elementcollector1

Ha! No. Short-term, low-dosage exposure isn't a problem. And what do you mean by 'best answer'?

Too much heat. Try a stovetop, or even a hot plate. As for activation, I don't think it does: carbon is carbon.

Uh-huh. Unless the OP can find a chemistry class, summer program, and/or internship, then I doubt all those requirements are going to be fulfilled. Speaking of which, it's less than a gram of hydrogen chloride. Why so scared? The worst they could do is get a bit 'stung' in the eyes and nose, if that. There's no need for such extensive training to handle such a simple chemical. OP, just wear gloves and glasses and you should be fine.

Well, we all have to learn somewhere, and how better to learn other than by asking questions?

Is this homework...? A simple solubility chart should tell you most, if not all, of the answers.

Depends. Are you doing this inside? If so, that much isn't going to kill you - just sting your nose a bit if you get too close. What reaction are you using?

My mistake, I meant NaOCl, not NaOH. The sodium hypochlorite oxidizes the H2S to sulfur, with salt and water as byproducts. Unless I'm missing something else?

If you're experimenting with pyrite, I would recommend reacting the pyrite with a strong acid such as sulfuric or muriatic to generate H2S, and then bubble this gas (warning: This stuff is pretty toxic and smells terrible) into NaOH solution, making sulfur and water. The sulfur should precipitate out as a yellowish-white fluff, and can be filtered off with a coffee filter and dried. Need I say "Don't hurt yourself"? Or "If you do hurt yourself, I'm not liable"?

The cations matter only if they form insoluble nitrates. Assuming you're using Group I or II compounds, that shouldn't happen. AgNO3 is a good test for different anions because it forms easily distinguished compounds for each anion (AgCl, Ag2CO3, AgOH, etc.), and these are usually insoluble (precipitate out of solution). To find what ions would not be easily detected, look up different silver compounds of the anions you're using and their solubilities. If the material is soluble to a reasonable degree, the test would not work, or at best would work poorly. If the material is completely insoluble, a precipitate should form and be identified.

I would guess, if the ratio of thionyl chloride to m-toluic acid is 1:1, then you instead use 2 thionyl chloride. In short, use twice as much as you should stoichiometrically.

H2SO4 is a strong acid in aqueous solution, not a base.

http://www.essind.com/ace/ace-peroxide-concentrate.html Found that thing at McLendon's again, and managed to track it down. At $20, it was prohibitively expensive. What's worse, the MSDS doesn't list the concentration. Could 'titrate' using MnO2, but that would require me to buy it.

I would say, if possible, abrade the rust off with sandpaper, and dip the part in WD-40 a few times. That should do it. If the part's too small, I would recommend an acid, such as vinegar, lemon juice, to dissolve the rust (may need heating), followed by oil. What is the part you're trying to remove rust from?

Depends on the concentration. Anyway, if you want hydrogen, electrolysis is the easiest way to go about it. Use sulfuric acid or an equivalent sulfate for the electrolyte, and capture the gas somewhere. Incidentally, what kind of catalyst would work to split ammonia? I feel like making some nitrogen.

MMO, lead dioxide or platinum anodes will produce perchlorate. Dissolve chlorate at room temperature, and electrolyze: Perchlorate precipitates out due to lower solubility (at least in the case of potassium).

Unfortunately, that second precipitation is likely aluminum hydroxide, as the finely powdered stuff can react (when fresh) with water. The electronegativity series only applies here in that Mg is more electropositive (less electronegative) than aluminum, and aluminum more so than copper. You're correct on the nature of the reaction, and as for the electron transfer, you'd need something like a balanced equation. 3CuSO4 + 2Al -> 3 Cu + Al2(SO4)3 Something like that? Now given the oxidation states you mentioned earlier, work out how many electrons got transferred.

You can have pure, anhydrous HCl, but it won't behave as an acid (as HCl is a gas dissolved in water). H2SO4 can be concentrated to 98.6% under standard conditions, and cannot be concentrated further due to disproportionation into sulfur oxides and water. Nitric acid can be concentrated to anhydrous, as far as I know. Acids are proton donors, in that they will donate an H+ to bases, who are proton acceptors (and thus the acid-base reaction). In short, depends on the acid. If it is a gas dissolved in water, such as HCl or HI, likely not: the compound will no longer behave as an acid. If the compound can be a liquid on its own, such as nitric and sulfuric acids, then yes, it can be made anhydrous.

For your copper sulfate solution, where did the sulfate come from? H2O2 is useless in this application. However, if you wanted to make copper sulfate chemically from sulfuric acid, H2O2, and copper metal, that would be useful.

In fact, I've actually turned that green solution to the blue Cr (2+) with copious amounts of Al foil. What do you plan to do with your solution?

On second thought, liquid chlorine. The danger of frostbite, absorbing chlorine into the skin, and breathing it! That's three dangers for the price of one!

Oh, sorry. There was a spam post just below yours, it must've gotten deleted before you saw it. Anyway, this is a question I'd like to know the answer of, in terms of pH. I'm thinking the pH would be an average of how much of either is in solution.

Mon dieu, I've seen plenty of spam over at the other site, but not at this one.

The salt and water form a eutectic mixture with a melting point below 0 degrees C.

Hydrolysis? FeCl3 is known to hydrolyze. I quote from the magical Wiki: "When dissolved in water, iron(III) chloride undergoes hydrolysis and gives off heat in an exothermic reaction. The resulting brown, acidic, and corrosive solution..." I think you may want to try recrystallizing. I've never worked with nitric acid or iron compounds before, but recrystallization might help. Also, as John said, nitric acid is going to be as good as ferric nitrate for etching silver.

Hydrides do exist, if I recall. According to Wiki, in fact, crystallization for all alkali hydrides follows the motif of NaCl (cubic). An interesting correlation... Of course, it behaves more like an extremely strong base, as well as hydrolyzing, and is used as an effective drying agent, so I suppose it's not entirely like table salt. There are also unstable versions of halogens in the +1 oxidation state, for example bromine in bromine chloride (BrCl). So, could hydrogen be an 'inverse halogen' of sorts?