Anodic oxidation of lead in aqueous carbonate solutions. II. Film formation and dissolution in the pH range 9 to 14

The anodic oxidation of lead has been studied in aqueous carbonate solutions at pH values in the range 9 ≤ pH ≤ 14. The solid films formed on the electrode surface have been identified by X-ray diffractometry. The dissolution and film formation processes have been studied by a number of electrochemical techniques at rotating disk electrodes. The anodic oxidation process can be divided into two distinct regions. In the main anodic oxidation process, a variety of surface phases is formed. The nature of the phase formed is pH-dependent. Plumbonacrite (Pb10O(OH)6(CO3)6) formation is observed at all pH values and predominates at pH ≥ 13. At pH < 11, cerussite (PbCO3) is the predominant phase, whereas at intermediate pH values, hydrocerussite (Pb3(OH)2(CO3)2) predominates. The dissolution rate of Pb2+ species from the electrode surface is directly proportional to the solubility of the predominant phase present. At more positive potentials, a reactivation, involving increased dissolution and a further stage of film formation, is observed. Litharge (PbO) is observed to grow underneath the initially formed basic lead carbonates. Dissolution occurs either by the field-assisted dissolution of the base-layer on the electrode or by metal dissolution through faults in the base-layer.

Anodic oxidation of lead in aqueous carbonate solutions. I. Film formation and dissolution at pH = 12

The anodic oxidation of lead has been studied using rotating disk electrodes, under voltammetric and potentiostatic conditions, in aqueous carbonate solutions at pH = 12. The solid films formed on the electrode surface have been identified by X-ray diffractometry. Two distinct surface layers are formed: a compact layer composed of two separate phases, plumbonacrite (Pb10O(OH)6(CO3)6) and hydrocerussite (Pb3(OH)2(CO3)2); and a more dispersed layer of hydrocerussite. The plumbonacrite in the compact layer is formed directly on the electrode surface by a solid-state growth process. The hydrocerussite component of this compact layer appears to form by conversion of the outer layers of plumbonacrite. The dispersed layer is formed by deposition from solution. The extent of deposition is controlled by the rate of transport of soluble Pb2+ species to the bulk of solution.

Anodic oxidation of UO2. V. Electrochemical and X-ray photoelectron spectroscopic studies of film-growth and dissolution in phosphate-containing solutions

The anodic oxidation of UO2 has been studied in aqueous phosphate solutions over the pH range 4 to 11, using a combination of electrochemical and X-ray photoelectron spectroscopic techniques. The early stages of oxidation, leading to the formation of a film of composition UO2.33, are unaffected by the presence of phosphate in the solutions. Phosphate concentrations [Formula: see text]prevent the formation of higher-oxide films that are present in phosphate-free solutions at higher oxidation potentials. Dissolution under steady-state conditions proceeds via a surface film of uranyl phosphate. For potentials [Formula: see text] (vs. SCE), the rate of dissolution to yield[Formula: see text] is controlled by charge-transfer kinetics. For potentials[Formula: see text], the rate-determining step is the chemical dissolution of the uranyl phosphate layer.

Kinetics of Dehydration of Aromatic Aldehydes Determined by the Rotating Disc Electrode Technique

The rate constant for the irreversible dehydration of a number of aromatic aldehyde hydrates has been measured. The hydrates were generated by constant potential electrolysis of the respective aromatic acid at a rotating disc electrode. The aldehydes are then further reduced to the alcohol. The rate constant can be calculated from the ratio of the concentration of intermediate aldehyde to the total concentration of products.The theory was further extended to include the reverse hydration of the aldehyde to the hydrate and applied in the case of m-chlorobenzaldehyde. A value for the equilibrium constant of K = 1.30 ± 0.25 was obtained for this reaction.