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Author: Subject: Early chemists obtaining their reagents from nature.
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[*] posted on 20-3-2010 at 21:38


{insert evil laughter here}

My point is that if someone says something like I'm just really trying to be totally independent from society and then people suggest hi-tech approaches, there's a bit of cognitive dissonance going on. If you're talking about making your own glassware, resorting to Lexan shields seems to be going the other way.

It's really difficult to "be totally independent", unless you're willing to step back several centuries, or expend a very large amount of effort. If you really want to be independent, you make your own glass and use that to make glassware. You have to hold the glass hot enough to be rather fluid for several hours, allowing bubbles to escape and some impurities to form a scum. The you lower the temperature a few hunderd C to blowing temperatures. Old school blowers used clay blowpipes, good for a half dozen pieces or so. If you go out and buy steel pipe you're hardly independent from society, same goes for the tungsten picks commonly used; if you're willing to by them you might as well save time and effort and buy finished glassware which will be of better quality than what you can make.

Metal pipe and wire is another example, and don't even mention ball bearings. As watson.fawkes said, it's a great way of learning how complex and interdependent technological society is. It's one thing to step back and do as much as you can with castoffs and salvage, but another to be able to cut the ties and keep them cut.


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[*] posted on 20-3-2010 at 23:37


I see your point but I don't think he is going to extremes. I think he just wants to make as much as he can. Maybe because they are coming down on us so much anymore or just wanting to learn how to create from scratch, or both. I mentioned the Lexan after reading agorot's concerns which were valid ones for sure. Nothing could ruin your day much more than acid flying in your face from home made defective wares. I took his quest to mean how can I do this or that while doing so within reason. Using what you can easily get if for no other reason than the time it would take to build some things such as your wire extruding machine or some such item. I am all for his quest if for no other reason than I have always wanted to do a lot of these things also. Within reason of course. The first Phosphorus production is one example of where I would have to draw the line, who in their right mind would want to smell a big vat of boiling urine? A guy has to have his limitations.




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[*] posted on 22-3-2010 at 07:44


Isn't there a mmorpg out where you can virtually be 'totally independent from society'. I remember i was a tailor in WOW, i made all sorts of cloths, clothing and useful textile items, and it was far easier that actually doing that shit for yourself. Maybe this is an alternative you can explore initially to see if you like it. Although producing the finished items was fun, having to hunt for 47 mountain sheep or something to make the cloth was really tedious and helped me to understand that it's far easier just to buy your clothing from China like everyone else does.
Another way to be totally independent from society is to isolate yourself somewhere inaccessible to others, like Mars. I think you would find some very interesting minerals there. This is my stupidest post yet, ...the skipper too....'




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[*] posted on 28-3-2010 at 06:20


I'm surprised no one has mentioned Kevin Dunn's "Caveman Chemistry" -- you can get lots of the material for free from his site. I paid the $15 to download a PDF of the whole ebook, and thought it was worth the read. Basically, he shows you how to go from naked Chimp, to tools, cloth, alkalis, soap, paper, glass, acids, explosives, pottery, dies, ethanol, batteries, metals, even some basic plastics, from the sort of stuff you acquire with a sharp-pointy-stick.

I’m interested in very similar lines to you – even if I don’t produce all of my own soap, bread, and gunpowder, I like knowing that I COULD if I wanted to. Buying magic stuff from the magic factory makes me feel helpless. If I KNOW how to make that stuff, well, then I’m just paying someone to do it more efficiently when I buy it. Whole different mental outlook.

Here’s what I’ve gathered, mostly from above book:
1) Fire can be made from scratch
2) Clay + Fire gets you containers that can survive very high heat, and lets you build a kiln
3) Wood ashes from your fuel gets you mild alkalis -- sodium carbonate, potassium carbonate
4) Calcium nitrate comes from compost or stale urine, mixed with (3), you can get sodium/potassium nitrate (Saltpeter)
5) Sulfer and saltpeter gets you Sulfuric Acid via the lead chamber process, there's a thread on this site.
6) Sulfuric acid and saltpeter gets you to nitric acid (also on this site)
7) Nitric acid gets you to several types of plastics, explosives, and plastic explosives
8) Heat limestone (CaCO3) in your kiln gets you to quicklime (CaO)
9) Add water gets you to slaked lime (Ca(OH)2)
10) 9 + 3 gets you to lye – NaOH
11) 10 + Fat gets you to soap
12) 10 + Fat + Salt, gets you to glycerin
13) Any sugar, and natural yeast, you get ethanol, and acetic acid if you want it
14) I think its possible to get to hydrochloric acid via the above, but don’t remember how to do it (guarantee you’ll find it on this site)
15) Many pure metals are obtained by heating minerals that contain them
16) Charcoal can be made from wood. Charcoal, a little forced air, and an old brake disk gets you a forge hot enough to melt iron – see check You Tube for “Making iron from dirt” for a video of a bunch of guys making a big iron bloom out of basically nothing.
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[*] posted on 21-4-2010 at 11:51
Formic Acid


Apparently chemists have always been a little crazy. I read that Formic acid was originally created by distillation of red ants.

