| Pages:
1
2 |
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
A reference on the creation (precipitation) of Magnesium oxalate: To quote from page 141 of "A dictionary of chemistry and mineralogy: with an account
of the ...", Volume 2, by Arthur Aikin, Charles Rochemont Aikin:
"The oxalat of magnesia, on account of the great insolubility of this earth, was formed by the addition of oxalat of ammonia to sulphat of magnesia.
It is remarkable that though the oxalat of magnesia is apparently equally insoluble in water as oxalat of lime, no turbidness appears on adding oxalat
of ammonia and sulphat of magnesia till the liquor is much reduced in bulk or heated, or till they have stood together for some hours. When the oxalat
of magnesia is once obtained, in either of these ways, or else by entire evaporation of the mixture and adding water to dissolve out the sulphat of
ammonia, it is a tasteless insoluble powder."
Link:
http://books.google.com/books?pg=RA1-PA141&dq=oxalat+of+...
From the above I would expect the chemistry of the preparation of H2SO4 from MgSO4 and Oxalic acid would proceed first to form the soluble Magnesium
hydrogen oxalate which subsequently breakdowns (from heating and/or time) to the insoluble Magnesium oxalate. Reference Wikipedia ( http://en.wikipedia.org/wiki/Oxalate ) to quote:
"Relationship to oxalic acid
The dissociation of protons from oxalic acid proceeds in a stepwise manner as for other polyprotic acids. Loss of a single proton results in the
monovalent hydrogenoxalate anion HC2O4−. A salt with this anion is sometimes called an acid oxalate, monobasic oxalate, or hydrogen oxalate."
Expected reactions:
MgSO4 + 2 H2C2O4 --> Mg(HC2O4)2 + H2SO4
MgSO4 + Mg(HC2O4)2 --> 2 MgC2O4 (s) + H2SO4
I would not be surprised if similar reactions occurred upon replacing MgSO4 with either Ammonium or Potassium or Sodium sulfate as the hydrogen double
salt (or bisulfate) exist. My speculation follows the apparent ability of Ammonium or Potassium or Sodium oxalate to dissolve a number of insoluble
oxalates forming soluble double salts. For example, per "A dictionary of chemistry and the allied branches of other sciences", Volume 4, by Henry
Watts, page 265 (link: http://books.google.com/books?pg=PA265&id=lYXPAAAAMAAJ#v... )
"Stannous oxalate dissolves in the oxalates of ammonium, potassium, and sodium, forming double salts."
Also:
"Neutral oxalate of ammonium dissolves oxalate of nickel, and the solution yields by evaporation green prisms of ammonio-nickel-oxalate. On adding to
the aqueous solution of this bait a small quantity of ammonia, a pale green precipitate is 'formed, consisting, according to Winckelblech (Ann. Ch.
Pharm. xiii. 278), of oxalate of nickel and nickel-ammonium"
And:
"Acid oxalate of potassium is used as a weak acid for scouring metals; also for removing ink-stains and iron-mould, the double oxalate of iron and
potassium being soluble in water."
[Edited on 9-5-2012 by AJKOER]
|
|
|
Formatik
National Hazard
Posts: 927
Registered: 25-3-2008
Member Is Offline
Mood: equilibrium
|
|
As long as sulfuric acid is present it could be expected to form equilibrium with the magnesium oxalate, forming this soluble acid oxalate. And so
hindering isolation. This would explain why I saw no precipitate. I remain skeptical that a route through magnesium sulfate could prove useful. I
won't pursue that reaction more. I've also tried some other reactions and described them below. Some success, some failure.
Perchloric acid attempt from oxalic acid (failed):
I've attempted to get perchloric acid through oxalic acid but this didn't work. The literature describes preparation of impure chloric acid from
solutions of sodium chlorate and oxalic acid frozen in a freezing mixture (Böttger, Lieb. Ann. 57 [1846] 138). Ammonium oxalate is about as soluble
as sodium oxalate, so I've attempted something similar with ammonium perchlorate and oxalic acid. The details are in the attached file below.
Another attempt of sulfuric acid from oxalic acid and CuSO4 (beware!):
I made another attempt with the copper sulfate and oxalic acid. But this time used larger amounts. This time I only filtered and siphoned the
filtrate, and did not evaporate and collect more solids. But this time I just boiled down the filtrate. Something very bad happened on boiling near
the end, all of the sulfuric acid and contents in the 600mL beaker ejected entierly! I think the sulfuric acid reacted violently with residual oxalate
(another crystallization would have been good) and the heating might have been too high.
