Precipitation and Dissolution of Iron and Manganese Oxides
by
Scot T. Martin
Division of Engineering and Applied Sciences 29 Oxford St., Pierce Hall, Room 122 Harvard University, Cambridge, MA 02138
E-mail:
[email protected] Web: http://www.deas.harvard.edu/~smartin
Submitted: September 2003
Chapter 4 of Environmental Catalysis (Editor, Vicki H. Grassian)
1
1. Introduction
2
Iron and manganese are the first and third most abundant transition metals in the
3
earth’s crust (5.6 × 104 ppm and 9.5 × 102 ppm, respectively) (1). The redox chemistries
4
of iron II/III and manganese II/III/IV have important roles and impacts in the
5
environment (2-5). In contrast, the second most abundant transition metal, titanium (5.7 ×
6
103 ppm), occurs only in the IV oxidation state. Abundant crustal aluminosilicates, such
7
as clays and quartz, also lack a dynamic redox chemistry in natural waters (6).
8
Redox chemistry strongly influences the precipitation and dissolution of Fe and
9
Mn solid phases (Figure 1) (7-10). Aqueous Fe(II) and Mn(II) are significant in natural
10
waters only in the absence of O2. Insoluble Fe(III) and Mn(III/IV) oxides form under oxic
11
conditions. Their solubilities limit the aqueous concentrations of Fe and Mn species.
12
The interconversions among redox states and physical states, while often
13
thermodynamically favorable, are frequently slow in the absence of catalysis. For
14
example, aqueous solutions of Mn(II) in the presence of O2 at pH = 8.4 are exoergic
15
toward oxidation, yet the uncatalyzed reaction proceeds slowly across years (11).
16
Surfaces, ligands, and other metals have varying levels of catalytic activity for Fe(II) and
17
Mn(II) oxidation (4, 12). As a second example, both the rate and the pathway of the
18
crystallization of supersaturated solutions is influenced catalytically by a range of factors,
19
including foreign surfaces, surface active ions, pH, and the magnitude of supersaturation
20
(4, 9). As a final example, iron and manganese oxide solids exposed to undersaturated
21
solutions have a free energy driving force favorable to dissolution, but the uncatalyzed
22
rates are often slow, especially in relation to the demand rate of biological organisms.
23
Catalyzed pathways, which usually entail the formation of surface complexes between
1
24
active ligands and oxide surface groups, are generally necessary to meet biological
25
demand. In these examples, the distinction is blurred between a true catalyst, which is
26
recycled during a reaction, and a stoichiomeric agent, which is consumed. As defined in
27
this chapter, a catalyst is any agent that increases the rate of a desired process, including
28
oxidation, precipitation, and dissolution, regardless of whether the agent is recycled or
29
consumed.
30
Quantifying and predicting Fe and Mn cycling is separately motivated within at
31
least four scientific communities (13):
32
(1) In the absence of more favorable terminal electron acceptors like O2 or NO3-,
33
Geobacter, Shewanella, and other microbes reduce Fe(III) and Mn(III/IV) oxides (14-
34
16). The reduction reaction couples with hydrocarbon oxidation to complete the
35
energy-producing metabolic pathways.
36
(2) In natural waters and soil zones having high dissolved oxygen concentrations, Fe and
37
Mn acquisition at concentrations sufficient for enzymatic function (17-20) challenges
38
microbes and plants (21-24). The enzymes developed when aqueous Fe(II) and
39
Mn(II) were abundant prior to oxygenation of the early Earth atmosphere (25).
40
Cellular FeT and MnT concentrations of 10-4 M (26) are 105 greater than typical ocean
41
waters (27) and 102 greater than typical acidic soil solutions (6). Organic ligands like
42
oxalate, which are common biological exudates, and siderophores, which are tailored
43
biological molecules, increase the dissolution rates of Fe(III) and Mn(III/IV) oxides
44
and increase Fe and Mn bioavailability (28-30).
2
45
(3) The carbon, nitrogen, sulfur, oxygen, and phosphorous geochemical cycles require
46
oxidation/reduction at many steps. Fe and Mn cycling is a key thermodynamic
47
regulator and kinetic catalyst in natural waters (2, 4, 5).
48
(4) The transport and fate of heavy metal pollutants in natural waters is strongly affected
49
by Fe and Mn oxide precipitation and dissolution (31). Heavy metals, especially in
50
sediments or groundwater, adsorb on Fe and Mn oxide surfaces (31-33). The heavy
51
metals are also incorporated in the Fe and Mn oxide matrix as impurities when
52
precipitation occurs or when new mixed metal/Fe and metal/Mn coprecipitates are
53
possible (31, 34-38). As such, the cycling of redox conditions in natural waters and
54
the associated precipitation and dissolution of Fe and Mn oxides lead to the cyclical
55
uptake and release of heavy metal pollutants (39-44). Movement of heavy metals
56
through soils is also increased when these metals adsorb onto iron oxide colloids (45-
57
47) and retarded when they adsorb onto iron and manganese oxide coatings on coarse
58
grains, such as quartz sand.
59
This chapter focuses on the rates and mechanisms of Fe and Mn oxide
60
precipitation and dissolution. The chapter’s organization is as follows. Thermodynamic
61
driving forces, illustrated by pE-pH diagrams, are shown to constrain oxidation,
62
precipitation, and dissolution (§2). Rates of aqueous Fe(II) and Mn(II) oxidation by O2
63
have both homogeneous and heterogeneous pathways (§3 and §4). Dissolution occurs by
64
proton-promoted, ligand-promoted, reductive, and synergistic pathways (§5). Modern
65
molecular techniques provide an increasing basis for mechanistic descriptions and
66
predictions of oxidation, precipitation, and dissolution (§6). Rather than reviewing all
67
available experimental techniques (48), three informative examples are chosen for §6, 3
68
including infrared spectroscopy, atomic force microscopy, and X-ray absorption
69
spectroscopy. Beyond the scope of this chapter are photochemical reactions (49-51),
70
biological reactions (16, 52-55), field studies (4, 56-58), and computational chemistry
71
(59).
72
2. Thermodynamic Driving Forces
73
The free energy driving forces relating the various Fe and Mn aqueous and solid
74
species (Figure 1) are represented in pE-pH diagrams (Figure 2) (9, 10, 60). The pE axis
75
represents the equilibrium partial pressures of O2 (increasing at high pE) and H2
76
(increasing at low pE) in the system, and hence the oxidizing or reducing power of the
77
environment. Specifically, pE is defined as:
78
1 pE = − pH − log PH2 2 1 pE = 20.77 − pH + log PO2 4
(1)
79
The dashed lines show the water stability region for 1 atm of gases. Above the top line,
80
water should thermodynamically form O2 if oxygen partial pressure is below 1 atm.
