Standard half potentials are always with reference to Standard Hydrogen half potentials (SHE) unless otherwise specified.  They can either be expressed in the reduction or oxidation half reactions at a given temperature.  Since these are all with reference to the standard Hydrogen half potential therefore the signs of + or - is always with reference with the SHE which is designated as a zero.  For any given standard reduction half-reaction, an E= + means that reduction will proceed as written, in this case reduction when paired with SHE.  An E= -, means that oxidation will be the normal pathway when paired with SHE. 

In the example given:  Zn2+ + 2e ---> Zn, E = -0.76 V vs. SHE.  The E is negative meaning the half reaction will not proceed if paired off with SHE.  Since SHE will have a greater tendency to be reduced. 

The interpretation of the negative (-) and (+) potentials as spontaneous and non-spontaneous goes beyond being "electrostatic"  and can be related to free energy. 

The signs are related to the change in Gibbs free energy. If a process has a negative (delta)G it must have a positive (delta)E and the reaction is spontaneus. Look at http://chemwiki.ucdavis.edu/Analytical_Chemistry/Electrochemistry/Electrochemistry_and_Thermodynamics

1.In any half cell reaction the "E" represents the redox potential of the couple. If the E is -ve, then the reaction is said to be non-spontaneous and if the E is +ve, then the reaction is said to be spontaneous (according to the equation -deltaG=nFE).

1) a positive or negative charge is related to whether a certain direction is spontaneous or not (favored or not) according to the Gibbs energy relationship

2) I dont think it is true, but need to check that

Thanks for your replies. Few years ago, I had correspondence/ discussion with three leading electrochemists. They were of the view that the electrode potential signs, as written in tables, represent the electrostatic sign of the electrode. For instance, if a single zinc rod were dipped in a zinc sulfate solution, it would acquire a negative charge. The sign of this charge could be detected using the very original experiments of Volta by his "condensing electroscopes". This made perfect sense. However, only one standard text by Bard is explicit about the invariant electrostatic sign of the electrode potential. Another leading electrochemist called it a common misconception perpetuated by textbooks by writing:

Zn --> Zn2+ + 2e E = +0.76 V,

from

Zn2+ + 2e ---> Zn, E = -0.76 V vs. SHE

which is as incorrect as writing

If H2O (l) --> H2O (g)   at +100 oC

H2O (g) --> H2O (l)   at -100 oC

Almost all texts examined so far sweep it under conventions. For interested readers, these so called conventions were called American and European conventions of the electrode potential signs representing G. N. Lewis and Ostwald's school of thought. Lets see what others have to say. Thanks.

Even though there is connection between the electrode potential and the charge of the electrode, in general one cannot predict the sign of an electrode just by looking its potential. More negative standard electrode potential show a tendency for the electrode to charge negatively and more positive standard electrode potential show a tendency for the electrode to charge positively. What the actual charge results, that depends on the conditions of the whole cell.

1) The negative sign for the electrode in the half cell 

Zn2+ + 2e ---> Zn, E = -0.76 V vs. SHE

means that a Zn electrode in contact with a Zn2+ would be negatively charged when forming an electrochemical cell if the second electrode were SHE, which has a 0 V electrode potential (i.e. in a cell formed by those electrodes, the Zn one would be negatively charged and the SHE would be positive charged). In general, the second electrode can be any electrode, not just SHE, and may have different potentials, so that the Zn can be positively or negatively charged (for instance, Zn would be positively charged if the potential of the second electrode is -1 V). Besides, SHE as having 0 V is just a convention, other reference electrodes can be used and the charge of the electrode is independent of electrode potential conventions.

2) If one has only one electrode (Zn dipped in zinc ions solution), inverting the order of the reactants and products cannot change the sign of the electrode. Under equilibrium conditions, it is exactly the same how the equation is written and the water analogy is a good one.

Many misunderstandings also come from confusing the meaning of the standard electrode potential (based on conventions such as 0 V for SHE and the sign convention based on the reduction or oxidation potential) and the potential difference in a cell which is what can be actually measured and related to the free energy of the reaction.

Thanks all. With the discussion so far, it is clear that some people take the electrode potential signs in purely thermodynamic sense and some others take it as an invariant sign = electrostatic sign of the single electrode wrt to SHE, which is assigned to 0 V for the sake of convenience. 

