Sciencemadness Discussion Board - Using Oxalic Acid to Make H2SO4,...

Using Oxalic Acid to Make H2SO4,...

AJKOER - 24-2-2012 at 14:55

Having noted the relative insolubility of many oxalates, I was wondering the potential use/feasibility of employing Oxalic Acid to make various other reagents, including H2SO4 and even HNO3. For example, with say FeSO4, the reaction could be:

FeSO4 + H2C2O4 --Heat --> FeC2O4 (s) + H2SO4

and similarly with CuSO4. Also, with nitrates:

Ca(NO3)2 + H2C2O4 = 2 HNO3 + CaC2O4 (s)

As support, see the details of a lab preparaton using Mohr's Salt, H2SO4 and Oxalic acid (in particular, the Part B synthesis on the third page). Based on this synthesis, if one were to add H2C2O4 to FeSO4, I would expect a yellow precipitate of Iron Oxalate upon gentle heating to boiling with stirring.

Please paste this Link:
"https://docs.google.com/viewer?
a=v&q=cache:xTMbuODMUQcJ:https://eee.uci.edu/programs/hongchem/6hMANprepfe.pdf+oxalic+acid+preparation+h2so4&hl=en&gl=us&pi
d=bl&srcid=ADGEEShSK5xAHcf4e2sc5yoZ0ioC_M76Hs-
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_ghyjX_dT3tESvQiPNppA"

There is also prior commentary on Sciencemadness:

Thread: H2SO4 Cheap way?

LINK:
http://www.sciencemadness.org/talk/viewthread.php?tid=14120#...

Quote: Originally posted by bbartlog

Even if it works, I don't think citric acid (or copper sulphate for that matter) are all that cheap compared to sulfuric acid. If you think it's a neat experiment, have at it; but if you think it's economical I have to wonder what the prices look like where you are.
I have also seen a claim in an old book that oxalic acid will do this. The insolubility of the oxalate (and maybe the citrate, dunno) might allow the formation of copper oxalate and sulfuric acid to be favored, but I'm skeptical. In the case
of citric acid, I know that calcium citrate is turned back into citric acid by addition of sulfuric acid, so I doubt the
reaction runs the other way for copper.

Quote: Originally posted by BromicAcid

Reminds me of something I read once on making very pure sulfuric acid. The method was to mix a saturated solution of calcium sulfate with oxalic acid, and then filter off the calcium oxalate. On the plus side an excess of oxalic acid would 'burn up' on concentration. On the minus side look at the solubility of calcium sulfate

Quote: Originally posted by Formatik

Try calcium chlorate and oxalic acid. Caclium oxalate is around as insoluble as BaSO4.

Thread:Subject: Nitric acid
http://www.sciencemadness.org/talk/viewthread.php?tid=401&am...

Quote: Originally posted by BromicAcid

Calcium oxalate is insoluble as all hell, maybe Ca(NO3)2 + H2C2O4 ---> 2HNO3 + Ca2C2O4 but I dont know how the oxalate would fair in such an increasingly acidic enviorment. Another worry would be Oxalic acid being precipiated out before the reaction is anywhere near complete due to it being a weak acid and the equilibrium being shifted due to increasing pH of the solution, anyways, it's just a thought, maybe more manageable then
the CaSO4 precipiate, maybe a constat system at high temperature could be done with slow addition of the nitrate to a oxalate solution with effient boiling off of the nitric acid as it is formed keeping equilibrium favored.

Quote: Originally posted by Mumbles

This is just a thought I've been having. Whenever people do the precipitation method they add all the Calcium nitrate at once and can get none or little Nitric acid. What if you added it in portions. Add an amount, then filter. Add another amount and filter again. This would increase the amount of nitric acid everytime. I'd imagine that a smaller amount of precipitate could be pressed harder to remove the remaining acid as well.

This has been running through my mind for a couple days. I've been trying to figure out what wouldn't work. I can't find
a reason for it not to. I know its far from perfect, but it may be a way to get a signifigantly larger amount of very high
purity Nitric acid in a relitively small amount of time.

For interesting properties of H2C2O4, see link:

http://books.google.com/books?id=AugDAAAAMBAJ&pg=PA759&a...

If anyone has experience using Oxalic acid, or would like to provide references, or opinions, please feel free to contribute to this thread.

[Edited on 25-2-2012 by AJKOER]

Endimion17 - 24-2-2012 at 15:34

Yeah, I've got some experience. It's too expensive to use it for making such dirt cheap thing as sulphuric acid.

But if it's from an academic standpoint, it's interesting.