81
Similarly, below the bottom line, water should thermodynamically form H2 if hydrogen
82
partial pressure is below 1 atm. Natural environments have a range of pE/pH values:
83
oceans (pE = O2 saturated, pH = 8), surface waters of lakes and rivers (pE = O2 saturated,
84
pH = 4 to 6), mine waters (pE = O2 saturated, pH = 1 to 3), groundwater and sediments
85
(pE = 0 to 3, pH = 7 to 9), and swamps (pE = 0 to -3, pH = 5 to 7) (9, 61, 62).
86
In the aqueous phase in the Fe system, there is an interchange among aqueous
87
Fe(II) and Fe(III) redox species with decreasing pE. There is also a shift among the
88
dominant hydrolysis species, such as from [Fe(H2O)6]3+ to [Fe(H2O)5OH]2+ with
89
changing pH. In contrast, [Mn(H2O)6]2+ is the unique dominant at the common pE and 4
90
pH values of natural waters. Hydrolysis begins only for pH > 10, while the stabilization
91
of aqueous Mn(III) requires strong ligands (63).
92
Figure 2 is drawn for 10-6 M FeT (top) and 10-6 M MnT (btm) at 25ºC, where FeT is
93
the sum of all species including for example Fe2+(aq), Fe3+(aq), FeOH2+(aq), and the
94
various solids. MnT is similarly defined. The lines in the diagram shift to form smaller
95
gray regions and larger white regions for decreasing FeT and MnT. The white pE-pH
96
regions show where FeT and MnT are thermodynamically speciated entirely in the
97
aqueous phase. For instance, if the prescription of 10-6 M FeT or MnT initially includes
98
solid species (e.g., as prepared in the lab or in an aquifer subject to seasonal pE-pH
99
cycles) in the white regions, then it is predicted that these solids will dissolve. However,
100
the dissolution rate is slow compared to the timescales of days, weeks, and months
101
usually relevant to the environment (§5).
102
The gray pE-pH regions conversely show where some (viz. near the boundary
103
lines) or most (viz. moving inward in the gray region) of 10-6 FeT or MnT should
104
thermodynamically include solid phases, with a small amount of aqueous Fe and Mn
105
species in equilibrium (e.g., 10-9 to 10-15 aq Fe and Mn). However, the precipitation rate
106
may be slow (§3 and §4). For instance, in the absence of catalysis, aq Mn(II) at pE = 10
107
and pH = 8 persists for years in solution, even though oxidation and precipitation of Mn
108
oxide solid phases is thermodynamically favored. By increasing pH, supersaturation
109
increases and Mn(II) oxidizes and precipitates. However, the first solid to form is
110
generally the least favored thermodynamically, a phenomenon described as Ostwald’s
111
rule of stages. For instance, Mn(III) solids such as MnOOH form initially even when the
112
free energy of formation for MnIVO2 is greater. The MnIV oxides form only after long
5
113
aging times of MnIIIOOH (64). For a similar reason, in the diagram for Fe oxides,
114
amorphous Fe(OH)3(s) is employed in the analysis rather than thermodynamically
115
favored but slowly forming hematite (α-Fe2O3) and goethite (α-FeOOH).
116
The pE-pH diagrams are useful for establishing the thermodynamically favorable
117
pathways for transformations. However, predicting the transformation rates requires a
118
kinetic analysis (§3, §4, and §5).
119
3. Rates of Homogeneous Oxidation
120
Aqueous Fe(II) and Mn(II) are oxidized by reaction with dissolved O2 (5, 12, 65,
121
66). Thermodynamic analysis shows that the species Fe(III) is thermodynamically
122
favorable for pE > 13.2 at acidic pH (Figure 2). However, the reaction of Fe(II) with O2 is
123
not observed in the absence of catalysis at low pH. Only for pH > 6 is the rate appreciable,
124
with a lifetime of approximately 1 day in water equilibrated with air at 25ºC at pH = 6
125
(Figure 3). For increasing pH, the lifetime rapidly decreases, being 30 min at pH = 7. For
126
more alkaline pH, the reaction is even faster, but the dominant pathway shifts from
127
homogeneous to heterogeneous (§4.2) because the reaction product Fe(OH)3(aq) rapidly
128
polymerizes to form am-Fe(OH)3(s), thus providing a reactive surface (67, 68).
129
The pH dependence of the reaction rate for pH < 7 is explained by the mechanism
130
shown in Table 1 (69). With increasing pH, the dominant aqueous species shifts from
131
[Fe(H2O)6]2+ to [Fe(H2O)5OH]+ to [Fe(H2O)4(OH)2]0, each of which having a
132
progressively more rapid water-exchange rate and faster bimolecular reaction rate
133
constant with O2 (70). The overall reaction rate is the sum of parallel pathways (Table 1
134
and Figure 3). The rate is also catalyzed by other dissolved species such as Cu2+, Fe3+,
135
Al3+, Co2+, and Mn2+ (71-73).
6
Compared to Fe(II)(aq), reaction of Mn(II)(aq) with O2 is at least 106 times slower
136 137
at circumneutral pH (Table 1 and Figure 3). Only for pH > 8 does the reaction rate
138
become appreciable. The reaction proceeds through the aqueous Mn(OH)2 species,
139
although the bimolecular rate constant of Mn(OH)2 with O2 is 105.2 lower than that of
140
Fe(OH)2. The reaction product (Mn(III)), in the absence of strongly complexing ligands,
141
rapidly polymerizes to form Mn oxide solids (67, 68), which catalyze further Mn(II)
142
oxidation (§4.2). Hence, separating homogeneous from heterogeneous pathways in
143
Mn(II) oxidation is difficult because they occur simultaneously under most experimental
144
conditions.
145
4. Rates of Heterogeneous Oxidation
146
The rate of Mn(II) and Fe(II) oxidation by O2 is catalyzed by metal oxide surfaces
147
(>S) (12, 72, 74-76). These surfaces are terminated by hydroxyl groups (>SOH), which
148
bind Mn(II) and Fe(II) as (>SO)2Mn and (>SO)2Fe. The inner-sphere surface complexes
149
promote rapid oxidation, just as OH ligands do for the homogeneous complexes (§3). The
150
catalysis occurs both on foreign surfaces (e.g., Mn(II) on FeOOH) (§4.1) and also for the
151
special case of autocatalysis (e.g., Mn(II) on MnOOH producing additional MnOOH)
152
(§4.2).
