Trace metal catalvsis in aauatic environments Once an inexplicable nuisance to chemical kineticists, it may now play a role in pollution control
Michael R!. Hoffmann W. M . Keck Laboratories California Institute of Technology Pasadena, Gal$ 91 125 Current research in the fate of pollutants in aquatic environments has been focused on pathways for pollutant transformation. Important pathways that have been identified for major organic and inorganic pollutants include redox reactions, autoxidations, and hydrolysis reactions. Many reactions in these general categories can be catalyzed by specific homogeneous or heterogeneous tramition metal complexes (1). Today, large-scale synthetic processes such as the autoxidation of acetaldehyde to acetic acid depend on the catalytic inlluence of soluble transition metals and specific organometallic complexes ( 2 ) . Historically, trace metal catalysis was an inexplicarble nuisance phenomenon for the kineticist. Impurities in solvents, reagents, and reactants frequently resulted in nonreproducible rates of chemical reactions and unusual empirical rate laws. Trace metal concentrations as low as 0.1-10 n M influence the rates of some peroxide reactions (3). For example, the decomposition of peroxymonosulfuric acid, a suggested intermediate ( 4 ) in the autoxidation (of SO2 dissolved in aqueous aerosols, is accelerated significantly by nanomolar amounts of some divalent first-row transition metals ( 5 ) . Similarly, ferrous ions promote the oxidat ion of a wide variety of aromatic substrates by hydrogen peroxide ( 6 ) .Other chemical reactions involving oxidationfreduction, hy0013-936X/80/0914-1061$01 .OO/O
drolysis, decarboxylation, transamination, and bromination have been shown to be sensitive to trace metal catalysis (7). The principal objectives of the article are to explore the possible roles that trace metals may play in the catalysis of aquatic reactions; to focus specifically on the kinetic and mechanistic aspects of hydrolysis, redox, and autoxidation reactions; and, finally, to explore the applicability of certain catalytic reactions for pollution control. The word “catalyst” has been defined in many ways. Ostwald (8) originally defined a catalyst as a substance that changes the speed of a chemical reaction without undergoing a chemical change itself. Bredig ( 9 ) expanded the definition to include substances that may be chemically altered but are not involved in a whole number stoichiometric relationship among reactants and products. Bell ( I O ) subsequently defined a catalyst for a homogeneous reaction to be a substance that appears in the rate expression with a reaction order greater than its stoichiometric coefficient. Regardless of its current colloquial definition, the common notion is that the function of a catalyst is to lower the activation energy of a given reaction. As illustrated in Figure 1, the catalyst in a homogeneous system will alter the reaction mechanism to one having a lower activation energy without changing the point of equilibrium for a reaction as predicted by the law of microscopic reversibility (3).Since the catalyst is involved in the rate-determining step, and consequently the molecular composition of the activated
@ 1980 American Chemical Society
complex (1I ) , a discussion of catalytic activity requires consideration of detailed reaction mechanisms. Therefore, the role of trace metal catalysis in aquatic reactions will be discussed in terms of postulated reaction mechanisms which are consistent with kinetic observations.
Hydrolysis reactions An important pathway for the transformation of organic esters in aquatic environments is the reaction known as hydrolysis or solvolysis. Most often hydrolysis reactions in natural systems are sensitive to specific acid or base catalysis ( 1 2 ) . In acid-catalyzed ester hydrolysis, the role of the proton, according to Bender and Brubacher (13), is to provide a reaction pathway of lower energy by withdrawing electrons and therefore weakening the bond to be broken. A mechanism for acid catalysis has been suggested by Bender ( 1 4 ) : 0
(1)
II
+ H+ Z
OH
(2)
R-C-OR’
I R-C-OR’
OH
I
R-C-OR’
OH
+ HzO e
I
R-C-OR’
I
OK+ OH
I
R-C-OR’
I OHz’
-
(3)
EOZH
+ R’OH + H+
where Equation 3 represents a summation of two steps involving proton transfer in addition to bond breaking. Volume 14, Number 9, September 1980
1061
FIGURE 1
Effect of a catalyst
Activation Activation energy, no catalyst with catalyst energy,
Reactants
AH of reaction
Products
for the rapid hydrolysis of a-amino acids, S, is shown in Figure 2 where there is a direct complexation of Cu2+ by the ester group. In general, the effectiveness of metal ion catalysts parallels their formation constants for complexation with the same substrate (i.e., kh CY p). Decarboxylation reactions are also catalyzed by a variety of metal ions. Prue (17) investigated the metalion-catalyzed decarboxylation of acetone-dicarboxylic acid:
Reaction coordinate
FIGURE 2
Hydrolysis reactions
\ / CU
/ \
\ /
H-2 0
cu
/ \
FjH2 0
"2
I
I II RCH -C - OR'
0
RCH -C
I
I
-OR'
OH2 cu "2
I
OH
1 I
RCH- C-OR'
-I _t
OH
NHz
RC-C-OH+R'OH
II
0
Proposed mechanism for the Cu2+ -catalyzed hydrolysis of phenylalanine ethyl ester is shown along with the cata!ytic order over the pH range of 7-8.
