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Author: Subject: Oxalic acid extraction?
Glaudge
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[*] posted on 31-7-2007 at 15:58
Oxalic acid extraction?


ok, been here a while with very little posts, but i finally have a question. i searched a little for this and got nothing much, but i was wondering if there is an easy way of obtaining oxalic acid, like extracting it from rhubarb plant, or is there an easyerway of obtaining it (14 yrs old so cant order the stuff from places)

yes i am aware of its toxicity, especially with it making rhubarb plant leaves one of the most poisonous plant parts on earth.

if this is already been descusd you have my permission to slap me and link me there. :D
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Xenoid
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[*] posted on 31-7-2007 at 17:21


Oxalic acid (Ethanedioic acid) is actually available in a number of OTC preparations.

Think - Fe stain removal by complexing!

Eg. Stains on wood and clothing, radiator flush etc.

Check out paint and hardware shops, auto shops etc. Read the contents label carefully!

Xenoid
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[*] posted on 31-7-2007 at 17:44


Rhubarb is, by far, not one of the most poisonous plants! Look up coniine (hemlock), aconite (pseudoaconitine, monkshood), abrin (jequirity), ricin (castor), and physostigmine (calabar bean) to name but a few.

But! Your isolation will involve steeping of extracting in water and precipitating the oxalate as in insoluble salt (Ca2+, for example). The monovalent salts are water soluble and may also present a useful route.

Attached is an oldy by goody that might help.

Cheers,

O3

Attachment: Kohman 1939 oxalic acid_01.pdf (1.1MB)
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[*] posted on 31-7-2007 at 18:40


thanks alot.

my long-term plan was to get ammonium oxalate for explosives de-sensitising (pyromaniac i love things that burn or go boom, but with no intentions on bodily harm)
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[*] posted on 31-7-2007 at 18:40


Coincidentially, I've been gathering oxalis (sounds familiar, eh?) from the yard. It's a weed, so grows pretty well. Seems to dry out to just about nothing though; yield will be quite low.

My intended prep was:
1. Dry out the plants.
2. Chop or grind to flakes; seperate stems (unless they contain some too)
3. Make tea. Estimate oxalate content and use sufficient amount of water to leach all.
4. Add calcium chloride (if necessary) and cool to precipitate calcium oxalate.
5. Add hydrochloric acid to precipitate oxalic acid.
6. If additional purification is needed, repeat steps 4 (using CaCO3) and 5.
Various steps of fractional crystallization can be inserted in this process as well.

I'm not sure if an acidic or basic extraction would have any value.

Tim




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Glaudge
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[*] posted on 31-7-2007 at 18:42


oops, u posted about 5 seconds after i replied


edit: any uses for (cooh)2 (i hope thats the formula) other than the ammonium oxalate with explosives that would be interesting to a 14-15 yr old?( ;) )



[Edited on 31-7-2007 by Glaudge]
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Ozone
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[*] posted on 31-7-2007 at 19:06


For higher yields check out skunk cabbage (eastern), a very neat plant that generates its own heat and "the plant of the day" (this everyday plant turns out to be quite cool):

http://faculty.ucc.edu/biology-ombrello/POW/dumbcane.htm

I think both spinach and parsely are good sources (if less interesting plants) as well.

Cheers,

O3




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Xenoid
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[*] posted on 31-7-2007 at 20:18


Quote:
Originally posted by Glaudge
edit: any uses for (cooh)2 (i hope thats the formula) other than the ammonium oxalate with explosives that would be interesting to a 14-15 yr old?( ;) )
[Edited on 31-7-2007 by Glaudge]


Well, you could heat it at 150oC, Oxalic acid decomposes to Formic acid (Methanoic acid) + CO2

.....mmmm, the smell of crushed ants!

Xenoid
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[*] posted on 31-7-2007 at 21:32


Screw the rhubarb.

Go get a prepaid visa card and order what you want and lie about your age.




Not all chemicals are bad. Without chemicals such as hydrogen and oxygen, for example, there would be no way to make water, a vital ingredient in beer.
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[*] posted on 31-7-2007 at 23:13


Plants are a poor source, but it's good lab practice on something that's not going to kill you unless you actually go about eating it in largish amounts. It can be enjoyable to wring some pure compound from a pile of dead plants.

As already mentioned, wood bleach is one of the most common OTC sources, often kept near the paint strippers in hardware stores.

Oxalic acid is used as a reducing agent, taking vanadium from the 5 to the 4 oxidation state for example.

