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deltaH
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| Quote: | | Do you have any idea at all what the polarization layer around an electrode is? |
Very much so, the electric
double layer at the cathode of a brine is made up of Na+ ions, not H+, (with some chlorides trailing behind that layer) that's why
it's hard for H+ to be reduced in neutral water at high rates without a platinum cathode, cause the Na+ layer is so concentrated!
| Quote: | | ...and that they're not naively valid when an electric field is present |
What exactly are you trying to say?
For H2O <=> H+ + OH-, are you saying that the electric field changes the activation energy for this or that it changes the equilibrium position
for this or something completely different, because I don't understand your point?
| Quote: | | have no interest in spending time educating you, |
I see, yet again, how you're not trying to educate me
(rolls eyes).
| Quote: | | above nonsense counts as a hazard to others. |
Indeed, electrochemical mechanisms are VERY dangerous...
perhaps we should have this thread moved to energetics.
[Edited on 21-10-2013 by deltaH]
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Let's try this from a different angle, from this wiki entry on glassy carbon, I quote the following:
"Glassy carbon electrode (GCE) in aqueous solutions is considered to be an inert electrode for hydronium ion reduction:[6]
H3O+(aq) + e- <=> H(aq) Eo = −2.10 V versus NHE at 25 °C
Comparable reaction on platinum:
H3O+(aq) + Pt(s) + e- <=> Pt:H(s) Eo = 0.000 V versus NHE at 25 °C
The difference of 2.1 V is attributed to the properties of platinum which stabilizes a covalent Pt-H bond.[6]"
Now these are electrode potentials measured under otherwise standard conditions and already you can see just how hard it becomes to
reduce H3O+ without an electrocatalytic cathode because without a catalyst, the first step is the formation of nascent hydrogen that is not covalently
bonded to anything. Granted, glassy carbon is a crap electrode, but I'm just proving a point how it can influence the reduction of hydronium ions and
that this has nothing to do with fields and the like.
So you see, even without bringing in kinetic effects, you already have a problem with reducing H3O+ on an noncatalytic electrode in neutral brines.
Then operate at appreciable rates and you make all these problems exponentially worse as you start introducing rate limitations and your concentration
of H3O+ is completely dwarfed by Na+!
[Edited on 21-10-2013 by deltaH]
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[Edited on 22.10.13 by bfesser]
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Nicodem
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Quote: Originally posted by deltaH  | So you see, even without bringing in kinetic effects, you already have a problem with reducing H3O+ on an noncatalytic electrode in neutral brines.
...
Then operate at appreciable rates and you make all these problems exponentially worse as you start introducing rate limitations and your concentration
of H3O+ is completely dwarfed by Na+!
Finally, the 2H+(aq) <=> H2(g) Estd. = 0V is defined only when H+ concentration is 1M, platinum is used as cathode and rates are miniscule!!!
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Oh please stop it with these long facepalm replies already. Why do you bother instead of just posting something relevant with a reference, or a review like this one?
You don't understand much about electrochemistry, neither do you bother to learn first, so stop pretending, misquoting wikipedia, misunderstanding the
concept of overpotential and abusing the misnamed "electrocatalytic electrode" concept.
You don't even understand that it is water that is being reduced, not the already present H3O+ or OH- (at the
magnitude of the electric field at the electrode surface, the self-dissociation of water is irrelevant). And if it was Na+ being reduced,
then the minimum potential for water electrolysis would be heavily electrolyte dependent, which it obviously is not unless one of its ions succumbs to
an electrochemical reaction (in which case the electrolyte is no good for water electrolysis anyway). The electrolyte is there to make the solution
conductive, that's all.
…there is a human touch of the cultist “believer” in every theorist that he must struggle against as being
unworthy of the scientist. Some of the greatest men of science have publicly repudiated a theory which earlier they hotly defended. In this lies their
scientific temper, not in the scientific defense of the theory. - Weston La Barre (Ghost Dance, 1972)
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BobD1001
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Well I just ordered 16oz of propylene glycol. I already have plenty of urea as well. Ill I need is the Zinc oxide. Does anyone have any novel
approaches to convert elemental Zinc into zinc oxide. Ive already found a couple methods online, just wanting to see if anyone has a creative method.
