UC

http://www.scienceforums.net/forum/showthread.php?t=14408 http://81.207.88.128/science/chem/compounds/copper_tetrammine_sulfate.html

Unless you purchase xylene from a chemical supplier, the hardware store material contains mostly the three xylene isomers, ethylbenzene, and likely small amounts of toluene and nonaromatic alkanes or alkenes. When heated with sulfur, the ethylbenzene is converted to styrene with the release of H2S (do it outside!). This polymerizes and when you try to dry the sulfur, small clumps of polystyrene will be left in the solid. Also, xylene has an incredibly potent smell. I'd do the whole thing outside with a good breeze.

http://theodoregray.com/periodictable/AlkaliBangs/index.html Ta-da! There are real cesium and rubidium + water videos linked from that page.

No, silver will preferentially plate out of solution at the cathode as it dissolves at the anode. The only thing limiting that would be the rate at which silver ions reach the cathode since ions don't magically teleport through solution.

You do know the difference between an amine and an amide, right? Have a closer look at the structures.

Is balls to the wall really that awesome? It sounds painful, or at least uncomfortable and socially awkward, especially if the wall is stuccoed.

The moving mirror needs to be moved at a precisely known and stable speed. In order for the fourier transform to work, the location of the moving mirror must be known as precisely as possible at all times. Good instruments use two additional light sources (one a laser) solely to keep track of the moving mirror, IIRC. You also need a beamsplitter, which has to be an IR transparent material polished to obscenely flat and parallel surfaces. There is some discussion on sciencemadness.org about building an FTIR without a beamsplitter. Personally, I think the beamsplitter is a better idea. Less parts than the series of lenses and mirrors you'd need without it. IR and near-IR sources are all non-coherent light. Globars, Nernst glowers. Heck, you can even use tungsten filament bulbs for near-IR. I believe the IR light is passed through a collimating lens though.

This whole thread and nobody has linked this paper yet? http://gagne.homedns.org/~tgagne/contrib/unskilled.html

If you need dilute sulfuric acid, consider looking for replacement battery electrolyte at car supply places. It's supposedly quite clean. The drain cleaners, while concentrated often have a plethora of tarry brown material, iron contamination, foaming agents, or buffers added limiting their usefulness. I take the following excerpt outlining the hazards of the peroxide route from user Formatic on sciencemadness.org. I strongly advise against it on the basis of his experiences. "A way to H2SO4 is from estimated amounts of H2O2 and SO2. H2O2 will oxidize SO2 even in the cold to form H2SO4. Sources of SO2 is by burning and roasting sulfides (pyrite (FeS2); sphalerite and ZnS; chalcopyrite (CuFeS2), galena (PbS), etc), sulfur, or decomposing sulfites, or thiosulfites with a dilute acid... I’ve decided to try the above oxidation, but aborted the procedure because the reaction got too violent. I added 110 g of a mixture of sodium hyposulfite and metabisulfite to a 500 mL flask, the flask had a rubber stopper with a 50 mL separatory funnel and also a tube running out of it. The tube lead into 50 mL of 35% H2O2 in a 100 mL graduated cylinder. Then the separatory funnel filled with 16.8% pure HCl. The acid was then let drip in slowly and portionwise with occasional stirring. At first the bubbling of SO2 proceeded smoothly for several minutes, and the reaction between H2O2 and SO2 is highly exothermic reaching around 105ºC at some points. Though after some volume reduction, at some point the SO2 generator did something unexpected, without any warning whatsoever e.g. effervescence, foaming, etc. as SO2 was bubbling into the H2O2, the stopper blew off violently from the flask and the tubing shot out of the graduated cylinder, the acid/peroxide mixture spattered all over even on my arms and over the gas mask. After washing off, I came back and tried an even slower addition, but even then the exact same thing happened. I thought maybe the acid was too strong and diluted it with around 2 times the volume with water. The same thing happened! So I halted the procedure."