1062
Environmental Science & Technology
Positively charged metal ions can function effectively as Lewis acids in these reactions and can be expected to catalyze reactions which are similarly catalyzed by Bronsted acids. For simple esters the proton is a more efficient catalyst due to its higher charge density. However, metals appear to be more catalytically active with substrates containing additional functional groups such as amino acid esters. An important factor in metal catalysis appears to be the ability of the metal ion to coordinate with the substrate and remove electron density from a site subject to bond breaking or nucleophilic attack (7). For example, Co(I1) and Cu(I1) have been shown to be effective catalysts for the hydrolysis of glycine methyl ester in the p H range of 7-8 ( 1 5 ) ;likewise Cu(I1) proved to be more effective than either H+ or OH- for the catalytic hydrolysis of phenylalanine ethyl ester (16). A mechanism consistent with the observed rate law:
0
II
HO+32H&XH$OzH
4
0
I
CHG-CH,
+ EOz
(5)
over a wide p H range and found that the deprotonated acid, A2-, was most strongly catalyzed by metal ions. H e found a linear free energy relationship between the rate constant for the catalytic pathway and the dissociation constant for the corresponding metal-malonate complex. In this case the activated complex was postulated to be similar to the bidentate chelate, malonate. The order of catalytic activity was reported as: Cu2+ > Ni2+ > Zn2+ > Co2+ > Mn2+ > Cd2+. Similar catalytic effects were reported by Gelles and Salama (18) for the decarboxylation of oxaloacetic acid for which the following order of catalytic activity was observed: Cu2+ > Zn2+ > Co2+ > Ni2+ > Mn2+ > Ca2+.
Metal ions catalyze other hydrolysis reactions ( 7 ) which may be of importance environmentally, such as the hydrolysis of tripolyphosphate: P30105-
+ H20
-+
HP2073-
HP042-
(6)
Catalytic activity is again correlated to the ability of a particular metal to form a complex with the substrate. In aqueous solutioin a t pH 7-9 the tripolyphosphate anion is extremely stable; however, with the addition of Cu2+, Zn2+, or Pb2+, the hydrolysis reaction is accelerated greatly ( I 9).
