Sciencemadness Discussion Board - Primordial chemicals

Primordial chemicals

12AX7 - 14-11-2005 at 15:11

I was thinking, what are the most basic, primordial reagents? Like, calcium hypochlorite is a nice oxidizer, but it has to be produced from something, either by chlorinating lime or electrolyzing a lime-containing solution. Going across the periodic table and thinking up various minerals of the elements, I remembered that manganese dioxide (the mineral pyrolusite( releases chlorine, a wonderful oxidizer, when added to hydrochloric acid. This chlorine can then be added to other solutions, lye or lime for example, to give common bleach or the calcium hypochlorite mentioned.

The same thought process can be applied to other chemicals: acids, bases, reducing agents and so on.

Without electrolysis, you can't free HCl from something like NaCl unless you add a stronger acid. The strongest, commonest acid is sulfuric acid, which can be produced by processing other natural minerals (copperas (ferrous sulfate), burning sulfur).

What else can you think of? How many things are made from it?

Try to think in nearly alchemical terms; no electrochemistry, well unless you want to smelt your own zinc and copper, which have different primary minerals besides.

I was having some fun with this thought process, I thought I'd start a discussion on it and see where it goes.

Tim

neutrino - 14-11-2005 at 16:07

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:

Ca2+ + SO42- --> CaSO4

Drying, we get solid MgCl2 . 6H2O. Finally,

2MgCl2 . 6H2O --heat --> Mg2OCl2 + H2O + HCl

Fleaker - 14-11-2005 at 16:26

Cinnabar is one, a reagent practically dug from the earth itself. That's rather primordial...just heat and collect the metallic mercury. (it's the sulfide, so heating in oxygen atmosphere yields sulfur dioxide).

nitroglycol - 14-11-2005 at 16:29

Quote:

Without electrolysis, you can't free HCl from something like NaCl unless you add a stronger acid.

Not quite true actually. Sodium bisulphate is considerably less acidic than HCl (the pKa of the bisulphate ion is 1.99, whereas that of HCl is around -7) but you can make HCl from NaCl and NaHSO4, because HCl comes off as a gas, shifting the equilibrium to the right.

Back on topic, though, it's a good question. KNO3 and NaNO3 are both naturally occuring minerals, so I suspect you could do some good stuff there. Solvents? Ethanol and methanol come to mind; not most hydrocarbons, since they are dependent on the petroleum industry.

[Edited on 15-11-2005 by nitroglycol]

The_Davster - 14-11-2005 at 17:23

Methanol and a few other(insugnificant) chems are produced from the destructive distilation of wood. Assuming copper can be processed from copper ore you can get formaldehyde.
There is also gold.
There is this volcanic lake somewhere that is concentrated sulfuric acid.
K2CO3 with a bit of sodium contamination can be leached from wood ashes.
Nitrates have already been mentioned, and assuming lead can be had from ore nitrites can be made.
I know there is a source for lime in nature, and with K2CO3 from ash, KOH can be had.
Ethanol as mentioned can be gotten from fermentation.

I can likely think of a few more, but for now, just be glad amateur chemistry has not been persecuted to such an extent to make us have to go back to the ways of the alchemists in obtaining our chems.

[Edited on 15-11-2005 by rogue chemist]

I like this thread

chloric1 - 14-11-2005 at 18:32

Well barium was not familiar to alchemist but definately primordial. First dig up natural barite. heat to about 1100 with a coal fire with bellows with fine charcoal. Add dilute hydrochloric acid to collect H2S directly or dissolve in water to get the hydrosulfide of barium. Boil with sodium carbonate solution to precipitate barium carbonate and decant the sodium sulfide. Take your barium carbonate and neutralize with nitric acid and use the nitrate to make barium oxide or peroxide. Or the nitrate can come from barium carbonate and hydrochloric acid and reacting the chloride with sodium nitrate in excess. Recrystallize the barium nitrate.

IrC - 14-11-2005 at 18:33

I think Cinnabar would be rare and hard to find, but nice if you could. Around here are entire mountains of chalk, easy to roast to the oxide, maybe even carbonic acid could be made from it. This whole idea could be a great chapter in Bromic's book project, how to make things from what is around us. You never know, it may be really handy some day, this is the main reason I love the very old chemistry and formula books.

12AX7 - 14-11-2005 at 20:29

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:

Ca2+ + SO42- --> CaSO4

Drying, we get solid MgCl2 . 6H2O. Finally,

2MgCl2 . 6H2O --heat --> Mg2OCl2 + H2O + HCl

Good point: some chlorides decompose a significant amount. Alkaline earth and transistion metal chlorides (and especially aluminum chloride hydrate) decompose, giving off HCl fumes with some amount of H2O, allowing muriatic acid to be distilled directly in low yield.

Sulfuric acid of course is had in higher yield from anhydrous iron, copper, or to a lesser extent zinc, sulfates.

Cinnabar is the primary mercury ore, you just have to find it -- mercury isn't very common in the Earth's crust after all, but just like silver, it's there. Mercury (elemental) and I think the oxide are also present to some extent.

Chalk, limestone and dolomite: all the calcium oxide and carbon dioxide you could hope for, most formations are several feet thick (up to a few hundred) and span thousands of miles. After calcining to yellow heat for an hour or three (big pieces may need days, cleaving into 1" slabs would help here) you're left with CaO and a lot of CO2 out the stack, which can be pumped and compressed, or bubbled into a solution for collection: read up on the Solvay process, which produces CaCl2 as a byproduct for well, given the materials CaCO3 + NaCl, I think you can guess why.

Oh, and let's not forget CaO is also the primordial alkali. Why?

The magnesium ought to be leachable, perhaps by dissolving a block of dolomite and precipitating CaCO3 with MgCO3 (from more dolomite? Dunno) to give seperate Ca and Mg (salt) products.

