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Everything posted by Melvin

  1. If you're going to distill the oil-ethanol solution, why not just distill the beer straight away?
  2. It would probably just react to give the salt of sodium and aspirin (sodium acetylsalicylate?) NaOH + CH3COOC6H4COOH --> CH3COOC6H4COONa + H2O I would think the NaOH would break apart the ester though (to sodium acetate and sodium salicylate)
  3. It depends on the concentration of the HCl that you added (its not pure HCl, which is a gas).
  4. It doesn't make sense to me why your Al/Cu cell gives such a low voltage (maybe oxidation of copper producing reverse voltage?). I was saying that seawater was slightly alkaline, which would make it more effective than a plain salt solution. I wasn't aware that you wanted a scale that big . The size of the plates only determines the current (bigger plates = faster reaction = more current). The voltage is determined only by what is reacting; in the case of the "4'x8' Mg and a 4'x8' stainless steel plates in the real sea" part of it would be the reaction of Mg (2.4v) and dissolved oxygen (-.397v). Combined gives you a theoretical maximum of 2.797v (note that since oxygen is reduced you negate the voltage). With the reaction using just the dissolved salts in the seawater, the reaction is between Mg (2.4v) and water (-1.23v), giving 3.63v (although the current is much less because the reaction is pretty slow). All of these values came from the electromotive series in the link I posted earlier. In reality, there will be resistance in the system so your actual voltage and current would be much less. The only real way to get exact values is to test it . Hope this helps a bit
  5. The graphite doesn't actually react; it just serves as a "return point" for the electrons. The Mg loses electrons, which flow through your circuit to the graphite, where the electrons react with the oxygen to reduce it to oxide (and probably the MgO reacts with water to make Mg(OH)2) I think the main problem with the Al/Cu cell is that the amperage is very low (because the reaction occurs slowly). I think if you used a more alkaline solution (which I believe seawater ranges from 7.5 to 8.5 in pH) you would get more current, so this could be feasible. The only real reaction here is with the aluminum, which reacts with the water to produce voltage. The more alkaline the solution, the faster this reaction occurs. The other electrode only serves to return the electrons to the water and produce hydrogen. Aluminum's value on the electromotive series is 1.67v (found at http://www.diracdelta.co.uk/science/source/e/l/electromotive%20series/source.html), so a cell like this could reach up to that voltage (although it probably won't due to some resistance). Magnesium dissolves in acid solutions, so if you used a very dilute acid solution with magnesium and carbon electrodes, you could be able to make a crude cell that could just be replenished with acid (and Mg when necessary). You could always make some kind of pump (non-electrical of course ) to disperse air into the water to replenish the oxygen supply.
  6. Na and Cl2 aren't used because such a cell would require a pure sodium anode, which would instantly react with water anyway and would be consumed too fast. It doesn't work with ions; it only would work with the pure elements. Since sodium has to be produced with electricity anyway, there's no sense in doing this. (To be honest, using the Mg doesn't seem great to me either, because Mg needs to be produced electrolytically as well)
  7. That's what I said. NaCl doesn't work because of the solubility of NaNO3. KCl works for KNO3, though.
  8. Hmm, I would think the barium salt would give a green flame (like barium salts typically do). Weird that yours didn't. KCl should give hints of purple with orange. Now, the way I would do it would be to soak strips of cardstock in solutions of different ions (Na, K, Ca, Ba, etc.) and let them dry out, then burn them to get the flame colors. Copper will also give green; copper with chloride gives blue or blue-green. Not sure about the pink, though EDIT: One thing I forgot is boric acid. Boric acid with methanol gives green flames, and boric acid with ethanol (I think) gives light blue flames.
  9. I thought the "aluminum foil in mouth" was electrochemical and requires that you have fillings. The fillings and the foil are electrodes, and saliva is the electrolyte. At least, that's what I remember...
  10. Just make sure you seal the connection well. The "something" was probably II (as in cobalt(II) chloride). That would give off chlorine at the anode (just like table salt would) instead of oxygen. Not sure where the person was going with that one. As long as you use good inert anodes (which you are), MgSO4 shouldn't give you any problems. The only time you would get different results is if you used a divided cell with MgSO4. In that case, you would get insoluble magnesium hydroxide AND hydrogen at the cathode, and sulfuric acid AND oxygen at the anode.
  11. Your reaction is a little off...it's actually NH4NO3 + NaOH --> NaNO3 + NH3 + H2O Be careful, though. As the video shows, the reaction is very exothermic and you could get sprayed by hot NaOH (which would suck). Ammonia is toxic as well.
  12. Thank you Ok, I will. I was more interested in just making it to see if I could. I might try making esters, but that makes me wonder about any impurities I did use splash-proof goggles and gloves, but yeah, I really need to use more safety equipment.
  13. The issue is where the wire (presumably Cu) contacts the graphite. If any copper is touching the solution, it will dissolve away pretty quickly to give insoluble Cu(OH)2. The pencil graphite is mixed with clay so will usually fall apart, so I use electrodes from D-size zinc-carbon (aka "Heavy Duty") batteries. These are nice, big graphite electrodes great for electrolysis. Just make sure the battery does not have the words "alkaline" or "lithium" on it. There are videos on youtube about it so you can search around. (You should still be careful with it and wear gloves when working with those batteries.)
