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Everything posted by xeluc

  1. unless you don't have a razor scooter... haha.. I used too.. Even had socks on it. But Aren't those a bit pricey for an aluminum source?
  2. Ok, well I'm gonna do some experiemtning and see what works best. After all, if it DOESN'T work, I can just re-disolve it ;-) Also, I believe CuCl is absolutly not soluble in alcohol. Unless the whole azeotrophism thing messes it up, I think that would work to keep it under. I'm ok with liquids in my samples..
  3. i know magnetite it, I dont' know about Hematite. Dude, it doesn't look like you have a whole lot anyhow, just scrap it and do it over, but make sure no copper is showing on your anode and you'll be fine.
  4. You should use the search function..... Me and Akcapr already gave out lot's of advice on rust synthesis. Check Here.
  5. :-\ Maybe not.. I am interested in thermite (Got the Iron oxide but never the Aluminum) and I'm not a pyromaniac...
  6. just look around your house! You can use many house hold items for stuff like this. Maybe a trip to a hardware store will help you.
  7. that is a VERY broad question. O2 can Oxidize.. so can H2O2. Flourine is a very powerful Oxidizer . I beleive that all Alkali Metals are good Reducers. Anything that you can think of that can take electrons form something is an oxidizer, and something that can give something electrons is a reducer.
  8. AHA! I know what the blue stuff on top is. If you notice on your first picture, you have the nail in the solution, but there is also copper wire exposed. This means that Copper is also oxidized into solution. the Light blue precipitate on top is Most probably Cu(OH)2 (Copper Hydroxide). As for the dark green stuff. I have also seen this when electrolising decently clean steel wool. I agree with Woelen, that it is a different kind of Iron hydroxide. Keep in mind that the brown stuff you called rust is actually Iron hydroxide. Apon drying it rearranges to become Iron Oxide So in the end, it's all the same. ALSO, I would think that no iron compound you could make is orange.... I'd say this is Copper (I) Oxide, as I have also made this, and have tested it and it tested positive for Cu Ions. what I don't understand is how you got this orange prec. first. You would think that the Iron would be oxidized first as Woelen said. Also, I dont understand why both Cu(OH)2 and Cu2O were formed... @Woelen: Does the green Hydroxide also decompose? My powersupply died so I can;t test this myself. Lol, does the picture say "supasweet" on that cylinder? haha.. sorry
  9. Aluminum foil is waaaay too coarse for thermite. To get powder of the grade you'd need, you'd need a ball mill to crush the foil you put into it. Ball mills aren't really expensive if you build them yourself.
  10. Well actually, I used one of your experiments, Woelen . In a container of conc HCl, I threw in a lot of Copper wire (better surface area for reacting) and some CuO to get Cu2+ Ions into solution (Cu(OH)2 or CuCl2 would work fine too, I just used what I had at the time). From there the Cu2+ Ions Oxidized the Cu into Cu+. After all traces of Cu2+ were eliminated I Transferred the HCl solution from my 20 Ounce bottle (Parents don't liek my chem experiments. They don't understand and therefor assume I'm making Meth ) to a more open container, where the Cu+ Ions would be oxidized further back to Cu2+ Ions. Then the now-brown solution was transferred and sealed back into teh plastic bottle. The bottle was sealed to aleviate excess HCl vapors. The Liquid would again becoem clear as Cu2+ ions were oxidized to Cu+. After each sucessive transfer, my solution became more and more conentrated wit Cu Ions. Originally I was going to use this as a source of CuCl2, but after beign enlightened by Woelen, I thought it'd be SO awesome to have a pure sample of CuCl. When the defining moment came, when I saw that glorious white precipitate come form a yellow/clear liquid, it was nice to say the least. It's 64 degrees F in my house right now, Oil prices in USA are too high for dad to turn on the heater . As a result, things evaporate slowly. Currently, I have a Dark olive green shell wiht some brown in it. When it is broken, white pure CuCl can be seen. Once this