But this is kind of off the direction... a great conversation as well but I was thinking more what you do with it. (Or maybe I'm wasting my time explaining; this is an interesting enough angle, at least.) Like, my oxidizer example: besides oxygen in air, about the only natural thing you have is MnO2, which can oxidize Cl- to Cl2. The MnCl2 can be precipitated and reoxidized with oxygen and fire or weathering and time (as happens naturally). That Cl2 gas can go on to do just about anything, up to and including things like permanganate (which can also be made from pyrolusite and a caustic fusion, in air), perchlorate, ferrate and so on.

And heck, primordial electrochemistry is worth thinking about, too. It's a lot of effort to mine, roast, smelt, distill and cast zinc anodes, but it can be done. Copper can be mined, roasted and smelted with a bit more ease, though it needs more fire to cast it (not a problem for a firetender such as myself

). Electrolyte, well that can be made from whatever, be it acid, base or a salt. Given enough surface area and a few cells, you can do all the standard electrochemistry, well assuming you can isolate platinum,

and make chlorates, persulfates and the ever most venomous fluorine, as well as the strongest reducers, the alkali metals (which can in fact be isolated by carbothermic processes!).

Organic chemistry of course all starts with organic chemicals, since it's a waste to start with CO2 and there's so much plant and animal life available to the desperate Mad Scientist. Finding the source of reagents (oh my, and glassware!

) for these synthesis might prove an interesting angle however.

Tim

chromium - 15-11-2005 at 02:21

If someone can give way to oxalic acid without nitric acid or NaOH then one can use it instead of sulfuric acid to decompose chlorides or nitrates to volatile acids.

If one has sulfuric acid then magnesium sulfate can easily be got from dolomite.

IIRC ferrocyanides and cyanides can be made from iron oxide and charcoal.

Acetic acid can easily be made by fermentation. When reacted with lime or calcium carbonate it gives calcium acetate that can be converted to acetone.

Lime can be chlorinated to calcium hypochlorite and this with acetone can be used to make chloroform.

nitroglycol - 15-11-2005 at 03:30

Quote:

Originally posted by chromium
If someone can give way to oxalic acid without nitric acid or NaOH then one can use it instead of sulfuric acid to decompose chlorides or nitrates to volatile acids.

IIRC, oxalic acid is found in rhubarb and wood sorrel leaves. I don't know about the concentration, though, or how pure you could get it under "primordial" conditions.

12AX7 - 15-11-2005 at 09:19

Quote:

Originally posted by chromium
If one has sulfuric acid then magnesium sulfate can easily be got from dolomite.

Or bitter wells where this occurs naturally -- hence the origin of epsomite (epsom salts).

Quote:

IIRC ferrocyanides and cyanides can be made from iron oxide and charcoal.

IIRC, urea, charcoal and iron. (Saw a synth on here somewhere.) Urea of course is boiled-down urine.

Tim

Microtek - 15-11-2005 at 10:07

Take a look at cavemanchemistry.com. It has many interesting experiments which could be of use to the stranded-on-a-deserted-island chemist.

edit: It seems that site has become commercialized; the experiments used to be freely availible - now you have to buy the book.

[Edited on 15-11-2005 by Microtek]

S.C. Wack - 15-11-2005 at 11:38

Ferrocyanide was made from leather, hair, blood, etc. as nitrogen source back in the day.
This is very stinky.
Obviously this can be dehydrated and heated in a closed crucible with or without potash to obtain KCN, which can be purified with distilled grain alcohol.

Iodide from burned seaweed.
Tin from SnO2 and coal if tinstone is around.
Sulfur from the local active volcano.
These things, like cinnabar, Chile saltpeter, coal, native copper, etc., are sadly rather nonexistent round here.

From burnt lime or potash heated with aged urine, ammonia. Ammonium carbonate from dry distillation of dried urine. From potash, urine, dirt, and time, KNO3. HNO3 and NH4NO3 from there. From ammonia dried with CaO, dry NH4NO3, and blown glass, a container of anhydrous ammonia in liquid form.

Numerous products that can be isolated fairly easily from large amounts of certain plant material.

Caveman chemistry dude may have beat me to it, not because he started work sooner, but because he settled for less. Don't look behind you.

Destructive distillation of large amounts of organic things is an art unto itself. See "coke manufacture..." in vol. 2, "bone oil" in vol. 1, and "wood, destructive distillation of" in vol. 7 of Thorpe's "A Dictionary of Applied Chemistry".

[Edited on 16-11-2005 by S.C. Wack]

mick - 15-11-2005 at 11:42

In William Prescott's (1796-1859) History of the Conquest of Peru, the Spanish were running out of black powder with no chance of being re-supplied. They made their own by sending an expedition down an active volcano to get the sulfur. I assume a suitable grade of carbon and potassium nitrate would be straight forward for the local expert to manufacture on site in those days.
I doubt it the people who descended the volcano would be very healthy after. I wonder what would have happened if they had not got their muskets working.
mick

chloric1 - 15-11-2005 at 12:44

Well, Mick, Lima would not be a modern megalopolis and there would not be mesitos. And the natives definately would not be speaking spanish. Would the world be better place? We never will know. It would be cool if all of Peru looked like Machu Piccu(sp?)

Sorry off topic. I think saltpeter occurs naturally in caves!

The_Davster - 15-11-2005 at 15:47

Methyl salicilate can be gotten from the wintergreen plant, which with base can be converted to salicylic acid.

SC Wack, do you have any books on destructive distillation of organic materials? I find the idea rather interesting. EDIT: not destructive distillation of reagents, but of natural plants and the like.

And electrolysis of various things should be able to be done as long as one can get metals to assemble the battery. Which in itself opens up a whole other range of compounds which can be made!

