Environ. Sci. Technol. 1904, 18, 450-456 Van Miller, J. P.; Lalich, J. J.; Allen, J. R. Chemosphere 1977, 6, 537-544. Gupta, B. N.; Vos, J. G.; Moore, J. A.; Zinkl, J. G.; Bullock, B. C. E H P Environ. Health Perspect. 1973, 5, 125-140. Blaser, W. W.; Bredeweg, R. A.; Shadoff, L. A.; Stehl, R. H. Anal. Chem. 1976,48, 984-986. Buser, H. R. J . Chromatogr. 1975, 107, 295-310. Buser, H. R.; Bosshardt, H. P. J. Assoc. O f f . Anal. Chem. 1976,59, 562-569. Rappe, C.; Buser, H. R.; Bosshardt, H. P. Chemosphere 1978, 7, 431-438. Buser, H. R.; Bosshardt, H. P. J . Chromatogr. 1974, 90, 71-77. Yamagishi, T.; Miyazaki, T.; Akiyama, K.; Morita, M.; Nakagawa, J.; Horii, S.; Kaneko, S. Chemosphere 1981,10, 1137-1144. Vos, J. G.; Koeman, J. H.; Van der Maas, H. L.; ten Noever de Brauw; de Vos, R. H. Food Cosmet. Toxicol. 1970,8, 625-633. Carreri, V. In ”Dioxin: Toxicological and Chemical Aspects”; Cattabeni, F.; Cavallaro, A.; Galli, G., Eds.; Spectrum Publications, Inc.: New York, 1978; pp 1-4. Baughman, R. W.; Meselson, M. Environ. Health Perspect. 1973,5, 27-35. Carter, C. D.; Kimbough, R. D.; Liddle, J. A.; Cline, R. E.; Zack, M. M.; Barthel, W. F.; Koehler, R. E.;Phillips, P. E. Science (Washington,D.C.) 1975, 188, 738-740. Olie, K.; Vermeulen, P. L.; Hutzinger, 0. Chemosphere 1977, 6,455-459. Buser, H. R.; Bosshardt, H. P.; Rappe, C. Chemosphere 1978, 7, 165-172. Eiceman, G. A,; Clement, R. E.; Karasek, F. W. Anal. Chem. 1979,51, 2343-2350. Cavellaro, A.; Bandi, G.; Invernizzi, G.; Luciani, L.; Mongini, E.; Gorni, A. Chemosphere 1980, 9, 611-621.
Bumb, R. R.; Crummett, W. B.; Cutie, S. S.; Gledhil, J. R.; Hummel, R. H.; Kagel, R. 0.;Lamparski, L. L.; Luoma, E. V.; Miller, D. L.; Nestrick, T. J.; Shadoff, L. A,; Stehl, R. H.; Woods, J. S. Science (Washington, D.C.) 1980, 210, 385-390. Chem. Eng. News 1979,57 (7), 23-29. Kimble, B. J.; Gross, M. L. Science (Washington, D.C.) 1980,207, 59-61. Junk, G. A.; Richard, J. J. Chemosphere 1981, 10, 1237-1241. Buser, H. R. Chemosphere 1979,8, 415-424. Buser, H. R. J . Chromatogr. 1975, 114, 95-108. Buser, H. R.; Bosshardt, H. P.; Rappe, C.; Lindahl, R. Chemosphere 1978, 7, 419-429. Kearney, P. C., presented at the 2nd International Workshop on Chlorinated Dioxins and Related Compounds, Arlington, VA, Oct 25-29, 1981. Dougherty, R. C. Biomed. Mass Spectrom. 1981,8,283-292. Robbins, J. A.; Edgington, D. N. Geochim. Cosmochim. Acta 1975, 39,285-304. ACS Committee on Environmental Improvement Anal. Chem. 1980,52, 2242-2249. Chrisp, C. E.; Fisher, G. L.; Lammert, J. E. Science (Washington, D.C.) 1977, 199, 73-75. “Compilation of Air Pollution Emission Factors”, 2nd ed.; U.S. EPA: Washington, DC, 1973. Hites, R. A,; Laflamme, R. E.; Farrington, J. W. Science (Washington, D.C.)1977, 198, 829-831. Miner. Yearb., U.S. Bureau of Mines, 1870-1980. “Production and Sales of Synthetic Organic Chemicals”; U.S. Tariff Commission: Washington, DC, 1919-1980.
Received for review J u n e 13, 1983. Accepted October 28, 1983. T h i s work was supported by the U.S. Department of Energy (Grant 8OEV-10449).
