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[*] posted on 1-1-2006 at 23:24


I have a pretty good description of coke making in Shreeve's "Chemical Process Industries." Modern coke ovens use ground up coal, have a wall temperature of about 2000F (~1100C), and cook for 17 hours.

One thing I am not doing that may be important is "quenching." This is done to cool off the coke immediately following the cook. Quenching is usually done with water but can also be done with CO2 or argon if heat recovery is desired. I've been just letting it cool in the furnace. This may be where it sucks in the oxygen as the CO over the coal decreases in volume. I may make another furnace run and add a cold water quench at the end.




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[*] posted on 29-1-2007 at 08:32


CaC2 can be made with charcoal heating, however oxygen is in order. This is made by heating KNO3, KNO2 can be reoxidized by catalytic oxidation of aqueous solution using charcoal as catalyst. With this heating Ca3P2 (white phosphorus can be made by heating charcoal with H3PO4 - obtained from apatite and sulfuric acid - make dihydrogenphosphate, filter, add more H2SO4 and filter again) can also be made, as well as CaSi2.
Quenching coke with water reminds me of water gas - this when burned with O2 also should give a high temperature.

[Edited on 29-1-2007 by Theoretic]




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[*] posted on 29-1-2007 at 10:34


Quote:
Originally posted by Theoreticwhite phosphorus can be made by heating charcoal with H3PO4 - obtained from apatite and sulfuric acid

White phosphorus can also be obtained by heating sand and putrefied urine. Heating white phosphorus in the absence of air should yield red phosphorus, I believe, although I have no idea of the temperature requirements for either of these. Blast furnaces are simple to design and build from local materials like clay provided care is taken to ensure no airbubbles are present.

As for the graphite production, a relatively low tech method is to layer some sand and charcoal followed by another layer of sand. The batch is then heated (usually by induction) which causes the charcoal to evaporate and and recrystalize with the graphite structure. Crude diamonds may also result although I think that is highly unlikely without sophisticated temperature control and elevated pressures. A (usually) undesireable byproduct of this method is silicon carbide which hasn't been mentioned yet so we can add that to the list.

Provided we can locate two metals from our immediate surroundings we can have a battery up and running. Two dissimilar metals (perhaps iron and copper since they are so readily available) separated by an electrolyte solution (salt water or vinegar) form a very crude battery with a voltage potential of about 1V. Multiple batteries may be connected in series to increase the voltage and/or in parallel to increase the amperage. Unfortunately most advanced applications (radio/induction heating) require AC which would need a generator or semiconductors to produce, but the DC batteries combined with the graphite would allow for electrolysis which in turn could yeild chlorates/perchlorates.

The ambitious among us might attempt to obtain silicon from sand and charcoal and find a way to produce a semiconductor so we can play on the computer when we're not smelting ores. Copper oxide(s) are also simple semiconductors if the silicon falls through.

This is a -great- thread. Someone should summarize the important content and have a mod sticky it. :)




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[*] posted on 18-10-2008 at 22:02


I went to a rock show today to see what primordal chemicals I could find for cheap. I picked up the following, all for about $5:

iron pyrite FeS.................3 pieces
galena PbS
chalcopyrite CuFeS2...........2 pieces
hematite Fe2O3
molybdenite MoS2

I don't know how pure my specimens are but the pyrite and galena look almost pure enough to be reagents. Possibly the hematite too.

There were nice specimens of stibnite and bismuthite also but they were too expensive for my purposes. Mostly I just like looking at them but will no doubt do some analyses.

Here's a picture of my loot. The polished hematite sphere is about 2cm in diameter.

loot.jpg - 86kB
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[*] posted on 18-10-2008 at 22:20


Hmm well this is pretty much about chemicals that can be had from the earth, or something simple like that right?

Well, I think Potassium Nitrate can be extracted from decaying organic matter, such as dung by adding Pot ash and hot water then filtering and recrystallizing. I don't believe that was mentioned. I remember reading this a few years ago, and was tempted to try it, now it seems I can't find the site that had the tutorials.

Once you get nitrates, you could reduce them to nitrites with carbon or something like that. Possibly cyanide also? I remember reading that somewhere on the forum.