[Edited on 16-11-2005 by rogue chemist]

IrC - 15-11-2005 at 18:18

"I think saltpeter occurs naturally in caves!"

Formed underneath piles of bat guanno (crap) by the reaction of the bat droppings with the minerals which make up the walls and floor of the cave, from millions of bats over many years. Caves were one of the major resources in the days of the wooden sailing ships for replenishing black powder stores.

[Edited on 16-11-2005 by IrC]

froot - 16-11-2005 at 13:09

Iron pyrite (FeS2) occurs in abundance in some regions.

Magpie - 16-11-2005 at 14:58

I agree that this making chemicals from nature's raw materials is at the heart of this forum.

I know where there is an inactive stibnite mine not too far from me. Some day I'll revisit it then make my own antimony.

12AX7 - 16-11-2005 at 19:07

Cool. If you find some nice sized and shaped crystals, I'd love to have some in my collection.

Plus maybe some ore for similar purposes

Tim

Magpie - 16-11-2005 at 19:22

12AX7 I'll bring back plenty of ore and you will be welcome to have some. Also the Sb. But don't be in a hurry as I won't be making a special trip to get it. The mine happens to be near some fine trout fishing and that would be my main reason for getting over that way.

Theoretic - 18-11-2005 at 04:08

"Sulfur from the local active volcano."
A true gem.

I can imagine Mr. Wack returning with his pockets full of sulfur, his clothes stained yellow, with multiple holes from lava droplets, a huge grin and 15 m/s (down the slope!), shouting "Sulfur! We have Sulfur! I am a true Mad Scientist!"... then the volcano erupts and further black smouldering holes appear in the rags-that-pass-for-a-labcoat... and he doesn't even notice.
Then for the next week he's bedridden with SO2 poisoning and prohibited from experimenting for that period.
"Iron pyrite (FeS2) occurs in abundance in some regions."
...which can be just left to sit there, and the oxygen and moisture will turn it to sulfuric acid and ferric sulfate:

4FeS2 + 15O2 + 2H2O => 2Fe2(SO4)3 + 2H2SO4

The ferric can be reacted with CuS ore to give copper sulfate:

CuS + 4 Fe2(SO4)3 + 4 H2O => CuSO4 + 8 FeSO4 + 4 H2SO4

Edit: the ferric can also be heated to total anhydrity and distilled to give SO3, this can be reacted with HNO3 obtained using your sulfuric to give nitrosulfuric acid <=> nitryl hydrogensulfate (NHS, yes I spot the pun) and this is a very good nitrating agent, giving higher yields than sulfuric/nitric. And plant oil/animal lard reacted with alkali (soda, potash, KOH, NaOH) gives fatty acids and glycerine. I can foresee NG! Nitrate cellulose with NHS and get NC, this together with NG gives blasting gel.
If one wants gunpowder AN/C is a good one, hygroscopic though. It's roughly as powerful as NG-based propellants, and I assume ANFO would be a better and non-hygroscopic propellant (again, plant oil can be used as fuel).

Or the sulfide could just be left there (as with pyrite, preferably moistened) to oxidize naturaly to CuSO4.

[Edited on 18-11-2005 by Theoretic]

stygian - 20-11-2005 at 18:50

Any ideas on primordial calcium carbide?

The_Davster - 20-11-2005 at 20:10

CaO + C and arc melting with a huge ammount of amps in a graphite crucible. I wonder if an arc welder could do this.....

woelen - 21-11-2005 at 01:18

Quote:

Originally posted by rogue chemist
arc melting with a huge ammount of amps

Primordial

?? I think I learned something new again about prehistoric times

?

But, the idea is nice, although I think it will be quite hard to do this at home.

stygian - 21-11-2005 at 05:57

an arc welder can melt an iron (1538C) electrode quite easily at a modest 100 amps or so. Using a carbon gouging rod and/or high current should easily be able to attain the 2000C needed. But I'm talking primordial here!

12AX7 - 21-11-2005 at 08:54

Yeah humm... no need for a crucible, it'll work fine just arcing inside a pile of the stuff, who cares. Graphite can be mined in a few places (I don't remember what conditions it forms in), but you'll be looking hard to find good enough samples to cut electrodes from. Then you need the insulated copper wire. Insulation as varnish, cotton (or other fiber) wrapping, etc.; copper from ore is simple enough, but then you also need the iron (preferrably tool steel!) dies to draw it through, which needs the high technique of making fine steel, plus drilling and reaming operations to form the die itself. Oh, and silicon steel for the dynamo- no idea where you'd get that without ferrosilicon, so, you'll have to settle with a core about twice as large, made of mild steel, to handle the magnetic field. Which then needs a rolling mill preferrably, which needs more tool steel and machining of course. (Read up on D. Gingery to get a feel for making machines from minimum resources.)

Once you have a dynamo and source of rotational energy (waterfall preferred, but you'll probably already have a low-pressure steam engine from your lathe's first project), you can connect them to your electrode clamps, graphite and blast away at a pile of charcoal and limestone (calcite preferred, since limestone is often actually dolomite and magnesium doesn't do you any good).

Tim

stygian - 21-11-2005 at 13:20

I read that charcoal and coal tar are baked to form electrode for carbide furnaces.

12AX7 - 21-11-2005 at 15:30

Yeah, but artificial graphite needs as much heat as calcium carbide to produce a reasonable product...so....yeah...

Tim

IrC - 21-11-2005 at 18:52

Is it all right with you people if I don't go totally caveman and just buy some items? I know it's cheating but I will keep it quiet if you do.

The_Davster - 21-11-2005 at 19:46

Cheater!! Caveman arc welder is dinosaurs on big hamster wheel hooked up to a generator

.