153
4.1 Mineral Surfaces
154
Reaction rates at surfaces (77, 78) are given either as the conversion rate per unit
155
surface area of the foreign surface (mol m-2 s-1) or as the conversion rate per liter of a
156
particulate suspension (M s-1) (13). The latter is the basic observable in experiments
157
employing particulates, whereas the former is a more intrinsic measure, which can be
158
estimated for a suspension of known loading (g L-1) and specific surface area (m2 g-1)
7
159
(Table 2). For single crystals, the conversion rate per unit surface is measured directly.
160
For comparing the relative importance of homogeneous versus heterogeneous oxidation
161
rates under a specific set of conditions, the heterogeneous rate expressed as (M s-1) is
162
more convenient because these units are the same as for homogeneous oxidation rates
163
(Table 1). Heterogeneous oxidation rates in natural waters commonly exceed
164
homogeneous oxidation rates.
165
Rate equations for heterogeneous oxidation are summarized in Table 2, where k is
166
the oxidation rate coefficient (M-1 s-1), [>SOFe2+] and [>SOMn2+] are the respective
167
binuclear surface concentrations of adsorbed Fe(II) and Mn(II) (mol m-2), KH is the
168
Henry’s law partition coefficient of O2 (M atm-1), PO2 is the partial pressure of O2 (atm),
169
A is the particulate surface area (m2), and V is the container volume of the aqueous
170
particulate suspension (L). [>SOFe2+] and [>SOMn2+], which depend on FeT/MnT and pH,
171
are typically quantified by measurements of adsorption isotherms. Examples of
172
heterogeneous rate coefficients are log k = 0.7 (M-1 s-1) for Fe2+/O2 reaction on FeOOH
173
and log k = -1.55 (M-1 s-1) for Mn2+/O2 reaction on Al2O3.
174
In many real-world applications, aqueous concentrations of Fe(II) and Mn(II) are
175
known rather than surface concentrations [>SOFe2+] and [>SOMn2+]. It is convenient to
176
recast the heterogeneous oxidation rate laws in terms of aqueous concentrations by the
177
use of adsorption isotherm equations. For incomplete surface coverage not sufficient to
178
significantly perturb the surface charge, surface concentration and aqueous concentration
179
are related by a simple mass action equilibrium law quantified by an intrinsic binding
180
constant βs of the surface complex. This more detailed rate law is provided in Table 2.
181
There, [>SOH] is the site density of OH groups on the metal oxide surface (mol m-2).
8
182
The first term gives the intrinsic chemical reactivity (i.e., the rate coefficient). The second
183
term relates the surface concentration of adsorbed manganese to the aqueous
184
concentration of manganese (i.e., a 3D to 2D transformation). The third term expresses
185
the aqueous oxygen concentration. The fourth term scales the specific surface reactivity
186
(mol m-2 s-1) to the particulate suspension reactivity (M s-1).
187
4.2 Autocatalysis
188
A special category of heterogeneous oxidation occurs when the product of the
189
oxidation further accelerates the reaction rate (79, 80). For example, the oxidation of
190
Mn(II) produces MnOOH(s), as follows:
191
homogeneous or → 4 MnOOH(s) + 8 H+ O2 + 4 Mn2+ + 6 H2O ⎯⎯⎯⎯⎯ heterogeneous
(2)
192
As the reaction proceeds, the MnOOH(s) surface area and hence the heterogeneous
193
reaction rate increase. The rate laws of autocatalysis (Table 3) are less precise than those
194
of heterogeneous reactions on foreign mineral surfaces. Detailed descriptions for the
195
autocatalysis pathways are hindered both by the complexities of separating homogeneous
196
from heterogeneous pathways and by limitations in characterizing the increasing mineral
197
surface area and the altering mineral phases during reaction.
198
5. Dissolution Rates
199
Iron and manganese oxide solids dissolve at the rates shown in Figure 4 (center;
200
right) when the contacting aqueous solution is strongly undersaturated (i.e., no back
201
reaction from precipitation). The dissolution rates depend on many factors (2, 4, 13, 81-
202
83). For instance, the rates increase with acidic pH. There are also several parallel
203
pathways having differing dissolution rates, including proton-promoted (§5.1) (slowest),
204
ligand-promoted (§5.2), reductive (§5.3), and synergistic (§5.4) (fastest). The
9
205
stoichiometries and the relative rates of these reaction pathways are given in Table 4. The
206
rates further depend on crystalline phase: amorphous Fe(OH)3 dissolves at least ten times
207
faster than γ-FeOOH. The rates also depend on initial chemical or physical preparation
208
and often pass through initial transients of rapid dissolution, which are at least ten times
209
faster than the steady-state dissolution rates (84, 85).
210
Although the rates depend on sample crystallinity and preparation and are quite
211
variable on a linear scale, on a log scale the differences among samples are less apparent.
212
With this caveat, organizational statements are possible. For example, γ-FeOOH
213
dissolves more slowly than amorphous Fe(OH)3; iron oxides dissolve more slowly than
214
manganese oxides; reductive dissolution is faster than ligand-promoted dissolution; and
215
proton-promoted dissolution is the slowest of all. The dissolution rates of iron and
216
manganese oxides can also be compared with the dissolution rates of other minerals
217
(Figure 4, left versus center and right). Iron oxides dissolve at rate comparable to chain
218
and sheet aluminosilicates. Manganese oxides dissolve at rates comparable to carbonates,
219
although the precise rate depends strongly on reductant concentration.
220
5.1 Proton-Promoted
221
Protons increase dissolution rates (86), which is rationalized by a catalytic role in
222
depolymerization. For example, dimers such as [Fe2(OH)2]4+, [Al2(OH)2]4+, and
223
[(VO)2(OH)2]2+ decompose at the compound rate: k = k1 + k2[H+], where k is the pseudo
224
first-order rate coefficient, k1 is the H2O reaction pathway, and k2 is the H+ reaction
225
pathway (87-90). Protons accelerate the rate by attaching to the oxygen in the hydroxyl
226
groups bridging the metals, thus removing electron density and weakening the bond
227
strength of the metal-oxygen linkage:
10
228 229
A mineral such as FeOOH or Fe2O3 is regarded as an infinite n-mer extension of
230
[Fe2(OH)2]4+. Mineral dissolution is then a stepwise depolymerization, and protons have
231
the role of weakening bonds and thus increasing dissolution rates (83). The log of the rate
232
is proportional to the log of the surface proton concentration (>SOH2+).
