White Yeti
National Hazard
Posts: 816
Registered: 20-7-2011
Location: Asperger's spectrum
Member Is Offline
Mood: delocalized
|
|
Really need feedback on a battery design.
Hello everyone.
I've been designing an electrochemical cell that uses two energy dense materials as oxidizing and reducing agents. Here's the theory:
At the anode, aluminium reacts with hydroxide anions:
Al + 3OH- ----> Al(OH)3 +3e- +1.66V
At the cathode, electrons come from the load and react with perchlorate anions and hydrogen ions in solution:
ClO4− + 2 H+ + 2 e− <-----> ClO3− + H2O +1.2V
Then once all the perchlorate ions are used up, another reaction kicks in:
ClO3− + 2 H+ + e− <-----> ClO2(g) + H2O +1.18V
Another reaction takes care of the chlorine dioxide to form hypochlorite:
ClO2(g) + H+ + e− <-----> HClO2(aq) +1.19V
Last but not least ,three final reactions reduce chlorine to +1, 0 and finally to -1:
HClO2(aq) + 2 H+ + 2 e− <----> HClO(aq) + H2O +1.67V
2 HClO(aq) + 2 H+ + 2 e− Cl2(g) + 2 H2O +1.63V
Cl2(g) +H2O(l) <-----> HCl(aq) + HClO(aq)
The theory looks simple, but making such a cell almost looks like a failed venture. What kind of separator do I use? Do I need to use a salt bridge?
Should I decrease the pH at the cathode to increase reaction rate? What kind of cathode do I use, graphite? Or copper wool or mesh?
I'm mainly throwing this idea out into the open. I don't have any perchlorate right now, and I'll have to wait quite some time before I can get any.
So please, feel free to try this out if you think this idea might work.
Thanks for reading.
[Edited on 10-2-2011 by White Yeti]
|
|
|
hissingnoise
International Hazard
Posts: 3940
Registered: 26-12-2002
Member Is Offline
Mood: Pulverulescent!
|
|
| Quote: | | The theory looks simple, but making such a cell almost looks like a failed venture. |
And looks totally bonkers since there are simpler routes to chlorine-water . . .
|
|
|
White Yeti
National Hazard
Posts: 816
Registered: 20-7-2011
Location: Asperger's spectrum
Member Is Offline
Mood: delocalized
|
|
I'm not looking for a simple route. Explain yourself.
This is a BATTERY DESIGN, not a chlorine generator. I thought I made that perfectly clear.
I originally thought of this design because both perchlorate and aluminium are substances that require much energy to be manufactured. Thus, reacting
them both in a battery would release tremendous amounts of stored energy. A simple parallel to this is flash powder (whether it uses chlorate or
perchlorate it doesn't matter).
|
|
|
hissingnoise
International Hazard
Posts: 3940
Registered: 26-12-2002
Member Is Offline
Mood: Pulverulescent!
|
|
Sorry WY, I skimmed through your post while still only half-awake . . .
|
|
|
Chordate
Hazard to Others
Posts: 108
Registered: 23-2-2011
Member Is Offline
Mood: No Mood
|
|
According to my table that "anode" reaction has a reduction potential of +2.31 volts, and the cathodes there are all smaller values. I think you need
to go check your math, because I don't think that cell is going to do what you want it to.
|
|
|
blogfast25
International Hazard
Posts: 10562
Registered: 3-2-2008
Location: Neverland
Member Is Offline
Mood: No Mood
|
|
Potassium perchlorate is very poorly soluble in cold water, potassium chlorate slightly better. Why not attempt to put together a cell with an Al
anode and a graphite cathode. Try a lightly alkaline electrolyte for the anode, a saturated, slightly acidic KClO3 solution for cathode electrolyte
and connect the electrodes over a voltmeter and a salt bridge? Just to see if you get any potential?
[Edited on 2-10-2011 by blogfast25]
|
|
|
AndersHoveland
Hazard to Other Members, due to repeated speculation and posting of untested highly dangerous procedures!
