Theophrastus

Given the temperature, pressure and density of a simple hydrocarbon, how would one go about find its identity, by using equations, and not empirical data.

I'm guessing yes, for obvious reasons...but just to be sure...

Actually, the green of copper roofs on old buildings is primarily a mix of copper carbonate and hydroxide, due to moisture and CO2 in the air, resulting in the characteristic green colour. Pure copper (II) oxide is black and insoluble. In terms of the colouration, the source of it is primarily copper ions, released by the electrodes. I recommend perhaps running the experiment for longer, and then bringing it to a halt, waiting to se what becomes of the solution as the water vaporizes. Usually (I've tried this one before, a fair number of times), one is left with a lot of leftover sodium chloride, a small quantity of greenish- bluish crystals (probably a smidge of copper chloride and hydroxide), and as well, a rather large quantity of dull orange copper powder. In your second experiment, is the grey compound suspended in the solution, or dissolved, as I would probably guess it to be aluminum hydroxide.

Hey, John, out of interest, would there be any way to extract the lactic acid, to get (regardless of crudeness) a relatively useable product? I think a while ago, YT, did something like this, extracting relatively pure (around 70% oleic acid) from olive oil by placing a jar of olive oil in the fridge overnight, then removing the insoluble mass present at the top. This is the very crude product. Then, adding sodium hydroxide, filtering off impurities, then adding hydrochloric acid finally leaves one with a somewhat usable oleic acid product. (Either way, the main impurities, are other aliphatic carboxylic acids like palmitic, stearic, and myristic acids)

Firstly I believe this is a repeat of an old thread if I'm not mistaken...hmm? As for uric acid, yes it can be found in urine, but in what purity? A wretched one, which would make it absolutely useless. I recommend that you start by buying one strong mineral acid (usually hydrochloric or sulfuric) and working your way up from there, as these can then further be used to make tons of organic, carboxylic acids, from their respective salts, when required and so forth. You can also then later, once you have enough technical, and theoretical experience, work with some even nastier (though useful) stuff like hydroiodic or nitric acid. You can easily buy hydrochloric acid at most hardware stores as "muriatic acid," and similarly, I believe I've heard sulfuric can be bought as "car- battery acid," though I've never done so myself. Happy experimenting. ps:I presume you know the risks associated with these stronger acids, if you don't, or are unsure, just consult MSDS (material safety data sheets; just type it in on google and you'll find your relevant querry) pps: and regarding horza's method, I know potassium permanganate (another strong oxidisers) to my own alarm, I found out very recently that it is sold quite commonly in certain pharmacies. Addition of potassium permanganate to alcohols wil yield the corresponding carboxylic acid, and the reaction runs to completion. In contrast, I find chromate would yield primarily an aldehyde product, unless you use excess oxidiser and heat it rather nicely, as it is a much weaker oxidiser. I believe dichromate would be more efficient in this respect. Chromates also happen to be carcinogenic, which... well... isn't quite fun.

What's a hickey? (tee hee!) *theo apologizes for throwing thread off topic with dirty comments*

Well yes' date=' but the rates of various enzymes may are rather slow, and variable; some taking much longer than others. (This could be exploited, given larger concentrations, but you'ld have to punch in the numbers) This can also depend on your choice of solvent, and what other substances are present. Aside from the physiological purposes in humans and animals, to which we are scarcely conscious of as it is, there are no other uses that I've heard of.

A rather simple question, but I'm hoping for some follow- up if an answer is given.

Out of interest, how would one go about reducing carboxylic acids, to aldehydes?

Sorry for the late post... but yeah, UC got it right. I'm using tylenol, actually (not the vile suspension, but the tablets), which where I come from is easy to come by, and relatively cheap. Joy.

Yeah, and even if the gunk were to be prevalent enough so as to form a visible, insoluble mass of sorts, the sheer quantity of impurities would make it rather useless. (Don't forget that dietary impurities (what you ate) can affect the composition of saliva as well) However, as horza said, unless your current was high enough, I find it unlikely that the enzymes would be damaged. Though they may perhaps pose a problem for enzyme function, as the process might reduce/ oxidize the metal ion cofactor (if altogether possible), or other such groups at the binding site, which can screw with your enzyme. So I would say that the process would probably affect function moreso than structure.