Then of course, you can get Sodium Chloride, and some other things probably from the seas.
Calcium carbonate (phosphate?) from eggshells or bones and Malonic acid in relatively high concentrations from beetroots.
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[*] posted on 21-10-2008 at 17:13


I can't seem to edit my last post, so I appoligize for the double post, but
I found a good site about extracting Nitre from organics: http://docsouth.unc.edu/imls/lecontesalt/leconte.html I have also posted it in another thread, but I figured it also went well with this topic.

Another thing is it is not the malonic acid that is found in beetroots, but the calcium salt.

One thing I wanted to say was also that it seems like sulfides are found in ash. I added a bit of Vinegar to some ash from my fire place, and for a moment or two noted the distinct smell of Hydrogen Sulfide. It after smelt very different, much like a perfume which I found to be really weird. Any explanation for this? I think perhaps the ligin in the wood being burnt could have vaporized, partially oxidized and then got absorbed into the ash. The CO2 bubbles from the Carbonates reacting with the Vinegar, or the vinegar replacing some of the organic acids could have brought out the smell. - sorry, that is a bit off topic, but I thought it was interesting.
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[*] posted on 21-10-2008 at 18:09


Vinegar has some interesting things in it, as it is. Apparently the ethanol is what's distilled, so there may be some congeners from that, and there's whatever's left from the acetobacter doing its thing, which is going to have its own effect.

I don't think you'll get any organics from ash (aside from elemental carbon, carbonate and, as you observed, maybe a few ppm sulfurousness), at least if it was well burned.

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[*] posted on 22-10-2008 at 16:43


Hmm, well vinegar + baking soda doesn't produce the smell like the ash + vinegar does. It also smelt very organic. I remember I enjoyed the smell quite a bit.

I guess I'll have to do it again sometime to make sure its coming from the ash, because it does seem weird that there would be organics in it like I seem to think.
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[*] posted on 22-10-2008 at 19:07


Quote:
Originally posted by neutrino
One I accidentally discovered: mixing calcium chloride and magnesium sulfate solutions, then heating the resulting paste to decomposition releases HCl. It goes something like this:

Ca<sup>2+</sup><sub></sub> + SO<sub>4</sub><sup>2-</sup> --> CaSO<sub>4</sub>

Drying, we get solid MgCl<sub>2</sub> . 6H<sub>2</sub>O. Finally,

2MgCl<sub>2</sub> . 6H<sub>2</sub>O --heat --> Mg<sub>2</sub>OCl<sub>2</sub> + H<sub>2</sub>O + HCl


I've recently done this using NaCl. The decomposition equation given in the Handbook of Inorganic Chemicals by Pradyot Patnaik forming also the basic salt is: MgCl2.6 H2O -> Mg(OH)Cl + HCl + 5 H2O.

I mixed powdered MgSO4.7 H2O and NaCl into a paste using water. And then heated them on a hotplate, after the water evaporated, and heating continued significant amounts of HCl evolved, also recognized by red litmus and the irritating odor.

But anyone have ideas on how to further work up the basic salt and use up the other Cl?
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[*] posted on 22-10-2008 at 19:29


Primordial atmosphere + lightning = amino acids, HCN, and other good stuff.

Native metals (Au, Hg, Cu, Pt, etc.).

NaCl, straight from the ground (or sea) in very high purity.

Sucrose (glucose and fructose) from the cane or beet, also in high purity. (see also honey)

Tartaric, malic, citric and aconitic acids.

Borax.

Glycerol.

Cis and trans polyisoprene.

Styrene (from styrax).

Lime from calcined limestone (maybe previously mentioned)

Cheers,

O3




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[*] posted on 22-10-2008 at 19:39


Also if we are looking into plant as sources as organic molecules, theres a huge number of possibilities. Everything from alkaloids, organic acids, etc.
It would be kinda interesting to see a list of all the plants that contain useful amounts of various chemicals, although that would take up a large amount of room and probably be best in its own thread :P

Quote:
Originally posted by Formatik
But anyone have ideas on how to further work up the basic salt and use up the other Cl?


Wouldn't just getting it hotter, or getting it wet again work?

