Equilibrium-hydrogen thermodynamic diagrams

The areas bounded by solid lines correspond to regions of thermodynamic stability of certain substances that are named in the diagram. This stability is relative. The dashed line a in the diagram corresponds to the equilibrium potential of the hydrogen electrode. Metallic zinc, for which the reaction lines are below the line for the hydrogen electrode, can be oxidized while hydrogen is evolved (see Section 2.4.1). 

The Hammond Postulate applies only if both forward reactions are fast. Obtain energies for the transition states leading to 1 -propyl and 2-propyl radicals (propane+F end andpropane+F center). Draw an energy diagram for each hydrogen abstraction reaction (place the diagrams on the same axes). Do these diagrams indicate that use of the Hammond Postulate is justified Calculate the barrier for each reaction, and calculate the relative concentrations of 1-propyl and 2-propyl radicals that would form at 298 K if each reaction were irreversible. Use equation (2). How does this (kinetic) ratio compare to the equilibrium (thermodynamic) ratio of these radicals ... 

Figure 1. Thermodynamic equilibrium in atmospheres of varying elemental proportions. The ternary diagram shows all compositions of systems containing carbon, hydrogen, and oxygen (each point represents 100% of the three components). Lower curves indicate the potential formation of solid carbon if equilibrium could be attained. Dashed curve holds at 500°K., the continuous one at 700°K. The upper lines indicate the asphalt threshold, the dashed one at 500° K., and the continuous one at 700° K. Above this threshold, thermodynamic equilibrium favors the formation of large proportions of polycyclic aromatic compounds ( asphalt ) ana a lesser increase of most of the other families of compounds. The dots through points A to C indicate the points used in the computations for Figure 2 (6).

Table 2.1 lists equilibrium ratios for the reduction of selected metal oxides [4], while Figure 2.2 provides a complete phase diagram for the reduction of iron oxide at different temperatures [3, 5], In order to reduce bulk iron oxide to metallic iron at 600 K, the water content of the hydrogen gas above the sample must be below a few percent, which is easily achieved. However, in order to reduce Cr2C>3, the water content should be as low as a few parts per billion, which is much more difficult to realize. The data in Table 2.1 also illustrate that, in many cases, only partial reduction to a lower oxide may be expected. Reduction of Mn2C>3 to MnO is thermodynamically allowed at relatively high water contents, but further reduction to manganese is unlikely. 

The skills developed to produce the equilibrium diagram Figure A.l, are now applied anew. Neither hydrogen nor carbon monoxide occur as free substances in nature, where they are immediately oxidized. They must be made and stored, at thermodynamic and economic cost. The reversible thermodynamics are assessed below, using as the basis of calculation a notional, electrochemical, equilibrium, steam reformer. Figure A.4, for comparison with the alternative practical and irreversible combustion-driven reformers. 

In general, if, for a given pH, the reversible potential of a metal is below the reversible potential of the hydrogen electrode, the metal can corrode by reaction with protons with the production of hydrogen. Keep in mind, however, that thermodynamic considerations are only valid for the prediction of equilibria and offer no information about the rate at which equilibrium is reached. Potential-pH diagrams tell us about the possibility of corrosion of a metal as a function of potential and pH, but not about the rate of corrosion. 

In Fig. 1.15 (bottom) curves /and m show the potential-pH relationships for the reversible hydrogen and oxygen electrodes at p 2 = Poi = f respectively. Within the area confined by the curves / and m, HjO is thermodynamically stable and Ph2 < i Poi < whereas below/and above m, Pn > 1 atm, and P02 > 1 atm, respectively see equations 1.11 and 1.12). Thus the diagram shows the solid phases of iron, the activities of metal ions and the pressures of hydrogen and oxygen gas that are at equilibrium at any given potential and pH when pure iron reacts with pure water. 

The following example illustrates how the stability or predominance diagram of water can be constructed from its basic thermodynamic information. Equation (4.33) describes the equilibrium between hydrogen ions and hydrogen gas in an aqueous environment ... 

Equation (4.33) and its alkaline or basic form, Eq. (4.34), delineate the stability of water in a reducing environment and are represented in a graphical form by the sloping line (a) on the Pourbaix diagram in Fig. 4.10. Below the equilibrium reaction shown as line (a) in this figure, the decomposition of H O into hydrogen is favored while water is thermodynamically stable above the same line (a). As potential becomes more positive or noble, water can be decomposed into its other constituent, oxygen, as illustrated in Eqs. (4.38) and (4.39) for respectively the acidic form and neutral or basic form of the same process. 

With aqueous solutions in pressurised cells, the temperature can be varied in a very broad range. Many fundamental electrochemical data have been obtained in this medium. Thermodynamic quantities such as activity coefficients of ions [252], equilibrium double-layer capacity [261], zeta potential of metals [233], potential-pH diagrams [206] or properties of the palladium-hydrogen electrode were determined [262]. Exotic systems, e.g. the solvation of rare earth atoms in liquid gallium [234], have been studied. Main research interests in subcritical aqueous solution were focused on the kinetics, reaction mechanism and transport properties. 

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