UC

What the bloody hell....

no.

I suggest you learn what the hell you're doing first. Don't touch organic chem, especially chlorinations, especially lachrimators, especially moisture and air sensitives until you've putzed around with inorganics for a while and then done basic organic stuff. It helps you get a feel for things. Proper glassware and a fume hood costs a fortune. Also, try doing some good old literature search. The original posts shows an incredible amount of ignorance. PCl3 cannot be made from red phosphorus at any rate and pool chlorine isn't "chlorine", but can be used to make it.

Thermite implies an aluminothermic reaction, which using sodium certainly is not. Lithium is a stronger reducing agent than sodium and lithium has a higher boiling point than sodium, so any attempt to make lithium with sodium is hopeless. You can often distill the product out of a molten bath of say, calcium, which uses Le Chatlier's principle to cheat standard reactivity series. Calcium is appreciably soluble in molten CaCl2 and as a result it is produced via some tricky engineering that pulls the forming metal up out of the CaCl2 bath as a rod, to prevent it from dissolving and being reduced at the other electrode. As for the original poster, there are better and less stupid ways to get BOOMS than alkali metals. I suggest you take that discussion elsewhere and try and keep your fingers.

References, if you please. Cyanic acid undergoes nucleophilic attack at the carbon by ammonia, followed by tautomerization, yielding urea. Simply adding something like NaOH will not reverse the reaction mechanism. The hydrolysis of urea proceeds like normal base-catalyzed amide hydrolysis, giving first sodium carbamate and ammonia. The carbamate readily undergoes a second hydrolysis reaction to give sodium carbonate and another equivalent of ammonia.

I'd actually go out on a limb and say there probably is some sodium in there. That reaction with water is quite too violent for magnesium, I would think. Magnesium is, however, not a stronger reducing agent than sodium at high temperatures. Sodium, however, is volatile and will boil off from the reaction site if enough heat is supplied. Reactivity series can be cheated by Le Chatlier's principle. This heat is probably being provided by the reaction of the hydroxide portion of NaOH with the Mg, forming Na2O, MgO, and H2. The flames gushing out of the pot are probably burning hydrogen and some sodium vapor. That pot looks to be aluminum, which is a horrible choice, since it reacts with sodium hydroxide. While impressive, the amount of sodium formed is probably very small and and it is finely divided, trapped in the debris. The reaction with water was far too fast for any significant amount of Na to be present. Splashing paraffin in is a horrible idea. Anyone who has ever worked with concentrated NaOH knows that it's nasty and molten NaOH, which you could expect to find in the reaction mix will instantly blind you if it gets in your eye and cause some pretty horrible wounds if it gets on your skin. The paraffin is also flammable, which introduces further hazards. A stream of chilled argon would be far superior, but if you can get that, you probably wouldn't be needing to make sodium, especially like this.

http://tinyurl.com/nom2x2

Hopefully you mean freshly prepared Fe2O3. The pottery grade stuff is calcined and extremely hard to dissolve. If you have pure Fe2O3, you'll only have Fe(III) in the resulting solution. You don't want to use up all the excess acid. FeCl3 is stabilized in solution by excess HCl. In water, FeCl3 disassociates to HCl and ferric hydroxychlorides. In dilute enough and neutral/basic solution, it proceeds all the way to ferric hydroxide and you get precipitate. Adding extra HCl drives the equilibrium to the left toward FeCl3 (hydrated).

Cuthber- Where do you buy your chicken wire? Around here, what we call chicken wire has holes about 4cmx4cm. Metal mesh also goes by the name "hardware cloth" if that'll help you in your search.

Fe (II) doesn't form complexes, at least not that I've heard of. You'll make lots of hydrous Fe(OH)2 which will oxidize before your eyes to rust. Fe (III) in solution will certainly form Fe(OH)3 of varying degrees of hydration the second ammonia hits it. Since NH4Cl decomposes into fumes, you would be quite likely to get iron oxide if an iron ammine complex were heated. Fe (II) is a crappy lewis acid. Strong lewis acids tend to be extremely hygroscopic, but crystallizing an anhydrous salt from aqueous solution is not unheard of. I would suspect that it is probably the dihydrate given that water is in relatively short supply in concentrated HCl. Bubble chlorine gas through the solution. Instant Fe (III). I'd tell you to add H2O2 to a solution of FeCl2 in HCl, but the iron ions catalyze the decomposition of H2O2 and you'll get a lot of hot oxygen and less chlorine than you'd like. Adding bleach instead of H2O2 works if you don't have a problem with sodium ion contamination.