Who comes up with ideas like that??

How would you catch enough red ants to distill? If you built some sort of trap, what would you do if you got both red and black ants mixed up??

And how do you kill them without releasing the active ingredients or getting it smeared all over your caveman club??

I thought it would be cool to extract metal from rocks too, so I bought some lepidolite and downloaded papers on processing. One guy wrote his masters research on lepidolite processing, but he spent about 6 pages screwing up because he couldn't find ethanol of sufficient dryness and that was the limit of my patience for reading of his efforts! I crushed some up and mixed something in it and watched with fascination as crystals separated and floated to the top.

Thus began my interest in chemistry. (and still haven't gotten into the habit of taking good notes - never have been able to)

Suzee

[Edited on 21-4-2010 by SWilkin676]
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[*] posted on 21-4-2010 at 12:39


or how bromine was first found from I believe sea weed?
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[*] posted on 21-4-2010 at 12:39


um there is a thing called suger cain man ill kill your ants and make other things on the way.

like oxalic acid :)

as for how did they kill the ants they boiled them.




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[*] posted on 21-4-2010 at 12:42


Quote:
I read that Formic acid was originally created by distillation of red ants. Who comes up with ideas like that??


It seems like it was a gradual progression. That anthills have a distinctive smell has no doubt been known since prehistoric times. Then they observed that the blue flowers of chicory turn red on exposure to these vapors (or to ants in general). Then they just squished the ants to get ant-juice out of them and saw that this, too, had the same result. At that point distillation must have seemed like a pretty obvious way (to those curious souls who had access to the necessary equipment) to attempt to isolate whatever active principle was involved.

Quote:
And how do you kill them without releasing the active ingredients


Pretty sure you'd *want* to squish them to release the active ingredients. Unfortunately a detailed description of the process in English does not seem to be available via Google books; there may possibly be an account by some guy Ettmueller but his stuff is in Latin...
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[*] posted on 21-4-2010 at 12:43


truly though I wish I had an ant farm for formic It would make my life so easy.

thats right all I have to do is leave some shit in my pc that is edable.

trust me I have done this by mistake.

I thought id fucked a formic reaction till I had a look at there roasted carceses in my pc.

pc did not last long funny how ants will kill solder and nickel legs when they get roasted.




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[*] posted on 21-4-2010 at 15:33


Quote: Originally posted by DDbiology2010  
Does anyone know of some books or resources that show you how to obtain reagents from nature. For instance, how did early chemists obtain their chemicals for their chemical reactions. Today, people just buy the chemicals from others. How do you start from scratch. Also, a book on how to build chemical glassware from scratch would be awesome too. Thanks.



My favorite —

Ammonium chloride


Mellor - Modern Inorganic Chemistry 8th ed. 1933
CHAPTER XXVIII
COMPOUNDS OF NITROGEN AND HYDROGEN

§ 1. Ammonia-Occurrence and Preparation.

History.-Ammonia was known to the early chemists, and Geber describes
the preparation of ammonium chloride by heating urine and common salt.
Hence the alchemists' term "—spiritus salis urinae",. Ammonium chloride
was first brought to Europe from Egypt, where it was prepared from the "
soot " obtained by burning camel's dung. The name ammonia seems to be
connected somehow with the Egyptian sun-god—Ra Ammon ; ammonium
salts must have been known to the early Egyptian priests. The term sal
ammoniac was one of the early names for ammonium chloride; the
equivalent term sal aarmoniacum" which appears in the translations of
Geber's writings, and which was used for some time afterwards, was pro-
bably a mis-spelling, since the term "salt of Armenia " "—sal arrmoniacum" -
was applied to common salt and to native sodium carbonate.

————
Mellor - Inorganic and Theoretical Chemistry [edited]

THE ALKALI METALS

§ 16- Ammonium Chloride

In 1705, L. Lernery first showed that ammonium chloride exists among the
products derived from volcanoes, where he found it admixed with sodium
chloride, and this fact was verified by F. Seras in 1737 and by F. de Bonde
in 1765. Ammonium chloride occures as a sublimate mixed with other
volatile maters in cavities in the neighbourhood of volcanoes and in
crevices in volcanic lava — e.g. at Etna, Vesuvius, Stromboli, Hecla, the
Sandwich Islands, etc. It has been also found in the vicinity of ignited coal
seams — e.g. at St. Etienne, Aveyron, Newcastle-on-Tyne, Bradley (Staffs),
Hurlet (Renfrewshire), West Wemyss (Fife), Arniston (Midlothian),
Bucharia, Kilauea (Hawaii), Waldenburg, Kattowitz, and at Duttweiler
(Prussia). It has been reported in guano from the Chincha Islands ; in
natural salt ; in the mother liquor of some brine springs — e.g. Halle—and
W. Diehl found about 0.01 per cent. of ammonium chloride in carnallite
from Stassfurt. It has also been found in small quantities in the secretions
and exudations of animals—e.g. the urine of the camel.