The purity of acid made this way should be alright for some purposes. CuSO4 has a solubility of 0.19g in 100g of 92.70% H2SO4 at 25 C (Solubilities of
inorganic and organic compounds, 2nd ed. (1919) by A. Seidell). CuSO4 should be the end-product copper salt and the white solid that was seen earlier
in the brown acid.
On another similar note, aqueous copper sulfate yields no precipitate or any reaction of note when added slowly into an excess of aqueous citric acid.
The reactivity of citrates might be the reason why there is no reaction. Oxalic acid can be boiled with nitric acid and is able to partially resist
the attack.
Hydrochloric acid from calcium chloride and oxalic acid:
Aqueous solutions of CaCl2 and oxalic acid yield an immediate fine white precipitate of calcium oxalate when mixed. The liquid part of the solution of
this eventually attacked and disintegrated aluminium foil evolving H2, whereas aqueous oxalic acid was unreactive. The calcium oxalate could be
filtered with filter paper.
Nitric acid from copper nitrate and oxalic acid:
Aqueous copper nitrate and oxalic acid when mixed immediately gave a light green-blue precipitate (same one as by copper sulfate). This is one way to
recover nitric acid from copper nitrate. Calcium nitrate should also work instead of copper nitrate. So, oxalic acid is yet another acid that will
liberate nitric acid from some nitrates.
Both hydrochloric and nitric acids unlike sulfuric acid are very easy to purify by distillation.
Attachment: perchloric-attempt.txt (861B)
This file has been downloaded 1051 times
[Edited on 11-5-2012 by Formatik]
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
Formatik:
Good stuff.
I was wondering, per my research, if AVOIDING Sodium, Potassium and Ammonium salts, which have an apparent propensity of forming double salts, may
produce better results in making HClO4 via Oxalate acid. In particular, Calcium perchlorate and Oxalate acid, as an example.
Also on the expected reactions:
MgSO4 + 2 H2C2O4 --> Mg(HC2O4)2 + H2SO4
MgSO4 + Mg(HC2O4)2 --> 2 MgC2O4 (s) + H2SO4
it may follow that a better procedure is certainly NOT to add MgSO4 slowly to aqueous Oxalic acid as this is the first reaction, but instead slowly
add hot aqueous H2C2O4 to MgSO4 in slight excess.
------------------------------------------------------------
Also, on my Oxalic acid dihydrate, I did receive a recall from Ebay asking me if I wanted a refund. Said no, however, I have since noticed (the
basement lab did get to 77 F with a recent warming spell) a gaseous release (CO2) and an apparent transformation into Formic acid (no longer forms any
insoluble precipitates). In the future, I will keep it cool and purchase not solely based on price. If shelf life for H2C2O4 is limited, I will also
avoid buying in bulk.
[Edited on 12-5-2012 by AJKOER]
|
|
|
Formatik
National Hazard
Posts: 927
Registered: 25-3-2008
Member Is Offline
Mood: equilibrium
|
|
Calcium perchlorate would be ideal for making perchloric acid from oxalic acid. But the problem is how to get it. One way is through calcium chloride
electrolysis (an unfavorable process according to this thread and this thread also). Another way would be through ammonium perchlorate and calcium hydroxide in aqueous media. Though this latter method has its
pitfalls as I'll explain below, since I've tried it.
The idea was to just boil the ammonia off and leave calcium perchlorate behind. Ammonium perchlorate was in aqueous solution with calcium hydroxide,
where calcium hydroxide was not solvated completley (this is the problem, but being done on purpose to see if it would work regardless). These two
upon mixing will give off a mediocre but not strong odor of ammonia. These two components were mixed in stoichiometric amounts. Boiling the mixture
gave off stronger amount of ammonia. However, despite further boiling only partial reaction occurred and this is problematic because of calcium
hydroxide's low solubility. My conclusion was that all calcium hydroxide must be entierly solvated in the water for this reaction to function, but
that means that a lot of water has to be used.
[Edited on 21-5-2012 by Formatik]
|
|
|
elementcollector1
International Hazard
Posts: 2684
Registered: 28-12-2011
Location: The Known Universe
Member Is Offline
Mood: Molten
|
|
Why not two-cell electrolysis of a sulfate? Seems that the sulfate ions would move much the same as chloride ions in a brine setup.
