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Na or NaO

  1. Apr 20, 2006 #1
    Can I get Na or NaO relatively easily and using not very complex/unfindable chemicals?
     
  2. jcsd
  3. Apr 20, 2006 #2
    it is probably not easy. thhe onlly way i know of off the top of my head is to do electrolysis on a molten form of something like sodium chloride.
     
  4. Apr 21, 2006 #3
    I'd go with electrolysis of the molten hydroxide, a lot lower temperature and you don't have to deal with nasty chlorine.
     
  5. Apr 21, 2006 #4

    mrjeffy321

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    Yes, electrolytically making Sodium metal from molten Sodium Hydroxide is much easier than molten Sodium Chloride.
    The only dificulty I found with this method in practice was trying o keep a delicate balance on the temperature of the reaction....too cold and the NaOH solidifies and will not conduct current....too hot and the Na dissolves into the NaOH and/or oxidizes with air much more quickly. I was able to produce a tiny ball (much less than a gram) of Na metal by the process which when I threw it into a pool of water, reacted in a way which you would expect Sodium to react.

    Once you produce the Sodium metal, it should naturally oxidize itself with the Oxygen from the air due to its high reactivity, you just need to be sure to keep it very very dry (including water vapor in the air) to avoide forming Sodium Hydroxide again.
     
  6. Apr 21, 2006 #5
    I've had similar problems. It would be ideal to have enough current in your cell that it would keep the NaOH molten and at a constant temperature, but this seems to be difficult to do in practice. What were the specifics of your cell? I basically followed instructions from http://www.sas.org/E-Bulletin/2001-10-05/chem/column.html using iron electrodes. The process works fairly well but only for very tiny amounts.
     
  7. Apr 21, 2006 #6

    mrjeffy321

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    Actually, that web site's description is not far off from what I did.

    I made a very "cheap-o" electrolytic cell to preform the reaction in.
    I used a tuna fish can as the cell body and made an small structure out of popcicle sticls to support two Iron electroldes made from the "wire" of coat hangers. As a power source, I used a computer power supply.
    The trick is to get the reaction started, once you get it going, the heat generated from the reaction and current passing through the cell should, in theory, keep it going.
    I partially melted some of pile of NaOH around the electrodes and plugged it all in, then continued heating (butane torch) it until I could see Na metal beginning to form.
    But my NaOH did not like to stay liquid very long, it kept solidifying and the electrolytic reaction stopped, I would need to continuously re-heat the NaOH to keep it going, but by doing this, the over all temperature of the cell would get too hot and the Na would quickly oxidize and/or dissolve into the NaOH (not to mention the tiny explosions of molten NaOH going off every once in a while [and me not wearing goggles]).
    In the end, I was able to produce a pretty good amount of NaOH as I watched it form on the electrode, but I was only able to extract out a very small glob of it (extracting the Na is the hardest part I think, it just doesnt want to stick to anything). When I tossed it in a pool of water, it hopped around on the surface and produced bubbles before quickly dissappearing. I was quite pleased with myself that day :approve: .
     
  8. Apr 21, 2006 #7

    Borek

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    I already told you: it is Na2O.
     
  9. Apr 24, 2006 #8
    Thanks guys.

    I wasn't really thinking of electrolysis when I first asked. However, this is what I ended up doing in my research: I tried to figure out some reactions. With NaCO3 I tried to find some answers:

    NaCO3 + H2SO4 --> NaHSO4 + HCO3- (What happens to the HCO3 anion?)

    NaCO3 + HCl --> NaOH + Cl + CO2 (I am not at all sure about this one. I had to re-do it a couple of times and I still don't know.)

    NaCO3 + H2O2 --> NaCO + H2O (This one I am pleased with, but I'm probably not right.)

    Then with CuSO4 I wanted to see if I could isolate the Cu:

    CuSO4 + H2SO4--> CuO2 + S2O4 + H2O2 (A lot of guessing work. Not at all sure on this one.)

    CuSO4 + 2HCl --> CuCl2 + H2SO4 (This one took a lot of re-doing as well and I have no clue about it.)

    CuSO4 + H2O2 --> Cu + H2SO6 (This one I think I might have found my goal, but once again it's scrappy.)
     
  10. Apr 24, 2006 #9
    Well I ended up doing dome theoretical experiments. With NaCO3 I tried to find some answers:

    NaCO3 + H2SO4 --> NaHSO4 + HCO3- (What happens to the HCO3 anion?)

    NaCO3 + HCl --> NaOH + Cl + CO2 (I am not at all sure about this one. I had to re-do it a couple of times and I still don't know.)

    NaCO3 + H2O2 --> NaCO + H2O (This one I am pleased with, but I'm probably not right.)

    Then with CuSO4 I wanted to see if I could isolate the Cu:

    CuSO4 + H2SO4--> CuO2 + S2O4 + H2O2 (A lot of guessing work. Not at all sure on this one.)

    CuSO4 + 2HCl --> CuCl2 + H2SO4 (This one took a lot of re-doing as well and I have no clue about it.)

    CuSO4 + H2O2 --> Cu + H2SO6 (This one I think I might have found my goal, but once again it's scrappy.)
     
  11. Apr 24, 2006 #10

    mrjeffy321

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    Alot of those reactions are incorrect.