AJKOER - 24-2-2012 at 17:27

Here is a good description on how to use oxalic acid. An interesting comment is that it is easy to find, use and the safest for the home as compared to other acids used for cleaning concrete and dissolving iron. To quote:

"Oxalic Acid
Anything that has the word "acid" sounds ominous. But oxalic acid is easy to find, use and the safest for the home. In fact it is found in many vegetables including spinach. It is used to dissolve the iron oxide (brown) stain on all minerals. Specimens collected at Phoenixville, Ellenville, Case Quarry, NH smoky quartz and many others clean up beautifully with oxalic acid. Zeolites do not respond as well, so you should test beforehand on small specimens to see how they react.

To make this as simple as possible I will give a step by step guide to its use. Do not take any shortcuts or make substitutions. Purchase a one pound box of Oxalic Acid (OA) powder at your local hardware store in the paint department or at a paint store. It is used as wood bleach and will be labeled as such. The most common brand is Rainbow.

Fill a plastic one gallon container 3/4 full with hot tap water. Pour in the OA crystals and stir for five minutes. Be careful not to inhale any powder when adding the crystals. Once the OA is dissolved top off the container to a full gallon. Label the container and put out of reach of children or pets.

When you are ready to use it place your specimens in a plastic container and add enough OA solution to cover. Set aside for several days.

After the iron color has disappeared then you can remove the specimens (with gloves on) and wash under running water for three hours. Then soak in clean water for a day changing the water as often as possible.

Heat speed up the reaction, as does agitation. If you have a hot plate and can set up outdoors or in an area with good ventilation the repeat step 4 but heat the solution to bath water hot (110 F). Never Boil! You will find that an hour in hot solution will usually do the trick."

Per another source:
"Oxalic Acid can be found and purchased on-line at the realmilkpaint.com for a very reasonable price of 6.99 per pound and is called Rainbow Oxalic Acid."

At realmilkpaint.com, the quoted price is $9.50 for 16 oz. As 1 ounce is 28.3495 grams, this is 453.6 grams, and as 1 mole is 126 grams, this constitutes 3.6 moles, so the cost per mole of Oxalic dihydate is $2.64. Now, for battery acid, say $5 for 4 lbs of 38% H2SO4. So, the price per mole (1 mole is 98.1 grams for pure H2SO4), my calculations give a cost per mole of about $0.71. So, H2C2O4 per mole is more expensive, but still relatively cheap. Note, H2C2O4 can be shipped whereas H2SO4 must be picked up as shipping is generally unavailable or very expensive.

[Edited on 25-2-2012 by AJKOER]

elementcollector1 - 27-2-2012 at 21:27

What of magnesium sulfate? It's relatively soluble, while magnesium oxalate is quite insoluble (0.138g/100mL). This could make sulfuric acid even dirt-cheaper than it already is if it works.

AJKOER - 28-2-2012 at 10:08

Actually, I was also thinking about the reaction:

MgSO4.7H2O + H2C2O4.2H2O ---> MgC2O4 + 9H2O + H2SO4

as it employs a common item, Epsom's Salt, found in most homes or easily purchased at most pharmacies.

Theoretically, if one were to dehydrate the Magnesium heptahydrate (Epsom's Salt) first, the acid could be even more concentrated. Some questions (until I perform this experiment), is how soluble is Magnesium oxalate in H2SO4, is the reaction reversible with strong heating decomposing the H2C2O4, can the reaction mix be boiled to concentrated the H2SO4 after removing the MgC2O4, or just distilled, do I have to use an excess of MgSO4 to avoid the formation of a soluble oxalate complex?

Worst case, dilute and cool to collect the MgC2O4 crystals and re-concentrate by distillation.

Advantages: a convenient, seemingly simple and safe route to pure H2SO4, as I doubt if there are significant impurities (especially heavy metals) in a health product. Also, if one immediately employs whatever acid one has prepared, there is a safety benefit in having one less dangerous compound (like battery acid) lying around.

Disadvantages: Unless one already has Epsom's salt, the expensive of your pure (dilute?) acid has just further increased. I calculate, for example, that 98% H2SO4 (from DuaDisel) for 950 ml with shipping to my home is only about $1.30 per mole, half the cost of buying a mole of Oxalic acid dihydate, although there could be quality differences.

[Edited on 28-2-2012 by AJKOER]

Formatik - 1-3-2012 at 00:13

If you are a fan of lesser known methods, I've used stoichiometric MgSO4 and NaOH in aqueous solutions, to precipitate Mg(OH)2 and leave mostly Na2SO4 behind. Then crystallizing the Na2SO4 and collecting.

I then used the Na2SO4 with hydrochloric acid to make sulfuric acid. It is such a laborious and energy-consuming process. The result from 7.4g dry Na2SO4 and 31mL 31% HCl was 2.3g H2SO4 (around 80-90% strength), so around less than half of the sulfate portion of sodium sulfate was converted to H2SO4.