Posts: 1986
Registered: 2-3-2011
Member Is Offline
Mood: No Mood
|
|
Perchlorate will not will not funtion as an effective oxidizer in such a cell (unless perhaps if heated to boiling).
The reaction of perchloric acid (at least under 70% concentration) with aluminum does not reduce any portion of the perchlorate. In this
regard, perchlorate is surprisingly resistant to reduction, as even nitrate is partially reduced in the same situation. Similarly, neither will Zn+HCl
reduce KClO4, although it reacts KNO3.
For this reason, it may be best to just use chlorate in the cell, rather than perchlorate. Sodium chlorate is also more soluble than potassium
chlorate.
I'm not saying let's go kill all the stupid people...I'm just saying lets remove all the warning labels and let the problem sort itself out.
|
|
|
White Yeti
National Hazard
Posts: 816
Registered: 20-7-2011
Location: Asperger's spectrum
Member Is Offline
Mood: delocalized
|
|
Thanks for the feedback.
Sounds like I need to use chlorate instead of perchlorate. I was actually planning on using sodium perchlorate, which is much more soluble than its
potassium counterpart.
Blogfast25: I see what you are saying, but what should I use for the salt bridge? It seems like the trickiest part of building such a cell is
controlling pH and the diffusion of OH- ions from one half cell to the other. The way I looked at it, was that water would react with (per)chlorate
and electrons at the cathode to form more water and OH- ions that would migrate to the anode and react with aluminium. So, do I build a salt bridge
made out of a hydroxide or will a weak base such as NaCO3 suffice?
|
|
|
AndersHoveland
Hazard to Other Members, due to repeated speculation and posting of untested highly dangerous procedures!
Posts: 1986
Registered: 2-3-2011
Member Is Offline
Mood: No Mood
|
|
Using a base in the salt bridge may prove rather difficult, considering that alkali chorate will not react with aluminum on its own. The solution will
likely need to contain some dilute acid. Ideally, dilute perchloric acid should be used, but nitric acid likely would also work, although it might
complicate the reactions. HCl should not be used since it is oxidized by chlorate. H2SO4 does not seem to readily dissolve aluminum metal.
I'm not saying let's go kill all the stupid people...I'm just saying lets remove all the warning labels and let the problem sort itself out.
|
|
|
blogfast25
International Hazard
Posts: 10562
Registered: 3-2-2008
Location: Neverland
Member Is Offline
Mood: No Mood
|
|
I just wouldn't worry about that so much just yet. If you can close the circuit with a charge carrying solution and your redox reactions work then you
should get a significant cell potential. That may not last very long because the actual amount of energy a cell can deliver depends enormously on its
construction. But it's perhaps better to try and produce 'proof of concept' before thinking about the patent office yet...
[Edited on 2-10-2011 by blogfast25]
|
|
|
White Yeti
National Hazard
Posts: 816
Registered: 20-7-2011
Location: Asperger's spectrum
Member Is Offline
Mood: delocalized
|
|
OK. I extracted some potassium chlorate from matches this morning. I'm going to try this out very soon. I'll share my results as soon as I have the
chance
Wish me luck!
|
|
|
White Yeti
National Hazard
Posts: 816
Registered: 20-7-2011
Location: Asperger's spectrum
Member Is Offline
Mood: delocalized
|
|
Results
The results are in for those who are interested:
Short circuit current:
Dry charge: 0mA
Wet charge: <1mA
Salt water charge: ~3mA
Salt water and potassium chlorate: ~8mA decreasing down to 5mA after 2-3 minutes.
Salt water, potassium chlorate and 3% hydrogen peroxide (just for the hell of it): Rises from 5-18mA and stabilises
Salt water, potassium chlorate H2O2 and iron(II) acetate, (to make dillute fenton's reagent): rises from 18-20mA
Conclusion, the crude design of the cell showed that adding chlorate marginally increases the performance of an aluminium battery. A more
sophisticated cell design is required to benefit from the oxidising power of chlorate to any extent.
|
|
|
![[*]](images/xpblue/default_icon.gif){kind=icon}
