A Polarographic Method for the Indirect Determination of Polarization Curves for Oxygen Reduction on Various Metals

A theoretical and experimental analysis of the self‐discharge of lead‐acid batteries shows that seven different reactions contribute to the process. The rate of each has been determined. It is shown that positive plate self‐discharge is due primarily to a reaction between PbO2 and grid metal. The rate of this reaction decreases with increasing acid concentration because of the passivating action of the PbSO4 layers formed. The passivating action is decreased by the presence of antimony in the grid, producing pores in the PbSO4 layers and providing further access of the electrolyte to the grid. At low acid gravities antimony compounds become almost insoluble and can act as passivators decreasing the reaction between grid metal and PbO2 . Other self‐discharge reactions in positive plates are the decomposition of water by PbO2 with evolution of oxygen and the oxidation of separator material in contact with PbO2 . Oxidation of hydrogen, coming from the negative plates, occurs at immeasurably small rates.Self‐discharge of the negative plates is due to the reaction between H2SO4 and lead, producing H2 and PbSO4 . This reaction is slow in absence of foreign substances because of the high hydrogen overvoltage on lead. However, contamination of the negatives with antimony decreases the hydrogen overvoltage and greatly accelerates the self‐discharge of negative plates. Furthermore, negative plates are self‐discharged by oxygen dissolved in the electrolyte. The rate of reduction of oxygen is so fast on a lead electrode in sulfuric acid that the reaction is diffusion controlled.

Section: / Imentioning

confidence: 74%

Self-Discharge Reactions In Lead-Acid Batteries

Rüetschi

Angstadt

1958

“…Since the limiting diffusion current of oxygen reduction to water was not observed, a slower reaction step must determine the rate of production of hexavalent molybdenum. Delahay and Stagg (13), who investigated oxygen reduction on molybdenum, suggested that oxygen reduction involves a peroxyderivative of molybdenum which subsequently reduces to water and a normal oxide state of molybdenum. The formation of the peroxyderivative occurs by the reaction of peroxide with a surface oxide.…”

Section: Discussionmentioning

confidence: 99%

The Electrochemical Behavior of Molybdenum

Wikstrom

Nobe²

1969

The anodic and cathodic polarization of molybdenum in 1N H2SO4 was investigated. In deaerated solutions, two stable rest potentials, 0.368 and --0.020v vs. NHE, were observed, depending on the activation procedure. In aerated solutions, only the rest potential of 0.368v was observed. From the more noble rest potential, an anodic Tafel slope of 44 _ 4 mv and a corrosion current of 1.5 _ 0.8 ~a/cm 2 were obtained. For the h.e.r., a Tafel slope of 70 _+ 3 mv and an exchange current of 0.10 _+ 0.04 ~a/cm 2 were obtained. Two distinct active-passive transitions were observed between the two rest potentials. The short-time galvanostatic transients exhibited two linear regions in the potentialtime responses and could be represented by an electrical model of two R-C networks in series. Capacitance values were determined from the slopes of the two linear regions. The first capacitance was attributed to MoO2 at the Mo-MoO2 interface and was a constant value of 3.75 ~f/cm 2 throughout the potential region investigated. The second capacitance was surmised to be a pseudocapacitance, and ranged from 200 to 600 ~f/cm 2. However, for sufficiently large polarization from the more noble rest potential, the second capacitance was too large to be measured.Few investigations have been made of the electrochemical behavior of pure molybdenum. Some early work (1-3) indicated that molybdenum passivated easily and that the end product of its electrochemical oxidation was hexavalent molybdenum. In acidic solutions, experimental rest potentials of 0.390v vs. NHE (4), 0.368v vs. NHE (5), and 0.000-0.058 pH v vs. NHE (8) have been reported. Anodic Tafel slopes of 70 mv (7) and 45 mv (4, 5, 8) have been determined for the dissolution reaction and cathodic Tafel slopes of 76-104 mv (6, 9, 10) have been reported for the h.e.r, on molybdenum. More recently, the anodic passivation (11) and the mechanism of the dissolution reaction (8) have been studied.As indicated above, two different stable rest potentials of molybdenum in acidic solutions have been reported. The anodic region above the more noble rest potential and the cathodic region below the more active rest potential have been investigated previously. It seems appropriate that one laboratory group investigate the electrochemical behavior of molybdenum in the region between these two stable rest potentials and in the potential regions previously investigated. ExperimentalThe test cell was a conventional one and made from Pyrex glass with a nominal capacity of 1 liter. The test solution was 1N H2804. Prior to introduction of a molybdenum electrode, the test solution was either aerated or deaerated (Air Reduction Prepurified N.,, 99.997%) for 24 hr.Planar (0.005 cm thick sheet, 99.9% Mo) and cylindrical (0.825 cm diameter rod, 99.9% Mo) molybdenum electrodes (Fansteel Metallurgical Corporation) were used. Both types of electrodes had an exposed area of 1 cm 2.Subsequent to fabrication, the general electrode surface preparation procedure was as follows:1. Mechanical polishing to 4/0 emery paper...

“…It was also desirable to determine whether hydrogen peroxide could be detected. Delahay and Stagg (9) reported that little hydrogen peroxide was produced by the reduction of oxygen on zirconium in phosphate buffered 0.2M KC1, pH 6.9.…”

mentioning

confidence: 99%

The Reduction of Oxygen on Passive Zirconium

Meyer¹

1963

The reduction of oxygen on film-covered zirconium was investigated by the measurement of the rate as a function of potential and concentration. The amount of hydrogen peroxide produced was measured and compared to the theoretical amount. The results showed two distinct polarization regions, a very low-current region with a Tafel slope of about 1.4 (2.3 RT/F) and an oxygen order of unity, and a higher current region with a much larger Tafel slope and a fractional oxygen order. Equations accounting for the general features of the data are derived based on a model which considers the separate potential drops across the film and the solution double layer.Zirconium is known to form a film, presumably of ZrO2, when in contact with 02 in aqueous solution. The standard potential (1) for the half-reactionZr + 2H20 = ZrO2-b 4H + ~-4e-is--l.43v. It is evident from this potential that ZrO2 is thermodynamically stable under potentials at which such oxidizing agents as 02, H202, and H + are easily reduced. Furthermore, measurements made in 02-saturated 0.1M NaeSO4 (pH ~ 4) showed that the rate of formation of ZrO2 at potentials in the vicinity of S.C.E. falls off to values on the order of 10 -s to 10 -9 amp/cm 2 after a few days (2). Therefore it should be possible to study reduction processes on film-covered zirconium without significant interference from the film-forming reactions. It is found, in practice, that determinations of reduction rates are reasonably reproducible on a given sample, and therefore it is safe to conclude that the effect of reduction experiments on the surface of the electrode is slight.In a previous paper (3), experiments were described concerning the reduction on zirconium of several oxidizing agents including oxygen. Because the primary purpose of that paper was to present and illustrate kinetic equations resulting from consideration of a two-barrier model for film-covered electrodes, specific mechanisms were not discussed. In the present paper, additional experiments concerning the reduction of oxygen on zirconium are described, and the mechanism of this reaction is discussed. This reaction was chosen because of its importance for the film-formation process in oxygenated aqueous solutions.The reduction of oxygen on metals has not received a great deal of attention. On mercury in acid solutions, Iofa et al.