Reduction and Dissolution of Manganese( I I I ) and Manganese(I V ) Oxides by Organics. 1. Reaction with Hydroquinone Alan T. Stone* and James J. Morgan W. M. Keck Laboratories of Environmental Engineering Science, California Institute of Technology, Pasadena, California 9 1125
Mn(1V) form sparingly soluble oxide/hydroxide solid phases, while Mn(I1) is soluble. For this reason, dissolution solubilized by reduction in anoxic waters are poorly unreactions 1 and 2 greatly enhance the mobility of mangaderstood. A study of the reduction and dissolution of manganese oxide suspensions by hydroquinone was unMnOOH(s) + 3H+ e- = Mn2+ 2Hz0 (1) dertaken to determine the rate and mechanism of the MnOz(s) 4H+ + 2e- = Mn2+ + 2H20 (2) solubilization reaction. Dissolution of the manganese(II1,IV) oxide suspension by hydroquinone in the pH range nese in natural systems. Although oxygenation of Mn2+ 6.5 < pH < 8.5 is initially described by the following emhas been studied extensively (I-3), reduction and dissopirical rate law: lution reactions are poorly understood. The purpose of d[Mn2+]/dt = k1(H+}0~46[QH2]1~0([MnO~]o - [Mn2+]) this work is to systematically explore the factors that influence how quickly manganese oxides are reduced and where [Mn2+]is the dissolved manganese concentration, dissolved under natural conditions. [QH,] is the hydroquinone concentration, and [MnO,lo is Natural organic compounds have been found to reduce the amount of manganese oxide added. The apparent a variety of inorganic species and are the most readily activation energy was found to be at +37 kJ/mol. Calcium available reductants in most natural systems. Soil fulvic and phosphate inhibited the reaction, by adsorbing on the acids have been shown to reduce Hg(I1) to Hg(O), Fe(II1) oxide surface. A model is proposed for the observed rate to Fe(II), and I2to I- (4). A number of studies have shown dependence, according to which complex formation bethat organic compounds with structures similar to natural tween hydroquinone and manganese oxide surface sites organics reduce and dissolve manganese oxides (5-7). occurs prior to electron transfer. Generation of radicals by reaction of manganese dioxide with hydroquinone was studied by Fukuzumi et al. (8,9) Introduction and Ono et al. (IO). The rate of semiquinone radical Within the pH range of natural waters, Mn(II1) and formation was found t o be first order with respect t o hydroquinone concentration and initial manganese dioxide *To whom correspondence should be addressed at the Departloading. A mechanism was proposed that involves hyment of Geography and Environmental Engineering, The Johns drogen atom abstraction from hydroquinone (IO). Hopkins University, Baltimore, MD 21218.
w The chemical processes by which manganese oxides are
+
+
+
450
Environ. Sci. Technol., Vol. 18, No. 6, 1984
0013-936X/84/0918-0450$01.50/0
0 1984 American Chemical Society
54Mn-Labeledmanganese oxides were recently used by Sunda et al. (11) to measure the rate of reduction and dissolution by natural marine organic compounds. The reaction was found to be photocatalyzed. The structures and functional groups of natural organics differ considerably depending upon their source. Humic substances collected from marine surface waters are primarily nonaromatic in character; it has been recently proposed that marine humics are primarily degradation products of fatty acids (12). Freshwater humics, in contrast, are more aromatic in character and contain a core structure of phenols and phenolic acids (13). The ability of organics in the different structures to reduce manganese oxides is expected to differ considerably. This is the subject of a future paper. Reductant molecules must be close to manganese oxide surface sites in order for reduction of the manganese oxide to take place. It is therefore expected that the rate of dissolution will be determined by how quickly surface reactions proceed, unless the reaction is fast enough to be diffusion controlled. The influence of chemical parameters on the reaction rate must therefore be addressed in terms of their influence on surface chemical reactions.