Elements Collected:52/87
Latest Acquired: Cl
Next in Line: Nd
|
|
|
pedrovecchio
Harmless
Posts: 38
Registered: 12-3-2012
Member Is Offline
Mood: No Mood
|
|
A reference was posted here
http://www.sciencemadness.org/talk/viewthread.php?tid=19639
about the preparation of sulfuric acid from oxalic acid. See the second post.
|
|
|
Formatik
National Hazard
Posts: 927
Registered: 25-3-2008
Member Is Offline
Mood: equilibrium
|
|
That thread has some nice old interesting references. Suitable cross-reference there. It has some similarities to the primordial chemicals thread, which is maybe not as systematic.
Concerning:
Nitric acid from calcium nitrate and oxalic acid:
Quote: Originally posted by Formatik  | | ... Calcium nitrate should also work instead of copper nitrate. So, oxalic acid is yet another acid that will liberate nitric acid from some
nitrates. |
The reaction of calcium nitrate and oxalic acid has been confirmed. Aqueous solution of calcium nitrate when mixed and swirled with aqueous oxalic
acid within a few moments turned cloudy and formed a white calcium oxalate precipitate fairly quickly.
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
Per a recent thread (see http://www.sciencemadness.org/talk/viewthread.php?tid=20109 ) to quote:
| 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. ......
|
as a word of warning when attempting to prepare HClO3 by the reaction of a chlorate and Oxalic acid, for example:
2 NaClO3 + H2C2O4 --> 2 HClO3 + Na2C2O4 (s)
in the event of excess Oxalic acid, however, the H2C2O4 subsequently acts as a reducing agent on the newly formed Chloric acid releasing equal amounts
of ClO2 and CO2 as reported by Chemrtd. Per one source, Chlorine-free Chlorine dioxide is formed as follows (see http://books.google.com/books?id=6wUmteTIc18C&pg=PA334&a... ):
2 HClO3 + H2C2O4 --> 2 ClO2 + 2 H2O + 2 CO2
Another source citing the reducing ability of Oxalic acid on HClO3 is a Chinese Patent (see http://osdir.com/patents/Chemistry-inorganic/Method-producin... ) to quote: "In contrast to expensive reducing agents previously employed such as
H2O2, oxalic acid, methanol, SO2 and saccharose, Urea is inexpensive, non-toxic and easily obtainable."
Given the toxic and explosive nature of ClO2, I am issuing this cautionary note as one should be slowly adding with stirring H2C2O4 to a solution of
excess NaClO3, and not the other way around (or elect to not perform this dangerous synthesis without proper safeguards in place).
[Edited on 7-8-2012 by AJKOER]
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
| Quote: Originally posted by Formatik | .....
Another attempt of sulfuric acid from oxalic acid and CuSO4 (beware!):
I made another attempt with the copper sulfate and oxalic acid. But this time used larger amounts. This time I only filtered and siphoned the
filtrate, and did not evaporate and collect more solids. But this time I just boiled down the filtrate. Something very bad happened on boiling near
the end, all of the sulfuric acid and contents in the 600mL beaker ejected entierly! I think the sulfuric acid reacted violently with residual oxalate
(another crystallization would have been good) and the heating might have been too high.
The purity of acid made this way should be alright for some purposes. CuSO4 has a solubility of 0.19g in 100g of 92.70% H2SO4 at 25 C (Solubilities of
inorganic and organic compounds, 2nd ed. (1919) by A. Seidell). CuSO4 should be the end-product copper salt and the white solid that was seen earlier
in the brown acid.
On another similar note, aqueous copper sulfate yields no precipitate or any reaction of note when added slowly into an excess of aqueous citric acid.
The reactivity of citrates might be the reason why there is no reaction. Oxalic acid can be boiled with nitric acid and is able to partially resist
the attack.
..........
|
Please do not attempt to push this reaction to the point where concentrated H2SO4 is formed. The violent ejection on heating, per recent research,
could be from the abrupt decomposition of unreacted Oxalic acid into H2O, CO and CO2. Source: Watts' dictionary of chemistry, Volume 3, by Henry
Watts, page 649 under 'Reactions'. To quote:
"-2. On heating with conc. H2SO4 or with P2O5 it is resolved into water, CO and CO2."
Link: http://books.google.com/books?pg=PA649&lpg=PA649&sig...
[Edited on 7-8-2012 by AJKOER]
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
| Quote: Originally posted by Formatik | .....