There is still nitrate present in your solution. As the concentration of nitric acid goes down, it's oxidizing power drops off and dilute acid will fail to attack additional copper. For the record, the endproduct of nitric acid oxidation of metal is usually NO2 (nitrogen dioxide). However, copper nitrate is much more soluble in water than the sulfate. Reduce down your solution and chill it and the crystals that form will be largely acid and nitrate free. Rinse them with a little cold water after removing them from the solution. To remove any trapped acid or nitrate (somewhat minimal, but possibly problematic), dissolve them in hot water and add a small amount of Copper (II) hydroxide, oxide, or carbonate (in reality, a basic carbonate since AFAIK you cannot prepare pure CuCO3). These will react with any excess acid. Filter the excess (insoluble) copper compound from the solution and chill to form very pure crystals of CuSO4*5H2O. Since the solution is fairly pure, you can boil it down and repeat to collect further crops of crystals. It's good practice to discard the last small amount of solution in any case since it contains the concentrated impurities (if any were present). An easier way to go about this whole process is to just dissolve the copper in moderately strong nitric acid (~30% works well, IIRC), neutralize and precipitate all the copper as basic carbonates with an excess of NaCO3 solution. Filter and wash the solid repeatedly with clean water to remove contaminants. Then slowly add dilute sulfuric acid to this until you just have a tiny bit of solid remaining. Filter this off, boil down and cool to crystallize. Repeat as above to collect copper sulfate pentahydrate.

Like a lot of things, what's currently economically viable is not what will be in the future. The average copper ore that we mine now is laughably poor in comparison to ore mined just a brief time ago historically speaking. Our techniques for concentrating the desired component efficiently improve with time and estimates may be based on current techniques. For example: http://en.wikipedia.org/wiki/Copper_extraction Furthermore, as limited supply drives the price up, the cost of processing lower and lower grade ore becomes reasonable. At some point extraction from sea water will cost just as much as treating the very low grade ore still available, and there are vast quantities in sea water.

http://en.wikipedia.org/wiki/Uranium#Resources_and_reserves Generally the cited links from wiki pages are also good resources (not always though). There are always thorium based systems as well.

Since you're talking about chalcogen-chalcogen bonds, you'd be in covalent bond land moreso than ligand complexes. Se(OH)4 exists, I suspect, at least in equilibrium with SeO2 in aqueous solution. ortho-selenous acid would be an appropriate (obsolete) name, IIRC.

One really nice side effect of this scam is that you can buy highly pure 35% H2O2 at "health food" stores. It's costy though.

I recently did a synthesis that involved heating benzaldehyde and pyrrole in refluxing propionic acid. Just...no. Very fortunately, the solvent was not butyric acid instead. I'll take my prize, now. I rather like the smell of dimethyl maleate, though it has a pretty low vapor pressure so I didn't know it smelled good until I spilled a drop on a glove and walked around wondering what that nice smell was for half an hour.

A small amount of variation happens sometimes. It probably has more to do with the instrument shimming and proper referencing than anything else. Concentration of sample may also be an issue as various kinds of intermolecular bonding change chemical shifts. I'd say if you have: 1) the right number of peaks 2) roughly appropriate ratios of integration area/peak height (integration is better, but sometimes things overlap a bit) for proton spectra (this goes down the toilet for carbon). I say roughly appropriate, because different protons have different relaxation times and the integration/height ultimately comes down to instrument paramaters versus these relaxation times. As an example, phenolic hydrogens consistently come up "short" with a relative integration of about 0.6 on the NMR I usually use. I could increase the signal collection time and the integration would increase to ~1.0 relative to other hydrogens, but it's not worth it unless it's for publishing. 3) roughly proper locations Then, you're fine.

Watch it or he'll ban internal rhyme as well.

Not until we get a punch-in-the-face-via-internet feature up and running. Our top scientists are working feverishly on this, I promise.

Fixed that for you.

Cold temperatures mean a lack of molecular kinetic energy. Temperature is molecular kinetic energy in fact. So anything spinning rapidly when cold is just ridiculous, unless it spins even more rapidly when hot. Similarly ridiculous is something releasing energy when cold for the same reasons. Secondly, why would spinning cause something to polarize? I'm apparently missing a chunk of logic here, because that makes no sense. Hydrogen gas is indeed susceptible to becoming an instantaneous or induced dipole, as are all other bonds for that matter. However, these have nothing to do with spinning or emitting energy.

Industrial HCl often has impurities in it. Pure material is completely colorless.

You can just buy ethyl lactate by the quart in some hardware stores. It's touted as a green/renewable paint stripper I believe. You'd have to cleave the ester, such as with NaOH, but such a solution, if lacking excess NaOH would be fine for preparing metal lactates.

I blame you.

You appear to not know what spam is. That was an attempt at humor.

Fire, and lots of it!