Oxidation of arsenate(III) with manganese oxides in water treatment

War. Res. Vol. 29, No. 1, pp. 297-305, 1995

Pergamon

0043-1354(94)E0089-O

Copyright © 1994 Elsevier Science Ltd Printed in Great Britain. All rights reserved 0043-1354/95 $7.00 + 0.00

OXIDATION OF ARSENATE(III) WITH M A N G A N E S E OXIDES IN WATER TREATMENT WOLFGANGDRIEHAUS,REINER SEITH and MARTINJEKEL* Technical University of Berlin, Department of Water Quality Control, Secr. KF 4, StraBe des 17. Juni 135, D-10623 Berlin, Germany (First received November 1993; accepted in revised form March 1994) Abstract--Arsenate(Ill) is the more toxic form of inorganic arsenic and its removal from drinking water is less effective as compared to arsenate(V). Arsenate(III) persists in aerated water, even at high pH, but is easily oxidized by manganese dioxides. The oxidation of As(III) follows a second order rate law with respect to As(Ill). The reaction rate is effected by the initial molar ratio of MnO2 to As(Ill), the respective rate order is 1.5. Calcium reduces the rate constant by the order of -0.2, whereas the pH has no influence (pH 5-10). There is no desorption of reduced manganese in the batch tests at high initial molar ratios. Flow-through-tests in small sand filters, loaded with manganese dioxide, showed a decreasing oxidation of As(Ill) in the first 60 h. Afterwards the oxidation increased again and was quantitative after 10 days. It is concluded that the enhanced oxidation after this initial period is caused by bacteria, as some ubiquitous bacteria are involved in manganese oxidation. An application of this oxidation technique in drinking water treatment should be possible, because As(III)-oxidation is efficient and the release of soluble manganese is low. Key words--arsenic, speciation, oxidation, manganese oxide, loaded filter, drinking water

INTRODUCTION Arsenic is a contaminant of well known toxicity and carcinogenicity in the drinking waters of some regions (WHO, 1981; Korte and Fernando, 1991). Arsenic concentrations in groundwater are usually below 5 # g l -~, but geochemical mobilization, contaminated soils in industrial regions and the application of arsenic pesticides lead locally to enhanced concentrations. Removal processes are needed in case of arsenic concentrations exceeding the maximum contaminant levels in drinking water. The drinking water standards are mostly set to 0.05 mgl -~, but lowering of these standards is discussed due to new results on the chronic toxicity of arsenic. The drinking water standard in Germany is 0.04 mg 1-~, but will be reduced to 0.01 mg 1-l in 1996. All elimination techniques, i.e. adsorption on activated alumina and coprecipitation with ferric salts, are effective for arsenate(V) (H2AsO4-, HAsO 2- ), but fail in case of arsenate(III) (H3AsO3) (Frank and Clifford, 1986; Jekel, 1994). As there is no simple treatment for the efficient removal of arsenate(Ill), an oxidation step is necessary to provide acceptable results in arsenic elimination. Arsenate(Ill) is mostly a hydrophilic, neutral species below pH 9, due to its pKa~ of 9.3 and is not accessible to the major removal mechanisms of anion sorption and anion exchange. *Author to whom all correspondence should be addressed.

The standard redox potential of the arsenate(III)arsenate(V) system is 0.56 V. This results in redox potentials of 0.06 V (99% As 3÷) and 0.23 V (99% As s÷ ) respectively at pH 7, as calculated by the Nernst equation. Although arsenate(Ill) might be easily oxidized by dissolved oxygen, it persists in aerated waters due to slow oxidation kinetics. Therefore, an oxidative treatment of arsenate(III) waters has to be included for effective arsenic removal, employing oxidizing agents or suitable catalysts. Frank and Clifford (1986) examined the oxidation with chlorine and found positive and reliable results. However, chlorine is not permitted as oxidizing agent in some countries, e.g. Germany (TrinkwV, 1990). Ozone also oxidizes arsenate(Ill). Boeckelen and Niessner (1992) found a rapid and quantitative oxidation even at a dose of 0.1 mg 1- ~. Due to its high oxidation potential, ozone is not feasible for a specific oxidation of arsenate(Ill), but leads to side reactions with natural organic matter. Investigations of the catalytic oxidation with the powdered activated carbon (PAC) Oxorbon St (Lurgi Co.), applied for the oxidation of ferrous iron, showed its principle feasibility for arsenate(Ill) oxidation (Gottschalk et al., 1992). The oxidation kinetics were slow and the required amount of PAC was rather high. Thus, an effective and simple application in drinking water treatment is unlikely. For the oxidation of arsenate(Ill) in drinking water treatment, a mild, preferably selective oxidizing agent or catalyst is needed, but effective catalysts were not