I am coming to another point, I would appreciate your comments. It is very well established in electrostatics that the potential bears the same sign as the charge. Now the absolute potential of the hydrogen electrode is very well known. It is around +4.44 V. Now by converting the electrode potential zinc half cell  vs SHE in absolute electrode potential term is:

E(Zn2+/Zn, absolute) = +4.44 + (-0.76) = +3.68 V.

There is an apparent dilemma, why is the absolute potential of zinc electrode positive, despite the fact that it is negatively charged?

Thanks.

Two different conventions are being used and thus the misunderstanding. The "absolute electrode potential" for Zn is related to the process

Zn → Zn++(solution) + 2 e−(gas)

Therefore, a positive value for the absolute potential of a zinc electrode shows that work must be done to take an electron away from the electrode, but it still says nothing on the charge of the electrode.  

Shouldn't the equation be the other way round. I believe that Wikipedia has it wrong on the section of absolute Hydrogen electrode.http://en.wikipedia.org/wiki/Absolute_electrode_potential

See for instance, http://pubs.acs.org/doi/pdf/10.1021/jp100402x. They say that the absolute hydrogen electrode potential (as written is) +4.44 V

H+ (aq) + e (gas) ---> 1/2 H2 (g)

rather than for:

1/2 H2 (g) --> H+ (aq) + e (gas)  as +4.44 V.

To avoid misunderstandings in this respect, it may be easier to consider the electronic energy (U) and then relate it to the electrode potential (E): 

U = -e E(abs) = -4.4 eV - e E

where E is the electrode potential vs SHE and E(abs) is the absolute electrode potential. U is the energy associated to the work done to take electrons from the electrode (at the Fermi level) to vacuum. Thus, energy values can be used for the chemical equations written in any form, instead of operating directly with the somewhat misleading electrode potentials. (See Fig. 2 in http://pac.iupac.org/publications/pac/pdf/1986/pdf/5807x0955.pdf )

Good questions. I also believe that there is confusion about the sign convention.

1) A negative electrode potential (in the scale based on SHE) of the Zn/Zn2+ electrode indicates that the electric potential of it is less than that of the SHE. Electric current flows from a point A to point B, when connected by a conductor,  if the electric potential at A is higher than that at B, It is NOT necessary that point A has a positive charge and point B has a negative charge, Hence, if you prepare a cell with a Zn/Zn2+ electrode and a SHE, then electric current flows from SHE to the Zn/Zn2+ electrode when they are connected with a conductor. (Electrons flow in the opposite direction).

2) It is not necessary that the zinc electrode is negatively charged. By looking at the sign of the electrode potential you cannot conclude that the zinc electrode is negatively charged. One has to do an experiment and find out the actual sign of  charge.

Interested reader can refer the attached file for bit more details on the subject of electrode potential which probably sheds more light on the issue.

Dear Dr. Bandarage, Thanks for adding good points

1. You raised an important point by saying that if point A is at higher potential than point B, it doesn't necessarily mean that they are oppositely charged. Both could be negative or positive but at different potentials.

2. I may not fully agree with your second statement. By looking at the tables of electrode potentials, one can judge the sign and charge of the electrode wrt SHE. A Zn2+/Zn would be electrostatically charged with a negative sign if it were dipped in its own solution such as zinc sulfate. Thus a single electrode dipped in zinc sulfate solution would be negatively charged just like a amber rubbed with fur. This was the original definition of charges by Benjamin Franklin. Please see Nernst theory of osmotic pressure. He called this phenomenon as "Elektrolytischen Losungdruck" =(electrolytic solution pressure).

You also raised a good point: "One has to do an experiment and find out the actual sign of charge.". Lets say we have a new cell design. What would be good experimental way (ideally electrostatic experiment) to determine the electrostatic sign of the electrode? Please exclude voltammeters for the time being.

Dear Dr. M. Farooq Wahab,

Thanks for digging deeper into the issue. I am a little puzzled by the claim that "By looking at the tables of electrode potentials, one can judge the sign and charge of the electrode wrt SHE". Probably, the point is whether one can 'talk" of a charge on an object with respect to a charge on another object (such a thing can be done with electric potential). Two types of charges may be defined using the charges on electrons and protons. If an object has excess of "electron type" charges, then we could say that the object has a negative charge. If an object has an excess of "proton type" charges, then it has a positive charge.

One may definitely say that an electrode with a negative electrode potential (in SHE scale) has a  negative charge only if the charge on SHE is zero (neutral). 