Boiling down HCl was bad enough, but distilling H2SO4 at atmospheric pressure is one of the worst things. I was using two Bunsen burners on one distillation flask just to distill such a very small amount of H2SO4. Distillation is necessary for removal of inorganic salts.

The process described in more detail here and here.

My comment above was also based on my experience of forming chloric acid from calcium chlorate and oxalic acid. Mentioned here. A precipitate did not form on mixing but on standing, I was hoping it would on mixing for stoichiometry since the calcium chlorate was impure. I decided to try the reaction after reading about old methods used to make impure chloric acid from sodium chlorate and oxalic acid mentioned earlier in the same previous thread. Chloric acid was also made from tartaric acid and potassium chlorate at versuchschemie (potassium bitartrate is fairly insoluble).

The oxalic acid and magnesium sulfate could be worth a shot, provided aqueous sulfuric acid doesn't react with oxalic acid and the solubility of magnesium oxalate is not increased much by sulfuric acid. The formation of an oxalate precipitate whether immediate or delayed, should put to rest any doubts.

weiming1998 - 2-3-2012 at 23:44

I have made 60-70mls of sulfuric acid using this method. I don't know about the purity and don't have the scales and equipment to titrate, but it has a much higher boiling point than water and is sort of oily (though not as thick as concentrated sulfuric acid) When it is dilute, it readily attacks iron when cold, but now it passivates. Copper sulfate does not form a blue solution in it at all.

But the strange thing is, it doesn't attack paper, and it produces a strange gas with KMnO4 that stings the inside of my nostrils. I'll try nitric acid using oxalic acid and copper nitrate next.

Edit: Oxalic acid and copper nitrate are somewhat successful, but it's a bit weird. It will only react with copper when heated (I do not think it's passivation, because even passivation will form a coating), which indicates a very dilute acid, but it took about 30 grams of CaCO3 to neutralize the 50mls. 2HNO3+CaCO3===>Ca(NO3)2+H2O+CO2. Pure nitric acid has a density of about 1.5g/ml, so 50ml is 75g. 126g of acid is required to neutralize 100 grams of CaCO3. Scaling it down, 75 is roughly 3/5 of 126, and 3/5 of 100g is 60g. 60g of CaCO3 is required. I report about 30g, so with an error variation of +- 10g, the nitric acid has to be at least 30% pure, not to mention a water-acid mix is less dense than pure nitric acid! A mix of at least 30% HNO3 will react with copper at room temp, not only when heated!

Thinking that making nitric acid by oxalic acid was an unsuccessful attempt, I tossed in a few chunks of KNO3 in my beaker of sulfuric acid, and heated it. It reacted, producing small amounts of gases that smells sharp, but can't be seen. Soon I took off the heat and got a pretty viscous mix (even without the salt contamination) that is about comparable in viscosity as concentrated drain cleaner H2SO4. I am now trying to nitrate some celluose in it.

So to sum it up, this is perhaps my most successful attempt in making H2SO4. Much better than the rest.

[Edited on 3-3-2012 by weiming1998]

[Edited on 3-3-2012 by weiming1998]

Formatik - 3-3-2012 at 10:57

Quote: Originally posted by weiming1998

I have made 60-70mls of sulfuric acid using this method.

What did you use to make the H2SO4? Did you see a precipitate form after mixing?

H2SO4 will not eat tissue paper as fast if it is contaminated by salts or has not been concentrated enough by boiling down. Remember, H2SO4 vapors are carcinogenic, so don't inhale.

Quote:

Edit: Oxalic acid and copper nitrate are somewhat successful, but it's a bit weird.

Note the solubility product of copper oxalate is close to barium sulfate. Nitric acid is easily purified by distillation. Copper sulfate and oxalic acid could then work to also form H2SO4.

Sodium oxalate has a fairly low solubility also. The solubility is significantly less than oxalic acid. So maybe it works with sodium sulfate to form H2SO4, mixing stoichiometric amounts then chilling the solution.

Potassium oxalate: 36g per 100mL at 20 C.
Sodium oxalate: 3.7g per 100mL at 20 C.
Oxalic acid: 9 to 10g/100mL at 20 C.
Source for solubility data (in H2O): GESTIS-Stoffdatenbank.

Perhaps potassium oxalate could also be used to make oxalic acid, by reacting it with hydrochloric acid and then chilling and filtering oxalic acid.

Be very careful with oxalic acid though unless you want a kidney stone, chronic skin absorption and chronic inhalation of oxalic acid are noted as giving kidney stones and urinary tract stones in some MSDS.

weiming1998 - 3-3-2012 at 15:15

Quote: Originally posted by Formatik

Quote: Originally posted by weiming1998

I have made 60-70mls of sulfuric acid using this method.

