Horace, My comments below, some things are still wrong
2009/11/25 Horace Heffner <hheff...@mtaonline.net>: > Gad. It still isn't right! Corrections below. I have vertigo at the > moment and can't think straight. I've actually done half of this > experiment, though decades ago, and it is interesting how the concentration > gradient wanders, it doesn't follow what you would expect for any kind of E > field. > > > On Nov 23, 2009, at 2:48 AM, Michel Jullian wrote: > >> See: http://sci-toys.com/scitoys/scitoys/echem/fuel_cell/fuel_cell.html >> >> I had no idea an ultraclean rechargeable battery could be done so simply! >> >> Supplies: >> <<- One foot of platinum coated nickel wire, or pure platinum wire. >> Since this is not a common household item, we carry platinum coated >> nickel wire in our catalog. >> - A popsickle stick or similar small piece of wood or plastic. >> - A 9 volt battery clip. >> - A 9 volt battery. >> - Some transparent sticky tape. >> - A glass of water. >> - A volt meter.>> > > It seems to me a small amount of lye would help the reaction along. No > matter, the intent is apparently not to create a working cell, i.e. generate > power, it is merely to generate a voltage. > > I see they sell the wire for $14.41 plus shipping. A bulk source for wire > and mesh might be: > > http://www.gerarddaniel.com/ > > > >> H2 and O2 are produced by short electrolysis runs, after which the >> bubbles clinging to the electrodes are catalytically recombined by the >> electrode surface material (platinum) to generate electricity :) >> >> 1/ The article features nice "explanations" of how it works, but how >> does it _really_ work? In particular, in the generating (fuel cell) >> phase, they don't say what makes the positive hydrogen ions climb >> "uphill" from the negative electrode to the positive one, anyone can >> explain this miracle? ;-) >> >> 2/ It seems to me a much higher capacity (and perhaps even practical) >> rechargeable battery could be made by using a hydrogen >> absorbing/desorbing material e.g. Pd for the negative electrode, and >> by making gaseous oxygen available at the anode. Storing the latter is >> not required of course, O2 from the air is fine... maybe a floating >> support which would keep a grid or flat serpentine shaped positive >> electrode at the surface of the water or just below? >> >> Michel > > The explanation looks bogus to me. I think the cell works by reversible > reactions, not recombination. > > Bockris states that conduction in an electrochemical cell in the volume > between the interface layers is almost entirely due to concentration > gradients. That is because almost all the potential drop is in the interface > layers themselves. The E field in the bulk of the cell is very small. > > I expect the cell actually operates by creating even *more* bubbles, not > consuming the gas already there in the form of bubbles. > > In the course of the brief electrolysis by battery, the volume of water > around the *anode* is preferentially filled with H3O+ ions, as the OH- ions > release their electrons and form O2 and H2O2, and the volume around the > *cathode* is filled with OH- ions as the H3O+ ions present at the cathode > surface are electrolyzed. This can actually be viewed by use of a dilute > electrolyte, plus a pH indicator like phenolphthalein, which is colorless in > acidic electrolytes, and pink in basic solutions. To do this first add the > (liquid) phenolphthalein to distilled water. Connect the battery. To view > the creation and migration of OH- ions: add a little bit of boric acid to > the water, and stir. Repeat the process until you can see the electrolyte > turns pink in the vicinity the *cathode* (- electrode) once the electrolyte > settles down. Boric acid was chosen because it is commonly available from > pharmacies. To view the creation and migration of H3O+ ions add a little > bit of lye to the water and stir. Repeat the process until you can see the > electrolyte is pink, but when the electrolyte settles down you can see the > volume around the *anode* (+ electrode) gradually turing clear. It can take > a little fooling around with concentrations to get the effect to work > quickly and dramatically. The diffusion occurs slowly but at a clearly > visible pace. I agree with the above paragraph now, but putting it right has broken your explanation for the generating phase two paragraphs below. > You can demonstrate the reversibility of the reactions by reversing the > battery. Note, however, that the diffusion occurs in a somewhat random > manner. It doesn't typically blossom out in a perfectly spherical or > cylindrical manner (depending on the electrode shape). Reversing the > reaction is thus not a perfect process either. I tried some of this decades > ago in a feeble attempt to make a display technology. I got a nice red > stream of ions coming from a copper anode in a basic solution. > > In any case I doubt it is actually recombination that causes the potential > at the electrodes. It is the presence of the high concentration of ions in > solution that makes the residual potential when the battery is disconnected. > The H3O+ ions take on electrons through the wire originally releasing > hydrogen at the site where the hydrogen was generated, No, the ions there are OH- ions > the anode, thus > making *more* hydrogen bubbles. Similarly, the OH- ions donate electrons to > make H2O2 and *more* O2 at the site where O2 was generated prior. No, the ions there are H3O+ ions. Want to give it another try? :) Michel > The meter is probably a 10 megohm meter, meaning registering the 2 V > potential requires generating 0.2 microamps of current, and thus 0.4 > microwatts of power. Not much of a fuel cell! > > It would be interesting to run the current for a while until a significant > concentration gradient can be visualized, and then disconnect the battery to > see what effect the current generated through the meter has on the visible > gradients. > > Note that the concentration gradients of H3O+ and OH- particles does not > necessarily require an E field to maintain them, provided there are other > radicals in the electrolyte. Salt buffers can be used to increase > conductivity without driving Ph to extremes. The presence of additional > radicals can balance the charges to neutral, or to match the E field in the > electrolyte. For example, if the electrolyte is NaOH, the Na+ can > redistribute to neutralize the charges. If boric acid is used, the B(OH)4- > radical will balance the charges. > > Best regards, > > Horace Heffner > http://www.mtaonline.net/~hheffner/ > > > > >