Well actually water makes up 99% of saliva's total composition, so your yield of other such proteins and whatnot, would be rather frugal. The remaining 1% consists primarilly of sodium, potassium, chloride bicarbonate phosphate ions and such, trace quantities of mucus, trace amounts of peroxide, thiocyanate, other antibacterials and such. The main enzymes are a- amylase, lingual lipase, lactoperoxidase, lactoferrin immunoglobin a, and various lysozymes, as well as trace amounts of tons of minor enzymes. It also contains various suspended human cells, and trace amounts of a strong painkiller; opiorphin. http://en.wikipedia.org/wiki/Saliva#Contents Here's a little more on the antimicrobial agents within saliva, if you're interested: http://jdr.sagepub.com/cgi/content/full/81/12/807

Ah, forgot to take note of that, that'll mean I'll have "sodium oxyphenylamine", which upon the addition of acid, will give me p- aminophenol.

A fun little thread that deserved some good- natured necromancy! I know this talk is all four years old, and whatnot, but it seemed twas cut off in an untimely fashion, for lack of a brilliant ending. Anyone with cool ideas who wants to join in the fun? Or maybe this year we'll get a new cast... Anyone got any worthy nominations? [Greek Chorus runs onstage]

And following up on fuzzwoods point, as its simply a matter of complexation, you can expect that the hydration of your resulting compound, will have little to do with your initial reactants (ie, just because the CUSO4 is pentahydrate, doesn't mean ZnSO4 shall have the same same degree of hydration, particularly given an aqueous solution). If I'm not mistaken ZnSO4 would likely take on the heptahydrate form, were it to be extracted from the solution.

Well, school work's kept me back for some time, but I've finally managed to respond. The experiment itself went rather successfully, and by the end of the week, the solution had turned a lovely yellowish- golden colour (due to the dissolved sodium terephthalate) with insoluble plastic materials (whose size had greatly decreased due to reaction) and small amounts of undissolved sodium terephthalate at the bottom. I then proceeded to empty out my initial 250ml flask, into a larger erlenmeyer flask (a little over 500ml), and added water, dissolving the remaining sodium terephthalate. (I actually accidentally went a little too far, and wound up diluting it a fair bit (thus the more bland colour of the solution in my second picture). I then filtered the solution, the filtration process being rather slow towards the end, taking about half an hour. I have yet to actually add the hydrochloric acid to precipitate the terephthalic acid, but will do so when I have some more free time. I was thinking of purifying my leftover ethylene glycol solution, after that, but it would have to be done outside, due to the dioxane formed. A tad unrelated, but how would one go about making high purity dioxane from ethylene glycol. (Relax, I know dioxane is nasty stuff and that in many cases, there are numerous ethers that are far safer and equally efficient in solvation. It's more of an out- of- interest thing than anything else.) ps: I've posted the photos of this synthesis on my profile.

What would be the products, in adding a strong base, like sodium hydroxide to acetaminophen? [ce] NaOH + HO-(C6H4)-NH-CO-CH3 -> [/ce]

(sigh) Lo and behold, a far simpler, more obvious, coherent, (not to mention, more likely correct) answer from UC!

While I am unable to pinpoint an exact compound, I would guess that your violet salt was probably a mixture of various isomers of iron complexes, such as the following complex (hypothetically speaking): [ce] Fe(HO-CH-CO-CH-CH2(OH)-CH2-OH)2 [/ce] (In this case, the CH's are positively charged, and bonded to the central iron atom) The gas released by the cathode was probably carbon dioxide (look up Kolbe Electolytic Synthesis) and of course, hydrogen gas, was released at the anode. Other than that, your remaining compounds are quite the mystery to me. While the black anodic precipitate sounds like magnetite to me ([ce] Fe3O4 [/ce]); how the iron got that way however, I have no idea. I would think that given oxygen gas, you would attain iron carbonate as a precipitate from the cathode, but again, electrolysis is a confusing process, with mixed results. I would keep some of the products though for analysis. In this respect, I'm sorry if my post was too late. If you still have the products, it might be interesting to do some ion tests on the precipitates. (If the yield was sufficient, saving some, if identified, for later use.) Cheer! -Theo ps: I'll get some steel nails, and some ascorbic acid from my local pharmacy and try the process out for myself. Can you give me some precise details on your execution? (voltage, current, temperature, electrolyte concentration, temperature, etc.) In doing so, I'm wondering if I can (a) destabilize my theorized complex with conceptrated acid, and then perform some cool ion tests, and other testing as to the nature of the ligands.