233
5.2 Ligand-Promoted
234
Ligands binding as inner-sphere complexes to the surface groups of iron and
235
manganese oxides increase dissolution rates (30, 83, 86, 91-93). The increase is
236
proportional to the ligand surface concentration and the ligand binding strength. The rate
237
law is: R = kL [>SL] where R is the dissolution rate (mol m-2 s-1), kL is the rate coefficient
238
(s-1), and [>SL] is the surface concentration of the ligand. For a homologous series of
239
surface structures, the rate coefficient is proportional to the adsorption strength of the
240
ligand, e.g., kL = 4.5 × 10-10 KL – 1.1 × 10-6 where KL is the Langmuir binding constant
241
for the series oxalate, glutarate, and malonate binding to hematite (93). The relationship
242
occurs because strong electron donation by the oxygen-ligand to a surficial metal atom
243
removes electron density in the bonds between the metal atom and the oxygen atoms of
244
the mineral lattice, thus weakening the bond and lowering the energy barrier for the
245
dissolution of the metal atom (82). In this regard, Mn and Fe cations bind much more
246
strongly to oxygen- than nitrogen-containing ligands (9). Furthermore, ligands effective
247
for promoting dissolution have two or more functional groups capable of chelation to
248
form inner-sphere bidentate mononuclear complexes. An example is oxalate (cf. Figure 6, 11
249
n = 0). In contrast, ligands forming bidentate binuclear complexes such as phosphate or
250
borate stabilize the surface against attack by H+ and H2O and thus reduce (inhibit) the
251
dissolution rate. Monodentate ligands such as acetate do not perceptibly affect the
252
dissolution rate.
253
5.3 Reductive
254
Reductants rapidly accelerate iron and manganese oxide dissolution (94-99).
255
Examples of reductants are ascorbic acid, hydrogen sulfide, and phenols. A reductant
256
typically forms an inner-sphere complex at the surface, though not always so. When an
257
electron is transferred to a Mn(III/IV) oxide, a surficial Mn(II) ion locked inside an oxide
258
lattice is formed. Because Mn(II) oxides are much more soluble than the corresponding
259
higher oxides (Figure 2), rapid Mn(II) depolymerization occurs, which is followed by
260
release to the aqueous phase of Mn(II). Fe(III) I oxides are similarly reduced, followed by
261
the release of aqueous Fe(II).
262 263
The Mn oxide solids are strong oxidants capable of oxidizing common organic matter:
264
MnO2(s) + 2 e- + 4 H+ → Mn2+ + 2 H2O, pE0 = 20.8, E10 2 = +1.23 V
(3)
265
MnOOH(s) + e- + 3 H+ → Mn2+ + 2 H2O, pE0 = 25.4, E10 2 = +1.50 V
(4)
266
H2CO(aq) + 2 MnO2(s) + 4 H+ → 2 Mn2+ + CO2(aq) + 3 H2O
(5)
∆pE0 = 18.4, E 0 = +1.09 V
267 268
The thermodynamic driving force of the reaction is often a good predictor of rate for a
269
homologous series of reactants (i.e., a linear free energy relationship). For example, for a
270
series of substituted phenols like p-methylphenol at 10-4 M and pH = 4.4, the Hammett
12
271
constant σ is a predictor of reductive dissolution rate R (mol m-2 s-1) of manganese oxide
272
(95): log10 R = −7.79 − 3.63 σ
273 274
The correlation holds because the Hammett constant is also a predictor of reduction
275
potential, which is the driving force for manganese oxide dissolution. Equation (6) is
276
derived over a ∆σ range of 0.66, so the effect on log R is substantial.
277
5.4 Synergistic
278
(6)
Some species interact cooperatively to increase the dissolution rate above the sum
279
of the individual dissolution rates. For example, Fe(III) oxides dissolve more rapidly in
280
the presence of aqueous Fe(II) and oxalate than in the presence of either separately (100).
281
The Fe(II)-oxalato aqueous complex is a strong reductant, which rapidly reduces the iron
282
oxide through reductive dissolution. Another example is the rapid dissolution of iron
283
oxides in the presence of a ligand-reductant pair, such as oxalate and ascorbate (101).
284
Ascorbate reduces Fe(III) to Fe(II) at the surface of the Fe(III) oxide, while oxalate forms
285
an inner-sphere complex at the surface and rapidly dissolves Fe(II). A final example is a
286
surface and aqueous ligand pair, such as oxalate and a trihydroxamate siderophore,
287
desferrioxamine B (DFO-B) (102). At low concentrations, oxalate adsorbs to the iron
288
oxide surface. However, for sufficiently low oxalate concentrations, the free energy
289
driving force toward dissolution is small because the total aqueous solubility of an iron
290
oxide is low. The DFO-B, however, binds aqueous iron very strongly. Addition of DFO-
291
B provides a sink for aqueous iron. Under these conditions, in one step oxalate dissolves
292
Fe(III) as a Fe(III)-oxalato complex, which rapidly hands off the Fe(III) to DFO-B.
293
Oxalate is then free to recycle to the iron oxide surface, leading to further dissolution.
13
294
The rates in all three examples are proportional to the surface concentrations of the
295
reactive species.
296
5.5 Master Equation
297 298 299
The overall dissolution rate R (mol m-2 s-1) is the sum of several process, as represented conceptually in the equation below (83): R = k H [>SOH +2 ]n + kOH [>SO- ]m + kaq + ∑ a k L ,a [>SL a ]
− ∑ b k I ,b [>SIb ] + ∑ c k R ,c [>SR c ] + ∑ d kM ,d [>SM d ]
(7)
300
for proton-promoted (H), hydroxide (OH), aqueous (aq), a ligands (L), b inhibitors (I), c
301
reductants (R), and d metal ions (M). Under tested applications, no more than two or
302
three terms dominate the overall dissolution rate. Hence, eq (7) is not to be interpreted
303
literally in its complexity.
304
6. Molecular Environmental Chemistry
305
The growth of molecular environmental chemistry in recent years provides a
306
renewed basis for testing and constraining models of mineral precipitation and
307
dissolution (48, 103-106). The left hand side of eqs (6) or (7) and the quantitative
308
formulations given in Tables 1 to 3 provide examples of macroscopic quantities, which in
309
these cases are the measurable dissolution rates. From the dependencies of these direct
310
macroscopic observables on factors such as pH or reductant concentration, molecular
311
descriptions of surface processes, such as given in the right hand side of eq (7), are
312
constructed by inference. These molecular inferences from macroscopic observables are
313
the common approach for the studies described in §2-§5. In the latest research (e.g., §6.1
314
to §6.3), direct measurement of the molecular quantities, which is always preferable to
315
inference, is an accelerating trend and provides a manner for further improving
14
316
mechanistic descriptions and thus our predictive capability of macroscopic quantities. Of
317
the numerous available techniques (48, 105), three illustrative examples include infrared
318
spectroscopy (§6.1), atomic force microscopy (§6.2), and X-ray absorption spectroscopy
319
(§6.3).