297

298

WOLFGANGDRIEHAUSet al.

known until now. Manganese dioxides have been shown to be relative strong oxidants in the environment, controlling the mobility (and toxicity) of iron, cobalt, chromium, and arsenic as well as of natural and anthropogenic organics (Postma, 1985; Crowther et al., 1983; Eary and Ray, 1987; Fendorf and Zasocki, 1992; Oscarson et al., 1980; 1983; Stone and Morgan, 1984). The redox reactions of arsenate(III) with manganese(IV)-oxides and manganese(III)-oxides at neutral pH can be written: H3AsO3 + M n O 2 .¢, HAsO42- + Mn 2+ + H20 E °=0.67V

(1)

H3AsO 3 + 2 M n O O H + 2H ÷ ,~- H A s O 2+ 2 M n 2÷ + 3 H 2 0

E ° = 0 . 9 5 V.

(2)

Oscarson et al. (1983) found a first order reaction rate in the redox reaction of various MnO2-modifications with arsenate(Ill), but the initial rate was somewhat higher. The use of c~MnO2, which resembles the naturally occurring mineral birnessite, leads to a faster oxidation of arsenate(Ill), compared to the orand fl-modification of MnO2. This can be explained by the low cristallinity and layered structure of the 6-modification, which has easily available reaction sites with M n 4+ and Mn 3÷ in the interlayer of the solid. M o o r e et al. (1990) investigated the influence of pH and temperature on the oxidation of arsenate(Ill) by c~MnO2 and suggested a structural based reaction scheme. The reaction rates increased slightly with decreasing pH, whereas increasing temperatures effected only the initial rates in the first 50 min of the experiments. Oscarson et al. (1983) as well as M o o r e et al. (1990) found a substantial lack of release of divalent manganese ions during the oxidation of arsenate(Ill). They explained this effect by adsorption to the manganese oxide or by the formation of an As(V)-Mn(II)complex. Since manganese dioxides are capable of oxidizing arsenate(Ill), they were assumed to be applicable to water treatment in arsenic removal processes. The investigations of this study were performed using the c~-modification of MnO2, as its chemical properties in natural and laboratory systems are well known (Postma, 1985; Crowther et al., 1983; Fendorf and Zasocki, 1992; Stone and Morgan, 1984). The subjects of this paper are the kinetics of arsenate(Ill) oxidation and the influence of pH and calcium and the applications in coated filters. METHODS

6MnO 2 was prepared by synproportionation, similar to the procedure described by Murray (1974). 7 ml of a 1 M Mn(NO3) 2 solution were slowly added to 0.41 of a stirred 35 mM KOH solution, containing 4.2 mmol KMnO4. Each precipitate, containing approximately 1 g MnO2, was washed several times with demineralized water, centrifuged and stored in demineralized water in polyethylene bottles at

