This study investigated the kinetics of the homogeneous oxidation of Mn(II) and the formation and disappearance of Mn(III) complexes. The rate of disappearance of the Mn(III)EDT A was partially dependent on pH, likely indicative of an unknown pH-dependent intermediate in the decomposition of the complex.
Introduction
Mn(ll)
Mn(IV)
This, too, was found to act via the oxidation of Mn(II) to Mn(III) by H2O2 and the stabilization of Mn(III) by organic ligands (41-43). A wide variety of Mn(III) complexes with aliphatic alcohols, polyalcohols, and carboxylic acids have been studied electrochemically, giving the reduction potential for each complex (61-62).
Redox Transitions
This study examined some unanswered questions regarding Mn(II) and Mn(III) redox transitions. One is the change in the rate constant for the disappearance of the Mn(III)Pi()7 species, and the other is the initial amount of Mn(III) formed. The concentration of the Mn complex can be expressed as total manganese.
One is that the reaction of permanganate with Mn(II) is slow, but eventually all the permanganate reacts to form Mn(III). The calculation was made for the disproportionation of the Mn(IIl)EDTA complex with the Mn(II)EDTA complex and MnO2. This study has examined the dissolution of an oxide containing Mn(III) in the presence of the ligands discussed in Chapter 4, viz.
In chapter 4 this ratio was insufficient for the stabilization of Mn(III)P2O7 complexes synthesized from MnO4- and Mn(II). Plot of average oxidation state of solids versus time for dissolution of MnOOH solids with P'i-)7. Plot of average oxidation state of solids versus time for dissolution of MnOOH solids by citrate.
Since such areas are usually anoxic, the source of Mn(l11) should come from Mn(III) oxides. This is indeed possible if the settling time of Mn(III) particles is faster than the reduction time or the oxidation time to Mn(IV).
Scope of This Study
Homogenous Oxidation
Mn(lll) Complexes
EXPERIMENTAL METHODS
Introduction
- Total Mn
- Mn(II)
- Oxidized Mn
A more specific approach for the detection of Mn(III) and Mn(IV) is provided by redox-based techniques. Therefore, the total concentration obtained with either of these techniques is equal to the concentration of Mn(III) plus twice the concentration of Mn(IV).
Methods Used
- Total Mn
- Oxidized Mn
- Mn(III)
- pH of Zero Point of Charge
Most observations; however, can be explained by a cycle in which the Mn(III)CIT complex undergoes an electron transfer in the inner sphere, producing Mn(II), which is then reoxidized to Mn(III) in the presence of oxygen.
Figures
HIGH TEMPERATURE Mn(II) AUTOXIDATION
Previous Studies
Plot of the average oxidation state of the solids versus time for dissolution of MnOOH solids with Pz()7.
Current Study
- Solubility Considerations
- Experimental Method
Results
- Extrapolation to Room Temperature
- Comparison and Discussion of Results
Tables
Figures
Mn(III) SOLUTION COMPLEXES
Introduction
- Pyrophosphate
- Citrate
- Ethylenediaminetetracetate (EDTA)
- This Study
Interestingly, at 1 g/l there was an initial drop in the oxidation state of the solids. There appears to be a slight drop in the average oxidation state of the solids from 3.0 to about 2.8.
Experimental
- Complex Preparation
- Experimental Monitoring
Results and Discussion
- Pyrophosphate
- Effect of pH
- Effect of Ligand Concentration
- EDT A
- Effect of pH
- Effect of Ligand Concentration
- Effect of Light and Ionic Strength
- Citrate
- Effect of pH
- Effect of Oxygen on Reformation of the Mn(l11) Complex
- Effect of Ligand Concentration
Conclusions
- Pyrophosphate
- EDT A
- Citrate
Mechanisms
Tables
Figures
MnOOH DISSOLUTION BY LIGANDS
- Mn(III) Solid Phases
- MnOOH Reactions
- This Study
- Solids Preparation
- Solids Characterization
- Experimental Procedures
Results and Discussion
- Pyrophosphate
- Effect of pH
- Effect of Solids Concentration
- Effect of Ligand
- Oxidation State
- EDTA
- Citrate
It does not seem unreasonable to assume that this is sufficient to achieve surface saturation. Further expressing Equation 5.6 in terms of the acid-base equilibria of the surface sites and the pyrophosphate in the pH range 6.2 to 8.4 would yield. No reduction in oxidation state occurs, so reductive dissolution is unlikely to occur, a conclusion corroborated by.
The dependence turns out to be to the approximately 1.4 power, if it is assumed that the formula weights of the solids in both experiments were 100 g/mol. This result shows that the rate constant for complete or nearly complete dissolution is independent of the concentration of solids, while the initial rate is dependent on the concentration of solids. Again, there appears to be a slight upward trend in the oxidation state, but this is difficult. distinguish it from random scattering.
This is not entirely surprising since the reduction of the Mn(III)EDT A complexes was found to be quite rapid in solution. Another interesting observation is that the pH 7.8 citrate dissolution experiment is the only one in which a significant portion of the solids remained undissolved. The dissolution of the solid is almost complete in this run, unlike the result at pH 7.8.