320
6.1 Infrared (IR) Spectroscopy
321
Inner-sphere surface complexes, whose concentrations appear as [>SX] in the
322
right hand side of eq (7), are directly observable by infrared spectroscopy (93, 107-111).
323
Infrared absorption occurs at the resonance frequencies of atomic vibrations. The number
324
of vibrations and their wavenumber positions depend upon the symmetry of the surface
325
complex and its local chemical environment.
326
Duckworth and Martin (93) employ a homologous series of dicarboxylic acids,
327
including oxalate, malonate, succinate, glutarate, and adipate, to study the effect of chain
328
length on the nature of the surface complexation structures formed on hematite. The
329
infrared spectra (Figure 5) lead to the deduction of the structure of the inner-sphere
330
surface complexes (Figure 6). Oxalate, malonate, and glutarate form bidentate surface
331
complexation structures, while succinate and adipate form monodentate structures. The
332
ligand-promoted dissolution rates are also measured, and the three bidentate ligands have
333
dissolution rates at least 10 to 100 times faster than the monodentate structures (Table 5).
334
These results nicely illustrate the gains possible from molecular level insights. For
335
instance, rationalizing the difference in the dissolution rate of glutarate versus that of
336
succinate is guesswork when only macroscopic measurements are available.
337
6.2 Atomic Force Microscopy (AFM)
15
338
The master equation for dissolution (eq (7)), which is derived from macroscopic
339
observations, assumes each surface site is equivalent. Under this treatment, the surface is
340
a terrace of infinite extent. In reality, manganese and iron oxide surfaces have
341
microtopogaphy, including for example terraces, steps, kinks, and pits (99, 112, 113).
342
Terrace ions locked into place by five bonds to nearest neighbors are released more
343
slowly than ions at kink positions having only three nearest neighbors (114). Moreover,
344
point defects, line defects, and other dislocations dissolve rapidly. A master equation
345
derived solely from macroscopic observations is thus inherently limited in its accuracy
346
and range of application.
347
The atomic force microscope (AFM) allows measurement of mineral surface
348
topography under aqueous conditions (Figure 7) (115, 116). A sharp tip with a curvature
349
radius of approximately 10 nm is rastered laterally across the surface, and vertical
350
deflections of the tip provide the topography of the underlying surface (117). The
351
microscopic data collected in time series allow new models of mineral dissolution to be
352
formulated and constrained (118). The dynamic range of the AFM, which is limited by
353
the scanning rate on the upper side and the tip lifetime on the lower side, is convenient
354
for studying minerals having dissolution rates of 10-6 to 10-10 mol m-2 s-1 (119). Hence,
355
manganese oxides are easily studied (Figure 4), whereas iron oxides dissolve too slowly
356
for AFM work, except under extreme conditions. Manganite (MnOOH) reductive
357
dissolution shows overall step retreat (e.g., circled region in Figure 7), while certain areas
358
of the surface are simultaneously reconstructing to minimize surface energy by filling in
359
pits (e.g., central arrow in Figure 7). Scaling of the complex time-dependent changes
16
360
shown in Figure 7 to the surface-averaged (i.e., macroscopic) dissolution rate is an
361
outstanding research challenge.
362
Atomic force microscopy is also useful for studying precipitation. Junta and
363
Hochella (120, 121) report on the heterogeneous oxidation of Mn(II) at hematite, goethite,
364
and albite surfaces by O2 for 7.8 < pH < 8.7. For the specified reaction conditions,
365
manganese oxide precipitates form preferentially at the steps and kinks of the substrates,
366
which could suggest a higher rate of heterogeneous oxidation at those microtopographic
367
features. Alternatively, the oxidation may occur widely on the terrace, and the oxidized
368
monomeric products may diffuse rapidly to more stable step and kink positions.
369
6.3 X-ray Absorption Spectroscopy (XAS)
370
X-ray absorption spectroscopy responds to the local coordination environment
371
and chemical oxidation state of the inner-sphere surface complexes of heavy metals on
372
iron and manganese oxides (122-126). X-ray absorption is modulated at energies near
373
and above an absorption edge because the final quantum mechanical state of the system,
374
which is a hole in the sample and a photoelectron in vacuum, is affected by scattering and
375
constructive/destructive interference of the ejected electron at its de Broglie wavelength
376
by coherent shells of neighboring atoms. Experiments must be conducted on synchrotron
377
beam lines to obtain the high intensity monochromatic X-rays necessary for XAS studies.
378
The interaction of arsenic with manganese oxide surfaces provides an example of
379
the important molecular information possible with XAS studies (127, 128). The
380
macroscopic observation is that As(III) reductively dissolves Mn(IV) oxides to form
381
As(V). Understanding the mechanisms and pathways in greater molecular detail and
382
hence quantifying the reductive dissolution rates in greater accuracy is important because
17
383
As(III) is more mobile in aquifers than As(V). X-ray absorption near-edge spectroscopy
384
(XANES) shows that when aqueous As(III) adsorbs to the surface of MnO2, the surface
385
complex is As(V). The extended X-ray absorption fine-structure (EXAFS) results show
386
that the inner-sphere complex is bidentate binuclear, (>MnIVO)2AsVOOH (Figure 8).
387
High-precision, quantitative structural information on local ordering is also derived from
388
EXAFS measurements: the As-Mn distance is 3.22 Å. The precise distances coupled with
389
computational modeling of optimized structures allow an investigator to rule out a host of
390
other possibilities, e.g., bidentate mononuclear binding, mixed Mn/As co-precipitates,
391
Mn(II)/Mn(III) structures, and As(III) structures. The only optimized structure providing
392
good agreement with the measurements (Figure 8, left) is shown in Figure 8 (right). XAS
393
studies show that bidentate binuclear structures are also formed by arsenic on iron oxides
394
(129-131).
395
7. Concluding Remarks
396
Because the uncatalyzed rates of Fe and Mn oxide precipitation and dissolution
397
are often slow, an understanding of the catalyzed pathways and mechanisms is necessary
398
for quantitative predictions of Fe and Mn transformations and the associated impacts in
399
the environment. Tables 1 to 5 and Figures 3 and 4 summarize the known rate
400
information. Aqueous Fe(II) and Mn(II) oxidation by O2 is catalyzed heterogeneously by
401
other mineral surfaces. The oxidized products polymerize and precipitate, first as higher
402
energy oxides having little long-range order and after aging as lower energy crystalline
403
oxides. When exposed to undersaturated aqueous conditions, the oxides dissolve at rates
404
dependent upon catalytic additives, including proton-promoted, ligand-promoted,
405
reductive, and synergistic dissolution pathways. Molecular techniques increasingly
18
406
provide a detailed mechanistic description of these processes. Under favorable
407
circumstances, these new descriptions are detailed enough that they can be modeled by ab
408
initio and semiempirical methods, allowing for direct comparison of experimental
409
observations and computational results and thus for the further improvement and
410
refinement of the latter.
