Environ. Sci. Technol. 1992, 26, 1361-1368
Surface-Catalyzed Autoxidation of Sulfur( I V ) in Aqueous Silica and Copper( I I ) Oxide Suspensions Devarakonda S. N. Prasad, Ashu Rad, and Krishna S. Gupte*
Atmospheric Chemistry Laboratory, Department of Chemistry, University of Rajasthan, Jaipur-302004, India Kinetics of the autoxidation of sulfur(1V) in unbuffered and acetate-buffered aqueous silica and copper(I1) oxide suspensions appear to be surface mediated, and the suggested mechanism requires the adsorption of both dioxygen and sulfur(1V) on the particle surface involving solid-liquid-gas interactions. Introduction Because of its direct bearing on atmospheric acid precipitation, the autoxidation of sulfur(1V) has been the subject of several studies (1-3). Nevertheless, the surface-catalyzed SOz autoxidation in aqueous suspensions of solid powdered catalysta such as carbonaceous particles (4,5),fly ash (6), glass (7),CdO (B), COO(9),atmospheric dust (IO),supported cobalt(I1) complexes (II), and a-Fe203 (12,13)and its various polymorphs (14)has been studied only recently. In a thorough investigation on a-Fez03,Faust et al. (13) provided evidence for the formation of surficial complexes, a-FeOSOf or a-FeO,SO-, in the low-pH range (1-3). Enhanced reactivity in the higher pH range (5-7) was attributed to a more favorable surface complexation of sulfite ion as in eq 1. An alternative mechanism involving co-adsorption of both sulfur(1V) and 02,as in eq 2, was also proposed. +
so:-
/OH
rFe,os022-
+ Hz0
(1)
We, in this paper, report the kinetics of the autoxidation of sulfur(1V) in aqueous suspensions of silica and copper(I1) oxide, because of their strong environnlental connection. Silica is a major constituent of suspended particulate matter (151,shales, sandstones, igneous and silicate rocks, and limestones (16). Copper oxide aerosol is emitted concurrently with SOz from copper smelters (17). Homogeneous copper(I1)-catalyzed sulfur(1V) autoxidation (18-22) and the Cu(I1)-S(1V) reaction (23-25) have been the subject of several studies. The heterogeneous reaction between CuO aerosol and SOz has also been studied (26, 27). Experimental Section Silicon dioxide (Nova Kem, particle size range 0.5-10 pm with 80% of them of 1.5-pm size), cupric oxide (particle size range 0.1-1.5 pm), and all other chemicals of reagent grade were used. All solutions were prepared in triply distilled water. The reactions were conducted in reaction flasks open to air and in room light. Acetate buffer was used for maintaining pH. The reactions were initiated by adding sodium sulfite to the buffer solution followed by addition of SiOzor CuO. The reaction mixture was stirred continuously and magnetically at 1600 f 100 rpm at which speed the reactions were shown not to be oxygen masstransfer controlled. The kinetics were followed by analyzing the aliquots periodically for sulfur(1V) iodometrically, using starch as an indicator. SiOz and CuO did not 0013-936X/92/0926- 1361$03.00/0
interfere in the estimation, as shown by control experiments performed in their presence and absence. Even the Cu(I1) ions, present in the solution owing to the dissolution o;f CuO in buffered solutions, did not cause any interference. Further it was also observed that in acetate buffers, when a dilute solution of KI is added to CuO suspensions o'r dilute Cu(I1) solutions, there is no immediate liberation of iodine. Apparently, the reaction between the acetate copper(I1) complexes and iodide ions is relatively slow and a large concentration of the 1: ter is needed for rapid release of iodine. The reaction rates were reproducible within &lo%. Sulfate was detected to be the only product, and no evidence was found for dithionate. Recovery of the sulfate as BaS04from the final product solution was -97 f 2% of the theoretical value in terms of eq 3. The rate of s1o*/cuo S(1V) + 1/202 SWI) (3) autoxidation was found to be the same in diffused room light as in the dark. Thus, a photocatalytic pathway can ble ignored in these studies. In both the cases, no reaction occurred in a nitrogen atmosphere over a period of 30 min.
-
Ilesul ts Preliminary Observations in Silica Suspensions. The possibility of catalytic activity of silica suspensions being due to leached trace metal ions was ruled out by the fact that the rate of S(1V) autoxidation in leachate and blank reactions was found to be almost the same (Figure 1). This conclusion was reinforced by the effect of EDTA, (0.01-1.0) X 10" mol dm+, which was unable to influence the silica-catalyzedreaction (Figure 1).However, the same [EDTA] inhibited the blank and leachate reactions, which are trace metal ion catalyzed. Kinetics in Buffered Silica Suspensions. With increasing buffer, while maintaining the ratio of [CH3COONa]/[CH,COOH] constant, the rate of autoxidation is inhibited. In all experiments described henceforth, overall [CH,COONa] was maintained at 0.15 mol dm-3, and wherever needed, [CH3COOH]was varied to obtain the desired pH. The uncatalyzed rate did not show any significant dependence on [buffer]; the observed [buffer] dependence is therefore almost wholly due to the elilica-catalyzed path. A log-log plot between Robs and [CH3COO-]yielded an order of -0.6 in [acetate]. The inhibition of apparent rate in the presence of increased [acetate] indicates competition for surface reactive sites. The metal oxides are known to coordinate organic ligands such as CH3COO- via an inner-sphere mechanism (28). Acetic acid is also known to be adsorbed on silica (29). The results of variation in [SiOz],as shown in Figure 2, fitted eq 4, where Runcatis the rate of the uncatalyzed = Runcat + kZ[sio21 (4) reaction. Values of kz and Runcatare given in Table I. When the dependence of Runcatand kz (Table I) on S(1V) is included, eq 4 modifies to eq 5. kl and k, were deterRobs = k~[S(IV)l+ k3[SiOzl[S(IV)12 (5) mined to be 1.73 X s-l and 1.9 X dm6 mol-I g-l sir1, respectively, at pH 5.14 and 30 "C.