Are you seriously thinking of making carbide this way? I swear there was a thread on CaC2 somewhere herehere
Looks like Chemoleo has already tried it. Be sure to post results! I want to try this but have no arc welder, I read about how to make one from 8 microwave transformers, so I am looking for some of those now...

stygian - 21-11-2005 at 20:28

*I'm* not. Obviously. Maybe I was misunderstood. What I meant to say was:

"I know arc welders can do it."

Then separately

"Is there some 'primordial way'?"

And if you allow electrochemistry, I found it, it's in the carbide thread.

Magpie - 21-11-2005 at 20:56

Does the formationof CaC2 from CaO and C require DC current and/or a high voltage or just the very high temperature generated by the electric arc? What I am getting at is if AC can be used why bother with the welder? Just buy a circuit breaker within the limitations of your house mains and some large cables.

[Edited on 22-11-2005 by Magpie]

stygian - 21-11-2005 at 21:28

It's just the heat needed. You cant JUST use a circuit breaker because without some kind of ballast you will overload the breaker instantly. Arcing through carbon should provide quite a bit of resistance though. Someone on roguesci a long time ago gave a real simple design. Ac plug, 2 wires, 2 electrodes, one wire having a nichrome coil in series for resistance (e.g. electric heater)

[Edited on 22-11-2005 by stygian]

IrC - 21-11-2005 at 23:46

That is incredibly power wasting though. I would look into cutting a section out of the core and making a way to slowly bring the core back together, this way you are controlling the magnetic field. In the 80's I used to build tesla coils up to 20 or more feet tall, and used pole pigs to power them. Taking an old variable core lincoln arc welder and wiring it as a variable reactor. You have a 240 volt primary, and here you put the primary in series with the pole transformers. A very heavy buss bar dead shorts out the arc welding secondary. In this way you are using controlled magnetic fields to energize the pole transformer from nada to as much amps as your mains supply can handle. Great use for the variable core arc welder and when building something you just take it out of line and put it back to arc welding use so you are out nothing!

Magpie - 22-11-2005 at 12:10

Are you guys saying that for a very crude arc furnace one might install a carbon electrode in a mound of finely divided and mixed CaO/C, locate the 2nd carbon electrode slightly above the mound (to provide an arc) and then as long as you had the appropriately sized inductor in series (or parallel?) you could make CaC2 using 240VAC?

stygian - 22-11-2005 at 13:16

The nichrome element in series is a resistor, effectively turning your 240VAC to a lower voltage, much higher current. I think.

mick - 22-11-2005 at 13:56

Back to the origins. A good quality set of bellows with charcoal might get the temperature up to making calcium carbide.
mick

stygian - 22-11-2005 at 13:59

Speaking of, just found an old roguesci thread that I remember. Someone tells of using a paint stripping heat gun to fire a clay flower pot full of charcoal. Supposedly does a quick job of liquifying cast iron. So, if you could heat air from a bellows somehow..

coke

Magpie - 28-12-2005 at 11:49

I took a crack at making coke yesterday. It was a failure.

I placed some "high volatile bituminous coal" (Utah) in 3 small crucibles each with the lid on. I then heated these crucibles for about 6 hours at 800-900C. Weight reduction was about 50% vs an expected 25-35%. There was a grey crust on the product which I took as inert cinder. Inside the crust the product was blacker but I could not get it to burn even after heating it to red heat in a bunsen burner. I take this to mean that I essentially had no carbon left.

My feeling is that it is a problem of scale in that I can't sufficiently keep my carbon from the oxygen in the atmosphere.

Does anyone have any ideas on what I could do to have success without building a huge oven?

12AX7 - 31-12-2005 at 20:37

I've heard coke can be hard to burn, Idunno. Burns nice in a blast furnace at any rate, but being well above white heat, that isn't very suprising

Next time, you might try heating in a stronger reducing atmosphere -- some charcoal or sawdust on top could help.

Tim

more coke tries

Magpie - 31-12-2005 at 21:16

I made a couple more runs with my coal, this time not grinding it up as I had the first time. I got the temperature up to 1100C for a few hours. Weight loss was about 50%. It looked like the other cinders so I didn't even try to get it started with the bunsen burner.

There is a blacksmith in my area who makes his own charcoal. I'm going to consult with him.

Maybe I need to cover the coal with dirt.

[Edited on 1-1-2006 by Magpie]

IrC - 31-12-2005 at 22:31

Your problem is oxygen. You cannot simply place a lid on and expect that no O2 will get in and burn the carbon. You need to build some type of furnace or crucible to go in one which can completely eliminate the chance of air being sucked in.

After 20 or 30 hours of heat you then quench the coke with water to avoid it burning. The whole process will work much better if you make very large amounts at a time.

http://www.lehigh.edu/~kaf3/cokedata/coking.html

[Edited on 1-1-2006 by IrC]

Magpie - 1-1-2006 at 10:38

Yes, I am now fully convinced that I'm not keeping out the oxygen sufficiently. This is a dramatic demonstration of the effect of scale on certain chemical processes.

I'm doing this more out of curiosity and stubborness than anything. I won't likely go to the trouble of building a large set up just to make some coke. I don't think I'll be blanketing my muffle furnace in argon either.

I am curious as to why it has to be heated for so long, however. The Lehigh write-up says something like 24-36 hours at what, 1100-1200C. Why would it take that long other than some of the coal chunks being very large, say around 1 foot (0.3m) in diameter, and the coal piles also being very large?

[Edited on 1-1-2006 by Magpie]

IrC - 1-1-2006 at 11:15

I'm sure you're right, that the time is related to material quantity. I suppose only experimentation would give you some sort of scaling factor. I do not believe it would be very hard to smooth the surfaces between the crucible and lid to make it air tight, especially if you had a weight with high melting temperature (crucible material?) on top of the lid during the process.