4°C. The X-ray diffraction analysis (Philips APD 1700, Cu Kct-radiation) of dried samples (80°C) gave broad, diffuse reflection peaks at 0.14/0.24/0.36/0.7l nm, which are characteristic for 6MnO2. The chemical composition was determined to K~7MnTO~4.3*3.7H20. The nominal oxidation degree was 1.92 +_0.03 (oxalic acidpermanganate-titration according to Hem, 1980). This corresponds to 12 20% of Mn 3+ in the 6MnO 2. The specific surface of the precipitate was 260 m 2 g-~ (EGME-method after Heilmann et al., 1965). The point of zero charge (PZC) was found to be below pH 3.5 by an alkalimetric titration, which is in agreement with Murray (1974). Stock solutions of As(III) and As(V) were obtained by Fluka (No. 11082) and Merck (No. 9939), respectively. For the oxidation and adsorption experiments l g of 6MnO2 was suspended in 0.51 of a 20 mM NaNO3 solution. Aliquots of this suspension were added to the batch tests. Each kinetic run was made in triplicate. The volume of the kinetic tests was 11 of a stirred 10 mM NaNO 3 solution, pH was adjusted with 0.1 M NaOH or 0.1 M HNO 3 and measured electrometrically with a glass electrode with liquid film diaphragma. The pH-drift during a 45 min experimental run was 0.4 units in maximum. An aliquot of the As(Ill) stock solution was added at the beginning of each kinetic run. The standard conditions of the kinetic studies were 92/~mol 1 ~6 MnO 2 and 6.7 #mol I-~ arsenate(Ill), giving a molar ratio of 14. Each run lasted for 45 min and 0.2/Lm filtered samples were taken at 5, 10, 20, 30, and 45 min and acidified to pH 2 with HCI to preserve arsenate speciation. Adsorption isotherm experiments were carried out by adding different amounts of As(llI) or As(V) to 0.09 mM suspensions of MnO:. pH was adjusted to defined values and corrected after 20h. After an equilibration time of 48 h, pH was measured and samples were filtered for determination of arsenic. The filtrates were analyzed for As(IlI) and total inorganic arsenic (AslII + V) by hydride generation atomic absorption with a Varian AA 400 spectrometer, equipped with a continuous flow hydride generator Varian VGA 76. All additions of chemicals were performed on-line to avoid any manual sample treatment. The procedure is described in Driehaus and Jekel (1992). Determination of manganese in solution was made by flame AAS (detection limit 0.02 mgl -~ or 0.4 #mol 1-~), or by graphite furnace AAS (practical detection limit 5 #g I i or 0.1 #mol 1-~). RESULTS AND DISCUSSION Adsorption o f arsenate(Ill) and arsenate(V) Although 6 MnO2 has a negative surface charge at pH > 4 , due to its pHpz c of <3.5, a significant adsorption of arsenate occurs. Adsorption isotherms are shown in Fig. 1. The curves are based on the Freundlich-isotherm with Freundlich exponents of 0.5-0.54 at pH 5.5 and 7.5. At pH 9.8, the corresponding value is 0.93, indicating a nearly linear dependence of the loading on the equilibrium concentrations of arsenic. The adsorbed species was assumed to be arsenate(V), as all initial arsenate(Ill) was oxidized during the equilibration time. A difference in adsorption of initial arsenate(III) and arsenate(V) was not observed, although Oscarson et al. (1983) got a minor adsorption of initial arsenate(V). Takamatsu et al. (1985) found increasing amounts of arsenate(V) adsorbed with the adsorption of Mn2+-ions. Manganese ions were produced during oxidation of arsenate(Ill), but did not enhance the adsorption of arsenate in the presently studied systems.

As(III)-oxidation with manganese oxides

three kinetic test runs in Fig. 2(b). Heterogenous reactions usually follow a first order or pseudo first order kinetic (Lasaga, 1981). But the measured decrease of arsenate(Ill) is also influenced by a fast initial adsorption. In addition, the reaction sites of MnO2 are not necessarily homogeneous and masking of reaction sites by the reaction products may reduce their number and activity. This results in a fast initial decrease, followed by a slower rate, which leads to a second order reaction with respect to As(Ill). Obviously, the reaction rate does not only depend on the initial 6 MnO2-concentrations, as expected for a surface controlled reaction. Run 2 and run 3 with similar 6 MnO2-arsenate(III)-ratios, but different initial concentrations, result in the same slopes and nearly identical values for the observed reaction constant k '. An empirical rate law for the studied system is thus given by

Oxidation kinetics

Reaction kinetics are described by a general rate law expression (Lasaga, 1981) d[A ] - k • [A ] ' , [B] ~ dt