0 1992 American Chemical Society
Environ. Sci. Technol., Vol. 26, No. 7, 1992
1361
30 40 50 60 Tim, min Flgure 1. Effect of EDTA on reaction rate at [S(IV)] = 5 X mol dm-3, pH 5.14 and t = 30 "C: (0)[SIO,] = 1.0 g dm-3; ( 0 )[SiOd = 1.0 g ~Irn-~, [EDTA] = 1 X 10" mol dm3; (e)[SIO,] = 1.0g dm- , [EDTA] = 1 X mol dm-3; (A) uncatalyzed reaction; (X) with leachate of 1.0 dm-3 SO,; (0)uncatalyzed reaction, [EDTA] = 1 X mol dm- . "0
10
20
f
lo-'
0 0
10
20
30 40 50 60 103[S(IY)1.[SiO,I,md g dni6
70
80
Flgure 3. Plot of Robs/[S(IV)]vs ([S(IV)][SiO,]) at pH 5.14 and t = 30 "C.
[SI02I, g dm-3 Figure 2. Variation of SiO, at pH 5.14and t = 30 "C. [S(IV)]: (A) 2X (0) 5X (8)7.5 X mol dm-3. Table I. Values o f Rumcat and pH 5.14 and 30 OC
Flgure 4. Variation of k, with [H+] at t = 30 "C.
k2in Silica Suspensions a t
103[S(IV)I, mol dm-3
107R""C.at,
mol dm-3 5-l
107k,, mol g-' s-l
2.0 5.0 7.5
2.2 7.1 10.3
0.66 3.70 10.50
Table 11. Values o f k la n d k3in S i l i c a Suspensions a t D i f f e r e n t pH and 30 "C
pH
The results of S(1V) variation were also in agreement with eq 5. When the results of all [S(IV)] and [Si02] variations were treated in terms of eq 6 (Figure 3), the best Rob,/[S(IV)I = k l -I- k3[S(IV)I[SiOd (6) fit kl and 12, values were found to be 1.78 X s-l and 1.7 x dm6mol-' g-l s-l, respectively, at pH 5.14 and t = 30 "C. When the pH was increased from 4.64 to 5.84, Robsincreased and an order of -0.94 f 0.07 in [H+] was determined. Values of kl and k3 at different pH values (Table 11) were obtained by varying [S(IV)] at different but fixed pH values. When the dependences of k, and k3 on [H+] (Figures 4 and 5) are included, the complete rate law can be written as eq 7. When kl or k3 were plotted Robs = k4[S(IV)][H+]-' + (k, + k,[H+]-1)[Si0,][S(IV)]2 (7) 1362
Environ. Sci. Technol., Vol. 26, No. 7, 1992
10-4[H+l-', dm3 mor'
4.64 4.87 5.14 5.23
lo%,, dm6 lo4&,s-l mol-' g-l s-l 0.79 1.61 3.65 5.30
0.7 0.6 1.5 1.4
pH
5.37 5.53 5.64
10Zk,,dm6 104kl, s-l mol-l g-I s-' 6.10 11.30 21.30
2.6 2.6 5.6
against [H+]- (Figure 4 and Figure 5) the values of k4, k5, and k6 were found to be 3.3 X lo* s-l, 2.1 X dm6 mol-' g-l s-l, and 8.1 X dm6 mol-' g-l s-l at 0.15 mol dm-3 acetate and 30 "C. When the buffer dependence of Robsis included, eq 7 becomes eq 8, which is equivalent to eq 7 through k, = Robs= k4[S(IV)][H+]-' i-
(kg/+ k,l[H+])-1[S(IV)]2[Si02] [CH3C00-]-0,6(8)
ki(CH3COO-)-0.6and k5 = Kgl (CH3COO-)-0.6.These relationships can be used to obtain k, and k, at other acetate concentrations, using the values reported here in at 0.15 mol dm-3 acetate. It must be distinctly understood that
Table 111. Robsat Different [S(IV)] and [SiOJ at pH 3.19 and 30 O C in Unbuffered Solutions
-
'yl
""il
/I /
50
1.6 3.5 5.0 6.65 9.0
1.0 1.0 1.0 1.0 1.0
0.375 0.79 0.82 0.85 1.16
5.0 5.0 5.0 5.0
1.0
2.5 5.0 7.5
0.82 1.66 3.92 5.66
9 0
87c
IO+[H+I-', dm mor' Flgure 5.
Variation of
k 3 with [H']
at t = 30 O C .
'W
6-
0
-
F 52
e
> 2 o L s
PH
Variatlon of pH in unbuffered soiutlons: [S(IV)] = 5 X mol dm-3, [SiO,] = 1.0 g dm-3, t = 30 O C .
-
Flgure 7. Plot of R,,I[S(IV)]
Flgure 8.
the rate laws in eqs 7 and 8 are applicable in acetatebuffered suspensions only and in the range of [S(IV)] = (1-8) x lo-, mol drn-,, [Si02] = 1-15 g dm-3, and pH 4.64-5.84 and under atmospheric oxygen pressure. From a plot of log Robsvs T', an overall empirical energy of activation value of 56 kJ mol-l was determined. Similarly, the energies of activation associated with k4,k,, and k, were determined to be 28.5, 70, and 84 kJ mol-l, respectively. Kinetics in Unbuffered Silica Suspensions. In the unbuffered suspension study, the desired initial pH was adjusted with perchloric acid. Initially, on increasing pH (2.27-3-68), the rate of the reaction decreases and attains a minimum value; a further increase in pH results in an increase in the rate (Figure 6). Additionally, above pH 4, an initial rapid drop in [S(IV)] also appears and its magnitude increases with increase in pH but is independent of [S(IV)]and [SiOz]as revealed by investigations at pH 6.24. For the investigation of SiOz and S(1V) concentration dependences, a pH value of 3.19, at which pH there was no initial rapid drop in [S(IV)], was selected. The results (Table 111) fitted eq 9. At pH 3.19 and t = Rcat = k7[S(IV)I[Si021 (9) 30 O C , the value of k7 was (1.4 f 0.1) X lo4 g dm-3 s-l. The energy of activation determined by plotting log k7 against T' was 83 kJ mol-l. The rate increase below pH 3.19 on decreasing the pH could be due to flushing out of SO, from the aqueous suspensions or due to greater reactivity of SOZ-Hz0.