  14. A saturated solution would give the highest conductivity, but as soon as the amount of water drops (is converted to H2/O2) the MgSO4 will start precipitating. Since the solubility of hydrated MgSO4 is around 71g/100mL, I would say use around 15g/100mL. This should give you plenty of conductivity without using too much.
  15. Should be pure...as long as you use an inert anode (best is Pt, graphite is okay)
  16. Ok, I wasn't sure about that. Thanks for clarifying. You don't need a saturated solution, just enough to make it conductive. Water decomposes at around 1.3v, but there is resistance in the system. 3v should be enough; the higher the amperage, the better.
  17. I think another common one is very dilute H2SO4... The one that I like is epsom salts. They're cheap, non-corrosive, and don't interfere with the reaction. If you do want to use NaOH, I would think you could definitely use less than a 1M solution, probably you could get away with less than a .5M solution. I'm not sure exactly, but you don't need much. Latex gloves would work. With table salt, you just need to keep it in a well ventilated area. You shouldn't be able to smell any chlorine unless you waft some of the anode gas; if you can smell it any other time, then you don't have enough ventilation. Also, electrolysis of table salt leads to the formation of hypochlorites (i.e. bleach).
  18. Well, electrolysis of pure, molten magnesium chloride gives you metallic magnesium. (MgCl2 MP 714C at wikipedia). This isn't really practical or safe to be done at home, especially since Mg is not too hard to find. A solution of magnesium chloride will just give you hydrogen at the cathode and chlorine/small amounts of oxygen at the anode. The Mg will not be reduced to metal, but simply converted to magnesium hydroxide, which precipitates. If you use the sulfate instead, you'll get only oxygen at the anode and hydrogen at the cathode. Plus, the hydroxide will not form since neither the magnesium or the sulfate ions are affected. By the way, magnesium does react with water. It's just slow. I placed a small amount of clean Mg wire into some water and you could see small bubbles of hydrogen forming (slowly). Since my dad is a welder, he can get Mg welding wire. I use that as my source (although I don't really use it very often).
  19. I didn't know you could even do that...although getting CO2 at those temperatures on your hands would be bad... I did some searching, and at http://hypertextbook.com/facts/2005/VictoriaPoon.shtml they claim that CO2 fire extinguishers can reach up to 5861 kPa. Searching around for lecture bottles of CO2, I found http://www.sigmaaldrich.com/catalog/ProductDetail.do?N4=295108|ALDRICH&N5=SEARCH_CONCAT_PNO|BRAND_KEY&F=SPEC that says CO2 lecture bottles are filled at 56.5 atm (5724.86 kPa). So it should be just as safe as a fire extinguisher when it comes to pressure.
  20. Okay, I did a simple test with the "sulfuric acid." I made some copper(II) hydroxide by mixing solutions of copper(II) sulfate and sodium hydroxide. I diluted the "sulfuric acid" to about 10% concentration. This caused the solution to get quite warm. I mixed a small amount of the copper(II) hydroxide with the diluted "acid," which gave a light blue solution. Just for comparison, I took pictures with that solution and a solution made with pure copper(II) sulfate. In both pictures, the left solution is the Cu(OH)2/"acid" and the right solution is the pure CuSO4. The left hand one looks slightly green the pictures, but in reality it was more bluish. The left solution may be simply a more diluted form of copper(II) sulfate, or it may be something that I don't know about.
  21. Most acids would dissolve FeO, giving salts. Depending on the acid, you might get partial oxidation to iron(III). Assuming that you don't get any iron(III), then the reaction is simple: FeO + 2H+ --> H2O + Fe+2.
  22. Okay, for some reason, this website won't let me edit my posts (and I'm not sure why not). I'll keep trying to edit it, but for now the best I can do is here: The previously mentioned procedure is dangerous as it involves toxic HCl and SOx gases and boiling, concentrated acids. It should only be performed with adequate ventilation using proper glassware. Protective eyewear, clothing, and gloves must be worn. Perform the experiment at your own risk.
  23. I'm sorry hermanntrude, I respect the hazmat policy. If you think anything I posted violates the policy, feel free to edit it out. As I mentioned, when I get more time, I will perform some more tests.
  24. I was using more than adequate ventilation; the only time I actually smelled a "strong HCl smell" was when I wafted some of the fumes. Otherwise, I didn't smell anything. The final solution was completely odorless and was quite viscous. The "SOx smell" was only observed during the boiling. Theophrastus, you're right that this doesn't work with lower concentrations of HCl. I used 30% HCl and carefully filtered off the sodium chloride before I started boiling it. When I get some time (school just started for me) I'll make some copper hydroxide and add some of the diluted product to it.
  25. Hmm, I seem to be digging up an older thread here... Anyway, I think this will be a good addition to it. I have been concerned with acquiring acids just like Theophrastus, especially sulfuric acid. I have plenty of HCl, but no sulfuric. Well, while searching around I found a chart that compares the solubilities of NaCl in HCl at http://pubs.acs.org/doi/abs/10.1021/je60084a015 and it shows that NaCl is pretty insoluble in concentrated HCl. So I mixed some excess conc. HCl with solid sodium bisulfate (normal sulfate would work but I didn't have any on hand), filtered off the insolubles, and boiled down the liquid. For most of the boiling, a strong HCl smell came off, but that became SOx smell and it gave off the characteristic white fumes as it became more concentrated. I ended up with thick, clear concentrated sulfuric acid . It discolors sugar immediately, and turns black within a few minutes.
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