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This is not a particularly elegant layout for the equation .... can anyone improve on it? David Martland 09:47 Dec 5, 2002 (UTC)

My chemistry textbook, Zuhmdal's Chemistry (third edition) seems to disagree with this equation, listing instead of "[red]/[ox]" the reaction quotient Q, which is equal to [red]^x/[[ox]^y, where x and y are the coefficients of the reduced species and the oxidized species, respectively, in the balanced redox reaction.

Am I simply misanalyzing something?

<<Next comment>> I've been seing variations in the Nernst equation. Some using log functions and some using ln functions. Also I've seen one where everything gets multiplied by 2.303 and I don't know which if any of them is correct!!! <</End comment>>

2.303 ~ 1/log_10(e) . This comes from the change of base of the log from e to 10. Ash Lightfoot 04:23, 14 March 2006 (UTC)[reply]

so what...? it doesn't matter whether I use log or ln?? For me this sounds like "it doesn't matter whether you multiply or divide". 167.83.109.22 (talk) 14:04, 16 July 2008 (UTC)[reply]

I agree with the comment that the equation should have Q above instead of concentrations. The concentrations imply that there are no protons in the redox equation, and that only one species of each half reaction is an ion. This does not have to be true. The use of Q is much more general. Olin 13:49, 17 March 2006 (UTC)[reply]

I think the page is fine.

This is a complex derivation for even PhD's in Material Science, so I don't think it CAN be improved upon. Let it be. 137.99.76.228 (talk) 15:40, 15 February 2008 (UTC)[reply]

The equation for biological membranes is misleading in two ways: - We normally express the Nernst equation at 37 degrees C rather than room temperature - It's ONLY [out]/[in] when you are considering CATIONS. If you're considering the effect of anions like chloride, you use [in]/[out].

Will update when I have time to check out my sources

Confuseddave 09:27, 8 February 2006 (UTC).[reply]

I don't agree with this comment--Nerst equations are used in many more fields than biology, and 25 degrees C is generally the standard. This stays true for galvanic cell testing, and any general corrosion or oxidation/reduction computation. So yes, if the equation is being used for a biological study, 25 degrees C may not be the correct temperature--but that's why there's a term for temperature (T). 137.99.76.228 (talk) 15:37, 15 February 2008 (UTC)[reply]

Activities not Concentrations

The Nernst Eq'n is first given with activities and then concentrations. The two are not equivalent, except at very high dilution. I believe the formulation with activities is correct, although the Nernst equation is sometimes given with a formal electrode potential and concentrations rather than a standard electrode potential and activities. However, if concentrations are used the conditions under which the formal potential was measured must be stated (supporting electrolyte, concentration of ions, etc.) Can someone sort this out (with appropriate references)? Ahw001 13:27, 23 March 2006 (UTC)[reply]

For more details see external link Ahw001 07:55, 24 March 2006 (UTC)[reply]

Formal Electrode Potential

Formal potential, E°', as applied in the Nerst equation, is the standard potential, set for the working electrode (WE) for a particular reversible reaction occurring at that WE. The formal potential is found half way between the two peaks in its cyclic voltammogram, where at this point the concentration of O (the oxidized substrate) and R(reduced substrate) at the electrode surface are equal. The difference between the formal potential E°’ and normal E°, is simply because in this case concentrations are being used, not activities. It is still the stardard set for the WE for a particular reaction.

[O]σ= [R]σ σ = at surface

"Electrochemical Methods" by A.J. Bard and L.R. Faulkner "A First Course in Electrode Processes" by D. Pletcher Ahw001 06:48, 28 September 2006 (UTC)[reply]

Too complex

"For simplicity, we will consider a solution of redox-active molecules that undergo a one electron reaction"

This is not simple unless you have just taken a chemistry course. Links, etc needed.

More on Q versus [red]/[ox]

This was alluded to by Olin and Confuseddave above, but I think there remains unresolved issues with the first equation in the article, which would be fixed by replacing ${\frac {a_{\mbox{Red}}}{a_{\mbox{Ox}}}}$ by Q, the reaction quotient. According to the (intro) chem that I know, this equation is a valid way to write the Nernst equation:

E=E^{0}-{\frac {RT}{zF}}\ln Q

But I'm less sure about

E=E^{0}-{\frac {RT}{zF}}\ln {\frac {a_{\mbox{Red}}}{a_{\mbox{Ox}}}}

For example, take the reaction

CATHODE: solid-A + electrons --> anion-A

ANODE: anion-B --> solid-B + electrons

Then the reduced species is [red]=[anion-A], and the oxidized species is [ox]=[anion-B], and clearly Q=[red]/[ox]. So no problem here; this case works.

But now suppose instead we have:

CATHODE: Cation-A + electrons --> solid-A

ANODE: solid-B --> Cation-B + electrons

then the oxidized species is [ox]=[cation-B] and the reduced species is [red]=[cation-A]. Then Q = [cation-B]/[cation-A], but [red]/[ox]=(1/Q)!

