Metal half-cell potentials

The metals must have sufficiently positive half-cell potentials that corrosive reactions that would change the potential to a corrosion potential, Ecorr, do not occur. This requirement, with few exceptions, restricts the metal component of the half cell to silver, mercury, and copper. For these metals, appearing in Table 6.1, corrosion due to hydrogen evolution will not occur since the metal half-cell potentials are above potentials for hydrogen evolution. Also, the kinetics of the reduction of any dissolved oxygen are sufficiently slow that the potential is shifted negligibly from that of the metal half cell. 

The value of this electric field depends on the electrochemical potentials of the electrode metals and is related to their position in the electrochemical series. Table l7.3 shows some values for some common electrode metals at romn temperature. Electrochemists have by convention adopted the hydrogen electrode as a standard of reference potmitial and assigned it a value of 0 V. All other metals have a ncmzero potential with respect to it Metal half-cell potentials depend on their electrochemical oxidation state, and they are usually arranged in a table showing their activity relative to others, such as seen in Table 17.3. 

Test method for porosity in gold platings on metal substrates by gas exposures Test method for half-cell potentials of uncoated reinforcing steel in concrete Method for detection of copper corrosion from petroleum products by the copper strip tarnish test... 

In Table 7-1 the relative tendencies of certain elements to react were listed qualitatively. We can give a quantitative measure of relative tendency to react, called standard reduction potential, as shown in Table 14-2. In this table, the standard half-cell potential for each half-reaction, as a reduction, is tabulated in order with the highest potential first. If we turn these half-reactions around, we change the signs of the potentials and we get oxidation potentials. We thus have half-reactions including both elementary metals and elementary nonmetals in the same table, as well as many half-reactions that do... 

The intimate relationship between double layer emersion and parameters fundamental to electrochemical interfaces is shown. The surface dipole layer (xs) of 80SS sat. KC1 electrolyte is measured as the difference in outer potentials of an emersed oxide-coated Au electrode and the electrolyte. The value of +0.050 V compares favorably with previous determinations of g. Emersion of Au is discussed in terms of UHV work function measurements and the relationship between emersed electrodes and absolute half-cell potentials. Results show that either the accepted work function value of Hg in N2 is off by 0.4 eV, or the dipole contribution to the double layer (perhaps the "jellium" surface dipole layer of noble metal electrodes) changes by 0.4 V between solution and UHV. 

Figure 2. Work function of polycrystalline Au electrode emersed from 0.1 M HC104 as a function of emersion potential. The work function of the clean metal surface was 5.2 eV (19). If the NHE absolute half-cell potential (with respect to s) is 4.45 V, the bottom line is equal to the solution inner potential, tfg. If it is 4.85 V, the upper curve through the points is equal to 0g. As always with WF, the Fermi level is taken as zero.

For the first two types, a table of metals relating their ease of oxidation to each other is useful in being able to predict what displaces what. Table 6.1 shows the activity series for metals, which lists the metal and its oxidation in order of decreasing ease of oxidation. An alternative to the activity series is a table of half-cell potentials, as discussed in Chapter 16. In general, the more active the metal, the lower its potential. 

Tables containing the same sequence of reactions as in Table 1, but without the voltage data, were in common use long before electrochemical cells were studied and half-cell potentials had been measured. If you read down the central column, you will notice that it begins with the sequence of metals Na, Zn, Fe, etc. This sequence is known as the activity series of the metals, and expresses the decreasing tendency these species to lose electrons- that is, to undergo oxidation.

Factors Involved in Galvanic Corrosion. Emf series and practical nobility of metals and metalloids. The emf. series is a list of half-cell potentials proportional to the free energy changes of the corresponding reversible half-cell reactions for standard state of unit activity with respect to the standard hydrogen electrode (SHE). This is also known as Nernst scale of solution potentials since it allows to classification of the metals in order of nobility according to the value of the equilibrium potential of their reaction of dissolution in the standard state (1 g ion/1). This thermodynamic nobility can differ from practical nobility due to the formation of a passive layer and electrochemical kinetics. 

The facility with which metal complexes bring about reactions 8.16 and 8.17 depends on several factors, one of the important ones being the half-cell potential (E°) of the M"+/M(n+1)+ couple. It should be remembered, however, that most E° values for metal ions have been measured in an aqueous environment. On complexation and in an organic liquid these values are expected to change substantially. The initial hydroperoxide required for metal-catalyzed decomposition, reactions 8.16 and 8.17, is normally present in trace quantities in most hydrocarbons. 

Iron or steel is often covered by a thin layer of a second metal to prevent rusting Tin cans consist of steel covered with tin, and galvanized iron is made by coating iron with a layer of zinc. If the protective layer is broken, however, iron will rust more readily in a tin can than in galvanized iron. Explain this observation by comparing the half-cell potentials of iron, tin, and zinc. 