Experimental Methods Unless otherwise stated, all solutions were prepared from reagent grade chemicals and deionized-distilled water (DDW) and filtered with 0.2-pm pore diameter membrane filters (Nucleopore Corp.) before use. All glassware was soaked in 5 N HNOBand thoroughly rinsed with DDW prior to use. Suspension Preparation. The preparation and characterization of the manganese oxide suspensions are summarized here; this is discussed in greater detail by Stone (14). A 2 X M NH40H solution was prepared with DDW and purged with nitrogen to remove oxygen. The solution was rapidly stirred, and enough stock MnC1, solution was added to form a suspension of 5 X lo4 mol/L Mn(OH),(s). A few minutes after the formation of the suspension, the purge gas was switched to pure oxygen, which was maintained for 100 min. Only minor amounts of settling of the dark brown suspension occurred over a period of several months. The particle size distribution was estimated by successive filtration using 0.1, 0.2, 0.4,0.6,0.8, and 1.0 Km pore diameter X 45 mm diameter membrane filters (Nucleopore Corp.). Within 10 days of preparation, most particles were between 0.2 and 0.4 pm in size. Over a period of several months, particles grew to over 1.0 pm in size, presumably by coagulation. Total manganese in suspension was determined by atomic absorption spectrophotometry (AAS) or by complexometric titration with ethylenediaminetetraacetic acid (EDTA) (15). Oxidizing titer was measured by using the Leucocrystal Violet method (16) and using permanganate titration (17). Table I lists values determined for three separate preparations. On the average, there are 1.3 equiv of oxidant/mol of manganese. X-ray diffraction patterns of oxide collected by filtration contain lines of the mineral phase feitknechtite (pMnOOH(s)) and considerably fainter lines of manganite (7-MnOOH(s)). IR spectra of the manganese oxides preparations are presented in Stone (14). The BET surface area of filtered and dried preparation N(8) was found to be 58 m2/g. The pH, (zero proton condition) values of suspensions N(7) and h(9) were determined by acidimetric titrations to be 7.4 and 6.9, respectively. Dissolution Experiments. All stock solutions were filtered with 0.2-pm pore diameter membrane filters
Table I. Characterization of Manganese Oxide Suspensions suspension N(7) N(8) N(9)
age, days 10 95 204 3 88 197 66
MnT: M 5.12 X 4.96 X 4.60 X 5.10 X 4.94 X 4.69 X 4.61 X
[oxid],* N
[oxid]/ MnT
stoichiometry
6.84 X 6.32 X 6.23 X 6.45 X 5.93 X 5.74 X 10-4f 5.90 X 104f
1.337 1.274 1.356 1.265 1.200 1.223 1.280
Mn01,67 MnO164 MnOl,6s Mn01,63 Mn01,60 Mn01,61 MnO,,,,
"MnT = total moles of Mn per liter of suspension. [oxid] = oxidizing titer in equivalents per liter. Determined by using atomic absorption spectrometry. Determined by using the Leucocrystal Violet colorimetric test. e Determined by using complexometric titration with EDTA. f Determined by using redox titration with standard oxalate and KMnOl solutions.
(Nucleopore Corp.) prior to use. Stock of hydroquinone (QH,; Sigma Chemical Co.) and 2,5-dihydroxybenzoicacid (2,5-DiOH;Aldrich Chemical Co.) solutions were prepared slight acidic and used within 24 h to minimize oxidation by oxygen. A carbon dioxide/bicarbonate buffer was used to maintain constant pH. Solution alkalinity was imposed by adding sodium bicarbonate. Solutions were constantly bubbled with a 1% carbon dioxidelbalance nitrogen gas mixture. This procedure kept the dissolved oxygen concentration to a minimum, as well as the pH constant. Reaction solutions were prepared by adding DDW and stock sodium nitrate and sodium bicarbonate solutions by pipet to a dried, 1-L jacketed beaker. Once stock manganese oxide suspension and organic substrate had been added, all solutions contained 5.0 X lo-, M sodium nitrate, unless otherwise noted. Stock manganese oxide suspension was added to each jacketed beaker, which were then sealed with a plexiglass faceplate. Suspensions were stirred with magnetic stir bars and bubbled with 200 mL/min of the carbon dioxide/ nitrogen mixture at least 1 h prior to the addition of organic substrate and throughout the dissolution reaction. Stock QH, and 2,5-DIOH were added by pipet. Suspension particles experienced an increase in ionic strength and a decrease in pH upon addition to the solution medium, causing flocculation. Experiments therefore measured the dissolution rate of freshly flocculated suspensions. Dissolved manganese ( [Mn2+]diss) and solid manganese oxide ([MnO,]) were distinguished by filtering aliquots of reaction solution with 0.2 pm pore diameter X 25 mm membrane membrane filters (Nucleopore Corp.). Before use, each membrane filter was rinsed once with 10 mL of 0.10 M HN03 and 3 times with 10 mL of DDW. Nitric acid was added to each filtered aliquot to make a 0.10 M solution and [Mn2+]diBs measured by using atomic absorption spectrophotometry (AAS). For reaction solutions of pH greater than 7.0, filtered aliquots contained less than 3% of manganese added as oxide suspension, prior to the addition of organic substrate. Separate experiments determined the amount of orthophosphate adsorbed by the manganese oxide suspension. Phosphate stock solution (made from KHzPOI and K2HP04)was added to the oxide suspension in 5.0 x M NaNO,, equilibrated with the carbon dioxide/nitrogen gas mixture, and then filtered with 0.2 Km X 45 mm membrane filters (NucleoporeCorp.). Filters were placed in 50-mL beakers contained 10 mL of 1.0 x lov2M ascorbic acid and sonicated. This treatment completely dissolved Environ. Sci. Technol., Vol. 18, No. 6, 1984 451