Recrystallised to good purity it can be used as a standard acid for titration, sulfamic acid is another OTC acid that can be used.
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[*] posted on 1-8-2007 at 00:37


Pure oxalic acid is OTC in hardware stores, generally right next to or near the HCl
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[*] posted on 1-8-2007 at 01:26


Oxalic acid could be also prepared with HNO3 and common sugar,

you just have to mix the acid, 65% is at least enough with the sugar in an erlenmeyer flask and then heat this mixture to the boiling point, but be careful, much NO2 escapes and the yields are low, just around 10%.

Not the best way to get Oxalic acid. So I also recommend you to buy this cheap and easy available chemical.
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[*] posted on 1-8-2007 at 01:59


I Like oxlaic aicd as a reducing agent. Mostly it is oxidized to Carbon dioxide/monoxide. It will reduce chlorates to chlorine dioxide in presense of dilute mineral acid. It will also reduce manganese dioxide in the presense of a strong acid to give the manganous salt of that respective acid. Alot of oxalic reductions require some heating to complete. I have done dry reduction of vanadium pentoxide with oxalic acid and came up with some nice blue tetraoxide:D Manganese dioxide can be found with pottery suppliers with vanadium pentoxide or obtained mixed with carbon in dry cell (nonalkaline) batteries.



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[*] posted on 1-8-2007 at 02:07


Quote:
I have done dry reduction of vanadium pentoxide with oxalic acid and came up with some nice blue tetraoxide:D

Hmmm.... sounds interesting. Did you use anhydrous oxalic acid, or the dihydrate? Up to what temperature did you need to heat the mix? How did you purify the blue VO2 (or V2O4) and how can you be sure that it is free of oxalic acid, and/or vanadium pentoxide?

This definitely sounds like something I want to try, but if you have experience with this, then I certainly want to use that.




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[*] posted on 1-8-2007 at 07:08


ok thanks, i'll look next time i'm at lowes
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[*] posted on 1-8-2007 at 08:14


woelen - Sidgwick says, without giving references

Quote:
(Vanadium dioxide) is usually made by heating the pentoxide with the trioxide, with oxalic acid, or with sulfur dioxide.


A slight excess of oxalic acid should hurt, and 95% alcohol will dissolve the excess.

Vanadium in the 4+ oxidation state does form complexes with oxalates, (NH4)2[VO(C2O4)2] 2H2O for example, that are blue and when cold do not test positive for oxalate. This suggests that if using a vanadate as the vanadium source, a large excess of oxalic acid is not a good itea, and that perhaps sticking to V2O5 is better.
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[*] posted on 1-8-2007 at 10:33


I tried to make vanadium (IV) oxide by means of heating with oxalic acid. The result was not satisfactory at all. I used the dihydrate, I only have that.

I did two experiments:
1) Take some V2O5 and a large excess amount of (COOH)2.2H2O. Grind both solids, such that a very fine yellow powder was obtained. Heat this yellow powder. A fairly vigorous reaction occurs. Soon, the material is dark blue, and a lot of noxious white smoke (I think sublimed oxalic acid) is produced. I continued heating until no more smoke was produced and then I heated a little more to be really sure that no oxalic acid remains. I finally obtained a very dark blue solid, which can easily be made into a fine powder.

2) Take some V2O2 and add a large excess amount of (COOH)2.2H2O and add some hydrochloric acid (30%). All of the solid dissolves (making funny crackling noise), and after quite some heating, I obtaiend a dark blue solution, really nice deep blue. This is an indication that I only have vanadium in the +4 oxidation state. Next, I continue heating. This first results in dense white fume, lateron real smoke is produced. When almost all liquid has evaporated, then first, a fairly light blue humid solid is produced. Still a lot of white smoke is produced. Again, heating was continued, until no more white smoke is produced. Also this experiment results in a very dark blue solid, just the same as in experiment (1). The chloride must have been driven off as HCl.

Now the bad part: When the dark blue solid is added to dilute hydrochloric acid, then a yellow/green solution is obtained. When the dark blue solid is added to concentrated hydrochloric acid, then a dark brown solution is obtained. After adding some sodium sulfite and gently heating, both liquids become beautifully blue. So, the black powder must contain quite a lot of vanadium in the +5 oxidation state. I think that this is due to oxidation by oxygen from the air.

So, simply heating the mix of oxalic acid and V2O5 in a beaker does not give the desired result.

@chloric1: Could you try adding some of your vanadium (IV) oxide to concentrated HCl. This is a sensitive test for vanadium in the +5 oxidation state. In conc. HCl, the very intensely colored brown VOCl3 is formed. Even in conc. HCl, vanadium in the +4 oxidation state does not form brown VOCl2. I did this countertest by adding some commercial VOSO4.xH2O to conc. HCl. So, your sample should dissolve with a bright blue color. Green, or even yellow/brown, means that it contains a lot of vanadium in the +5 oxidation state.

