Environ. Sci. Technol. 1988, 22, 275-282
Kinetics and Mechanism of Autoxidation of 2-Mercaptoethanol Catalyzed by Cobalt( I 1)-4,4',4",4"'-Tetrasulfophthalocyanine in Aqueous Solution Plng-Sang K. Leung and Mlchael R. Hoffmann"
Environmental Engineering Science, W. M. Keck Laboratories, California Institute of Technology, Pasadena, California 9 1125 The kinetics of the catalytic autoxidation of 2mercaptoethanol to 2-hydroxyethyl disulfide in the preswere ence of cobalt"I)-4,4',4'',4'''-tetrasulfophthalocyanine investigated. The experimental rate law Y = -d[RS-] /dt = kob,d[CoTSP]~[RS-] was found, with lt31K21
Itobsd
+
k32K22K1'
aH+
=
+ k33K23KiK2/ 2 ..
au+
where k3iis the rate constant of the electron-transfer step of the ith catalytic center, Kl, and K2i are the equilibrium constants of oxygenation and substrate complexation of the ith catalytic center, respectively, and K{, K i , Ka(, and K,,' are the apparent first and second acid dissociation constants of the catalyst, HOC2H4SHand HOC2H4S-,respectively. A nonlinear least-squares fit of the experimental data to the above rate law gave values of k31K21 = (1.5 f 0.05) X lo4 M-l s-l, k32K = (1.3 f 0.1) X lo4 M-l d , k3&, = (7.5 f 0.2) X lo3 M-! s-l, pK{ = 10.8, and p K i = 10.97 a t 25 "C. The active catalytic center is a dimer that is bridged by 2-mercaptoethanol. Hydrogen peroxide and 2-mercaptoethanol radicals were identified as reaction intermediates. Introduction Oxidation by reagents such as molecular oxygen, chlorine, ozone, hydrogen peroxide, and permanganate is one of the primary modes by which pollutants are removed from industrial and domestic waste waters. With the exception of 02,chemical oxidants are commercially cost intensive in addition to being inherently hazardous under certain conditions. These factors invariably contribute to the cost of water treatment processes. In contrast to the highly reactive oxidizing agents, molecular oxygen usually requires the presence of a catalyst in order to insure complete and rapid reaction with reductants. Autoxidation reactions of many inorganic and organic compounds are accelerated in the presence of first-row transition metal ions and complexes in which the metal center has access to more than one stable oxidation state ( I , 2). A considerable amount of research has been directed toward determining the mechanisms of transition metal catalysis in the reactions of O2 with a wide variety of reduced sulfur compounds such as SO2 (3), H2S ( 4 ) , and mercaptans, RSH (5-7) in aqueous solution. Sulfur-containing compounds are ubiquitous contaminants in waste waters discharged from mining facilities, pulp and paper mills, tanneries, and oil refineries (8, 9). Furthermore, the catalytic oxidation of S(-11) and S(1V) in aquatic environments plays an important role in the natural sulfur cycle. We have previously examined the catalytic properties of homogeneous and heterogeneous phthalocyanine complexes of Fe(II), Mn(II), Co(II), Ni(II), and Cu(I1). Phthalocyanines are macrocyclic tetrapyrrole compounds that readily form square-planar coordination complexes 0013-936X/88/0922-0275$01.50/0
in which the metal atom is bonded to the four pyrrole nitrogen atoms. Structurally related porphyrin and Schiff bases form analogous complexes with first-row transition metals. The catalytic properties of divalent and trivalent metal-phthalocyanine, -porphyrin, and -Schiff base derivatives have been compared to those of catalase, peroxidase, oxidase, and oxygenase enzymes (10-14). In addition, these compounds represent suitable models with which to study the catalytic effects of trace metals in the aquatic environment, because of their similarities to the structure of naturally occurring pigments such as chlorophyll. Autoxidation of mercaptans (15) may occur as follows:
32102
+ RSH(3p,-3d,)
-
+
+ Hf
02- RS'
(1)
or 'A,02
+ RSH,,,J
-+
02- + RS'
+ H+
(2)
-
In the first reaction, the half-filled T ; orbital of groundstate oxygen overlaps with an excited sulfur atom (3p1 3d, transition) on the mercaptan, whereas in the second reaction, a direct overlap of empty T ; orbital of an excited singlet-state oxygen with the filled 3p, orbital of the sulfur atom is required. Both reactions have relatively large activation energies. Cobalt(I1)-tetrasulfophthalocyanine (CO~TSP) has been reported to be an effective catalyst for the autoxidation of 2-mercaptoethanol(16) and cysteine (17). However, the reaction kinetics are understood only at a qualitative level. Simonov et al. (18) proposed a linear free-radical mechanism for the catalytic autoxidation of simple mercaptans, while Dolansky and Wagnerova (5) have proposed an ordered ternary complex mechanism in which O2and RSH are simultaneously bound to the active catalytic center. In spite of these studies, no attempt has been made to describe in detail the stoichiometry, product distribution, reaction intermediates, and the effect of pH, ionic strength, and temperature on the kinetics of the catalytic autoxidation of mercaptans in aqueous solution. In this paper, we describe the detailed kinetics and mechanism of autoxidation of 2-mercaptoethanol as catalyzed by Co(I1)-
4,4',4'',4'"-tetrasulfophthalocyanine.