Redox reactions Transition metals that have several stable oxidation states are frequently active as catalysts involving electron transfer. Catalysis i n many redox reactions results from a special combination of relative rates when an active metal exists in more than one oxidation state ( f 1). The catalytically active metal can be oxidized and reduced cyclically a t a rapid rate with a concomitant catalytic influence on a slower reaction. Redox reactions such as the oxidation of Cr(II1) 1.0 Cr(V1) by S 2 . 0 ~ ~ are often thermodynamically favorable but kinetically slow ( 2 0 ) . The oxidation of Cr(II1) proceeds according to the following ftoichiometric equa- The independence of the reaction rate tion: on the concentration of Cr(II1) suggests that the product of the de2Cr3+ 3s20g1- 7 H 2 0 Cr207 6S04214H+ ( 7 ) composition of the activated complex, AgS208-, is the active oxidant. with a AGO = -530 kJ/mol. In the From an environmental standpoint, absence of catalytic influences the re- the oxidation of Cr(II1) by 0 2 is a reaction rate is imperceptible. Upon the action that is thermodynamically faaddition of Ag+ the reaction proceeds vored under certain p H conditions but a t a measurable rate according to the kinetically hindered. Conceivably, this following rate law ( 2 1 ) : reaction is also sensitive to trace metal catalysis. If so, the notion of safe disposal of Cr(II1) in aquatic systems must be examined in light of its potential catalytic conversion to toxic (8) Cr(V1) under suitable oxidizing conFor redox reactions, the rate law is ditions. often first order in oxidant and zero Other redox reactions, such as the order in reductant when catalyzed by oxidation of oxalic acid by chlorine, are trace metals. The best evidence that sensitive to catalysis by metals with trace metals are involved in a reaction multiple oxidation states. Mn(II1) mechanism is an unusual empirical catalyzes the oxidation as follows: rate law (3). 81 For the oxidation of Cr(III), Yost Mn3+ C ~ 0 4 ~ - MnC204+ (21) proposed the following alternative rate-determining steps: ( 1 1)
+
+
+
+
+
S20s2-
+ AgfF-fkk- 1l 2S042-
MnC204+
+ Ag3+
(9)
or k S20g2- Ag+&
+
+
.k-1
S042- .SO4-
k
Slow
Mn2+
+ C 0 2 + COz-.
co2-. + c12
+ Ag2+
(10)
CI-
+ Mn2+
+
-
c02
(12)
+ c1. + c1-
Mn3+
+ C1-
(13) (14)
+ c12
+
2c02
+ 2cl-
(15) where Equation 15 represents the overall stoichiometry or the summation of Equations 1 1- 14. This reaction sequence is an example of catalysis by a metal that has two stable oxidation states differing by two electrons (Le., Mn(I1) and Mn(1V)). The intermediate oxidation state, Mn(III), is unstable with respect to disproportionation ( 2 2 ) with K lo9:
+
-
2Mn3+ 2 H 2 0 e Mn2+ MnO2
+
+ 4H+
(1 6)
Although Mn3+ is highly unstable, it may be sufficiently persistent under suitable conditions to expect a significant catalytic influence. Another mode of metal catalysis for redox reactions is the phenomenon of induced catalysis. This occurs when a relatively fast reaction between an oxidant and reductant forces a slow reaction between one of the former reactants and another substance ( 3 , 11). Again, the reaction of oxalic acid and chlorine will be used to illustrate this catalytic phenomenon. The oxidation of C ~ 0 4 ~by- Cl2 is accelerated by the reaction of Fe(I1) and Cl2. A mechanism proposed by Taube ( 2 3 ) is as follows: Volume 14, Number 9, September 1980
1063
- + + + + c204-.+ + + - + + - + + +
Fe(I1) Cl2-
+ Cl2
HC204H + 2C1Cl2
Fe(II1)
C12(17)
C204-a
( 1 8)
2co2 (19) Cl2 c204- c12- 2 c o 2 (20) c2042c12c02 Cl2(21) Cl2Cl22c1Cl2 (22) This catalytic scheme differs fundamentally from the Cr(III)/S20g2-/ Ag+ reaction in that the apparent catalytic metal, Fe(II), is not regenerated. Metals such as Cu(II), Ni(II), and Co(I1) also exert a positive catalytic effect on the above reaction. c12-