Electrode potentials are defined using SHE as the reference. (The choice is arbitrary as far as charges are considered.) Similarly, one may define another "ELECTRODE POTENTIAL" using some other electrode as the reference.; for example standard  Zn/Zn2+. Then on that scale the "standard electrode potential" of Zn/Zn2+ will be zero. What does that mean in terms of charge on Zn/Zn2+?

Another very interesting point you raised is that a zinc rod dipped in a zinc sulphate solution is negatively charged. I fully agree that it is a possibility which would not contradict the observed value of standard electrode potential of Zn/Zn2+. But how do you argue that it is so? What would you get if you apply the same arguments to a Cu/Cu2+ electrode which has a positive electrode potential (in SHE scale)?

Dear Dr. Bandarage,

Thanks for your input. You raised interesting points. Lets look at the problem from a historical perspective. The fundamental practical definition of charges is still the same given by Benjamin Franklin in the 16th century. A glass rod rubbed with silk acquires a positive charge and amber rubbed with fur acquires a negative charge. The electron has been given a negative sign because it behaves like amber rubbed with fur on a electroscope not vice versa. Lets stick to Franklin's definition for positive and negative charges for the time being.

A) The charge on SHE is not neutral and it never is. You are totally right, that if we had Zn2+/Zn as a reference, we could easily assign it as 0V and quote other potentials with respect to Zn2+/Zn cell. Now, this does not imply that Zn2+/Zn is electrostatically neutral. The actual point is that electrode potentials as they appear in modern tables imply a electrostatic sign wrt SHE. A leading electrochemist said that this happened by chance that modern tables are written as reduction potentials as recommended by IUPAC. By chance they also reflect the electrostatic sign of the electrode.

On the other hand, you will notice that some people take these signs in thermodynamic sense as well (in terms of delta G). This was the school of thought of Lewis and Randall (people who shaped modern thermodynamics in the US) following Gibbs.

You wrote:

"Another very interesting point you raised is that a zinc rod dipped in a zinc sulphate solution is negatively charged. I fully agree that it is a possibility which would not contradict the observed value of standard electrode potential of Zn/Zn2+. But how do you argue that it is so? What would you get if you apply the same arguments to a Cu/Cu2+ electrode which has a positive electrode potential (in SHE scale)?"

B) Your last paragraph is the gist of my discussion and I have been trying to search for an (historical) answer for the last 6 years!  Note that ordinary galvanic cell will not respond to electroscopes-the tools used for determining signs of the charges. In short, it was Volta's condensing electroscope which helped him determine the sign of cells he made.

Most of the early interesting work in electrochemistry is in German. Unfortunately this is limiting my quest for reading the very original work of Nernst and Ostwald who  pioneered and formulated these questions. I am attaching two pages from a leading electrochemistry text of its time (in German, 1922). See the images please (Fig 50 and Figure 51). The author seems to address your last question in terms of Nernst theory of electrolytic solution pressure. Hope I can find some translation/ translator but the images are still worthwhile. The footnote on pg 152 is also worthwhile. The translation of which is available in the article

http://pubs.acs.org/doi/abs/10.1021/j150185a005?journalCode=jpchax.2

Please read this article for further discussion. I can email a copy, if you don't have access to this one.

I haven't heard of any way to know the actual charge of an electrode without the use of modern electrochemical instrumentation. However, I think that a Coulomb's torsion ballance might work. The experimental set-up would be a little strange, but it could be possible to determine the charge of a single electrode.

As far as I could trace by contacting science historians, they all pointed to Volta's condensing electroscope. This electroscope in principle, can allow to determine the electrostatic sign of a single electrode or both electrodes in voltaic cells. I am sure this was such a routine experiment that Faraday in his book on Electricity & Magnetism writes that it can be shown using a gold leaf electroscope to show that zinc in a voltaic pile is negatively charged. He doesn't bother to write further details since I assume that it must be such a common experiment- well known to 18th century scientists.

As a modern electrochemist, how would you experimentally determine the electrostatic sign of a voltaic cell (assuming no info is available in tables) of a single electrode using modern instruments?

As a side note, Bard published some papers on electrostatic electrochemistry on insulators in Nature. He reduced copper using an electrostatically charged insulator rod- an similar experiment was done by Faraday hundreds of years ago. Sometimes old ideas are very useful.

