Metal-hydrogen potential

In galvanic coupling, titanium is usually the cathode metal and consequently not attacked. The galvanic potential in flowing seawater in relation to other metals is shown in Table 10. Because titanium is a cathode metal, hydrogen absorption may be of concern, as it occurs with titanium complexed to iron (38). 

The dipole density profile p (z) indicates ordered dipoles in the adsorbate layer. The orientation is largely due to the anisotropy of the water-metal interaction potential, which favors configurations in which the oxygen atom is closer to the surface. Most quantum chemical calculations of water near metal surfaces to date predict a significant preference of oxygen-down configurations over hydrogen-down ones at zero electric field (e.g., [48,124,141-145]). The dipole orientation in the second layer is only weakly anisotropic (see also Fig. 7). 

Three kinds of equilibrium potentials are distinguishable. A metal-ion potential exists if a metal and its ions are present in balanced phases, e.g., zinc and zinc ions at the anode of the Daniell element. A redox potential can be found if both phases exchange electrons and the electron exchange is in equilibrium for example, the normal hydrogen half-cell with an electron transfer between hydrogen and protons at the platinum electrode. In the case where a couple of different ions are present, of which only one can cross the phase boundary — a situation which may exist at a semiperme-able membrane — one obtains a so called membrane potential. Well-known examples are the sodium/potassium ion pumps in human cells. 

When such a polyfunctional electrode is polarized, the net current, i, will be given by ii - 4. When the potential is made more negative, the rate of cathodic hydrogen evolution will increase (Fig. 13.2b, point B), and the rate of anodic metal dissolution will decrease (point B ). This effect is known as cathodic protection of the metal. At potentials more negative than the metaTs equilibrium potential, its dissolution ceases completely. When the potential is made more positive, the rate of anodic dissolution will increase (point D). However, at the same time the rate of cathodic hydrogen evolution will decrease (point D ), and the rate of spontaneous metal dissolution (the share of anodic dissolution not associated with the net current but with hydrogen evolution) will also decrease. This phenomenon is known as the difference effect. 

Sometimes anodic protection is used, in which case the metal s potential is made more positive. The rate of spontaneous dissolution will strongly decrease, rather than increase, when the metal s passivation potential is attained under these conditions. To make the potential more positive, one must only accelerate a coupled cathodic reaction, which can be done by adding to the solution oxidizing agents readily undergoing cathodic reduction (e.g., chromate ions). The rate of cathodic hydrogen evolution can also be accelerated when minute amounts of platinum metals, which have a stroug catalytic effect, are iucorporated iuto the metaf s surface fayer (Tomashov, 1955). 

Fig. 4. Catalytic activities of metals (as potentials measured at 10-4 A.cm-2) for anodic oxidation of different reductants. Er thermodynamic oxidation-reduction potentials of reductants. H2 reversible hydrogen electrode potential in solution used to study oxidation of each reductant. Adapted from ref. 38.

Scheme 17.3 Mechanisms of C02 insertion into a metal-hydrogen bond. L represents a potentially dissociable ligand. Ancillary ligands are not shown.

Some researchers have begun to explore the possibihty of combining transition metal catalysts with a protein to generate novel synthetic chemzymes . The transition metal can potentially provide access to novel reaction chemistry with the protein providing the asymmetric environment required for stereoselective transformations. In a recent example from Reetz s group, directed evolution techniques were used to improve the enantioselectivity of a biotinylated metal catalyst linked to streptavidin (Scheme 2.19). The Asn49Val mutant of streptavidin was shown to catalyze the enantioselective hydrogenation of a-acetamidoacrylic acid ester 46 with moderate enantiomeric excess [21]. 

The determination of equilibrium (standard) potentials is rather problematic for several reasons for instance, hydrolysis and disproportionation reactions, the existence of a large number of structural forms (e.g. a-, fi-, y-, 5-Mn02), strong dependence on pH and ionic exchange processes, and the instability of the species in contact with water (e.g. Mn-metal-hydrogen evolution, Mn04 oxygen evolution however, these processes are rather slow). 

When the symmetry factor was introduced by Volmer and Erdey-Gruz in 1930, it was thought to be a simple matter of the fraction of the potential that helps or hinders the transfer of an ion to or from the electrode (Section 7.2). A more molecularly oriented version of the effect of P upon reaction rate was introduced by Butler, who was the first to apply Morse-curve-type thinking to the dependence of theenergy-dis -tance relation in respect to nonsolvent and metal—hydrogen bonds. 

The impure metal dissolves easily in mineral acids and in fluoroboric, sulfamic am trifluoromethylsulfonic acids to give Cr2+ solutions, but oxidation of Cr2+ by hydrogen ion (equation 6), °(Cr3+, Cr2+) = —0.41 V) even in an inert atmosphere is catalyzed by thi impurities and various ions.71 Indefinitely stable chromium(II) solutions can be obtained fron the pure (electrolytic) metal (99.5% or better), although the reaction with acid may need to b< initiated by heat and the inclusion of some metal previously attacked by acid. The use of ai excess of metal, which can be filtered off, ensures that little acid remains. In near neutra solution the hydrogen potential is lowered and the Cr2+ ion is stable. In alkaline condition brown Cr(OH)2, which slowly reduces water, precipitates.73,73... 

The electrode potential is defined as the potential difference between the terminals of a cell constructed of the half-cell in question and a standard hydrogen electrode (or its equivalent) and assuming that the terminal of the latter is at zero volts. Note therefore that the electrode potential is an observable physical quantity and is unaffected by the conventions used for writing cells. The statement. . . the electrode potential of zinc is —0.76 volts. . . implies only that a voltmeter placed across the terminals of a cell consisting of standard hydrogen electrode and the zinc electrode would show this value of potential difference, with the zinc terminal negative with respect to that of the hydrogen electrode. An electrode potential is never a metal/solution potential difference , not even on some arbitrary scale. 

Table 8.1 presents estimates of the differential formation energies for Af—H(s) bonds [A(—AGbF)J at various metal electrodes.3 On the basis of these and Eq. (8.7), approximate formal reduction potentials [ ,m h3o+/w-h(s)] have been estimated for the reduction of H30+ (unit activity) at die several metal electrodes (Table 8.1). In our view direct evaluations of A(—AGBF)M H(S) via electrochemistiy will provide useful insights to metal-hydrogen bonds and metal-catalyzed hydrogenations.

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