jdurg

Heh. So we'll ignore the fact that the precipitate would wind up blocking the arteries and veins in her brain and vital organs and kill her instantly?

Well there's an enzyme in livers that catalyze the decomposition of it, so even in low concentrations it will decompose very quickly. At higher concentrations it would violently decompose in the blink of an eye. Also, even at low concentrations H2O2 isn't stable. That's why there's the pressure release caps on ALL bottles of H2O2. This is so that the decomposition products can escape from the container and not create a physical explosion as opposed to a chemical one. For the really high concentrations, they typically don't even store it. There's a certain limit to how high of a concentration the manufacturers will make. If you need higher, you pretty much have to make it yourself.

The f and g orbitals are complex shapes that really can't be described. I think if you look on webelements you may be able to find representations of them, but they certainly don't look like typical shapes.

Pretty much anything can cause H2O2 to decompose at the higher concentrations. Mechanical shock or temperature changes can result in the decomposition. The Activation Energy is really small for the reaction 2H2O2 => 2H2O + O2, so not a whole lot is needed to get it going.

The stuff is really 'complexly simple'. At first it seems utterly complex. Then you learn about it and it seems really simple. Then you learn more and it becomes utterly complex again. It's really remarkable.

There's also the high probability that the ink was not a water-based ink so water won't do anything to it.

The temperature range, as far as I can remember, is any temperature at all. So this will happen at room temperature. (I know because I've done it before). If air leaks in, you may still get a reaction going as long as the concentrations are high enough. H2 is just so damned light that it will quickly escape if a hole opens up.

Yeah, highly concentrated H2O2 isn't technically classified as an explosive compound, but it pretty much produces the same results.

Hey Woelen, here's another little 'experiment' to test out with your red selenium. You mentioned how you had made a bunch of red selenium and that it stayed red fro a long, long, long time. My question to you is, at what temperature was your selenium kept? My reason for asking is that apparently the transition temperature at which red selenium begins to form gray/black selenium is about 45 degrees Celsius, and at higher temperatures the conversion takes place much more rapidly. (45 Celcius being equal to about 113 Fahrenheit). A friend experimented with some red Se and said that when he heated a vial up slowly to 45 degrees C it started to darken and turn gray. At 55 degrees C, it pretty much instantly turned black. When I originally bought my ampoule of selenium, it was somewhat red but was quickly darkening. I guess what happened is that during shipping, the temperature of the ampoule was raised above that 45 degrees C mark and that started the change. In your experiments, you stated that the one you kept wet seemed to stay a brighter shade of red. I'm guessing that the high specific heat capacity of the water allowed the Red Se to stay red longer than the sample that was kept outside of the water.

Yup. Iron typically doesn't exhibit this reoxidation again in a typical thermit reaction because the reaction of iron and atmospheric oxygen isn't all that intense. The concentration of O2 in the atmosphere isn't high enough to really get the molten iron going. Still, some of it does oxidize and you can see that on the iron once it cools. With something like thorium, or zirconium, or magnesium, however, the reaction with oxygen under normal atmospheric conditions is VERY vigorous. (Hence why Zr and Mg have been used in flashbulbs, and why Thorium is used in lantern mantles). There's no need for an increased concentration of O2 to get those metals to burn violently.

1): not much. 2): not much, though contamination may do something.

Yes, thorium is radioactive. It's melting point is 1750 degrees Celcius and its boiling point is 4790 degrees C. I don't have a CRC handbook with me at the moment, so I cannot calculate the enthalpy of reaction for 3ThO2 + 6Al => 2Al2O3 + 3Th. Just keep in mind that Thorium is MUCH more reactive than iron is. As a result, as the liquid thorium forms it will most likely ignite in the air and form the thorium oxide again. I would test this out using praseodymium oxide first as the chemistry of Pr is pretty close to that of Th.

There's a very good chance that the resistance to the flow of electricity has heated up the surface of your magnesium ribbon. Mg WILL react with water to form hydrogen gas at elevated temperatures. With the electrolysis going on, there's a very good possibility that any oxide coating on the Mg was removed causing it to react much quicker than expected. I'm willing to bet that the hydrogen you got was a result of the Mg reacting with the water, or some acids that were forming where the Cl was coming out.

Yeah a boiling metal is dangerous, but it quickly cools down and liquifies again. It's just that the sudden boil can cause molten metal to be flung all over the place and start a huge fire. The premise behind a thermit reaction is that you take a very reactive metal, and a metal oxide. The very reactive metal, in most cases aluminum, will take the oxygen from the less reactive metal's oxide forming aluminum oxide and the less reactive metal. The heat generated is of such a great magnitude that the metal formed is a liquid. This can be done with virtually any metal and metal oxide. The HUGE problem is that the degree of reaction varies widely. Something like the Al/Fe-oxide thermit is actually VERY tame, believe it or not, yet it's still an unpredictable and violent reaction.

The big problem with perchloric acid that is not exaggerated is the reactions between HClO4 and many organic compounds. Perchloric acid accidents typically happen when the acid is used in a fume hood not designed for perchloric acid. While HClO4 is not volatile at room temperature, virtually all reactions involving a concentrated acid don't remain at room temperature. Remnant organics in these fume hoods may then get a chance to react with the perchloric acid and some unstable organic compounds could form. That's the main thing behind the danger of perchloric acid. The toxicity, as woelen pointed out and that perchloric acid producer/seller's article pointed out, is not really there.

I'd say use some hydrogen peroxide solution (3%) and let the shoes soak in there for a while.

Thermite temperatures typically get into the thousands of degrees celcius. Hence why iron is produced as a liquid, and some other common metals boil.

I remember learning the OIL-RIG one as well. Then there's the AN-OX, RED-CAT for remembering what happens at the anode and what happens at the cathode.

"Feed me Seymour............" Sorry. Habit. Nice plants though.

I know. I'm just being super anal-retentive today. I mean, if you really wanted to get down to it, a plasma of lithium nuclei would make an INCREDIBLY potent oxidizer because it would want to take electrons from just about everything. (Since it's just a bunch of naked nuclei).

Typically when someone asks what the strongest oxidizing agent is, however, they aren't asking about an electrolysis apparatus. You can't go onto your chemical supply shelf and pull a bottle of electrolysis out and mix that with another reagent and suddenly have the reaction occur. With fluorine, chlorine, oxygen, etc. you can do just that.

When I worked in a research lab, we cleaned EVERY piece of glass equipment with a solution of nitric acid. Nitric acid will dissolve just about anything, and I don't know of a single metal nitrate that isn't soluble.

For inorganic mercury compounds, yes, it would. Elemental mercury won't leach through the gloves, and the inorangic salts won't leach through them. Organic mercury compounds, however, are a COMPLETELY different story. Organic mercury will typically move right through a latex glove, as will most organic compounds.