The potential of zero charge can be determined experimentally for a given electrode and then, by comparison with the potential measured, the sign of the charge can be known. However, I have never thought of a general and simple way to know the charge of any given electrode. I think it is not easy to carry out. It is actually a simple question with a complex answer. Perhaps a way to find out the sign of the charge of any electrode could involve the comparison with a mercury electrode under certain conditions (an experiment related to those by Lippmann used to measure the potential of zero charge of mercury), but I am not certain that this would work in general. By the way, very interesting stories on Faraday's experiments.

Thanks Dr. Ybarra for bringing in interesting ideas. Your suggested experiment was done by Ostwald in 1901. The paper is again in German and the journal is not even online. It is  Zeit. Physikal. Chem., 1901, 36, 91. I have access to the English abstract published elsewhere. He apparently did what you just suggested above. He called those potentials as his absolute electrode potentials. I guess these type of experiments (null point of metals) allowed 19th century physical chemists to assign signs to their measured electrode potentials. As far as I know, neither Nernst or Ostwald employed Volta type condensing electroscopes to determine the correct electrostatic sign of the electrodes.

PS. Faraday use metallic electrode connected to electrostatic machines to do electrolysis. This was an example of single electrode electrolysis of KI...which a modern student can never imagine because the books do not train them to think outside the box. Electrochemistry is usually buried under sign conventions. Bard made it to Nature magazine by employing charged insulators for electrolysis for and called it "Electrostatic Electrochemistry"

Thank you for bringing thought-provoking questions.

Dear Dr.  Wahab

Thanks for giving that very nice old article in J. Chem. Phys. I enjoyed reading it. I am making some comments based on my understanding of it.

It indicates the confusion on what sign to be used when reporting a “potential” (more correctly Inter-facial Potential Difference) associated with an electrode, at that time. I guess, armed with a clear (operational) definition of Electrode Potential we have now come out of it.

The paper reports some values of inter-facial potential differences of electrodes and suggests using the same sign as the charge on the metal. Experimental details are not given. Hence, it is not clear whether the reporters of these values actually determined the charge on the electrode. If one uses another electrode in measuring the said inter-facial potential difference, then the ambiguity we discussed earlier (introduced by the charge on the second electrode) gets in. However, it may be possible to determine the charge on an electrode by observing the direction of the electric field at the interface. It is not clear whether they did something similar to that.

The statement, “if we dip a piece of copper into a copper sulphate solution, we usually consider that some copper ions are precipitated on the copper charging the copper positively, and leaving the solution charged negatively” indicates copper getting a positive charge in a copper sulphate solution is an assumption for the author.

Dear Prof. Bandarage,

You wrote " The statement, “if we dip a piece of copper into a copper sulphate solution, we usually consider that some copper ions are precipitated on the copper charging the copper positively, and leaving the solution charged negatively” indicates copper getting a positive charge in a copper sulphate solution is an assumption for the author.

This is a correct assumption by the author, about the actual charge of the electrode. We may not agree with the mechanism. If you see the reference no. 1 by Ostwald in the paper I sent you, you will see that Ostwald has shown that a zinc rod dipped in _any_ acid he tested, it was found to be negatively charged. Similarly, a copper rod was found to be positively charged in all cases. Ostwald made use of the null point measurements using Lippman capillary electrometers. Eventually after 50 years, Stockholm IUPAC convention went back to the Ostwald school of thought of assigning the same sign as what Ostwald did centuries ago.

Dear Dr Farooq,

I would really like to see the original work by Ostwald, so if you manage to get it, I would appreciate if you can send me a copy. I think that the sort of experiment I described, using the Lippmann capillary electrometer, probably just compares the charge on an electrode with that of a mercury one, and I don't think that would be a general method valid for any given electrode. I think that it could be equivalent to measure the electrode potential against any reference electrode (e.g. saturated calomel electrode) and then compare that value with the well-established value for the potential of zero charge for a mercury electrode. I am assuming a lot of things and the original work would help.

This also brings an intriguing question. It is known that the potential of zero charge is different for every metal. So if one has two electrodes of different metals at the potential of zero charge and then one measures the potential difference between them, this potential difference would be not null. But how is that possible if both have zero charge?

Dear Dr. Ybarra ,

Thanks for that interesting point. Does this difference has something to do with the work function of the metal?

Thank you, Prof. Bandarage. It is related to the work function or the Fermi level. The point I find more interesting with the questions posed by Dr Farooq is that I can understand the problem in terms of the Fermi level or the electrochemical potential, but there is obviously some confusion when one tries to combine that with electrostatics.