1O2[CuO1 ,g dm-' vs [CuO] at pH 5.23 and t = 30 O C .
Kinetics in Buffered CuO Suspensions. In this case also, an increase in [buffer] inhibited the rate of autoxidation, but only slightly. The kinetics order in [CH3COO-] was determined to be -0.2. Hence, in all kinetics experiments, [CH,COONa] was held at 0.16 mol dm-3, and [CH,COOH] was altered for varying pH. The results of CuO variation at different [S(IV)] and at pH 5.23 and t = 30 "C fitted eq 10, which yielded k8 and k, values of 5.3 X s-l and 5.5 X dm3 g-l s-l, respectively. Robs=
(k8 ikg[CUO])[S(IV)]
(10)
The Robsvalues for [S(IV)] variation at different CuO were plotted against S(IV),and from the slopes of resulting linear plots, values of (It8 + k,[CuO]) were obtained. These slopes were then plotted against [CuO] to obtain values of kg and k,, which were 1.4 X 10" s-l and 4.75 X dm3 g-l s-l, respectively, at pH 5.23 and t = 30 "C. A good agreement in the value of kg determined from [CuO] and [S(IV)] variations is seen although there is some difference in the value of k,. Equation 10 can be rearranged as eq 11,where Runcat(=k,[S(IV)]) is the rate of the CuO-cata(11) Rcat = Robs - Runeat = kg[S(IV)l[CuOl lyzed reaction. A plot of Rcat/[S(IV)]vs [CuO] is linear (Figure 7), testifying to the validity of eq 11. When the pH was varied (4.35-5.64), the rate of the reaction increased. In order to analyze the [H+] dependence of the catalyzed path (eq l l ) , the values of R, were obtained, by subtracting from Robs the values of Runcat determined separately. A plot of Rcat/[CuO][S(IV)] vs ["]-I is linear with a nonzero intercept (Figure 8). Thus,
Environ. Sci. Technoi., Voi. 26,No. 7, 1992
1363
Table V. Values of Robsand Second-Order Rate Constant, k,,,for CuO-Catalyzed Reaction in Unbuffered Solutions at 30 "C
103[S(IV)I, mol dm-3 2.0 4.0 2.0 2.0 2.0 2.0 2.0 2.0 2.0
OO
'
'
5
lb
'
1'5
' 2 0
' 2 5 ' 3 0 ' 3 5 1O4[H*I~',drn3 mol-'
' 4 b '
L -
Varbtion of Rmt with hydrogen ion concentration in buffered CuO suspensions at 30 O C . Flgure 8.
2.0 2.0 2.0 2.0 2.0 4.0 6.0 9.0
[CUOI;
1 0 7 ~ ~ 103k12, ~ ~ ~
g dm
PH
mol dm-.3s-l
g-I dm3 8-l
0.05 0.05 0.05 0.05 0.03 0.05 0.07 0.11 0.12 0.15 0.18 0.20 0.25 0.05
4.21 4.21 4.80 5.49 5.28 5.28 5.28
2.27 3.68 2.00 4.23 1.26 2.30 3.00 4.30 4.76 6.47 8.33 9.00 11.9 2.02 3.03 4.30 8.45
2.4
0.05
0.05 0.05
5.28 5.28
5.28 5.28 5.28 5.28 5.28 5.28 5.28 5.28
1.8
2.0 4.2 2.1 2.3 2.1 2.2
2.0 2.2 2.3 2.2 2.4 2.0 1.5 1.4 1.9
Table VI. Rate Parameters for Homogeneous Cu(I1)-Catalyzed Autoxidation of Sulfur(1V) in Acetate-Buffered Solution at 30 "C
107k,~ 104[CU(II)], 1O3[S(IV)l 106[H+], 1 0 7 ~ , ~ , , (eq 14), dm2,91 mol-0.97 s-l mol dm+ mol dm-3 mol dm-3 mol dm-3 5-l I
-2
c
o,50LiLL~-10 20
30
40
50
time, rnin
Rate profiles for CuO-catalyzed [S(IV)] autoxidation in unbuffered solutions at 30 OC: [S(IV)]= 2 X mol dm-3, [CuO] = 0.05 g dm-3;(0) pH 6.27,(0)pH 6.83,(A)pH 5.2. Flgure 9.