I
2-Mercaptoethanol and other mercaptans are often found in sour refinery wastes. The catalytic autoxidation of 2-mercaptoethanol and other mercaptans by Co"TSP may be used for the selective elimination of these malodorous and sometimes hazardous compounds from waste water. Experimental Procedures Reagents. 2-Mercaptoethanol stock solutions were prepared with redistilled 99% reagent-grade 2-mercaptoethanol (Aldrich). Co(")-"4',4'',4'"-tetrasulfophthalocyanine, abbreviated as CoTSP, was synthesized according to the procedure described by Boyce et al. (3). Buffers were prepared with reagent-grade sodium hydroxide (Mallinckrodt), sodium chloride (J. T. Baker), and sodium bicarbonate (Mallinckrodt). Water used to prepare the buffers and reagent solutions was obtained from a Milli-Q water purification system (Millipore), had a resistivity of
0 1988 American Chemical Society
Envlron. Sci. Technol., Vol. 22, No. 3, 1988
275
18 MQ cm, and was irradiated with ultraviolet light to remove any trace organics that may have been present. Kinetic Measurements. Kinetic measurements were made on a Hewlett-Packard Model 8450 spectrophotometer. A minimum of 100 data points was collected for each kinetic measurement. At least three measurements were made for each value of kobsd. Data were analyzed on-line with an IBM-XT computer. Constant temperature was maintained at 25.5 "C with a Haake water bath. A sodium chloride stock solution was added into the appropriate buffer to establish the desired ionic strength (g)at 0.4 M. pH was determined with a Beckman Altex 4 71 pH meter and Radiometer glass electrode. Dissolved oxygen levels were established by dispersing N2and O2gas mixtures into the CoIITSP-containing buffer. The reaction was monitored at 233 nm [the absorbance maximum for HOC,H4S(as)]. Pseudo-first-order conditions of [O,] >> [HOC2H4SH]~ < 2.2 X M were employed for all kinetic runs. Stoichiometry. Reaction stoichiometry with dissolved oxygen in excess was determined by measuring residual dissolved oxygen after 99 % of 2-mercaptoethanol had reacted in alkaline solution. Dissolved oxygen was analyzed by the azide modification iodometric method (19). Product Identification. Reaction products were extracted into chloroform. The extracted sample was concentrated to dryness under a stream of N2 gas. The residual was redissolved in 2 mL of chloroform and isolated by high-performance liquid chromatography (HPLC) (Hewlett-Packard1084B; reversed-phase C18column). The sovlent carrier was 15% methanol in deionized doubly distilled water. The oven temperature was set at 40 "C. Products were identified by comparison to standards. Intermediate Identification. Hydrogen peroxide or peroxide ion was identified by the horseradish peroxidase fluorescence method (20). Samples (2 mL) were withdrawn from the reaction vessel at 60-s intervals and were immediately neutralized to pH -7 with concentrated hydrochloric acid. Afterwards the fluorescence reagent, which consisted of potassium hydrogen phthalate, p-hydroxyphenylacetic acid, and peroxidase, was added to the sample. Before the actual measurement was made, sodium hydroxide was added to raise the pH to 110 to enhance the fluorescence intensity. The excitation wavelength was 350 nm; the fluorescence intensity was measured at X = 400 nm on a Shimadzu RF-540 spectrofluorophotometer. 2-Mercaptoethanol radical was identified by monitoring the formation of ascorbate radical in a reaction mixture containing ascorbic acid. Ascorbate is known to react with mercaptan radical, producing the corresponding ascorbate radical (RS' AHA'- RS- Ht) (21, 22), which absorbs light at X = 360 nm with €360 = 3300 M-' cm-'. Hydroxide Ion Identification. The hydroxide ion production was identified by continuously monitoring the pH of the unbuffered reaction solution.
