Metal-reaction equilibrium potential dependence

As described in Chapter 3, and expressed by Nemst s equation (3.9), the equilibrium potential depends on the activity, or practically on the concentration of the species that take part in the reaction. The reduction of the oxygen eoneentration that oeeurs close to the metal surface when the electrode is polarized in the eathodie direetion, leading to the concentration profile shown in Figure 4.4, causes therefore a change in the equilibrium potential. This equilibrium potential shift is regarded as an overvoltage ... 

Although from the thermodynamic point of view one can speak only about the reversibility of a process (cf. Section 3.1.4), in electrochemistry the term reversible electrode has come to stay. By this term we understand an electrode at which the equilibrium of a given reversible process is established with a rate satisfying the requirements of a given application. If equilibrium is established slowly between the metal and the solution, or is not established at all in the given time period, the electrode will in practice not attain a defined potential and cannot be used to measure individual thermodynamic quantities such as the reaction affinity, ion activity in solution, etc. A special case that is encountered most often is that of electrodes exhibiting a mixed potential, where the measured potential depends on the kinetics of several electrode reactions (see Section 5.8.4). 

From Eqs. (1222) and (12.23), it is clear that the corrosion current depends upon the exchange currents (i.e., available areas and exchange-current densities), Tafel slopes, and equilibrium potentials for both the metal-dissolution and electronation reactions. To obtain an explicit expression for the corrosion current [cf. Eq. (12.22)], one has first to solve Eqs. (12.22) and (12.23) for A0corr. If, however, simplifying assumptions are not made, the algebra becomes unwieldy and leads to highly cumbersome equations. 

The extent of deposition of a metal M introduced as M"+ on a Cu catalyst (or on a metal of lower standard potential) depends on the redox reaction equilibrium constant ... 

A slightly more complicated diagram is that for A1 as shown in Fig. 9. Many valve metals have such fairly simple E-pH diagrams. Lines a and b are present, as always. The solid lines represent three reactions oxidation of A1 metal to Al3+, the equilibrium between Al3+ and A1203, and the equilibrium between A102 and A1203. The latter two are purely chemical reactions, so there is no potential dependence. The oxidation of A1 at low pH does not involve water, whereas it does at higher pH. 

One of the first questions one might ask about forming a metal complex is how strong is the metal ion to ligand binding In other words, what is the equilibrium constant for complex formation A consideration of thermodynamics allows us to quantify this aspect of complex formation and relate it to the electrode potential at which the complex reduces or oxidizes. This will not be the same as the electrode potential of the simple solvated metal ion and will depend on the relative values of the equilibrium constants for forming the oxidized and reduced forms of the complex. The basic thermodynamic equations which are needed here show the relationships between the standard free energy (AG ) of the reaction and the equilibrium constant (K), the heat of reaction, or standard enthalpy (A// ), the standard entropy (AS ) and the standard electrode potential (E for standard reduction of the complex (equations 5.1-5.3). 

The O2/H2O system is very slow so that the exchange current a I equilibrium is extremely low (10 /10 A cm 2) as a consequence, any other reaction at the electrode will hamper its study and that could be the reaction of impurities or other redox reactions involving the electrode itself. The so-called noble metals are not really inert and do interact with oxygen a platinum surface in contact with an O saturated solution adsorbs oxygeti as an electronically conducting monolayer but can be further oxidized to PIO, PtO . A detailed analysis of these phenomena, which falls outside the scope of the present review, can be found elsewhere [311. A platinum electrode, when a complete electronically conducting monolayer of I l—O is formed at the surface of the metal, behaves as an ideally inert electrode in such conditions, rest potentials dependent on pO2 and pH can be measured during a few hours, close to... 

A Gf for an anodic reaction is decreased and A Gt increased by applying a positive voltage with respect to the equilibrium condition (Fig. 7.2b). Accordingto this model, the potential dependence of the interfacial current is caused by the potential dependence of the rate constants. As will be shown later, in this aspect metal electrodes behave completely different from semiconductor electrodes. It also becomes clear from Figs. 7.1 and 7.2 that a variation of the concentration ratio leads to the same effect as that caused by an application of a voltage to the cell. This is reasonable because an increase of tVej speeds up the reaction rate, i.e. the barrier height must be smaller as shown in Fig. 7.1a. 

During each run, the membrane is electrochemically loaded with hydrogen from the left side with a constant electrolysis current of 30 mA. The initial potential of the palladium/palladium hydride (vs. Ag/AgCl reference electrode) on carbon dioxide reaction side at the start of the run is defined as E. This potential depends on the hydride content of the membrane and the equilibrium between the metal hydride/bicarbonate solution. 

Formation of ions during equilibrium electrochemical reaction depends on the pH of the solution and electrode potential. The relationship between electrode potential and pH of the solution can be represented by a phase diagram that is known as the Pourbaix diagram [5]. If a metal is made anodic in an aqueous solution, several reactions can occur depending on the change in free energy. For example, if zinc is made anode in water, the following possible reactions may take place ... 

The reactions described above all involve the displacement of a carbanion from the cobalt atom of methylcobalamin. These reactions occur under aerobic conditions with rate constants in the order of milliseconds. It is apparent that metals which react by electrophilic attack on the Co-C bond (SE2 mechanism) occur with the more oxidized state of the metal, i.e., Pb, Tl, Hg, Pd, which have standard reduction potentials greater than 0.8 volts. Because of the "base on - base off" equilibrium these reactions are pH dependent. 

It can be seen that the corrosion current and potential depend on both the equilibrium potentials for the hydrogen evolution reaction and metal dissolution calculated from the Nernst equation, and the kinetic parameters, the exchange currents and the Tafel slopes. Table 9.1 shows the corrosion currents calculated for some typical values of these parameters it is also important to note that even a... 

Open Circuit Potential. Metal immersed in an aqueous solution develops an electric potential at its surface called open circuit potential (OCP) which is a characteristic of the metal solution system. The magnitude of OCP is measured with respect to reference electrode with the help of high impedance voltmeter and potentiostat is used to polarise or displace equilibrium potential of specimen in the negative (cathodic) or positive (anodic) direction with reference to OCP. This is manipulating the rates (ionic currents) of respective cathodic and anodic half-cell electrochemical reactions. The electrochemical potential of a metal in a certain solution is dependant on the type of the metal, the composition of the solution and its pH, oxygen content and temperature [104, 105]. 

The two-step synthesis (subsequent reduction of the precursor salts) is more likely to result in phase-separated catalysts, possibly core-shell catalysts. Typically a colloid of a particular metal is first formed then subsequently reacted with the salt of the second metal. Numerous examples of such reactions exist in the literature [95, 96]. The deposition of the seeond metal onto the preformed colloid can be driven by a redox reaction if the seeond metal is more noble than the metal of flic preformed colloid, i.e., depending on the difference in the free energy (AG), which is determined by the difference in the equilibrium potentials, Ei and E2 of the redox reactions of the two metal/salts, as follows [97] ... 

Electrochemical reaction kinetics is essential in determining the rate of corrosion of a metal M exposed to a corrosive medium (electrolyte). On the other hand, thermodynamics predicts the possibility of corrosion, but it does not provide information on how slow or fast corrosion occurs. The kinetics of a reaction on a electrode surface depends on the electrode potential. Thus, a reaction rate strongly depends on the rate of electron flow to or from a metal-electrolyte interface. If the electrochemical system (electrode and electrolyte) is at equilibrium, then the net rate of reaction is zero. In comparison, reaction rates are governed by chemical kinetics, while corrosion rates are primarily governed by electrochemical kinetics. 

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