Table IV. Values of k,,and k l l for CuO-Catalyzed Autoxidation at Different Temperatures and at [-02CCH,] = 0.16 mol dm-3
103klo,dm3 g-' s? 10sk,,, dm3 g-'
30
t , "C 35
40
1.66 3.70
2.71 6.13
3.40 8.72
the complete rate law for the catalyzed path can be written as eq 12. Values of k,, and kll, determined at different Rcat
= (k10 + k~~[H+I-')[CuOl[S(IV)I
Environ. Sci. Technol., Vol. 26, No. 7, 1992
2.0 2.0
2.0 2.0 2.0 2.0 2.0
2.0 2.0 2.0 2.0 2.0 2.0 2.0 2.0 2.0
2.0 2.0 4.0 6.0 8.0
5.00
5.00 5.00 5.00
2.0
2.0
5.25 5.25 5.25 5.25 5.25 5.25 5.25 5.25 5.25 5.25 5.25 5.25 5.25 5.25 5.25 5.25 5.25 5.25 5.25 5.25 5.25 9.55 17.0
2.34 1.94 1.25 1.70 1.92 1.44 1.46 1.25 2.30 1.52 5.00 5.70 6.00 7.00
7.20 7.80 8.20 5.20 11.9 31.3 65.4 2.20 1.40
8.1
8.0 7.9 8.2 8.0 8.2 8.2 7.7 5.2 6.8 8.5 5.8 6.8
(12)
temperatures, are given in Table IV. From these values, energies of activation for k,, and k,, paths were found to be 57 and 67 kJ molP, respectively. When the dependence of Rcaton [buffer] in terms of [acetate] is included, eq 12 becomes 13. Rcat = (klo' + kl1'[H+]-l) [CUO][S(IV)][CH3COO-]+'.2 (13) Kinetics in Unbuffered CuO Suspensions. The influence of pH was studied in the range 3.21-6.83. The disappearance of S(1V) with time was normal between pH 3.21 and 5.49. However, in the region pH >6, an initial rapid drop in [S(IV)] was observed (Figure 9). The CuO and S(1V) dependences (Table V) investigated at pH 5.28, at which there was no complication,fitted the experimental dm3 8-l s-l a t rate law 14. k12 had a value of 2.44 X Robs = kl2[S(IV)I [ c u o l (14) 30 "C and pH 5.28. The values of Robsbetween pH 4.21 1364
0.001 0.002 0.004 0.007 0.01 0.03 0.05 0.07 0.10 0.30 3.75 6.30 8.80 14.0 18.8 24.0 30.0 5.00 5.00
and 5.28 are nearly the same and only the value at pH 5.49 is seen higher (Table V). Homogeneous Copper(I1) Catalysis. For the sake of comparison, the homogeneous Cu(I1) ion catalysis of S(1V) autoxidation in acetate-buffered solutions has also been investigated briefly. [Cu2+]lower than 1 X mol dmn3 were ineffective in catalyzing the reaction (Table VI). The kinetics results of variation of [Cu(II)] = (0.1-3) X mol mol dm+, and pH were in dm-3, [S(IV)] = (2-8) X agreement with eq 15. The value of k13 was found to be (7.5 f 0.8) X los7 dm2,91 s-l at 30 "C. Robs
= k13[CU(II)]".~~[ S(IV)] 1.75[H+]--l
(15)
Homogeneous Copper(I1) versus Heterogeneous CuO Catalysis. The log of the solubility product of CuO is reported to be -20.35 (30). The concentration of Cu(I1) ions, [Cu(II)],, present in suspensions owing to partial dissolution of CuO, was determined by stirring suspensions containing different amounts of CuO for -30 min, filtering
Table VII. Solubility of CuO, and Rates for S(1V) Autoxidation at Different CuO and Copper(I1) Concentrations ([S(IV)] = mol dm-3,pH 5.23, t = 30 "C) 2X
10[CUO], g dm-3
104[CU(II)I,," mol dm-3
0.3 0.5 0.7
3.0 2.8 3.5 4.3 6.0 7.6 9.0 0.7 2.2
1.1
1.5 2.0 2.5 1.1
2.5
104[Cu(II)l,,,b mol dm-3
107RIC(IIp:l mol dm- s-
mol 107RIC"(Iylaq'; dm- s-
Buffered Solutions 3.77 3.33 6.29 6.12 8.80 8.16 13.8 13.3 18.9 15.7 25.0 21.5 31.3 27.0
4.80 4.50 5.21 5.30 5.55 5.53 5.60
5.0 5.7 6.0 7.0 7.2 7.8 8.2
Unbuffered Solutions 13.8 4.3 31.3 12.0
0.90 1.05
0.9 1.0
1o1Robs,C
mol dm-3 s-'
[Cu(II)], represents the copper ion concentration present in solution containing CuO after 30 min of stirring. [Cu(II)], is the copper ion concentration that would be obtained if the CuO shown in Column 1 dissolves fully. CRobsis the rate of the CuO-catalyzed heterogeneous reaction. dR~Cu(II)l, is the rate of a homogeneous Cu(I1)-catalyzed reaction having a copper ion concentration shown in column 2. eRICu(II)laq is the rate of a homogeneous Cu(I1)-catalyzed reaction having a copper ion concentration shown in column 3.
out CuO, and determining [Cu(II)],. Although with increase in [CuO] the values of [Cu(II)], increase, the increase in the latter is proportionatly less (Table VII). The rates of S(1V) autoxidation at a given [CuO] and at the corresponding [Cu(II)], (Table VII) were determined separately. For example, at 0.15 g dm-3 [CuO], Robfor the heterogeneous reaction is 1.57 X lo4 mol dm-3 s-l, while for the homogeneous reaction at corresponding [Cu(II)],, Le., 6.0 x IO4 mol dm-3, it is 5.5 X lo-' mol dm-3s-l. Thus, the value of Robs is seen to be substantially higher than Rrcu(II)lB. Obviously, had the catalysis been solely due to dissolved Cu(I1) ions, the rates with 0.15 g dm-3 [CuO] and with 6.0 X lo4 mol drn+ [Cu(II)], should have been same. This is not the case, and the similar situation prevails at other CuO and Cu(II), concentrations also. This suggests that CuO catalysis is not solely due to Cu(I1) ions alone, but the surface catalysis also plays a prominent role. Now we look the problem from a different angle. In an extreme and most unfavorable case, if it is assumed that the whole of CuO taken in suspension gets dissolved completely and immediately, the [Cu2+]present in solutions owing to the dissolution would be equal to [Cu(II)],. The rates at various CuO and C U ( I I )concentrations ~~ were determined separately and are given in Table VII. These results show the CuO-catalyzed rates, Robs, to be substantially higher than the corresponding homogeneous [Cu(II)],,-catalyzed rates. The foregoing analysis implies that, even initially, the [Cu2+] present in the reaction mixture is equal to [Cu(II)], or to [Cu(II)], . This is not the case. In fact, in the initial region, [CufiI)], is of the order of mol dm-3 or less. As such low [Cu(II)] is not effective in catalyzing the autoxidation (Table VI), it is safe to assume that the interference from the homogeneous Cu(I1) catalysis in the initial region would be minimal. This is also supported by the quite diverse nature of the rate laws in eqs 12 and 15 for CuO- and homogeneous Cu(I1)-catalyzed reactions. The effect of EDTA (Figure 10) and mannitol (Table VIII) on the homogeneous and heterogeneous catalyzed reactions is also revealing. While the addition of EDTA introduced an induction period in the heterogeneous reaction, it completely inhibited the Cu(I1)-catalyzed reaction. While the homogeneous reaction suffers only a slight deceleration in the presence of mannitol, the heterogeneous reaction, apart from a slightly stronger rate of deceleration, is also attended by an induction period, the manifestation of which may be linked the free-radical scavenging (32)and to coating of the catalyst surface (33).