Ammonium chloride is formed during the mixing of equal volumes of
ammonia and hydrogen chloride gases. When the two gases meet a white
cloud of ammonium chloride appears, but not, say H. von Helmholtz and F.
Richarz if the gases be previously dried. H. B. Baker showed that
combination does not occur if the gases be thoroughly dried, and that a
minute quantity of water is necessary for the reaction If the dried mixed
gases be confined in a vessel fitted with platinum plates with opposite
electrical charges, the two gases are separated-the ammonia collects about
the negatively charged plate, and the hydrogen chloride at the other
electrode. The action is not efectrolytic since no discharge occurs.

Ammonium chloride is also formed by the action of hydrochloric acid on a
soln. of ammonia or ammonium carbonate; J. G. Gentele r, made it by the
double decomposition of ammonium bicarbonate and sodium, magnesium,
calcium, and other chlorides; H. J. E. Hennebutte and E. Mesnard, and A.
Dubose and M. Heuzey, made it by the action of ammonium bicarbonate or
sulphate on the double chloride of iron and calcium; and it is made by the
action of soln. of ammonium sulphate and sodium chloride ; when the soln.
is cone. the crystals of sodium sulphate separate out and thev are removed
by suitable shovels; the cone. soln. of ammonium chloride which remains is
purified by crystallization. Ammonium chloride can also be obtained by
sublimation from a dry intimate mixture of the same two salts. A. French
made it by the joint action of air and steam on a mixture of salt, pyrites, and
carbon or organic matter :

2NaCI+4H2O+SO2+C+N2=2NH4Cl+Na2SO4+CO2.

Ammonium chloride has been observed as a product of many reactions — e.
g. the thermal decomposition of ammonium perchlorate, hydroxylamine
hydrochloride, or hydrazine dihydrochloride, N211602 ; the action of
hydrogen chloride on anhydrous azoimide: 3N3H+HCI=NH4CI+4N2; the
action of ammonia on chloramine: 3NH2C'+2NHS=3NH4Cl+N2; etc. J.
Raschen and J. Brock patented a process in which nitrosyl chloride, NOCl,
mixed with hydrogen is passed over heated platinized asbestos :
NOCl+3H2=NI14Cl+H20- Various proposals have been made to recover
the ammonium chloride formed in the ammonia-soda process when
NaCl+(NH4)HCO3=NaHCO3+NH4Cl+N2

Ammonium chloride was formerly obtained in Egypt as a product of the
combustion of camels dung, which always contains some sodium chloride;
the ammonium chloride was isolated as a sublimate from the soot. In India
dung was mixed with salt and similarly treated. Other nitrogenous products
can be treated in a similar way. The aq. liquids which collect during the
distillation of nitrogenous organic substances, which contain chlorides, also
contain ammonium chloride in soln. with other ammoniacal products. For
example, the ammonia liquor of gasworks, coke-oven plants, shale works,
and blast furnaces is a soln. of ammonia together with a great many salts of
ammonium—e.g. ammonium carbonate, sulphide, sulphate, cyanide, etc.—
and if the coal contains sodium chloride—salty coal—the gas liquor is
almost certain to contain some ammonium chloride. In any case, if the
ammoniacal liquor be neutralized with hydrochloric acid, the ammonium
salts are in a great measure converted into an impure ammonium chloride.
M. Adler converted the ammonium salts in the liquor into the chloride by
treatment with calcium chloride A. WuIfing used ferrous chloride.

It is also practicable to drive off the ammonia from the ammoniacal liquor
by treatment with milk of lime, and to pass the evolved ammonia into a
vessel called a saturator containing hydrochloric acid, cooled by water. If
the, acid is more cone. than corresponds with a sp. gr. 1.1, it loses some
hydrogen chloride when hot. Tile saturator is made of stoneware or some
resistant material since hot hydrochloric acid attacks lead. The liquid in the
saturator contains about 25 per cent. of ammonium chloride, and it is
pumped into a large wooden tank lined with lead. A coil of lead pipe heated
by steam is immersed in the liquid until a film of crystals forms on the
surface. The liquid is then decanted to a leaden vessel, where it is allowed
to crystallize; the crystals are removed, and the liquid run back to the
evaporator along with some fresh soln. The soln. of ammonium chloride is
not allowed to come in contact with iron, for during evaporation some
ammonia is lost, and the acid liquid attacks iron. The ammonium salts can
also be converted into the sulphate by treatment with sulphuric acid, and
subsequently the sulphate converted into chloride, as indicated above.
Ammonium chloride is also made by neutralizing with ammonia the spent
pickling liquor from galvanized iron works which contains a large
proportion of ferrous chloride ; or by treating with ammonium carbonate, or
a mixture of ammonia and carbon dioxide, the soln. of calcium chloride
obtained as a by-product in the ammonia-soda process. On evaporation,
crystals of ammonium chloride are obtained after removing the precipitated
ferric hydroxide, in the former case, and the calcium carbonate in the latter
case.

