Rate of Mn

Despite the uncertainty about the rate of incorporation of °Thxs into Mn deposits, Th chronology can still provide some useful information about their growth rates. For instance, °Thxs has been used to suggest short-term changes in growth rates of Mn crusts (Eisenhauer et al. 1992) and to check the Be chronologies of crusts used to reconstruct the radiogenic isotope history of seawater (Abouchami et al. 1997 Frank 2002). 

Figure 6. An example of the use of to assess the growth rate of Mn nodules taken from Krishnaswami et al. (1982). Both panels show the same °Thxs data from nodule RN Vitiaz from the Southern Indian Oeean. Errors on the activities are within symbol size. The lower panel shows the hxs activity, while the upper panel shows the same data normalized to the Th activity. Note that both profiles show a general exponential decrease which can be used to assess the growth rate using the relationship that °Thxs ° = 230j jj imtiai g-X23ot showu ou both panels are for a steady growth rate of 1.15 mmMyr.

Eisenhauer A, Gogen K, Pemicka E, Mangini A (1992) Climatic influences on the growth rates of Mn cmsts during the Late Quaternary. Earth Planet Sci Lett 109 25-36 Elderfield H, Rickaby REM (2000) Oceanic Cd/P ratio and nutrient utilization in the glacial Southern Ocean. Natme 405 305-310... 

The rate of Mn transfer between the EXC and the ERO fractions decreased after 15 days, with only partial recovery of ERO-Mn to its initial and presaturation value. The missing Mn resided in the OM and the RO fractions, and apparently it was held in non-accessible sites, which were slow to equilibrate with the solution. This required a much longer period of time to be reoxidized. 

Experiments examining the influence of calcium and phosphate on the reductive dissolution of manganese oxides by hydroquinone have, in fact, shown inhibition by adsorbed ions (33). As the total phosphate in solution is increased, the rate of Mn + release diminished in proportion to the phosphate surface coverage. 

The rates of Mn(II) oxidation in natural waters, although slow, are typically orders of magnitude faster than the rate of oxidation of Mn(II) in solution (8,12). It has been suggested that the enhanced rate of Mn(II) oxidation in natural waters is due either to bacterial oxidation (13-16) or to the "catalytic" effects of surfaces such as metal oxides (8, 17-19). The existing evidence suggests that in certain environments bacterial mediation of the reaction is important (13-15). But in many cases the relative importance of bacterial and abiotic "catalysis" in natural waters has not been clearly defined. 

This paper discusses the oxidation of Mn(II) in the presence of lepidocrocite, y-FeOOH. This solid was chosen because earlier work (18, 26) had shown that it significantly enhanced the rate of Mn(II) oxidation. The influence of Ca2+, Mg2+, Cl", SO,2-, phosphate, silicate, salicylate, and phthalate on the kinetics of this reaction is also considered. These ions are either important constituents in natural waters or simple models for naturally occurring organics. To try to identify the factors that influence the rate of Mn(II) oxidation in natural waters the surface equilibrium and kinetic models developed using the laboratory results have been used to predict the... 

All the ions studied, except phthalate, inhibit the oxidation of Mn(II) to some degree. The relative extent to which these ions (at the concentrations indicated) affect the rate of Mn(II) oxidation is as follows ... 

As shown in Table V the rate of Mn(II) oxidation in 0.7M NaCl is about 3 times slower than in 0.7M NaC104, but the amount of Mn(II) adsorbed is only about a third less in the chloride solution. In the absence of surface complexation constants in these electrolyte matrices no model calculations can be made. 

It has been shown elsewhere (26) that in natural waters the degree of enhancement of Mn(II) oxidation predicted on the basis of model calculations is as follows y-FeOOH > a-Fe00H > silica > alumina. It has also been shown that the rate of Mn(II) oxidation is strongly influenced by pH, y-FeOOH concentration, temperature and ionic strength. Depending on the conditions, the predicted half-life 1/2 = ln 2/ki ) f°r Mn(II) oxidation may vary from a few days to thousands of years. By way of example, at pH 8, p02 0.21 atm, 25°C in waters containing 4(iM y-FeOOH and 0.2uM Mn(II), the half-life for oxidation is about 30 days. 