Autoxidation reactions The term autoxidation refers to general reactions of oxidizable materials with molecular oxygen although earlier definitions were more restrictive. In general, reactions of triplet ground-state oxygen with singlet spin-state reductants are slow in the absence of catalytic influences because they involve a change in spin multiplicity and because a significant amount of bond deformation or alteration of a permanent nature may occur. Most often autoxidations are undesirable for commercial materials because of resulting damage and reduced lifetimes. Autoxidations are greatly accelerated by trace metals such as C u ( I I ) , Co(II), F e ( I I ) , Mn(II), Ni(II), and many of their coordination complexes. The autoxidation of H2S and HSin both fresh and seawater systems has been shown to be sensitive to trace metal catalysis by Mn(II), Cu(II), Fe(II), Ni(II), and Co(I1) (24, 25). A general characteristic of trace metal catalysis, as mentioned above, is the occurrence of nonreproducible reaction rates and unusual empirical rate laws. The kinetics of the autoxidation of sulfide have been studied by a number of investigators under slightly different experimental conditions. These results are summarized in Table 1. From comparison of the reported half-lives in Table 1 as measured by different investigators, it is obvious that the reaction may be sensitive to a variety of catalytic influences in addition to trace metals. These include general base, microbial, and surface catalysis. The reported half-lives vary from days to minutes under similar concentration conditions and the observed kinetic orders vary from one investigator to another. 1064
Environmental Science & Technology
Oxidation
503‘-
Termination
so; so;
f
so; + so; +
k4
2s04‘-
5_ k
S2O62-
+
0 2
ks_
S2O62-
+
202
Radical mechanism rate laws Thermal reaction k 2 slow;
k5
Termination
l’[~Y1’*~SO~z~CO~l~L
(73
(1) v = d [SO&] = k z k l k dt k 3 slow; k6 Termination (2) v
=
k3
Photochemical reaction (3)v
(2) lI2
[M+]
”* [S032-] 3’2
=(y)
[S03’-l [02)lfZ
lr*
Hoffmann and Lim ( I ) have studied the catalytic autoxidation of sulfide over the p H range 5-1 2. The catalytic effects of Co(II), Ni(II), and Cu(I1)-4, 4’, 4”, 4”’-tetrasulfophthalocyanineas shown in Figure 3 were determined. The reaction kinetics were characterized in terms of a rate law of the general form:
which reduces to a rate law with an apparent zero-order dependence on [02] when [O,] >> [HS’] > KC and K B > [HS-1:
v = k[Catalyst][HS-]
(24)
Mechanistically, the rate law of Equation 23 combined with spectroscopic information indicates that the reaction proceeds via the formation of a tertiary activated complex in which 0 2 and HS- are reversibly bound to the metal complex as shown below: 81
Co(1I)TSP + 0 2 $ CO(III)TSP-02. Co(III)TSP-02.
(25)
+ HS-t..
$2
CO(1I)TSP
+ HS02(26) k
+ HS- e Slow Co(1I)TSP + HS02-
Co(III)TSP-02.
(27)
HS02-
+0 2
k’
HS04-
(28)
Rapid
The rate constant, k, over a wide range of pH was found to be 1.0 f 0.2 X lo5 M-2 min-I with a catalytic turnover number > 100 000. For example, the half-life for the uncatalyzed reaction is approximately 50 h, whereas the
half-life when [CoTSP] = 0.1 nM was reduced to 5 min ( I ) . The mechanism shown above is a nonradical enzymelike mechanism. Frequently, autoxidation reactions proceed via free radical reactions initiated by a suitable metal catalyst, and their rate laws are characterized often by nonintegral reaction orders in catalyst concentration. Catalytic autoxidation of SO2 by transition metals dissolved in submicron water aerosols has been suggested ( 4 , 24-28) as a viable nonphotolytic pathway for the rapid formation of sulfuric acid in relatively humid atmospheres. However, present experimental evidence in the laboratory and field shows considerable disagreement on the relative rates of SO2 transformation to H2SO4 in microdroplets. Inconsistent reaction rates, rate laws, and pH dependencies have been observed for aqueous solution kinetic studies of the reaction of S02eH20 and oxygen as pointed out by Hegg and Hobbs (27), Larson et al. ( 4 ) , and