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deltaH
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| Quote: | | Oh please stop it with these long facepalm replies already. |
As this is voiced as a request, I must
respectfully decline.
| Quote: | | Why do you bother instead of just posting something relevant with a reference, or a review like this one? |
I
don't see how posting such a broad treatise on electrolysis supports any specific arguments you have made. I on the other hand have posted a specific
link with a specific quote that is referenced to support my argument that if the cathode is inert but conductive, reduction of H3O+ becomes very hard
indeed and your argument that why would it reduce Na+ when it can 'simply' reduce H+ at 0V is simply not true under the conditions that brine
electrolysis is typically carried out with.
| Quote: | | You don't understand much about electrochemistry,so stop pretending |
Your attempts to encite and enflame me
are not working.
| Quote: | | misquoting wikipedia |
quoted word for word and referenced
| Quote: | | misunderstanding the concept of overpotential and abusing the misnamed "electrocatalytic electrode" concept. |
it would help if you explained yourself here, otherwise it's just an opinion.
| Quote: | | You don't even understand that it is water that is being reduced |
I have repeatedly stated that that is what
happens overall, but that at high rates, the pathway by which is might proceed may not be by the reduction of hydronium ions under the conditions of
brine electrolysis. YOU stated that 2H+ => H2 has a Estd of 0V and so why would Na+ be reducing. That's what started this whole discussed and my
argument was that mechanistically, this pathway is very minor under the conditions at which this process runs, i.e. neutral pH and ordinary fairly
inert cathodes.
| Quote: | | And if it was Na+ being reduced, then the minimum potential for water electrolysis would be heavily electrolyte dependent, |
No, I clearly stated that this is only significant at high rates and high rates is not where the minimum potential is measured. It
is exactly because the easy pathway is exhausted at high rates that the reaction also starts to occur on less favourable pathways. At minimum
potential, the easiest pathway dominates and no Na+ is reduced.
| Quote: | | The electrolyte is there to make the solution conductive, that's all. |
I disagree but I concede this requires
an experiment, so I propose the following that I will try to do as soon as I can:
I happen to have some choline chloride solution, which to my mind will perform poorly at high rates as there is no way that a quaternary ammonium ion
can mechanistically reduce at the cathode, so that would FORCE direct water reduction on an electrode that is not electrocatalytic and so be the least
efficient.
I will use graphite electrodes and carry out with 9V to ensure high rates (plenty of overpotential) and measure current using say 1M Choline chloride.
Then I will repeat the experiment with identical setup and 9V, but use 1M NaCl and 1M HCl and measure current for each of those. I hope to show that
the current will be highest for 1M HCl and comparable to NaCl and smallest for ChCl.
If you're done flaming, feel free to comment on such an experiment's methodology/construct which would be helpful.
[Edited on 21-10-2013 by deltaH]
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Mesa
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| Quote: Originally posted by BobD1001 | | Well I just ordered 16oz of propylene glycol. I already have plenty of urea as well. Ill I need is the Zinc oxide. Does anyone have any novel
approaches to convert elemental Zinc into zinc oxide. Ive already found a couple methods online, just wanting to see if anyone has a creative method.
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Burn it.
Throw MnO2 at it(Zinc batteries work on this concept, the outer casing is made of fairly pure zinc. I pulled a dead battery apart today and the entire
outer casing disintegrated into white flakes.)
Dissolve in HNO3, then heat the resultant salt to 125*C(more efficient than burning the base metal, careful of the fumes though.)
