zephler1
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So much oxalic acid - what to do with it?
I've often read that it is used in illicit labs, but for what? 2 carbons, two of the same functional groups do not look that attractive...or am I
missing something? Is it just used for formic acid production?
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TheChemiKid
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If you want to, you should make some iron oxalate, then decompose to pyrophoric Iron.
http://www.youtube.com/watch?v=_2HHuUMkg58
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Zephyr
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Here's a thread on a viable method of making sulfuric acid for oxalic acid.
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DraconicAcid
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Make transition metal complexes, such as K3[Fe(C2O4)3].
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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bfesser
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Thread Moved
14-1-2014 at 06:28 |
TheChemiKid
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You could make formic acid, that is a useful chemical.
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AJKOER
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When working with Oxalic acid be aware of possible undesirable reactions between products when employing an excess of H2C2O4. For example, the action
of Oxalic acid on an aqueous chlorate can form Chloric acid and then ClO2. Here is extract from a prior thread (see http://www.sciencemadness.org/talk/viewthread.php?tid=20109 )
Quote: Originally posted by chemretd  | | A "safe" way of generating chlorine dioxide, mixed with an equal volume of carbon dioxide was published (I can't remember where) as warming a mixture
of potassium chlorate and oxalic acid on a water- bath. In my spotty youth (MANY years ago), I tried this with sodium chlorate. It generated a green
gas alright, but with a very alarming LOUD crackling noise. I suspect that this safe method is a serious accident waiting to happen.
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Here is the reaction equations and source for Chemretd comment:
2 NaClO3 + H2C2O4 --> 2 HClO3 + Na2C2O4 (s)
and with an excess of Oxalic acid:
2 HClO3 + H2C2O4 --> 2 ClO2 + 2 H2O + 2 CO2
so the net reaction is:
2 NaClO3 + 2 H2C2O4 --> 2 ClO2 (g) + 2 H2O + 2 CO2 (g) + Na2C2O4 (s)
See http://books.google.com/books?id=6wUmteTIc18C&pg=PA334&a...
as was reported by Chemretd. This reaction is surprising to some and may precipitate an accident.
Another example of this possible dual role of Oxalic acid was reported by Watt's (see page 649 http://books.google.com/books?pg=PA649&lpg=PA649&sig... ) as the reduction of HOCl by Oxalic acid. So the addition of very dilute Oxalic
acid (and not concentrated H2C2O4 which can produce a very violent reaction with a hypochlorite and an explosion with a dry hypochlorite) to even pure
NaOCl would first form HOCl, which could subsequently reacts with the excess H2C2O4 to form HCl and CO2 (and then some Cl2). My take on the reaction
sequence:
H2C2O4 + 2 NaOCl + xH2O --> 2 HOCl + xH2O + Na2C2O4 (s)
2 HOCl + 2 H2C2O4 --> 2 HCl + 4 CO2 (g) + 2 H2O
2 HCl + 2 HOCl <----> 2 Cl2 + 2 H2O
Yet another example occurs when trying to create highly concentrated H2SO4 by the action of Oxalic acid on a sulfate. Using an excess of H2C2O4 to
drive the reaction together with heating indeed results in a very high strength of sulfuric acid. However, at a certain point, the excess H2C2O4 is
violently (and usually unexpectedly) decomposed (into CO2, CO and water vapor, see http://escholarship.org/uc/item/1s96t1nf#page-5 ) by the now newly formed concentrated H2SO4. In other words, a large violent ejection of highly
corrosive acid occurs (see comments by Formatik at http://www.sciencemadness.org/talk/viewthread.php?tid=18963&... ).
----------------------------------------------------------------------------------------------------
Also be wary of some of the properties of special oxalates, like Silver Oxalate (see http://en.wikipedia.org/wiki/Silver_oxalate ), in conjunction with the anhydrous Oxalic acid. If one where to sprinkle a very small amount of Ag
powder on a much larger amount of moist H2C2O4, by design, one could form a minute amount of a heat, friction and shock sensitive explosive (namely,
Silver Oxalate). Now, as the amount formed is small, one would think no major risk. Wrong! Assuming you are not working with the dihydrate acid
(H2C2O4.2H2O), relatively unknown to many, Oxalic acid can be made to detonate. See "Bretherick’s Handbook of Reactive Chemical Hazards", Sixth
Edition, Volume 1, page 269, link: eng.monash.edu/materials/assets/documents/resources/ohs/bretherick-vol1.pdf to quote:
"CHETAH, 1990, 184
Surprisingly, even the dihydrate is apparently detonated by a 50 g tetryl booster,
the anhydrous acid is thermally less stable and thus probably more sensitive. There
is, however, no history of explosion."
That is, H2C2O4 can function in the capacity of a secondary explosive, while the dihydrate form of the acid may be a little more resistant. Bottom
line, a little Ag on a large amount of H2C2O4 could precipitate a big event, especially if heated in a closed container, or accidentally shocked.