Quote:

Edit: Oxalic acid and copper nitrate are somewhat successful, but it's a bit weird.

1, Yes, I did see a milky white precipate form and yes, it is contaminated by salts so that's why it did not go through my filtering paper. I stopped boiling it as soon as I see the white vapours coming out. The problem is, I don't have a distillation setup. If I did, I can probably distil nitric acid using NaHSO4+KNO3 (a mixture of water vapour and bright red NO2 comes out instead of the acid.)

2, Sodium sulfate has quite a low solubility as well (4g/100ml
at room t, but if the temperature 32oC or above is achieved, it quickly reaches 40g/100ml solubility. I suppose warm water will do.
3, That would probably work, but where would you get your potassium oxalate from (apart from a chemical supply store)
4, I did not know that there is skin absorption of oxalic acid, so I will be more careful. The very sharp and piercing smell of boiling oxalic acid solution on a stove is the reason why I stopped boiling the solution at all.

[Edited on 3-3-2012 by weiming1998]

AJKOER - 4-3-2012 at 17:38

Thanks for the observations Weiming1998.

I manage to purchase Oxalic acid online at Ebay after not being able to find it at Home Depot or Lowe's (also got some weird looks). I eventually paid $14 for 2 pounds with free shipping, so the cost per mole for 99% pure Oxalic acid dihydate is $1.94. I previously noted that for 950 ml of 98% H2SO4 from DuaDisel would cost with shipping about $1.30 per mole.

My most interesting project, at some point in the future, would be to make HClO3 from H2C2O4 + KClO3 (Potassium chlorate from KCl + Magnesium chlorate). The Mg(ClO3)2 is from adding MgSO4 to HOCl (from NaClO and dilute H2C2O4), warm to 70 C, at which point the Magnesium hypochlorite does rapidly disproportionate into the chloride and chlorate. Hopefully, upon adding H2C2O4 to the KClO3 crystals, I should form HClO3.

[Edited on 5-3-2012 by AJKOER]

weiming1998 - 4-3-2012 at 18:23

Quote: Originally posted by AJKOER

Thanks for the observations Weiming1998.

I manage to purchase Oxalic acid online at Ebay after not being able to find it at Home Depot or Lowe's (also got some weird looks). I eventually paid $14 for 2 pounds with free shipping, so the cost per mole for 99% pure Oxalic acid dihydate is $1.94. I previously noted that for 950 ml of 98% H2SO4 from DuaDisel would cost with shipping about $1.30 per mole.

My most interesting project, at some point in the future, would be to make HClO3 from H2C2O4 + KClO3 (Potassium chlorate from KCl + Magnesium chlorate). The Mg(ClO3)2 is from adding MgSO4 to HOCl (from NaClO and dilute H2C2O4), warm to 70 C, at which point the Magnesium hypochlorite does rapidly disproportionate into the chloride and chlorate. Hopefully, upon adding H2C2O4 to the KClO3 crystals, I should form HClO3.

[Edited on 5-3-2012 by AJKOER]

How about the direct addition of oxalic acid to a magnesium chlorate solution? Potassium chlorate is less soluble than potassium oxalate, so it is not a favoured equilibrium of the reaction. Sodium chlorate, magnesium chlorate or even calcium chlorate would probably work wonders.

Anyway, there is one problem with the production of H2SO4 using oxalic acid. When I heat, then cool the solution, large amounts of spiky crystals form at the bottom of the beaker. It takes as much as 3-4 filterings to get rid of it completely. It forms, and messes up the reaction with KMnO4, even when I add an excess of MgSO4 instead of oxalic acid.

AJKOER - 4-3-2012 at 20:47

OK, on the chlorate question, the Mg(ClO3)2 is formed along with MgCl2 per the disproportionation of the Mg(ClO)2. As such, if I add H2C2O4 to this solution, I could get HClO3 and HCl, which could further react. For example:

HClO3 + HCl <----> HClO2 + HOCl

HClO3 + HClO2 ---> 2 ClO2 + Cl2 + 2 H2O

Transforming into the less soluble KClO3 and extracting allows the formation of a more stable HClO3 solution (at least, I hope).

Sorry, but I am not exactly clear on the Sulfate salt being reacted with the H2C2O4 to make H2SO4 (MgSO4, Na2SO4, CaSO4,...). Also, do not boil an aqueous solution of H2C2O4 in the presence of a strong acid (like H2SO4) as it may start to decompose (to H2O, CO and CO2) or a strong oxidizer (to H2O and CO2). This may reverse a reaction as Oxalic acid's undergoes a gaseous decomposition. Note the direction for use above (that I supplied when employing as a rust remover) clearly states do not boil. That may be due to either a REDOX reaction involving H2C2O4 in the presence of Fe2O3, or the impact of heat on the soluble complexes formed to dissolve the rust.

