Electrodes hydrogen ions reduction

Cathodic protection also can be accomplished by lowering the electrode potential to E M, the equilibrium potential for the metal to be protected, by an external power source. The circuit used to accomplish this is the same as shown in Fig. 2.12. With slight modification, it is again shown in Fig. 4.25 in which the metal to be protected is iron and the cathodic reaction supporting corrosion is either hydrogen-ion reduction, oxygen reduction, or both. 

It is the preconcentration period that enhances the sensitivity of this technique. In the preconcentration phase precise potential control permits the selection of species whose decomposition potentials are exceeded. The products should form an insoluble solid deposit or an alloy with the substrate. At Hg electrodes the electroreduced metal ions form an amalgam. Usually the potential is set 100-200 mV in excess of the decomposition potential of the analyte of interest. Moreover, electrolysis may be carried out at a sufficiently negative potential to reduce aU of the metal ions possible below hydrogen ion reduction at Hg, for example. Concurrent H" " ion reduction is not a problem, because the objective is to separate the reactants from the bulk electrolyte. In fact, methods have been devised to determine the group I metals and NEC " ion at Hg in neutral or alkaline solutions of the tetraalkylammonium salts. Exhaustive electrolysis is not mandatory and 2-3% removal suffices. Additionally, the processes of interest need not be 100% faradaically efficient, provided that the preconcentration stage is reproducible for calibration purposes, which is usually ensured by standard addition. 

Activation Polarization. Activation polarization is caused by a slow electrode reaction. The reaction at the electrode requires an activation energy in order to proceed. The most important example is that of hydrogen ion reduction at a cathode, H + - H2. For this reaction, the polarization is called hydrogen... 

In practice, electrode reactions other than hydrogen ion reduction are used to construct practical reference systems. Potentials determined using these halfcells can be related back to the hydrogen reference or absolute potential scale if... 

If it is assumed that the hydrogen and catalytically active ions are reduced independently, and each accordii to its own rules [93], it follows from Fig. 17 that change in catalsdic current density (lower part of Curves 2-4) leads to a considerably larger shift in the potential of the electrode than hydrogen ion reduction, and with increase in potential the catalytic wave is therefore simply overlapped by the normal hydrogen-ion reduction current. 

Although Table 2.16 shows which metal of a couple will be the anode and will thus corrode more rapidly, little information regarding the corrosion current, and hence the corrosion rate, can be obtained from the e.m.f. of the cell. The kinetics of the corrosion reaction will be determined by the rates of the electrode processes and the corrosion rates of the anode of the couple will depend on the rate of reduction of hydrogen ions or dissolved oxygen at the cathode metal (Section 1.4). 

The formal potential of a reduction-oxidation electrode is defined as the equilibrium potential at the unit concentration ratio of the oxidized and reduced forms of the given redox system (the actual concentrations of these two forms should not be too low). If, in addition to the concentrations of the reduced and oxidized forms, the Nernst equation also contains the concentration of some other species, then this concentration must equal unity. This is mostly the concentration of hydrogen ions. If the concentration of some species appearing in the Nernst equation is not equal to unity, then it must be precisely specified and the term apparent formal potential is then employed to designate the potential of this electrode. 

Haapakka and Kankare have studied this phenomenon and used it to determine various analytes that are active at the electrode surface [44-46], Some metal ions have been shown to catalyze ECL at oxide-covered aluminum electrodes during the reduction of hydrogen peroxide in particular. These include mercu-ry(I), mercury(II), copper(II), silver , and thallium , the latter determined to a detection limit of <10 10 M. The emission is enhanced by organic compounds that are themselves fluorescent or that form fluorescent chelates with the aluminum ion. Both salicylic acid and micelle solubilized polyaromatic hydrocarbons have been determined in this way to a limit of detection in the order of 10 8M. 

The standard cell potential for the reduction of hydrogen ions to hydrogen gas is, by definition, 0.00 V. This potential is for the standard hydrogen electrode, SHE, which is the reference to which we compare all other cell potentials. All metals above hydrogen on the Activity Series will displace hydrogen gas from acids. (See Chapter 4) Metals below hydrogen will not displace hydrogen gas. 

Further, these anodic and cathodic reactions can occur spatially at adjacent locations on the stuface of a metal electrode rather than on two separated metal electrodes as shown in Fig. 11-1, where the anodic dissolution of iron and the cathodic reduction of hydrogen ions proceed simultaneously on an iron electrode in aqueous solution. The electrons produced in the anodic dissolution of iron are the same electrons involved in the cathodic reduction of hydrogen ions hence, the anodic reaction cannot proceed more rapidly than that the electrons can be accepted by the cathodic reaction and vice versa. Such an electrode at which a pair of anodic and cathodic reactions proceeds is called the mixed electrode . For the mixed electrodes, the anode (current entrance) and the cathode (current exit) coexist on the same electrode interface. The concept of the mixed electrode was first introduced in the field of corrosion science of metals [Evans, 1946 Wagner-Traud, 1938]. 

Fig. 11-1. Mixed electrode model (local cell model) for corrosion of metals i = anodic current for transfer of iron ions i = cathodic current of electron transfer for reduction of hydrogen ions.

The potential of a mixed electrode at which a coupled reaction of charge transfer proceeds is called the mixed electrode potential , this mixed electrode potential is obviously different from the single electrode potential at which a single reaction of charge transfer is at equilibrium. For corroding metal electrodes, as shown in Fig. 11—2, the mixed potential is often called the corrosion potential, E . At this corrosion potential Eemt the anodic transfer current of metallic ions i, which corresponds to the corrosion rate (the corrosion current ), is exactly balanced with the cathodic transfer current of electrons for reduction of oxidants (e.g. hydrogen ions) i as shown in Eqn. 11-4 ... 

Figure 11-7 shows the polarization curve of an iron electrode in an acidic solution in which the anodic reaction is the anodic transfer of iron ions for metal dissolution (Tafel slope 40 mV/decade) the cathodic reaction is the cathodic transfer of electrons for reduction of hydrogen ions (Tafel slope 120 mV /decade) across the interface of iron electrode. 

The maximum reduction limit for the thermodynamic stability of water is defined by the reduction of hydrogen ions in solution with the formation of diatomic gaseous molecules H2 (normal hydrogen electrode) ... 

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