The silver-silver chloride electrode is an example of a metal electrode that participates as a member of a redox couple. The silver-silver chloride electrode consists of a silver wire or rod coated with AgCl(s) that is immersed in a chloride solution of constant activity this sets the half-cell potential. The Ag/AgCl electrode is itself considered a potentiometric electrode, as its phase boundary potential is governed by an oxidation-reduction electron transfer equilibrium reaction that occurs at the surface of the silver ... 

The potentials of half-cells that are constructed using amalgams are important whenever one is interested in performing reactions in a media in which hydrogen evolution at the metal-solution interface might interfere with the reaction of interest. Consequently, half-cell potentials for alkali metal amalgams are important to synthetic chemists and electrochemists alike. The standard potential for a metal amalgam in contact with a solution of its monovalent cation ( °(M" /M(Hg))) can be related to that of the pure metal in contact with its monovalent cation (EJ ) as shown in Eq. (14)... 

The problem of drifting cell potentials seems to arise from the fact that no defined reference potential exists at the interface between the internal metal wire and the polymeric membrane. Some workers postulate an oxygen-dependent half-cell potential at this interface with the required oxygen being able to permeate the PVC film (S2). Clearly, the question as to what is happening at the interfoce of the metal and the membrane needs to be resolved before CWEs of the above type gain widespread acceptance as useful analytical devices. 

We have seen in Section 26.2.1 that thermodynamics (i.e., equilibrium half-cell potentials) can be used to determine which of two half-cell reactions proceeds spontaneously in the anodic or cathodic direction when the two reactions occur on the same piece of metal or on two metal samples that are in electrical contact with one another. The half-cell reaction with the higher equilibrium potential will always be at the cathode. Thus, under standard conditions any metal dissolution (corrosion) reaction with an E° less than 0.0 V vs. SHE will be driven by proton reduction while metal dissolntion reactions with an E° less than -e1.23 V vs. SHE will be driven by dissolved... 

A pure metal can be anodic only if its equilibrium half-cell potential, E M, is less than the half-cell potential of some cathodic reaction, E x, such that the total cell potential (Ex - E" ) causes current to flow as in Fig. 1.6, that is, current away from the anode area as ions in the solution. A few representative anodic reactions are listed in Table 1.3 along with their standard equilibrium half-cell potentials. 

The initial consideration in analyzing an existing or proposed metal/environment combination for possible corrosion is determination of the stability of the system. According to Eq 1.18, the criterion is whether the equilibrium half-cell potential for an assumed cathodic reaction, E x, is greater than the equilibrium half-cell potential for the anodic reaction, E M. A convenient representation of relative positions of equilibrium half-cell potentials of several common metals and selected possible corrodent species is given in Fig. 1.7. To the left is the scale of potentials in millivolts relative to the standard hydrogen electrode (SHE). The solid vertical lines identified by the name of the metal give... 

In the derivations of Eq 3.14 and 3.19 for the metal oxidation current density, iox M, and the metal-ion reduction current density, ired M, it was not necessary to restrict the half-cell potential to its equilibrium value. Deviation from E M will occur if the potential of either the metal or the solution is changed, resulting in an overpotential defined in general by Eq 3.1. More specifically, small deviations are associated with charge-transfer polarization, and the overpotential is designated as ... 

In Chapter 2 (in the section Interface Potential Difference and Half-Cell Potential ), the equilibrium half-cell potential for the metal reaction, E M, was defined relative to potentials < > as follows ... 

This case is illustrated in Fig. 4.21. The only change from Case 1 is the position of the oxidation curve for metal A, which is now placed sufficiently positive that its equilibrium half-cell potential is above that for... 

Cathodic protection is the process whereby the corrosion rate of a metal is decreased or stopped by decreasing the potential of the metal from Ecorr to some lower value and in the limit to E M, the thermodynamic equilibrium half-cell potential. At this potential, iox M = ired X[ = i() xi- and net transfer of metal ions to the solution no longer occurs. This is the criterion for complete cathodic protection (i.e., E = E m). 

For metals such as titanium and chromium, the active peak in the anodic polarization curve may occur below the half-cell potential for the... 

The standard reduction potential (E°) provides a measure of the stability of a metal in a particular oxidation state. The E° value is the voltage generated in a half-cell coupled with the standard hydrogen electrode (SHE), which itself has a defined half-cell potential of 0.0 V. Put simply, the more positive is E0 the more difficult is it for metal oxidation to a hydrated metal ion to occur. Alternatively, we could express it by saying that the less positive is °, the more stable is the metal in the higher oxidation state of its couple... 

The net effect of either complexation or chelation of a metal ion is to reduce the concentration of the metal ion at the electrode. This changes the potential and can be determined from the Nemst equation. Usually, a half cell potential such as... 

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