The purification of ammonium chloride. Crude sal ammoniac is usually
contaminated with iron or tarry matters, and in consequence, the colour
varies from yellow to red; it can be purified by heating it in thin layers on an
iron plate hot enough to drive off the water and free acid, and to carbonize
most of the tarry products, The grey mass is then sublimed. The sublimation
is conducted in cast-iron pots lined internally with firebricks, and covered
with a lid made of slightly concave plates. The salt to be sublimed is well
dried, and heated. Tile pots hold about hall a ton, and the sublimation
occupies about five days. The sublimate forms a solid fibrous crust about 4
inches thick. The crust is easily detached from the lid ; it is then broken up,
separated from adhering dirt, and packed for the market in barrels or sacks.
W. Hempel proposed converting the crystalline salt into hard stone-like
masses by press. between 50o and 100o.

Robert Hunt
Ure's Dictionary of Arts. Manufactures, and Mines
Longman's, Green, and Company
London 1878


AMMONIUM CHLORIDE. Commonly called Sal-Ammoniac. (Sal ammoniac, Fr.;
Salmiak, Ger.) The early history of this salt is involved in much uncertainty. It
would appear that the sal ammoniacus of the ancients was, in fact, rock salt. The
earliest knowledge of the compound has been claimed both for the Arabians and
the Egyptians; but the late Dr. Royle remarked, that 'the salt must have been
familiar to the Hindoos [sic] ever since they have burnt bricks, as they now do,
with the manure of animals, for some may usually be found crystallised at the
unburnt extremity of the kiln.'

This salt is formed in the solid state by bringing in contact its two gaseous consti-
tuents, hydrochloric add and ammonia. The gases combine with such force as to
generate, not only beat, but sometimes even light. It may also be prepared by
mixing the aqueous solutions of these gases, and evaporating till crystallisation
takes place.

When ammoniacal gas is brought into contact with dry chlorine, a violent reaction
ensues, attended by the evolution of heat and even light. The chlorine combines
with the hydrogen to produce hydrochloric acid, which unites with the remainder
of the ammonia, forming chloride of ammonium, the nitrogen being liberated. The
same reaction takes place on passing chlorine gas into the saturated aqueous
solution of ammonia.

Manufacture of Chloride of Ammonium from Camel's Dung.--In Egypt--which un-
doubtedly was the great seat of the manufacture of this salt from the beginning
of the thirteenth to the middle of the seventeenth century, and whence all the
European markets were supplied--the following is the process by which it is
obtained:--

The original source was the urine and dung of the camel, which are dried by
plastering them upon the walls, and burning, other fuel being very scarce in that
country. A fire of this material evolves a thick smoke, charged with chloride of
ammonium, part of which is condensed with the soot.

In every part of Egypt, but especially in the Delta, peasants are seen driving
asses loaded with bags of that soot, on their way to the sal-ammoniac works.

Here it is extracted in the following manner:-- Glass globes, coated with loam,
are filled with the soot, pressed down by wooden rammers, a space of only two
or three inches being left vacant, near their mouths. These globes are set in
round orifices formed in the ridge of a long vault or large horizontal furnace flue.
Heat is gradually applied by a fire of dry camels' dung, and it is eventually
increased till the globes become obscurely red. As the chloride of ammonium is
volatile at a temperature much below ignition, it rises out of the soot in vapour,
and gets condensed into a cake upon the inner surface of the top of the globe. A
considerable portion, however, escapes into the air; and another portion
concretes in the mouth, which must be cleared from time to time by an iron rod.
Towards the end, the obstruction becomes very troublesome and must be most
carefully attended to and obviated, otherwise the globes would explode by the
uncondensed vapours. In all eases when the subliming process approaches to a
conclusion, the globes crack or split; and when they come to be removed, after
the heat has subsided, they usually fall to pieces. The upper portion of the mass
is separated, because to it the white salt adheres; and, on detaching the pieces
of glass with a hatchet, it is ready for the market. At the bottom of each balloon a
nucleus of salt remains, surrounded with fixed pulverulent matter. This is
reserved, and, after being bruised, is put in along with the charge of soot in a
fresh operation.

The sal-ammoniac obtained by this process is dull, spongy, and of a greyish hue;
but nothing better was for a long period known in commerce. Fifty years ago [ca
1828], it fetched 2s. 6d. a pound; whereas now [ca 1878], perfectly pure
sal-ammoniac may be had at one-fifth of that price.

Manufacture of Sal-Animoniac From Bones and other Animal Matter.--

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[*] posted on 21-4-2010 at 15:37


Quote: Originally posted by DDbiology2010  
Does anyone know of some books or resources that show you how to obtain reagents from nature. For instance, how did early chemists obtain their chemicals for their chemical reactions. Today, people just buy the chemicals from others. How do you start from scratch. Also, a book on how to build chemical glassware from scratch would be awesome too. Thanks.