411 412
Acknowledgments
413
STM is grateful for support received from the Chemical Sciences, Geosciences, and
414
Biosciences Division of the Office of Basic Energy Sciences in the U.S. Department of
415
Energy and the American Chemical Society Petroleum Research Fund.
19
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25
109. SJ Hug and B Sulzberger, In situ Fourier transform infrared spectroscopic evidence for the formation of several different surface complexes of oxalate on TiO2 in the aqueous phase, Langmuir 10: 3587-3597, 1994. 110. ST Martin, JM Kesselman, DS Park, NS Lewis, and MR Hoffmann, Surface structures of 4-chlorocatechol adsorbed on titanium dioxide, Environ. Sci. Technol. 30: 2535-2542, 1996. 111. SJ Hug, In situ Fourier transform infrared measurements of sulfate adsorption on hematite in aqueous solutions, J. Colloid Interface Sci. 188: 415-422, 1997. 112. WK Burton, N Cabrera, and FC Frank, The growth of crystals and the equilibrium structure of their surfaces, Philos. Trans. R. Soc. London. Ser. A 243: 299-358, 1951. 113. C Eggleston, S Higgins, and P Maurice, Scanning probe microscopy of environmental interfaces, Environ. Sci. Technol. 32: 456A-459A, 1998. 114. IV Markov, Crystal Growth for Beginners, Singapore: World Scientific, 1995, 422 pp. 115. AJ Gratz, S Manne, and PK Hansma, Atomic force microscopy of atomic-scale ledges and etch pits formed during dissolution of quartz, Science 251: 1343-1346, 1991. 116. PE Hillner, AJ Gratz, S Manne, and PK Hansma, Atomic-scale imaging of calcite growth and dissolution in real time, Geology 20: 359-362, 1992. 117. VJ Morris, AR Kirby, and AP Gunning, Atomic Force Microscopy for Biologists, London: Imperial College Press, 1999, 332 pp. 118. OW Duckworth and ST Martin, Connections between surface complexation and geometric models of mineral dissolution investigated for rhodochrosite, Geochim. Cosmochim Acta 67: 1787-1801, 2003. 119. PM Dove and FM Platt, Compatible real-time rates of mineral dissolution by atomic force microscopy (AFM), Chem. Geol. 127: 331-338, 1996. 120. JL Junta and MF Hochella Jr., Manganese (II) oxidation at mineral surfaces: A microscopic and spectroscopic study, Geochim. Cosmochim. Acta 58: 4985-4999, 1994. 121. J Junta, MF Hochella, Jr., and D Rimstidt, Linking microscopic and macroscopic data for heterogeneous reactions illustrated by oxidation of manganese (II) at mineral surfaces, Geochim. Cosmochim. Acta 61: 149-159, 1997. 122. EE Koch, ed. Handbook of Synchrotron Radiation, Amsterdam: North-Holland, 1983, pp. 123. DC Koningsberger and R Prins, ed. X-ray Absorption: Principles, Applications, Techniques of EXAFS, SEXAFS, and XANES, New York: Wiley, 1987, 673 pp. 124. KF Hayes, AL Roe, GE Brown, KO Hodgson, JO Leckie, and GA Parks, In situ Xray absorption study of surface complexes: Selenium oxyanions on αFeOOH, Science 238: 783-786, 1987. 125. GE Brown, Jr. and GA Parks, Sorption of trace elements on mineral surfaces: modern perspectives from spectroscopic studies and comments on sorption from the marine environment, Int. Geol. Rev. 43: 963-1073, 2001. 126. PA Fenter, ML Rivers, NC Sturchio, and SR Sutton, ed. Applications of Synchrotron Radiation in Low-Temperature Geochemistry and Environmental Science, Washington, D.C.: Mineralogical Society of America, 2002, 579 pp. 26
127. BA Manning, SE Fendorf, B Bostick, and DL Suarez, Arsenic(III) oxidation and arsenic(V) adsorption reactions on synthetic birnessite, Environ. Sci. Technol. 36: 976-981, 2002. 128. AL Foster, GE Brown, and GA Parks, X-ray absorption fine structure study of As(V) and Se(IV) sorption complexes on hydrous Mn oxides, Geochim. Cosmochim. Acta 67: 1937-1953, 2003. 129. GA Waychunas, BA Rea, CC Fuller, and JA Davis, Surface-chemistry of ferrihydrite. 1. EXAFS studies of the geometry of coprecipitated and adsorbed arsenate, Geochim. Cosmochim. Acta 57: 2251-2269, 1993. 130. GA Waychunas, JA Davis, and CC Fuller, Geometry of sorbed arsenate on ferrihydrite and crystalline FeOOH: Reevaluation of EXAFS results and topological factors in predicting sorbate geometry and evidence for monodentate complexes, Geochim. Cosmochim. Acta 59: 3655-3661, 1995. 131. S Fendorf, MJ Eick, P Grossl, and DL Sparks, Arsenate and chromate retention mechanisms on goethite. 1. Surface structure, Environ. Sci. Technol. 31: 315-320, 1997. 132. B Simon, Dissolution rates of NaCl and KCl in aqueous-solution, J. Cryst. Growth 52: 789-794, 1981. 133. E Busenberg and LN Plummer, A comparative study of the dissolution and crystal growth kinetics of calcite and aragonite, U.S. Geol. Surv. Bull. 1578, 1986. 134. HU Sverdrup, The Kinetics of Base Cation Release due to Chemical Weathering, Lund: Lund University Press, 1990, 245 pp. 135. RA Wogelius and JV Walther, Olivine dissolution at 25ºC: Effects of pH, CO2, and Organic-Acids, Geochim. Cosmochim. Acta 55: 943-954, 1991. 136. JI Drever, The effect of land plants on weathering rates of silicate minerals, Geochim. Cosmochim. Acta 58: 2325-2332, 1994. 137. JV Walther, Relation between rates of aluminosilicate mineral dissolution, pH, temperature, and surface charge, Am. J. Sci. 296: 693-728, 1996. 138. AA Jeschke, K Vosbeck, and W Dreybrodt, Surface controlled dissolution rates of gypsum in aqueous solutions exhibit nonlinear dissolution kinetics, Geochim. Cosmochim. Acta 65: 27-34, 2001. 139. OW Duckworth and ST Martin, Role of molecular oxygen in the dissolution of siderite and rhodochrosite, Geochim. Cosmochim. Acta, accepted July, 2003.