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papaya
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| Quote: Originally posted by Nicodem |
You don't even understand that it is water that is being reduced, not the already present H3O+ or OH- (at the
magnitude of the electric field at the electrode surface, the self-dissociation of water is irrelevant). And if it was Na+ being reduced,
then the minimum potential for water electrolysis would be heavily electrolyte dependent, which it obviously is not unless one of its ions succumbs to
an electrochemical reaction (in which case the electrolyte is no good for water electrolysis anyway). The electrolyte is there to make the solution
conductive, that's all. |
Not to defend any side, I just don't understand one thing here - "The electrolyte is there to make the solution conductive, that's all", but to
continuously conduct DC current there must be redox reactions present at the electrodes, otherwise the cell will quickly polarize and stop to conduct
(capacitor behavior, if DC is applied), isn't it? If the electrolyte is not involved in the processes on electrodes(so only water is electrolyzed),
then it'll not affect solution's conductivity for the reason I mentioned above and thus the conductivity must appear the same as for pure water (I'm
not speaking about conductivity measurements with special instrument, where high frequency AC is employed thus no polarization occurs and no
electrolysis also, consider only the DC case). Also the fact that with mercury cathode one obtains Na amalgam by electrolysis of NaCL solution
indicates that somehow Na+ can be reduced from solution.
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deltaH
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I wouldn't advise that as you might not obtain a material with a very high specific surface area
and activity. Your best best would be to follow a specific procedure that is known to work. Would probably entail something like dissolving your zinc
in dilute hydrochloric acid, then precipitating it in one fast go with the addition of stoichiometric amounts of dilute sodium hydroxide. Finally, let
it stand for a while, decant, filter and wash + repeat before drying at moderate heat.
[Edited on 21-10-2013 by deltaH]
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@papaya Interesting point, I haven't thought about it in that way. I must admit, this is also challenging my preconceptions and
understanding of the role of electrolytes, which is why I am very curious to conduct an experiment exploring this.
I couldn't at first glance find a decent paper where electrolytes were compared on the basis I'm thinking of. Usually these investigations into
electrolysis are always conducted using amazing electrodes because they want to maximise efficiency and in such a case, you just need your electrolyte
to provide the maximum ion conductivity.
All the mechanism takes place catalytically on the electrodes in the case of platinum and such and so probably the electrolyte plays no role in such
cases other than to conduct the ions to and from.
I want to know what happens when Jo Bob (you and me) does this using simple electrodes we can get a hold of. I think in that case the choice of
electrolyte may have a significant effect.
[Edited on 21-10-2013 by deltaH]
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[Edited on 22.10.13 by bfesser]
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Nicodem
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| Quote: Originally posted by papaya | | "The electrolyte is there to make the solution conductive, that's all", but to continuously conduct DC current there must be redox reactions present
at the electrodes, otherwise the cell will quickly polarize and stop to conduct (capacitor behavior, if DC is applied), isn't it?
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The electrolyte is there in order to prevent the polarization - without it the cell would be only a leaking capacitor. You must have forgotten that
that during the electrolysis of water there is a net flow of current from one electrode to the other because anions form at the cathode (2H2O + 2e-
=> H2 + 2OH-) and cations at the anode (2H2O => O2 + 4H+ + 4e-). In other words, the redox is
occurring at two physically separated places, the oxidation at one side, the reduction at other. The electrolyte serves as a bridge connecting the
reaction by moving the charge.
| Quote: | | If the electrolyte is not involved in the processes on electrodes(so only water is electrolyzed), then it'll not affect solution's conductivity for
the reason I mentioned above |
No, the nature of the electrolyte highly affects the conductivity because its ions move trough the solution.
| Quote: | | Also the fact that with mercury cathode one obtains Na amalgam by electrolysis of NaCL solution indicates that somehow Na+ can be reduced from
solution. |
Mercury electrodes have a huge overpotential for hydrogen (the highest of all metals). Overpotential is highly current density dependent and given
enough current, this can be raised enough for the sodium amalgam formation. Actually, at high enough brine concentrations and low sodium
concentrations in the amalgam, the hydrogen overpotential is higher than the sodium overpotential and thus the formation of H2 can be
almost completely suppressed. See the short explanation in Electrochemical Engineering (H. Wendt, G. Kreysa), page 294.