Magpie - 1-1-2006 at 15:18

My procedure was to put a small lump of coal (if the lump of coal is too big it will push off the lid as it swells) in each of 3 crucibles, set them in the hot furnace, and then put the lids on and close the furnace door. Within a minute great quantities of volatiles and smoke come out at the door seal, some of it catching on fire (doing this in a hood or outside is advised).

With the above procedure it would be impossible to seal the lids as those volatiles have to get out. You might be able to just leave the lids off until this rapid volatilization is finished and then place close fitting lids on for the rest of the heat, but I'm skeptical.

As a side note, coal is neat stuff. I'd never really played with it before. A very thin metal-like film forms on the inner surface of the crucible. I'm assuming it's a condensation product. I'm curious as to what this is. If I wanted to know bad enough I could do a series of wet chemical qualitative analyses. This sounds like too much work, however.

12AX7 - 1-1-2006 at 15:40

Probably either glassy carbon or graphite deposited parallel to the surface (hence shiny). Brauer mentions synth with graphite oxide from solution, or cracking alkanes at 800C+. You would have the latter

Curiously, Brauer says a tinge of oxygen helps to make a shiny deposit, but on the other hand, you have a pretty graphite-like material to start with, lots of aromatics and such being distilled, probably forms graphite much easier.

Tim

IrC - 1-1-2006 at 17:39

Put the coal in the crucible, then cast a lid with castable refractory around a ceramic exit tube and one way valve of some sort. After it is hardened at lower temperatures crank up the heat.

Magpie - 1-1-2006 at 19:36

IrC: Along that line you could probably just build a small iron retort out of black iron pipe and lead the tube to a water seal. At T= 1100C I would think that the iron would not suffer.

IrC - 1-1-2006 at 20:04

I started to mention iron earlier as a screw on lid would be great. I decided not to as I just was not sure about the temperatures involved. Then again they do use these big metal furnaces in industry so why not for small scale also.

Magpie - 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.

Theoretic - 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]

Levi - 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.

Magpie - 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

kclo4 - 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.

kclo4 - 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.

12AX7 - 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.

Tim

kclo4 - 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.

Formatik - 22-10-2008 at 19:07

Quote:

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?

Ozone - 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

kclo4 - 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

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.

[Edited on 22-10-2008 by kclo4]

not_important - 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.

Formatik - 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]

chloric1 - 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.

Formatik - 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.

Formatik - 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.

Liedenfrost - 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.

Formatik - 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.

Liedenfrost - 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?

kclo4 - 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

Formatik - 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.

Liedenfrost - 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

In conclusion, some real world experimentation is in order.

[Edited on 19-11-2008 by Liedenfrost]

Formatik - 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]

Formatik - 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.

Attachment: oxalic.pdf (154kB)
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Formatik - 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.

Formatik - 24-5-2009 at 17:16

What all can be done with NaCl?

NaCl when glowed with boric acid and air is let in at the same time gives off Cl2 (Schulze, J.pr. Ch. [2] 21 [1880] 407).

Metallic Fe reacts with NaCl in a vac. at 800 deg. and above: Fe + 2 NaCl = FeCl2 + 2 Na (Hackspill, Ann. Chim. [10] 5 [1926] 228). KCl, KBr, KI also react with Fe at 900 to 1000 deg. forming the Fe-halogenide but yielding little potassium (KF gives a bit better yield), both also volatilize at the same time. Other salts can be reduced to obtain good yield like K2SO4 at 1000 deg. (by-products: Fe2O3, FeS, SO2, O2, after 2hr heating, 80% yield of K, larger filings of Fe can be used). K2CO3 and Fe is an explosion risk since CO is a by-product, this can combine with K to form explosive compounds (so the same thing when reducing K2CO3 with C). See Gmelin.

Melting together S with NaCl forms sodium sulfide and SCl2 (DE49628). This is a bit more complicated, the patent has a schematic where molten sulfur acts on molten alkali chloride, where then both of those mentioned compounds form at the same time. They stress temp. needs to be high enough otherwise you get sulfide and Cl2. The vapor SCl2 is also contaminated with S. For which reason they recommend large excess of alkali chloride.

Cr2O3 or Mn2O3 heated to red glow with NaCl, causes them to decompose when air or O2 is lead into it, forming Cl2, and with water vapor - HCl evolution (Hargreaves, Robinson, E.P. 508 [872]; Ber. 5. [1872] 1064).

Leading steam into a melt of NaCl and silicate, gives the reaction: 2 NaCl + H2O + Na2SiO3 = 2 HCl + Na4SiO4 (DE 382216 [1916]).

The outcome of the action of NaCl on SiO2 depends on the atmosphere, in which the mixture of both is heated according to:

a) 4 x NaCl + y SiO2 + x O2 = 2 x Na2O, y SiO2 + 2 x Cl
b) 2 x NaCl + y SiO2 + x H2O = x Na2O, y SiO2 + 2 x HCl
c) 4 HCl + O2 = 2 H2O + 2 Cl2

In a stream of dry air, only reaction a occurs. In moist air, a, b, and c occur together. In moist nitrogen, only b occurs. In moist air, reaction b is predominant at up to 1000 deg. the extent of the action is small however, but at 600 deg. the reaction becomes noticeable. Increasing the temperature and moisture content increases the extent of the reaction. The deciding factor is the size of the contact surface of the reacting material (Gmelin, Na, Sys. 21, 329).

With water: overheated water vapor will react with NaCl to form NaOH and HCl. The decomposition through water comes to a stop when for 8 moles NaCl, 1 mole NaOH has formed. If water vapor is led over NaCl in a heated platinum crucible, a considerable reaction occurs under 500 deg. The reaction starts clearly around 700 deg., the speed of which increases with the temp. In a porcelain crucible the decomposition isn't quantitative (Gmelin).