Conclusions
- Pyrophosphate
- EDTA
The complex of Mn(l11) with the P2074- species is the longest-lived of the complexes studied, with first-order constants on the order of 10-7 s-1. All dissolved MnOOH was found to remain in solution as Mn( III ) species P2O7. To get an idea of the quantitative concentrations and fluxes of Mn(III) possible a simple box model was constructed.
Comparing the results of the run with Mn(III)L and the one without shows a significant difference. The area affected by such a cycle will depend on the time scale of loss of the Mn(III) complex. Most of the Mn in desert varnish is in the +4 oxidation state, but there is evidence of some Mn below the +4 oxidation state (14).
For EDT A, the main challenge lies in the mechanism of reaction and identification of the products. Detection of degradation products will provide insight into the mechanism of the reduction reaction. For all three ligands investigated, more experiments can be done to give a better idea of the rate law.
Another very interesting question is the possibility of the reaction marked by the dashed line in Figure 6.1 b. Since MnO2 is the most stable form and one of the most common forms of Mn, it would be important if Mn(III) complexes could be formed from them.
Mechanisms
Tables
Figures
SUMMARY AND CONCLUSIONS
Important Findings of This Study
- Mn Oxidation
- Mn(III)
- MnOOH Dissolution
Rate expressions and rate constants for the disappearance of Mn(III) complexes are given in Table 6.1 for the three ligands studied. Although the disappearance rate of Mn(Ill)P207 species was observed to be dependent on both pH and total P20f concentration, the data are too limited to establish a hard law on the rate. The dependence of the loss rate of (P2074-1 appears complex, possibly the result of different Mn(IIl)P207 complexes with different susceptibilities to hydrolysis of the ligand.
This is believed to be a result of the slowness of inner-sphere electron transfer of the complex compared to electron transfer from an outer-sphere electron donor, the excess EDT A. These fractional dependences are at least partly due to competition between the reduction of the Mn(III) complex and the oxidation of the Mn(II) citrate complex. The rate constants and rate expressions for the dissolution of MnOOH by the ligands used in this study are shown in Table 6.2.
This may be the result of the unprotonated ligand binding to the surface and releasing a Mn(III)P2Or species upon solution. The decomposition was found to be totally reductive with no evidence for Mn(III) in solution. The fractional dependence is probably caused by the adsorption of EDTA to the surface, similar to pyrophosphate.
Implications of This Study
- Surface Waters
- Ground Waters
- Other Potential Evironments for Mn(III) Complexes
Because Mn(II) oxidation has been shown to proceed through a Mn(III) oxide phase, these Mn oxides represent another possible source of Mn(III) complexes. Again, the formation of Mn(III) complexes should occur in the surface microlayer or in areas of high biological activity. Another area in natural systems where large amounts of Mn(III) stabilizing ligands are likely present is soil water and sediments.
The processes considered were: oxidation of Mn 2+ to MnOOH, oxidation of MnOOH to MnO2, oxidation of Mn(II)L to Mn(III)L, reduction of MnOOH and MnO2 to Mn 2+, dissolution of MnOOH with a ligand and reduction of Mn(III)L in Mn(II)L. The rate constant used was for the reduction of Mn(IIl,IV) solids with hydroquinone. Although not all ligands promote the oxidation of Mn(Il)L to Mn(III)L as much as citrate, and although a micromolar concentration of such a ligand may seem unlikely, it should be noted that even if the concentration of the ligand were 2 orders of magnitude lower, it would Mn(III)L still the dominant oxidized species of Mn.
It could be argued that biological oxidation of Mn(II) would yield a higher MnOOH concentration. Groundwater represents another environment where potentially significant amounts of Mn(III) stabilizing ligands are present. Significant amounts of Mn(III) oxides are likely present in groundwater, either as minerals or as oxidation products formed in more oxygen-rich environments.
Directions for Future Research
- Mn(II) Oxidation
- Mn(III) Solution Complexes
- MnOOH Dissolution
Analysis of the products both in the presence and absence of oxygen can also reveal the mechanism of degradation of EDT A. If product peaks could be identified, it would be useful in determining both mechanism and stoichiometry. The actual oxidation pathway of the citrate would be crucial in determining how many cycles are required to destroy the citrate and whether the reaction is indeed catalytic.
If we want to approach the existence of such complexes in natural systems, it will be useful to know the reaction rates for different types of ligands. Perhaps the most useful area of research in this area would be the development of more sensitive techniques for the detection of Mn(III). Since natural levels of Mn are usually in the submicromolar range, such techniques would be crucial for conducting any studies in the natural environment or when attempting to simulate natural systems in the laboratory.
It would also be interesting to examine the surface of the solids during the reaction to determine the identity of surface ace complexes formed, if any, or at least to determine the dissolution pattern of the surface. It should be of interest to investigate the reaction of MnO2 with a ligand, such as citrate, which has the ability to both reduce Mn(IV) to Mn(III) and then to stabilize Mn(III). This study has found that Mn(III) complexes can form both from solution phase Mn(II) and Mn(III) solids.
Tables
Figures