+
-
+
IONIC STRENGTH = 0 . 4 M TEMP. :2 5 . 5 * C kObLd 17.6 x IO-'SEC-'
H
r 2 c o e f f . = 0.999 chi-SQUARE : 0 , 0 5 7 2 % OF REACTION F I T :91.06 INITIAL ABSORBANCE : 1 . 2 2 ABSORBANCE AT INFINITY :0.22
0.6
0.4
0.2
o 720
360
IO80
1440
1800
TIME (seconds1
Flgure 1. First-order kinetic plot of absorbance vs time for the oxidation of 2-mercaptoethanol by oxygen as catalyzed by Co(I1)-tetrasulfophthaiocyanine.
[ C O ~ T S P ] ; ~ . ~x 10-7t.4
4.2
-
[ HOCpH45-]0.3.3 TEMPERATURE 0
pHi10.5
0
pH'I3.O
x 10-4M 25.5'C
IONIC STRENGTH
3.4 -
s
0 40 M
IONIC STRENGTH: 0.45 M 0
0
'
t
3
OO
20
40
80
60
% OF
100
O2
Figure 2. Determination of the reaction order with respect to [O,] for the conditions given above.
IONIC STRENGTH i O . 4 M
-
TEMP,
:
25 5-C
+
Results Under pseudo-first-order conditions ([O,] >> [HOCH2CH,SHIT),plots of In ( A , - A,) vs t were linear (r21 0,999) for 50-90% of the reaction. This linearity of the pseudo-first-order relationship in mercaptoethanol indicates that the rate of autoxidation is first order with respect to thioethanol. Figure 1compares the calculated exponential fit with the actual kinetic data. The dependence of IzoM on [02J was determined at both pH 10.5 and pH 13.0 (Figure 2). Results of these experiments showed an independence of the rate on the concentration of dissolved oxygen. Plots of k o h d vs [CoTSPITfrom 0.64 to 2.6 pM as shown in Figure 3 yield straight lines (91 0.99) with 276
EXPERIMENTAL KINETIC OATA CALCULATE0 CURVE
Environ. Sci. Technol., Voi. 22, No. 3, 1988
oo!o
I
I
I
0.6
I .2
1.8
,
I
I
2 4
3.0
[ C O T S P ] ~(M)x106 h
Figure 3. Determination of the reaction order with respect to [Co"TSP], for the conditions given above.
slopes of 8795, 11288, and 7076 M-l s-' at pH 9.7,10.7, and 13.5, respectively. Thus, the rate of autoxidation is first order in both HOC2H4SHand catalyst but apparently zero order with respect to dissolved oxygen. In the absence of buffer, the pH, as shown in Figure 4, rose smoothly during the course of reaction when pHo -10. The increase of pH indicates that hydroxide ion is a product under these Conditions. The stoichiometry of the reaction under the conditions given in Table I was established to be 1mol of oxygen to 4 mol of HOC2H4SHas shown in eq 3, when dissolved oxygen was in excess with respect to the anticipated
9.0
=
7.2
IONIC STRENGTH :O 4 M
cn
COTS SF]^: 3 . 2 ~IO-'M
0
9.50
I
I*'.
p.0.4 M TEMPERATURE
0
25.5*C
:
[CoTSPIT = 3.2 x IO-'M
9'82 1
I
400
0
I
1200 TIME (seconds)
SO0
X
I
.O
CoTSP ADDED
10.25 -
-:
TEMPERATURE
5.4
3
25.5.C
d RS-+ASCORBIC ACID 12x10'4M1
I
I
(600
2000
Figure 4. Plot of pH vs tlme for the unbuffered autoxldatlon of 2mercaptoethanol. COTS SF]^
Table I. Stoichiometry Determination of Catalytic Autoxidation of 2-Mercaptoethanol by
Co(II)-4,4',4",4"'-Tetrasulfophthalocyaninea initial oxygen, mmol
final oxygen, mmol
6.8 X 1.36 X 1.36 X 6.4 X 1.36 X loe2 6.5 X 5.9 X 1.36 X
amt of oxygen used, mmol
amt of thioethanol reacted, mmol
ASbJ
6.8 X 7.2 X 7.1 X 10" 7.8 X
3.2 X 3.2 X 3.2 X lo-' 3.2 X
4.7 4.4 4.4 4.1
I
0.34
+ 2Hz0
I
ConTSP
0.321
-
- BEFORE
-.-
1
ACIDITION OF R S 1
0.0 SEC.
~ T .. C30.0 O OSEC. QEC.