2 .o m
'E
0
--E"
5 z
1.5
m
0
1.o
0
10
20
30
40
50
60
lime, min
Flgure 10. Effect of EDTA on CuO- and Cu(I1)-catalyzed reactions: [S(IV)] = 2 X mol ~ l m - pH ~ , 5.28, t = 30 OC, [CuO] = 0.07 g dm-3, [CuSO,] = 4 X lo-, mol dm-3, [EDTA] = 4 X lo-, mol dm3; (0)CuSO, with EDTA, (X) CuSO, without EDTA, 0) CuO with EDTA, (A)CuO without EDTA.
Table VIII. Influence of Mannitol on CuO and Cu(I1) Ion Catalyzed Autoxidation of S(1V) [S(IV)] = 2 X mol dm-3,pH 5.15, t = 30 "C
[CuO], 103[CuS04], 102[mannitol], 1O7RObs, induction g dm-3 mol dm-3 mol dm+ mol dm-3 s-l period, min Heterogeneous Reaction 8.20 0.05 7.66 0.50 2.94 1.00 2.52 1.50 2.11 2.50 1.64
0.10 0.10 0.10 0.10 0.10 0.10 1.0 1.0 1.0 1.0 1.0 1.0
5 8 9 10 10
Homogeneous Reaction 4.12 0.10 3.21 0.50 2.07 1.00 2.00 1.50 2.00 2.50 2.11
The high pHZpC (9.5) of CuO (34) will make X u (OH)(OH),+ (eq 16) be the dominant form in the pH range
of this study. The surface hydroxyl groups are likely to adsorb Cu(Hz0):+ and its various acetato-sulfito comEnviron. Sci. Technol., Vol. 26, No. 7, 1992
1365
plexes, through coordination. However, the cation adsorption is not expected to be high as the experimental pH in our study is much less than pHZpC(35). Discussion In aqueous media the hydroxylation of silica is well established. Hydroxylated silica is known to undergo the following acid-base equilibria:
The pH of the zero-point charge, pHZPC,is reported to be 2.8 (36). Interestingly, the minimum in the rate in unbuffered suspensions occurs around pHzpc (Figure 6), at which the electrical double-layer effects on HS03- would be most pronounced. S(1V) in the pH range of the study would be largely present as HS03- ( 3 3 , with only a fraction as S032-and S02-H20. The observed [H+] dependence requires the former two forms to be reactive. The effect of stirring speed on the reaction rates and the absence of any reaction without 0, require the involvement of oxygen as a reactant. As the reaction appears to be surface catalyzed, it is reasonable to assume that both S(1V) and 0, are adsorbed on the particle surface. Incidentally, the adsorption of a number of diverse species on silica surface is well-known (38). Evidence for the adsorption of sulfur(1V) on CuO also exists (26, 27). Moreover, silica is also reported to catalyze the oxidation of Fe(I1) to Fe(II1) (29). In sulfur(1V) autoxidation in soot suspensions (4),the energy of activation (49 kJ mol-l) was in the neighborhood of the activation energy (50 kJ mol-l) for the chemisorption of oxygen. The initial rapid reaction noted in this case was shown to be due to the reaction of preoxygenated carbon particles with S(IV). Degassing of the carbon particles led to the disappearance of the initial rapid reaction. Brodzinsky et al. (4) adduced these two findings as evidence of the adsorption of oxygen on carbon particles. A similar mechanism is likely to operate in our case too, as the energies of activation are of similar magnitude and an initial rapid reaction in unbuffered suspensions has also been noted. Based on the above arguments, the following mechanism in terms of the Hinshelwood-Langmuir model (39)may be proposed.
+ SO?- rapid (18) Kl9 SiO, + 0, eSi02-02 rapid (19) Km Si02.02+ HS03- e SiO2.O2.HSO3- rapid (20) K21 Si02.02+ S032-e Si02.0243032- rapid (21) Kzz Si02.02.HS03- + HS03-e Si02.02(HS03-), rapid HS03- _ffs Ht
(22) KZ.3
Si02-02.S032+ HS03- e Si02.02-S032-.HS03-
2SiO, + HS05Si02.02.S032-2Si02 + Sob2-
SiO2.O2.HSO3-
Si02-02-HS03--S0322 SiO, 1363
rapid (23) slow slow
+ S20Y2-+ OH-
Environ. Sci. Technol., Vol. 26, No. 7, 1992
(24)
SiO2*O2(HSO3-), SO5,-
+ H20
+
-
+ Hf
S20Y2+ H20
HS05-
slow (27)
rapid
2S042- + 2H+
(28)
rapid
(29)
HSOc + S(1V)
BS(V1) rapid (30) Peroxomonosulfate (40) and -disulfate (41)ions, implicated in eqs 24-27, have been proposed as intermediates earlier. The formation of S2072- finds support from the surface species SiP207,which is reported to form when H3P04 is adsorbed on silica (29). Assuming that the equilibria in the proposed mechanism (eqs 18-23) are rapid and that the equilibrium constants have low values, the rate law 31 at constant [O,] can be derived. Robs = Runcat + ( ( A+ B[H+]-’)[S(IV)]+ (C + D[Ht]-1)[S(IV)]2)[Si02] (31) where A = k14&&19[021 B = k1&21K1&d[021 c = k1&2&&19[021 = k16K23KZlKl&d[021 Apparently, in unbuffered suspensions with the contribution from Runcst,C and D appear to be insignificant, and so eq 31 reduces to 32. (32) Robs = ( A + B[Hf]-1)[S(IV)I[Si021 On the other hand, in buffered suspensions, A and B appear to be insignificant and therefore eq 31 modifies to 33.