27
List of Tables Table 1.
Rate equations for homogeneous phase oxidation of aqueous Fe(II) and Mn(II) by O2 (5, 12). The product Mn(OH)2+ rapidly polymerizes, forming a Mn oxide precipitate such as MnOOH(s).
Table 2.
Rate equations for catalytic heterogeneous oxidation of aqueous Fe(II) and Mn(II) by O2 on mineral surfaces (5, 12, 78). The surface group >SOM2+ is binuclear (i.e., (>FeO)2Fe2+).
Table 3.
Rate equations for the autocatalytic heterogeneous oxidation of aqueous Fe(II) and Mn(II) by O2 (5, 12, 99). The oxidation products of Fe(II) and Mn(II) precipitate and provide increasing particulate surface area loading and hence greater reactivity with time. As compared to Table 2, the formulation in Table 3 is less precise due to the poor characterization and the time evolution of the precipitate product.
Table 4.
Stoichiometry and relative dissolution rates of the proton-promoted, ligandpromoted, reductive, and synergistic pathways (2, 4, 83). Photoreductive pathways are omitted.
Table 5.
Physical data and experimental results for the ligand-promoted dissolution of iron oxide (hematite) by dicarboxylic acids (93). The numbers in parentheses give the carbon chain length, n, in -OOC(CH2)nCOO-. Also given are the Langmuir binding constants, K (M-1), of the inner-sphere surface complexes, which are determined by analysis of infrared spectra. Conditions: 5 mM acetate buffer, pH = 5.0, 2 g L-1 hematite, 25°C.
28
List of Figures Figure 1.
Common dissolved and precipitated iron and manganese species in several oxidation states.
Figure 2.
pE-pH diagrams of iron and manganese species (5, 10, 60). The water stability field is shown for PO2 = 1 atm and PH2 = 1 atm.
Figure 3.
Homogeneous phase oxidation rates of aqueous Fe(II) and Mn(II) by O2 (5, 12).
Figure 4.
Comparison of pH-dependent dissolution rates of iron and manganese oxides with other common minerals (86, 99, 132-139). Rates apply when the aqueous solution concentrations are far from equilibrium, i.e., strongly undersaturated. (Ascorbic acid rates are shown for reductive dissolution.)
Figure 5.
Infrared spectra of the surface-adsorbed complexes of dicarboxylic acids on hematite at pH = 5.0 (93). ∆ν is the difference between νas(CO2) and νs(CO2). For comparison, gray bars indicate the absorption regions of νas(CO2) and νs(CO2) vibrations of the aqueous species. The bidentate structures cause faster ligand-promoted dissolution than the monodentate structures (Table 5).
Figure 6.
Proposed surface chemical reactions of oxalate, malonate, succinate, glutarate, and adipate to a hematite surface hydroxyl group at pH = 5.0 (93). The bidentate structures cause faster ligand-promoted dissolution than the monodentate structures (Table 5).
Figure 7.
AFM deflection-mode micrographs of surfaces changes observed in situ at 298 K during exposure of manganite (γ-MnOOH) to ascorbic acid at pH =
29
3.2 (10 mM NaNO3) (99). (A) Prior to exposure to reductant. (B) Same as A after exposure to 1 mM reductant for 95 min. In height images (not shown), z-scale is 125 nm. Figure 8.
Structural information for surface complexes from synchrotron-based extended X-ray absorption fine structure (EXAFS) measurements (127). (left) Radial structure functions (not phase corrected) for As(III)- and As(V)treated synthetic birnessite (MnO2) and 1.0 mM As(V) solution. Dashed lines are the fits to the experimental RSF data and peaks correspond to As-O and As-Mn atomic shells around the As atom. (right) Structural diagram of MnO2 crystallite showing possible linkages between an arsenate ion (As(V) tetrahedron) and a pair of edge-linked MnO6 octahedra. As-Mn interatomic distance is 3.22 Å, suggesting a bidentate binuclear complex. (Adapted from ref (127).)
30
Homogeneous Fe2+ Oxidation Fe2+ + O2
log k0 = –5.1 (M-1 s-1)
→ Fe3+ + O−2
Fe(OH)+ + O2 → Fe(OH)2+ + O −2
log k1 = +1.4 (M-1 s-1)
Fe(OH) +2 + O 2−
log k2 = +6.9 (M-1 s-1)
Fe(OH)2 + O2
→
[Fe(OH)+] = K1[Fe2+]/[H+]
log K1 = –9.5 (M)
[Fe(OH)2] = β2[Fe2+]/[H+]2
log β2 = –20.6 (M2) log KH = –2.9 (M atm-1)
[O2] = KH PO2 − ( d [Fe 2+ ] dt )
homo
= k0 [O 2 ][Fe 2+ ] + k1[O 2 ][Fe(OH) + ] + k2 [O 2 ][Fe(OH) 2 ] = ( k0 [Fe 2+ ] + k1 K1[Fe 2+ ] /[H + ] + k2 β 2 [Fe 2+ ] /[H + ]2 ) K H PO2
Homogeneous Mn2+ Oxidation
Mn(OH)2 + O2
→
Mn(OH) +2 + O 2−
[Mn(OH)2] = β2[Mn2+]/[H+]2 − ( d [Mn 2+ ] dt )
Table 1.
homo
log k2 = +1.7 (M-1 s-1) log β2 = –22 (M2)
= k2 [O 2 ][Mn(OH) 2 ] = ( k2 β 2 [Mn 2+ ] /[H + ]2 ) K H PO2
-2 -1
-1
-1
Table 2.
Detailed formulation
Rate of release of particulate suspension (M s )
Rate per mineral surface area (mol m-2 s-1)
Heterogeneous Mn2+ Oxidation
Detailed formulation
Rate of release of particulate suspension (M s )
Rate per mineral surface area (mol m s )
Heterogeneous Fe2+ Oxidation
) hetero
(
= k[>SOFe 2+ ]( mol / m2 ) K H PO2
)
hetero
=
hetero
PO2
(
( m2 )
)( A
)
)
V( L )
)
= k[>SOMn 2+ ]( mol / m2 ) K H PO2
H
V( L )
hetero
− ( d [Mn 2+ ] dt )
=
)(K
H
PO2
( m2 )
)( A
( m2 )
)( A
)
V( L )
V( L )
) 10-12.7 M, and [>SOH] = 1.6 × 10-5 mol m-2.