You'll need to shield it from oxidation if you want ferrous chloride crystals. FeCl2 oxidizes to Fe(III) ion very easily and if you don't have an excess of HCl around to trap it as FeCl3 and keep it from hydrolysing, you will get nothing but rust. If you want FeCl2, degrease some steel wool and throw into concentrated HCl in a covered beaker (It spatters a *lot*). A white powder settles out afterward, which is some form of FeCl2. I'm not sure if this would be anhydrous or a hydrate.

A two-fer: <Aerv> I'm bored of WoW, honestly :/ <- It finally happened!!!! The world must be ending. _______________________________________ <AzurePhoenix> I found your thread on glass stirring rods to be both engaging and stimulating <UnintentionalChaos> lol <UnintentionalChaos> kinky <AzurePhoenix> ... this is rare for me but i cant even muster some form of followup mockery * UnintentionalChaos wins <AzurePhoenix> this is proof of the evil nature of alcohol <UnintentionalChaos> bwahaha * AzurePhoenix beats Chaos with a sack of bottles <UnintentionalChaos> the pain is worth it

But the sodium will cause thermal burns as well and you risk a hydrogen explosion...in your mouth.

Glass stirring rods may not be anywhere near the most expensive equipment the average amateur chemist owns, but they are nonetheless ubiquitous. I also think it's nice to have at least one piece of equipment that you made in some way. Current ebay prices put the typical 6mm wide, 12 inch (305mm) long stirring rod between $0.75 and $1 (USD) each before shipping. Instead of spending my money on a dozen or so stirring rods, I purchased a large quantity of 6mm diameter clear borosilicate rod for glass working. I won 5lbs of rods for only $5 (which comes to 51 pieces that are 24 inches (610mm) long) and picked them up since the seller was fairly local. Stirring rods that I make from this glass cost me slightly under $0.05 each. They look like this as obtained, with crude broken ends: Normally, glass working (especially of borosilicate) is done with high-temperature torches, but for the simple application of rounding the end of the rods, a plain propane plumbing torch is sufficient. To round the ends, hold the rod vertically in the flame like this: The glass does not soften too fast with such a cool flame, but after a minute or so, it has liquefied enough to ball up. After cooling, it looks like this: Repeat on the other end of the rod. Of course, these rods are 24 inches long, which would make for a very inconvenient stirring rod. Half of this length is, however, a very nice size. Locate the center of the cooled rod and make a mark. You could just use brute force to break the rod, but if you have a fine metal file, you can do better. Place the rod against a hard surface and slowly "saw" with the edge of the file perpendicular to the glass rod until you make a small notch in the glass like this: Slowly rotate the rod while continuing to "saw," using the notch as a guide until you make it all the way around. It will look like this when you finish: I recommend wrapping a towel around the rod before you break it to prevent any glass chips from going flying. Apply a moderate amount of force to the spot where you scored the glass and it will give way with a more or less clean break. Mine, shown here, didn't want to cooperate with a perfect, flat break, but you can see that the break propagated mostly. along the scored line: Repeat the procedure with the torch to round off the new ends. Repeat and you have yourself a lifetime supply of stirring rods for pennies. Making a glass spatula The typical laboratory spatula is stainless steel, which is acceptable for most commonly encountered chemicals, but fails for things like copper (II) chloride or iodine when even traces of moisture are present. Plastic spatulas are also available, but for 5 cents, It's easy enough to make one out of glass. For this, you'll need a pair of cheap pliers without teeth, which have a reasonably large surface area to the jaws. I use a pair of 8" linesman's pliers that I took a shop grinder to. Heat up the end inch or so of glass until it starts to get soft, then insert the pliers into the flame (hence the need for cheap, as this will ruin the temper) and use them to flatten the glass, reheating between each squeeze. When working with such a low temperature flame, removing the glass item causes it to harden almost immediately. The metal on the pliers is a conductor of heat, and thus a bad choice for glass working, but still usable if you take care to reheat the piece often. Use the jaws to reshape the spatula as you please and put a bend in the end of it like this: I wouldn't call them pretty, but they will do their job with the more aggressive solids in your chemical library. The small amount of rust on the right spatula was from the pliers and is embedded in the glass, where it can do no harm.

