Talk:Standard electrode potential

Request title change

Requested title change to "Standard Reduction Potential" from "Reduction Potential" to link shortform the commonly used SRP to this page

—The preceding unsigned comment was added by Ventulus (talk • contribs) 21:28, 12 April 2007 (UTC).[reply]

"Reduction potential" is no longer the preferred term: IUPAC recommends either "electrode potential" or "redox poten≝tial". I have linked the line on SRP through to this page. Physchim62 (talk) 03:47, 13 April 2007 (UTC)[reply]

Merge/expansion

These two articles should be left separate. Reduction potential on its own is much easier to apply to a biological system as I have discovered. I don't think I have to say much since someone has made all the necessary comments below. Rylkiwanuka 18:08 16 September 2007 —Preceding signed but undated comment was added at 17:10, 16 September 2007 (UTC)[reply]

I am currently working on a new version of this article at Standard electrode potential/Temp. Physchim62 (talk) 06:14, 28 November 2005 (UTC)[reply]

I agree that the two should be merged. There is little that one would want to say on either subject without a very large degree of overlap Ahw001 09:28, 26 March 2006 (UTC)[reply]

I think it's better to mantain the articles separated because many people are in confusion with the different ways to relate to the same concept. 26 august 2006

I also think they should be separated, as ORP is used as a field parameter in water quality assessment and ground water analysis. For those uses, a general definition or suggestion of implication of Eh is more important than explaining how the standard was achieved. Also, field conditions never have the standard constants, like temperature, which might throw off users of the ORP recorded in fie24.172.82.130 11:46, 11 January 2007 (UTC)ld equipment like water quality meters, etc. 24.172.82.130 11:46, 11 January 2007 (UTC)[reply]

I suggest keeping separated. I've added information regarding practical measurements of ORP, which seems like good focus for a general article on ORP, that would then link to more specific discussion of standard electrode potential. Jjotter 22:18, 2 February 2007 (UTC) Actually reduction potential referes to the system. e.g. it can be applied to the biological systems as well. But when we say standard electrode potential, then it become difficult to apply the concept of redox potential through standard electrode potential to general systems.[reply]

Actually reduction potential referes to the system. e.g. it can be applied to the biological systems as well. But when we say standard electrode potential, then it become difficult to apply the concept of redox potential through standard electrode potential to general systems.so i think these two sections should be kept separate.

A question

Can someone please explain to me why Li has a lower value than K (for example), but K is more electropositive? Thanks... -postglock 13:01, 30 December 2005 (UTC)[reply]

Absolute potential of each half-cell?

I'm changing this odd statement: there is no way to measure the individual potentials of the electrodes in isolation. There are ways to measure electrostatic potentials, but they're not simple and not very accurate. --Wjbeaty 03:07, 15 November 2006 (UTC)[reply]

With reference to what? earth? but then, whose earth? The absolute potential of a half-cell is undefined, and this has some significant implications for thermodynamics. I will invite other comments before reversing your edit, as you obviously believe what you're saying, even if I think you're wrong! :) Physchim62 (talk) 15:44, 16 November 2006 (UTC)[reply]

Earth? Huh? You do understand how the halfcell potentials are created, right? The electrolyte is a conductor, as is the metal, while the Helmholtz layer at the metal/electrolyte interface acts both as a charge pump and as a capacitor dielectric. This capacitor is charged up to a particular potential diff. This absolute halfcell potential appears between the capacitor's two plates: between the electrode and the electrolyte. Since we can't connect a common voltmeter to the electrolyte without introducing huge artifacts, a non-contact voltage measurement would be required. And these measurements aren't very accurate. But just because imprecise and exotic voltmeters are needed, that doesn't mean that the electrolyte/metal capacitor has an undefined voltage.

For example, a quick online search turns up various papers which put the absolute halfcell potential of the SHE at around 4.45v. Since chemists of the 1800s had no way to measure this value, yet had a great need to organize the list of halfcell potentials, they apparently decided to declare the SHE halfcell potential to be zero, even though it's nowhere near zero.

So, which thermodynamic implications arise if we accurately measure the value of the SHE halfcell potential? --Wjbeaty 08:52, 8 December 2006 (UTC)[reply]

Provided half-cell potentials are not being used in isolation but combined to give overall cell potentials, there are no thermodynamic implications - the reference point is arbitrary 84.92.241.186 15:52, 29 April 2007 (UTC)[reply]

E naught

Shouldn't it be an 'E' with a superscript 'o' horizontally crossed through it?