Standard half potentials are always with reference to Standard Hydrogen half potentials (SHE) unless otherwise specified.  They can either be expressed in the reduction or oxidation half reactions at a given temperature.  Since these are all with reference to the standard Hydrogen half potential therefore the signs of + or - is always with reference with the SHE which is designated as a zero.  For any given standard reduction half-reaction, an E= + means that reduction will proceed as written, in this case reduction when paired with SHE.  An E= -, means that oxidation will be the normal pathway when paired with SHE. 

In the example given:  Zn2+ + 2e ---> Zn, E = -0.76 V vs. SHE.  The E is negative meaning the half reaction will not proceed if paired off with SHE.  Since SHE will have a greater tendency to be reduced. 

The interpretation of the negative (-) and (+) potentials as spontaneous and non-spontaneous goes beyond being "electrostatic"  and can be related to free energy. 

Dear Dr. Ybbara, With some digging I managed to find the original two Ostwald's papers, but before that I would suggest that you read the following short article for the sake of context of Ostwald's first paper. The articles have been seen via personal message.

http://pubs.acs.org/doi/abs/10.1021/j150185a005?journalCode=jpchax.2

I cannot understand most of it because there is no translation available, but the tables do give some clue what he was doing. I noticed that he used "Kapillar elektrometer" and "Tropf elektrode" and this is certainly Lippman's electrometers and his dropping Hg electrode. It was a hot topic of its time leading to couple of Nobel prizes.

(a) You have raised very interesting ideas related to the hypothesis that the charge of the test electrode must be with respect to mercury; this is 100% possible but we need to read what he was saying. I feel that the sign of the charge must be dependent only on the metal/ solution interface. Lets think of a thought experiment: A single rod of zinc is dipped zinc sulfate solution without contacting any other metal. What is the electrical state of charge of the zinc rod? Is it negatively charged or is it positively charged?

BTW, have you seen "charge sniffers?  http://scitation.aip.org/content/aapt/journal/ajp/79/2/10.1119/1.3531961

We need something of that sort for this simple thought experiment.

(b) The question "So if one has two electrodes of different metals at the potential of zero charge and then one measures the potential difference between them, this potential difference would be not null. But how is that possible if both have zero charge" is very interesting and thought provoking.  I feel it is the same old question of Volta i.e. contact electricity that is two dissimilar metals will produce a potential difference. The very thought of them being at PZC is very intriguing. What is your opinion? 

Thanks for your input and giving new directions.

Regards.

Dear Dr. Ybarra,

The table of values you have given on the potentials of  metals with zero charge is interesting. It has a bearing on the thread of the discussion here. I guess one could understand that it should not be zero (in general) if you look at the definition of electric potential. When the positive charge you bring from infinity comes closer, the free electrons on the metal moves closer to the positive charge leading to an attraction. Hence, the metal will do work on the positive charge. This will lead to a negative potential on the metal. But I am puzzled by the positive values of metals like copper, Au, Pt etc.

Dear Dr. Wahab,

Coming back to the fundamental question of charge on Zn in the Standard Zn/Zn2+ Electrode, you may be able to determine it (without performing experiments) in principle, if you can find Gibbs free energy change for the reaction Zn(s) = Zn2+(aq) + 2e . (If it is negative, then the charge on Zn has to be negative since the reaction is spontaneous.) We need the chemical potentials of Zn(s) and Zn2+(aq). One may imagine finding them in a book of constants. But we also need the chemical potential of electrons on neutral Zn. Where can we find that?

Electrochemists have worked for quite a long time on the electrical double layer on electrodes. They me be knowing something about these issues.

Most of the questions we've been dealing here (the measurement of charge at an electrode, the early experiments by Lippmann, the potential of zero charge and its connection with the work function, Fermi levels and the thermodynamics of electrode processes) are fundamental but not easy at all. I found many answers in the first volume of the "Comprehensive Treatise of Electrochemistry", edited by Bockris, Conway and Yeager. However, it is difficult to make a rigorous yet simple account of the whole subject. If I ever get to find an elegantly simple treatment, I will post it here. 

Dr. Ybarra,

Thanks. I also felt that this seemingly simple problem is not easy to solve. During free time I am looking at ways to determine the electric field at the surface of a metal. Surface enhanced Raman spectroscopy (SER) is being used to get information on the charge on metallic nano-particles.