…there is a human touch of the cultist “believer” in every theorist that he must struggle against as being
unworthy of the scientist. Some of the greatest men of science have publicly repudiated a theory which earlier they hotly defended. In this lies their
scientific temper, not in the scientific defense of the theory. - Weston La Barre (Ghost Dance, 1972)
Read the The ScienceMadness Guidelines!
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deltaH
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| Quote: | Mercury electrodes have a huge overpotential for hydrogen (the highest of all metals). Overpotential is highly current density dependent and given
enough current, this can be raised enough for the sodium amalgam formation. Actually, at high enough brine concentrations and low sodium
concentrations in the amalgam, the hydrogen overpotential is higher than the sodium overpotential and thus the formation of H2 can be almost
completely suppressed. See the short explanation in Electrochemical Engineering (H. Wendt, G. Kreysa), page 294.
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I quoted Wiki's electrode potential for glassy carbon, an inert electrode, as being -2.1V, for the reaction
H3O+ <=> H(aq), I presume this was determined in a 1M H3O+ solution as is the norm and that ?this does matter?, the concentration dependency of
which can be calculated using the Nernst equation.
By your own arguments, this starts coming close to Na+'s reduction potential of -2.71V in a similar way than the mercury amalgam does (though even
higher). Surely when you factor in the 10-7 H3O+ concentration of neural brines, why would you not be coming close to the reduction of Na+?
[Edited on 21-10-2013 by deltaH]
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papaya
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Thanks for the answer, I believe you must be knowing what you say, but the following is little complex for me to understand (or imagine)
| Quote: Originally posted by Nicodem |
The electrolyte is there in order to prevent the polarization - without it the cell would be only a leaking capacitor. You must have forgotten that
that during the electrolysis of water there is a net flow of current from one electrode to the other because anions form at the cathode (2H2O + 2e-
=> H2 + 2OH-) and cations at the anode (2H2O => O2 + 4H+ + 4e-). In other words, the redox is
occurring at two physically separated places, the oxidation at one side, the reduction at other. The electrolyte serves as a bridge connecting the
reaction by moving the charge. |
So let's look at this more from the point of the physics. Total current = charge transferred per time. Only charge carriers in the solution are H+
and OH- that form at anode/cathode in redox processes as you wrote above - I said ONLY, because for example Na+ and SO4 2- while they exist in
solution and will move in the electric field they'll NOT give their charges to anyone else thus they CANNOT transfer charge from one electrode to the
other (otherwise we deal with Na2SO4 electrolysis, which you deny). Now, one situation - you have pure water, another - electrolyte added, do you say
that in the presence of electrolyte H+ and OH- will move FASTER to opposite electrodes (because I = q/t and "I" is higher with electrolytes, only
"real" charge carriers being H+ and OH- ), or do you say more H+ and OH- will form at electrodes per time that will move with the same speed?
Note I'm not wanting to disprove what you say, just want to clarify things for myself, and electrochemistry interpretations usually are confusing to
me, I try to hold it simple.
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watson.fawkes
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| Quote: Originally posted by papaya | | Only charge carriers in the solution are H+ and OH- that form at anode/cathode in redox processes as you wrote above - I said ONLY, because for
example Na+ and SO4 2- while they exist in solution and will move in the electric field they'll NOT give their charges to anyone else thus they CANNOT
transfer charge from one electrode to the other |
If it's charged and mobile, it will move under the action
of the electric field and thus be a charge carrier. If it's not a net charge carrier, then it only acts as a transient charge carrier. As a rule, the
more charge carriers, the higher the conductivity.
Current flow does not require the bulk mass flow of its charge carrier. In a copper wire, the charge carrier are electrons, and no one expects the
copper to move for there to be a current. If a water molecule ionizes in the bulk electrolyte, it is not the case that the two component ions need to
move individually to the electrodes to complete their half reactions. These ions will collide with other water molecules, combining and resplitting in
such a way that the charge can move faster than the component atoms.