Seems like these to reactions could happen:
Mg(OH)Cl <=> MgO + HCl

2 Mg(OH)Cl <=> Mg(OH)2 + MgCl2

.. which then the MgCl2 could react with water to produce the HCl and Mg(OH)Cl again. To me it seems like the first reaction more likely to happen.

Although over all it seems like a good waist of time and energy to get the last Cl out, simply because NaCl is so easy to get. :P

[Edited on 22-10-2008 by kclo4]
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[*] posted on 22-10-2008 at 21:48


There is a type of cement made using MgO and MgCl2, which react to form the oxychloride. See http://en.wikipedia.org/wiki/Sorel_cement

Heating the oxychloride drives off HCl, leaving MgO, can't do that in solution.
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thumbup.gif posted on 22-10-2008 at 22:13


Roasting it further is also what I had in mind and looks to be the best answer so far, but the question I've got is wether the necessary temperature is in the lower range like the reaction I've done above, or not.

Here we are. According to USP5243098, which references Kirk-Othmer for thermal kinetics and mechanism of decomposition of the magnesium chloride hydrates:

MgCl2.6 H2O <- -> MgCl2.4 H2O + 2 H2O at 95-115ºC;
MgCl2.4 H2O <- -> Mg(OH)Cl + HCl + 3 H2O at 135-180º (missing a Cl in their equation);
MgCl2.H2O <- -> Mg(OH)Cl + HCl at 186-230º;
Then: Mg(OH)Cl <- -> MgO + HCl at 230º.

[Edited on 22-10-2008 by Formatik]
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[*] posted on 23-10-2008 at 15:33


Honestly I would try the lowest energy route possible since HCl is usually readily available. I would collect all vapors in a receiver and use this for making chlorides as you need them. One modification would might be "neater" would be magnesium sulfate with potassium chloride in solution to make a slurry of potassium sulfate and adding alcohol to further precipitate potassium sulfate. The resulting alcohol magnesium chloride could be distilled to reclaim alcohol and obtain the hexahydrate for your HCl exploration. The potassium sulfate is usefull for making glass, alum, and as a hydroponic ingredient.



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[*] posted on 29-10-2008 at 03:21


I didn’t think of boiling off an alcoholic solution, that could be a good idea. I was also originally thinking of isolation of the MgCl2, though later separation of the roast mixture is not too hard, since MgO is nearly insoluble in water, where warm water could be used with to extract the sulfate.

I’ve been reading a bit in the older Gmelin and looking at non-electrolytic methods of generating Cl2 which don't use the oxidation of HCl acid, and they mention glowing a mixture of salt, MnO2 and single hydrated MgSO4 gives chlorine: 2 MgSO4 + MnO2 + 4 NaCl = 2 Na2SO4 + 2 MgO + MnCl2 + Cl2. Described in Compt. Rend. 41, 95; Ann. Pharm. 96, 104, or using a mixture of 2 At* MgCl2 to 1 At. MnO2, where Cl2 escapes and MgO and MnCl2 are left over (Chem. Centr. 1863, 254; J.B. 1862, 659).

By roasting of iron vitriol and table salt in air in glowing heat (J. Pharm. [3] 17,443; J.B. 1850, 273); or using a mixture of pyrite, table salt, and iron oxide (Dingl. 173, 129; Techn. J.B. 1864, 153 and 171). J. and W. Allen heat pyrite mixed with salt in apparatuses which allow for precise control of air inflow, where Cl2 escapes and useable Na2SO4 is also left behind (Chem. Centrl. 1871, 249).

Heating copper chloride up to the point where it glows red, where 3 At. yields 2 At. Cl2 (Dingl. 136, 237; 162, 448; Techn. J.B. 1861, 177). Laurens mixes crystallized copper chloride with 1/2 part sand, this is then dried and heated to 250 to 300º, where Cl2 forms, the residue then has some HCl acid added and is then let sit in air, where it converts again to copper chloride (Repert. chim. appl. 3, 110; J.B. 1861, 898).

*At. likely just means parts.
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[*] posted on 4-11-2008 at 12:23


Bases:

KOH:

1 mol. K2CO3 is heated to red glowing with 2 mol. iron oxide (silicic acid-free) in iron retorts, and then the residue, which is called a "potassium ferride" (compound of potassium oxide with iron oxide) is extracted with hot water under stirring to form also again the iron oxide (DE 21593). The same happens here if the K2CO3 is substituted with Na2CO3.