[Edited on 5-3-2012 by AJKOER]

weiming1998 - 4-3-2012 at 22:08

Quote: Originally posted by AJKOER

OK, on the chlorate question, the Mg(ClO3)2 is formed along with MgCl2 per the disproportionation of the Mg(ClO)2. As such, if I add H2C2O4 to this solution, I could get HClO3 and HCl, which could further react. For example:

HClO3 + HCl <----> HClO2 + HOCl

HClO3 + HClO2 ---> 2 ClO2 + Cl2 + 2 H2O

Transforming into the less soluble KClO3 and extracting allows the formation of a more stable HClO3 solution (at least, I hope).

Sorry, but I am not exactly clear on the Sulfate salt being reacted with the H2C2O4 to make H2SO4 (MgSO4, Na2SO4, CaSO4,...). Also, do not boil an aqueous solution of H2C2O4 as you get H2 and CO2, possibly reversing the reaction as the Oxalic acid decomposes.

[Edited on 5-3-2012 by AJKOER]

MgSO4 is used in the reaction with oxalic acid, and there is an excess of it. Maybe the crystals are precipated MgSO4 hydrate. I get your idea of using KClO3 now, but it would not work because it is less soluble than potassium oxalate, thus if any H+ and ClO3- ions combine, the precipation of KClO3 will push the reaction backwards. How about extracting NaClO3 from boiled down bleach residue by solubility differences (NaClO3 is much more soluble than NaCl, especially in hot water), recrystalize it a few times, then add oxalic acid to it. I would think that it should work much better (sodium oxalate will appear as almost insoluble in a concentrated solution of NaClO3). If you don't want to boil down bleach, try reacting Ca(ClO)2 and Na2CO3 together.

AJKOER - 5-3-2012 at 09:08

Weiming: Good catch, I was surprised by how soluble the Potassium Oxalate is relative to the KClO3.

Possible answer: work with the MgCl2/Mg(ClO3)2 solution, cool and try to separate out the MgCl2. However, I am not too crazy about this route as one mole of HCl (from any residual MgCl2 reacting with H2C2O4) can consume two moles of HClO3 to form ClO2/Cl2.

I have an idea to add soluble Lead (II) acetate ( Pb(C2H3O2)2 ) to the Magnesium chloride/chlorate solution. The Lead chloride created is insoluble:

MgCl2 + Pb(C2H3O2)2 --> Mg(C2H3O2)2 + PbCl2 (s)

Mg(ClO3)2 + Pb(C2H3O2)2 <--> Mg(C2H3O2)2 + Pb(ClO3)2

Adding H2C2O4, we could have chloride free HClO3 (and some Acetic acid):

2 H(+) + C2O4 (-2) + 2 ClO3 (-1) + Pb (+2) --> PbC2O4 (s) + 2 H(+) + 2 ClO3 (-1)

[Edited on 5-3-2012 by AJKOER]

plante1999 - 5-3-2012 at 09:52

Quote: Originally posted by AJKOER

Weiming: Good catch, I was surprised by how soluble the Potassium Oxalate is relative to the KClO3.

Possible answer: work with the MgCl2/Mg(ClO3)2 solution, cool and try to separate out the MgCl2. However, I am not too crazy about this route as one mole of HCl (from any residual MgCl2 reacting with H2C2O4) can consume two moles of HClO3 to form ClO2/Cl2.

I already make HClO3 by the oxalic process:

http://www.sciencemadness.org/talk/viewthread.php?tid=16330&...

I was suprised to see how much HClO3 was oxidizing... It oxidized at RT many thing like paper.

[Edited on 5-3-2012 by plante1999]

AJKOER - 5-3-2012 at 10:54

Plante1999: That's great, but apparently you did start with a chlorate (actually, a good idea as it cheap, pure and available) and hopefully, it was a chloride free salt.

My synthesis is based on an old commercial process (see http://www.sciencemadness.org/talk/viewthread.php?tid=18452 ) that uses the fact that Mg (and also Zn) hypochlorite has the ability to quickly disproportionate into the chlorate, a plus for the home chemist without access to KClO3 or just likes to do it himself.

I am also proposing (perhaps an original idea) the use of Lead (II) acetate to facilitate the creation of hopefully a chloride free HClO3 from a chloride/chlorate solution using Oxalate acid as well.