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Mesa
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As with papaya, the vast majority of this stuff is way over my head so forgive me if the next question seems a tad stupid, but given glass becomes
conductive when molten, does the previous poster's argument indicate that sodium would be obtainable by simply heating glass with a blowtorch and
applying DC current?
Edit: Theoretical Sodium borohydride from a borosilicate glass melt?(Wikipedia hints at powdered borosilicate glass forming NaBH4 when reacted with
sodium hydride, so I assume it'd need to be done in an atmosphere of hydrogen.)
[Edited on 22-10-2013 by Mesa]
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elementcollector1
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| Quote: Originally posted by Mesa | As with papaya, the vast majority of this stuff is way over my head so forgive me if the next question seems a tad stupid, but given glass becomes
conductive when molten, does the previous poster's argument indicate that sodium would be obtainable by simply heating glass with a blowtorch and
applying DC current?
Edit: Theoretical Sodium borohydride from a borosilicate glass melt?(Wikipedia hints at powdered borosilicate glass forming NaBH4 when reacted with
sodium hydride, so I assume it'd need to be done in an atmosphere of hydrogen.)
[Edited on 22-10-2013 by Mesa] |
For the latter procedure, one must obtain sodium hydride first.
For the former, a blowtorch is not sufficient to make large quantities of glass truly 'molten'. Besides, there are much easier ways to obtain Na
(namely, the electrolysis of the molten hydroxide, which can be melted with a blowtorch but really probably shouldn't - see the thread in
Technochemistry).
Elements Collected:52/87
Latest Acquired: Cl
Next in Line: Nd
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papaya
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| Quote: Originally posted by watson.fawkes |
Current flow does not require the bulk mass flow of its charge carrier. In a copper wire, the charge carrier are electrons, and no one expects the
copper to move for there to be a current. If a water molecule ionizes in the bulk electrolyte, it is not the case that the two component ions need to
move individually to the electrodes to complete their half reactions. These ions will collide with other water molecules, combining and resplitting in
such a way that the charge can move faster than the component atoms. |
Completely agreed here - they don't have to travel the whole way rather will split next water molecule to a similar species, but what is the
electrolyte doing then? And what is transient charge carrier you mentioned - seems it can explain things when you open brackets, thanks.
<hr width="80%" />
| Quote: Originally posted by Mesa | As with papaya, the vast majority of this stuff is way over my head so forgive me if the next question seems a tad stupid, but given glass becomes
conductive when molten, does the previous poster's argument indicate that sodium would be obtainable by simply heating glass with a blowtorch and
applying DC current?
Edit: Theoretical Sodium borohydride from a borosilicate glass melt?(Wikipedia hints at powdered borosilicate glass forming NaBH4 when reacted with
sodium hydride, so I assume it'd need to be done in an atmosphere of hydrogen.)
[Edited on 22-10-2013 by Mesa] |
First question - I think it can work, and once glass becomes conductive the high current applied may be sufficient to keep the glass very hot and
conductive. Watch this made in russia experiment https://www.youtube.com/watch?v=p0E9Pd4CUBc
I can even speculate that the glass even doesn't have to be in molten state - it's enough to become ion-conductive, because I've heard about one
experiment, when they turn on a light bulb, wait it to become hot, the apply a secondary high voltage DC - negative to the filament and positive to
the glass from outside (don't know details), so after some time bulb becomes sodium "plated" from inside(electrons bombard conducting glass from
inside), has anybody heard about this?
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deltaH
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Beautiful papaya, it never occurred to me that one could do this!
The fascinating thing about such a setup is that I would guess that at the cathode you would be producing sodium metal which probably quickly reacts
with air to form sodium oxide and dissolves back into the glass, at the anode you probably produce oxygen gas. Now does this mean you can construct a
hypothetical oxygen generator with such a setup if you make a much bigger cell? Air(oxygen) in on the left, pure O2 out on the right?