the oxide suspension. The clear solutions were then diluted and analyzed for orthophosphate by using the Molybdenum Blue Technique (18). Results and Discussion Preliminary experiments examined how effectively the filtration technique separates reduced, dissolved manganese from oxidized, particulate manganese. During one dissolution experiment, aliquots filtered with 0.1- and 0.2-pm membrane filters were taken shortly after one another and their manganese concentrations compared. The 0.2-pm filter retained 90% or more of the manganese retained by the 0.1-pm filter during the first 75% of the reaction. Thus,the filtration technique using 0.2-pm filters underreports the amount of remaining oxide, but only by 10% or less. fi2+, produced by the reaction with hydroquinone, may be adsorbed on the remaining oxide and therefore retained by the filter. To check this possibility, suspensions were prepared containing 2.8 X M oxide suspension, 3.0 X 10” M MnC12,and 5.0 X M NaN03. At either pH 7.20 or pH 7.91 (using the carbon dioxide/bicarbonate buffer) loss of Mn2+by adsorption was less than 2% of the amount added. The effect of laboratory lighting on the reaction rate was examined by performing duplicate experiments under fluorescent laboratory lights and under darkroom lights. Rates of dissolution were the same. Transport-controlled reactions should show a dependence on stirring rate. Increasing the stirring rate from 100 to 1100 rpm did not increase the rate of dissolution. In order to measure the effect of dissolved oxygen, duplicate experiments were performed with an oxygen-free carbon dioxidelnitrogen gas mixture and one containing 22% oxygen. No difference in the dissolution rate was observed. The reactivity of a manganese oxide suspension may change over the several months in which it is used in experiments. Changes upon aging may be due to gradual recrystallization of the oxide or lowering of oxide surface area caused by coalescence of particles. Duplicate dissolution experimentswere performed on the same suspension after 19, 33, 63, and 112 days of aging. The dissolution rate decreased 40% during this period, or by approximately 0.4% per day. For this reason, the relative suspension age must be considered when experiments performed on different days are compared. Order with respect to Suspension Loading. Experiments described in this section were performed with a 15-fold excess of hydroquinone reductant, and therefore, the hydroquinone concentration can be considered constant. [Mn2+]dissis an operational definition of dissolved manganese which does not necessarily reflect the amount of manganese oxide that has been reduced, only the amount of manganese not retained by the 0.2-pm membrane filter. Replacing hydroquinone by an equal concentration of catechol results in an 8-fold increase in the dissolution rate (discussed in Stone (14); unpublished results). The dissolution rate is controlled, therefore, by the rates of surface chemical reactions and not by diffusion. The lack of a dependence on stirring rate and the magnitude of the apparent activation energy support this conclusion. The rate of dissolution depends upon the manganese oxide suspension loading ([MnO,lo) as shown in Figure 1. By use of the method of van’t Hoff (19), the log of the initial dissolution rate plotted vs. the log of the suspension loading gives a line having a slope equal to the order. A slope of 1.03 is derived from the experiments shown in Figure 1, indicating that the initial dissolution rate is first 452
Environ. Sci. Technol., Vol. 18,No. 6, 1984
5
0
20
15
IO
Time (Minutes )
Flgure 1. Effect of varying the manganese oxide suspension loading. Reaction conditions: 4.29 X lo4 M QHp,5.00 X lo-* M NaNO,, paH 7.7, Pco2 = 1.1 X lo-’ atm, and 1.0 X M alkalinity, 25 ‘C. 0,
I
I
I
1
0 I .44 X ~ 2 . 1 0 2.07 X o 3 . 5 1x V 4.2 I X
I
[
M MnOx]O 100 ~ -5~ lO-5M I O - 5 ~ 0-5M
0
0 0
A A A
A
0
0 0
0 0
0
0
V n
V -2.5
I
0
I
I
I
I
I
5
IO
15
20
25
Time
( Minutes)
Figure 2. Equation 5 is used to fit experimental data from different suspension loadings (Figure 1). Sets of points from separate runs have been offset for sake of comparison.
