UC

The corrosiveness of HF is only due to the powers of the fluoride anion. Chloride ion has just about 0 oxidation potential. The oxidizing power of nitrate should have nothing to do with how "strong" it is. The statement you make about the H2O2 is a load of bollocks. he's saying that if you make HCl and H2SO4 oxidizing, you will see that you are mistaking oxidizing power for acidity. Please look up the Hammett acidity function. At any rate, the strongest acid that can exist in aqueous solution is [ce] H3O^+ [/ce]. To what degree the acidic protons of a given acid choose to protonate water to make hydronium ion is the main concern.

no. There's an almost 100% chance that is just cold water with chlorine dissolved in it.

That's what gloves are for.

You do realize that this has been beaten to death as a subject on sciencemadness? I believe it's the anhydrous material that is truly dangerous, and I'm not sure that anyone there has been able to produce it, though I don't read energetic materials forum much. Also, the formation of the ammine complex requires that the solution have a reasonable amount of NH3 in it. Mixed with strong HCl, this is extremely unlikely. Splash some HCl in some dark blue tetraamminecopper (II) solution. The color of the complex (and the complex itself) vanishes upon acidification. The NOx formed will end up destroying a good portion of the ammonium ion anyway, via nitrous acid. IIRC, NH4NO2 *can* be made with ammonia and hydrogen peroxide, but is only "stable" in the cold, decomposing on heating to nitrogen and water.

for plain hydrogen? The energy is well above visible light and will rip apart all sorts of bonds, not just hydrogen. precious metal catalysts are useful for catalytic hydrogenation. I suspect you make something more along the lines of a mess or poisoned catalyst if you put it in a metal salt solution. As was said, simple bubbling will not achieve your desired reduction, nor will any easy method in aqueous solution. You're better off using another metal or something like hydrazine to reduce the copper. Alternatively, you precipitate copper hydroxide, roast to the oxide, and heat that in a hydrogen atmosphere to make copper powder. The real question is, why bother. Electrolysis readily plates copper sponge onto the cathode, and bright copper if the solution is acidic and contains appropriate levelling agents. I'm under the impression that you were after H2SO4 though, in which case, give up. If it were that simple, it would be widely known.

I would greatly advise against this. I believe someone over on sciencemadness managed to get clean bromine out of 1-Bromo-3-chloro-5,5-dimethylhydantoin (which is what I think those tablets are) but they used precise amounts of appropriate acid and distilled the bromine off. Halogens are all pretty horrible and I doubt you have the proper equipment to handle them. On the other hand, there is a fairly simple way to make bromine on woelen's site that even a beginner could probably follow.

The gases are soluble. Halogens are all at least slightly soluble.

It's some sort of mixed valency Cu(I),(II) complex involving chlorides. It's fairly common to see this. Please do this outdoors, if you aren't already. Aqua regia evolves dangerous NOCl fumes as well as NOx. Due to the ammonium ion, the web of reactions occuring is even odder. Here's a short rundown of some of what's going on: [ce] 4H^+ + 2NO3^- + Cu -> Cu^+^2 + 2NO2 + 2H2O [/ce] [ce] Cu^+^2 + Cu -> 2Cu^+ [/ce] (This is aided by chloride complexation of the Cu(I) ion as soluble CuCl2^-) [ce] Cu^+ + 2H^+ + NO3^- -> Cu^2^+ + H2O + NO2 [/ce] [ce] 3NO2 + H2O -> 2HNO3 + NO [/ce] [ce] NO + NO2 + 2H2O <-> 2HNO2 [/ce] [ce] HNO2 + NH4^+ -> N2 + 2H2O + H+ [/ce] [ce] HNO2 + HCl <-> NOCl + H2O [/ce] [ce] HNO3 + 3HCl -> Cl2 + 2H2O + NOCl [/ce] [ce] Cl2 + Cu -> Cu^2^+ + 2Cl^- [/ce] [ce] Cl2 + 2Cu^+ -> 2Cu^2^+ + 2Cl^- [/ce] [ce] 2NO + O2 -> 2NO2 [/ce] I'm sure I've missed a few.

Chalk isn't usually chalk anymore IIRC. I think it's almost always gypsum ([ce] CaSO4*2H2O [/ce]) nowadays. I precipitated mine from boiling CaCl2 solution using sodium carbonate, but only because it's an enormous pain in the arse to try to get CaCl2 back out of solution. If you have the time, you can dissolve eggshells or seashells in acid, filter, and add carbonate to precipitate CaCO3. Good quality limestone is pretty pure CaCO3. The common impurities are MgCO3 (form a complete solid solution, IIRC) and iron oxides.

