The Standard Hydrogen Potential

In any galvanic cell an electrical potential difference exists between the two electrodes. Potential difference is measured in volts. The two sides of a galvanic cell are called half- 

A series circuit is the one used in flashUghts and other devices that use batteries—if two or more batteries are used together, they are coimected + to -. (If circuit elements are cormected + to + and - to -, it is called a parallel circuit. We can ignore parallel circuits in this discussion because galvanic cells always have series circuits.) 

When we cormect a voltmeter to the electrodes of a galvanic cell, we are measuring the electrical potential difference of the total cell, which is the sum of the potential differences of the two half-cells. Since we cannot measure the potential difference of a half-cell by itself, chemists have chosen to assign the standard hydrogen electrode (SHE) a potential difference of 0.00 V. The standard hydrogen electrode consists of a platinum electrode immersed in a 1 M HCl solution with hydrogen gas bubbling into the solution at a pressure of 1 atmosphere (atm). 

If the hydrogen electrode is the anode, then the half-reaction taking place is 

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Noble Potential—a potential that is more cathodic (i.e., positive) than the standard hydrogen potential. 

The redox potential is generally referred to the standard hydrogen potential (SHE), which has an exactly defined energy, E y, relative to the energy of the free electron in vacuum or at infinity. Thus, electrode potentials of redox couples can be expressed on the absolute energy scale according to... 

The standard hydrogen potential is the reference potential against which all half-reaction potentials are assigned. 

It is obvious that this reaction can only lead to metal dissolution, if the metal electrode potential is negative from the hydrogen electrode potential. This is the reason for the classification of metals into noble metals (the equilibrium potential is more positive than the standard hydrogen potential) and non-noble metals (the equilibrium potential is more negative than the standard hydrogen potential). The kinetic of the total process can be described by the Butler-Volmer equation for the two partial reactions. 

Hj. All values refer to the standard hydrogen potential (SHE). Reproduced with permission from Santos et al. [68], figure 1. 2011, Wiley. 

The hydrogen source comes from the water. Theoretically, any metal with "activity" higher than that of H2 could react with water to form H2. In other words, any metal with a standard reduction potential lower than 0 V (the standard hydrogen potential) could react with water to form H2. Let us take the reaction between Na and water as an example ... 

The calculation of electrode potential E) from Eq. (10.9) requires a series of standard electrode potential (Eq) data. It is well accepted that the standard hydrogen potential is standard electrode potential in the world. Its electrode reaction is ... 

Standard Hydrogen Electrode The standard hydrogen electrode (SHE) is rarely used for routine analytical work, but is important because it is the reference electrode used to establish standard-state potentials for other half-reactions. The SHE consists of a Pt electrode immersed in a solution in which the hydrogen ion activity is 1.00 and in which H2 gas is bubbled at a pressure of 1 atm (Figure 11.7). A conventional salt bridge connects the SHE to the indicator half-cell. The shorthand notation for the standard hydrogen electrode is... 

The standard electrode potential for zinc reduction (—0.763 V) is much more cathodic than the potential for hydrogen evolution, and the two reactions proceed simultaneously, thereby reducing the electrochemical yield of zinc. Current efficiencies slightly above 90% are achieved in modem plants by careful purification of the electrolyte to bring the concentration of the most harmful impurities, eg, germanium, arsenic, and antimony, down to ca 0.01 mg/L. Addition of organic surfactants (qv) like glue, improves the quaUty of the deposit and the current efficiency. 

Laboratory experiments have shown that IGSCC can be mitigated if the electrochemical potential (ECP) could be decreased to —0.230 V on the standard hydrogen electrode (SHE) scale in water with a conductivity of 0.3 ]lS/cm (22). This has also been demonstrated in operating plants. Equipment has been developed to monitor ECP in the recirculation line and in strategic places such as the core top and core bottom, in the reactor vessel during power operation. 

Fig. 6. Band edge positions of several semiconductors ia contact with an aqueous electrolyte at pH 1 ia relation to the redox (electrode) potential regions (vs the standard hydrogen electrode) for the oxidation of organic functional groups (26,27).

Other Coordination Complexes. Because carbonate and bicarbonate are commonly found under environmental conditions in water, and because carbonate complexes Pu readily in most oxidation states, Pu carbonato complexes have been studied extensively. The reduction potentials vs the standard hydrogen electrode of Pu(VI)/(V) shifts from 0.916 to 0.33 V and the Pu(IV)/(III) potential shifts from 1.48 to -0.50 V in 1 Tf carbonate. These shifts indicate strong carbonate complexation. Electrochemistry, reaction kinetics, and spectroscopy of plutonium carbonates in solution have been reviewed (113). The solubiUty of Pu(IV) in aqueous carbonate solutions has been measured, and the stabiUty constants of hydroxycarbonato complexes have been calculated (Fig. 6b) (90). 

Solution Potential. The standard electrode potential of aluminum (A1 + 3e) is —1.66 V on the standard hydrogen scale and —1.99 V... 

The thermodynamics of electrochemical reactions can be understood by considering the standard electrode potential, the potential of a reaction under standard conditions of temperature and pressure where all reactants and products are at unit activity. Table 1 Hsts a variety of standard electrode potentials. The standard potential is expressed relative to the standard hydrogen reference electrode potential in units of volts. A given reaction tends to proceed in the anodic direction, ie, toward the oxidation reaction, if the potential of the reaction is positive with respect to the standard potential. Conversely, a movement of the potential in the negative direction away from the standard potential encourages a cathodic or reduction reaction. 

Thus, because the standard cell potential for reaction 15 is positive, the reaction proceeds spontaneously as written. Consequendy, to produce chlorine and hydrogen gas, a potential must be appHed to the cell that is greater than the open-circuit value. This then becomes an example of an electrolytic process. 

If electron flow between the electrodes is toward the sample half-cell, reduction occurs spontaneously in the sample half-cell, and the reduction potential is said to be positive. If electron flow between the electrodes is away from the sample half-cell and toward the reference cell, the reduction potential is said to be negative because electron loss (oxidation) is occurring in the sample halfcell. Strictly speaking, the standard reduction potential, is the electromotive force generated at 25°C and pH 7.0 by a sample half-cell (containing 1 M concentrations of the oxidized and reduced species) with respect to a reference half-cell. (Note that the reduction potential of the hydrogen half-cell is pH-dependent. The standard reduction potential, 0.0 V, assumes 1 MH. The hydrogen half-cell measured at pH 7.0 has an of —0.421 V.)... 

Some typical half-cell reactions and their respective standard reduction potentials are listed in Table 21.1. Whenever reactions of this type are tabulated, they are uniformly written as reduction reactions, regardless of what occurs in the given half-cell. The sign of the standard reduction potential indicates which reaction really occurs when the given half-cell is combined with the reference hydrogen half-cell. Redox couples that have large positive reduction potentials... 

It must not be assumed that the protection potential is numerically equal to the equilibrium potential for the iron/ferrous-ion electrode (E ). The standard equilibrium potential (E ) for iron/ferrous-ion is -0-440V (vs. the standard hydrogen electrode). If the interfacial ferrous ion concentration when corrosion ceases is approximately 10 g ions/1 then, according to the Nernst equation, the equilibrium potential (E ) is given by ... 

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