Equilibrium electrode potential difference

Considering Equation (4), it is clear that one has to operate a galvanic cell at the maximum possible cell voltage in order to maximize the electric energy yield. As shown previously (Figure 3.1.6), the maximum value of U is the equilibrium electrode potential difference E or E°. Thus, one may formulate the fundamental relationship between chemical and electric energy ... 

Here Eq is the equilibrium electrode potential difference between the two electrodes of the galvanic cell. It is now possible to choose the standard reference electrode as the hydrogen electrode, for which the chemical potentials jXj of the constituents and H2 are taken to be zero at standard temperature and pressure. The equilibrium condition 7.14 then becomes... 

Equilibrium electrode potentials are readily established when metal electrodes are in contact with melts. However, two difficnlties arise in attempts to measnre them suitable, sufficiently corrosion-resistant reference electrodes must be selected, and marked diffusion potentials develop at interfaces between different melts. 

According to Eq. (3.21) and taking into account that the electrode potential differs by a constant term from the metal-solution Galvani potential, we thus have an expression for the equilibrium potential of this electrode and, at the same time, for the equilibrium potential, redox of this redox system ... 

If current passes through an electrolytic cell, then the potential of each of the electrodes attains a value different from the equilibrium value that the electrode should have in the same system in the absence of current flow. This phenomenon is termed electrode polarization. When a single electrode reaction occurs at a given current density at the electrode, then the degree of polarization can be defined in terms of the over potential. The overpotential r) is equal to the electrode potential E under the given conditions minus the equilibrium electrode potential corresponding to the considered electrode reaction Ec ... 

In contrast to the equilibrium electrode potential, the mixed potential is given by a non-equilibrium state of two different electrode processes and is accompanied by a spontaneous change in the system. Besides an electrode reaction, the rate-controlling step of one of these processes can be a transport process. For example, in the dissolution of mercury in nitric acid, the cathodic process is the reduction of nitric acid to nitrous acid and the anodic process is the ionization of mercury. The anodic process is controlled by the transport of mercuric ions from the electrode this process is accelerated, for example, by stirring (see Fig. 5.54B), resulting in a shift of the mixed potential to a more negative value, E mix. 

Initially, both electrodes are at equilibrium. Since the anode has accumulated electrons and the cathode has depleted electrons, electrons begin to flow from electrode from the anode to the cathode. The thermodynamic driving force for the electron flow is the electrode potential difference, which for the fuel cell reaction is 1.23 Y at standard conditions. In addition to electron flow, H + ions produced at the anode diffuse through the bulk solution and react at the cathode. The reaction is able to continue as long as H2 is fed at the anode and 02 at the cathode. Hence, the cell is not at equilibrium. The shift in electrode potential from equilibrium is called the overpotential (>/). 

If a hydrogen electrode be immersed in each solution since the Aystem is in equilibrium, the potential difference between these electrodes must be zero. Similarly for the chlorine ion, there will be a zero potential difference between two silver-silver chloride electrodes immersed in the two media. 

A simple way of visualizing the procedure is to represent all equilibrium potentials on a single vertical axis (Fig. 7.176). Corresponding to any activity ratio aA/aD, there is an equilibrium electrode potential for the interface relative to the SHE. The same is true for the other activity ratio a A/a D. Thus, the separation between any two points yields the potential difference across a cell, with the activity ratios corresponding to the points. 

If both electrode processes operate under standard conditions, this voltage is E°, the equilibrium standard electrode potential difference. Values of E and E° may be conveniently measured with electrometers of so large an internal resistance that the current flow is nearly zero. Figure 3.1.6 illustrates the measurement and the equilibrium state. The value of E° is a most significant quantity characterizing the thermodynamics of an electrochemical cell. Various important features of E and E° will be addressed in the following chapters. 