Therefore, I propose switching the equation in the intro of the article to the version with Q, and taking out the version with [red]/[ox] altogether. Unless I'm confused. Thanks! --Steve (talk) 23:56, 4 May 2008 (UTC)[reply]

UPDATE: Oh I see, this is the Nernst equation for how a reduction potential depends on concentration. In that case, I'd say the first sentence is wrong when it says "cell (or half-cell)", rather than just half-cell, and ditto for the definition of z. It would be an improvement, I think, to start with the Q version, and have the reduction-potential version as a special case. But at the very least, it could be made much clearer that this is the equation for reduction potential, not for the E of a whole cell or anything else. Anyway, would anyone object to a Q-centric opening? --Steve (talk) 01:27, 5 May 2008 (UTC)[reply]

UPDATE 2: Having heard no objections in the last week, I've edited the intro accordingly. Comments? --Steve (talk) 04:09, 13 May 2008 (UTC)[reply]

Example Calculation

This page would be made clearer with an example calculation. 58.6.101.63 (talk) 01:01, 29 October 2008 (UTC)[reply]

And also it could list the reverse potential of the 4 most abundant ions in neurons at rest. XApple (talk) 00:32, 12 January 2009 (UTC)[reply]

Sign

What is the reference for the Nernst_equation#Nernst_potential? Is it the inside potential with respect to the outside or vice versa?

Bless sins (talk) 23:05, 6 October 2010 (UTC)[reply]

Notation for the activities

I find it weird that the equation for the half-cell reduction potential is written like this:

E_{\text{red}}=E_{\text{red}}^{\ominus }-{\frac {RT}{zF}}\ln {\frac {a_{\text{Red}}}{a_{\text{Ox}}}}

Because it says that $a_{\text{Red}}$ stands for 'activity of the reductant' and $a_{\text{Ox}}$ for 'activity of the oxidant'. This is strange because the equation refers to only half of a redox reaction, which means only the reductant OR the oxidant is present. It would be clearer if the article made reference to a model reduction half-reaction:

Ox + n electrons --> Red

Also, in the first statement of this article: "Nernst equation is an equation that can be used (in conjunction with other information) to determine the equilibrium reduction potential of a half-cell in an electrochemical cell" shouldn't it say that the Nernst equation can be used to determine the potential of a half-cell, be it at equilibrium or not?

— Preceding unsigned comment added by 178.199.110.247 (talk) 21:32, 30 November 2011 (UTC)[reply]

Involvement in cold fusion - reasoning by J. Huizenga

text brought from talk cold fusion

One of the critics (J Huizenga) of cold fusion said something about F&P missinterpreting Nernst equation. Some details should added concerning the kind of missinterpretation involved according to Huizenga's view.--5.15.200.238 (talk) 18:28, 30 August 2013 (UTC)[reply]

That's from Huizenga 1993, pp. 33, 47. From page 33. "Based on his interpretation of the Nernst equation (taught in college freshman chemistry courses), Pons concluded that the deuterium pressure in the palladium cathode was equivalent to a hydrostatic pressure of approximately 10²⁷ atmospheres! It seems that it was this incorrect conclusion which led Fleischmann and Pons to believe that, in a palladium cathode, the deuterium nuclei would be forced together close enough to fuse. The Nernst equation, applicable under equilibrium conditions, was used to relate the overpotential in an electrochemical cell to deuterium fugacity. If this simple, but erroneous procedure is followed, a large overpotential does give a high deuterium fugacity. The use of the Nernst equation, however, for the overall deuterium evaporation reaction under conditions of large values of the overpotential for estimation of the pressure of deuterium in the palladium is inappropriate [J. Y. Huot, J. Electroche. Soc 136 631 (1989)] The actual hydrostatic pressure of deuterium in the palladium is many orders of magnitude less than estimated by this faulty procedure."

Another book cites this paragraph from Huizenga, and adds: "'The general message is clear: in using our pre-equilibrium arguments to establish Eq. (2.43), we must ensure that the pre-equilibrium is actually operative if the equation is to be meaningful. If large currents are being drawn, this may not be the case. Accordingly, the reader might note the need for caution in applying the analysis derived in Section 2.8. He/she might also wish to review he discussion of the Nernst equation in Chapter 1 of this book." ^[1]

There was a comentary in a German journal[1] (search "Nernst")

Ebert K; Nachr. Chem. Tech. Lab. 37 (1989) 470 (in German).

"Elektrochemisch induzierte Fusion von Deuterium" (Electrochemically induced

fusion of deuterium).

** Early comment, reporting on the initial F&P press conference and the paper

in JEC. The article is not very critical, raising only a slight doubt as to

the applicability of the Nernst equation to an overvoltage (the famous 0.8eV).