Apparently also, H2C2O4 and Urea heated in a close vessel can violently explode, so something else to avoid (same reference, page 270).
Otherwise, enjoy your Oxalic acid taking appropriate safety precautions given its toxic nature (www.setonresourcecenter.com/msdshazcom/htdocs//MSDS/S/sunbur... ).
[Edited on 15-1-2014 by AJKOER]
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TheChemiKid
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You could try to sell some of it to people who may be interested. 
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papaya
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| Quote: Originally posted by AJKOER |
Also be wary of some of the properties of special oxalates, like Silver Oxalate (see http://en.wikipedia.org/wiki/Silver_oxalate ), in conjunction with the anhydrous Oxalic acid. If one where to sprinkle a very small amount of Ag
powder on a much larger amount of moist H2C2O4, by design, one could form a minute amount of a heat, friction and shock sensitive explosive (namely,
Silver Oxalate). Now, as the amount formed is small, one would think no major risk. Wrong! Assuming you are not working with the dihydrate acid
(H2C2O4.2H2O), relatively unknown to many, Oxalic acid can be made to detonate. See "Bretherick’s Handbook of Reactive Chemical Hazards", Sixth
Edition, Volume 1, page 269, link: eng.monash.edu/materials/assets/documents/resources/ohs/bretherick-vol1.pdf to quote:
"CHETAH, 1990, 184
Surprisingly, even the dihydrate is apparently detonated by a 50 g tetryl booster,
the anhydrous acid is thermally less stable and thus probably more sensitive. There
is, however, no history of explosion."
That is, H2C2O4 can function in the capacity of a secondary explosive, while the dihydrate form of the acid may be a little more resistant. Bottom
line, a little Ag on a large amount of H2C2O4 could precipitate a big event, especially if heated in a closed container, or accidentally shocked.
Apparently also, H2C2O4 and Urea heated in a close vessel can violently explode, so something else to avoid (same reference, page 270).
Otherwise, enjoy your Oxalic acid taking appropriate safety precautions given its toxic nature (www.setonresourcecenter.com/msdshazcom/htdocs//MSDS/S/sunbur... ).
[Edited on 15-1-2014 by AJKOER] |
Very interesting, I've heard about silver oxalate, but not the acid itself. I guess anhydrous acid can be made by heating of dihydrate ? If oxalic
acid is explosive itself (and sensitive), also it lacks of some oxygen, it must be a good sensitizer for ammonium nitrate, isn't it?
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woelen
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I think that these explosive properties of oxalic acid are only interesting from a theoretical point of view. If real practical explosives could be
made with oxalic acid, then we would have heard of them and then there would be people around who tried them. I know of the silver oxalate (and
similarly, copper(I) oxalate may also have explosive properties).
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AJKOER
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I agree that Oxalic acid may be a poor example of a secondary explosive (barring further research and those willing to field test it). However, if
there is a barrel of H2C2O4 in a room on fire, containing other energetically unstable compounds, I would not dismiss the possibility of a
potentially more serious incident.
A point of correction in my scenario, apparently Oxalic acid does not directly attack Silver (or Gold, for that matter) per this source "A system of
chemistry: in four volumes, Volume 2, page 487 by Thomas Thomson and Thomas Cooper link to a Google book: http://books.google.com/books?id=mv1YAAAAYAAJ&pg=PA487&a... . However, it is claimed to dissolve small portions of Ag2O and, I suspect, act
similarly on Silver tarnish (Ag2S). As such, this may degrade the ability of Ag2C2O4 to form accidentally in significant quantities (and to act,
therein, as a primer for the secondary, dry Oxalic acid) by the direct interaction of Silver and the acid. Interestingly, this tends to supports
Woelen's comment on why we haven't heard of it.
------------------------------------------------------
Another point I should add, my supply of Oxalic acid degraded over the summer months. It was stored in a basement and I would be surprised if the
temperature reached much over 75 F. Impurities in my purchased H2C2O4.2H2O my also proved to be its undoing. Bottom line, buying large quantities at a
single time may not be wise. By all means, keep it cool and dry.
------------------------------------------------------
[EDIT] Further research, see "Explosives and Their Power" by Marcellin Berthelot, page 364, at http://books.google.com/books?id=mXFBAAAAIAAJ&pg=PA364&a... which discusses the explosive nature of metallic oxalate, including Oxalic acid
itself, from the perspective of their heat of formation. Bottom line, the equation:
M2C2O4 --> 2 CO2 + M2
must be exothermic. In the case of H2C2O4, in a solid state, it is not. However, when in a gaseous state, the equation above is marginally exothermic
(meaning potentially explosive). As such, H2C2O4 is expected to be a very low yield secondary explosive. Values for Copper oxalate and Silver oxalate
are also given and discussed.
[Edited on 16-1-2014 by AJKOER]
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ScienceSquirrel
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I have a 500g container of oxalic acid dihydrate in my chemical cupboard.
