erik-k
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Making a Cu-Al battery
Hi - new member here, looking for help with my attempt to make a battery based on Copper (II) Chloride and Aluminum. Copper plates out, chloride ions
go to the Aluminum side, Al dissolves into AlCl<sub>3</sub>, and sends electrons to Cu side. The only problem is that rather than
generating the two volts it "should" if one goes by standard reduction potentials, it outputs more like .85 to .9 when measured. The aluminum was from
flashing, which I heated enough to destroy it's coating and then abraded it's surface clean. The copper chloride was from a solution I'd made that was
concentrated to the point of being brown, and the electrode a piece of wire, sanded to assure a clean surface. The salt bridge was paper, folded 4
sheets thick and soaked before use. So, any suggestions about what I'm doing wrong? The voltage deficit is on par with dissolved oxygen and acid
hydrogens being reduced to water, but I'm not terribly sure about this.
I'm thinking about using ethanol instead of water as the electrolyte, as both Copper (II) and Aluminum chlorides are quite soluble (50g/100g) in it.
Am I correct in thinking it would solve the problem of the solvent taking part in reactions?
Thanks for help in advance!
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Nerro
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I forgot how to calculate it (nernst?) but I do remember it's also dependant on the temperature and concentrations of both reagents.
#261501 +(11351)- [X]
the \"bishop\" came to our church today
he was a fucken impostor
never once moved diagonally
courtesy of bash
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Elawr
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Couple of problems with aluminum: first is the oxide coating which reforms almost immediately on exposure to O2. Mercury will form an amalgam with the
aluminum and keep the oxide coating from protecting the aluminum, thus revealing the second problem. Aluminum is so reactive that will try to give
its electrons to H20 molecules rather than behave and do what you want it to - give electrons to your copper salts and giving you a nice electric
current. High or low pH will also remove the protective oxide coat, but you still have the same problem.
If you could work around this, you could have a super battery, what with all the energy aluminum releases when oxidized. You need an electrolyte
harder to reduce that your aluminum. I don't know about ethanol, may not be very conductive, but then I've never tried.
Pretty interesting problem your working on there, I'd be eager to find out what you learn!
[Edited on 15-12-2006 by Elawr]
[Edited on 15-12-2006 by Elawr]
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Nerro
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Would you be prepared to test this with several solvents?
Methanol, ethanol, glycerol for instance.
And how about CuSO4 and zinc, zinc doesn't go all crazy in water, the reaction with water will occur but very very slowly.
The reaction with Zn should give you 0.91V, so if you make a layercake battery from it you could have quite enough power from that.
[Edited on Fri/Dec/2006 by Nerro]
#261501 +(11351)- [X]
the \"bishop\" came to our church today
he was a fucken impostor
never once moved diagonally
courtesy of bash
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12AX7
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You also need a "gravity cell". Take clean water (with a little electrolyte for conductivity) and add copper sulfate and an inert electrode to the
bottom. Put the aluminum/zinc on top so it's as far as possible from the dense copper-rich layer.
Tim
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garage chemist
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Why do you use a solution of copper salt as your electrolyte? It's the worst electrolyte you can use because copper will plate out on the more
reactive metal even when no current is flowing, and therefore your battery will rapidly discharge itself even when no current is flowing.
And I agree that aluminium is a bad choice for a battery because of its oxide layer and rapid reaction with water if attempts to remove the oxide
layer chemically are made.
Magnesium is the better choice here, giving even more voltage than aluminium theoretically should.
Magnesium is all around us, just like aluminium- Mg firestarters being the most prominent example.
You will have to use a magnesium salt or other neutral solution as electrolyte.
A small battery from magnesium ribbon and copper wire in NaHSO4 solution (bad choice, I know, due to self- discharging- but it makes possible the
highest current output) is powerful enough to make a small lightbulb glow dimly.
[Edited on 15-12-2006 by garage chemist]
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erik-k
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I mentioned ethanol as an alternative because both copper and aluminum chlorides are soluble in it, so creating a conductor shouldn't present a
problem in that respect. Copper chloride is soluble in methanol, too, but I'm still looking for info on AlCl<sub>3</sub> vs methanol.