I don't know how to summarize this in the article. --Enric Naval (talk) 19:39, 27 September 2013 (UTC)[reply]

Currently, the article is so biased against the possibility that CF might be real that addressing a single issue (viz. whether Huizenga's criticism would in fact stand up if one looked properly into the details) seems hardly worth the trouble. This article should really be re-written from the start, with a view to including the totality of the evidence on both sides. --Brian Josephson (talk) 09:10, 28 September 2013 (UTC)[reply]

"Currently, the article is so biased against the possibility that CF might be real" because the reliable sources overwhelmingly say it is not real. As has been mentioned, we summarise what the reliable sources say. We are not aiming to balance minority views with the mainstream. Fringe theories are put into perspective with the mainstream. I won't link the policies and guidelines as you have already admitted you have no intention of familiarising yourself with them, IRWolfie- (talk) 10:37, 28 September 2013 (UTC)[reply]

It would be interesting to know just what article of the Wikicreed legitimises the extent of bias that is apparent in this article. --Brian Josephson (talk) 14:06, 28 September 2013 (UTC)[reply]

The relevant section on minority views is at Wikipedia:Undue_weight#Due_and_undue_weight. - MrOllie (talk) 15:01, 28 September 2013 (UTC)[reply]

Thanks for the reference. Yes indeed; it is very easy of course to interpret wikicreed in a biased way (metabias one might call this), some people can do that with their eyes closed! --Brian Josephson (talk) 16:37, 28 September 2013 (UTC)[reply]

Original derivation and history

From the beginning the equation was based on a implausible concept used by the author that beeing the dissolution pressure of metals which had astronomically high values--5.15.176.81 (talk) 15:50, 29 September 2013 (UTC)[reply]

Limitations

I introduced the term "unpoised electrode" which I found in internet, but I'm not a chemist. A professional look would be welcome (if this is a commonly used term). Perhaps some link to kinetic equation would be nice in this point ("other effects" mentioned in the article). Also probably so called "anodic index" is related to this section - see galvanic corrosion article. — Preceding unsigned comment added by 159.205.140.205 (talk) 20:34, 3 November 2013 (UTC)[reply]

Formal Electrode Potential

Formal potential, E°', as applied in the Nerst equation, is the standard potential, set for the working electrode (WE) for a particular reversible reaction occurring at that WE. The formal potential is found half way between the two peaks in its cyclic voltammogram, where at this point the concentration of O (the oxidized substrate) and R(reduced substrate) at the electrode surface are equal. The difference between the formal potential E°’ and normal E°, is simply because in this case concentrations are being used, not activities. It is still the stardard set for the WE for a particular reaction. [O]σ= [R]σ σ = at surface

Formal Electrode Potential

Formal potential, E°', as applied in the Nerst equation, is the standard potential, set for the working electrode (WE) for a particular reversible reaction occurring at that WE. The formal potential is found half way between the two peaks in its cyclic voltammogram, where at this point the concentration of O (the oxidized substrate) and R(reduced substrate) at the electrode surface are equal. The difference between the formal potential E°’ and normal E°, is simply because in this case concentrations are being used, not activities. It is still the stardard set for the WE for a particular reaction. [O]σ= [R]σ σ = at surface

^ Richard G. Compton; Craig E. Banks (2011). Understanding Voltammetry. World Scientific. p. 57. ISBN 978-1-84816-586-1. Retrieved 27 September 2013.

[ComptonBanks2011-1] Richard G. Compton; Craig E. Banks (2011). Understanding Voltammetry. World Scientific. p. 57. ISBN 978-1-84816-586-1. Retrieved 27 September 2013.

[1]

@@ Line 155: / Line 155: @@
 == Limitations ==
 I introduced the term "unpoised electrode" which I found in internet, but I'm not a chemist. A professional look would be welcome (if this is a commonly used term). Perhaps some link to kinetic equation would be nice in this point ("other effects" mentioned in the article). Also probably so called "anodic index" is related to this section - see [[galvanic corrosion]] article.  <span style="font-size: smaller;" class="autosigned">— Preceding [[Wikipedia:Signatures|unsigned]] comment added by [[Special:Contributions/159.205.140.205|159.205.140.205]] ([[User talk:159.205.140.205|talk]]) 20:34, 3 November 2013 (UTC)</span><!-- Template:Unsigned IP --> <!--Autosigned by SineBot-->
+== Formal Electrode Potential ==
+Formal potential, E°', as applied in the Nerst equation, is the standard potential, set for the working electrode (WE) for a particular reversible reaction occurring at that WE. The formal potential is found half way between the two peaks in its cyclic voltammogram, where at this point the concentration of O (the oxidized substrate) and R(reduced substrate) at the electrode surface are equal. The difference between the formal potential E°’ and normal E°, is simply because in this case concentrations are being used, not activities. It is still the stardard set for the WE for a particular reaction.
+[O]σ= [R]σ σ = at surface
 == Formal Electrode Potential ==