Platinum-water potential

Presently, the simulation of real metal surfaces with defects of all sorts is hampered by the restriction of system size and also because only little quantitative information is available concerning the specific electronic structure near these defects and its consequences for the binding of water molecules and ions. However, first steps into this direction were made recently Siepmann and Sprik [51], who investigated the influence of surface topology on water/electrode systems, and by Nagy and Denuault [98, 99], who extended the platinum-water potential to defect surfaces. However, since they did not take into account the electronic structure of the metal, their approach is purely geometrical. 

Fig. 9. Potential versus pH diagram for the system platinum/water at 25°C (after Lee [109]).

At a platinum electrode, potentials of about 3 V (versus Ag/AgCl) can be reached, and the limiting reaction is the discharge of the supporting electrolyte. The useful range in cathodic direction is very dependent on the water content. 

Figure 7.2. Measurement of the platinum electrode potential in relation to the electrode current for the Fe +-Fe redox pair at pH 2 in water, under conditions of (a) [Fe ] = [Fe ] = 10 M, and (b) [Fe ] = [Fe ] = 10 M. (Adapted from W. Stumm and J. J. Morgan. 1981. Aquatic Chemistry. 2nd ed. New York Wiley. Used with permission.)...

V. These values were interpreted in terms of potential-pH diagrams for the platinum/water systems. Such an interpretation is invalid in principle in that the potential-pH diagrams formally apply to equilibrium conditions where irreversible electrolysis events (including the electrode oxidation) are not involved, in contrast to the present case of electrochemical reactions on the electrodes driven by the passage of a current. 

In 2007 and 2008, Sugino et al. [201] and Otani et al. [202] investigated biased platinum-water interfaces. Sugino et al. found that an orientation of the water molecules emerged due to the negative bias potential of the water-Pt(l 11) interface... 

Thus, the data obtained with the EQCM on Pt are consistent with either reaction A or C shown in Table 1. In order to be able to distinguish between these two processes, one needs to know the orientation of water molecules on platinum at potentials where surface oxide formation starts E 0.75 V vs. SHE). These potentials are positive with respect to the pzc, which is in the range of 0.15-0.30 vs. SHE in acid solutions [134,135] and it is likely that water molecules are oriented with oxygen toward the metal. Thus, the most likely process of electrochemical oxidation of a Pt surface is reaction C in Table 1, which produces Pt(OH)2. This is a prime example showing how the results obtained by the EQCM can complement those derived from classical electrochemical experiments to gain information regarding the processes taking place at the metal/solution interface. 

In the case of the Pt(lOO) surface the interaction potential is derived from semiempirical quantum chemical calculations of the interactions of a water molecule with a 5-atom platinum cluster [35]. The lattice of metal atoms is flexible and the atoms can perform oscillatory motions described by a single force constant taken from lattice dynamics studies of the pure platinum metal. The water-platinum interaction potential does not only depend on the distance between two particles but also on the projection of this distance onto the surface plane, thus leading to the desired property of water adsorption with the oxygen atoms on top of a surface atom. For more details see the original references [1,2]. This model has later been simplifled and adapted to the Pt(lll) surface by Berkowitz and coworkers [3,4] who used a simple corrugation function instead of atom-atom pair potentials. 

Figure 23.15. Potential-pH diagram for the platinum-water system at 25 °C...

Titanium has potential use in desalination plants for converting sea water into fresh water. The metal has excellent resistance to sea water and is used for propeller shafts, rigging, and other parts of ships exposed to salt water. A titanium anode coated with platinum has been used to provide cathodic protection from corrosion by salt water. 