In freshwater, Mn(II) oxidation is slightly slower than in 0.1M NaClO. The difference between the Mn(II) oxidation rate in freshwater and 0.1M NaCIO, is greatest at pH 8.5, at this pH the rate of Mn(II) oxidation is only 40% lower in the freshwater than in 0.1M NaClO. In the estuarine-water at pH 8.5 the rate of Mn(II) oxidation is 20 times slower than in 0.1M NaCIO,. The speciation calculations indicate why the model predicts the oxidation is slower than in natural waters (see, for example Table VII). 

The importance of bacteria in mediating Mn(II) oxidation in certain environments is evident. But, the mechanisms whereby bacteria oxidize Mn(II) are poorly understood. Some bacteria synthesize proteins or other materials that enhance the rate of Mn(II) oxidation (.52). Other strains of bacteria require oxidized manganese to oxidize Mn(II) (53), suggesting that they may catalyse the oxidation of Mn(II) on the manganese oxide surface. Other bacteria may catalyse the oxidation of Mn(II) on iron oxide surfaces, as iron is associated with manganese deposits on bacteria collected in the eastern subtropical North Pacific (54). 

The rates of Mn(II) removal in some natural waters are similar to the Mn(II) oxidation rates predicted on the basis of these laboratory studies. However, in other cases the rate of manganese removal in natural waters is much faster than that expected on the basis of this work. In these systems significant manganese removal may occur as the result of adsorption, bacterially mediated oxidation, or biological uptake. 

The Co/Mn/Br now eliminates the bottleneck caused by the presence of Co(in)s. The steady state concentration of Co(III) is also much lower caused by its rapid reduction by Mn(II). This reduces carboxylic acid decomposition. We have measured the rate of Mn(III) oxidaion of bromide in the presence and absence of p-xylene and do not find any difference in rate. Hence the system also eliminates the slow Co(III) + chlorotoluene reaction. This sequence of reactions is overall faster and more selective than either the thermal or cobalt catalyzed oxidation of m-chlorotoluene. 

Intestinal absorption studies of Mn-MP were undertaken in an effort to assess the viability of the metalloporphyrin as an oral hepatobiliary agent [101, 102]. Mixed micelles of Mn-MP complexed with monoolein and taurocholate were administered to rats, resulting in liver image enhancement 68% above baseline levels six hours after administration [101]. In pigs, the mixed micelle preparation showed variable enhancement over 24 hours. Observation that Mn-MP interacts with oleic acid vesicles [103] led to investigations of the effect of oleic acid on the absorption rate of Mn-MP from the small bowel into the circulatory system [102,104]. The increase in absorption of the complex was mediated by a decrease in the relaxivity of the metalloporphyrin resulting from the interaction with the lipid vesicles. 

A similar effect of pH on dissolution rates of Mn(III/IV) oxides was observed by Stone (1987b) with substituted phenols. In this study, phenols with alkyl, alkoxy, or other electron, donating substituents were more slowly degraded. Stone (1987b) even found that p-nitrophenol, the most resistant phenol studied, reacted slowly with Mn(III/IV) oxides. 

It is interesting to compare the expected rates of Mn oxidation via abiotic mechanisms with the rates expected from the biological kinetic rate law described above. Abiotic Mn oxidation rates at pH 8.03 were measured in seawater by von Langen et al. (1997) who reported a first-order rate constant of l.lxlO-6 (normalized for Po2 = 1 atm and T = 25°C). At this pH and for similar conditions, the cell concentration of L. discophora required to obtain the same rate would be only 0.30 pg/1 (Zhang et al., 2002) (i.e., approximately 3x10s cells/1). It is reasonable to assume that cell populations of Mn-oxidizing bacteria far greater than this would be possible in natural environments. Even smaller population sizes would be required to match abiotic rates (if they could be measured) at lower pH values. 

In Fig. 1 the observed rate constant is plotted versus pH. Ill The rate of formation of Mnftpps) in the absence of cadmium(II) is independent of pH, while in the presence of cadmium(II) the rate depends on pH. Cadmium(II) accelerates the rate of Mn(tpps) formation at pH higher than 6. The following mechanism was proposed for the formation of Mnftpps) in the presence of cadmium(II). 1 ... 

