Corrosion thermodynamics cell potential

Each reactant and product appears in the Nemst equation raised to its stoichiometric power. Thermodynamic data for cell potentials have been compiled and graphed (3) as a function of pH. Such graphs are known as Pourbaix diagrams, and are valuable for the study of corrosion, electro deposition, and other phenomena in aqueous solutions.Erom the above thermodynamic analysis, the cell potential can be related to the Gibbs energy change... 

Tower, Stephen. All About Electrochemistry. Available online. URL http //www.cheml.com/acad/webtext/elchem/. Accessed May 28, 2009. Part of a virtual chemistry textbook, this excellent resource explains the basics of electrochemistry, which is important in understanding how fuel cells work. Discussions include galvanic cells and electrodes, cell potentials and thermodynamics, the Nernst equation and its applications, batteries and fuel cells, electrochemical corrosion, and electrolytic cells and electrolysis. 

Aluminium-air cells are attractive in principle because of the high thermodynamic electrode potential and theoretical capacity of aluminium. Because of corrosion, however, the only commercial power sources based on this couple are either reserve systems or are mechanically rechargeable, i.e. the anodes are replaced after discharge. A typical cell may be written as... 

What energy drives an overall corrosion reaction, e.g., Fe + 2H+ — Fe2+ + H2 It is, of course, the free energy of the overall corrosion reaction. It must be negative in value for the direction indicated, for otherwise thermodynamics forbids it to occur. This negative free energy of the overall reaction can be converted into the corresponding cell potential, AG = -nFE (n is 2 in the above reaction because two electrons are involved in the constituent electrochemical reaction, Fe - Fe2+ + 2e, and 2H+ + 2e —>... 

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. 

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... 

Figures 4.3(a) and (b) are sections in the zx-plane showing the distribution of potential (( )) in the solution as cross sections of imaginary surfaces in the solution of equal potential (isopotentials) and the distribution of current as current channels with cross sections defined by traces of the surfaces. ..(n - l),n, (n + 1)... perpendicular to the isopotentials. These traces are located such that each current channel carries the same total current. Figure 4.3(a) applies to an environment of higher resistivity (e.g., water with specific resistivity of 1000 ohm-cm) and Fig. 4.3(b) to an environment of lower resistivity (e.g., salt brine, 50ohm-cm). The figures are representative of anodic and cathodic reactions, which, if uncoupled, would have equilibrium half-cell potentials of E M = -1000 mV and E x = 0 mV and would, therefore, produce a thermodynamic driving force of Ecell = E x - E M = +1000 mV. This positive Ecell indicates that corrosion will occur when the reactions are coupled. For the example of Fig. 4.3(a), the high solution resistivity allows the potential E"m at the anode to approach its equilibrium value (E M = -1000 mV) and, therefore, allows the potential in the solution at the anode interface, < )s a, to approach +1000 mV (recall that (j)s = -E"M). The first isopotential above the anode, 900 mV, approaches this value. The solution isopotentials are observed to decrease progressively and approach 0 mV at the cathode reaction site.

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). 

The reversible potentials can be used to predict the corrosion tendency of the metal when the metal and the electrolyte are under standard thermodynamic conditions described in Chapter 2, Section 2.12.2. Table 6.1 is written as reduction reactions following the guidelines suggested by the International Union of Pure and Applied Chemistry (lUPAC) during the Stockholm Convention in 1953. The procedure for estimating half-cell potential is presented in Chapter 2. In an electrochemical ceU, the electrode with a smaller standard potential in Table 6.1 undergoes oxidation and transfer electrons to the electrode with a larger standard potential, which is reduced at the interface. In redox sys-... 

One thought you may have about corrosion is that we cannot expect standard conditions in real world applications. What must we do to account for the differences that arise when standard conditions are not present This is a very important question with clear ties to thermodynamics. The equation that describes cell potentials under nonstandard conditions is called the Nemst equation ... 

Another basis for thermodynamic prediction of corrosion reactions is the equation AG = —nFE, where n is the number of electrons transferred in the reaction, F is Faraday s constant (96,500 C, which is equivalent to 1 mole of electrons), and A is the cell potential. A positive cell potential will lead to a negative AG , and hence the reaction will proceed spontaneously as written in Eq. (2). As the overall cell reaction proceeds, the voltage difference between the electrodes drops as the free energies of the products and reactants approach each other. At steady state, both electrodes have the same voltage and no charge transfer occurs. 

Given the thermodynamic nature of the measurement and our current understanding of potential measurements, that is probably the best that we can do presently on interpresentation of half cell potentials. A fuller description of interpretation methods for chloride contamination is given in Vassie (1991). For all its limitations, the half cell is a very powerful diagnostic tool for corrosion investigation. Its main problem is that some people rely on too simplistic interpretation. Half cell potentials are not corrosion rate measurements. Corrosion rate measurement is discussed below. 

The present section illustrates how calculations from basic thermodynamic data can lead to open-circuit cell potential in any condition of temperature and pressure. Chemical power sources, with the exception of fuel cells, are all based on the corrosion of a metal connected to the negative terminal. The aluminum-air power source, that owes its energy to the corrosion of aluminum in caustic, was chosen for this example because of the relative simple chemistry. 

The half-cell potentials in the standard emf series are thermodynamic parameters that are valid only at equilibrium corroding systems are not in equilibrium. Furthermore, the magnitudes of these potentials provide no indication as to the rates at which corrosion reactions occur. 

Primarily connected to corrosion concepts, Pourbaix diagrams may be used within the scope of prediction and understanding of the thermodynamic stability of materials under various conditions. Park and Barber [25] have shown this relevance in examining the thermodynamic stabilities of semiconductor binary compounds such as CdS, CdSe, CdTe, and GaP, in relation to their flat band potentials and under conditions related to photoelectrochemical cell performance with different redox couples in solution. 

The physical properties of lithium metal were given in Table 4.4. Despite its obvious attractions as an electrode material, there are severe practical problems associated with its use in liquid form at high temperatures. These are mainly related to the corrosion of supporting materials and containers, pressure build-up and the consequent safety implications. Such difficulties were experienced in the early development of lithium high temperature cells and led to the replacement of pure lithium by lithium alloys, which despite their lower thermodynamic potential remained solid at the temperature of operation and were thus much easier to use. 

Since Ecell is defined to be positive for a spontaneous reaction, this equation correctly expresses a decrease in Gibbs function, which is the thermodynamic criterion for a spontaneous reaction at constant T and P. It is evident that if AGreact can be calculated from AG j- data, the potential of a cell arranged for reversible operation can be determined conversely, experimental measurements of Ecen permit calculation of AGreact. Both types of calculations are useful in electrochemical work and, thus, in the analysis of corrosion. 

The pressure to stop the Ni corrosion cell is much less than the pressure to stop the iron corrosion ceU. Iron is more thermodynamically active with standard potential of -0.44 V vs. SHE when compared with nickel with standard electrode potential ofe° = -0.250V vs. SHE. 

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