Speaking of which, before NaCl becomes heavily restricted, where can one get or extract it? Besides the ocean and salt flats, which might be only ready useful sources if you are near enough.

JohnWW - 24-5-2009 at 17:45

I would think that NaCl, along with KCl, MgSO4, CaSO4, Ca(HCO3)2, and various minor solutes in sea water such as NaF, LiCl, RbCl, CsCl, etc., would become "restricted" or "controlled" chemicals at about the same time as H2O does similarly. This is in view of the fact that most of the world's water is sea-water, containing these in solution. People will be made to pay a huge amount of dollars every time they want a drink of water, which will have an excise tax imposed on it. Access to rivers, lakes, and the sea will become strictly controlled by huge armies of cops, to make sure that no-one helps themselves to free supplies. Rainwater collection will be prohibited.

But I think that other household chemicals, such as starch, cellulose, sucrose, C2H5OH, CH3OH, iso-C3H7OH, NaOCl, Ca(OCl)2, NaHCO3, Na2CO3, citric acid, and tartaric acid, will become "restricted" or "controlled", and subject to huge excise taxes like that already on C2H5OH, well before NaCl and H2O do. Especially the organic substances, because they can be either nitrated to make explosives, or fermented and distilled to make C2H5OH.

[Edited on 25-5-09 by JohnWW]

hissingnoise - 25-5-2009 at 08:01

The Totalitarians are, as we speak, looking with suspicion at air as a source of precursors like O2 (potent oxidiser), nitrogen (NH3) among others, and various acidic oxides.
Large sections of society could be gainfully employed in restricting public access to this ubiquitously dangerous mixture of substances.
Wholesale metering of air, to determine personal usage would be perceived to progress the WOD/WOT alliance.
Known chemophiliacs could be monitored for their lifestyles and their stepping out of line could result in a stint in Gitmo or Australia (once it's depopulated to absorb the influx). . .
Australia here we come!

littlepop - 27-5-2009 at 07:18

Quote: Originally posted by The_Davster

Methyl salicilate can be gotten from the wintergreen plant, which with base can be converted to salicylic acid. [rquote]

Willow bark is where the first salicylic acid came from. I don't know the concentration, but it does help a headache.

[Edited on 5-27-2009 by littlepop]

Formatik - 1-6-2009 at 22:42

Black poplar (Populus nigra L.) leaves contain salicin. Willow tree bark is rich in salicin. Wood sorrel (Oxalis acetosella L.) leaves contain oxalic acid, or also in the mashed leaves of Sorrel (Rumex acetosa L.). Oxalic acid exists in quite a few plants, generally not as a free acid though. The Benzoin Tree (Styrax benzoin) bark is a source of benzoic acid. Mannite in the sweet exude from the stem of Fraxinus Ornus and Fr. rotundfolia. Sodium carbonate from the burned plants near the ocean, potash from land plant ashes. Although plants can be a source of several basic chemical compounds, these are usually more economical to produce by other means (energy costs and isolation from other products). Books like:

The Organic Constituents of Plants and Vegetable Substances and Their Chemical Analysis
http://books.google.com/books?id=wRgAAAAAQAAJ

An Introduction to the Chemistry of Plant Products
http://books.google.com/books?id=nY8HAQAAIAAJ

have more of that kind of information, and if you dig hard enough, you might find more specific analytical and extraction procedures.

Magpie - 29-9-2009 at 18:00

I just aquired the 1927 edition of "Laboratory Experiments in Organic Chemistry," by Roger Adams and J. R. Johnson. It's mostly a repeat of procedures I already have but there are some interesting ones new to me:

1. This is a first year, first semester book but the student gets to make n-valeronitrile using 39g of "powdered sodium cyanide (handle with great care)."

2. Syntheses of fluorescein and eosin dyes.

3. and my favorite: "Hippuric Acid from Urine." The procedure says "before retiring at night, ingest 5g of pure sodium benzoate (or ammonium benzoate)..." "Collect the overnight urine voided the next morning and isolate the hippuric acid as described below."

http://en.wikipedia.org/wiki/Hippuric_acid

Those old time chemists really knew how to party.

chloric1 - 30-9-2009 at 06:28

Quote: Originally posted by Formatik

Quote:

I read almost 20 years ago, probably in Merck index, that salicylic acid is decomposed to phenol by heating. I remember heating it and maybe a third or half sublimed on the sides of the container unchanged and the other portion formed a clear liquid that smelled strongly of phenol

. I did not analyze so the experiment needs repeating. If I do not order salicylic acid then I will have to buy some generic aspirins and extract the salicylic acid. Note the ASA is easily deacetyled but boiling with 10% H2SO4 or so I read.

entropy51 - 30-9-2009 at 07:05

Preparation of phenol by heating salicylic acid and CaO was one of the experiments in the chemistry sets from the 1950's.

Now they'd arrest you for putting those two chemicals in a "toy".

chloric1 - 30-9-2009 at 15:20

Quote: Originally posted by entropy51

Preparation of phenol by heating salicylic acid and CaO was one of the experiments in the chemistry sets from the 1950's.

Now they'd arrest you for putting those two chemicals in a "toy".

A hence freaks like me are hydrolysizing powdered aspirins with battery acid

entropy51 - 30-9-2009 at 15:26

Yes, I think amateur chemists born after 1970 are definitely at a disadvantage.

Probably at least half my chem inventory was bought down at the drugstore in the 1960's. And cheaply too!

[Edited on 1-10-2009 by entropy51]

UnintentionalChaos - 30-9-2009 at 15:44

Quote: Originally posted by chloric1

Quote: Originally posted by entropy51

Preparation of phenol by heating salicylic acid and CaO was one of the experiments in the chemistry sets from the 1950's.

Now they'd arrest you for putting those two chemicals in a "toy".