( C + D[H+l-’)[S(IV)12[Si021 (33) The rate eqs 32 and 33 are identical to experimentally observed eqs 7 and 9. Based on the mechanism proposed for autoxidation in a-Fe203suspensions (13),an analogous detailed mechanism involving a concerted 4-electron transfer may be proposed as follows: Robs
= Runcat
OH >SiLOH
+
+
HS03-
H+
==
O ‘H
OS02H >SiLOH \OH O , SOH , >Si-OH
+
HS03-
+
+ O2 +
H’
H+
O , SOpH >Si-OH O ‘H OS02H >Sic, O ‘ S0,H
OS0,H
H20
(34)
+
H20
(35)
$.
H20
(36)
‘OS02H
O , SOZH >Si-o,+ ‘OS02H >SiCOH O ‘ S0,H
+
OS0,H i >SiL02+
\OS02H
+
,0S03H H ~ O1 , >S~-OH ‘OSO~H
+ 2H20
* 2504‘-
+
+
OH >Si/OH
+
H+
(37)
4H+
(38)
O ‘H
When acetate is present, it competes for a nearby surface site and thus shuts off this pathway as follows: /OH >Si
-OH \OS02H
+ CH3COO- + H+
>Si
O , Zc -OH
3
+ H20
(39)
\OS02H
Analogous to the silica case, the following mechanism may be proposed for autoxidation in CuO suspensions.
(25) slow (26)
SiO,
Ka
CuO + O2 eCuO-O2 Cu0.0,
+ S03H- eCu0-O26O3H4
1
rapid rapid
(40) (41)
+
K42
CuO-O2 SO?- F= CUO.O~.SO~~-rapid (42) Cu0.O2.SO3H- 5 CuO
+ HS06-
cuo.02.so32- -5%CUO + 50:-
slow slow
(43) (44)
which is followed by the steps in eqs 28 and 30. In terms of >Cu(OH)(OH2)+,which is the dominant hydroxylated form of CuO under our conditions (34),a mechanism similar to that written for silica may also be given. The mechanistic steps in eqs 40-44 would lead to the experimentally observed rate law. The proposed mechanisms would require the adsorption of S(1V)to be described by a Langmuir type of adsorption isotherm due to the limitations of surface reactive sites. However, the absence of saturation kinetics suggests the value of adsorption equilibrium constants in our case to be quite low. The absence of any significant reaction between CuO and S(1V) in the absence of O2rules out the possibility of a reduction-oxidation mechanism (42,43) being operative and makes the alternative associative mechanism, which is based on the adsorption of O2 and its subsequent dissociation either in a separate step or concerted with S(1V) (41, 42), more probable. It is possible to explain the present results by assuming the formation of surficial complexes akin to -02~C~11(H20)4SV032and -02.C~~'~(H20)4SIv032-,which have been proposed for homogeneous Cu(I1)-S(IV)-O2 system (25).Probably the adsorbed S(IV) transfers its electrons to matrix-Cu, and S(V), so formed, interacts with adjacently adsorbed 02,takes back an electron from matrix-Cu, and leaves off as S052-. In the case of silica, there is no possibility of electron transfer between surficial silicon and an adsorbed species. The role of the silica surface appears to be limited to the reactant activation through chemisorption. A comparison of the present results with those on COO (9) and Co203(44) shows the reactivity order to be CuO > Co203> COO> SiOz. The major kinetics and mechanistic differences in transition metal oxide and Si02-catalyzed reactions arise from the order in S(IV),which is one in the former and two in the latter, and from the possibility of intervention by multiple oxidation states in the former and its absence in the latter. In the case of Co203,the electron-transfer mechanism appears to be the same as proposed in the case of CuO. For COO,the inaccessibility of the Co(1) state makes a similar mode of electron transfer less probable. Obviously, activation through electron transfer must follow a reverse course; Le., at first adsorbed O2extracts an electron from matrix-Co2+and gets reduced into 02-,which may interact with adjacently adsorbed S(IV)to ensure further reactions. In the case of CuO, both activation modes are possible as both Cu(1) and Cu(I1) states are accessible (25) and this probably explains the high rate of CuO-catalyzed autoxidation. Berresheim and Jaeschke (45), in their study on the kinetics of SO2 removal by aerosols of CuS04,CuCl,, and Cu(NOJ2, ascribed the catalytic activity of aerosols to the activity of dissolved metal ions present in the aqueous phase, although the correlation was quite poor. Quite likely, surface catalysis might be important in their case too.
Conclusions The kinetics results suggest the particle surface to be the major cause of catalysis of SO2autoxidation in silica and CuO suspensions. This is in agreement with previous indications that wetted aerosol particles may play an im-
portant role in adsorption and subsequent oxidation of SO2 (45). Although the low water content of aquated haze aerosol particles makes this path a poor competitor to homogeneous reactions in cloudwater, the CuO catalysis may be of importance in smelter plumes, e.g., from Sudbury, Canada, and Khetri, India.