Beilke and Gravenhorst (28). Redox reactions sensitive to homogeneous trace metal catalysis exhibit reaction rate laws that are often first order in one reactant and zero order in the other. The autoxidations of SO2 and H2S in aqueous solution fall into this general category of redox reactions. For example, oxygen concentrations in water are reduced to near zero levels in mass-transfer studies of aeration equipment by the addition of S032and an appropriately soluble Co(I1) salt. The reduction of oxygen proceeds according to the stoichiometry: HzO * SO2
+ ' / ~ O L H2S04
(29) In one experimental study ( 2 9 ) , a nonintegral rate Ilaw, with a zero-order dependence on oxygen
-d[HzO
+
SO*]/dt = U M
= ~ M [ C O ( I I ) ] : ~ [ S O ~ ~ -(30) ]~
was observed for the catalytic effect of hexaquo Co(I1) A suitable mechanism for the thermal initiation of a free radical reaction is shown in Figure 3 based on the original mechanism proposed by Backstrom ( 3 0 ) . Using the steady-state and long-chain length approximations for a free radical, closed-sequence catalytic process (31), two different rate laws are predicted. When the k3 step is assumed to be the rate-determining; step and kg the termination step, t'he rate law given in Equation 2 of the figure is obtained mathematically. However, when the reaction is initiated by UV radiation or the thermal reaction is limited by the k2 step, different rate laws are predicted. Laboratary studies of the liquid-phase oxidation of SO2 by 0 2 have suggested many different rate laws, reaction rates, and p H dependencies under similar conditions. In general, the majority of experimental results show that the homogeneous aqueous-phase reaction rate has a n overall reaction order of 3 with first-order dependencies on [M+], [S03*-], and [H+]. Some free radical mechanisms postulated by various investigators are inconsistent with experimentally determined reaction orders. Bassett and Parker (33) reported evidence for forlmation of a discrete complex of manganese, oxygen, and two sulfite ions in their study of the autoxidation of SO3*-. Matteson et al. (32) also postulated a n intermediate complex of 02-Mn(II)-S02 for SO2 oxidation in aqueous aerosol droplets. In both these studies, the anticipated radical termination product, dithionic acid, was not observed. These results suggest an enzymelike pathway given
in Equations 25-28. A hypothetical mechanism involving the nominal activation of molecular 0 2 by a suitable metal catalyst followed by complexation and oxidation of sulfite would be as follows:
+ O ~ ~ C O ( I I I ) O *(31) . BI Co(III)O2 - + S032Bl
Co(I1)
-
C o ( I I I ) 0 2 . SO3*- (32)
-
C o ( I I I ) 0 2 SO3*-
k3
+ 0 2 CO(II)-O2 c0(11) + RH ~ C ~ ( I I ) R H C0(11)02 + Co(1I)RH + Co(I1) + Co(II1)R + HO2 Co(I1)
P1
k
k
C0(11)02 R H + Co(I1I)R
+ H02.
Similar Co(I1)-catalyzed mechanisms for the autoxidation of benzylCo(1II)O SO42- (33) thiol (37),cumene ( 3 8 ) ,acrolein (38), and 2-mercaptoethanol(39) have been P4 Co(1II)O SO4*- S032proposed. Certain transition metal complexes are capable of reversibly 2 - s 0 3 C ~ ( I I I ) 0 So42- (34) binding and activating molecular oxygen ( 4 1 - 4 3 ) . The commonly ack5 2 - S 0 3 c ~ ( I I I ) OSO4*~ cepted mode of 0 2 attachment to Fast Co(I1) complexes involves dioxygen Co(I1) 2 s 0 4 2 - (35) apically bound in a bent configuration where Equation 33 would be the rate- in which a d-electron from Co(I1) is determining step. Using a determinant transferred to 0 2 , producing 0 2 - bound procedure for a system of nonhomo- to Co(II1). It is this attached supergeneous linear equations obtained oxide that is thought to be the active from steady-state considerations (35), catalytic site. Catalytic