Because on the cathode side you consuming oxygen from air (provide electrode is exposed enough off course) and on the anode you are releasing very
pure oxygen?
Electrochemical oxygen purification 
[Edited on 22-10-2013 by deltaH]
<hr width="80%" /> | Quote: | | ...split next water molecule to a similar species... |
What did you mean by this?
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[Edited on 22.10.13 by bfesser]
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Mesa
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| Quote: Originally posted by elementcollector1 | For the latter procedure, one must obtain sodium hydride first.
For the former, a blowtorch is not sufficient to make large quantities of glass truly 'molten'. Besides, there are much easier ways to obtain Na
(namely, the electrolysis of the molten hydroxide, which can be melted with a blowtorch but really probably shouldn't - see the thread in
Technochemistry). |
Yes I realise this, I just performed that experiment today(unsuccessfully, poor soldering job on electrode caused it to fall apart.) I was only
asking to get a better idea of the scope of non-aqueous electrolysis.
As for the requirement of sodium hydride in my theoretical synthesis, wouldn't a flow of hydrogen over the sodium as its forming on the electrode give
at least a quantitatable(I checked google, turns out this is a real word ) yield
for proof of concept?
Still all theoretical, I'm having enough trouble with the molten salt electrolysis as it is. And I find pyrometallurgy much easier and more
accessable given how much more forgiving it is in terms of temperature ranges.
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watson.fawkes
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It's ionizing and deionizing, or rather forming transient neutral
species with short lifetimes. Remember that a liquid is a mess of thermal motion as well. All it takes is a geometric bias for a given ion to form a
neutral species nearer to one electrode and to split apart nearer to the other for it to act as a current carrier.
An energetic difference between an ionic type of bonding and a covalent one is the activation energy of transitioning between bonding and non-bonding
states. Ionic bonding has much lower activation energy, so you'd expect more transient species to form as a result of random collisions. (You could
put hydrogen bonding along this continuum as well, with ionic in the middle.)
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papaya
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Aha, turns out it's some sort of ELECTRONIC conduction in the solution, where the electron exits from cathode reducing water to OH- and from the
latter it sometimes springs onto other molecule/ion and this way travels until gets the anode- more ions, higher probability to catch a "bus" to the
next station! Because I can't call it something else than electronic conduction what is described (don't get angry if I misinterpreted something),
prove me wrong.
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deltaH
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@Papaya
In nicodem's general review link he posted, there's this graph:
Notice that the current is very small until you hit the minimum potential for electrolysis to kick off... after that the current increases
proportionally, so I don't think 'electronic conduction' occurs significantly in the electrolyte alone. Possibly the activation energy for electrons
hoping on and off ions in much larger than say in a metal where there is a 'sea' of loosely bound electrons.
[Edited on 22-10-2013 by deltaH]
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Nicodem
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| Quote: Originally posted by papaya | | Aha, turns out it's some sort of ELECTRONIC conduction in the solution, where the electron exits from cathode reducing water to OH- and from the
latter it sometimes springs onto other molecule/ion and this way travels until gets the anode- more ions, higher probability to catch a "bus" to the
next station! Because I can't call it something else than electronic conduction what is described (don't get angry if I misinterpreted something),
prove me wrong. |
No, no and no. It is not the electrons that transmit the current (except in the electrode and its surface). Like I already said, the ions of the
electrolyte do this - that's why they are needed.
There really is no need to imagine wild things. You can imagine it as a simple train of charged particles (anions and cathions) each traveling to the
corresponding electrode in order to maintain the local neutrality. The locomotive that forces them to travel is the electric field. This way every
increase in charge at the electrode due to a molecule of water being electrolyzed is compensated by the slight movement of the equally charged
electrolyte ions toward the opposite electrode. The ionic movement is dictated by the neutrality demand. Without this, the electrodes would just
polarize due to the first few H+ and OH- formed at the electrode. The charge would not be able to be transmitted trough the
medium due to lack of carriers and the electrolysis would stop immediately - you would have a capacitor.