Hydrochloric acid from calcium chloride and oxalic acid:
Aqueous solutions of CaCl2 and oxalic acid yield an immediate fine white precipitate of calcium oxalate when mixed. The liquid part of the solution of
this eventually attacked and disintegrated aluminium foil evolving H2, whereas aqueous oxalic acid was unreactive. The calcium oxalate could be
filtered with filter paper.
|
Calcium chloride, in my opinion, should not be the chloride of choice here in preparing HCl from H2C2O4. My concerns rest on research indicating how
Calcium oxalate and HCl interact under varying conditions. To quote from Watt's, page 254 at http://books.google.com/books?pg=PA254&id=lYXPAAAAMAAJ#v... :
"Calcium oxalate is precipitated as a white powder whenever a soluble calcium-salt is mixed with oxalic acid or an alkaline oxalate, provided there be
no strong mineral acid present in large excess. It is insoluble in water, acetic acid, and solution of sal ammoniac, nearly insoluble in free oxalic
acid, and sparingly soluble in lactic acid, but it dissolves with tolerable facility in hydrochloric or nitric acid, whence it is precipitated by
caustic alkalis or alkaline carbonates."
Also per page 255:
"A solution of calcium oxalate in hot hydrochloric acid, deposits crystals of the salt CaC2O4.H20 (E. Schmid). According to Souchay and Lenssen, this
salt is deposited on cooling, when oxalate of calcium is added at 100°, to hydrochloric acid of specific gravity less than 1.1, in quantity
sufficient to saturate it; but if the solution is not saturated, it deposits after some time, square prismatic crystals consisting of
CaC2O4.3H2O.—By adding oxalate of calcium to warm hydrochloric acid, of specific gravity 1-10 or higher, double salts are obtained in scaly
crystals, consisting of oxalate and chloride of calcium."
If the Ca2C2O4 happens to dissolve, the reaction creating HCl is reversible. A better and less expensive route is perhaps via dry NaCl as per Watt's
(http://books.google.com/books?pg=PA649&lpg=PA649&sig... ), "Oxalic acid expels HCl when heated with dry NaCl."
Note, these issues (solubility and double salts) are not concerns working with Ca(NO3)2 and H2C2O4 as to quote Watt's: "With nitric acid, oxalate of
calcium behaves in the same manner as with hydrochloric acid, excepting that it is insoluble in strong nitric acid, and therefore does not yield any
oxalato-nitrate (Souchay and Lenssen). According to Schmid, a solution of calcium oxalate in hot nitric acid, deposits monoclinic laminae of the
monohydrated salt, the last mother-liquors, however, yielding free oxalic acid."
[Edited on 7-8-2012 by AJKOER]
|
|
|
Formatik
National Hazard
Posts: 927
Registered: 25-3-2008
Member Is Offline
Mood: equilibrium
|
|
@AJOKER:
#1. Concerning sodium chlorate and oxalic acid:
This reaction is not really dangerous from what I understand. I have looked at the original reference and it mentions no hazards. Just mixing the
sodium chlorate with saturated oxalic acid in aq. solutions, then chilling with a freezing mixture (HCl acid and Glauber's salt) and filtering.
There is also no dry heating of ingredients, never do this! The literature notes using the method of heating aqueous solutions of chlorate with oxalic
acid has been used to make ClO2 available for bleaching purposes without the explosion hazard.
#2. Concerning violent observations during H2SO4 preparation.
This is what I speculated happened also in the thread up above by pedrovecchio. The time it did not happen was after several crystallizations and a
milder heating.
#3. Hydrochloric acid from oxalic acid:
The distillation of HCl acid would not be difficult. But if the method as described by Watt's works by dry heating the oxalic acid with NaCl, it would
be preferable. That mean yet another basic compound that forms HCl when heated with salt on top of NaHSO4, Epsom salt.
|
|
|
tetrahedron
Hazard to Others
Posts: 210
Registered: 28-9-2012
Member Is Offline
Mood: No Mood
|
|
i can see at least two problems when using organic acids for the type of displacement described in this thread
1. organic acids like oxalic, tartaric etc are weak and only partially dissociate in a concentrated solution. the equilibrium becomes
even less favorable as the organic salt starts precipitating and the inorganic acid (nearly fully dissociated) causes a rapid increase in the
concentration of H+
2. particularly with oxidizing anions such as nitrate and (per)chlorate, but also e.g. hot sulfate, the decomposition of the organic
substrate becomes relevant (even more so as the pH sinks)
so in the end it's a matter of equilibrium, and you'll more or less invariably end up with a product contaminated with organic stuff and/or the
original cation, depending on the ratio of educts. this might or might not be an issue depending on the application.