A hence freaks like me are hydrolysizing powdered aspirins with battery acid

I prefer to use base, as the hydrolysis is very rapid and homogenous. After fairly brief heating, acidify, filter, and recrystallize from boiling water. I did the recrystallization in a metal bowl once, and the liquid that sloshed up the sides got significantly hotter and charred....it reeked of phenol.

"3. and my favorite: "Hippuric Acid from Urine." The procedure says "before retiring at night, ingest 5g of pure sodium benzoate (or ammonium benzoate)..." "Collect the overnight urine voided the next morning and isolate the hippuric acid as described below.""

Is it wrong that I want to try this?

[Edited on 9-30-09 by UnintentionalChaos]

chloric1 - 30-9-2009 at 17:28

Well the hydrolysis of asprin with diluted H2SO4 was to isolate the salicylic acid. The sulfuric acid is a catalyst to break off acetic acid from the aspirin. It can then be heated with CaO. I would presume none of the salicylic acid will sublime out

.

@entropy- Yes I have ALWAYS said that being born after 1970 that I missed out! I was even saying that in the late 1980's when I first started my chemistry explorations. The positive side of this is since the dawn of my hobby I have been looking on how to either buy OTC reagents with out fuss or how to make my own. The stepping up of restrictions during the 1990's was my basic training so to speak. Thats good becuase these days even the very early 1990's seem like a golden era compared to today. Where are going to get a gallon of CCl4 for $20? Or better was a kilogram of PCl5 for $50!! In 1992 this was what was availabe from Chem-Lab supplies.

Wow I totally deraled this thread

Sorry! BTW does the phenol distill from the CaCO3 residue or do you just lixivate with water or alcohol?

watson.fawkes - 30-9-2009 at 17:35

Quote: Originally posted by UnintentionalChaos

"3. and my favorite: "Hippuric Acid from Urine." The procedure says "before retiring at night, ingest 5g of pure sodium benzoate (or ammonium benzoate)..." "Collect the overnight urine voided the next morning and isolate the hippuric acid as described below.""

Is it wrong that I want to try this?

No, I thought that one looked totally great. When you publish pictures of your experiment, though,please omit those of urine collection.

entropy51 - 30-9-2009 at 17:37

Quote:

Is it wrong that I want to try this?

If you use food grade Na benzoate and have normal kidney function it won't be wrong at all. We'd love to hear about it!

Actually I think eating cranberries might have the same effect, IIRC they contain lots of hippuric acid.

@chloric1, if I were to try to collect the phenol I would probably try steam distilling it out. Phenol does steam distill! You would need to need to acidify it before steam distilling, of course.

[Edited on 1-10-2009 by entropy51]

[Edited on 1-10-2009 by entropy51]

Magpie - 30-9-2009 at 17:38

@chloric1: nice to see you using the archaic term "lixivate."

Quote:

Is it wrong that I want to try this?

No, go for it. I feel the human body is an underutilized bioreactor for home chemistry. (Just kidding)

JohnWW - 30-9-2009 at 18:46

Quote: Originally posted by Magpie

I just aquired the 1927 edition of "Laboratory Experiments in Organic Chemistry," by Roger Adams and J. R. Johnson. (cut)

That is the First Edition of the book. It went through at least seven editions, the last appearing to be the Seventh in 1979, after which the authors died.

See http://www.orgsyn.org/obits/johnsonjr.pdf
for the obituary of Prof J R Johnson, 9 Aug. 1900 - 25 May 1983. No edition of the book appears to have ever been scanned and uploaded.

UnintentionalChaos - 30-9-2009 at 18:56

Quote: Originally posted by entropy51

Quote:

Is it wrong that I want to try this?

@chloric1, if I were to try to collect the phenol I would probably try steam distilling it out. Phenol does steam distill!

[Edited on 1-10-2009 by entropy51]

Well, the heating with CaO (Ca(OH)2?) is what achieves decarboxylation. At those temperatures, I suspect it might distill from the mix, much like benzene does from the fusion of benzoate and NaOH. Either that, or it remains trapped as calcium phenoxide. In that case, I would add water and steam distill.

Would you mind listing the extraction procedure, Magpie?

[Edited on 10-1-09 by UnintentionalChaos]

JohnWW - 30-9-2009 at 19:14

Quote: Originally posted by Formatik

(cut)The Organic Constituents of Plants and Vegetable Substances and Their Chemical Analysis
http://books.google.com/books?id=wRgAAAAAQAAJ
An Introduction to the Chemistry of Plant Products
http://books.google.com/books?id=nY8HAQAAIAAJ
(cut)

The first of these books, which has these publication details:
Author: Wittstein, Georg Christoph, 1810-1867; Müller, Ferdinand von, 1826-1896, tr
Publisher: Melbourne, MʼCarron, Bird & Co., 1878
can be downloaded from http://www.archive.org/details/organicconstitu01wittgoog ; the full URLs are:
http://ia351402.us.archive.org/0/items/organicconstitu01witt... 20.2 Mb and
http://ia351402.us.archive.org/0/items/organicconstitu01witt... 12.8 Mb

The second one, which has these publication details:
Author: Haas, Paul, 1877-; Hill, Thomas George, 1876-
Volume: 1
Subject: Botanical chemistry
Publisher: London, New York, Longmans, Green, 1928; and
Author: Haas, Paul, 1877-
Volume: 2
Publisher: London ; New York ; Toronto : Longman's, Green and co, 1929
can be downloaded from:
http://www.archive.org/details/introductiontoch01haas
http://www.archive.org/details/introductiontoch02haas .
The full URLs are:
http://ia311036.us.archive.org/0/items/introductiontoch01haa... 21.7 Mb and
http://ia311036.us.archive.org/0/items/introductiontoch01haa... 10.2 Mb and
http://ia311009.us.archive.org/3/items/introductiontoch02haa... 10.3 Mb and
http://ia311009.us.archive.org/3/items/introductiontoch02haa... 5.1 Mb

Magpie - 30-9-2009 at 19:38

Quote:

Would you mind listing the extraction procedure, Magpie?