,EC
9
0.16
___+
2(HOCzH,S)2
+ 40H-
(3)
An oxidizing radical was found to be an intermediate of the autoxidation. Figure 5 shows the increase of absorbance at X = 360 nm after the addition of ascorbic acid to the reaction mixture. Such an increase implies that the 2-mercaptoethanol radical HOC2H4S' is an intermediate of the reaction. Production of hydrogen peroxide, the other reaction intermediate, is shown in Figure 6. Therefore the overall 1:4 stoichiometry for pH >10 is given by the sum of 0 2 + 2HOC2H4S2H20 (HOCzH4S)z + H2Oz 20H- (4) and H2Oz + 2HOCzHdS(HOCzH4S)z + 20H- (5) The detailed kinetics and mechanism of the oxidation of thioethanol by hydrogen peroxide (i.e., eq 5) have been reported by Leung and Hoffmann (23). The change of the CoTSP spectrum in the visible region during the course of reaction is shown in Figure 7. ConTSP has two characteristics peaks at X = 626 and 668
+
WITH 02 AT I ATM.
p H : 11.4 IONIC STRENGTH = 0 . 4 M TEMP. = 25.5'C
A02
stoichiometric requirements. By comparison of chromatographic retention times with standards, the product of the catalytic autoxidation of 2-mercaptoethanol was determined to be 2-hydroxyethyldisulfide. Thus,the overall stoichiometry for pH >10 is 2
[ 0 2 ] 0 1 ~= ~SAT. ,
tI
#Experimental conditions: p = 0.01 M, pH 11.8, temperature = 25 "C, [COTSP]~ = 5.2 X lo4 M. bAS = amount of 2-mercaptoethanol reacted: AO, = amount of 0,reacted.
4HOC2H4S- + 0
= 1.3x 10-6
[HOCpH4S-]O- 4.6 x 10-4M
-
-
+
0 08
0 00 500
560
620
680
A 800
74c
WAVELENGTH (nm)
Flgure 7. Timedependent changes In the visible spectrum of Co'ITSP during the autoxidation of 2-mercaptoethanol.
nm that correspond to a dimer and to the sum of a monomer plus an oxygenated monomer, respectively (24). The increase of absorption at X = 626 nm and the concomitant decrease of absorption at X = 668 nm after the addition of the thioethanol suggests that the dimer concentration is enhanced during the reaction. The resulting dimer is likely to be bridged by the anion of 2-mercaptoethanol. The above kinetic information can be illustrated with the mechanism proposed in Figure 8. The rate of disappearance of 2-mercaptoethanol is given by the sum of four terms: v = -d[RSH],/dt = 2 ( ~ + l V Z + v3 + ~ 4 ) (6) Environ. Sci. Technol., Vol. 22, No. 3, 1988 277
RS'
t
RS'
-
RSSR
Figure 8. Schematic diagram illustrating the proposed mechanism for the catalytic autoxidation of 2-mercaptoethanoi.