(3)

for a reaction scheme: aA + b B ~ c C

+dD

where [A ] and [B ] are the activities of the reactands and the exponents ~ and fl give the reaction order with respect to reactands A and B. In contrast to elementary reactions, in heterogeneous, surface catalysed reactions the exponents ~t and fl are generally different from the stoichiometric factors a and b. For the reaction of arsenate(III) with manganese oxides [equation (1) and (2)], the depletion of As(Ill) should depend on both, As(Ill)- and MnOzconcentration. The decrease of the arsenate(III)-concentration and the formation of arsenate(V) during a kinetic run are shown in Fig. 2(a). The arsenate(Ill + V)-concentration remains constant after two minutes, while arsenate(Ill) is further oxidized and arsenate(V) is released into the solution. Manganese oxides adsorb arsenate very quickly. In this way, the strong, initial decrease of As(III) is an effect of adsorption and oxidation. Manganese concentrations in solution were not plotted, as they were determined to be always below 0.1 #moll -l during the experiment. The decline of arsenate(III) follows a second order reaction kinetic, as shown by the linearized plot of

0,1

A d s o r b e d As,

299

d[As nl ] - dt

= k * [AsUi] =

(4)

Integration o f [4] w i t h c~ = 2 leads to

[AsIII]-I _ [AslII]g' = k • t

(5)

Including the obvious influence of the initial molar ratio gives d[AsIIl]d----~- k ' * [Asm]= * ( [6MnOzlO[Asm]o J'~t~

where the reaction order • is 2 with respect to arsenate(III), and fl gives the exponent of the initial

mol per mol

0,01

0,001

0, 0 0 0 1

. 0,1

.

.

.

.

.

.

.

.

.

.

.

.

.

1

.

.

. 0

E q u i l i b r i u m conc., jJmol 1-1 Fig. 1. Adsorption of arscnate(III) and arsenate(V) on 6 MnO 2 at various pH-values (10 mmol l-~ NaNO 3, curve fit according to the Freundlich-isotherm). WR 29/I--T

(6)

300

WOLFGANG DRIEHAUSet al.

As, pmol 1-1

14 12 10

/

8

As(lll+V) O

As(Ill) 4

As(V) /

-i3

o

2'0

/C

1,6

(b

1, 4

)/ -

3'0

4'0

s'o

6O

Reoction time, min.

1/Co , I/pmol

1,2 1 0,8 0,6 0,4 0,2 0

0

10 '

2'0

40 3b Reaction time, min. '

50 '

60

Fig. 2. (a) Decrease of arsenate(IlI) and arsenate(III + V) and formation of arsenate(V) during a kinetic run. (b) Linearized second order rate plots of runs with different initial MnO2- and arsenate(III)concentrations at pH 7. Run 1:MnO2 115 #moll -I, As(III) 6.7 #tool l-I; run 2: MnO 2 57/~mol l -I, As(III) 6.7/~mol 1-~; run 3:MnO2 115 #moll -l, As(III) 13.3/xmoll -I. molar ratio, k ' is the observed rate constant in kinetic tests, where more than one parameter, in general c (AsIII), was thought to be not constant. All computations of the arsenate(III) decrease gave a second order fit and regression coefficients R 2 better than 0.95. Figure 3 shows a plot of log k ' vs the log of the initial molar ratio. The exponent fl was estimated by

linear regression to be 1.5 at initial ratios between 17 and 90. At ratios below 17 it changes dramatically to higher values. This might indicate a change in the reaction mechanism probably caused by a heterogenity of reaction sites in the structure of 5MnO2. The manganese oxide used in the present study contains some Mn(III) (12% of total manganese), as indicated by a nominal oxidation degree of 1.92.

