Surface science studies of model fuel cell electrocatalysts

The document begins with the illustration of the most widespread misunderstandings in the literature about the physical meaning of absolute electrode potential. The correct expression for this quantity is then de— rived by a thermodynamic analysis of the components of the emf of an elec— trochemical cell. It is shown that in principle three reference levels can be chosen to measure an absolute value of the electrode potential. Only one of these possesses all the requisites for a meaningful comparison on a con— mon energy scale between electrochemical and physical parameters. Such a comparison is the main problem for which the adoption of a correct scale for absolute electrode potentials is a prerequisites. The document ends with the recommendation of a critically evaluated value for the absolute potential of the standard hydrogen electrode in water and in a few other protic solvents. The \"electrode potential\" is often misinterpreted as the electric potential difference between a point in the bulk of the solid conductor and a point in the bulk of the electrolyte solution (L4) (Note a). In reality, the transfer of charged particles across the electrode/electrolyte solution interface is controlled by the difference in the energy levels of the species in the two phases (at constant T and p), which includes not only electrical (electric potential difference) but also chemical (Gibbs energy difference) contributions since the two phases are compositionally dissimilar (refs. 1,2). The value of the tjq of a \"single\" electrode, e.g. one consisting of an electronic conductor in contact with an ionic conductor, is not amenable of direct experimental determination. This is because the two metallic probes from the measuring instruments, both made of the same material, e.g. a metal M1, have to be put in contact with the bulk of these two phases to pick up the signal there. This creates two additional interfaces: a M1/solution interface, and a M1/electrode metal interface. The experimental set-up can be sketched as follows: M1 SIMIMI (1) where M is the metal of the electrode under measure, S is the electrolyte solution, M1 is the metal of the \"connections\" to the measuring instrument and the prime on M indicates that this terminal differs from the other one (M1) by the electrical state only. It is expedient to replace the M1/S interface with a more specific, reproducible and stable system known as the reference electrode. It ensues that an electrode potential can only be measured against a reference system. The measured quantity is thus a relative electrode potential. For the specific example of cell (1), the measured quantity E, the electrode potential of M relative to M1 (Note b), is conventionally split into two contributions, each pertaining to one of the electrodes: EEM_EM1 (2) EM and EM1 can be expressed in their own on a potential scale referred to another reference electrode. In this respect, the hydrogen electrode is conventionally taken as the universal Note a: This quantity, known as the Galvani potential difference between M and 5, has been defined in ref. 3. Note b: In accord with the IUPAC convention on the sign of electrode potentials, all electrode potentials in this document are to be intended as \"reduction potentials\", i.e. the electrode reaction is written in the direction of the reduction (refs. 3,4). 956 Absolute electrode potential (Recommendations 1986) 957 (for solutions in protic solvents) reference electrode for which, under standard conditions, E°(H/H2) = 0 at every temperature (Note c). Since EM as measured is a relative value, it appeals to many to know what the absolute value may be: viz. , the value of EM measured with respect to a universal reference system not including any additional metal/solution interface. Actually, for the vast majority of practical electrochenilcal problems, there is no need to bring in absolute potentials . The one outstanding example where this concept is useful is the matching of semiconductor energy levels and solution energy levels . However, from a fundamental point of view, this problem comes necessarily about in every case one wants to connect the \"relative\" electrode potential to the \"absolute\" physical quantities of the given system. On a customary basis, since the electrode potential is envisaged as the electric potential drop between M and S, the cell potential difference for system (1) is usually written as the electric potential difference between the two metallic terminals: EMi M1 (3) Since three interfaces are involved in cell (1), eqn.(3) can be rewritten as: E (M{ M) + (M S) + (S Mi) (4) Comparison of eqn. (4) with eqn. (2) shows that the identification of the absolute electrode potential with (M S) is not to be reconmended because it is conceptually misleading. Since M' and M are in electronic equilibrium, then (ref. 3): (4M ) = ('/F pr/F) (5) where the right hand side of eqn. (5) expresses the difference in chemical potential of electrons in the two electrode metals. Substitution of eqn.(5) into eqn.(4) gives: E = (p ii'/F) (E'q p'/F) (6) The two exoressions in brackets do not contain quantities pertaining to the other interfaces. They can thus be defined as single electrode potentials (Note d). Since eqn. (6) has been obtained with the two electrodes assembled into a cell, it is possible that terms common to both electrodes do not appear explicitly in eqn. (6) because they cancel out ultimately. The relationship between the truly absolute electrode potential and the single electrode potential in eqn.(6) can thus be written in the form (Note e) (ref. 5): EM(abs) = EM(r) + K (7) where K is a constant depending on the \"absolute\" reference system, and