It has been there for a few years at least; stored in the dark at 5 -25 C and it is still a white, free flowing crystalline solid.
I suspect that making it explode requires extreme provocation.
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blogfast25
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| Quote: Originally posted by AJKOER | [EDIT] Further research, see "Explosives and Their Power" by Marcellin Berthelot, page 364, at http://books.google.com/books?id=mXFBAAAAIAAJ&pg=PA364&a... which discusses the explosive nature of metallic oxalate, including Oxalic acid
itself, from the perspective of their heat of formation. Bottom line, the equation:
M2C2O4 --> 2 CO2 + M2
must be exothermic. In the case of H2C2O4, in a solid state, it is not. However, when in a gaseous state, the equation above is marginally exothermic
(meaning potentially explosive). |
Self contradictory twaddle, most of it at least.
Firstly, Wolfram alpha cites the Heat of Formation for gaseous and solid oxalic acid as – 731.8 and – 829.9 kJ/mole respectively.
Secondly, conversion of metal oxalates by pyrolysis is likely to lead to metal oxides because these have almost always strongly negative heats of
formation. It’s true that in some case metal can be obtained, in others mixtures of metals and oxides. See e.g.:
http://researchspace.ukzn.ac.za/xmlui/handle/10413/3288
“M2C2O4 --> 2 CO2 + M2
must be exothermic.”
Prove it.
M2???
[Edited on 18-1-2014 by blogfast25]
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AJKOER
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Blogfast:
The author's cited general decomposition reaction does not, as written, apply to Oxalic acid, more generally to metallic oxalates of heavy metals (the
topic of that chapter).
Per this source "Theoretical Study of Thermal Decomposition Mechanism of Oxalic Acid", J. Phys. Chem. A, 1997, 101 (14), pp 2702–2708, link: http://pubs.acs.org/doi/abs/10.1021/jp9638191 the products for Oxalic acid are actually H2O, CO and CO2:
H2C2O4 ---> H2O + CO2 + CO
But, relative to the possible detonation of H2C2O4 itself, the reference also states to quote:
"The unimolecular formation of CO2, CO and H2O from oxalic acid via a concerted transition state has an activation barrier of only 42 kcal/mol,
indicating it is a more favorable unimolecular decomposition channel."
Now, as you noted, Wolfram alpha cites the Heat of Formation for gaseous and solid oxalic acid as – 731.8 and – 829.9 kJ/mole respectively, so I
attribute 98.1 kJ/mole gain from moving the Oxalic acid from a solid to a vapor. As per above, only 42 is needed for the activation barrier, so the
detonation of the vaporized H2C2O4 becomes possible (as the decomposition reaction is now apparently marginally exothermic, as was claimed by the
author, and confirmed by actual test detonation with a tetryl booster).
-------------------------------------------------------------
My calculations (source: https://www.google.com/search?q=heat+of+formation+table&... ):
H2O -285.8
CO2 -393.5
CO -110.5
Sub Total: -789.8
H2C2O4 -829.9
Net: -40.1
as compared to the study's citation of -42.
[Edited on 29-1-2014 by AJKOER]
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blogfast25
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Remember, we’re going H2C2O4(s) === > H2O(l) + CO(g) + CO2(g) or broken down using Hess:
H2C2O4 === > H2 + 2 C + 2 O2
Add:
H2 + ½ O2 === > H2O
C + ½ O2 === > CO
C + O2 === > CO2
Sum total is the decomposition to water, CO and CO2 (first equation).
For the reverse of the formation of solid H2C2O4, one has + 829.9 kJ/mol. Add to this the heats of formation of H2O, CO and CO2:
+ 829.9 – 285.8 – 110.5 – 393.5 = + 40.1 kJ/mol (not minus).
Slightly endothermic. I don’t know what the authors understand by ‘activation barrier'.
But Entropic effects could be very important here, because gases are generated. That could tip the ΔG (= ΔH – TΔS) to the negative
and that would make the reaction THERMODYNAMICALLY favourable. I'll see if my CRC lists any entropy values for oxalic acid.
What do you understand by ‘gain’? There is nothing gained here: the + 98.1 kJ/mol needs to be PUT IN to go from the solid state to the gaseous
state (the heat of vaporisation, a positive value, always).
[Edited on 30-1-2014 by blogfast25]
Edit:
My CRC gives standard Entropies (S) as follows:
H2C2O4(s) = + 109.7 J/mol K
CO(g) = + 197.7 J/mol K
CO2(g) = + 213.8 J/mol K
H2O(l) = + 70.0 J/mol K
For H2C2O4(s) === > H2O(l) + CO(g) + CO2(g)
The Standard Entropy change is thus: ΔS = - 109.7 + 70.0 + 197.7 + 213.8 = + 371.7 J/mol K or at 298 K: T ΔS = + 110.8 kJ/mol
With the estimated value for ΔH above, ΔG = ΔH - T ΔS then becomes negative, albeit fairly weakly.
[Edited on 30-1-2014 by blogfast25]
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