From lime: in an iron or silver vessel, 1 part K2CO3 in 12 parts H2O* is brought to boiling, and then quicklime slurry is added until the filtered liquid doesn't effervesce with excess acid, this usually needs about 2/3 parts lime from the lime slurry. Then let the solution sit in well covered vessel, remove the lye solution from the residues, and then boil it off in a silver vessel until the remaining oily KOH as a whole begins to entierly evaporate as white fog. More KOH can be extracted from the residues with H2O. In this procedure, unsolubilized lime powder doesn' interact properly and goes without effect. With insufficient water, the -CO3 is removed only partly, with 4 parts H2O to 1 part K2CO3 not at all; where actually conc. KOH will remove the -CO3 from CaCO3 (Gmelin, 7 Aufl., Bd. II, 13-14). *The new Gmelin also emphasizes that a specifically 12% K2CO3 solution brings the highest yield. I've tried this reaction with aq K2CO3 and a clear Ca(OH)2 solution and got the expected CaCO3 ppt. For the filtering, one could wait until the CO3 settles and then decant most of it, or use glass filtering since KOH attacks paper.

The CaO can be obtained by roasting CaCO3 (limestone, marble, chalk, etc.) which decomposes @ ~800 deg.: CaCO3 = CO2 + CaO, and K2CO3 from extraction procedures involving the workup of burned wood or plant ashes.

NH3:

A good number of organic nitrogenous compounds break down when heated, esp. with alkalis to form NH3, some also organic bases. In Beilstein it's mentioned that aq KOH and urea give no NH3 in the cold, but on boiling with alkalis or acids, it eventually decomposes to CO2 and NH3. Annalen 123 [1862], 77 briefly mentions that boiling urea with aq KOH forms K2CO3 and NH3.

There are some other sources saying heating with Na2CO3 solution works also, like the Chemical News and Journal of Industrial Science, 61 [1890] 79, where a dilute aq urea solution with a few cc of strong Na2CO3 aq is then distilled, forming free NH3.

I've mixed an aq urea solution with aq NaOH, and these didn't begin forming any NH3 until after a while of boiling, and the formation seems only gradual. I've also done it with hot aq conc K2CO3 and conc urea solutions, heating these to boiling quickly forms noticeable NH3, distilling this mixture into a cooled receiver containing some H2O yielded a moderatly strong NH3 smelling solution, the vapors of which when the solution was cool gave NH4Cl fog with HCl acid vapors.
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[*] posted on 10-11-2008 at 01:26


I'm looking for the ancient way (Hilaire Rouelle in 1773) of retrieving urea from urine, I found an article on how he did it but it only says he boiled urine dry*, obviously this is not the whole story.

The problem with theoretically precipitating the urea is that other salts might come with it, such as the ammonium sodium hydrogen phosphate that they used to prepare white phosphorus ages ago with.
As I'm looking for relatively pure urea I was thinking of nitrating the dried urine with HNO3
and then basifying the urea nitrate with NaOH soln... I've read that the average piss contains about 30g of urea and would love to know how to retrieve this.

*http://www.experiencefestival.com/a/Urine/id/578774

Unfortunately I do not have a means of experimenting as I am living away from home at the moment.
I am quite surprised that there isn't more information about Rouelle's method on the net.

By the way I love this thread and thought I might add, I've read that Phenol can be distilled from heating equal parts CaO and salicylic acid in the Golden book of chemistry.
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[*] posted on 10-11-2008 at 23:25


There was this thread on extracting urea from urine some time ago: https://sciencemadness.org/talk/viewthread.php?tid=2415 It can be done to form it from extraction from the nitrate or oxalate. No need for hydroxides. Beilstein also mentions these in the 3 Aufl., Bd. I, 1290-91.
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[*] posted on 11-11-2008 at 17:20


I'm sorry Farmatik but that link does not work.
I did try and use the search tab before posting my question but I was unable to find any threads that describe a process.