[Edited on 5-3-2012 by AJKOER]

AJKOER - 31-3-2012 at 06:15

Quote: Originally posted by AJKOER

Having noted the relative insolubility of many oxalates, I was wondering the potential use/feasibility of employing Oxalic Acid to make various other reagents, including H2SO4 and even HNO3. For example, with say FeSO4, the reaction could be:

FeSO4 + H2C2O4 --Heat --> FeC2O4 (s) + H2SO4

I came across a YouTube video on this very reaction in a more dilute mode for the preparation of Iron oxalate:

http://www.youtube.com/watch?v=7QzAmvzgK2M&NR=1&feat...

Note, the suggested synthesis is just to wait about an hour for the Oxalate to form, and no continuous application of heat.

The product is, of course, dilute H2SO4, but does simply reducing the amount of water achieve a successful reaction with the formation of more concentrated H2SO4?

[Edited on 31-3-2012 by AJKOER]

weiming1998 - 31-3-2012 at 16:24

Quote: Originally posted by AJKOER

No, having less water does not produce a more concentrated acid. More of the reactants just doesn't react. The water needs to be in a large enough amount so that both the oxalic acid and the sulfate can dissolve in water. I thought of having less water=more concentrated acid before too, but it just gave me a less and a more contaminated yield. Slow heating gives a much larger yield too. I have a video on the synthesis of sulfuric acid with magnesium sulfate and oxalic acid http://www.youtube.com/watch?v=5cnqWepbVhY&feature=g-upl...

In the video, for the sake of filming, I only heated it for 10 minutes, so my yield is lower than my previous attempt (50 mls instead of just over 100mls of boiled-down acid). You should heat it for longer through, an hour on very gentle heat is probably perfect.

chemicalmixer - 2-4-2012 at 18:26

For some reason your video is unavailable at the moment. You mention that using too much oxalic can ruin the sulfuric acid product, since oxalic is a strong reducing agent, however it seems to me that using an excess of oxalic would be advantageous, since otherwise your H2SO4 will wind up being contaminated by trace amounts of MgSO4. Why not just use an excess of oxalic to ensure that all of the Mg ions are precipitated, and then boil the acid to concentrate it, thereby destroying any excess oxalic in the process?

Also, wouldn't this procedure work better with gypsum, since Ca oxalate is very highly insoluble, while Mg oxalate is slightly soluble? Plaster of paris would be a good and cheap source of gypsum, although it would likely first have to be added to an excess of DH2O in order to convert it from the hemi- or anhydrous form into the dehydrate, before reacting it with the oxalic acid solution.

This would be a great procedure for making H2SO4 from recycled products, if one could utilize old scraps of sheetrock as a gypsum source, along with oxalic acid that's been formed by the electrolytic oxidation of spent ethylene glycol antifreeze.

[Edited on 3-4-2012 by chemicalmixer]

weiming1998 - 3-4-2012 at 01:31

But gypsum, though, is only very slightly soluble (about a few grams, no more than 10, per litre.) while MgSO4 is fairly soluble (around 71g/100ml for the heptahydrate), while Mg oxalate is about 0.1g/100ml. So the Ca2+ ions might not dissolve in the water enough to react with the C2O42-, making yields low. You can try though, and see if that is a better idea.

As for the oxalic acid, it probably will decompose in boiling H2SO4, but it might react with the acid first (hot H2SO4 is a weak oxidizing acid) to form a mix of SO2, H2O and CO2, something like this: H2C2O4+H2SO4===>2CO2+2H2O+SO2, decreasing the yield.

As for MgSO4 contamination, eventually the acid will become so concentrated that CuSO4 placed into the acid doe not change the acid's colour to even a slight blue, so even if there was MgSO4 contamination, it would only be trace, that will hardly affect any reactions, and could be removed by drying the acid completely with a sulfate dessicant.

AJKOER - 3-4-2012 at 09:43

True CaSO4 is not very soluble in water, but may be more soluble, for example, in salt (NaCl) water. If so, the final products would be H2SO4 and HCl. Hopefully, if so desired, the volatile HCl could be removed.
-------------------------------------

On another topic, I have been investigating the property of the Oxalate salts created. In particular, the heating of Iron Oxalate produces a very fine pyrophoric Iron powder.

See http://www.youtube.com/watch?v=_2HHuUMkg58

Silver Oxalate (the dry salt is a heat, shock and friction sensitive explosive) is, in a suitable medium, capable of forming a fine Silver particle upon decomposition. In fact, per "Thermal decomposition as route for silver nanoparticles" by S. Navaladian, B. Viswanathan, R. P. Viswanath, and T. K. Varadarajan, to quote: "The thermal decomposition of Ag2C2O4 was quicker in ethylene glycol medium than the aqueous medium" with formation of an Ag nanoparticle colloid. The decomposition reaction is given by:

Ag2C2O4 ---------> 2 Ag(s) + CO2 (g)
..................N2; Heat

Source: http://203.199.213.48/50/1/naval.pdf

See also: http://www.youtube.com/watch?v=K-jba4qHBw8 ).