Example: for Mn2+/O2 reaction on FeOOH, log k = -0.16 (M-1 s-1), βs =
k β s [Mn 2+ ][>SOH]( mol / m2 ) [H + ]2
hetero
= k[>SOMn 2+ ]( mol / m2 ) K H PO2
(
log k = -1.55 (M-1 s-1) for Mn2+/O2 reaction on Al2O3
Examples: as above
(
)
)(K
( m2 )
)( A
log k = -0.16 (M-1 s-1) for Mn2+/O2 reaction on FeOOH
− ( d [Mn 2+ ] dt )
Examples:
− d [>SOMn 2+ ]( mol / m2 ) dt
(
hetero
(
= k[>SOFe 2+ ]( mol / m2 ) K H PO2
k β s [Fe 2+ ][>SOH]( mol / m2 ) [H + ]2
(
− ( d [Fe 2+ ] dt )
Example: as above
− ( d [Fe 2+ ] dt )
Example: log k = 0.7 (M-1 s-1) for Fe2+/O2 reaction on FeOOH
(
− d [>SOFe 2+ ]( mol / m2 ) dt
Heterogeneous Autocatalytic Fe2+ Oxidation
Autocatalytic rate (M s-1)
⎛ d [Fe 2+ ] ⎞ −⎜ = k[FeO x (s)][Fe 2+ ] ⎟ ⎝ dt ⎠ autocatalytic FeOx is poorly characterized and can contain several mineral phases (e.g., amorphous Fe(OH)3 or FeOOH).
Heterogeneous Autocatalytic Mn2+ Oxidation
⎛ d [Mn 2+ ] ⎞ −⎜ = k[MnO x (s)][Mn 2+ ] ⎟ ⎝ dt ⎠autocatalytic Autocatalytic rate (M s-1)
Example: ks = 5 × 1018 M-4 day-1 where MnOx is poorly characterized and can contain several mineral phases (e.g., MnOOH, Mn3O4, MnO2, or birnessite).
Table 3.
Iron (III) Oxide Dissolution
Proton-Promoted Slowest
FeOOH + 3 H+ → Fe(III)(aq) + 2 H2O Ligand-Promoted (FeOOH)>FeIII-OH + L- + H+ → (FeOOH)>FeIII-L + H2O → FeOOH + Fe(III)-L(aq) Examples: L- = oxalate, malonate, citrate, Reductive FeOOH + e- + 3 H+ → Fe(II)(aq) + 2 H2O Examples: see below for manganese. Synergistic
Fastest
(FeOOH)>FeIII-L + e- → FeOOH + Fe(II)(aq) + LExamples: L- = oxalate, malonate, or citrate; e- = ascorbate or Fe2+
Manganese (III, IV) Oxide Dissolution
Reductive MnOOH + e- + 3 H+ → Mn(II)(aq) + 2 H2O MnO2 + 2 e- + 4 H+ → Mn(II)(aq) + 2 H2O Examples: e- = ascorbate, hydroquinone, dithionite (S2O42-), H2S, pyrogallol
Table 4.
Table 5.
4.21
-
OOC(CH2)4COO-
4.34
-OOC(CH2)3COO-
Glutarate (3)
Adipate (4)
4.42
2.85
1.25
pKa1
OOC(CH2)2COO-
-
OOC(CH2)COO-
-
OOCCOO-
-
Formula
Succinate (2)
Malonate (1)
Oxalate (0)
Ligand
5.64
5.43
5.42
5.70
4.27
pKa2
5.0
5.0
5.0
5.0
5.0
pH
< 2.8 × 10-12
3.3 × 10-11
< 2.8 × 10-12
1.4 × 10-11
1.0 × 10-10
Rate (mol m-2 s-1)
< 4.1 × 10-7
4.8 × 10-6
< 4.1 × 10-7
2.0 × 10-6
1.5 × 10-5
Rate Constant, k (s-1)
N/A
7200 ± 700
2700 ± 300
3000 ± 300
30000 ± 3000
Langmuir Binding Constant, K (M-1)
Figure 1.
Figure 2.
Figure 3.
Figure 4.
1.0
Oxalate ~ = 400 cm-1 0.8 Dn Bidentate nas(CO2) 0.6 Structure 1711 cm-1 0.4
nas(CO2) 1701 cm-1
ns(CO2) 1305 cm-1
0.2 0
Malonate ~ -1 0.8 Dn = 282 cm Bidentate 0.6 Structure
nas(CO2) 1631 cm-1
0.4
ns(CO2) 1349 cm-1
0.2
Arbitrary Units
0
Succinate ~ = 219 cm-1 0.8 Dn Monodentate 0.6 Structure 0.4
nas(CO2) 1628 cm-1
nas(CO2) 1547 cm-1 ns(CO2) 1409 cm-1
0.2 0
nas(CO2) Glutarate ~ = 219 cm-1 1632 cm-1 0.8 Dn Bidentate 0.6 Structure 0.4 0.2
ns(CO2) 1450 cm-1
0
Adipate n (CO2) ~ = 221 cm-1 as 0.8 Dn 1635 cm-1 Monodentate 0.6 Structure 0.4
nas(CO2) 1540 cm-1 ns(CO2) 1404 cm-1
0.2 0
Figure 5.
1950
1850
1750 1650 1550 1450 Wavenumber (cm-1)
1350
1250
O
OH
n=0
Fe
O C
+ OH2+
-
O
O
C
C
+ H3O+
Fe
O-
O
C
O
O
O
n=1
Fe
O
O
OH
C
+ OH2+
-
CH2
O
C
C
Fe
O
O
OH
+ 2H2O
CH2 C O -
O
O C
n=2
O
O
OH
C
OH2+
(CH2)2
C
-
-
+ H2O
Fe OH2+
OH
O
C O
O
+
Fe
(CH2)2
O
n=3
Fe
O
O
OH
C
+ OH2+
-
(CH2)3
Fe
C
CH2 CH2
O
OH
O
C
O
C
+ 2H2O
CH2
O -
O
O C
n=4
Fe OH2+
Figure 6.
O C
+ -
O
C
O
O
OH
(CH2)4
(CH2)4
C
-
+ H 2O
Fe
OH
O
OH2+
B
A
1 µm
Figure 7.
1 µm
5.00
3.00
As-O Shell 1.70 Å
Mn-Mn = 2.86 Å
As-Mn Shell 3.22 Å
As(III)-MnO 2
Mn4+
Mn4+
As
As
n=
As(V)-MnO 2
3.2 2Å
2.00
1.00 As(V) Solution 0.00 0.00 1.00 2.00 3.00 4.00 5.00 6.00 Distance (Å)
Figure 8.
Bidentate Binuclear Surface Complex >(MnIVO)2AsVOOH
-M
Fourier Transform Magnitude
4.00
Data Fit
= Oxygen