Now, you just have a superscript letter 'o'. —The preceding unsigned comment was added by 60.241.120.20 (talk) 22:54, 9 December 2006 (UTC).[reply]

Well? Does a symbol even exist for E naught? 60.241.129.19 02:38, 12 January 2007 (UTC)[reply]

Either is correct! See Quantities, Units and Symbols in Physical Chemistry, pp. 49 (note 10), 53. Physchim62 (talk) 15:20, 11 May 2007 (UTC)[reply]

Contradiction in definition of Standard Electrode potential ??

First the text says ...

Since the electrode potentials are conventionally defined as reduction potentials, the sign of the potential for the metal electrode being oxidized must be reversed when calculating the overall cell potential. Note that the electrode potentials are independent of the number of electrons transferred -that is, they are set to one mole of electrons transfered- and so the two electrode potentials can be simply combined to give the overall cell potential even if different numbers of electrons are involved in the two electrode reactions

Then we have the following text ..

Since the table of standard electrode potentials is defined for a transfer of one mole of electrons, care must be made in determining an electrode potential using two other electrode potentials. Adjustments have to be made for the number of electrons being transfered.
For example:
(eq1) Fe3+ + 3e- --> Fe(s) is listed as -0.036V
(eq2) Fe2+ + 2E- --> Fe(s) is listed as -0.44V
to get a third equation:
(eq3) Fe3+ + e- --> Fe2+ (listed as +0.77V)
one would need to take eq1 and multiply the voltage by 3, reverse eq2 (changes the sign) and multiply the voltage by 2. Adding those two voltages together gives the standard potential for eq3.

Surely that is saying that the calculated electrode potential of the cell does depend on moles of electrons transferred - which contradicts the first bit I quoted.

Ideas anyone?

Clive long (talk) 11:51, 13 May 2008 (UTC) Clive long[reply]

I think that everything after the "that is" in the first paragraph you quote is incorrect. The cell potential is related to the Gibbs free energy change of the cell reaction, and so is dependant on the equation of the cell reaction. Standard electrode potentials are independant of the nulber of electrons transferred, but you must take into account the relative numbers of electrons for each half-reaction when calculating the cell potential. I shall wait a few days to see if anyone contradicts me, then change the article! Physchim62 (talk) 16:36, 15 May 2008 (UTC)[reply]

I think both statements are more or less correct. The problem is that the "auto-balancing magic" only works when you are doing a full cell (that is, no electrons hanging around in the equation!). For example, 2 Fe³⁺ + Fe -> 3 Fe²⁺ has a cell potential of 0.77 - -0.44 = 1.21 V. Note that the 0.77 is not multiplied by 2. But now, if you want to combine two half-reactions to get another half-reaction, as in the case given by Clive, where the electrons don't cancel out, you need to do the accounting explicitly. Remember that what is additive is the ΔG; if you convert the E's to ΔG's, add them up, and then convert back, you should get the right answer (ΔG = -nFE, where n is the number of electrons). You can also do the explicit accounting for a full cell reaction and get the correct result, which you'll see includes multiplying and dividing by n (which does nothing to the E).

Put another way, the equation Fe3+ + 3e- --> Fe(s) really means 1/3 Fe3+ + 1e- --> 1/3 Fe(s). You cannot subtract from this equation 1/2 Fe2+ + 1e- --> 1/2 Fe(s) and get Fe3+ + e- --> Fe2+ without including stoichiometric coefficients. What you get instead is the reaction 1/3 Fe3+ + 1/6 Fe -> 1/2 Fe2+, which is the "auto-balanced equation" for the full cell. --Itub (talk) 16:54, 16 May 2008 (UTC)[reply]

Effect of neutron beams on the potential

It seems that some results have been reported by Bagotski on the influence of neutron beams on the electrode potential.--188.27.144.144 (talk) 13:13, 30 October 2013 (UTC)[reply]

Link to article by Bagotski from Electrochimica Acta : Electrochimica Acta.--188.27.144.144 (talk) 11:09, 20 February 2014 (UTC)[reply]

Further Reading... for real?