order with respect to suspension loading. This is the result expected for surface-dependent reactions, since the oxide surface area is proportional to the amount of suspension added. The amount of oxide remaining at time t is the amount originally added minus the amount of manganese dissolved:
The initial dissolution rate is proportional to the amount of oxide loading, which suggests that the rate at any time during dissolution is proportional to the amount of remaining oxide:
where kexpdis the experimentally determined rate constant. Integration of eq 4 gives In (tMnO,lo - [Mn2+ldiss)- In ([MnO,lo) = -kexptlt
(5)
If this equation is valid, a plot of In ([MnO,lo - [Mn2+]&,) vs. time should give a straight line. Figure 2 plots the data from Figure 1 in this manner. Until about 50% of the oxide is dissolved, the plots are quite linear. After this
Table 11. Results of Experiments Used To Find Order with Respect to [QH,]
trial M
[MnOzlot M 2.89 x 10-5 2.89 x 10-5 2.87 x 10-5 2.87 x 10-5 2.87 x 10-5 2.10 x 10-5 2.10 x 10-5
PH
R
7.80 7.83 7.71 7.81 7.82
Y QP
7.72 7.71
L U
Q
[QHzI,M 8.57 x 10-4 5.71 x 10-4 4.29 x 10-4 2.86 x 10-~ 2.86 x 10-4 4.29 x 10-4 2.14 X 10"'
[QHzI/ [MnOxlo
min-'
29.7 19.8 14.9 10.0 10.0
9.93 x 10-2 7.02 X 6.09 X 3.72 X lo-' 3.55 x 10-2
4.61 X 3.37 x 2.57 X 1.75 X 1.68 X
lo2 102 lo2 lo2 lo2
53 53 66 63 63
20.4 10.2
5.34 x 10-2 2.57 X lo-'
2.26 X lo2 1.08 x 102
75 96
age,
kP, min-1.M-0.46
kexptl,
days
Table 111. Results of Experiments Used To Find pH Dependencea
set
trial
pH
[MnOxJo, M
[QHz],M
[alkalinity],M
1
QN
6.49 6.89 7.29 7.83
2.10 X 10" 2.10 X lo-' 2.10 x 10-5 2.10 x 10-5
2.14 X lo-' 2.14 X lo4 2.14 x 10-4 2.14 x 10-4
6.98 7.29 7.71
2.87 2.87 2.87
7.87 8.25 8.58 9.06
2.81 x 10-5 2.81 x 10-5 2.81 X 1.87 X
QE QO
QP
v
2
W U QD
3
&C QB T
X
X X
pco,,
age,
kexptl,
atm
min-l
days
C1.0 X 9.8 x 10-4 3.6 X 1.0 X lo-'
1.1X lo-' 1.1 x 10-2 1.1 X 1.1 X lo-'
8.56 X lo-' 7-68 x 10-2 4.88 X lo-' 2.57 X lo-'
96 84 96 96
4.29 X 4.29 X 4.29 X
9.6 X 10"' 3.0 X 1.0 X
1.1 X 1.1 X lo-' 1.1 X
1.39 X lo-' 1.04 X lo-' 6.09 X lo-'
69 69 69
4.29 x 10-4 4.29 x 10-4 4.29 X 10"' 4.29 X lo4
2.9 x 10-3 5.7 x 10-3 1.0 x 10-2 1.0 X lo-'
2.1 x 10-3 2.1 x 10-3 2.1 x 10-3 3.2 X
8.52 x 5.33 x 3.67 x 2.48 X
79 79 79 66
10-2 10-2 10-2
lo-'
aResulta of linear regression: set 1,slope = -0.44, corrl = 0.962; set 2, slope = -0.49, cord = 0.997; set 3, slope = -0.45, corrl = 0.995. 2.7)
I
I
-0.8
I
2.6 -
I
I
I
I
I
I
-,.2 -I
-
2.5 -
-
a 2.4
-
-
1 1
0 01
2.3-I
.4-
-I
.6 -
2.2 -
2.1 -
-3.8
-3.6
-3.4
-3.2
-3
- I .8.