where v1 = k31[(02-Co11TSP-RS--Co11TSP-RS-)-6], v2 = K32[ (02-Co"TSP-RS--Co1'TSP-RS-)-'], ~3 = k33[ (02CO~TSP-RS--CO~TSP-RS-)~], v4 = rate of oxidation from Hz02 cycle, and k3i = rate constant of the electron transfer step of ith catalytic center. The first, second, and third catalytic centers are formed by the combination of two protonated CoTSP monomers, a protonated CoTSP monomer and a deprotonated CoTSP monomer, and two deprotonated CoTSP monomers, respectively. Since [H202] is very small, u4 can be neglected; thereby, eq 6 becomes v = 2(Vl
+ v 2 + v3)
(7)
By using the total concentration of thiol, [HOC2H4SHIT = [HOC2H4SH]+ [HOC2H4S-]+ [-OC2H4S-],to express [RSH]T in term of [HOC2H4S-],eq 7 can be written as -d[RS-] 2 ( ~ 1 ~2 ~ 3 ) y = - (8) dt UH+ Kaz' 1+,+Ka1 a ~ + Substitution of the corresponding equilibrium relationships for the binding of molecular oxygen and complexation of the substrate by the first catalytic center into v1 yields
with respect to both [RS-] and [CoTSPIT1and that v1 is independent of [O,]. Similar expressions can be derived for both v2 and v3. Thus, under the assumption that Kli[02]>> KliKZi[O2][RS-] and Kli[02]>> 1,the total rate of autoxidation becomes v = ( ~ ~ ~ K ~ ~ [ C O+ T Sk32K22[CoTSP]TZ P]T~ + ~~~K~~[C~TSP]T /1~ + ) [aH+/Kal' RS-] Ka2//aH+ (12) With the total catalyst concentration expressed as [COTSPIT= [COTSPIT~ [COTSP]T~ + [COTSP]T~,eq 12 can be rewritten as
+
v =
+ +
v1
= I~~~K~~KZ~[(CO'ITSP-RS--COIITSP)-~] [O,] [RS-] (9)
where K,, and K2, are the equilibriumconstants for oxygen binding and substrate complexation by the first catalytic center, respectively. Expressing the concentration of the first catalytic center in term of the total concentration, [COTSP]T~= ~([(CO"TSP-RS--CO~'TSP)-~] + [(OzCO"TSP-RS--CO"TSP)-~] [ (02-Co"TSP-RS-COI'TSP-RS-)~]),yields k31KllK21[COTSPIT1[021[RS-] v1 = (10) 2(1 + Kii[Ozl + KiiKzi[021[RS-l) If we assume K11[02]>> KllKz1[02][RS-]and Kll[021 >> 1,then the rate expression of eq 10 can be reduced to the following form: v1 = l / z ~ ~ 3 1 ~ 2 1 ~ ~ ~ ~ ~ ~ 1 (11) , 1 ~ which shows that the reaction rate should be first order
+
278
Environ. Sci. Technol., Vol. 22, No. 3, 1988
1 + -K1' + - K{Ki aH+
(1 +
aH+'
aH+ + Kal'
")
UH+
(13) where It3i and K2i are the rate constants for the rate-limiting electron-transfer steps and the equilibrium constants of substrate complexation by the ith catalytic center, respectively, Ka{ and Ka2'are the apparent acid dissociation constants of HOCzH4SH and HOC2H4S-,respectively, whereas K{ and Kz' are the apparent acid dissociation constants of reactions 14 and 15 involving the pyrrole nitrogens, where C = Co'ITSP and R = HOC2H4. K'
(H20-C-RS--C-H20)-6 (Hz0-C-RS--C-Hz0)-6 (H20-C-RS--C-HzO)4
+ H+ (14)
Kz'
e (HzO-C-RS--C-HzO)-'
+ H+ (15)
'Itobsd in this case is given by k31K21
hobad
=
~ ~ - 1 ~
+
k32K2ZKl' aH+
K1' Kl'K2' 1+-+aH+
a"+'
+ k33K23K;Kz' UH+2 (16) Kal'
UH+
Table 11. Comparison of Valuesa PH
( C O T S P ] ~: 6 . 5 x I O - ' M IONIC STRENGTH: 0.4 M
10.8
9.9
11.7
12.6
13.5
PH
Flgure 9. Calculated and observed profiles for kobsd vs pH for the catalytic autoxidation of 2-mercaptoethanol.
Values of Ka2/,k31K21, and k3&3 can be obtained by analyzing the kinetic data at pH 510 and at pH 212. At pH 510, assuming that 1 >> Ki/aHt >> K1'K2//aH+2 and (1 aH+/Kai' Ka2//aH+)N (1 + aH+/K,[), eq 16 reduces to
+
+
k31K21
kobsd
= 1 + aH+/Kal'
and k31K21 is calculated to be 1.5 X lo4 M-l. At pH 212, assuming K ~ ' K ~ / / u H>>+Ki/aH+ ~ >> 1, eq 16 becomes
and pKa2/and k33K23 are calculated to be 13.5 f 0.3 and 7.5 X lo3 M-' at p = 0.4 M and T = 25.5 "C. The constants k32K22 and K i were obtained by the nonlinear least-squares analysis of the following function (25): n
where
wts(i) = a(i)-l
(20) and ~ ( i=)sample standard deviation of data point i. The constants k31K21 and k33K23 were further refined by the same function, while the values of kcdcd were determined according to eq 13. Values obtained by this method were k31K21 (1.5 f 0.05) X lo4 M-' s - ~k32K22 , = (1.3 f 0.1) X lo4M-l s-l, k3& = (7.5 f 0.2) X lo3 M-l s-l, K1' = (1.7 f 0.8) X M, and K2/ = 1.1 X M at 25.5 OC. Values of the calculated pseudo-first-order rate constants (kdcd) are compared to the experimental values in Table I1 and in Figure 9.