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The absolute electrode potential: an explanatory note (Recommendations 1986)

Underpotential deposition at single crystal surfaces of Au, Pt, Ag and other materials

Mechanism and electrocatalysis in the direct methanol fuel cell

The division of thermodynamic solvation free energies of electrolytes into contributions from individual ionic constituents is conventionally accomplished by using the single-ion solvation free energy of one reference ion, conventionally the proton, to set the single-ion scales. Thus, the determination of the free energy of solvation of the proton in various solvents is a fundamental issue of central importance in solution chemistry. In the present article, relative solvation free energies of ions and ion−solvent clusters in methanol, acetonitrile, and dimethyl sulfoxide (DMSO) have been determined using a combination of experimental and theoretical gas-phase free energies of formation, solution-phase reduction potentials and acid dissociation constants, and gas-phase clustering free energies. Applying the cluster pair approximation to differences between these relative solvation free energies leads to values of −263.5, −260.2, and −273.3 kcal/mol for the absolute solvation free energy of the proton in me...

/pdf/single-ion-solvation-free-energies-and-the-normal-hydrogen-26dyitnfeg.pdf

Single-Ion Solvation Free Energies and the Normal Hydrogen Electrode Potential in Methanol, Acetonitrile, and Dimethyl Sulfoxide

The effect of adsorbing submonolayers of bismuth and tin onto single-crystal, polycrystalline and dispersed platinum electrodes on the oxidation of formic acid and methanol was investigated. While the oxidation of formic acid on Pt(111) was found to be enhanced by adsorbed sub-monolayers of bismuth, the oxidation of methanol was inhibited by all adsorbed coverages of both tin and bismuth on all electrodes. The adsorbed tin was found to be more stable on the dispersed Pt electrodes than on the smooth electrodes.

Effect of Bi and Sn adatoms on formic acid and methanol oxidation at well defined platinum surfaces

The relation of the “medium effect” or “solvent activity coefficient” to the directly measurable real free energy of transfer is discussed. Measurements of the latter quantity for solutions in aqueous solvents containing methanol, ethanol, dioxan, acetone, ethylene glycol, formamide or dimethyl sulphoxide are reported and discussed. The surface composition of the solvent mixture is accessible and this information can be used to set reasonable limits on the possible value of the “medium effect”. Thus it can be shown that some assignments of single ion medium effects based on over-simplified models of solvation are inadequate. The assignment based on the neglect of liquid junction potentials between solvents is often surprisingly good.

The medium effect for single ionic species

The influence of adsorbed toluene-p-sulphonate ions on the rate of electroreduction of periodate ions at mercury has been studied as a function of electrode potential and bulk concentration of the adsorbing ions in solutions of constant ionic strength. The ratio of the reduction rate in the presence of adsorbed anions to that in their absence at constant potential was found to be a function of the surface concentration of the adsorbed species independent of electrode potential, bulk concentration of the adsorbing species and their orientation on the electrode surface. An analysis of the decelerating effect of the adsorbed ions in terms of an electrostatic model of the interface which considers discreteness-of-charge effects is presented. The results of the analysis indicate that the reacting anion pairs with non-adsorbed cations in the vicinity of the reaction site before entering the activated state.

Roger Parsons

Papers

Effect of Bi and Sn adatoms on formic acid and methanol oxidation at well defined platinum surfaces

The medium effect for single ionic species

Effects of adsorbed ions on simple electrode reactions. Part 1.—Reduction of periodate ion in the presence of toluene-p-sulphonate anions