activity of metal a rate law similar in form to Equation phthalocyanines in aqueous solution 23 is obtained. This equation can be has been reported by Cook (13 , 4 4 ) for reduced under certain conditions to give a rate expression which is zero the autoxidation of HI and benzaldeorder in oxygen and first order in both hyde and Wavnerova and co-workers ( 4 5 , 46) for the autoxidation of hycatalyst and sulfite concentration. droxylamine and hydrazine. Sheldon and Kochi ( 3 5 ) have summarized a sequence of reactions that may be involved in the catalytic au- Pollution control applications toxidation of various hydrocarbons. Transition metals and associated For a generalized hydrocarbon, R H , a probable sequence of simultaneous complexes may prove to be useful from a commercial standpoint for improved reactions would be as follows: M2+ 0 2 e M3+-02. (36) pollution control systems. The catalyzed autoxidation of H2S and SO2 ~ 3 + - 0 ~ RH . and other reduced sulfur compounds M 2 + R* HO2. (37) in the presence of homogeneous (47, 48) or heterogeneous ( 3 9 )organomeM3+-02 R H tallic complexes may provide conveM~+-o*H R. (38) nient and economical methods for sulfur pollution control. The principal M3+-02 RH products of the metal phthalocyanM3+ R- H02 (39) ine-catalyzed autoxidation of H2S and or mercaptans in sour refinery distillates were reported to be colloidal sulfur, M3+-02 * R H M3+ R- HOz- (40) polythionates, and sulfate (49). Conceivably, similar catalytic systems where M 2 + is a catalytically active could be used for the autoxidation of metal. A characteristic of most au- selected organic pollutants for industoxidations is the formation of reactive trial wastewater treatment. peroxide intermediates as shown in Improved catalytic ability, elimiEquations 37-40. nation of recovery problems, and In commercial autoxidations, M*+ longer catalyst lifetimes may be is frequently Co(I1) or an associated achieved with supported organomeorganometallic complex. For example, tallic catalysts (50-52). Mass et al. Uri (36) proposed a mechanism for (39) have reported on enhanced cataCo(I1) stearate catalysis of the au- lytic autoxidation of mercaptoethanol toxidation of methyl linoleate that is by Co(I1) tetraminophthalocyanine consistent with the generalized se- attached to cross-linked polyacrylquence above: amide. Autoxidation of Klaus plant Slow
+
+
+
-+ -+
+
+
+
-
-+ + - + + + +
+ +
Volume 14, Number 9, September 1980
1065
(30) Backstrom, H. Z. Physik. Chem. 1934, 258, 122. (31) Boudart, M., "The Kinetics of Chemical
FIGURE 4
A possible application
L
Effluent
4
+
Influent
SO?
-
effluents or SO2 in flue gases by supported organometallic complexes such as the phthalocyanines may provide an economical method for emission reduction. Figure 4 illustrates a hypothetical countercurrent reactor with fixed beds of solid supported catalyst that could be used for SO2 scrubbing with the production of H2S04 as an alternative to limestone slurry scrubbers which produce a n unusable solid reaction product. Research along these lines is currently underway in this laboratory.
References (1) Hoffmann, M. R.; Lim, B. C. Enuiron. Sci. Technol. 1979,13, 1406. (2) , . Parshall. G. W. Science 1980. 208. 1221. (3) Edwards, J. 0. "Inorganic Reaction
New
York,
1965. (4) , , Larson. T. V.: Horike. N. R.: Halstead. H. Atmos. hnuiro'n. 1978,12, 1597. (5) Ball, D. L.; Edwards, J. 0.J . Phys. Chem. 1958,62, 343. (6) Walling, C. Acct. Chem. Res. 1975, 8, 125. (7) Basolo, F.; Pearson, R. G. "Mechanisms
of Inorganic Reactions," 2nd ed.; Wiley: New York, 1967. (8) Ostwald, W. Physik. 2. 1902,3, 313. (9) Bredig, G. Ulmann's Enzyke Technol. Chem. 1909,6, 670. (10) Bell, R. P. "Acid and Base Catalysis," 2nd ed.; Cornel1 University Press: Ithaca, 1971. 1066
(11) King, E. L. "Catalysis in Homogeneous
