also a bit unusual to have a green ferric compound.

[Edited on 15-6-2014 by blogfast25]




View user's profile Visit user's homepage View All Posts By User
numos
National Hazard
****




Posts: 261
Registered: 22-2-2014
Location: Pasadena
Member Is Offline

Mood: No Mood

[*] posted on 15-6-2014 at 15:45


Ooo... something new to try, do you know if sodium oxalate could be used instead of potassium? Would Na3[Fe(C2O4)3] have the same properties?

Otherwise... I have sodium oxalate, but I have no Potassium bases to make potassium oxalate out of.

Also, this isn't just directed at you, but when you use H2O2 and other aqueous substances (eg. Ammonia) , can you give the conc.? 3%? 30%? 95%?




View user's profile View All Posts By User
Justin Blaise
Hazard to Self
**




Posts: 82
Registered: 5-10-2011
Location: Parts Unknown
Member Is Offline

Mood: No Mood

[*] posted on 15-6-2014 at 19:31


I once neutralized a small amount of CoCO3 with aqueous oxalic acid. I boiled the mixture until it stopped producing CO2 and I was left with a salmon colored precipitate of what I believed to be cobalt (II) oxalate. The color was not much different than that of the original carbonate, so it was a little disappointing.
View user's profile View All Posts By User
numos
National Hazard
****




Posts: 261
Registered: 22-2-2014
Location: Pasadena
Member Is Offline

Mood: No Mood

[*] posted on 15-6-2014 at 22:19


I tried the Iron and Copper complexes with Sodium oxalate, and it definitely does work, the copper looks identical to the picture.



View user's profile View All Posts By User
sasan
Hazard to Self
**




Posts: 92
Registered: 22-2-2014
Location: TEHRAN / IRAN
Member Is Offline

Mood: Radiative

[*] posted on 15-6-2014 at 22:35



B&F oxalatoferrates have fluorescence property,but not strong and beautiful like sodium flourscein and pyranine and they will decomposes under UV light:(

There is manganese oxalate complex too and it is more pretty(deep red maroon color)but the production is difficult and highly depends on the concentration and temprature.I go for this complex today but I dont think if it will work or not

I tried nickel complex too,but I didn't get a good result,maybe I used wrong amounts.I will retry this later

I don't know about the vanadium complex,but I think there is cobalt (III) complex exists(maybe treating cobalt salts with pottasium oxalate and hydrogen peroxide solution with stochiometric amounts.)

justin blaise use my procedure for cobalt complex,right now I don't have enough cobalt salt,make potassium oxalate by treating your oxalic acid with potassium carbonate(not potassium hydroxide)and make cobalt salt with your cobalt carbonate.

numos I have sodium oxalate too,but it is very less soluble comparing to potassium oxalate,and for making oxalate complexes you need concentrated solution and you can't make concentrated solution of oxalate salt with your sodium oxalate,treat your sodium oxalate with hydrochloric acid to obtain oxalic acid precipitate(I know oxalic acid is soluble in water,but I think it is insoluble in HCl(aq) and in sodium chloride solution),then combine your oxalic acid with potassium carbonate in right amounts to gain potassium oxalate solution and dry it to form K2C2O4.2H2O

about the concentration of the hydrogen peroxide,for every reaction I use excess hydrogen peroxide solution with concentration of ~30%

I hope I answered your questions




View user's profile View All Posts By User
papaya
International Hazard
*****




Posts: 594
Registered: 4-4-2013
Member Is Offline

Mood: reactive

[*] posted on 16-6-2014 at 01:24


I must say that lately I prepared copper (II) oxalate (which must be regarded as a complex), though unwillingly, with a rather unexpected route and without using any oxalic acid.
Take copper metal, dissolve in 50% nitric acid - there must be enough excess of acid left in solution. Then add ethylene glycol and let it stay for a day. I tried this on a test tube scale - don't scale up, it may be not safe, since after HOURS an exothermic runaway with NOx evolution occurs after which light blue precipitate starts to form (supposedly copper oxalate dihydrate). After 24 hours I filtered that, washed with water and dried at RT. It's a beautiful light blue fine powder, the color seems very interesting under daylight - must be a little flourescence.
But the most striking thing to me was how it decomposes when heated on the foil - RED fumes are evolved with NO residue left at all - the fume consists of probable nano-sized copper/copper oxide particles(not NO2).. If the fumes hit the flame - it gets colored intense green! Just try!

[Edited on 16-6-2014 by papaya]
View user's profile View All Posts By User
woelen
Super Administrator
*********




Posts: 7764
Registered: 20-8-2005
Location: Netherlands
Member Is Offline

Mood: interested

[*] posted on 16-6-2014 at 01:41


Some oxalates and oxalato complexes are explosive!