For a readable description of the production of phosphorus from urine &c. —

John Emsley
The 13th Element : The Sordid Tale of Murder, Fire, and Phosphorus
John Wiley & Sons
2000
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[*] posted on 22-4-2010 at 06:19


Quote: Originally posted by DDbiology2010  
Does anyone know of some books or resources that show you how to obtain reagents from nature. For instance, how did early chemists obtain their chemicals for their chemical reactions. Today, people just buy the chemicals from others. How do you start from scratch. Also, a book on how to build chemical glassware from scratch would be awesome too. Thanks.



Early chemists — a date range would be useful.

You preparing a list of the chemicals they used
would be good start in answering the question.
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[*] posted on 22-4-2010 at 10:30


You won't get anything of high purity, but you could try some of the methods that the grandfather of all chemistry used:

http://en.wikipedia.org/wiki/J%C4%81bir_ibn_Hayy%C4%81n
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[*] posted on 29-4-2010 at 15:41


Quote: Originally posted by The WiZard is In  
... History.-Ammonia was known to the early chemists, and Geber describes
the preparation of ammonium chloride by heating urine and common salt.


Boiling regular urine with K2CO3 or Ca(OH)2 will also form some NH3 since it has urea, though diluted with steam and any volatiles that may be present in the urine, e.g. volatile sulfur compounds from vegetables. Hence, boiling it down also prior to base treatment will get rid of those.

Also, some description from me of heating aqueous urea with K2CO3 or NaOH can be found in this thread: http://www.sciencemadness.org/talk/viewthread.php?tid=4800

[Edited on 30-4-2010 by Formatik]
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[*] posted on 29-4-2010 at 15:52


Quote: Originally posted by Skyjumper  
or how bromine was first found from I believe sea weed?


Close. Bromine was isolated from sea water which remained
after the crystallization of salt from the salt marshes of
Montpellier by A. J. Balard 1824.

Araon J Ihde
The Development of Modern Chemistry
Harper & Row 1964
Dover Reprint

Bernard Courtois (1777-1838), the discoverer of iodine, was the son of a saltpeter
manufacturer in Dizon. The elder Courtois assisted Guyton de Morveau when the latter
lectured on chemistry at the Dijon Academy. After a pharmaceutical apprenticeship,
Bernard was given an opportunity to study under Forncroy at the École Polytechnique
of which Morveau was director. For a time young Courtois was active in pharmaceutical
circles, but he joined his father when the saltpeter business was faced with financial
difficulties.

At that time, the ashes of seaweed collected along the coasts of Normandy arid Brittany
served as a source of sodium and potassium salts. One day in 1811 young Courtois
observed clouds of purple vapor rising from mother liquor which had been acidified with
sulfuric acid. The vapors, which had air irritating chlorine-like odor, condensed on cold
objects in the form of dark crystals with a metallic luster. A study of the properties of
these led Courtois to suspect that he had discovered a new element. However, the
press of business activities and the inadequacy of his laboratory facilities caused him to
turn over his chemicals to Charles-Bernard Desornies arid Nicolas Clement, two
chemist friends. These men reported the new substance in 1813. Courtois became
active in the manufacture of iodine, but others investigated its chemistry. Davy and
Gay-Lussac independently established it as an element. Its relationship to chlorine was
immediately apparent, and tire, oxygen-free nature of hydrogen iodide was generally
accepted.


--
djh
---------------------------
The Chymists are a strange Class of Mortals, impelled
by an incomprehensible Impulse to take their Pleasure
amid Smoke and Vapour, Fume and Flame, Poison and
Poverty - Yet among all these Evils, I seem to live so
sweetly that may I die if I would change places with the
Persian King!
Johann Beccher

Acta Laboratorii Chymica Monacensis, seu
Physica Subterranea, (1669)

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[*] posted on 7-5-2010 at 23:59


Quote: Originally posted by ninefingers  
Jules Verne describes his castaways in The Mysterious Island cooking pyrite (Iron Sulfate) ,water, and copper sulfate that they dug out of the soil. These distilled into S03 and H20; forming H2SO4. Then they cooked manatee fat into soap with wood ashes; leaving glycerin in the bottom of the pot. The Sulfuric acid and saltpeter from seagull dung made, of course, nitric acid for their nitroglycerin.

Also, burning sulfur with saltpeter produces sulfuric acid.

I dig old synths, too. :) I have a lot of bone meal (Calcium Phosphate et al) that I am going to calcine in a crucible to get Phosphorus, I hope.:(

[Edited on 3-20--10 by ninefingers]


CuSO4 decomposes to SO3 at 650 °C, with the addition of water it becomes sulfuric acid. You can do electrolysis of the CuSO4 solution and then boil it to concentrate it but be careful as it could release SO3. You will have h2so4.