Discussion The mechanism presented above is similar in some respects to the mechanisms proposed previously by Dolansky et al. (5) and Boyce et al. (3). However in our case, the binding of molecular oxygen precedes complexation of the substrate. A t the relatively high pH of the system, the deprotonation of the pyrrole N s and the hydrolysis of an apical water to give a hydroxide trans to the site open for oxygen binding favors O2addition. This pathway is consistent with the observed zero-order dependence on dissolved oxygen, which also suggests that the majority of the catalyst is present as an oxygenated adduct &e., [CoTSPIT = oxygenated adduct). Another major difference between the mechanism proposed here and those of Dolansky et al. and Boyce et al. is that the catalytic center appears to
9.17 9.38 9.60 9.79 10.19 10.36 10.49 10.74 10.98 12.14 12.48 12.80 13.04 13.24
kob#d
kobsd X
with
lo3, M-'
5.0 f 0.5 7.3 f 0.2 9:1 f 0.1 10.1 f 0.1 12.1 f 0.05 12.7 f 0.2 12.4 f 0.3 11.8 f 0.1 11.4 f 0.4 7.2 f 0.5 6.5 f 0.2 5.5 f 0.2 5.5 f 0.1 5.0 & 0.05
ko&d
5-l
at Various PH kcdCdX
io3, M-'
5-l
5.29 7.04 8.90 10.35 12.38 12.71 12.73 12.23 11.22 7.38 6.75 5.96 5.17 4.37
"Experimental conditions: p = 0.4 M, temperature = 25.5 "C, [O&ss = saturated with pure o2gas at 1 atm, [COTSP]~ = 6.4 X lo-'' M, and buffer system = NaCl, NaHCO?, and NaOH.
be a dimer rather than a monomer. The general spectral features of the oxygenated CoTSP complex have been reported previously by Gruen and Balgrove (24). As RS-was added to the catalyst solution, significant changes in the CoTSP visible spectrum were observed; these were consistent with the formation of a dimer as the reactive species. This spectral change was also noted by Beelen et al. (26) in their study of the autoxidation of 2-mercaptoethanol catalyzed by CoTSP. According to the spectral changes shown in Figure 7, most of the catalyst dimerized within the first minute of reaction. This implies that [COTSPITzz 2[dimer] and that a half-order dependence instead of a first-order dependence on [CoTSPITcatalyst should be observed if a monomeric species is the principal reactive species. According to the mechanism of Figure 8, several possible reactions can be considered as the rate-determining step. They are the binding of molecular oxygen, the complexation of the substrate by the oxygenated dimer, and the subsequent electron-transfer steps. However, it appears that one of the two sequential electron-transfer steps is most likely to be the rate-limiting step since a zero-order dependence on both the substrate at low ionic strength and dissolved oxygen at all ionic strengths was observed. Furthermore, both oxygen binding and substrate complexation should be fairly rapid ligand substitutions. The rate of ligand substitution for first-row transition metals is frequently independent of the nature of the ligand. The observed rates for each metal ion are similar to their corresponding water-exchangerates; therefore, the rate of ligand substitution for Co2+should be close to its waterexchange rate. The water-exchangerate for Co2+is * lo6 M-l s-l (27). Thus, the characteristic rate constant for substrate complexation and dioxygen binding of the cobalt(I1) centers should also be on the order of lo5. This value is 1 order of magnitude larger than our kobsdvalues. The observed independence of the rate of autoxidation on [02]requires both the rate of oxygenation to be fast and the equilibrium constant of the oxygenation to be quite large (Le., KIi 2 lo3). Jones et al. (IO)have argued that the binding of a fifth ligand in an apical coordination Co(I1) porphyrin is a prerequisite for oxygenation. Therefore, the binding of the RS- may result in an increase in both the rate and equilibrium constant of oxygenation. The bonding of oxygen to cobalt complexes can be explained in terms of u donation of electrons from the sp2 lorre pair on dioxygen to d,z orbital on the cobalt, coupled with the ?r-back-bondingfrom the cobalt d,, or dyzorbitals Envlron. Sci. Technol., Vol. 22, No. 3, 1988 279
Table 111. Correlation of Oxygenation Equilibrium Constant of Monobridged Dioxygen Cobalt Complexes with Basicities of Ligands"
ligands TEP'