Copper oxalate upon heating in an inert atmosphere liberates Cu and, in air stream for some 2 hours, produces a very fine CuO with a high surface area (useful for catalytic purposes). See: "Copper Oxalate synthesis and decomposition" at http://info.tuwien.ac.at/struchem/files/poster_cuox.pdf

Zinc oxalate behaves a little differently. See "Mechanism and kinetics of thermal decomposition of zinc oxalate" by Barbara Małecka, Ewa Drozdz-Ciesla, Andrzej Małeck. Per the paper: "Thermal decomposition of ZnC2O4 in helium atmosphere in isothermal and non-isothermal conditions was studied. Decomposition of ZnC2O4 is a single-stage reaction with ZnO as final product". From this, my take on the decomposition reaction:

ZnC2O4 ---> ZnO + CO + CO2

Link:
http://144.206.159.178/FT/1034/593171/12218059.pdf

With respect to nickel oxalate, again a very porous and fine particle product. See: "Rapid nickel oxalate thermal decomposition for producing fine porous nickel metal powders: Part 1: Synthesis" by C.S. Carney, C.J. Gump and A.W. Weimer. To quote: "The product powder was characterized as 2–40 μm diameter microcontainer particles comprised of nano-sized nickel primary particles with diameters of 20–70 nm. The production of porous elemental nickel powder via the aerosol flow thermal decomposition of nickeloxalate results in powder with acceptable electrical properties and is more benign and potentially cheaper than the current commercial process for nickel powder production."

Link: http://www.sciencedirect.com/science/article/pii/S0921509306...

Here are the details on the decomposition of calcium oxalate monohydrate. Apparently, three mass losses are observed:

- CaC2O4.H2O --> CaC2O4 + H2O
- CaC2O4 --> CaCO3 + CO
- CaCO3 --> CaO + CO2

http://www.setaram.com/files/application_notes/AN134-Decompo...

Sodium oxalate appears to behave similarly forming Na2CO3.

With respect to Magnesium oxalate, see "Magnesium oxide nanocrystals via thermal decomposition of magnesium oxalate", by Fatemeh Mohandesa, Fatemeh Davarb and Masoud Salavati-Niasari.

Link:
http://www.sciencedirect.com/science/article/pii/S0022369710...

With respect to Lead oxalate, the thermal decomposition is given by:

3 Pb(OOC)2 ---> 2 PbO + Pb + 4 CO2 + 2 CO

Source: http://pubs.rsc.org/en/content/articlelanding/1948/jr/jr9480...

[Edited on 3-4-2012 by AJKOER]

chemicalmixer - 3-4-2012 at 11:45

Quote: Originally posted by weiming1998

But gypsum, though, is only very slightly soluble (about a few grams, no more than 10, per litre.) while MgSO4 is fairly soluble (around 71g/100ml for the heptahydrate), while Mg oxalate is about 0.1g/100ml. So the Ca2+ ions might not dissolve in the water enough to react with the C2O42-, making yields low. You can try though, and see if that is a better idea.

Even though the solubility of gypsum is low, the solubility of CaC2O4 is much lower, so, if given enough time to react, nearly all of the Ca should eventually precipitate as CaC2O4 which can be filtered off, leaving behind dilute H2SO4.

For instance, boiling a slurry of CaSO4 in Na2CO3 solution (CaSO4 + Na2CO3 --> CaCO3 + Na2SO4) will eventually convert all of the gypsum and washing soda to limestone and Glauber's salt, since CaCO3 is much less soluble than CaSO4.

[Edited on 3-4-2012 by chemicalmixer]

weiming1998 - 4-4-2012 at 03:44

Quote: Originally posted by chemicalmixer

Quote: Originally posted by weiming1998

Yes, that is true. Calcium oxalate is much less soluble than CaSO4. I will try the synthesis with CaSO4 now, and see if I could create any sulfuric acid.

[Edited on 4-4-2012 by weiming1998]

AJKOER - 8-4-2012 at 11:06

I just prepared some dilute H3PO4 by adding dry Oxalic acid dihydrate to a hot solution of Tri-Sodium phosphate (commonly known as TSP where it is sold in hardware stores). I have to heat the distilled water to get all the TSP to dissolve. Make sure you perform this reaction in a vessel with excess capacity as the reaction was vigorous with foam which quickly dissipated. Reaction:

2 Na3PO4 + 3 H2C2O4.2H2O --> 3 Na2C2O4 (s) + 2 H3PO4 + 6 H2O

The H2C2O4 should be in excess to avoid the formation of NaHC2O4, sodium hydrogen oxalate. The solution very rapidly clears leaving a thick white deposit of Sodium oxalate, which is easily separated.