Here you go:

"Measure the volume of urine and for every 100 cc. add 25 g. of pulverized ammonium sulfate and 1.5 cc. of cond. sulfuric acid. Stir thoroughly until the ammonium sulfate is dissolved, and set aside for several hours in a cool place or in an ice-bath, until the hippurcic acid has separated as completely as possible. It is advantageous to allow the solution to stand in a cool place overnight, since the hippurcic acid separates slowly.

Filter the crystals with suction and wash with a little ice-water. Transfer the crystals to a beaker and purify by crystallization from about 50 cc. of hot water, with addition of 1-2 g. of decolorizing charcoal. To aid in the removal of colored impurities it is advantageous to boil the hot solution for about 10 minutes with decolorizing charcoal before filtering the solution. Allow the solution to stand for at least an hour, and cool in an ice-bath before filtering the crystals of purified hippurcic acid. Calculate the percentage yield of the hippurcic acid isolated (assuming a complete conversion of the ingested benzoate into hippurcic acid)."

another way to Cl2

Formatik - 14-10-2009 at 16:25

Quote: Originally posted by Formatik

... thermal kinetics and mechanism of decomposition of the magnesium chloride hydrates ...

I thought there was all that was to its basic thermal decomposition. But some time ago I reading Darstellung von Chlor- und SalzsÃ¤ure by Nikodem Caro it was giving some procedures and citing DE51084, which states if MgCl2.6H2O is heated in an airstream at a temperature not exceeding 120 deg., then the chloride hydrate can be made up to 80% waterfree without it melting, nor that HCl or Cl2 is lost. Then in this condition it can be subjected to higher temps. without melting. And then by further heating nearly all of the water can be driven off. But since in the first heating phase one heats very close to the melting point, the first dehydration phase is a careful operation.

Then the patent states to make it go smoother what they do is take a hot MgCl2.6H2O solution, a given amount of e.g. 50% of anhydrous MgCl2 is added. By cooling you get a solid mass which is broken into pieces. The latter can be heated in a continuous apparatus to 300 to 400 deg. without melting. At this temp. the pieces are subjected to an airstream (dried by H2SO4, CaCl2, etc.) to drive off water, and so on. Eventually the anhydrous MgCl2 is subjected to fire and thus liquefied at red glow and then subjected to air stream, Cl2 is given off and MgO also forms. Rest is described in cited patent.

12AX7 - 15-10-2009 at 15:57

Crazy. But I suppose it makes sense: MgO has such an insanely high HoF, it's highly likely to form as a result of any possible reaction, at least given enough delta T.

Tim

Formatik - 27-3-2010 at 19:40

Acetone forms by the dry distillation of citric acid (Robiquet, Berz. Jahresber. 18, 502 and Ann. 25, 138), and by distillation of sugar, starch, or gum with CaO (Fremy, Ann. 15 [1835], 279). Fremy's descriptions of the distillations are given below. And the part of the paper dealing with it is also attached below.

Basically:
* The sugar and lime need to be very fine and they both must be as extremely intimately mixed as possible. The smallest particle of sugar should react with the smallest part of the lime. Because the unreacted sugar particles decompose and yield a volatile oil with a burnt sugar odor that normally forms by regular distillation.
* The ratio of 1 part sugar and 8 parts lime has proven to be the best. At least 500g of sugar is recommended.
* The mixture is then distilled in an apparatus which has to be at least as double as large as the mixture since the volume increases and the mixture rises. At this point the flame is removed. The reaction will continue on its own until it stops by itself. If the mixture was intimate enough, only little flammable gas is evolved (don't do it in an enclosed area or heat too strong) and an oily substance of varying composition distills over. It has an etheral odor and amber color.
* The oil shaken with water causes it to partly solubilize. The water soluble portion heated in a water bath at 70-80 C, is where acetone distills over. To get the acetone pure (Bp. 56 C), it needs to be rectified several times in the water bath.
* The oil which remains insoluble when shaken with water is obtained pure by rectification. At first acetone comes over, later the pure oily compound goes over. To be obtained entirely pure, shake again with water, rectify, and leave it in contact with CaCl2 to remove all water. The oily compound is colorless when pure and has a pleasant odor, is soluble in alcohol and ether, insoluble in water. It boils by 84 C. The formula determined for this was C6H10O. The material is called "metaceton". I'm not sure what exactly that is. I've looked through Beilstein a bit but found no compound with that formula and boiling point.
* The same said of the distillation with sugar, is said to apply for distillation using starch or gum instead, with the same compounds forming. Starch forms more "metaceton" than acetone. And gum forms a lot more acetone than metaceton.
* The distillation likely gets messy, so it's best to use unloved apparatuses. The yield of acetone probably sucks as not all sugar is converted to acetone some to metaceton and gases, but still an interesting method of formation.

Attachment: Ann. 15, 279.pdf (389kB)
This file has been downloaded 644 times

not_important - 27-3-2010 at 20:03

you can't find it because it doesn't exist

[Edited on 28-3-2010 by not_important]

J_ChemSoc_RSC_V56_1_p487.png - 40kB

Formatik - 27-3-2010 at 20:09

I thought there was something fishy about that determination. The boiling point and formula found for acetone were pretty spot on though. Thanks.

not_important - 27-3-2010 at 20:33

Propaldehyde is useful, reduce it to n-propanol or oxidisr to propionic acid. And 'dimethylfurfuran' may be http://en.wikipedia.org/wiki/2,5-Dimethylfuran