I also tried to Google ''site:sciencemadness.org urine'' and many of the links on the first page of hits seemed to be broken too.

As a side note related to my first post, Do you think the production of Carbolic acid from the distillation of Salicylic acid and Calcium oxide is practical, that is, do you think it might rival the coal process for small scale manufacture?
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[*] posted on 12-11-2008 at 15:25


Quote:
Originally posted by Liedenfrost
I'm sorry Farmatik but that link does not work.
I did try and use the search tab before posting my question but I was unable to find any threads that describe a process.

I also tried to Google ''site:sciencemadness.org urine'' and many of the links on the first page of hits seemed to be broken too.

As a side note related to my first post, Do you think the production of Carbolic acid from the distillation of Salicylic acid and Calcium oxide is practical, that is, do you think it might rival the coal process for small scale manufacture?


You probably are having the same problem as me, just remove the s in https - then it should allow you to see the site :)
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[*] posted on 15-11-2008 at 00:27


Quote:
Originally posted by Liedenfrost Do you think the production of Carbolic acid from the distillation of Salicylic acid and Calcium oxide is practical, that is, do you think it might rival the coal process for small scale manufacture?


This is the first time I've heard of CaO and salicylic acid. Beil. II 952 mentions that heating salicylic acid with 3 or more moles KOH to 250 deg. doesn’t change the acid. With 4 moles KOH there is at 300 deg. a partial decomposition of salicylic acid to CO2 and phenol, but with 6 moles KOH the acid remains unchanged at even 300 deg. (Ost, J.pr. [2] 11, 392). Heating 1 mol salicylic acid with 6-7 moles NaOH to 300 deg. decomposes most of the acid into CO2 and phenol, using 4 moles NaOH the decomposition is nearly complete. However, with 8 moles NaOH most of the acid remains unchanged (Ost). Distilling calcium salicylate Ca(C7H5O3)2, gives besides phenol in the distillate, also small amounts diphenyl oxide (Goldschmiedt, Herzig, M. 3, 133). I don’t know how these compare to the coal process.
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[*] posted on 18-11-2008 at 20:10


Quote:
Originally posted by Formatik
Distilling calcium salicylate Ca(C7H5O3)2, gives besides phenol in the distillate, also small amounts diphenyl oxide (Goldschmiedt, Herzig, M. 3, 133)


Diphenyl Oxide/phenoxybenzene is insoluble in water* but phenol is quite soluble.
I hope I'm not wrong in saying that distilling into a container of dH2O would therefore easily separate the two and make for quite pure phenol.
* http://www.thegoodscentscompany.com/data/rw1004531.html

Not important mentions that copper ions help some decarboxylations in the following thread, it would indeed be interesting to see if copper(I or II) oxide helped the yield in making phenol.

http://www.sciencemadness.org/talk/viewthread.php?tid=6282&a...




With the fear that I am derailing this thread I have added what may be a chem-lite method of producing ammonia gas.

As discussed in the following link Sweat contains urease which is a enzyme that catalyzes(under moderately basic conditions) the break down of urea into CO2 and NH3
http://www.crscientific.com/experiment2.html

As this is a wise old forum this has been discussed before with the usual addition of extra urea (attained from ones urine as discussed above perhaps)
Since the gym sock is full of urease enzymes one could produce quite a lot of ammonia without heating.
The generic pathway to ammonia from urea is fairly straight forward but I thought that posting the biological method would be an addition to the thread.

Also, I will regurgitate an old thread discussion that may be suitable for the Home Chemist to turn primordial substances into useful chemicals.

DeAFX gives a good account of Tannic acid --> Gallic acid
http://www.sciencemadness.org/talk/viewthread.php?tid=4051&a...

Although I don't know how the following reaction is supposed to occur, The 1911 Encyclopedia Britannica claims that Gallic can also be ''produced by heating an aqueous solution of di-iodosalicylic acid with excess of alkaline carbonate, by acting on dibromosalicylic acid with moist silver oxide,''

http://www.1911encyclopedia.org/Gallic_Acid


Of course one would be only making gallic acid to... treat your psoriasis and external haemorrhoids :o

In conclusion, some real world experimentation is in order.