Acknowledgments We are thankful to Dr. R. E. Huie, National Institute of Standards and Technology, Gaithersburg, MD, for help in numerous ways and to the reviewers for making several useful suggestions.
Literature Cited (1) Calvert, J. G., Ed. SO2, NO and NO2 Oxidation Mechanisms: Atmospheric Considerations;Acid Precipitation Series; Butterworth: Boston, 1984; Vol. 3. (2) Schwartz, S. E., Ed. Trace Atmospheric Constituents, Properties, Transformationsand Fates, 1st ed.; John Wiley: New York, 1983. (3) Busar, R. B., Lodge, J. P., Jr., Moore, D. J., Eds. Sulphur in the Atmosphere; Pergamon Press: Oxford, UK, 1978. (4) Brodzinsky, R. J.; Chang, S. G.; Markowitz, S. S.; Novakov, T. J . Phys. Chem. 1980,84, 3354-3358. (5) Rogowski, R. S.; Schryer, D. R.; Cofer, W. R., 111; Edhal, R. A,; Munna Valli, S. In Heterogeneous Atmospheric Chemistry; Schryer, D. R., Ed.; American Geophysical Union: Washington, DC, 1982; pp 174-177. (6) Cohen, S.; Chang, S. G.; Markowitz, S. S.; Novakov, T. Environ. Sci. Technol. 1981, 15, 1498-1502. (7) Rani, A.; Prasad, D. S. N.; Jain, U.; Gupta, K. S. Indian J. Chem. 1991,30A, 756-764. (8) Rani, A.; Prasad, D. S. N.; Bhargava, R.; Gupta, K. S. Bull. Chem. SOC.Jpn. 1991,64, 1955-1961. (9) Prasad, D. S. N.; Rani, A.; Madnawat, P. V. S.; Bhargava, R.; Gupta, K. S. J.Mol. Catal. 1991,69, 393-405. (10) Rani, A.; Prasad, D. S. N.; Madnawat, P. V. S.; Gupta, K. S. Atmos. Environ. 1992, 26A, 667-673. (11) Hong, A. P.; Boyce, S. D.; Hoffmann, M. R. Environ. Sci. Technol. 1989,23, 533-540. (12) Faust, B. C.; Hoffmann, M. R. Environ. Sci. Technol. 1986, 20,943-948. (13) Faust, B. C.; Hoffmann, M. R.; Bahnemann, D. W. J. Phys. Chem. 1989,93,6371-6281. (14) Leland, J. K.; Bard, A. J. J. Phys. Chem. 1987, 91, 5076-5083. (15) De, A. K. Environmental Chemistry; Wiley Eastern, New Delhi, 1989; Chapter 6, p. 119, Chapter 7, p 236. (16) Vander Weijden, C. H.; Middelburg, J. J. Water Res. 1989, 23, 1247-1253. (17) Eatough, D. J.; Christenson, J. J.; Eatough, N. L.; Hill, M. W.; Major, T. D.; Mangelson, N. F.; Post, M. E.; Ryder, J. F.; Hansen, L. D. Atmos. Environ. 1982, 16, 1001-1015. (18) Fuller, E. C.; Crist, R. H. J. Am. Chem. SOC.1941, 63, 1644-1650. (19) Barron, C. H.; O'Hern, H. A. Chem. Eng. Sci. 1966, 21, 397-404. (20) Mishra, G. C.; Srivastava, R. Chem. Eng. Sci. 1976, 31, 469-471. (21) Lunak, S.; El-Wakil, A. M.; Veprek-Siska, J. Collect. Czech. Chem. Commun. 1978,43, 3306-3316. (22) Backstrom, H. L. J. 2.Phys. Chem. 1934,25B, 122-138. (23) Carlyle, D. W.; Zeck, 0. F. Znorg. Chem. 1973,12,2978-2983. (24) Zeck, 0. F.; Carlyle, D. W. Inorg. Chem. 1974,13,34-38. (25) Conklin, M. H.; Hoffmann, M. R. Environ. Sci. Technol. 1988,22,891-898,899-907.
(26) Kent, S. A.; Katzer, J. R.; Manogue, W. H. I d .Eng. Chem. Fundam. 1977,16,443-451. (27) Chen, L. C.; Peoples, S. M.; McCarthy, J. F.; Amdur, M. 0. Atmos. Environ. 1989, 23, 149-154. (28) Stumm, W.; Kummert, R.; Sigg, R. Croat. Chem. Acta 1980, 53, 291-312. (29) Iler, R. K. The Chemistry of Silica;John Wiley: New York, 1979; pp 650-704. Envlron. Sci. Technol., Vol. 26, No. 7, 1992
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(30) Kotrily, S.; Sucha, L. Handbook of Chemical Equilibria in Analytical Chemistry; Ellis Horwood: Chichester, UK, 1985. (31) Veprek-Siska, J.; Wagnesova, D. M.; Eckschlangger, K. Collect. Czech. Chem. Commun. 1966, 31, 1248-1255. (32) Byerley, J. J.; Fouda, S. A.; Rempel, G. L. J. Chem. SOC., Dalton Trans. 1975, 1329-1338. (33) Graedel, T. E.; Weschler, C. J. Rev. Geophys. Space Phys. 1981, 19, 505-539. (34) Parks, G. A. Chem. Rev. 1965, 65, 177. (35) Schindler, P. W. In Adsorption of Inorganics at SolidLiquid Interfaces: Anderson, M. A., Rubin, A. J.,Eds.; Ann Arbor Science: Ann Arbor, MI, 1981; pp 1-49. (36) Huang, C. P. In Adsorption of Inorganics at Solid-Liquid Interfaces;Anderson, M. A., Rubin, R. J., Eds.; Ann Arbor Science: Ann Arbor, MI, 1981; pp 183-217. (37) Sillen, L. G.; Martell, A. E. Stability Constants of Metal Ion Complexes;Special Publications; The Chemical Society: London, 1964; No. 17; 1970; No. 25. (38) Shenk, J. E.; Wefer, W. H., Jr. J.-Am. Water Works Assoc. 1968, 60, 199-203.