As(IlI)-oxidation with manganese oxides Mn(III) and some Mn(IV) are located in the interlayers of the structure (Golden et al., 1986). Moore et al. (1990) suggested, that the interlayer manganese is easier available for redox reactions than the manganese(IV) in the octahedral sheets, which cover the interlayer. In the octahedral sheets, manganese(IV) is surrounded by six oxygen atoms, and this coordination is thought to be more stable. At a low initial molar ratio of &MnO2 to arsenate(Ill), most of the interlayer manganese would be reduced by arsenate(Ill) and the more stable manganese(IV) of the octahedral sheets might be consumed. Therefore, it is conceivable that the rate of the oxidation increases with the initial molar ratio, giving a reaction order of greater 1. On the other hand, the redox reaction proceeds very quickly with an excess of &MnO2. Such conditions would occur in filters for oxidative treatment of arsenate(Ill) containing water. Under those conditions, reaction kinetics should not be a limiting factor for the application of this technique. The very slow oxidation kinetics, reported by Oscarson et al. (1983) and Moore et al. (1990) are probably caused by using dried substrates of MnO 2, which exhibits undoubtedly reduced surface activities. Data from these authors are not directly comparable to the systems investigated in this study. For further kinetic studies, regarding the influence of pH, calcium, phosphate, and sulphate, concentrations of 92 # M &MnO2 and 6.7 # M arsenate(Ill) (initial molar ratio 14) were chosen, resulting in an

301

average reaction constant of 0.0211#mol -t min -~ and a variation coefficient of 21% (n = 11) at pH 7.

Influence of pH, calcium, phosphate, and sulphate As indicated by reaction [1], pH should not affect the redox reaction between arsenate(Ill) and MnO:, as no proton consumption occurs at pH > 7. At pH < 7, one proton is consumed. If the electron transfer is the rate determining step, an enhanced rate constant with decreasing pH would be expected. Results of kinetic runs in the pH-range of 4-10 in Fig. 4 show a minor influence of pH on the reaction rate. A definite influence of pH is only obvious at pH-values below 5. This results in a reaction order of zero with respect to H+-activity. For this surface catalyzed redox reaction, transport processes, e.g. diffusion and adsorption/desorption, rather than the electron transfer itself should be rate determining. The influence at a pH below 5 could be attributable to the behavior of reduced manganese. At the chosen molar ratio, soluble manganese is released into solution at pH 4, whereas at higher pH, its adsorption is almost quantitative. In experiments with an initial molar ratio of 3, the reaction rate was much higher at pH 6 than at 7, and manganese was released to a higher extend at pH 6. Adsorbed manganese masks likely reaction sites and diminishes the reaction rate. Though this, the adsorption behavior of manganese causes effects of pH. Calcium, the most frequent divalent cation in drinking water, reduces the oxidation rate. The

log k'

...~'°"" °~.o.,,,.-'°" ,:

"J~'*°°°'°"**'*

.~ ....... ".........slope: 1.53 -1

.4-

-2

o,a

i

1,2

i

1,4 log [MnO2]/[As,I]

i

1,6

i

118

Fig. 3. Log-log plot of k ' vs the initial molar ratio ([MnO2]o 57-460/zmoll -t, [As !11]o 1.3-13.3/~ tool 1- t ).

302

WOLFGANGDRIEHAUSet al.

log k' -1

-1, 2

-1, 4 + X

-1, 6 +

•

X

•

-1, 8 ¸

-2

3

~

~

~

~

~

~

1'0

pH Fig. 4. Log-log plot of k ' vs pH. Different symbols indicate independent experimental series.

results from six kinetic runs are illustrated in Fig. 5. [Ca 2+] was varied between 1 and 3 mM at pH 7 and between 0.3 and 1 mM at pH 6 and 8. The lower rate constants at pH 6 could be explained by the fact that

the kinetic runs at the different pH values are independent series and rate constants are not always fully comparable from one series to another. As shown for the initial molar ratio, a reaction order

log k' -1 pH 6 -1, 2 pH 7 )1(

pH 8

-1, 4

-1, 6

Y~

-1, 8

215

~ _

-2

-4

-3, 5

O, 196

-3

•

-2, 5

log [ca] Fig. 5. Variation of log k' with log [Ca2+].

