| Pages:
1
2 |
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
Using Oxalic Acid to Make H2SO4,...
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-
NnQJIgcvke7lxmp5G_LJVNe_EKRppLsJ8DSO0Ia7_jUFQ8i_Vv0rJDbFWozUtddlUAU5Hx9bZfp3hZrTLqpO0Jhuh7Tnj3maeZkcxdCNJ&sig=AHIEtbSdavdaF
_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  |
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
International Hazard
Posts: 1468
Registered: 17-7-2011
Location: shores of a solar sea
Member Is Offline
Mood: speeding through time at the rate of 1 second per second
|
|
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
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
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
International Hazard
Posts: 2684
Registered: 28-12-2011
Location: The Known Universe
Member Is Offline
Mood: Molten
|
|
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.
Elements Collected:52/87
Latest Acquired: Cl
Next in Line: Nd
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
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
National Hazard
Posts: 927
Registered: 25-3-2008
Member Is Offline
Mood: equilibrium
|
|
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
National Hazard
Posts: 616
Registered: 13-1-2012
Location: Western Australia
Member Is Offline
Mood: Amphoteric
|
|
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
National Hazard
Posts: 927
Registered: 25-3-2008
Member Is Offline
Mood: equilibrium
|
|
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
National Hazard
Posts: 616
Registered: 13-1-2012
Location: Western Australia
Member Is Offline
Mood: Amphoteric
|
|
| Quote: Originally posted by Formatik |
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. |
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
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
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
National Hazard
Posts: 616
Registered: 13-1-2012
Location: Western Australia
Member Is Offline
Mood: Amphoteric
|
|
| 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
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
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
National Hazard
Posts: 616
Registered: 13-1-2012
Location: Western Australia
Member Is Offline
Mood: Amphoteric
|
|
| 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
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
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
International Hazard
Posts: 1936
Registered: 27-12-2010
Member Is Offline
Mood: Mad as a hatter
|
|
| 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]
I never asked for this.
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
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
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
| 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
National Hazard
Posts: 616
Registered: 13-1-2012
Location: Western Australia
Member Is Offline
Mood: Amphoteric
|
|
| Quote: Originally posted by AJKOER | | 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] |
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
Harmless
Posts: 17
Registered: 19-3-2012
Location: north america
Member Is Offline
Mood: friction sensitive
|
|
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
National Hazard
Posts: 616
Registered: 13-1-2012
Location: Western Australia
Member Is Offline
Mood: Amphoteric
|
|
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
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
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
Harmless
Posts: 17
Registered: 19-3-2012
Location: north america
Member Is Offline
Mood: friction sensitive
|
|
| 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
National Hazard
Posts: 616
Registered: 13-1-2012
Location: Western Australia
Member Is Offline
Mood: Amphoteric
|
|
| Quote: Originally posted by chemicalmixer | | 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] |
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
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
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
National Hazard
Posts: 927
Registered: 25-3-2008
Member Is Offline
Mood: equilibrium
|
|
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.
|
|
|
| Pages:
1
2 |
|
![[*]](images/xpblue/default_icon.gif){kind=icon}