Here is good page with short explanations:
http://chemed.chem.purdue.edu/genchem/topicreview/bp/ch20/fa...
There are some animated tutorials on electrolysis around. Here is one, but it only poorly addresses the role of the electrolyte:
http://www.youtube.com/watch?v=4bdFKyVM8fk
Try searching for a better tutorial, there must be some around, though the other I found were pretty much pathetic or simplistic to a degree of
unfactuality.
…there is a human touch of the cultist “believer” in every theorist that he must struggle against as being
unworthy of the scientist. Some of the greatest men of science have publicly repudiated a theory which earlier they hotly defended. In this lies their
scientific temper, not in the scientific defense of the theory. - Weston La Barre (Ghost Dance, 1972)
Read the The ScienceMadness Guidelines!
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watson.fawkes
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| Quote: Originally posted by Nicodem | | It is not the electrons that transmit the current (except in the electrode and its surface). Like I already said, the ions of the electrolyte do this
- that's why they are needed. |
If this were true, then there would be zero steady state current when an
electrolytic cell is run at significantly lower than threshold voltage, say, 10 mV. Every such cell, though, has a mixture of both ionic and
electronic conduction. Which is dominant depends upon lots of conditions: electrolyte composition, electrical field strength, temperature. In a 10 mV
potential cell, the majority carrier in the steady state is indeed electronic. Ions act as a minority carrier (and not a non-carrier) because of their
thermal motion.
In an ideal electrolysis cell, the bulk electrolyte has zero resistance, all the potential drop occurs in the polarization layers, only the ionic
additions to the solvent react at the electrodes, etc. In that situation you would indeed have zero electronic conduction. In a well-run practical
electrolysis cell, you have ions as the majority carrier and electrons as a minority carrier. In a not-so-well run one, there's more current than
ionic conduction can provide (rate limited by ion drift velocity), and you get electronic conduction and electrolysis of the solvent.
The solvent itself also does provide some amount of electronic conduction. The conductivity of very pure water is small, around 10^-7 S/m, but that of
insulating oils is around 10^-14 S/m and lower. You might be able to use tritiated water introduced at one electrode to measure the ratio of
electronic to ionic conduction in such a situation, but I don't know if such work has been done.
Edit: added the units to conductivity values
[Edited on 2013-10-23 by watson.fawkes]
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Nicodem
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| Quote: Originally posted by watson.fawkes | | If this were true, then there would be zero steady state current when an electrolytic cell is run at significantly lower than threshold voltage, say,
10 mV. Every such cell, though, has a mixture of both ionic and electronic conduction. |
You are correct. The conductance that obeys Ohm's law is expected to be electronic in nature. This background current is however independent of the
electrolysis reaction (it does not participate) and is generally magnitudes lower than the electrolysis working current. I should have stated that I
was referring to the current involved in the electrolysis (the current that is generated due to H+ and OH- formation at the
electrode). But then again, I'm not so sure anymore that papaya was asking about the electrolysis or in general, in which case you explained it
better.
…there is a human touch of the cultist “believer” in every theorist that he must struggle against as being
unworthy of the scientist. Some of the greatest men of science have publicly repudiated a theory which earlier they hotly defended. In this lies their
scientific temper, not in the scientific defense of the theory. - Weston La Barre (Ghost Dance, 1972)
Read the The ScienceMadness Guidelines!
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Upsilon
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How well does PC work as a solvent in comparison to water when dissolving halide salts? Being able to dissolve certain substances is one thing, but to
what extent is another.
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elementcollector1
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Just put a bunch of (hydrated) yellow CeCl3 into some pure propylene carbonate. So far, it does not appear to be dissolving, but time will tell. Would
cooling things down make it dissolve more?
Elements Collected:52/87
Latest Acquired: Cl
Next in Line: Nd
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blogfast25
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Stirring. Mild heating. That should do it.
But water is the enemy of electroplating REs. And 6 mol of water per mol of CeCl3 is A LOT!
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