6
I
I
I
I
I
I
6.5
7
7.5
8
8.5
9
5
L0910[QH2]
Figure 3. Determinatlon of the order of the reactlon wlth respect to hydroquinone. The order is found from the slope of log k, plotted vs. log [OH,].
time, points curve progressively upward, indicating deviation from eq 5. Alteration of the surface area and chemical characteristics of the oxide by extensive dissolution are probably responsible. Values of the experimental rate constant (kexpu)reported in this work were found by fitting a line through points collected during the first 50% of the reaction plotted as In ([MnO,lo - [Mnz*]dia,)vs. time. Subsequent changes in dissolution rate will not be addressed. Kexptlvalues calculated on duplicate runs differ by less than 5 % . Second-order rate constants (k,) are calculated when hydroquinone is not in excess by assuming that hydroquinone
Figure 4. Effect of pH on the experimental rate constant, kexpfl. Reaction conditions for the three sets are given in Table 111.
is a two-equivalent reductant. Experimental Rate Law. The order of the reaction with respect to hydroquinone was calculated from experiments summarized in Table 11. Values of the log of the experimental rate constant corrected for pH (k,) were calculated by using eq 5 and 6 and plotted vs. the log of the hydroquinone concentration (QH,). This is shown in Figure 3. The slope, calculated for successive pairs of points, is the order. The calculated order is close to 1.0 for all but the pair points a t the highest hydroquinone concentrations, for which the order is 0.8. The order calculated at higher hydroquinone concentrations may be still smaller. Environ. Scl. Technol., Vol. 18, No. 6, 1984
453
-I H
In
P
Time
( Minutes)
Flgure 5. Dissolution of manganese oxide suspensions by p-benzoquinone. p-Benzoquinone concentration: (0)7.14 X lo4 M; (A)2.00 X M; (0)5.00 X M; (0) 1.00 X M. Conditions: 2.85 X M [MnO,],, 5.00 X lo-' M NaNO,, 1.0 X lo-' M alkalinity, paH 7.8,and Pco, = 1.0 X lo-* atm, 25 "C.
The influence of pH on the dissolution rate is illustrated in experiments summarized in Table 111. Figure 4 plots the log of kexptl vs. paH (activity scale); lines are drawn between points from experiments performed under similar conditions. The three groups of data do not plot on the same line because Pco2,[QH,], and suspension age differ. The dissolution rate may be a complex function of pH. Linear regression of the values in Table 111 gives a fractional order with respect to [H+)of 0.46. This empirical result can be used to express the influence of pH on kexptl (eq 6). k, is the experimental rate constant corrected for changes in paH.
0
454
Environ. Sci. Technol., Vol. 18,No. 6, 1984
1
1
-3
-2
-I
L o g t o (Cone ) i M )
Figure 6. Effect of calcium and phosphate on kexpt, for dissolution in M) hydroquinone. Conditions: 2.13 X M excess (2.14X [Mn0,la, 5.00 X lo-' M NaNO,, and Pwp = 1.0 X atm, p*H 6.80, 25 'C.
Table IV. Adsorption of Phosphate and Its Effect on the Dissolution Rate" trial AH AC AD AE AG
AF
kexptl = k,(H+]0.46
(6) p-Benzoquinone is the most likely oxidation product of the reaction of hydroquinone with manganese oxide. UV spectra of filtered aliquots confirmed that it was the principal product (see Stone (14)). p-Benzoquinone can itself dissolve the manganese oxide suspension, but at a rate much slower than hydroquinone, as shown in Figure 5. Dissolution by p-benzoquinone is an autocatalytic reaction; the rate of dissolution increases with time during early stages of the reaction. The amount of p-benzoquinone generated during reactions of manganese oxide with hydroquinone was not sufficient to influence the reaction rate. Mn2+is also a reaction product, and an experiment was performed to observe its effect on the reaction rate. Two experiments were performed that were identical except for the addition of 9.3 X lo4 M Mn2+to one suspension prior to the addition of hydroquinone. Initial rates of dissolution in the two experiments were the same, indicating that the additional Mn2+ did not influence the reaction rate. The apparent activation energy of the reaction was deM [MnO,lo, 2.1 termined under conditions of 2.1 X X M QH,, 5.00 X M NaNO,, 1.0 X M alkalinity, and a p8H of 7.0, at temperatures between 5 and 25 "C. By use of an Arrhenius plot, the apparent activation energy was found to be +37 kJ/mol, sufficiently high to rule our transport control of the reaction rate. This value is close to the activation energy of +33 f 2 kJ/mol determined by Ono et al. (10) for the reaction between manganese dioxide and hydroquinone. The effect of changes in ionic strength on the dissolution rate by hydroquinone (at pH 6.9) and 2,5-dihydroxybenzoicacid (at pH 6.6) was quite small (14). The preceding experimental information can be combined to give the following empirical rate law: d[Mn2+Idi8,/dt= k,{H+]0.46[ QHz]l~O([MnO,]o - [Mn2C]di88)*.o (7)
1
-4
X
[phosphate]~, [phosphate),d,, mol/L mol/L
0.00 5.00 X lo4 1.00 x 1.00 X 3.00 X 1.00 X lob3
0.00 2.29 X 3.15 x 8.17 X lo-' 1.23 X low7 1.73 X lo-'
kz(X), L/(mol.s)
k2(X)/ kz(0)
4.39 x 3.42 X 3.17 x 2.24 X 1.78 X 1.33 X
0.779 0.722 0.510 0.405 0.303
102
lo2 lo2 lo3 lo2 lo2
Conditions: 2.00 X M QH2,2.89 X M [MnO,lo, 5.00 M NaNOR,paH 6.80, and Pen, = 1.0 X 25 O C .
