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Posts posted by UC

  1. Largely practical knowledge.


    For example, NO2 solvated in water disproportionates to give nitric acid and nitric oxide.


    [ce] 3NO2 (aq) + H2O (l) -> 2HNO3 (aq) + NO (g) [/ce]


    Almost everything is explainable in terms of lewis acid-base interactions, and electronegativity can help steer reaction predictions, but guarantees nothing. For example, sulfur tetrafluoride is a toxic gas that reacts violently with water. Sulfur hexafluoride is a dense, essentially inert gas.

  2. Nickel is much lower on the electromotive series than zinc, and the reaction is less energetically favorable- in other terms, slower.


    Outside or in a fume hood, you could add nitric acid or sodium nitrate (though you will need to remove the sodium chloride from the product) to make the oxidizing mixture known as aqua regia (or poor man's aqua regia for the sodium nitrate version, potassium nitrate would be okay as well- don't use ammonium nitrate). This gives off all sorts of nasty fumes (hence outside), but will even dissolve gold, which normally does not react at all with hydrochloric acid.

  3. http://en.wikipedia.org/wiki/Oxygen


    See the thing on the right that says freezing point? That's for atmospheric pressure.


    Mercury is regularly distilled to purify it. Good reagent grade material is triple-distilled.


    Iron can certainly be vaporized, but would react readily with oxygen if not under vacuum or inert gas blanket.




    Same deal. Note how extremely high the temperature is. This would probably best be done using an electric arc furnace.

  4. Think of a solution as ion soup. You're making tetraamminecopper(II) ions, so the anion doesn't really matter- unless it causes the complex to precipitate. If a copper salt is insoluble but its tetraammine form is, the solid will slowly dissolve in the ammonia.

  5. You'd be better off with a solution of bromine in a pressure equalizing dropping funnel. Appropriate solvents would be glacial acetic acid, chloroform, or dichloromethane, depending on the reaction solvent. You can always add it straight in under strong stirring, but you'll need better control over the flow rate than for solutions.

  6. A *sealed* bag for an aqueous solution? That's a *lot* of CaCl2 and effort for a small amount of ammonium chloride.


    You're better off evaporating the solution with a hot plate or just by air exposure (very slow) and then cooling it.


    ammonium chloride is very soluble in water and when concentrated, the vapor pressure of the solution is very low.


    About 2 years ago, I neutralized almost a gallon of 5% ammonium hydroxide with a slight excess of HCl, and left the solution in a cut-off polyethylene milk jug with paper towels over the top to keep out dust and bugs. The heat in the summer evaporated much of the liquid, and when winter came, the very gradual onset of cold precipitated out gorgeous crystals of ammonium chloride. I filtered these off and left the wet, somewhat acidic NH4Cl in a dessicator over NaOH pellets, which do double duty of trapping the water vapor and excess HCl.


    Since I've lowered the concentration by removing crystals and increased the vapor pressure, after the coming summer, I should be able to harvest another crop. Sure, I could do it with a hotplate and be done in a day or two, but where's the fun in that :P

  7. Uh, I don't know how easy A. Niger would be to get a culture of.


    The balanced reaction may look simple, but it's carefully controlled in steps by enzymes. Burning sugar or anything of the sort would produce precisely a mess. That process is also thermodynamically downhill, so there's a nice payoff of ATP for the organism, I suspect.

  8. 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.

  9. 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.

  10. 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."

  11. 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.

  12. 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.

  13. I have heard this commented on before, that we do not have much uranium left in the world. This topic may already have a thread. I tried searching for it to no avail, so I apologize in advance if it does. Anyway could someone elaborate on this?




    Generally the cited links from wiki pages are also good resources (not always though).


    There are always thorium based systems as well.

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