CPKb 1% KO, 35.8 15.8 EPYDEN~ 30.6 14.7 4-IMDIEN' 29.1 12.6 PYDIENf 21.6 11.4 TRPY (PHEN)g =13 6.3 TRPY (BPY)h =12 5.4 "Data obtained from ref 31. bxpK = sum of the basicities of ligands bound to Co metal center. 'TEP = tetraethylenepentamine. EPYDEN = 2,6-bis(1,4-diaza-5-hexyl)pyridine.e 4-IMf PYDIEN = 1,9DIEN = 1,9-di-4-imidazolyl-2,5,8-triazanonane. di-2-pyridyl-2,5,8-triazanonane. TRPY (PHEN) = 2',5'-di-2pyridylpyridine (1,l-phenanthroline). TRPY (BPY) = 2',5'-di-2pyridylpyridine (2,2'-bipyridyl). f
into the a* dioxygen orbitals. The ligands coordinated trans to O2could compete for the same a-electron density on the cobalt; therefore, the strength of the Cc-02 bonding is very sensitive on the a-accepting or a-donating ability of the trans axial ligands. Good a-electron acceptors will decrease the electron density on the metal, resulting in a weaker Co-02 bond, whereas, a good a-electron donor will increase the electron density on the Co center available for a-back-bonding on O2and thereby stabilize the Co-02 bond. Several investigators have reported that the greater electron density on the metal results in a stronger metal-oxygen bond (28,29). Sulfur ligands (30)are generally good a donors and therefore they should enhance the rate of oxygenation and strengthen the Co-oxygen bond. However, little thermodynamic data on the stability of mixed complexes of Co(I1)-O2 and mercaptans are available. This may be due to the relatively rapid oxidation of mercaptans to disulfides. The rate constants for dioxygen binding by a variety bridged Co-chelate complexes have been reported to be large (IO). In addition, Martell (31) has shown a linear correlation of the sum of the basicities of the ligands bound to the cobalt metal center with the stabilities of the corresponding monobridged dioxygen Co-complexes (Table 111). All of these complexes show very large stability constants for oxygen binding that range from IO6 to Since the sum of the basicities (CpK) of the ligands on CoTSP-RS-) > 18, it is plausible that the oxygenation equilibrium constant is at least greater than lo3 and possibly 110'. The catalytically active dimer probably consists of two CoTSP monomers linked by a single RS- ligand at the metal center. However, two different molecules can be considered as possible monomer bridges. The first possibility involves a dioxygen bridge to yield p-peroxo-binuclear CoTSP center, while the other would be the binuclear complex bridged by the thioethanol. Assuming that the former dimer is the reactive species, the reaction cycle would be Co"TSP-02
+ RS- 2RS--Co"TSP-02
&
2(RS--Co"TSP-02) RS--Co"TSP-02-Co1'TSP-RS-
+0 2
(22)
where K3 and K4 are the equilibrium constants of reaction Envlron. Scl. Technol., Vol. 22, No. 3, 1988
-
+
is extremely slow. Since the bonding of RS- trans to dioxygen should enhance the rate of oxygenation, the rate of the formation of the corresponding binuclear oxygenated adduct would be even slower. Formation of a psuperoxo complex is too slow to be consistent with our kinetic observations. It may be argued that the steric interaction between the two macrocyclic rings in our proposed RS-bridged diiner would be too severe to allow dimerization to occur, but it has recently been shown that two CoTSP molecules can be bridged by the smaller N atom of an amino group (33). Thioethanol may bridge the two Co2+centers through either sulfur or oxygen. However, sulfur should preferably serve as the bridging site due to its better electron-donating ability and higher polarizability than oxygen. A rate-limiting electron-transfer step has been proposed previously for the catalytic autoxidation of both SO2- and H2Sby CoTSP. Hoffrnann and Hong (34)have reported electron spin resonance (ESR) evidence for the formation of superoxide-like ternary CoTSP complex during the catalytic autoxidation of sulfite. Results from other ESR studies (IO) have also shown that superoxide-like Co complexes are formed during the course of catalytic autoxidation. Additional information obtained from magnetic susceptibility data and EPR spectra shows that for lowspin d7 complexes, Co(II)L4(B),where L4 and B are the equatorial and axial ligands, respectively, there is only one unpaired electron, and more than 80% of the electron density resides in the a* dioxygen orbital (IO). Therefore, we believe that RS--CO"TSP-RS--CO~~'TSP-O~*is the principal reactive species in the catalytic cycle. The effect of pH on the rate of autoxidation, as shown in Figure 8, suggests that the pyrrole nitrogens are protonated in ConTSP. Two possible acid-base equilibria can participate in the catalytic cycle to give the observed pH effect. They are the deprotonation of hydrogen ion either at the pyrrole nitrogens (eq 14 and 15) or at the bound water molecules (eq 25), where C = Co"TSP. K
