The chemical activities of Ox- and Red-forms are diminishing in opposite directions in the electrochemical series1-4. It was concluded, that a chemical equilibrium of conjugated forms takes place in the center of the series5. The direction of spontaneous half-reactions, leading to formation of double layers, is due to instability of Red- or Ox-forms, since the energy of formation the intermediate bound electrons neinterm in them is const5-7.The transfer from the pair of weak metal - copper Cu ( Eo Cu2+/Cu = + 0.34 V SHE1-4) to the pair of weak non-metal- iodine I2 ( Eo I-/I2 = + 0.535 V SHE1-4 ) is in the center of the series. In formal half - reaction of reductions Cu2+ + 2e = Cu, ions Cu2+ are stable in solutions, copper has some metal activity8. In formal reduction half-reaction ½ I2 + einterm = I-, iodine has non-metal activity and ions I- are stable in solutions8. Hence, spontaneous half-reactions of oxidation Cu → Cu2+ + 2einterm and reduction ½ I2 + einterm → I- take place. Thus, the chemical equilibrium of Red- and Ox-forms in half-reaction is between these spontaneous half-reactions at the average potential ~+0.44 V SHE.The different pairs Men+/Me and Mem+/Me of not active metals (Cu, Ru, Te, Po) are in the same center of the series1-4. These pairs have the close standard potentials: Eo Cu2+/Cu = + 0.34 V SHE1-4 and Eo Cu+/Cu = + 0.52 V SHE1-3; Eo Ru2+/Ru = + 0.455 V SHE2,3 and Eo R u3+/Ru= + 0.38 V SHE2; Eo Te2+/Te= +0.40 V SHE2 and Eo Te4+/Te = +0.568 V SHE2,3; Eo Po2+/Po = = +0.368 V SHE1,2 and Eo Po4+/Po = + 0.73 V SHE1 (+0.775V SHE2; +0.76 V SHE3). Hence, the pairs Men+/Me and Mem+/Me of each metal should have the potential determining ions with the close energy formation in solutions. However, these ions are also known as thermodynamic stable and unstable in solutions (for example, Cu2+ and Cu+; Ru2+, Ru3+ and Ru4+; Te2+ and Te4+; Po2+ and Po4+)8. Hence, their spontaneous oxidation and reduction half-reactions (due to instability of metal or its ions) take place by both sides from the chemical equilibrium. The average potential calculated by the sum of these standard potentials is equal + 0.472 V SHE (+0.477 V SHE).Moreover, the standard potentials of the redox pairs Men+ / Mem+ formed by different ions of these metals are also close to their Eo Men+/ Me and Eo Mem+/Me ( Eo Cu2+/Cu+ = +0.159 V SHE 1-3; Eo Ru3+/Ru2+ = +0.249 V SHE1-3; Eo Ru4+/ Ru3+ = +0.49 V SHE9; Eo Te4+/ Te2+= +0.736 V SHE, calculated by Latimer equation; Eo Po3+/Po2+ = +0.330 V SHE2; Eo Po4+/Po2+ = + 0.9 V SHE3). The average potential found from the sum of these verified redox potentials is equal to +0.477 V SHE. The standard potentials Eo of such pairs are differed for more active metals (for example, Eo Fe 3+ / Fe = - 0.037 V SHE, Eo Fe 2+ / Fe = - 0.473 V SHE, Eo Fe 3+ /Fe +2 = + 0.771 V SHE)1-4.The closeness of the standard potentials Eo of the pairs Men+/Me, Mem+/Me and Men+/Mem+ pairs exists only for weak active metals and non-metal which outer electrons are not bound by strong metallic or non-metallic bonds. The standard potential ~+0.475 V SHE10 of the chemical equilibrium of half-reactions is confirmed by analysis of other half-reactions of different kind5,6. The electrode potential ~+0.475 V SHE can therefore be used as datum for measuring the potential differences in electrical double layers due to spontaneous half-reactions on electrodes. By such an approach, the changes of the Gibbs energies in spontaneous half-reactions are calculated by the differences E equilibrium - (+0.475 V SHE), where E equilibrium are the equilibrium potentials of electrodes. Hence, the change of the Gibbs energy in the spontaneous half-reaction of the standard hydrogen electrode ½ H2 → H++ einterm (0 V SHE) is equal - 0.475 eV11,12.