[Edited on 19-11-2008 by Liedenfrost]
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[*] posted on 24-11-2008 at 21:13


Since we are on the topic of carboxylic acids:


Oxalic acid: There are ways to prepare oxalic acid from wood. Procedure (Beil., 3 Aufl. Bd. I, 639): the same parts of wood cuttings (wooden chippings from brown coal or wood shavings) with KOH and NaOH are heated to 240-250 deg., then later extract with water, and evaporate to sp. gr. of 1.35. Upon cooling, sodium oxalate crystallizes, meanwhile all potassium stays in solution as potash. The sodium salt is boiled with caustic lime, and then the calcium oxalate decomposed with H2SO4. If in this procedure, only NaOH is used without the KOH, a lot less oxalic acid is obtained (Possoz, J. 1858, 242; Thorn, J.pr. [2] 8, 182). Aq. oxalic acid can be boiled down and then largely crystallized by freezing. The preparation of oxalic acid from wood has also been discussed in this forum: http://sciencemadness.org/talk/viewthread.php?tid=1859 In this thread another member also shared their experience, it could be better to use sawdust. I've done the method using moderate HNO3 and sugar, this is also pretty basic but it makes a very large amount of NO2 gases.

Formic acid: Formates: by oxidation of starch, sugar, or albuminates with MnO2 and dilute H2SO4 (Döbereiner, Gilbert's Ann. 71, 107; A. 3, 144; C.Gmelin, P. 16,55). Moist KOH absorbs CO at 100 deg. forming potassium formate (Berthelot, A. 97, 125). At 190-220 deg. moist CO is absorbed lively from NaOH + Ca(OH)2 forming formate. Above 220 deg. decomposition of the formed formic acid occurs, giving H2 and carbonate (Merz, Tibirica, B. 13, 23; Fröhlich, Geuther, A. 202, 317).Formation by reduction of carbonic acid: ... by adding zinc and zinc carbonate into hot KOH solution (Maly, A. 135, 119), etc. - Beil., I, 3 Auf., 393. Formic acid: obtained in large amounts by heating glycerin with oxalic acid (Berthelot, A. 98, 139): C3H5(OH)3 + C2H2O4 = C3H5(OH)2(CHO2) + CO2 + H2O = CH2O2 + C3H5(OH)3 + CO2. Prep.: in a retort the same parts oxalic acid and syrupy glycerin (or better mannite) are heated on a waterbath (Lorin, J. 1870, 644; 1875, 505): If the CO2 evolution becomes less, so a new amount of oxalic acid is added (since glycerin replenishes), etc. until one eventually obtains a distillate of formic acid of 55%. If anhydrous oxalic acid is used, the acid which distills over is about 75% in content. One solubilizes anhydrous oxalic acid in the warm 75% pure acid, then after cooling, decant the liquid and then distill the acid. A 99% pure acid is obtained by neutralizing the acid with NaHCO3 and heating the dry sodium salt with equal amounts of anhydrous oxalic acid in a waterbath (Lorin, Z. 1865, 692; A.ch. [4] 29, 367; Bl. 25, 520; 37, 104). For concentrating, the commercial acid is distilled with H2SO4 in a vacuum at the highest 75 deg. (Maquenne, Bl. 50, 662). A small part of it remains with the H2SO4. Detailed procedure for formic acid preparation: http://www.erowid.org/archive/rhodium/chemistry/formic.acid....

Acetic acid: The aqueous liquid made by distilling wood is poured off from the tar and distilled. The distillate is saturated with CaO. Then this is distilled to get the methanol. The caclium acetate is evaporated until solid, weakly roasted to destroy mixed-in resinous material. Then distilled with H2SO4 or HCl. For preparation of anhydrous acetic acid, dry sodium acetate is distilled with H2SO4. Lime tree, willow, white beech get the most (6.1-6.3%), firs, spruce, pine the least (2.4-2.8%) acetic acid. Hardwoods get more acetic acid than coniferous woods. Log wood more acid than wood from the branches, and the latter more than the outer edges. Fast decomposition of the wood reduces the yield of acid considerably (Senff, B. 18, 65). Also: Völckel, A. 86, 66. From the raw pyroligneous acid there is very little formic acid, but actually more propionic acid and higher homologues (up to capronic acid) in smaller amounts present. If this acid is bound with NaHCO3, then the mentioned acids remain in the mother liquor of the sodium acetate (Z. 1869, 445). n-butyric acid, n-valeric acid, and two crotonic acids and an acid C5H8O2 have been obtained this way also (B. 11, 1356). Vinegar made from wine also mentioned, normal alcohol forms no acetic acid, but occurs there where germs can propagate, there is where acetic fermentation occurs (Pasteur, . 1861, 726; 1862, 475). For forming of the Mycoderma (the converting bacteria) phosphates are needed (of K, Mg and NH4). If these are absent, like in pure ethanol, then no fungus growth or acetic fermentation is possible. Fermentation only occurs at the surface of the liq.etc. - Beil. I, 3. Auf., 398.

