The Correlation between the Electronegativities and the Standard Potentials
Authors: Chernomorskii Alexandr I. 
Published: July 2021

Electronegativity [1.2] can be defined as the ability of an atom in molecule to draw covalent bonding electrons to itself from the intermediate state in which electrons are equally bound with each atom. The equal bonds of electrons with atoms would be compensated in the intermediate state. However, there should be saved the definite simple non-specific electrostatic (coulomb) attraction of bonding electrons with each polarized atom in this intermediate state [3-5]. Such a non-specific intermediate bond of electrons is revealed between thermodynamic non-stable specific bonds of electrons in reductant and thermodynamic stable specific bonds of electrons in oxidants. Due to this reductants are able spontaneously give up electrons in chemical interactions. Oxidants can only accept electrons in chemical interactions if their own electrons are bound thermodynamically stable [3]. The intermediate electron bond is revealed in the correlations [6-10] between the standard potentials
E

o and electronegativities χa
of metal and non-metal electrodes (Fig.1.2) and between the standard potential
E

o and the electron work functions
W

e of metals (Fig.3) [1,2].The charges of potential-dependable ions are shown in Fig.1-3. In the coordinates χa -

E

o, with the increase of the empirical electronegativies χa
the reducing activity of typical metals (in the line Na --- Te) decrease and then the oxidation activity of halogens increase (in the line I ---- F). On the other side, with increase of χa
the oxidation activity of cations-oxidants decrease (in the line Au+ --- Te4+) and then the reducing activity of anions-reductant (S2-, Se2-) increase. These relations are intersected at the standard potential ~ +0.5 V SHE and the intermediate χa,interm
~ 2 that is between electronegativities for metals (tellurium Te) and non-metals (iodine I), Fig.1. In the coordinates χ -

E

o, two average lines are also intersected at ~ +0.5 V SHE (Fig.2). At that, the deviations appear for halogens and transition metals in the cordinates χ -

E

o. There is the maximum of the correlation between the electron work functions of metals
W

e and the standard potentials
E

o of metal electrodes. This maximum is characterized by the same standard potential ~ + 0.5 V SHE and the electron work functions
W

e = ~ 5.2 eV between the known
W

e = ~ 5.0 eV for tellurium and
W

e = ~ 5.4 eV for iodine. This electron work function
W

e = ~ 5.2 eV is the highest one for typical metals (reductants) and the least one for non-metals (oxidants). It is concluded that the intermediate electronegativity (Fig.1-3) and the standard potential ~ +0.5 V SHE should characterize some electrode form (neither metal nor non-metals with the intermediate bond of outer electrons). The intermediate electron bond is not enough unstable for spontaneous giving electrons and not enough stable to accept electrons by such a hypothetical form. Such a form would not be able to reduce or oxidize itself water molecules, i.e., the electrode interaction would be absent. The chemical equilibrium of half-reactions in the center of the electrochemical series (Fig.1-3) is thermodynamically equal to the absence of the electrode interaction. The standard potential ~ +0.5 V SHE is close to Billiter potential +0.475 V SHE [3, 8]. It was determined that the intermediate electron state between thermodynamic stable and unstable electron states is characterized more exactly by the ionization energy ~5.18 ÷ 5.19 eV. [3,10]. The explanations of the correlation in the coordinatesw χ -

E

o are given. The changes of the electronegativity χa -

χa,interm
define the formation and accepting of intermediate non-specific bound electrons in half-reactions - halves of summary oxidation-reduction reactions and changes of free energies in them [4,5].
References
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1988,2, 36-39.
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1982,12, 25- 27.
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(10) A.I.Chernomorskii, The Intermediate Electron Bond and Half-Reactions, Scientific Resources, New York, 1999.

Figure 1