The absence of any significant increase in the rate of Mn(IV) reduction under anaerobic conditions is contrary to the generally accepted view that such conditions are essential (see Alexander, 1961). It seems necessary to define the precise Eh-pH conditions at the bacterium- and Mn02(s)-water interfaces, since O2 availability is not necessarily reflected in the Eh level. 

Manganese may be soil- or foliar-applied, but soil application is the most common method for correcting Mn deficiencies. Optimum rates of Mn to be applied depend on soil pH, Mn source, soil organic matter content, and method of incorporation with soil. Soluble Mn salts quickly revert to unavailable forms when applied to Mn-deficient soils, so the method of application is important. The optimum rates reported in the literature vary widely, cotton responding to as little as 0.2 g Mn M" when side-dressed. Onions required an optimum rate of only 0.9 g of Mn m" when banded, but 6.8 gm" when broadcast (Murphy and Walsh, 1972). Foliar application is a very efficient way to correct Mn deficiency much lower rates are required than for soil application. 

Nesbitt et al. (1998) used XPS data collected on the As(III)/bimessite system to formulate a two-step mechanism for the process whereby As(ni) donates one electron to each of two surface Mn(IV) octahedra. These Mn(FV) octahedra are reduced, forming Mn(III), which is subsequently attacked by more As(in), and reduced to Mn(II). Below pH 6, Mn(II) is not appreciably sorbed by Mn(IV) hydroxides, and therefore is preferentially partitioned into solution. The rate of Mn(IV) reduction to Mn(III) is faster than the transformation of Mn(III) to Mn(II) as a result, the authors suggested that a short-lived Mn(III)-rich hydroxide phase forms on the HMO surface during the reaction (Fig. 11). 

Oxidation Rate of Mn near the Sediment-Water Interface. We used a box model approach to calculate rates of the Mn redox cycle from sediment-trap data and in situ sampling techniques (peeper and lander experiments). Figure 1 depicts an overview of the relevant processes and length scales. Sedimentation rates of particles were determined with sediment traps at 81 time intervals of about 2 weeks during 1988-1991 and at three depths (z = 20, 81, and 86 m). 

Burial Rate of Mn. Analysis of sediment cores provided a closer look at the dynamics of Mn within the sediments. Figure 5 combines results from three sediment cores taken at the deepest site of Lake Sempach. Two cores were taken in January 1991 and one in May 1988. The arrows indicate the flux-averaged Mn concentrations of settling material at depths of 20 and 81 m, respectively. The time scale in Figure 5 was calculated with 137Cs dating (41) by using Ps = 1.84 g/m2 per day as the sediment accumulation rate. 

The effects of solution composition (aside from reactant concentrations, pH, and background electrolyte concentrations) on reaction kinetics should also be studied. Using different but chemically related reactants and competitive effects from other species in solution are examples of this approach. For example. Stone and Morgan (1984) examined the effect of different organic compounds on dissolution rates of Mn oxides. 

Fig. 7-6. Variation in rate of Mn released from a Mn(III/IV) oxide with increasing concentrations of p-methyiphenol in 0.001 M acetate buffer. (Reprinted with permission from. Stone, 1987, Copyright 1987, American Chemical Society.)...

The interconversion of the various oxidation states of Mn in natural waters is influenced by UVR through its effects on reactions involving ROS [Chapter 8] and natural phenols, photoinduced charge transfer reactions, and microbial processes. The oxidation of Mn + is slow at pH < 8.5 in the absence of a catalyst. The oxidation of Mn(ii) is faster on metal oxide surfaces than in homogeneous solution in the pH range of 8 to 9 [217], and its oxidation also can be biologically mediated in the environment [153]. In comparison to bacteria-free waters, the oxidation rate of Mn(ii) in seawater is increased dramatically by catalysis on bacterial surfaces. However, even with such catalysis, its half-life still is of the order of weeks to months in open ocean waters [153]. 

In all cases the Mn(II) complexes, I, exhibit a pure first-order rate of loss of Mn(II) ion from the ligand at a fixed pH. The dissociation of the ligand does exhibit a first-order dependence in This observation allows one to describe the kinetics of ligand dissociation by its second-order dissociative rate constant, diss where multiplication of the value determined for a complex by the [H ] gives the first-order rate of Mn(II) release from the complex at any pH, and hence allows one to conveniently calculate the for the complex at any pH. To convert the reported value to a half-life at any pH, Eq. 3 applies. 

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