-2

As(III)-oxidation with manganese oxides

303

///,"°"

0,25

/,////"

m

pH 6 0,2

D

0

E E

,/

pH 7

D

0,15

/,,"

/,//" /'

pH 8

-6

,/"

e~ ..Q

/'

•

0,1

rm

,,

•

,//

/'

(3 (-

/,/

El

D i m 1~3

/

,,'~: ,"

0,05

• •

• •m

,, 4"" m

•

rzl ,j"

Umm I u

0

0,05

I

0,1 0,15 Idn reduced, mmol

!

0,2

0,25

Fig. 6. Adsorption of reduced manganese in the presence of arsenate. MnO2 0.94 mmol 1-% As(Ill) 0.3 mmol 1-l. could also be computed for the initial molar concentration of calcium. The respective equation is given by k ' = const. * [Ca 2~]-°'2 (pH 6-8).

(7)

The oxidation kinetic is not influenced by phosphate (0.03-0.1 mM) and sulphate (3 mM) at pH 7. The adsorption of arsenate was not affected by phosphate, suggesting that there was no competing anion sorption. Release o f soluble manganese

Soluble Mn 2÷ is produced during the oxidation of arsenate(Ill). Determinations of manganese in the filtrates of short term batch tests showed only negligible amounts, which were below 0.2/zmol 1- ~or even below the practical detection limit of 0.1/zmoll -~. This lack of manganese release can be attributed to the high adsorption capacity of manganese oxides for soluble manganese and probably to the formation of an arsenate-manganese ion complex (Oscarson, 1983). Experimental data showed that the adsorption capacity for Mn 2+ is more than 0.55 mol per mol 5MnO 2 at pH 6.7, whereas Morgan and Stumm (1964) estimated a capacity of 0.5 mol per mol at pH 7.5. Hem and Lind (1983) found an increase of the oxidation degree of freshly precipitated manganese oxides on aging at pH < 7 due to disproportionation of Mn 3+. Therefore, a precipitation and oxidation of manganese ions could be a favorable mechanism to explain the deficit of soluble manganese during the oxidation of arsenate(Ill). The adsorption-desorption behavior of manganese

at an initial molar ratio of 3 is shown in Fig. 6. The release of manganese is high and depends strongly on pH. The adsorption capacity is much lower than in systems without arsenate(Ill) and reaches only 0.1 mole Mn 2÷ per mol 5MnO2 at pH 7. Adsorbed manganese is even remobilized at pH 6, accompanied by further reduction and dissolution of MnO2. It is not conclusive from our results, whether reduced manganese forms a manganese-arsenate(V) complex, as discussed by Oscarson et al. (1983) and Moore et al. (1990). Nevertheless, these authors made their investigations at initial MnO2-arsenate(III) ratios of nearly 2, where the formation of such complexes might be feasible. Calculations of data from tests with loaded filters showed that the complexation of manganese ions by arsenate(V) should be of minor importance. Oxidation in a loaded filter

The examined oxidation process was tested in a flow through reactor, realized in the technique of a preloaded filter. The filter had an area of 3.8 cm 2 and a depth of 25 cm. It was filled with crushed quartz (0.5-1mm) and loaded with 1.2mmol (105mg) 6MnO2. The inlet water was Berlin tap water with a pH of 7.8, an ionic strength of 16.5 mmol 1-~ and an oxygen content near saturation. As(Ill) was continuously dosed to the inlet. The filter rate was 5 m h - ' , A significant release of solid manganese oxide occurred in the first day of the run with manganese concentrations up to 0 . 4 m g l -~. Throughout the following run time, values for solid manganese were below 2 0 # g l -I.

304

WOLFGANG DRIEHAUS el al.

~mol r-1

1,8,

1,4 1,2 1'

0,8, 0,6, 0,4' 0,2'

Inlet/~(III+V)

(

Outlet AI(III+V)

i

/ lO

0

20

30

40

50 50

100

150

200

250

300

350

400

450

500

~mol r 1

1,6 1,4, 1,2:_

1,

~

--'~

Al(lll) oxidized

~,

0,8, 0,6, 0,4' 0,2', 0

Inlet As(Ill)

Mn, releu~l

10

20

30

40

50 50 100 150 Run Time, h

200

250

300

350

400

450

500

Fig. 7. Concentrations of arsenate(III + V) and oxidized arsenate(III) and soluble manganese in the inlet and outlet of the loaded filter.