Copper(II) is borderline. It very easily loses CO2 and leaves behind a very fine powder of copper. Ferrous oxalate is somewhat more difficult to decompose, but it can be used to make pyrophoric iron powder. Silver oxalate explodes on heating, expelling CO2 with great violence and leaving fine silver powder behind.




The art of wondering makes life worth living...
Want to wonder? Look at https://woelen.homescience.net
View user's profile Visit user's homepage View All Posts By User
papaya
International Hazard
*****




Posts: 594
Registered: 4-4-2013
Member Is Offline

Mood: reactive

[*] posted on 16-6-2014 at 01:47


Quote: Originally posted by woelen  
Some oxalates and oxalato complexes are explosive!

Copper(II) is borderline. It very easily loses CO2 and leaves behind a very fine powder of copper. Ferrous oxalate is somewhat more difficult to decompose, but it can be used to make pyrophoric iron powder. Silver oxalate explodes on heating, expelling CO2 with great violence and leaving fine silver powder behind.


Sure, but copper one, at least the dihydrate didn't show any energetic properties, even when heating confined in a foil it decomposed the same way.
View user's profile View All Posts By User
woelen
Super Administrator
*********




Posts: 7764
Registered: 20-8-2005
Location: Netherlands
Member Is Offline

Mood: interested

[*] posted on 16-6-2014 at 02:08


As I said, copper(II) oxalate is borderline. It can decompose quickly, but not explosively. I can imagine that if it is really confined, that it can lead to an explosion, due to pressure buildup. Confining it in a little ball of Al-foil is not sufficient. The produced gas is not formed quickly enough to rip the foil apart, it quickly escapes through the folds in the foil.




The art of wondering makes life worth living...
Want to wonder? Look at https://woelen.homescience.net
View user's profile Visit user's homepage View All Posts By User
bismuthate
International Hazard
*****




Posts: 803
Registered: 28-9-2013
Location: the island of stability
Member Is Offline

Mood: self reacting

[*] posted on 16-6-2014 at 02:49


I just made some and found that it decomposes almost instantly under flame. I've been wondering (and experimenting with to find out) if complexes would have energetic properties.
Would an ethylenediamine complex be possible? That might be explosive.

[Edited on 16-6-2014 by bismuthate]




I'm not a liar, I'm just an enthusiastic celebrant of opposite day.
I post pictures of chemistry on instagram as bismuthate. http://iconosquare.com/bismuthate
or this viewer if you don't have an instagram (it sucks though) http://web.stagram.com/n/bismuthate
View user's profile Visit user's homepage View All Posts By User
PHILOU Zrealone
International Hazard
*****




Posts: 2880
Registered: 20-5-2002
Location: Brussel
Member Is Offline

Mood: Bis-diazo-dinitro-hydroquinonic

[*] posted on 16-6-2014 at 05:23


Quote: Originally posted by woelen  
Some oxalates and oxalato complexes are explosive!

Copper(II) is borderline. It very easily loses CO2 and leaves behind a very fine powder of copper. Ferrous oxalate is somewhat more difficult to decompose, but it can be used to make pyrophoric iron powder. Silver oxalate explodes on heating, expelling CO2 with great violence and leaving fine silver powder behind.

Mercury oxalate is even better for the explosive purpose...at least it crushes more sand in sand test than silver oxalate.




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)
View user's profile View All Posts By User
PHILOU Zrealone
International Hazard
*****




Posts: 2880
Registered: 20-5-2002
Location: Brussel
Member Is Offline

Mood: Bis-diazo-dinitro-hydroquinonic

[*] posted on 16-6-2014 at 05:38


Quote: Originally posted by bismuthate  
I just made some and found that it decomposes almost instantly under flame. I've been wondering (and experimenting with to find out) if complexes would have energetic properties.
Would an ethylenediamine complex be possible? That might be explosive.

[Edited on 16-6-2014 by bismuthate]

Oxalate is a weak explosophoric group. The driving force is the oxydoredox potential of:
-the oxydation of the oxalate dianion (reducer) into 2 gaseous CO2 molecules:
(-)O2C-CO2(-) --> 2 CO2(g) + 2e(-)
-the reduction of the oxydising metalic cation into metal
Cu(+) + 1e(-) --> Cu
Cu(2+) +2e(-) --> Cu
Ag(+) + 1e(-) --> Ag
Hg(+) + 1e(-) --> Hg
Hg(2+) + 2e(-) --> Hg

The heat of reaction (explosion) is not that much with comparison with other energetic materials...mostly because the carbon in oxalate is already almost fully oxydised.

Thus adding highly defficient OB complexants (like ethylene diamine) will temper the already poor explosive properties.