kl = 8 X lo3 ( L / m ~ l ) l . * ~ ( l /at s ) 25 , "C, I = 5.0
M X NaN03. The rate law can also be expressed in terms of oxide surface area by using the BET surface area measurement: d [Mn2+]di8s/ dt = kA(H+}o.46[ QH2]l.oAMnO,
k~ = 1 X 10' ( L / m ~ l ) ~ . ~ ~ ( L / r n ~at) (25 l / s"C, ) I
(8)
= 5.0 X M NaN03. AM,Oz is the oxide surface area per liter, in the units of square meters per liter. Inhibition by Calcium and Phosphate. Specifically adsorbing cations and anions may lower the reaction rate by blocking oxide surface sites or by interfering with the release of Mn(I1) into solution. A series of experiments were performed in which stock calcium nitrate or potassium phosphate solutions were added to the oxide suspension and equilibrated with the carbon dioxide/nitrogen mixture prior to the addition of organic reductant. kexp(X)/kexp(0) is the experimental rate constant determined in the presence of a specified amount of calcium or phosphate divided by the rate constant measured in their absence. At the same pH and adsorbate concentration, phosphate inhibits the reaction much more dramatically than calcium, as shown in Figure 6. In M phosphate at pH 7.68, the dissolution rate is only 25% of the rate in the absence of phosphate. Inhibition by phosphate was slightly greater at pH 7.68 than at pH 6.86. Phosphate inhibited the reaction with 2,5-DIOH to the same extent as with hydroquinone. The amount of inhibition at pH values of 6.75 and 7.65 were the same, however, for reaction with 2,5-DIOH (14). Inhibition experiments were also performed at much lower concentrations of hydroquinone. Results are summarized in Table IV. A different manganese oxide suspension (N(9))was for these experiments, so results cannot
lo-,
\
0.251
/
MI-
/
I'
I
c
3
I-. 0.50
boundary layer; (ii) formation of a surface complex between the adsorbate and oxide; (iii) charge transfer within the surface complex; (iv) desorption of oxidized organic substrate; (v) movement of reduced Mn(I1) from the crystal lattice to the adsorbed layer; (vi) desorption of Mn2+;(vii) diffusion of products away from the surface. The dissolution rate is increased when hydroquinone is replaced by catechol (14). Steps that do not involve the organic substrate are, therefore, not rate limiting. Surface Site-Binding Model. If either surface complex formation or charge transfer is rate limiting, then the rate equation will resemble an adsorption isotherm. This will be demonstrated by using a surface site-binding model based on a similar treatment by Castellan (24). Consider the following two reactions, in which hydroquinone (QHJ forms a surface complex with manganese oxide prior to electron transfer: =MnOH EMnQH
ka
kl + QH, r EMnQH + H 2 0 k-i
+
Mn2+(aq) oxidized hydroquinone
(9)
(10)
where =MnOH and =MnQH are surface species. S , is defined as the total number of surface sites at time t and BB the fraction of sites bound with hydroquinone. 8~ = [=MnQH]/S,
(11)
The dissolution rate is equal to the rate of reaction 10. d[Mn2+]/dt = k2[=MnQH] = k2S,BB
(12)
Any change that lowers the total number of surface sites or the fraction of sites bound to the substrate will lower the dissolution rate. S, can be considered constant for initial rates of dissolution. Applying a steady-state approximation (d[=MnQH]/dt = 0) to eq 9 and 10 yields a rate equation expressed in terms of bulk hydroquinone concentration:
The assumption is made that S, = [=MnOH] + [-MnQH]. If the rate of adsorption is slow relative to the rate of surface reaction, eq 13 is first order with respect to [QH,]. This is unlikely because of the dramatic effect that minor changes in reductant structure have on the dissolution rate (14).