Theo- Heptane is a very good solvent for iodine and less nasty than benzene. You'll find that the volatility of iodine makes separation nearly impossible though. You're better off slowly adding a sulfite solution to the tincture just until it goes clear, neutralizing with NaHCO3 solution (since the reduction generates HI), boiling off the alcohol, and then oxidizing the iodide back to iodine with chlorine generated in-situ. This part is actually what ope was trying to suggest above, but I've removed ethanol from the equation. Acidify the liquid with HCl and use a dropper to add bleach with swirling until you don't get any further precipitate.

The standard lab test for phosphate is to use ammonium heptamolybdate solution. On addition to either phosphate or arsenate, an intense yellow coloration or precipitate is developed. This is due to the formation of complex phosphomolybdic acids. Barium nitrate or chloride will give a white precipitate for both phosphate and sulfate. Once you know it must be one of the two, use the molybdate to test further. Theo has the others correct. Bicarbonate is very hard to distinguish from carbonate, although, you should be able to add a very small amount of calcium chloride solution into the bicarbonate solution without precipitate. Upon heating or standing, you'd get familiar calcium carbonate. Calcium and magnesium bicarbonate are what make hard water "hard."

You clearly have no idea at all what you're talking about. Kindy shut it. To the author of the thread: The way to approach this kind of problem is to backtrack through simple reactions. The molecule you want to make has bromines on adjacent carbons. This is the result of the textbook addition reaction of bromine across a ___________. There are quite a few ways to make a _________, but there is probably one you have learned which utilizes triphenylphosphine. I'll let you run with the rest of the problem. Hopefully that nudge got you going in the right direction

No, salt has chloride anion. Chlorine gas was used to kill tons of people in WW1. At low concentrations, however, it acts merely as an irritant. Plus NaCl is pretty toxic, all things considered. It would just be hard to eat enough to kill you because of the taste. When I was less careful with my experiments, I managed to almost gas myself once or twice. The smell clings in your nose for hours and I was coughing for quite a while. Unless you're stupid enough to add acid to the bleach, you won't reach levels like I inhaled.

Teflon is PTFE or polytetrafluoroethylene. Have you considered just ampouling the samples? with a little practice, this can be done somewhat crudely with a plain old borosilicate test tube and a propane plumbing torch. The halogens won't permeate solid glass. To do this with iodine, just keep the flame away from the iodine. For bromine, you should submerge the bottom of the tube in ice-salt brine to minimize volatility. You can buy teflon sheeting and cut out a cap liner to stick in a tube with a screwcap. Tightening the cap converts this into an inert compression fitting. It will very slowly leak over time, but is readily openable if necessary. If you don't need a shelf stable sample, seal the iodine in a HDPE jar with a tight fitting lid, place it inside a second jar containing sodium sulfite and sodium bicarbonate to absorb escaping iodine, and place the whole shebang in the freezer to keep volatility to an absolute minimum. Bromine is also best stored in the freezer inside a teflon-capped tube with the sulfite/bicarbonate secondary container if not ampouled. Alternatively, you can buy some FEP or PFA plastic bottles, which are also fluoropolymers and entirely resilient to attack by the halogens.

http://macro.lsu.edu/HowTo/solvents/toluene.htm For general chemistry, truly dry toluene is not necessary, but specialized reactions require truly anhydrous conditions. I believe refluxing over sodium with benzophenone as an "indicator" is the standard drying method.

It's a lovely mix of chlorine and chlorine oxides. Nothing you really want to be breathing on purpose, but a little won't kill you.

What kind of levels are you expecting? I suspect that you will need a very sensitive test for this project. Samples should probably be digested in Kjeldahl flasks with nitric and if available, perchloric acids under extended heating, before quantitative transfer to a volumetric flask and analysis.

Modern pennies are zinc, which has a low boiling point. You can probably figure out the rest. Use some pre-1982 pennies. I bet it doesn't work (they're solid copper). This is pretty ridiculously dangerous, what with flying thermite and zinc and zinc oxide fumes, which can cause metal fume fever.

lol. this post is priceless. can I frame it and put it on the wall so I can laugh every day?

Electronegativity has everything to do with it. As you may have guessed, covalent and ionic are not absolute terms. It is a range. The bigger the separation in electronegativities, the more polar they become, until the bonds become so polar that they can disassociate. The general rule is metal + nonmetal = ionic, nonmetal + nonmetal = covalent, but this isn't always the case. If you look, I think you'll find that a Pb-C bond is less polar than the C-H bonds in hydrocarbons, which we often think of as the prime example of being nonpolar. If you take metals to mean the alkali and alkaline earth metals, than yes, the rule applies.