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
Tetrahedron:
To quote Wiki on Oxalic acid:
"Oxalic acid is a relatively strong acid, despite being a carboxylic acid:
C2O4H2 → C2O4H− + H+; pKa = 1.27
C2O4H− → C2O42− + H+; pKa = 4.27"
Also, your implied comment on the behavior of Oxalic acid in Nitric acid, is largely incorrect. Per Watt's boiling Nitric acid only slowly oxidizes
Oxalic acid to CO2.
Source: "Watts' dictionary of chemistry", Volume 3, by Henry Watts, page 649.
Link:
http://books.google.com/books?pg=P649&lpg=PA649&id=Q...
Also, the thread "Oxalic acid and nitric acid" at http://www.sciencemadness.org/talk/viewthread.php?tid=20109
|
|
|
tetrahedron
Hazard to Others
Posts: 210
Registered: 28-9-2012
Member Is Offline
Mood: No Mood
|
|
i stand corrected on the oxidation by nitric acid. indeed, oxalic acid is obtained by the oxidation of sugars using nitric acid, so it makes sense
that oxalic acid is a relatively stable end product
H2SO4 -> H+ + HSO4- has a pKa = -3. this severely inhibits the second dissociation of oxalic acid necessary to precipitate Cu, i.e. taking the
reaction to completion becomes exponentially slow, i.e. the end product will always be contaminated (either by Cu or OA)
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
For the record, if heating Oxalic acid and NaCl produces HCl gas per Watt's, then heating anhydrous H2C2O4 and NaI (NaBr) may also be a path to
anhydrous HI (HBr).
Just two more acids to add to the list.
|
|
|
tetrahedron
Hazard to Others
Posts: 210
Registered: 28-9-2012
Member Is Offline
Mood: No Mood
|
|
this route for HBr might just work, but for HI probably not. HI is very susceptible to oxidation to I2, it decomposes in the presence of O2,
accelerated by light. according to Brauer you would use hypophosphorous acid instead (a reducing agent). that's what makes red phosphorus so important
=D
edit. actually Brauer reports I2 + H3PO2, the wikipedia reference is wrong. H3PO2 might be overkill, but it probably works too
[Edited on 3-10-2012 by tetrahedron]
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
I suspect that Oxalic acid acting on Sodium iodide will also form Hydroiodic acid, as per this source page 182 (https://docs.google.com/viewer?a=v&q=cache:mE02fxIC8rcJ:... ), even Citric acid is apparently so capable of creating aqueous HI.
Now, even if some the HI is reduced to I2 (say by air), then in the presence of excess H2C2O4, upon shaking to capture the Iodine, it is possible that
Hydroiodic acid could be reformed. See the implied net ionic equation at http://www.sciencebuddies.org/science-fair-projects/project_... , namely:
C2O4- + I2 → 2 I- + 2 CO2
|
|
|
Kemisten
Harmless
Posts: 2
Registered: 4-5-2011
Member Is Offline
Mood: Cool
|
|
Oxalic acid dihydrate and calcium nitrate tetrahydrate when heated yields nitric acid, although there is considerable decomposition.
Experimental:
25g of oxalic acid and 45g of calcium nitrate was added together in a distilling flask and heated. Quickly after the salts melted and yellow fumes
were given off, and then strong white fumes. After about 10 minutes droplets of clear liquid started to condense, at this point the salts were boiling
away and the flask filled with blood red NO2. After 30 minutes I stopped and had recieved 17ml of WHITE FUMING NITRIC ACID .
I'm using a home rigged glass/PTFE hybrid liebeg condenser.
I used agricultural grade calcium nitrate prills which should have been powdered if my ballmill wasn't broken and hardware store oxalic acid.
The picture shows how it looked after 10-15 minutes.
I suspect that I had too much heat which caused the severe decomposition. I will use an oilbath next time!
I'll do some more testing on ratios, heat, addition of water and I just ordered 5kg of labgrade calcium nitrate powder.