Citric acid: Beil., I 835: It occurs commonly in fruits, roots, leaves, etc. Prep.: one allows lemon juice to ferment, saturate it with lime, heat the solution to boiling, filter it boiling hot, then react the precipitated calcium citrate with H2SO4. 100 parts of lemons yields 5.5 parts citric acid. Detailed method starting from lemon juice and chalk (7.56 L lemon juice got about 453 g citric acid, 5%): http://www.erowid.org/archive/rhodium/chemistry/citricacid.t...

Benzoic acid: hippuric acid (benzoyl-glycocine) or the urine of cows or horses which contains this acid is boiled with strong HCl acid: CH2(NH.C7H5O).CO2H + H2O = CH2(NH2).CO2H + C7H5O.OH. This reference says that this was applied to large scale production of benzoic acid. They also described preparation from the bark of Styrax Benzoin. Horse, cow, sheeps urine, buh - Beil II, 1182 talks about hippuric acid: field-fed sheeps produce about 30 g per day, a human about 1g per day, with partially vegetable nutrition: 2.5g per day. A reason why hippuric acid occurs higher in these animals is because they are herbivores and consume foods that are rich in the phenolic substances, the content increases by consumption of foods, especially fruits, esp. cranberries, yellow plumbs, prunes, reine-claudes, cinnamon, etc. The procedure is also only briefly described in II, 1137: urine from horses, cows, etc. is evaporated away to ½ to 1/3 of the volume, filtered and then mixed with HCl acid, then the precipitated hippuric acid which forms after standing in the cool for a while is filtered, mixed with HCl acid and boiled for ¼ hour. But now to find a better source for hippuric acid. The separation is described here: 1 part of hippuric acid is boiled for 30 min with 4 parts conc. HCl, the glycine hydrochloride and benzoic acid forms, addition of water then causes greater portion of benzoic acid to precipitate and is then filtered. Mitscherlich also produced benzene by distilling benzoic acid (obtained through the one plants' gum) with CaO.

Lactic acid: Formation: by boiling of glucose with NaOH (B. 4, 346). By heating sucrose with Ba(OH)2 to 150 deg. (Bl. 25, 189). By heating lactose with KOH to 40 deg. (J.pr. [2]24, 503). By the melting of glycerol with KOH (B. 11, 1167), etc. Prep.: 6 parts of sucrose sugar are mixed with 1/144 parts tartaric acid and dissolved in 35 parts boiling water. After 2 days, one adds 1/18 parts of rotten German hand cheese, 8 parts soured milk and 2.5 pts ZnO is added. The mixture remains at 40 to 45 deg., for 8-10 days, under a lot of stirring. Then everything is heated to boiling and then filtered, the precipitated zinc salt is recrystallized. It is decomposed with H2S, and then the free lactic acid is separated from the mixed-in mannite through shaking with ether (Bensch, A. 61, 174; Lautemann, A. 113, 242). Another way: a mixture of 500 g sucrose sugar, 250 g H2O, and 10 ccm H2SO4 (3 parts H2SO4, 4 parts H2O) is heated to 50 deg. for 3 hrs, then let cool, and under cooling, 50 ccm portions of a total 400 ccm 50% aq NaOH solution is added. The liquid gets heated to 60-70 deg. up until it does not reduce Fehlings solution. Then it is cooled, and then the H2SO4 necessary to neutralize the NaOH is added (3 pts H2SO4, 4 pts H2O). By cooling, shaking, addition of Glauber's salt crystals, the precipitation of Na2SO4 crystals is accelerated. After 12-24 hrs, the Na2SO4 is precipitated with alcohol (93%), then half of the liquid is saturated with ZnCO3, it is filtered boiling hot and the other half of the alcoholic solution is added to the filtrate. After 36 hrs standing, the precipitated zinc lactate is filtered off (Kiliani, B. 15, 699, vgl. B. 15, 136). - Beil, I, 552.