is all dried, I will end up re-inserting it into my CuCl-synthesizer and try again. So Woelen, can i have a reasonably pure sample of CuCl with my method stated in my above post? EDIT: It seems as if you answered my question in your previous post. So, why do YOU not do this? I am sure you would like a more-pure sample... If you have tried for so long I can reasonably say you would have tried this yourself, and if it had worked, you would have a picture of THAT instead of the olive green sample you have. So..... ??? EDIT 2: Why is it that I answer my own questions right after I post them. I re-read your post and you mentioned drying the with acetone. Well, I don't understand how that would help you.. Other than causing it to dry faster than when in water therefor maybe minimizing corrosion? Anyhow, after reading that part, it seemed like you were geared torwards a dry specimen (can I say specimen or is that only for live things..). In which case, my method would not be satisfactory for your needs. Hey, what if you kept the cap pf a vial full of water/CuCl slightly loose, boiled hte water away and as soon as the last of the visible water boiled away, you sealed the vial. It would be a slight vaccum with H2O Vapor, but if done right, little or no O2. Also, your vial would be seamingly dry, save maybe a LITTLE moisture that would condense... But I think that's pretty close. Who knows, just a thought. Don't knwo if it's feasable but I cant see why not. Other than maybe the CuCl goign everywhere during boiling. But the end result would be white or only very slightly oxidized CuCl...
  11. hm... I noticed that for a few hours while I was waiting for my CuCl precipitation to fall tothe bottom, the Powder was snow white. Within 10 minutes of decanting however, things were a different story. My crap is a lot darker than yours.. What if I put some CuCl that has not been dried (white) ins a Vial and heat up a tub of water to drive out disolved gases, then screw on the cap of the vial under water, to eliminate any air inside. The point here is not to make something where I can use CuCl for experimentation, I jsut want a pure sample. When it is created, it is EXTREMLY pure, theres nothing saying I can;t at LEAST get a product simlar to yours Woelen. I'm thinking however that CuCl will also react with water. Of course this must be a much slower process as I havn't seen this take place. A MSDS did say it reacts with moisture. Maybe only in the presence of Oxygen though? Can someone second this?
  12. yes, I beleive I will do that. Thanks borek.
  13. I just made some beautiful white CuCl. It is currently under water and seems decently stable. However, the container I used to decant liquid got some CuCl on the edges and it has turned a light olive Green color. I have to assume that the rest will too when dry. What is the best way of keeping my CuCl white?
  14. I can only hope to accomplish what you have, heh. anyway thanks a lot.
  15. Sorry, on elast thing and I'm out of your hair. if you could keep everything out of reach of Oxygen, if you added NaOH to the CuCl2(-) complex would you get CuOH? I cant find any ino on it, the closest thing is like CuCl2OH or something. I'm trying to make Cu2O, using the limited chems I have.... I was asking about the OH because I know most Cu+ compounds are unstable and I was hoping the CuOH would drop out water and become Cu2O, but I don't even think CuOH even exists...
  16. Ok, I'm fairly certain this is impossible to balance. I also now don't know how the oxidation state could even be changed. It seems more likely that instead of that happening, Cu2O just disolves with the help of HCl to make the Cu+ Complex and water. Cu2O + 2HCl --> 2CuCl2(-) + H2O. Interestingly (And logically now that I think about it) Everything works out. Like the Cl's I mean. At first I put Cu+ in the reaction and I'm like.. so where do the Cl's go... And then I remembered the complex that it is in.. It all makes sense now. I really needed to figure that out on my own. So then your Cu+ complex just turns into the Cu2+ complex With the help of O2. If I had some Cu2O I could test this hypothesis, but I have none.. Anyway... Thanks for helping to understand all of this, unless I'm still wrong. In which case....hah...