I'll have to repeat this experiment, but now with a protective atmosphere. How I am going to do that, I must still think about.




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[*] posted on 1-8-2007 at 12:25


Well, first off I want to mention that I performed this in 2003 so my memory might be a bit sketchy. What I do remember is that the mixture when heated first is a liquid slurry and is olive green. After further heating the white fumes evolve and a dry sky blue powder remains. This dissolves in dilute sulfuric acid to give a blue vanadyl solution. I did not do the chloride route but I was almost certain that I read that tetravalent vanadium forms green chloro complexes in HCl. Could be wrong.



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[*] posted on 1-8-2007 at 12:44


Sky blue powder? My resulting powder is very dark blue, almost black. I also indeed get a liquid slurry,which is green, and then blue. But on further heating, when driving off all smoke, it finally becomes black. Do you think I am heating too strong or too fast? I heat in a glass beaker, directly above a flame, so the material might become quite hot.

EDIT: If your powder is sky blue, then I doubt you have VO2 (or V2O4). According to literature it is black. Some books say it is very dark blue, but definitely not sky blue.

MSDS's say the same. An example:

[url]http://www.espimetals.com/msds's/vanadiumoxidev2o4.pdf[/url]

So, I think that I made VO2, but very impure material. You probably made some oxalato complex.

EDIT2: Can't get the URL correctly working. Please copy the stuff between [ url] and [ /url] in the address bar of a web browser.

[Edited on 1-8-07 by woelen]




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[*] posted on 1-8-2007 at 15:41


Ok I checked my references and they all list Vanadium dioxide as being blue/black or very dark blue. THe black material could be trivalent oxide which is formed when the pentoxide is heated in hydrogen or carbon monoxide, both of which exist in the exhaust of the propane flame. Liek I said before my memory might not be clear on this and it might have been dark blue. I do know for a fact that it dissolved in dilute sulfuric acid to give a clear blue solution. I used a crucible and it seemed like the only strong heating was decompising the green oxalate complex. Stirring was definately necessary which helped prevent lumping of final product. Supposedly the dioxide is aphoteric and should also dissolve in alkalies. The trivalent is decidingly basic and dissolves into acids with difficulty. I might have to do this again probably next week because I want to electrolyze vanadyl ions with a lead cathode in sulfate medium. I might add Ammonium sulfate so I may crystallize a more stable alum. I might pour naptha or mineral spirits on top the electrolyte to keep air out.



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[*] posted on 13-8-2007 at 18:25


hmm, i looked on a certian site, and said silver oxalate is a low power explosive (hmm, contrary to the NH4 ion's use)
that has similar behavior to black powder (burns under norm conditions, explodes in confinenent)
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[*] posted on 14-8-2007 at 01:14


When silver oxalate could explode, does the same applies to copper oxalate?
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[*] posted on 14-8-2007 at 05:53


Per, i`ve heard in some book which among with the sugar/HNO3 method there are a synthesis of oxalic acid involving NaOH and cellulose/wood , but the book doesnt give more info about..

someone know about it? if this is feasible at home? if yes, this would be very interesting because of the vey easy acquisition of substances..




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[*] posted on 14-8-2007 at 06:15


Heating a mixture of NaOH or KOH and sawdust in thin layers so as to have good exposure to air gives the corresponding oxalate. There's more to it than that, the heating pattern is important as incorrect heating leads to low yields heavily contaminated with coloured material.

Yes, it can be done on a smaller scale, a large cast iron skillet can be used for the fusion. Workup is a little trick, as it involves filtering strong lye solutions, but nothing more exotic than that. Yields aren't great, but if you can get the hydroxide cheaply it should be OK.

Best yields are on a large scale, a long channel that the fusion mixed moved down over several hours gave high yields. The commercial methods were able to recover some formate as well. They also were set up to recover the excess hydroxide, and regenerate it.

Here are some docs on it.


[Edited on 14-8-2007 by not_important]

Attachment: OxalicAcid.zip (845kB)
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[*] posted on 14-8-2007 at 15:33


Quote:
Originally posted by Ozone
Attached is an oldy by goody
Attachment: Kohman 1939 oxalic acid_01.pdf


Oxalic acid is a liver poison
and I've been eating vegetables all my life
H O L Y . . SMOKES ! WHO NEW !

I'm chowing down on beef only from now on.

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