n-butyric acid: occurrence: bound on glycerin in cow butter. Butter contains 2% in the form of butyric acid glycerin, by rancid butter, part of the acid becomes free. Also present in pyroligneous acid, sweat, etc. Prep. by fermentation of calcium lactate: 5 kg of rice or starch are boiled with 60 L H2O for a few hours. After cooling, 60g malt and 2 L milk are added, 1 kg of finely cut meat, and 2 kg chalk are added. The mixture is let stand for several weeks at 25 to 30 deg. under intermittent stirring (Grillone, A. 165, 127). If the gas evolution has stopped, it is filtered, and the filtrate is boiled. Calcium butyrate precipitates, but the acetate and capronate stay in solution. The calcium salt is filtered hot, and decomposed with conc. HCl, or the solution of the calcium salt is precipitated with soda, the filtrate evaporate, the sodium salt is decomposed with H2SO4, and raw butyric acid is fractioned. For purifying, the butyric acid is solubilized in H2O, filtered from the oily capronic acid, neutralized with CaO, and the solution of this salt is evaporated. The precipitated Ca salt is removed, and decomposed with HCl acid (Lieben, Rossi, A. 158, 146). Another: 100g of potato starch has 2 L of 40 deg. H2O poured over it, 0.1 g potassium phosphate is added, 0.02 g MgSO4, 1 g NH4Cl, 50 g CaCO3 and a trace of Bacillus subtilis. After 10 days standing at 40 deg., 1 g alcohol, 34.7 g butyric acid, 5.1 g acetic acid, 0.33 g succinic acid is obtained. The Bacillus subtilis is obtained by moving a hand full of hay in 1/4 L H2O and 5 minute long boiling of the strained liquid (Fitz, B. 11, 52). - Beil. I, 421.

A project I'm interested in, is forming succinic acid from oleic acid (which in turn was made from almond oil, olive oil, or butter, etc. by turning it into soap, purifying,etc.), then oxidize with moderate to weak HNO3, as mentioned in the Handbook of Chemistry by Leopold Gmelin here.

[Edited on 24-11-2008 by Formatik]
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[*] posted on 10-4-2009 at 09:32


Quote: Originally posted by Formatik  
Oxalic acid: There are ways to prepare oxalic acid from wood. Procedure (Beil., 3 Aufl. Bd. I, 639): the same parts of wood cuttings (wooden chippings from brown coal or wood shavings) with KOH and NaOH ... The preparation of oxalic acid from wood has also been discussed in this forum: http://sciencemadness.org/talk/viewthread.php?tid=1859


This method also releases CH4 and H2 so it needs good ventilation. There is a more specific procedure here in Fownes' Manual of Chemistry.


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[*] posted on 19-4-2009 at 13:03


I was reading in Gmelin about the alkali peroxygen compounds, to find Na2O2 can be made by glowing NaOH or Na2O in contact with air or oxygen, or by glowing NaNO3 according to Gay-Lussac and Thénard. Also a mixture heated to red glow made of NaNO3 and CaO or MgO at 300-500 deg. absorbs O from air which is lead into it (DE 82982). KO2 can be made by glowing continously the oxide or hydroxide in dry oxygen (H. Davy).

After reading the short extracts, 10g NaOH was put in a stainless steel bottle and then heated to faint-red glow (kept mainly below dark-red glow), and then dry, CO2-free air lead into the bottle. After over 1 hour, the heat was removed and it was let cool. The hot NaOH strongly attacked the stainless steel to form a black and also fluorescent green mass. It's difficult to find a decent vessel to conduct this in, without it being attacked. The only thing I can think of is Pt, but that's not practical or basic.




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