------------------------------------

I also reacted a dilute NaClO/NaCl solution with dry H2C2O4.2H2O. Again, the reaction was very vigorous with foam and an obvious presence of Cl2 per the reactions:

H2C2O4.2H2O + 2 NaClO --> Na2C2O4 + 2 HOCl

H2C2O4.2H2O + 2 NaCl ---> Na2C2O4 + 2 HCl

2 HCl + 2 HOCl ---> 2 Cl2 (g) + 2 H2O

or the net reaction:

2 H2C2O4.2H2O + 2 NaClO+ 2 NaCl --> 2 Na2C2O4 + 2 Cl2 (g) + 6 H2O

However, the amount of Na2C2O4 precipitate formed was below expectation. This could be due that some of Oxalate was in solution as NaHC2O4 or there could have been a side reaction involving the direct decomposition of H2C2O4:

H2C2O4.2H2O + NaClO + H2O --> NaCl + 2 CO2 + 4 H2O

This speculation on the side reaction is supported by warnings that H2C2O4 is said to have explosive (or rapid decomposition) reactions with strong oxidizers.

[Edited on 8-4-2012 by AJKOER]

Formatik - 7-5-2012 at 17:44

Quote: Originally posted by Formatik

Copper sulfate and oxalic acid could then work to also form H2SO4.

Sulfuric acid from copper sulfate and oxalic acid:

This works! Few minor complications though.

Concentrated stoichiometric aqueous solutions of copper sulfate and oxalic acid dihydrate yields a very fine light green copper oxalate precipitate and eventually a clear slight yellow and when more concentrated pale green solution. This solution was found to contain H2SO4.

Because the copper oxalate was so fine, it passed through filters. So the yellowish solution was let sit a few hours until clear and siphoned in part. The siphoned solution was evaporated for a few days and the solid removed. Then the solution was boiled down until thick fumes formed (H2SO4).

The boiled down liquid had the color brown and gave a further crystallization, which was a white solid that somewhat stuck together. It seems the white solid can be removed by further siphoning. The warm liquid when dripped onto paper tissue ate holes through it under carbonization.

Sulfuric acid from calcium sulfate and oxalic acid:

Calcium sulfate is also reported to work. There was a thread recently on here somewhere mentioning a reference where calcium sulfate was said to work, though I doubt the claim the acid is of good purity.

But we need to consider that copper sulfate is much more soluble than calcium sulfate and the solubility product of calcium oxalate is 2.32E-9, whereas copper oxalate is 4.43E-10. So copper sulfate is better suited for those reasons. It's possible solubility of both oxalates is increased in acid solution.

Sulfuric acid from magnesium sulfate and oxalic acid:

When I mixed stoichiometric concentrated solutions of oxalic acid and magnesium sulfate, there was no precipitation (no magnesium oxalate formed). After standing for several hours there was crystallization (not sure if that was any reaction, but I suspect unreacted crystallization of reagents), and the liquid was always heavily contaminated with solids. Even after several crystallizations, and finally boiling the liquid down left only solids and no acid. If this route does form any sulfuric acid (which isn't entirely clear to me, it would be much more contaminated than the acid from the above copper route.

AJKOER - 9-5-2012 at 06:08

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 - 10-5-2012 at 21:52

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 970 times

[Edited on 11-5-2012 by Formatik]

AJKOER - 11-5-2012 at 06:58

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 - 21-5-2012 at 12:19

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 - 21-5-2012 at 21:14

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.

pedrovecchio - 24-5-2012 at 23:50

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 - 26-5-2012 at 15:20

Quote: Originally posted by pedrovecchio

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.

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 - 6-8-2012 at 12:11

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 - 7-8-2012 at 04:34

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 - 7-8-2012 at 06:27

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 - 23-8-2012 at 00:01

@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 - 3-10-2012 at 05:03

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 - 3-10-2012 at 06:23

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 - 3-10-2012 at 07:41

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 - 3-10-2012 at 09:14

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 - 3-10-2012 at 12:54

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 - 4-10-2012 at 12:49

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 - 12-7-2013 at 17:45

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

[<a href="u2u.php?action=send&username=bfesser">bfesser</a>: "CN"→calcium nitrate</sub>]

[Edited on 7/13/13 by bfesser]

RingoStarr - 10-8-2013 at 20:08

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 - 18-8-2013 at 16:59

@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 - 18-8-2013 at 23:04

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 - 19-8-2013 at 08:43

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%.