It is more likely that adsorption and desorption reactions are fast relative to electron transfer. When this is true, eq 9 becomes a preequilibrium step (with PQN = kl/k-J and the rate equation becomes d[Mn2+l - ~ z S , ( K & ~ Q H , I ) -dt
K~H[QHzI +1
(14)
When surface coverage is low, eq 14 is first order with respect to [QH,]. The order drops to zero as the surface becomes saturated with organic substrate. A change in pH may change the proton level of hydroquinone-oxide surface complexes, as well as the total surface coverage by hydroquinone. Consider the following surface complexes: =MnQH2+, =MnQH, and EMnQ-. Equation 12 must be reformulated to account for different rates of electron transfer within each kind of complex. d[Mn2+]/dt = k2[=MnQHz+]+ k,q=MnQH]
+ k,'/[=MnQ-]
(15)
The overall pH dependence of the dissolution reaction is Environ. Sci. Technol., Vol. 18, No. 6, 1984
455
determined by the relative magnitudes of the terms in eq 15. If adsorption and desorption are fast relative to electron transfer, then the relative distribution of surface species can be modeled by using equilibrium expressions (25).
+ H+ KL =MnOH = =MnO- + H+ R Z =MnOH + QH, + H+ = rMnQHz+ + H 2 0 =MnOH2+= =MnOH
(16) (17)
KbH2 (18)
ZMnOH
+ QH,
= =MnQH
+ H20
K&
=MnOH + QH2 = =MnQ- + H+ + H 2 0
K&
(19) (20)
Values of the equilibrium constants are not known. =MnOH is the dominant unbound surface species at pH values near pHzpc,and its abundance varies little upon small changes in pH. Equation 15 can be rewritten by using eq 18-20: d[Mn2+]/dt = [aMnOH][QH2](lz2K~~,[H+] kz’K&
+
+ k,”K&/[H+]) (21)
According to this equation, the order with respect to [H+] is +1, 0, and -1 for reaction via complexes =MnQH2+, =MnQH, and =MnQ-, respectively. As dissolution occurs, the number of oxide surface sites changes. The ratio of available surface sites to remaining manganese oxide is defined as f(t) and may be a function of time. f(t) = St/[MnOxlt
(22)
In the special case where f(t) is a constant, S, is proportional to [MnOXlo- [Mn2+Idiss.The relation between the amount of remaining oxide and the dissolution rate observed in eq 7 may arise from this special case. Agreement with the Surface Site-Binding Model. The formation of hydroquinone/oxide surface complexes prior to electron transfer is postulated to account for (i) large differences in dissolution rate caused by small changes in reductant structure (14), (ii) inhibition of the reaction by adsorbing calcium and phosphate ions, and (iii) high rates of dissolution relative to possible rates of Mn(111) and Mn(1V) release from the oxide surface. Two limiting cases of eq 13 can account for the observed first-order dependence with respect to [QH,]. More research is needed to verify that the surface site-binding model described in this paper best represents the reaction on the oxide surface. In particular, measurements of surface coverage by hydroquinone should be made and compared directly to rates of dissolution. The empirical order with respect to (H+)of 0.46 lies between the results expected for reaction via =MnQH,+ and via =MnQH. This order may, therefore, arise from the sum of the terms given in eq 21. Protons may influence the reaction rate in other ways as well such as aiding the release of adsorbed Mn(I1) or oxidized organic substrate from the surface.
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Environ. Sci. Technol., Vol. 18, No. 6, 1984
Conclusions
Results show that the rate of dissolution of manganese oxide suspensions is not diffusion controlled and that reduction of Mn(II1) and Mn(1V) must occur on the oxide surface. The experimentally derived rate equation for the reaction is consistent with a surface site-binding model, in which hydroquinone forms a surface complex with manganese prior to electron transfer. Evidence suggests that the rate of the surface chemical reaction, rather than the rate of surface complex formation, is rate limiting. Inhibition by specifically adsorbing calcium and phosphate ions gives further support to the surface site-binding model. Registry No. MnOz, 1313-13-9; Mn203, 1317-34-6; QH,, 123-31-9.
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