Dear Prof. Ybarra and Prof. Bandarage,

You are right that metal electrolyte interface is theoretically very challenging. However, it seems that older physical chemists were pretty comfortable by connecting electrostatics and electrochemistry. Can we discuss, step by step simpler questions? Readers can take one point at a time. 

1. Lets start with zinc and silver. The work function of zinc is lower than silver. If zinc and silver are connected (contact electrification) would zinc become positively charged and silver negative? This is one of the earliest experimental observations of Volta, when work functions and electrons did not exist. I am sure this was not an erroneous result since I cannot find a correction. 

https://www2.chemistry.msu.edu/faculty/harrison/cem483/work_functions.pdf

2. Now if we imagine, a zinc rod dipping in Zn2+ solution in one beaker, and a silver rod dipping in Ag+ solution, if we make a closed circuit by using salt bridge and external connection, Zinc terminal is negatively charged and Silver terminal is positively charged. Why is this so?

3. It was very well known, in a Lippman capillary, that mercury drop is positively charged wrt to the acid? It seems that Ostwald and his contemporaries were very comfortable with this idea. I am unable to find an experimental set up to determine the just the sign of the electric charge of a single metal dipping in a solution.

Thanks all for your input and time.

Dear Prof. Wahab,

Your problem is very intriguing. I must thank you for posing it. As you have pointed out the subject is very old and there were no work functions and electrons then. Yet the scientists make remarkable progress. It is good to have a clear idea of what we mean by "positively charged" and "negatively charged". As discussed earlier, in question 2 the claims about the "charged nature" of metals is uncertain. The claim appears to be based on the direction of flow of electric current. Electric current can flow form a "positively charged" object to a "positively charged" object OR from a "negatively charged" object to a 'negatively charged" object OR from a "positively charged" object to a "negatively charged" object. The way I see it, the condition is that the electric potential should be higher on the object where the electric current originates.

I'm not sure that these problems can be properly understood if we try to think the way early physical chemist thought, from a merely electrostatic point of view. I have tried to summarize an approach to these questions in the framework of electrochemical potentials, which can be used to answer the first two questions. Electrochemical potential can be finally related to the electrode potential, and that's a quantity with which electrochemists feel at easy.  I am uploading my notes with the hope it might be of some help; if they are found useful, I will complete them later.

The Lippmann experiment can also be better understood using electrochemical potential, including the experimental determination of the sign of the Hg electrode, though I haven't included it for the moment.

Thanks Prof. Ybarra. It is indeed a useful theoretical treatment. Please continue. With that, we should also be able to find an experimental way for case II in your notes. Your case II proves that a potential difference develops when the metal is dipped in an electrolyte. The ideas of double layers were known a century ago. What experimental and observations did these people have which made them think about the double layers at the metal-electrolyte interface and consequently how did they determine the sign of the electrode wrt to the electrolyte? I will also try to find more about it.

Standard electrode reduction potential for different metals are tabulated for:

Now the negative sign of standard reduction potential for any metal indicates that an electrode when joined with standard hydrogen electrode (SHE) will not act as cathode and reduction reaction half reaction will not be spontaneous. Instead the said metal will act as anode and oxidation (loss of electrons) will occur spontaneously on this electrode.

For example, standard reduction potential of zinc is -0.76 volts. So, when zinc electrode is joined with SHE, it will not act as cathode as reduction half reaction will not occur spontaneously, instead Zinc electrode will acts as anode (-ve electrode) i.e., oxidation occurs on this electrode spontaneously. The +ve sign of standard reduction potential indicates that the electrode when joined with SHE acts as cathode and reduction occurs on this electrode spontaneously.

Positive sign of reduction potential physically signifies non-spontaneous reduction reaction and negative sign indicates spontaneous reduction reaction. From electrochemical point of view this is justified as below:

ΔG=−nFE∘ Gibbs free energy of an electrochemical cell.

E is electrode potential

ΔG is Gibbs free energy, The difference between the enthalpy of a system and the product of its entropy and absolute temperature; a measure of the useful work obtainable from a thermodynamic system at constant temperature and pressure.

F = Faraday Constant

n is number of electrons transferred in the balanced equation

Negative values of E∘ (Reduction potential) will lead to positive values of ΔG and vice versa. ΔG is positive, the reaction is non-spontaneous, and when ΔG is negative, the reaction is spontaneous. So, we can summarize as: