Dissolution rate of oxides

Colmer and Hinkle (14) identified T. ferrooxidans in acidic mine waters. Subsequent studies by Silverman et al. (15,16) confirmed that T. ferrooxidans could be utilized to oxidize FeSo in coal in 3 to 4 days, and the rate of oxidative dissolution was a function of the particle size and rank of the coal. Dugan and Apel (4,5) showed that a mixed culture of T. ferrooxidans and T. thiooxidans was most effective at a pH of 2 to 2.5 when the nutrient was enriched with NH " -. They reported 97% removal of pyritic sulfur from a coal sample with 3.1 weight percent sulfur. Norris and Kelly (17) reported that other acidophilic bacteria, Leptospirillum ferrooxidans in mixed cultures with T. thiooxidans, was effective for FeS2 removal. ... 

Sand et al. (2001) suggested that Fe(III) ions or protons (H ) are the only chemical agents that dissolve a metal sulfide, and that bacteria have functions to regenerate these ions and to concentrate them at the mineral/water or the mineral/ bacterial cell interface. The chemical reactions require the presence of surrounding bacterial cells. The concentration of reactants in this nano-meter-thick layer causes the observed acceleration of rates of oxidative dissolution of sulfide minerals. 

The chemical oxidation of metal sulfides is controlled in part by the dissolution of sulfide minerals under acidic conditions and by the presence of oxidants (DO, Fe ) that lead to the disruption of sulfide chemical bonds. Bacteria can have a significant effect on the rate of oxidative dissolution of sulfide minerals by controlling mineral solubility and surface reactivity. Metal-enriched waters and solutions rich in sulfuric acid that form in association with mining can be directly linked to microbial activity. The majority of studies to date have focused on the reactivity and kinetics of sulfide minerals in the presence of A. ferrooxidans and L. ferrooxidans, and in some cases A. thiooxidans (Singer and Stumm, 1970 Tributsch and Bennett, 1981a,b Sand et al., 1992, 2001 Nordstrom and Southam, 1997 Sasaki et al., 1998 Edwards et al., 1998, 1999, 2000 Nordstrom and Alpers, 1999a Banfield and Welch, 2000 Tributsch, 2001). Additional studies have been conducted on other species of bacteria and archea (Edwards et al, 1998, 1999, 2000). 

These equations indicate that the base iron fimctions as a reducer to accelerate the dissolution of iron oxides. Because it is difficult to determine the endpoint for the dissolution of fouling oxides, an inhibitor is generally added for safety purpose. Both anodic and cathodic inhibitors could be added to retard the corrosion of the bare metal after dissolution of the fouling oxides. Figures 10.1 and 10.2 illustrate the action that could be played by either an anodic inhibitor (Fig. 10.1) or a cathodic inhibitor (Fig. 10.2). It can be seen that although the anodic inhibitor retards the anodic dissolution of iron at the endpoint, it concurrently decreases the rate of oxide dissolution permitted by the chemical system. 

On the other hand, pit initiation which is the necessary precursor to propagation, is less well understood but is probably far more dependent on metallurgical structure. A detailed discussion of pit initiation is beyond the scope of this section. The two most widely accepted models are, however, as follows. Heine, etal. suggest that pit initiation on aluminium alloys occurs when chloride ions penetrate the passive oxide film by diffusion via lattice defects. McBee and Kruger indicate that this mechanism may also be applicable to pit initiation on iron. On the other hand, Evans has suggested that a pit initiates at a point on the surface where the rate of metal dissolution is momentarily high, with the result that more aggressive anions... 

Some metals and alloys have low rates of film dissolution (low /p) even in solutions of very low pH, e.g. chromium and its alloys, and titanium. In these cases the value of /p is quite low, and although it increases as the temperature increases, a maximum is reached when the solution boils. The maximum current is below and breakdown does not occur. However, in certain alloys, e.g. Cr-Fe alloys, the protective film may change in composition on increasing the anode potential to give oxides that are more soluble at low pH and are therefore more susceptible to temperature increases. This occurs in the presence of cathode reactants such as chromic acid which allow polarisation of the anode. 

Duncan and Frankenthal report on the effect of pH on the corrosion rate of gold in sulphate solutions in terms of the polarization curves. It was found that the rate of anodic dissolution is independent of pH in such solutions and that the rate controlling mechanism for anodic film formation and oxygen evolution are the same. For the open circuit behaviour of ferric oxide films on a gold substrate in sodium chloride solutions containing low iron concentration it is found that the film oxide is readily transformed to a lower oxidation state with a Fe /Fe ratio corresponding to that of magnetite . 