Among the different surface atom positions illustrated on Fig. 2.8, the kink site position, or the half crystal position, as introduced independently by Kossel [2.12] and Stranski [2.13], has a special significance for the definition of the equilibrium conditions (vapor pressure, equilibrium concentration, equilibrium electrode potential, etc.) of the infinitely large (bulk) crystal. 

We now estimate the anode/cathode potential difference during current flow for driven and self-driven cells. First, we define the cell potential (V) as the difference in the anode and cathode potentials. For a self-driven cell, the cell voltage at a given current density will be less than the difference in equilibrium electrode potentials (AEJ due to the presence of various overpotentials. 

During current passage, the electrode potential differs from its equilibrium value, E, since the concentration of metal... 

If an electrochemical reaction is perturbed from the equilibrium state, the relative stabilities of the species in the reaction are changed. The manifestation of the perturbation is the measured electrode potential, which differs from the equilibrium electrode potential for the reaction. If the measured electrode potential is positive with respect to the equilibrium potential, the reaction given by Equation 4 proceeds irreversibly from left to right, i.e., the reduced form of the chemical species is unstable while the oxidized form of the species is stable. The converse is true when the... 

Mixed potential systems with the cathodic partial process under transport control and the anodic partial process under activation control is typical of many corrosion systems. For the cathodic partial process to be under transport control. Equation 44 must be unity or larger. This occurs when the absolute value of the difference between the equilibrium electrode potential of the cathodic partial process and the corrosion is on the order of one volt. This condition prevails for most metals of interest in corrosion studies if oxygen... 

The E° and E° values shown in parentheses are the equilibrium electrode potentials for reactions (142) and (143) at 25°C and with reactants and products at unit activity. In practice, chlor-alkali cells operate at different temperatures and concentrations, hence, the El and ° terms should be properly corrected using the Nemst equation. Typical conditions encountered in diaphragm-type chlor-alkali cells are as follows ... 

The minimum cell potential is determined by the thermodynamics of the overall chemical change in the cell, and this can only be reduced by changing an electrode reaction (e.g., replacing H2 evolution by O2 reduction in a chlor-alkali cell). The equilibrium cell potential difference, is related to the Gibbs free energy change for the cell reaction, AGj.gn by... 

Self-discharging at the positive electrode can be driven by the electrode potential difference. For example, the equilibrium potential of an Pb02/PbS04 electrode is 0.6 V more positive than that of an oxygen H2O/O2 electrode. This potential difference can make the following reaction happen, consuming Pb02 ... 

Equilibrium Electrode Potential The difference in potential between an electrode and an electrolyte when they are in equilibrium for the electrode reaction which determines the electrode potential. 

Fig. 6.12 (a) Schematic potential energy curve of ORR at an electrode potential of 0 V vs. RHE black line), 1.23 V vs. RHE (equilibrium electrode potential, red line), and at 0.8 V vs. RHE LGJe, blue line). Energy difference of every step at 0 V vs. RHE is marked with AGj. (b) Dependency of AGi and AG4 on the dissociated adsorption energy of oxygen. The fourth step limits the activity of ORR on the Pt(l 11) facet. The dashed line indicates the equihbtium electrode potential of ORR. Modified with permission from ref [53]... 

The reason why it is not possible to measure electrode potentials is of an operational nature. An absolute electrode potential is the potential difference settled at equilibrium between the electrode metal and the ionic solution where it dips. The measurement necessitates introducing another electrode into the solution. As a result, a supplementary electrode potential exists. The measurement with the help of a voltmeter cannot give another value than that of an electrode potential difference. 

In fact, this strong pH-dependence of the equilibrium Galvani potential difference can be observed with the above construction. Such a large pH-dependence can even be found with a suspension of Raney nickel in contact with a normal platinum electrode. Attempts with other metals, which are known not to catalyze reaction (R2), yielded irreproducible results. This example serves to illustrate the requirement of an unobstructed phase transfer, which in this case was coupled with a chemical reaction. 

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