According to wikipedia, this is how you can get blue vitriol:

MgSO4(aq) + 2 H2O + Cu(s) → H2(g) + Mg(OH)2(s) + CuSO4(aq)


But does this work with every sulphate compound I am not sure, I should try that with this:

calcium sulphate + water + copper ==> hydrogen + calcium hydroxide + copper (ii) sulphate
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[*] posted on 8-5-2010 at 01:27


There is this concept referred to as the electromotive series or electrochemical series.

Code:
____________________________________________________ | | | Electromotive Series of the Metals | | from Lange's Handbook of Chemistry, Eighth edition,| | Handbook Publishers Inc., Sandusky, Ohio, 1952. | | | | Metal Voltage | | | | Magnesium -2.34 volts | | Beryllium -1.70 | | Aluminum -1.67 | | Manganese -1.05 | | Zinc -0.76 | | Chromium -0.71 | | Iron -0.44 | | Cadmium -0.40 | | Nickel -0.25 | | Tin -0.14 | | Lead -0.13 | | Hydrogen -0.13 | | Copper +0.34 | | Silver +0.80 | | Palladium +0.83 | | Platinum +1.20 | | Gold +1.42 | |____________________________________________________|


Any metal below hydrogen will not react with proton sources (acids, H2O) to form H2. Copper is below hydrogen, it will not release H2 from water or acids.


You left out a very important part of that Wiki section:
Quote:
It can also be prepared by electrolysis of magnesium sulfate ...

the electric current is what drives the reaction, which actually is not that simple one given but rather two reactions, SO4<sup>2-</sup> reacting with Cu and 2 'protons' to form CuSO4, and Mg<sup>2+</sup> and water reacting at the cathode with 2 electrons to form H2 and Mg(OH)2.

Low solubility sulfates such as CaSO4 will not work very well, their solutions have only a low concentration of ions and are not very conductive. Ca(OH)2 is also slightly soluble, and so would react with the CuSO4 in solution. Al2(SO4)3 would work better.

The temperature of decomposition of most sulfates is high enough that the SO3 mostly breaks done into SO2 and O2, this has been covered in one of the currently active sulfuric acid threads if you want more information..


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[*] posted on 8-5-2010 at 01:47


It looks it's not possible then, thanks for explaining it to me :)
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[*] posted on 8-5-2010 at 01:55


You can make yourself glasware there iscouple of books about Scientific Glassblowing
here is some videos:
http://www.youtube.com/watch?v=QQnmSz8A5yc&playnext_from...
http://www.youtube.com/watch?v=8m-jEywA-T0&playnext_from...
http://www.youtube.com/watch?v=8m-jEywA-T0&playnext_from...
http://www.youtube.com/watch?v=Gh95--jOpx8&playnext_from...
http://www.youtube.com/watch?v=oNBC9F_d3QM&playnext_from...
http://www.youtube.com/watch?v=qmgftOMneXc&playnext_from...
http://www.youtube.com/watch?v=FHOyAcCqXQA&playnext_from...

I had some books about scientific glacblowing. In my university is Scientific Glassblower too (he made retort for me).
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[*] posted on 8-5-2010 at 06:47


Quote: Originally posted by Random  
It looks it's not possible then, thanks for explaining it to me :)


No, that's not what I said. What I implied is that you simple do not bother to really research these ideas you keep popping up, and just run with your first generally rather incomplete impression.

Most of of the reagents you are after have long discussions related to them, read those might help understand possible issues, your idea may have already been discussed, and generally one or more of the old threads is the proper place to be asking questions. In this case you might learn that SO3 is difficult to absorb in water, and that SO2 + O2 has potential for making H2SO4. You'd also see that I gave a reference for information on the decomposition by heat of a number of sulfates.

Wikipedia is best thought of as a Cliff Notes reference. The information given is frequently a quick overview and is rather incomplete if you actually want to perform a process given; the information given is just plain wrong at times as well. If something there looks interesting, you need to read any references given and/or chase down better detailed information elsewhere.

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[*] posted on 8-5-2010 at 10:54


Quote: Originally posted by not_important  
Quote: Originally posted by Random  
It looks it's not possible then, thanks for explaining it to me :)


No, that's not what I said. What I implied is that you simple do not bother to really research these ideas you keep popping up, and just run with your first generally rather incomplete impression.

Most of of the reagents you are after have long discussions related to them, read those might help understand possible issues, your idea may have already been discussed, and generally one or more of the old threads is the proper place to be asking questions. In this case you might learn that SO3 is difficult to absorb in water, and that SO2 + O2 has potential for making H2SO4. You'd also see that I gave a reference for information on the decomposition by heat of a number of sulfates.

Wikipedia is best thought of as a Cliff Notes reference. The information given is frequently a quick overview and is rather incomplete if you actually want to perform a process given; the information given is just plain wrong at times as well. If something there looks interesting, you need to read any references given and/or chase down better detailed information elsewhere.