  17. Before I say anything, I just want you to know that this is what I think is going on, so maybe you should wait till someone agrees or corrects me before you take this... Yeah, that works for single replacement reactions, but double replacement reactions are a little more difficult. I know that if you have two solutions and mix them together and one or even two of the four possible compounds (2 of which being what you started out with of course) is insoluble, a reaction will occur. Ex: CuCl2 (aq) + NaOH (aq) --> Cu(OH)2 + NaCl (aq) . I believe that when all possible compounds are soluble, then maybe there wouldn't BE a reaction? the way I see it, you just have lots of Ions floating around, so really if you mix 2KCl (aq) + 2NaOH (aq) and you then evaporate the water, you might get KCl + NaCl + NaOH + KOH. So it's like an Equilibrium. This of course only applies to aqeous solutions. If you have solid reactants, then I supose you'd figure out each Ions tendency to steal away a different Ion. Like.. Na = Good at displacing Cu = Not good. F = I think good at displacing, It's sure reactive.. OH = Dunno, but it's not as good as F So I THINK NaOH + CuF would cause a double replacement reaction. I'll need someone to second this though...
  18. I have also experienced this happening, not with Gold slabs though. I always thought it was a suction of some sort since the to things were so smooth. I now know differently, Interesting.
  19. Ouch, that was a stupid mistake on my part... Well what about part 2 of my post?
  20. well I'm glad you figured it all out. The fact that H2SO4 is used in dehydration led me to believe that. Anyhow, as your usually the one who comes to the rescue when I don't understand something, I figured I'd just post this here. (Note: Nothing is balanced just to simplify things (Actually, I did a little..)) I believe that a solution of HCl and CuCl2 with Cu will Oxidize Cu into Cu+, then O2 Combines with Cu+ into Cu2O and the HCl Oxidizes that into CuCl2 and H2O. I worked out the charges, and they ar ethe same on both sides of the reaction. The Cu+ changes to Cu2+ and the O2- changes to OH-. So does that mean that the O is reduced? It's harder for me to understand since it ends up with an H there also. It'd make more sense for a Peroxide anion to be formed, but that would create H2O2, not H20. Could you clarify this? EDIT: I now understand that if a Peroxide Ion was formed, there would not be a conservation of charge.. Also, What if there was no HCl, just CuCl2 and H2O? I believe one of 2 things could happen. 1: Nothing.. Maybe you need an acidic PH for the Oxidation of Cu2+ to Cu+ to occur.. Dunno. 2: Cu2+ Oxidizes Cu to 2Cu+, 4Cu+ will bind with O2 to create 2Cu2O. This will then precipitate out of solution. Are any of these right? Thanks..
  21. Is the dark compound stable when all H2SO4 is boiled away or do you think this complex is only stable under acidic conditions... I understand you said it couldn't be dehydrated CuCl2, but I looked at your pictures and the brown particals resembled my dehydrated CuCl2. Of course, your probably right about the pictured brown stuff being too dark, I don't have any more CuCl2 at the moment, althoguh I am brewing some up . Here's something, do you get the same results with Copper Sulphate? If so, you can rule out Chlorine as having anything to do with the reaction. Also, I'm jsut guessing that CuCl2 is insoluble to H2SO4, so if you mix CuCl2 and some other acid other than H2SO4 that isn't a solvent for CuCl2, if the brown stuff appears, you can say that teh sulphate Ion plays no part. Your a really smart guy though and I'm sure you've thoguht this out much more thoroughly than me, I'm jsut trying to bounce some ideas. It seems reasonable that you could figure out what's in this compound by testing for when it is present, as for another acid that is insoluble with CuCl2, no idea. EDIT: assuming that any water in the H2SO4 didn't hydrate the CuCl2, you could try adding Anhydrous Copper Chloride and seeing how much the compound darkens, if it does.
  22. Regardless, you end up with Fe2O3 one way or the other.
  23. a super saturated solution is a solution that has more solute than can normally be disolved in the solvent at a given temperature. Disolve sugar in 90 degree C water until its saturated. Since more sugar can be disolved at a higher temp. When the temp is slowly lowered to room temperature, the solution becomes supersaturated. The disolved solute in supersaturated solutions is very unstable however, and excess agitation can actually cause a precipitation of the excess sugar, making it saturated.
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