At least half of this is nonsense. If you electrolyze NH4Cl, you will make plenty of NCl3. Any NCl3 at all is too much. End of discussion, unless you'd like someone to seriously hurt themselves trying any of this.

http://www.cem.msu.edu/~reusch/VirtualText/intro1.htm <- online organic chem text http://www.orgsynth.org/ <- excellent resource http://www.periodictable.com/ <- because it's awesome

I assume you mean more than just sputtering...can you give us any more specifics on what you're trying to do?

no, just metabisulfite is a good reducing agent and is acidic. [ce] Fe2O3 [/ce] is fairly easily reduced to Iron (II) ions under these conditions and dissolves. Metabisulfite is just a dehydrated form of bisulfite and yes, [ce] SO2 [/ce] will escape an aqueous solution of it.

http://www.sciencemadness.org/talk/viewthread.php?tid=2103 http://www.sciencemadness.org/talk/viewthread.php?tid=9797 </thread>

<nitpick mode> *mass* is density multiplied by volume. Technically, newtons would be the SI equivalent of weight, not grams. </nitpick mode>

Well, sodium chloride doesn't dissolve in alcohol, but that's besides the point. The easiest way to do this is by azeotropic drying or using molecular sieves. An alkali metal hydroxide dissolved in alcohol exists in the following equilibrium: [ce] KOH + MeOH <-> KOMe + H2O [/ce]. If the alcohol is truly dry, the equilibrium will lie to the right unless the concentration of KOH is very high (and therefore the concentration of water would be high). If you can remove the water from the reaction mix, then Le Chatlier's Principle states that the equilibrium will push to the right. Either you use molecular sieves to adsorb the water or you add a solvent that has an azeotrope with water, but not MeOH, which will boil off below the boiling point of methanol. simply distill off the azeotrope to drive the proportion of methoxide up. Also a problem with the original idea is that alcohols are not nearly as polar as water and as such, even ionic compounds that dissolve in them don't disassociate into ions nearly as well. These solutions are in most cases not good conductors.

salter, use the chem notations: [ce] CuCl4^2^- [/ce] for example, delivers [ce] CuCl4^2^- [/ce] Magnesium hydroxide would be formed in aqueous solution, not the oxide, and it is white colored. The black may be nonreactive contaminants in the Mg though. I also think you mean cathode instead of diode Very different things.

No, all you will produce in aqueous solution is magnesium hydroxide. You need molten salts of some kind or a mercury cathode to get metal.

No, there is a reason nitrating mixtures for polynitrated benzene rings contain a vast amount of concentrated or fuming H2SO4. The ring is extremely deactivated, especially in this case. Any chance or reaction would require a high concentration of nitronium ion, and probably a fair amount of heat.

Roasting metal sulfides takes an enormous amount of heat and is appropriately done in a furnace. H2O2 is generally stabilized with assorted phosphate buffers or phosphoric acid, which means that these will be contaminants in yoru final product. By far, the easiest way to make lots of SO2 is to acidify and gently heat bisulfite or metabisulfite (available fairly cheaply from wine making supply places) solutions. Ideally you'd use a non-volatile acid, perhaps cheap drain cleaner H2SO4 that is too dirty for direct use. The gas will have water vapor in it, but I doubt that is an issue if you're trying to just make aqueous acid in the end. If the seals on your apparatus aren't good enough, you will definetly smell it. The SO2 causes a choking, burning sensation and always makes me sneeze a lot.

This is actually well known, although NaNO3 is usually used for other reasons. In solution, you have ions. The salts are not very important. As long as there are NO3-, H3O+, and Cl- ions around, the mixture will work. It is known as "poor man's aqua regia." Using NH4NO3 is a bad move since chlorine might be formed in the mix and this would make chloramines or NCl3 and also because the ammonium cation precipitates things like Hexachloroplatinate ion, preventing dissolution. Merged post follows: Consecutive posts merged This isn't the issue so much as iodine stains. HI is very easily oxidized by air to I2 and solutions often have a brown tinge from this.

BAD idea. This nets you NCl3, a very powerful, dangerous, and unstable compound. Think NI3, but liquid and not safe when wet. You can get one or the other out, but not both at once. Either use H2SO4 for the HCl or use NaOH solution for the NH3.

Certain wild mushrooms can be extremely dangerous if there is also alcohol in your system.

Do a google search for chlorate cells and read about the electrode choices (especially the anode, which is under much harsher conditions) for them. These are the rugged kind of electrodes you want.

Throw it in concentrated HCl. If it dissolves, it's Cu2O. If it doesn't dissolve, it's Cu. If it looks like this: http://en.wikipedia.org/wiki/File:Synthesizing_Copper_Sulfate.jpg it's definetly copper.