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)
View user's profile View All Posts By User
sasan
Hazard to Self
**




Posts: 92
Registered: 22-2-2014
Location: TEHRAN / IRAN
Member Is Offline

Mood: Radiative

[*] posted on 16-6-2014 at 06:43



please don't go far from the topic,it is about the oxalate complexes not explosives or etc;)

papaya" potassium oxalato copprate is not fluorescent,I have UV tube I tried that under UV,but is has no fluorescence property,it is just very shiny and crystaline under lights,the ferrioxalate is somewhat fluorescent.

The cobalt complex is too unstable to make by amateur chemists,It is very photosensitive and very airsensitive(unlike cobalt (II) hydroxides that will oxidizes to (III) in presence of air:o)and reduced to cobalt (II) and destroys the potassium trisoxalatocobaltate(III)

The manganese complex is similar to cobalt,its complex exists just for Mn(3+),and in presence of air will reduced to Mn(2+)(Unlike manganese (II) hydroxide that will oxidize to Manganese(III)oxide:o) and destroys the potassium trisoxalatomanganate(III) compelx,but it is more easier to make manganese complex comparing too cobalt

any ideas?

Woelen,in your website I saw the potassium bis&trisoxalatochromates(III) pictures,how did you make that?I tried my procedure but I had a very low yield.


View user's profile View All Posts By User
Praxichys
International Hazard
*****




Posts: 1063
Registered: 31-7-2013
Location: Detroit, Michigan, USA
Member Is Offline

Mood: Coprecipitated

[*] posted on 16-6-2014 at 07:53


Anyone try tetraaminecopper II oxalate?

Maybe take the same route as the persulfate derivative.... Mixing an excess of ammonium oxalate (4.45g/100ml @20C) solution with aq. ammonia, add copper sulfate solution, hopefully precipitating the complex instead of just copper(II) oxalate (~2x10-10g/100ml @20C). I can't find solubility data on ammonium bioxalate so I do not know for sure if this is going to work. I'll try it this week if nobody else gets around to it.




View user's profile Visit user's homepage View All Posts By User
bismuthate
International Hazard
*****




Posts: 803
Registered: 28-9-2013
Location: the island of stability
Member Is Offline

Mood: self reacting

[*] posted on 16-6-2014 at 11:43


Quote: Originally posted by Praxichys  
Anyone try tetraaminecopper II oxalate?

No but I plan to soon. I'll post the results when done.




I'm not a liar, I'm just an enthusiastic celebrant of opposite day.
I post pictures of chemistry on instagram as bismuthate. http://iconosquare.com/bismuthate
or this viewer if you don't have an instagram (it sucks though) http://web.stagram.com/n/bismuthate
View user's profile Visit user's homepage View All Posts By User
Praxichys
International Hazard
*****




Posts: 1063
Registered: 31-7-2013
Location: Detroit, Michigan, USA
Member Is Offline

Mood: Coprecipitated

[*] posted on 16-6-2014 at 15:21


I just tried tetraaminecopper II oxalate with no success so far.

1. I mixed up a CuSO4 solution with some oxalic acid solution and immediately got a milky, pale blue precipitate which I presume is copper oxalate. Adding 10% ammonia solution slowly darkened the solution until it was deep blue, exactly like tetraaminecopper II sulfate. Mixing with 50% MeOH resulted in a precipitation of some deep blue crystals. I scaled it up and I am letting it precipitate right now to investigate. My gut tells me that this is just tetraaminecopper II sulfate. It looks exactly like it.

2. Oxalic acid was dissolved with aqueous ammonia until the solution remained basic, indicating an excess of ammonia. A solution of copper II sulfate was added, giving the same result as solution 1. A MeOH precipitation was not tried as the solution was fairly dilute at this time and appeared identical to solution 1.

3. Some basic copper II carbonate was added as a powder to an oxalic acid solution, producing some bubbles and a pale teal precipitate, appearing that not all of the carbonate had reacted. Extra oxalic acid was added to be sure all copper carbonate was consumed. The color did not change. It is speculated that the carbonate was encapsulated by the formation of oxalate, or perhaps teal is the actual color of copper II oxalate. Addition of aqueous ammonia to the solution produced a medium-blue cloudy mixture which will need scaling up and filtering for examination.




View user's profile Visit user's homepage View All Posts By User
thesmug
National Hazard
****




Posts: 370
Registered: 17-1-2014
Location: Chicago, Il (USA)
Member Is Offline

Mood: No Mood

[*] posted on 16-6-2014 at 21:27


I think you are right about getting the sulfate ions back. You might try eliminating the sulfate before attempting to complex with ammonia. I'm unsure of how you might do this.



