All reactions, and particularly redox processes, occur via the lowest-energy pathway that is available (mechanistically feasible) to the system. Metal electrodes are transformed anodically via electron removal from (a) solvent molecules [e.g. Ag(s) + 6H(2)O - e(-) --> Ag-1(OH2)(6)(+)]; (b) electrolyte anions [e.g. Ag(s) + Cl- + e(-) --> Ag-1 Cl(s)]; or (c) Lewis-base ligands [e.g. Ag(s) + 4NH(3) - e(-) --> Ag-1(NH3)(4)(+)]. The same is true for reduced transition-metal complexes [e.g. Fe-II (bpy)(3)(2+) - e(-) --> Fe-III(bpy)(3)(3+) ; ligand-centered oxidation]. In the absence of ligands, most oxidations are mediated (catalyzed) via the electron-transfer transformation of water to protonated-hydroxyl radicals ((H2+O.)), which couple with metal (or unsaturated carbon) centers to form covalent bonds [e.g. Ag(s) +6H(2)O - e(-) --> Ag-1(OH2)(6)(+) --> Ag-1-OH(s) + H-11 + O-5]. Most reductions are mediated (catalyzed) via the electron-transfer transformation of water (or hydronium, ion) to hydrogen atoms (H-.), which couple with unsaturated centers (or functional groups; e.g. -OH of the substrate molecules to form covalent bonds; e.g. H-OH in the case of Ag-1-OH to produce silver metal. The oxidation of metal electrodes involves electron removal (within the interface) from a solvent molecule or basic constituent (ligand) rather than from the valence-electron shell of the metal [e.g. Ag(s) + Cl- - e(-) --> Ag-Cl(s), Edegrees = +0.22 V versus NHE; Cl- - e(-) --> [Cl-.], Edegrees = +2.47 V]. The difference in oxidation potential for the free ligand in the absence of the metal electrode and in its presence is a measure of the metal-ligand differential bond energy [e.g. for Ag-Cl(s), Delta (-DeltaG(BF)) = -DeltaEdegrees x 23.06 kcal (eV)(-1) = 51.9 kcal mol(-1)]. (C) 2002 Elsevier Science B.V. All rights reserved.