Hydrogen-ion reduction curve

Fig. 5.17 Net polarization curves, N, resulting from the metal anodic curve, M, and the sum cathodic curve, SC, for the oxygen-reduction and hydrogen-ion-reduction curves. Curves M and SC are from Fig. 5.16. pH = 1. PQ = 0.05 atm...

Fig. 6.4 Schematic experimental polarization curves (solid curves) assuming active-passive behavior for the individual metal-oxidation curve and Tafel behavior plus limiting diffusion for the individual dissolved-ox-ygen and hydrogen-ion reduction curves (dashed curves)...

Fie. 3.18 Illustration of the effect of exchange current density on the polarization curve for oxygen reduction in aerated environments of pH = 0.56 and Pq2 = °-2 atm- Curves converge to the same diffusion limit and are identical when the hydrogen ion reduction is the dominant reaction. 

Fig. 3.19 Cathodic polarization curves for 100 and 10,000 ppm Fe3+ (as FeCI3) on platinum in nitrogen-deaerated solution. The increase in current density at 400 mV (SHE) is due to a velocity effect in introducing nitrogen sparging into the solution. The limiting current density is increased by a factor of about 100 on increasing the concentration from 100 to 10,000 ppm. The increase in current density near-100 mV (SHE) is due to hydrogen ion reduction resulting from a decrease in pH dueto Fe3+ hydrolysis.

Curves for hydrogen-ion reduction are based on experimental values of the polarization parameters governing the polarization curves. The anodic polarization curves for iron show a dependence on pH due to the influence of hydrogen ion concentration on the kinetic steps in the iron oxidation. Based on Ref 7 and 8... 

All of the curves in Fig. 5.6 start in the active dissolution potential range and hence do not show the complete polarization curve for the iron extending to the equilibrium half-cell potential as was done in Fig. 5. 4. This extension was shown as dashed lines and the equilibrium potential was taken as -620 mV for Fe2+ = 10 6. Qualitatively, the basis for estimating how the active regions of the curves in Fig. 5.6 would be extrapolated to the equilibrium potential can be seen by reference to Fig. 4.16. There, the corrosion potential is represented as the intersection of the anodic Tafel curve and the cathodic polarization curve for hydrogen-ion reduction at several pH values. It is pointed out that careful measurements have shown that the anodic Tafel line shifts with pH (Ref 6), this shift being attributed to an effect of the hydrogen ion on the intermediate steps of the iron dissolution. 

Fig. 5.19 Potentiostatic polarization curve for pure chromium in hydrogen-saturated (deaerated) 1 N H2S04at25 °C. Dashed section is a cathodic "peak" where the hydrogen-ion reduction dominates over the passive chromium oxidation. Redrawn from Ref 9...

In deaerated 1 N H2SO4 (pH = 0.56), hydrogen-ion reduction is the cathodic reaction with the cathodic polarization curve intersecting the iron, nickel, and chromium curves in the active potential region. Hence, active corrosion occurs with hydrogen evolution, and the corrosion rates would be estimated by the intersections of the curves. The curves predict that the titanium will be passivated. However, the position ofthe cathodic hydrogen curve relative to the anodic curves for titanium and chromium indicates that if the exchange current density for the hydro-... 

If it is assumed that the hydrogen and catalytically active ions are reduced independently, and each accordii to its own rules [93], it follows from Fig. 17 that change in catalsdic current density (lower part of Curves 2-4) leads to a considerably larger shift in the potential of the electrode than hydrogen ion reduction, and with increase in potential the catalytic wave is therefore simply overlapped by the normal hydrogen-ion reduction current. 

In Sections 10.11-10.16 it is shown how the change in pH during acid-base titrations may be calculated, and how the titration curves thus obtained can be used (a) to ascertain the most suitable indicator to be used in a given titration, and (b) to determine the titration error. Similar procedures may be carried out for oxidation-reduction titrations. Consider first a simple case which involves only change in ionic charge, and is theoretically independent of the hydrogen-ion concentration. A suitable example, for purposes of illustration, is the titration of 100 mL of 0.1M iron(II) with 0.1M cerium(IV) in the presence of dilute sulphuric acid ... 

Figure 11-7 shows the polarization curve of an iron electrode in an acidic solution in which the anodic reaction is the anodic transfer of iron ions for metal dissolution (Tafel slope 40 mV/decade) the cathodic reaction is the cathodic transfer of electrons for reduction of hydrogen ions (Tafel slope 120 mV /decade) across the interface of iron electrode. 

AgCl electrode) and the cathodic current due to the reduction of hydrogen ion begins to flow at about -1.1 V. Between the two potential limits, only a small current (residual current) flows. In curve 2, there is an S-shaped step due to the reduction of Cd2+, i.e. Cd2++2e +Hg <=t Cd(Hg). In DC polarography, the current-potential curve for the electrode reaction is usually S-shaped and is called a polaro-graphic wave. 

At first, no current flows through the cell until a decomposition potential Ed- is reached. At that point, the current begins to flow (Curve A). We also observe that gas bubbles are formed at the working electrode, and that current fluctuates somewhat randomly. Two chemical processes are taking place at the electrodes. At the W electrode (which is now the cathode), electrons are transferred from the electrode to the hydrogen ions, H+. Thus, the reduction takes place at the cathode, according to the following electrochemical reaction. 

Minor differences between the three electrolyte solutions are also observed. First, electrolyte number 3 only shows a peak maximum in the current-potential curves at potentials higher than 8 V. However, this is very clear because its pH value is smaller, indicating that this electrolyte solution possesses a higher buffer capacity against consumption of hydrogen ions in the vicinity of the fibre surface, avoiding hydrogen gas formation and Ni(OH)2 precipitation. Secondly, at a potential of 4V, no deposition occurred in electrolyte solution number 3, indicated by the absence of an increase in the measured electrical current and confirmed by XPS data. Additionally in this case, the lower pH plays an important role because of the lower pH value, the applied potential difference does not overlap with the potential window in which the reduction of Ni(II) occurs. Therefore no deposition is observed. 

The shape of the reduction curve indicates that the reduction of cupric ion is autocatalytic. This conclusion is confirmed by the almost complete disappearance of the induction period when cuprous acetate is added, strikingly illustrated by a comparison of curves I and II in Fig. 4 (5). In both cases the quantity of hydrogen absorbed is close to the theoretical value for the reduction of Cu++ to Cu+. There is evidence that the presence of aniline in the quinoline is effective in providing some cuprous ion from cupric to initiate sufficient catalyst to start the reaction. 

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