Nevertheless, Ta5+ and Nb5+ interact with aqueous media containing fluorine ions, such as solutions of hydrofluoric acid. On the other hand, as was clearly shown by Majima et al. [448 - 450], the increased hydrogen ion activity can also significantly enhance the dissolution rate of oxides. The activity of hydrogen ions can be increased by the addition of mineral salts or mineral acids to the solution. 

Andersen et al. predicted that similar results would be expected for the corrosion of other multivalent metals oxidizing via lower oxidation states. They also pointed out that their interpretation was consistent with the kinetics of the corrosion of copper in oxygenated HCl solutions. Here the final product is Cu and thus there is no vulnerable intermediate. In consequence, the rate of copper dissolution from either Nj-saturated or 02-saturated HCl solutions was the same at a given potential in conformity with the additivity principle. 

Under the effect of oxidizing agents, a metal may become passivated even when not anodically polarized by an external power source. In this case, passivation is evident from the drastic decrease in the rate of spontaneous dissolution of the metal in the solution. The best known example is that of iron passivation in concentrated nitric acid, which had been described by M. V. Lomonosov as early as 1750. Passivation of the metal comes about under the effect of the oxidizing agent s positive redox potentiaf. 

Sometimes anodic protection is used, in which case the metal s potential is made more positive. The rate of spontaneous dissolution will strongly decrease, rather than increase, when the metal s passivation potential is attained under these conditions. To make the potential more positive, one must only accelerate a coupled cathodic reaction, which can be done by adding to the solution oxidizing agents readily undergoing cathodic reduction (e.g., chromate ions). The rate of cathodic hydrogen evolution can also be accelerated when minute amounts of platinum metals, which have a stroug catalytic effect, are iucorporated iuto the metaf s surface fayer (Tomashov, 1955). 

Such reactions processes are responsible for the transition from PS formation to electropolishing with increasing potential as typically revealed in an i-V curve.18 PS formation can occur when the surface is not or only partially covered by oxide. Once the whole surface is covered with an oxide film further reaction can only proceed through the formation of oxide followed by its dissolution. Further increasing the potential will only result in an increase of oxide film thickness. On the other hand, increasing HF concentration will increase the dissolution rate of oxide. The presence of oxide on the silicon surface in the PS formation region and its increase with potential have been experimentally observed.98... 

Formation rate of oxide film relative to its dissolution rate... 

Rate of direct dissolution of silicon relative to indirect dissolution via oxide formation and dissolution... 

The Rate of reductive Dissolution of Hematite by H2S as observed between pH 4 and 7 is given in Fig. 9.6 (dos Santos Afonso and Stumm, in preparation). The HS" is oxidized to SO. The experiments were carried out at different pH values (pH-stat) and using constant PH2s- 1.8 - 2.0 H+ ions are consumed per Fe(II) released into solution, as long as the solubility product of FeS is not exceeded, the product of the reaction is Fe2+. The reaction proceeds through the formation of inner-sphere =Fe-S. The dissolution rate, R, is given by... 

Rates of reductive dissolution of amorphous manganese (111,1V) oxide particles decrease as the electrode half-wave potentials of the substituted phenols (as reported by Suatoni et al., 1961) increase (4.8 x 10 5 M total manganese, pH 4.4). 

Pyrite Oxidation. The oxidation of Fe(ll) minerals by Fe3+ is also of importance in the oxidation of pyrite by 02. This process is mediated by the Fe(II)-Fe(III)system. Pyrite is oxidized by Fe3+ (which forms a surface complex with the pyrite (cf. formula VI in Fig. 9.1) (Luther, 1990). The rate determining step at the relatively low pH values encountered under conditions of pyrite dissolution is the oxygenation of Fe(II) to Fe(III) usually catalyzed by autotrophic bacteria (Singer and Stumm, 1970 Stumm-Zollinger, 1972). Thus, the overall rate of pyrite dissolution is insensitive to the mineral surface area concentration. Microbially catalyzed oxidation of Fe(II) to Fe(III) by oxygen could also be of some significance for oxidative silicate dissolution in certain acid environments. 

Dissolved iron(III) is (i) an intermediate of the oxidative hydrolysis of Fe(II), and (ii) results from the thermal non-reductive dissolution of iron(III)(hydr)oxides, a reaction that is catalyzed by iron(II) as discussed in Chapter 9. Hence, iron(II) formation in the photic zone may occur as an autocatalytic process (see Chapter 10.4). This is also true for the oxidation of iron(II). As has been discussed in Chapter 9.4, the oxidation of iron(II) by oxygen is greatly enhanced if the ferrous iron is adsorbed at a mineral (or biological) surface. Since mineral surfaces are formed via the oxidative hydrolysis of Fe(II), this reaction proceeds as an autocatalytic process (Sung and Morgan, 1980). Both the rate of photochemical iron(II) formation and the rate of oxidation of iron(II) are strongly pH-dependent the latter increases with... 

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