Reactions in a Liquid Phase," In "Catalysis: Fundamental Principles"; Emmett, P. H., Ed.; Reinhold: New York, 1955; Vol. 11, Part 2. (12) Wolfe, N. L.; Zepp, R. G.; Paris, D. F. Water Res. 1978,12, 561. (13) Bender, M. L.; Brubacher, L. J. "Catalysis and Enzyme Action"; McGraw-Hill: New York, 1973. (14) Bender, M. L. Chem. Revs. 1960, 60, 53. (15) Kroll, H. J . Am. Chem. SOC.1952, 74, 2036. (16) Bender, M. L.; Turnquest, B. W. J. Am. Chem. SOC.1957, 79, 1889. (17) Prue, J. E. J. Chem. SOC.1952,2331. (18) . . Gelles. E.: Salama. A. J . Chem. SOC. 1958,3689. (19) Thilo, E. Aduan. Inorg. Chem. Radiochemistry 1962,4, 31. (20) Fronaeus, S.; Ostman, C. 0. Acta. Chem. Scand. 1956,10, 320. (21) Yost, D. M. J . Am. Chem. SOC.1926,48, 152. (22) Cotton, F. A,; Wilkinson, G. "Advanced '
Benjamin:
1951,1540. (34) Laidler, K. J.; Bunting, S. J. "The
Chemical Kinetics of Enzyme Action"; Clarendon Press: Oxford, England, 1973. (35) Sheldon, R. A,; Kochi, J. K. Oxidation Combustion Revs. 1973,5, 135. (36) Uri, N. Nature 1956,177, 1177. (37) Dance, I. G.; Conrad, R. C.; Cline, J. E. Chem. Comm. 1974,13. (38) Hara, T.; Ohkatusu, Y.; Osa, T. Bull. Chem. SOC.Japan 1975,48, 85. (39) Mass, T. A.; Kuijer, M.; Zwart, J. Chem. Comm. 1976,87. (40) Collman, J. P. Accts. Chem. Res. 1977, 10, 265. (41) Basolo, F.; Hoffmann, B. M.; Ibers, J. A. Accts Chem. Res. 1975,8, 384. (42) Jones, R. D.; Summerville,D. A,; Basolo, F. Chem. Reus. 1979, 79, 139. (43) Cook, A. H. J . Chem. SOC.1938,1761. (44) Cook, A. H. J. Chem. SOC. 1938,1768. (45) Wagnerova, D. M.; Schertnerova, E.;
A potential application of polymer-supported organometallic catalysts for SO2 stack-gas scrubbing is illustrated above. The active catalyst, C, is supported on appropriate solid supports and piaced in fixed beds in a countercurrent flow reactor.
Mechanism";
Processes": Prentice-Hall. Inc.: Enelewood Cliffs, N.J:, 1968. (32) Matteson. M. J.: Stober. W.: Luther, H. Ind. Eng. Chem. Fund. 1969,8, 677. (33) Bassett, H.; Parker, W. G. J . Chem. SOC.
Environmental Science & Technology
Inorganic Chemistry"; Wiley-Interscience: New York, 1972; p. 849. (23) Taube, H. J . Am. Chem. SOC. 1946,68, 611. (24) Krebs, H. A. Biochem Z. 1929, 204, 343.
(25) Chen, K. Y.; Morris, J. C. J. Sanit. Eng. Diu., ASCE, 98, 215. (26) Cheng, R. T.; Corn, M.; Frohliger, J. 0. Atmospheric Enuiron. 1971,5, 987. (27) Hegg, D. A,; Hobbs, P. V. Atmospheric Enuiron. 1978,12, 241. (28) Beilke, S.; Gravenhorst, G. Atmospheric Enuiron. 1978,12. 231. (29) Chen, T.; Barron, C. H. Ind. Eng. Chem. Fund. 1972,11, 466.
Veprek-Siska, J. Collect. Czech. Chem. Commun. 1974,38, 3096. (46) Wagnerova, D. M.; Schertnerova, E.; Veurek-Siska. J. Collect. Czech. Chem. Commun. 1973,38, 756. (47) Urban, P.; Rosenwald, R. H., U.S. Patent 3 530 063, Sept. 23,1970. (48) Pikl, J., US.Patent 2 860 116, Nov. 11, 1958. (49) Urban, P. Chem. Abstr. 1962, 56,
P6251f. (50) Pittman, C. U.; Smith, L. R.; Hanes, R. M. J . Am. Chem. SOC.1975,97, 1742. (51) Robinson, A. L. Science 1976, 194, 1261. (52) Hartley, F. R.; Vezey, P. N. Adv. Organomet. Chem. 1977,15, 189.
b Michael R. Hoffrnann received his B.A. degree in chemistry f r o m Northwestern University in 1968 and his Ph.D. from Brown University in 1973 with a degree program in chemical kinetics. After receiving his Ph.D., Dr. Hoffmann was appointed a Research Fellow at the California Institute of Technology under the N I E H S postdoctoral training program. His research has been focused on applied environmental kinetics in the areas of chemical catalysis, oxidation of reduced surfur compounds, microbial catalysis, and ligand substitution. From 1975 to 1979, Dr. Hoffmann was an assistant professor of environmental engineering in the Department of C i d and Mineral Engineering at the University of Minnesota. In the summer of 1980, he moved to the California Institute of Technology where he was appointed an associate professor of encironmental engineering science.