Inlet and outlet concentrations of the first filter segment over a period of 21 days are shown in Fig. 7. Adsorption equilibrium for arsenate was reached after 10 h. The oxidation of arsenate(III) decreased in the following period to values of 30%. Release of soluble manganese increased and was equal to the amount of arsenate(III) oxidized between 40 and 60 h. It was expected, that the oxidation of As(III) would further decrease and reach a steady state at a low level. But the oxidation of arsenate increased after 60 h again and arsenate(III) was quantitatively oxidized after 260 h, while release of soluble manganese decreased to concentrations of 0. l # M. This obvious change in the oxidation behavior of the tested filter might be attributed to the seeding of bacteria, which are probably involved in the oxidation of manganese. Two reaction pathways may occur in the biologically aided arsenate(III)-oxidation in the MnO2-1oaded filter: (i) manganese oxides are the primary oxidant, which are then regenerated and precipitated by bacteria; (ii) bacteria directly oxidize arsenate(III), i.e. by extracellular enzymes.

Table 1. Balance data from the filter Operating time

0-10

10-56

56-260

260-497

h

Inlet As(lII + V) Inlet As(lII) Adsorbed As(Ill + V) As(III) oxidized Released Manganese

21.8 20.6 8.8 20.4 0

l l2 104 4.6 60 31.8

364 312 5.2 174 80

677 517 I 1.0 485 58

/~mol /Jmol /~reel #reel #reel

Mass balances for the filter were computed for four periods (Table l). Following the equation [1], we assume that 740 #moi of oxidized arsenate(III) in the filter produced an equal amount of divalent manganese. In the intervals up to 260 h, nearly 50% of reduced manganese were released. When the oxidation is in steady state after 260 h, total release of manganese was only 12% with respect to oxidized arsenate(III). The mass balances as well as the period of 10-12 days until reaching a steady state indicate, that bacteria may contribute to the oxidation of arsenate(III) in the presence of manganese oxides. There is no evidence for assuming only abiotic reactions, since the filter was fed with tap water and bacteria involved in the oxidation of divalent manganese are ubiquitous. Bacteria often occur in both the iron and the manganese removal filters in water treatment facilities and enhance the elimination of both elements (Mouchet, 1992). CONCLUSIONS

The h-modification of manganese dioxide is an effective oxidizing agent for the treatment of waters containing arsenate(III). The reaction rate depends strongly on the initial molar ratio of the reactands, but was not influenced by pH-variations between 5 and 10 at an initial molar ratio of MnO2/As(III) of 14. Calcium has only a minor effect on the oxidation. The deficit of soluble divalent manganese, that is produced by the redox process, is mainly attributed to adsorption on the oxide surface.

As(llI)-oxidation with manganese oxides This oxidation technique was successfully employed in a preloaded filter. One unexpected effect of this experiment was the increase of As(liD-oxidation and the decrease of manganese concentration after 60 h, which can not be explained by an inorganic reaction mechanism. The contribution of bacteria in this redox-reaction with manganese oxides is likely. Arsenate(Ill) may be directly oxidized by bacteria or reacts with biologically precipitated manganese oxides. The influence of other reductants commonly occurring in drinking water, i.e. Fe(II) and humic substances should be considered, as they probably decrease the As(III)-oxidation and lead to higher concentrations of soluble manganese. Although manganese is not toxic, enhanced amounts o f this element after an oxidative treatment are not tolerable, as they may cause some problems in the drinking water supply, like turbidity and precipitation in pipes and leads to impaired taste of the water. Nevertheless, nothing is yet known about the long term stability of reaction filters containing manganese dioxides, but it seems that they provide a simple oxidative treatment of arsenate(Ill) containing waters. The removal of arsenate(V) itself requires a second process, for example the adsorption on activated alumina or the coprecipitation with ferric salts. Acknowledgements--We gratefully thank Mrs H. Schmeis-

ser for the assistance in AAS-analysis and appreciate the financial support of the investigations by the "Federal Ministry for Research and Technology" (Grant No. 02 WT 9002/7).

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