(39) Laidler, K. J. Chemical Kinetics, 3rd ed.; Harper and Row: New York, 1987; pp 229-275. (40) Srivastava, R. D.; McMillan, A. F.; Harris, I. J. Can. J. Chem. Eng. 1968,46, 181-184. (41) Betterton, E. A.; Hoffmann, M. R. J. Phys. Chem. 1988, 92, 5962-5964. (42) Berg, J. V. D.; Dillon, A. J. V.; Meijden, J. V. D.; Geng, J. W. In Surface Properties and Catalysis by Non-Metals; Bonnalle, J. P., Derauane, E.; Eds.; Reidel Publishing Co.: Drodrecht, The Netherlands, 1982; pp 493-532. (43) Raitert, V. A.; Golodets, G. J.; Pyatnitzkii, Yu. L. Proceedings of the 4th International Congress on Catalysis; Moscow Academiai Kiado: Budapest, 1971; p 466. (44) Prasad, D. S. N. Ph.D. Thesis, University of Rajasthan, Jaipur, India, 1990. (45) Berresheim, H.; Jaeschke, W. J. Atmos. Chem. 1986, 4, 311-333.
Received for review October 7,1991. Revised manuscript received November 15,1991. Accepted February%, 1992. This work was supported by an Indo-U.S. Subcommission Research Project.
Rare Earth Distributions in Catalysts and Airborne Particles Michael E. Kltto,+ David L. Anderson,*9$Glen E. Gordon, and Ilhan Olmez*
Department of Chemistry and Biochemistry, University of Maryland, College Park, Maryland 20742 Zeolite cracking catalysts used by petroleum refineries were analyzed for 38 elements. Concentration patterns of rare earth elements (REEs) in 10 zeolite catalysts show an enhancement of light REEs relative to the crustal abundance pattern, resembling those measured in refineries emissions. Release of zeolite catalyst material from fluidized catalytic crackers and incorporation of zeolite catalysts into refined oil provide new atmospheric elemental signatures for tracing emissions from refineries and oil-fired power plants, respectively. Though both have enhanced La/REE ratios, emissions from these two sources can be distinguished by their La/V ratios. Three-way catalytic converters of newer automobiles contain REEs and may, thus, be a significant source in some cities.
Introduction Catalysts. With the introduction of fluid catalytic crackers (FCCs) in the 1940s, traditional acid-treated clay catalysts were replaced by more stable synthetic silicaalumina gels. Recognizing the need for improved properties (activity, selectivity, and stability) in the catalysts, research efforts led to the invention of zeolite cracking catalysts in the late 1950s (1). Over 90% of the catalytic cracking units in the United States were using zeolite catalysts by 1969 (2). Cracking catalysts containing rare earth mixtures are consumed by the petroleum refining industry to produce light-weight hydrocarbons, such as gasoline and fuel oil (3). These zeolite cracking catalysts dominate the worldwide FCC demand, with the amorphous aluminosilicates still comprising a significant share of FCC usage in Europe and the Middle East ( 4 ) . The United States is responsible for -70% of global FCC zeolite consumption. Data from the ‘Present address: -Wadworth Center for Laboratories and Research, New York State Department of Health, Albany, NY 12201. Present address: FDA Laboratory, NIST, Building 235/B125 Gaithersburg, MD 20899. 8 Present address: Nuclear Reactor Laboratory, Massachusetts Institute of Technology, Cambridge, MA 02139. 1388
Environ. Sci. Technol., Vol. 26, No. 7, 1992
US.Bureau of Mines ( 5 , 6 )show the correlation between rare earth oxide (REO) demand and REO-containing FCC utilization (Figure 1). Petroleum catalysts have accounted for an average of 40% of the rare earth element (REE) consumption in the United States in the last 20 years, varying from a high (65%) in 1983 to a low (31%)in 1986. The recent decrease in use of REOs in petroleum catalysts is a result of the phase-out of leaded gasoline. Because of the need to replace alkyllead compounds with hydrocarbons of higher octane number, refiners are now using catalysts that contain reduced concentrations of REOs 0.5-2% (7). Zeolites typically comprise -15% of an FCC catalyst’s composition, the remainder being the aluminosilicate structure and an inert filler. A cavity-filled structure gives zeolites their cation-exchange and reversible dehydration properties (8). Synthetic zeolites used by the petroleum industry are somewhat distinguishable by their Si/ Al ratio, which is important in determining the hydrothermal stability for cracking (9, 20). Alkali and alkaline earth metals present in crude oils deactivate zeolite catalysts by poisoning (acid-site neutralization) and degradation (hightemperature fluxing). Sources of metal impurities and their effect on catalyst activities have been reviewed by Letzsch and Wallace (11). Recently developed synthetic zeolites have greater tolerances (up to several thousand ppm Na), while also upgrading octane numbers and limiting coke production (22). Cracking-catalyst activity levels depend on the accumulation of product coke on the catalyst. Thus, selective cracking and control of the zeolite’s pore size (to minimize coking deactivation) have dominated research interests. FCCs are circulated through a regeneration stage, using hot air and steam, to remove coke deposita prior to their reuse. However, they gradually lose their effectiveness and must be withdrawn from circulation and replaced with fresh catalyst. Also, as noted below, some of the catalyst material apparently gets into the fuel burned by oil-fired plants. Bastnasite and monazite ores are the major commercial sources of REEs (13,14). The former is enriched in light
0013-936X/92/0926-1368$03.00/0
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