The standard potential for the anodic reaction is 1.19 V, close to that of 1.228 V for water oxidation. In order to minimize the oxygen production from water oxidation, the cell is operated at a high potential that requires either platinum-coated or lead dioxide anodes. Various mechanisms have been proposed for the formation of perchlorates at the anode, including the discharge of chlorate ion to chlorate radical (87—89), the formation of active oxygen and subsequent formation of perchlorate (90), and the mass-transfer-controUed reaction of chlorate with adsorbed oxygen at the anode (91—93). Sodium dichromate is added to the electrolyte ia platinum anode cells to inhibit the reduction of perchlorates at the cathode. Sodium fluoride is used in the lead dioxide anode cells to improve current efficiency. 

Cathodic protection applications in fresh water include use of ferrite-coated niobium , and the more usual platinum-coated niobium . Platinised niobium anodes have been used in seawater, underground and in deep wells " and niobium connectors have been used for joining current leads Excellent service has been reported in open-seawater, where anodic potentials of up to 120V are not deleterious, but crevice corrosion can occur at 20 to 40V due to local surface damage, impurities such as copper and iron, and under deposits or in mud ... 

By virtue of the high breakdown potential of the oxide film (approximately 155 V in sea water and 280 V in low conductivity water of pH = 7) tantalum has found use as a substrate for platinum in impressed-current cathodic-protection anodes, which can be used at high impressed voltages (50 V) and high current densities. However, because of its lower cost, niobium is preferred for systems that have to operate at high voltages... 

Prepare 250 mL of 0.02 M potassium dichromate solution and an equal volume of ca 0.1 M ammonium iron(II) sulphate solution the latter must contain sufficient dilute sulphuric acid to produce a clear solution, and the exact weight of ammonium iron(II) sulphate employed should be noted. Place 25 mL of the ammonium iron(II) sulphate solution in the beaker, add 25 mL of ca 2.5M sulphuric acid and 50 mL of water. Charge the burette with the 0.02 M potassium dichromate solution, and add a capillary extension tube. Use a bright platinum electrode as indicator electrode and an S.C.E. reference electrode. Set the stirrer in motion. Proceed with the titration as directed in Experiment 1. After each addition of the dichromate solution measure the e.m.f. of the cell. Determine the end point (1) from the potential-volume curve and (2) by the derivative method. Calculate the molarity of the ammonium iron(II) sulphate solution, and compare this with the value calculated from the actual weight of solid employed in preparing the solution. 

Two platinum electrodes are immersed in sulphuric acid of suitable concentration containing the nitrate ion to be determined and a potential of about 100 millivolts is applied. Upon titration with 0.4M ammonium iron(II) sulphate solution there is an initial rise in current followed by a gradual fall, with a marked increase at the end point the latter is easily determined from a plot of current against volume of iron solution added. The concentration of water should not be allowed to rise above 25 per cent (w/w). The temperature of the solution should not exceed 40 °C. 

Therefore, the following method was suggested and realized (the scheme is shown in Fig. 17). A 1.5 M solution of KCl or NaCl (the effect of preventing BR solubility of these salts is practically the same) was used as a subphase. A platinum electrode was placed in the subphase. A flat metal electrode, with an area of about 70% of the open barriered area, was placed about 1.5-2 mm above the subphase surface. A positive potential of +50 -60 V was applied to this electrode with respect to the platinum one. Then BR solution was injected with a syringe into the water subphase in dark conditions. The system was left in the same conditions for electric field-induced self-assembly of the membrane fragments for 1 hour. After this, the monolayer was compressed to 25 mN/m surface pressure and transferred onto the substrate (porous membrane). The residual salt was washed with water. The water was removed with a nitrogen jet. 

Glassy carbon electrodes polished with alumina and sonicated under clean conditions show activation for the ferrl-/ ferro-cyanlde couple and the oxidation of ascorbic acid. Heterogeneous rate constants for the ferrl-/ ferro-cyanlde couple are dependent on the quality of the water used to prepare the electrolyte solutions. For the highest purity solutions, the rate constants approach those measured on platinum. The linear scan voltammetrlc peak potential for ascorbic acid shifts 390 mV when electrodes are activated. 

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