If this works, I'll order 25kgs of oxalic acid 
<!-- bfesser_edit_tag -->[<a href="u2u.php?action=send&username=bfesser">bfesser</a>:
"CN"→calcium nitrate</sub>]
[Edited on 7/13/13 by bfesser]
|
|
|
RingoStarr
Harmless
Posts: 9
Registered: 16-12-2012
Member Is Offline
Mood: No Mood
|
|
There is a nice video on youtube here on how to make nitric acid with calcium nitrate fertilizer and sulfuric acid. The overall equation for the reaction is:
2 [5Ca(NO3)2.NH4NO3·10H2O] + 11 H2SO4 → 22
HNO3 + 10 [CaSO4·2H2O] + (NH4)2SO4 + 8H20
I'm wondering if oxalic acid reacts with ammonium nitrate to form nitric acid or ammonium oxalate or if the ammonium nitrate is unaffected.
Case 1 → nitric acid is produced:
2 [5Ca(NO3)2·NH4NO3·10H2O] + 11
[H2C2O4·2H2O] → 22 HNO3 + 10
[CaC2O4·H2O] + (NH4)2C2O4·2H2O +
31H2O
Case 2 → ammonium nitrate as a spectator:
2 [5Ca(NO3)2·NH4NO3·10H2O] + 10
[H2C2O4·2H2O] → 20 HNO3 + 10
[CaC2O4·H2O] + 2NH4NO3 + 30H2O
I'm after approximately 70% nitric acid just like in the youtube video. Any help to determine if case 1, case 2 or neither case is correct is much
appreciated.
|
|
|
Formatik
National Hazard
Posts: 927
Registered: 25-3-2008
Member Is Offline
Mood: equilibrium
|
|
@Kemisten. It would also be interesting to a density reading on the acid, since calcium nitrate and oxalic acid are usually full of water (as
hydrates, especially the ones used in your synthesis), so that I wouldn't expect that "fuming nitric acid" would form, but a less concentrated acid.
@RingoStarr. It is probably more like case number 2. I haven't had any success forming perchloric acid using ammonium perchlorate (failed attempt
described above). And I have before mixed small amounts of saturated aq. solutions of NH4NO3 and oxalic acid, chilled this with ice for several
minutes and poured the solution, the solution didn't attack copper at room temperature after standing for over 30 minutes.
|
|
|
Kemisten
Harmless
Posts: 2
Registered: 4-5-2011
Member Is Offline
Mood: Cool
|
|
| Quote: Originally posted by Formatik | @Kemisten. It would also be interesting to a density reading on the acid, since calcium nitrate and oxalic acid are usually full of water (as
hydrates, especially the ones used in your synthesis), so that I wouldn't expect that "fuming nitric acid" would form, but a less concentrated acid.
|
Nitric acid
As stated in wikipedia, HNO3 begins to fume at a concentration above 86%, so i guess the acid is just around that percentage. I do not
believe the acid made from this process can be over 90% concentration due to the water in the reactants.
But, this process is still very interesting because both calcium nitrate and oxalic acid are cheap and very available. The acid can then be
concentrated by other methods later on.
I'll check the density from freshly made acid soon.
|
|
|
Formatik
National Hazard
Posts: 927
Registered: 25-3-2008
Member Is Offline
Mood: equilibrium
|
|
This nitric acid is much less concentrated. The reaction should be:
Ca(NO3)2.4 H2O + H2C2O4.2 H2O = CaC2O4 + 2 HNO3 + 6 H2O
Another reaction is responsible for the red fumes:
4 HNO3 = 4 NO2 + O2 + 2 H2O
We don't know how much decomposition of nitric acid may have occurred. So being very generous here and even assuming a complete conversion according
to the first equation gives masses of 126.02g nitric acid and 108.06g water, which is 53.8% nitric acid.
The concentration of the nitric acid might be even less though considering some nitric acid decomposition, and nitrogen dioxide loss. If the starting
ingredients both have less water, then a higher acid concentration should be obtainable.
The fuming nature of nitric acid can be vague. I have seen it fume at less concentration. Humidity can play a role in fumes of acids, so can
undissolved fumes in containers. White fuming nitric acid used in the professional context usually means nearly anhydrous nitric acid (99.9%). In the
amateur context, it could mean any high strength nitric acid that fumes white (I've used it that way too  ). Then it's better to say whitish fuming nitric acid and include density measurements to avoid any mix-up. Though
fuming nitric acid usually means very high concentrated, above regular "concentrated nitric acid".
Fumes are definitely not an accurate indicator of concentration, a titration needs to be done, or at the very least a density reading to give a rough
idea. But based on the above, concentration is not going to be much above 40-50%.
|
|
|
| Pages:
1
2 |
![[*]](images/xpblue/default_icon.gif){kind=icon}