Well, all of those ideas are because I have very limited access to the chemicals so I am trying to make them from the simplest substances. Of course, I am trying some of them currently.

As for the easiest way to make sulphuric acid is maybe electrolysis of water solution of CuSO4 like in this video:

http://www.youtube.com/watch?v=5dUSF9Gl0xE
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[*] posted on 19-4-2011 at 09:07


Quote: Originally posted by quicksilver  
Cinnabar when "roasted" yields mercury (I have NOT tried this) but did get some samples.[Edited on 19-3-2010 by quicksilver]

Was the cinnabar red?, if so **BE WARNED** If you hold it in your hand for a minuite or two, apon opening your hand up, you will see silver beads of mercury. So keep it in a suitable glass tube when handeling.

I know, I sell the stuff.

I love Nervada, it has a wealth of mineral specimens.
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[*] posted on 19-4-2011 at 10:57


Quote: Originally posted by bbartlog  

Pretty sure you'd *want* to squish them to release the active ingredients. Unfortunately a detailed description of the process in English does not seem to be available via Google books; there may possibly be an account by some guy Ettmueller but his stuff is in Latin...


----------
A system of chemistry,
Volume 2 By Thomas Thomson
1810

[OCR'd by Google.com/books]

This acid is first. mentioned in the Philosophical History.
Transactions for 1671, in a paper by Mr Ray, giving an account of
the observations of Mr Halse, and the experiments of Mr Fisher, on
the acid juice which is spontaneously given out by ants, and which
they yield when distilled *. Mr Fisher compares this liquor with
vinegar, but points out some differences between them. Scarcely
any addition was made to these facts nil Margraff published a
dissertation on the subject in the Berlin Memoirs for 1749, in which
he describes the me- . ' thod of obtaining the formic acid from the
formica rvfa, or red ant, and points out its properties with his
usual precision and method f. A new dissertation was published on
the same subject by Messrs Arvidson and Oehrn in 1782, in which
the discoveries of Margraff were confirmed, and many new
particulars added. Hermbstadt's paper on the same subject
appeared in Crell's Annals for 1784. His researches were directed
chiefly to the purification of the formic acid. He demonstrated that
the juice of ants contained several foreign bodies, and among
others, that a portion of malic acid might be detected in it. Richter
published experiments ou formic acid about the year 1793,
pointing..... &c., &c.

http://tinyurl.com/3rjtvdo

Someone -- should pop down to the BNL or some such and
pull the original paper. Or order it through la Intralibrary Loan.
Just tell the librarian that We will never find a cure for cancer
if I don't have this ....
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[*] posted on 19-4-2011 at 11:24


You can find loads of books on torrent sites. More than you can ever read. I looked for some a while ago, now I 've got 6000 of them, haven't even had time to read the titles...
search terms like 'chemistry', 'organic', 'survival', 'banned books', military manuals' 'paladin press', 'TEOTWAWKI', 'how to', etc... should give some results.
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[*] posted on 23-4-2011 at 17:31
Potassium salts from wood ashes


Following is a description of my semisuccessful attempt at extracting potassium carbonate and potassium sulfate from wood ashes. If you don't like long posts, move on to the next one...

Having heated my house with wood last winter, I ended up with a full 55 gallon drum of wood ashes. I decided to try to extract some pearlash from these.
I screened some of the ashes into a garden cart through 1/4" hardware cloth to remove some of the larger contaminants (mostly unburned charcoal) and weighed out 15kg of the resulting ash into a 30 gallon HDPE tub. To this I added 30kg of water. The resulting mass (which resembled freshly mixed concrete) I stirred two or three times a day over a period of three days, then let settle for a day.
Unfortunately not much settling occurred and only a small volume (~3 liters) of clear liquid could be decanted from the top of the tub. Seeing now why all of those 19th century texts describe leaching by filtering water *through* the ashes, I changed my approach and transferred the mass (in two portions) into a five gallon bucket with holes in the bottom and a layer of straw, then used 15kg further of water to leach more salts from the ash.
This yielded 19kg of a slightly yellowish liquid, sp gr 1.02 and pH 12. Obviously, more than half of the 45kg of water used remained trapped in the wet gray mass of ash, and thus the ley contained less than half of the available soluble matter; but as I didn't want to deal with boiling down a significantly larger volume of ley, I decided to discard the wet ash rather than leach it further.
This liquid presumably contained potassium carbonate, sulfate, and chloride (see 'The soap maker's handbook' from 1912, by Carl Deite, Alwin Engelhardt, F. Wiltner, p155, for one discussion of the soluble constituents of wood ash).
The ley was heated in a stainless steel stockpot at somewhat below boiling (about 80C) for several days until the volume was reduced by ~80%. The 3.8kg of amber colored liquid resulting had a specific gravity of 1.09 and a pH of 13. This was filtered and transferred to a smaller stainless steel pot and boiled down with a further ~70% reduction in volume. At this point small crystals began appearing on the surface, boiling was stopped, and on cooling a crop of light tan colored crystals was obtained. These were filtered out and washed twice with methanol (IPA would have been better, but methanol was what I had handy). On drying, 53g of a salt, presumed to be potassium sulfate, was obtained.
I did to few tests to confirm that it was not some other salt:
- a spatula added to 10ml of concentrated HCl gave no effervescence, so it wasn't a carbonate
- 3g were dissolved in 30ml of water (this required warming the water slightly, which is broadly consistent with the solubility of potassium sulfate), and this solution was added to a similar solution of CaCl2; after a few minutes a white precipitate formed (presumed: K2SO4 + CaCl2 -> CaSO4 + 2KCl). This suggests that the compound was not (at least not mostly) a chloride.
- the pH of 1g of the salt dissolved in 10ml of water was 9. This is *not* consistent with potassium sulfate (a neutral salt) but I attribute it to incomplete washing of the sample before drying, so that a small amount of KOH or K2CO3 still contaminated the sample
- neither heating it to 300C nor subsequently leaving it to sit in damp air for a couple of days changed the mass of the sample (49g), so it was not a hydrate salt nor a hygroscopic material.