To be honest I have no clue exactly what occured, but I think you really ought to use a pure nitrogen (or nitrogen + inert gas (ie helium)) environment next time, as I'm pretty sure that the fireballs (pretty neat ) were the result of aerial impurities. My first thought was that magnesium nitride reacts with moisture in the air to produce ammonia gas. While ammonia by nature, is not very flamable in absence of a catalyst, and I'ld be pretty surprised if that were true, as you stated that you had a rather low oxygen supply. Of course, given the heat, as you opened the lid, it is possible that the still- hot ammonia might burn. Hot ammonia gas, is also known to burn when reacting with magnesium metal: another possibility. [ce] Mg3N2 + 6H2O -> 3Mg(OH)2 + 2NH3 [/ce] [ce] 4NH3 + 3O2 -> N2 + 6H2O [/ce] [ce] 2NH3 + 3Mg -> Mg3N2 + 3H2 [/ce] In the third reaction, in absence of oxygen, the hydrogen might later burn when exposed to air (via the opening of the lid). I also thought about the results of small quantities of hydrazine (which as you know burns very violently), given the production of ammonia, however, ammonia is usually converted to hydrazine in presence of a strong oxidiser, and thus would be very unfavourable or nonexistant. (If ammonia were the only reactant, and temperature the only influence) [ce] 2NH3 + H2O2 -> H2N-NH2 + 2H2O [/ce] I'm not sure how favourable such a reaction would be, as magnesium hydride decomposes (or reacts exothermically, in presence of water or oxygen) at 300 degrees celsius, to hydrogen and magnesium. (given a pressure of 1 bar) I also thought that under certain conditions, magnesium may only partially reduce the ammonia, to produce hydrazine. (however this is purely hypothetical) Like I said I'm not sure, but what I've described above, may be a possibility.

Alright, a week it is then. I suppose I might as well take the time to add some more sodium hydroxide and PETE, if it's gonna sit that long. I'll be leaving town on Tuesday, and returning the next weekend, so that'll be ample time for the reaction. Let's see what happens.

Well the molten lye thing's out of the picture, so no worries John. And yes, I was thinking of something on the lines of Kolbe Electrolytic Synthesis, but you raise a very good point (Yay you!); I could instead decarboxylate the terephthalic acid later formed, with molten lye, rather than winding up with naphthalene, and other unwanted biproducts, via the Kolbe electrolytic synthesis. (Not fun.) However, if I do try out this method, it'll only be in a while, as I'm still in the process of making a new condenser/still for refining alcohols, benzene and whatnot, and benzene has a rather low boiling point (or at least certainly lower than sodium hydroxide). ps: I've started the experiment, with 150mL of 50% ethanol, and have added some sodium hydroxide and PETE flakes. I'll run the solution for about a day, then post the results. When I'm done, I'ld add some HCl, and see what my terephthalic acid yield was. Happy experimenting! -Theo

Thanks, yeah that's it exactly. Thanks for the note; at least I know it would be a hell of a lot easier (no molten lye, and therefore ethylene glycol vapours), but like I said, I don't have any ethylene glycol as it is, (I was actually planning to use the batch I make to either convert to formaldehyde, or glyoxal) so I guess I'll use a solution of ethanol and water, and see how it turns out. I was actually planning to then try to electrolytically decompose the terephthalic acid (I'd add hydrochloric acid to the sodium terephthalate formed, and the terephthalic acid is nearly insoluble in water and ethanol) to make benzene, which for me would be a hell of a lot cheaper and easier, than buying it. Yay synthesis!

Thanks UC, I'll go buy a stainless steal pot for this one (or something of the sort). And no, there won't be any splashes. The only thing I'm worried about is ethylene glycol vapour, which boils before sodium hydroxide melts. I thought of using an alcohol- based solent, as it would be a lot safer, but I don't think the yield would be as good. I could try and plug any possible holes, but I'm worried of what'll occur if the ethylene glycol escapes. Of course I will be reared incapable of seeing what is going on, but it's just safer that way. I guess I'll probably run it for about 15 minutes, then leave it for an hour to cool, in an ice bath. Joy. Any thoughts? Concerns? Reasons to cry?