    Remember, Sodium likes to have a +1 charge on each atom (completing its octet), so Sodium Carbonate should be written as,
    Na2CO3.
    There is no way you are going to find a chemical reaction you can do in aqueous solution with a Sodium compound which will produce Sodium metal...it is just too reactive a metal (even if you do, it will react with the water and produce NaOH).

    An easier was to produce elemental Copper from solution would be to dispace it using a more reactive metal. Copper is prettly far down the activity series (below Hydrogen), so it shouldnt be hard to find a metal which will displace it, for example Zinc, Aluminum, Magnesium, Iron, ...
     
  12. Apr 24, 2006 #11
    that's Na2CO3 and you'll get carbonic acid, H2CO3 which is unstable and will decompose into water and carbon dioxide.

    This is analogous to the baking soda and vinegar reaction which I'm sure you've done. There's no chlorine (and that would be Cl2 by the way) or sodium hydroxide formed! This is also pretty much the same reaction (except with HCl) as the above one.

    Na2CO3 + HCl --> NaCl + CO2 + H2O

    How did you get this? I am pretty sure no reaction occurs and I highly doubt NaCO exists.

    No reaction would occur. All you would have is a bunch of ions in solution.

    The reactants and products on this equation are reversed. The equilibrium in your current equation lies far to the left. If you heat up a mixture of CuSO4 and HCl, the HCl will just distill off. Since H2SO4 has a higher boiling point than HCl, you can, however, mix H2SO4 and CuCl2 and upon heating HCl will distill off you'll be left with CuSO4.

    No reaction would occur. I suggest that you read a general chemistry book or look at some online tutorials. Just because an equation balances, doesn't mean that you'll get those products or that the reaction will occur. Nevertheless, these equations that you are attempting to solve are fairly simple and with some knowledge of single and double displacement reactions, you'll be well on your way.

    To make Cu from CuSO4, I agree with mrjeffy321, displace it with a more reactive metal. Try throwing a nail into your solution.

    I don't know if you were just guessing with the H2O2 reactant. The only related thing I can think of is that H2O2 used in combination with HCl will speed up the dilution of many metals such as iron and copper.

    As far as making sodium, it is rather dangerous and as pointed out already it cannot be made from aqueous solution in any way. Your best option is electrolysis of molten sodium hydroxide (there are several ways to make this by the way). You could also try some thermite reaction to produce sodium, but this method is probably even more dangerous and requires much more experience.

    Aqueous no. But I still have hopes for electrolysis in other solvents. There are some nonaqueous electrolysis possibilities for lithium, at least.
     
  13. Apr 25, 2006 #12

    mrjeffy321

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    Well Lithium is an even more reactive metal than Sodium (several elements higher up on the activity series than Na), but it just doesnt give off as much energy when it reacts, or doesnt react as fast, or something since a Li + H2O reaction seems to be much less energetic than a Na + H2O reaction.

    What kind of solvent were you thinking of? In order to dissolve the NaCl (or other Sodium Salt) you would need some type of polar solvent that at the same time would be harder to reduce than the Sodium in solution, not an easy combo to find.

    This would be a very interesting thermite reaction for sure, but it would be hard to preform. In order for the thermite reaction to proceed, you would need some Sodium salt (hopefully an Oxide if we want to stay true to classic thermite) and a more reactive metal. This buisness of needing a more reactive metal wont be easy as we would need to use something like Potassium or Lithium, metals in their own right are hard to get in their elemental form due to their high reactivity.
     
  14. Apr 25, 2006 #13
    Well, I was using a textbook. My level of chemistry (Grade Eight) is not very advanced. Thanks for averything.
     
  15. May 10, 2006 #14
    I have read that small yields of lithium can be produced by electrolyzing a solution of LiCl in dimethyl sulfoxide. I am not sure if such a reaction would work with sodium. Basically what you need is a solvent that can dissolve sodium ions and that does not react with metallic sodium.

    If you want to learn about sodium "thermite" reactions, I suggest you look at this http://www.sciencemadness.org/talk/viewthread.php?tid=2105. Several members of that board successfully produced small amounts of sodium.
     
  16. May 19, 2006 #15
    Come on

    Why worry about making it. its cheap from any specialist chemical provider and you don't need a license or any thing.
    in england so don't know any current american suppliers but its much easier than wasting your time with pointless incorrect reactions.
     
  17. May 19, 2006 #16
    investigation

    :surprised if of course its for an investigation and you have to find any easy way than i just thought why not try adding a small amount of solid pottassium to a solution of NaOH. the solution formed will be KOH and Solid Na will precipitate out.
     
  18. May 19, 2006 #17
    how do you plan on doing that without defined chemicals???
     
  19. May 19, 2006 #18
    Unfortunately, a simple displacement reaction will not work. Any Na formed will immediately react with the water in the solution to form the hydroxide. You'd have to do this with molten NaOH and have solid K on hand. With that much effort you might as well just electrolyze molten NaOH.

    In my opinion, buying sodium would be "pointless". Sure, anybody can go out and purchase lots of chemicals, but that defeats the entire purpose of amateur chemistry. The point is to learn chemistry with a hands-on approach and see the reactions for yourself. And you're correct, sometimes you'll run into incorrect reactions like yours above, and you'll have to overcome both chemical and practical problems. Those experiences and failures are invaluable, however, and they expand your knowledge base and contribute to the whole learning process.
     
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