Returning my attention to the remaining ley (now at sp gr 1.27), I repeated the previous step of boiling until crystals began to form and then cooling it. This time, only 10g of crystals were obtained on cooling. Further, the pH had actually dropped to 12 or so; I attribute this to the gradual conversion of the small amount of KOH in the ley to K2CO3 by atmospheric CO2, which apparently had now been completed after many hours of uncovered heating.
This crop of crystals did show a slight effervescence on addition of a spatula to an HCl solution, so I concluded that the K2CO3 was beginning to crystallize out. However this small crop might well have been a mixture of compounds and I discarded it rather than attempt further analysis.
The liquid was now boiled to dryness, with constant stirring so as to form a tractable bunch of granules rather than a rock-hard lump. 340g of a light brown crumbly salt were obtained.
This was transferred to a thin-walled steel canister with a tight-fitting lid that had a small hole punched in it. Pearlash was historically purified by heating to drive off or burn the impurities (see for example US patent #1, though in that case the ash is burned before leaching). The canister was placed in the glowing coals of my wood stove and a fairly hot fire (dry wood) built around it.
Unfortunately, my fire was hot enough to melt K2CO3 ... I would have thought it could reach 700 or 800C, but was somewhat surprised to see it reach 900C (melting point of potassium carbonate) without me using coal or any sort of forced air. Further, the canister I used was unsound in that it had a rolled seam at the bottom. Not only does fusion of the salt impair purification, but a third or more of it leaked out as well.
After cooling, I obtained a gray hard mass, which I broke out of the canister with a hammer. Black impurities coated it where it had touched the steel. The fragments weighed a total of 216g.
These were dissolved in an equal mass of cold water and then filtered, yielding 336g of a very pale yellow liquid of sp. gr. 1.47 (yes, the impurities were a lot of the remaining mass). I boiled this in a 500ml beaker to reduce the volume from 230ml to 150ml, then let it cool gradually (wrapped and covered in cloth). White crystals precipitated and the remaining liquid now had a specific gravity of 1.53, which, while not *quite* high enough for a truly saturated solution of pure K2CO3 at 5C (see: http://www.armandproducts.com/pdfs/k2so3P33_46.pdf for data), was close enough to be reassuring.

Crystals were filtered and 60g of crude K2CO3 hydrate obtained. Unfortunately this was left to sit on a watch glass overnight, and in the damp conditions resulted in deliquescence of the material with some of the resulting puddle pouring over the edges of the watch glass.
I recombined the now-mostly-liquid K2CO3 hydrate with the remaining solution and boiled it to dryness over a burner, with occasional stirring to prevent the formation of a single solid cake. This yielded 106g of white granular material, presumably K2CO3. In addition to being deliquescent it effervesces vigorously when a spatula is added to HCl, and yields a solution of pH 11 when 1.4g are dissolved in 100ml of water. While it is no doubt mostly K2CO3 I expect that there are still some impurities, maybe KCl or else other carbonates. Nonetheless it would pass for a pretty good pearlash circa 1845...
Overall this is a rather disappointing yield given the amount of K2CO3 that I estimate was present in the original ash (maybe 4%, i.e. 600g... though some references say you can get 10% from the ash of leafy/juicy plants, this was almost all hardwood ash, which is not as rich in potassium). Still I was pleased to get something given the number of mistakes made in the course of this extraction.

First picture shows crystals after boiling down to 150ml. Second shows later boiling to dryness. Third is potassium carbonate (106g) on the left, potassium sulfate on the right. Fourth is the K2CO3 on the watch glass before it turned itself back into a puddle of liquid.


Crystalline_K2CO3.jpg - 17kB Dried_K2CO3.jpg - 26kB K2CO3_K2SO4.jpg - 22kB

Pre-deliquescence.jpg - 38kB
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