Another interesting thing is that if I use something other than water, the reduction potentials could change. <a
href="http://pubs.acs.org/cgi-bin/abstract.cgi/inocaj/1962/1/i01/f-pdf/f_ic50001a030.pdf?sessid=6511">Apparently</a>, the reduction of copper
(II) to copper is +.575V in ethanol unless I read the top paragraph wrong. No such luck finding similar data on Al in Ethanol, though.
Regarding the need for the solvent to be more reducing than aluminum, I think ethanol fits the bill. Wikipedia's sodium article describes ethanol's
reaction with Na as "much slower than water's," implying that ethanol's reduction potential falls close enough to Sodium's for the reaction to not be
violent. This probably places it at least at Aluminum's.
The gravity cell idea is also a good one. I noticed when I was filtering some CuCl<sub>2</sub> that the drops of more concentrated
solution tended to sink rather than mix, but thought nothing of it at the time. So next time, I'll put the copper solution on bottom.
As for why I'm using a copper salt, because they're easy to make from scratch. Electrolyze a solution of sodium bicarbonate with a carbon anode and
copper cathode and you'll have copper carbonate flaking off the metal. Then just dissolve in the acid of choice to make copper salts and let the
CO<sub>2</sub> bubble away.
Won't a copper solution want to plate itself out on any metal with a lower reduction potential? I agree that Al is a rather bit of a pain to use as a
battery, in any case - hence my desire to use ethanol. It's true that Mg would generate almost twice the total energy output (there are commercial
Mg-Cu reserve batteries, 1.6V), but I don't have any Mg on hand ^_^. I also notice that these batteries have the same ~1.1V deficiency as I see in my
cell (2.0 - .85). But imagine replacing the Al with Mg, and having the cell output hit 2.7V!
Thanks for all the replies - I'll try and get some ethanol and do a test cell. Do you think common 95% be good enough, or do I have to find absolute
ethanol?
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woelen
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Keep in mind, that dissolving CuCl2 in methanol or ethanol might not yield a conductive solution. CuCl2 is quite covalent, and in those solvents it
might be you simply get molecules of CuCl2. You can see that easily. If a solution of CuCl2 becomes yellow/brown, then it dissolves as molecules CuCl2
(I know it does in acetone).
I noticed that even the hydrated salt Cu(H2O)2Cl2 (which itself is a complex with 4-coordinated copper) can dissolve as anhydrous compound in many
solvents. The water apparently is expelled from the complex and taken up by the solvent.
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erik-k
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First test using Ethanol - Progress but no cigar yet.
I just did my first test of a cell using ethanol - An improvement, but the alcohol brings it's own problems.
I managed to make a fairly pure solution of copper chloride in water which I then heated it to drive the water off, leaving a coating of dark orange
crystals (dehydrated CuCl<sub>2</sub> . I added ~3 to 5mL of the
alcohol to dissolve the crystals (neither of the contaminants is soluble in ethanol), and got a nice blue-green solution, indicating that the molecule
did break up.
I made a small paper "boat" to act as the salt bridge, and put in a few salt crystals and a few drops of ethanol to make the aluminum side of the
cell. This is when the problem became evident: Ethanol's adhesion and cohesion are completely different from water's, and it spread through the sheet
of paper at least 20 times faster than the water did in my previous test. So, I need a new material to play salt bridge.
When I put the aluminum side in the copper end cap that held the CuCl<sub>2</sub> solution, my voltmeter briefly registered a potential of
about 1.2V. If nothing else, this alone is an accomplishment. The potential fell rapidly as the solutions flowed through (and possibly over) the paper
barrier. What else would make a decent salt bridge? Cardboard soaked in Ethanol? Perhaps an ammonium chloride bridge, arranged like [Cu metal /
CuCl<sub>2</sub> solution | paper | NH<sub>4</sub>Cl | paper | AlCl</sub>3</sub> / Al metal]? Ammonium chloride
isn't particularly soluble in Ethanol (.76g/100g), so it would be easy to have the middle bridge packed with NH<sub>4</sub>Cl grains.
Well, I'm making progress at least.
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