Good eyes
View user's profile View All Posts By User
bismuthate
International Hazard
*****




Posts: 803
Registered: 28-9-2013
Location: the island of stability
Member Is Offline

Mood: self reacting

[*] posted on 17-6-2014 at 02:26


Barium nitrate?
Or you could just filter it and then try to dissolve it in ammonia.




I'm not a liar, I'm just an enthusiastic celebrant of opposite day.
I post pictures of chemistry on instagram as bismuthate. http://iconosquare.com/bismuthate
or this viewer if you don't have an instagram (it sucks though) http://web.stagram.com/n/bismuthate
View user's profile Visit user's homepage View All Posts By User
Praxichys
International Hazard
*****




Posts: 1063
Registered: 31-7-2013
Location: Detroit, Michigan, USA
Member Is Offline

Mood: Coprecipitated

[*] posted on 17-6-2014 at 04:34


This morning yielded a tiny but recoverable amount of deep blue precipitate which held its color on drying. The sample was heated with decomposition releasing ammonia first and remaining a pale blue powder reminiscent of anhydrous copper II sulfate. However, further heating to ca. 400°C yielded some tiny particles of what appeared to be copper metal, indicating that the precipitate might be some amino-complexed double salt of copper II sulfate and oxalate.

I think the most promising method might be to make up a pile of clean and dry copper oxalate dihydrate, then gas a thin water slurry of it with ammonia with vigorous stirring until everything is dissolved, followed by crashing out into MeOH or EtOH. This will take considerable experimentation to determine the correct solubility ratios. I'm going to a concert tonight so I will not be able to try this until tomorrow.




View user's profile Visit user's homepage View All Posts By User
DraconicAcid
International Hazard
*****




Posts: 3967
Registered: 1-2-2013
Location: The tiniest college campus ever....
Member Is Offline

Mood: Semi-victorious.

[*] posted on 17-6-2014 at 08:25


Quote: Originally posted by thesmug  
I think you are right about getting the sulfate ions back. You might try eliminating the sulfate before attempting to complex with ammonia. I'm unsure of how you might do this.

Just add enough oxalate to precipitate the copper, filter out the copper(II) oxalate, then dissolve in aqueous ammonia.




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.
View user's profile View All Posts By User
papaya
International Hazard
*****




Posts: 594
Registered: 4-4-2013
Member Is Offline

Mood: reactive

[*] posted on 17-6-2014 at 09:54


Quote: Originally posted by DraconicAcid  
Quote: Originally posted by thesmug  
I think you are right about getting the sulfate ions back. You might try eliminating the sulfate before attempting to complex with ammonia. I'm unsure of how you might do this.

Just add enough oxalate to precipitate the copper, filter out the copper(II) oxalate, then dissolve in aqueous ammonia.


What is the structure of the supposed complex, any data available?
View user's profile View All Posts By User
Justin Blaise
Hazard to Self
**




Posts: 82
Registered: 5-10-2011
Location: Parts Unknown
Member Is Offline

Mood: No Mood

[*] posted on 17-6-2014 at 19:44


Found some info on preparing oxalate complexes of some trivalent metals in Inorganic Syntheses Volume 1, page 35.

Although more facile methods to make the chromium and iron complexes exist, the preparation of the cobalt complex seems interesting. If I can procure some lead dioxide I'll be sure to give it a shot.

"C . POTASSIUM TRIOXALATOCOBALTIATE
K3[Co(C2O4)3].3H2O
H2C2O4 + CoCO3 -> CoC2O4 + H2O + CO2
2CoC2O4 + 4K2C2O4 + PbO2 + 4HC2H3O2 ->
2K3[Co(C2O4)3] + 2KC2H3O2 + Pb(C2H3O2)2 + 2H2O

Twenty-three and eight-tenths grams (0.2 mol) of cobalt
carbonate is dissolved in a solution of 25.2 g. of oxalic
acid (H2C204*2H20) and 73.7 g. of potassium oxalate
(K2C2O4-H20) in 500 cc. of hot water. When the solution
has cooled to 40°C., while it is vigorously stirred, 23.9 g.
of lead dioxide (see synthesis 16) is added slowly, followed
by 25 ml. of glacial acetic acid added a drop at a time.
The stirring is continued for an hour, during which time
the color changes from red to deep green. After the
unused lead dioxide is filtered out, the trioxalatocobaltiate is
precipitated by the addition of 500 ml. of alcohol. The
material appears as emerald-green needles, which are
sensitive to both light and heat; The yield is 70 g. (71 per
cent). "




Attachment: inorganic-synthesis volume 1.pdf (7.4MB)
This file has been downloaded 1303 times
View user's profile View All Posts By User
 Pages:  1  

  Go To Top