H20-C-RS--C-OH2 HZO-C-RS--C-OH-
(21)
According to the above mechanism, the following rate law will be obtained:
280
21 and 22, respectively. According to eq 23, the rate should be second order in [Co"TSPIT and reciprocal first order in [O,]. If 1>> K3[RS-] or K,[RS-] >> 1,either a secondorder dependence or a zero-order dependence on [RS-] should be observed. However, these predicted kinetic orders for the substrate, catalyst, and dioxygen do not agree with the experimental observations. Wagnerova et al. (32)have studied the oxygenation of C O ~ T Sin P alkaline solution (pH >12) and observed that the peak at 670 nm was the only peak in the spectrum after 5 days of oxygenation and that the peak normally attributed to the dimer at 630 nm was not observed over a 5-day period. This implies that the reaction H2O-COTSP-02 + H20-COTSP-02 H~O-COTSP-O~-COTSP-OH~ 0 2 (24)
+ H+ (25)
We have found by potentiometric titration that the pK, of the bound water molecules on the dimer H20-Co"TSP-OH2 is >>14.7. Since sulfur is a good electron donor, the incorporation of sulfur into the dimer would further enhance the electron density on the cobalt, and consequently, the cobalt would withdraw less electron density from the bound water molecules, As a result, the hydrogen ion would be more difficult to deprotonate from the bound water molecules, and thereby the pK5 should be 114.7. Thus, the effect of eq 25 is ignored for the autoxidation. Berezin (35) has noted that protonation of metal-
phthalocyanine complexes occurs preferentially on the ring nitrogens. Since phthalocyanine is a highly conjugated macromolecule, deprotonation of a single nitrogen would change the electron density on the ring and thereby could alter the electron density around the Co center and result in the change of the rate of electron transfer. If the CoI'TSP is deprotonated, kobsdwould be as given by k31K21
kobsd
1
+ aH+/Ka{ + ka2//aH+
(26)
According to eq 26, the observed pH dependence is determined by the deprotonation of HOC2H4SHand HOC2H4S-;however, pK,; for the 2-mercaptoethanol, which was estimated to be >13 by titration, is too high to yield the observed dependence. The estimated value of pK,; is reasonable since the pK,/ of the OH group of ethanol is N 15 (36). Thus, deprotonation of a pyrrole nitrogen is likely. The values of pK{ and pKi were determined to be 10.8 and 11.0, respectively. These values agree well with the pK,'s for deprotonatioh of primary, secondary, and tertiary amines (37). The observed production of ascorbate radicals during the course of the reaction provides strong evidence for the formation of RS' as an intermediate. Ascorbate radicals may have been produced by one of two alternative pathways. The first pathway (21,22) would be k , = 6.8 X lo8 M-I RS- A'H+ RS' AH(27)
+
+
RS* + RS-.
X
lo0 M-I
In summary, the rapid oxidation of 2-mercaptoethanol and other reduced sulfur compounds in the presence of Co'ITSP suggests that these and related catalysts (3, 4) may be used in pollution control applications such as SO2 stack-gas scrubbing, post-Klaus plant scrubbing of H2Sand SOz, sweetening of sour oi1 refining waste waters that contain high concentrations of RSH, elimination of excess rocket fuel wastes (N2H4), odor control in waste wrrter treatment facilities, and odor and corrosion control in sanitary sewers. With a growing likelihood of emission control requirements for reduced sulfur compounds from power plants and oil refineries, alternatives to limestone slurry scrubbers will be needed.
Acknowledgments We gratefully acknowledge the assistance of the Pollution Control Processes/Environhental Engineering Section (US.EPA) and Donald Carey. We are also grateful to Detlef Bahnemann and Eric Betterton for their interest in this work and for their extremely valuable assistance. Glossary hbed kcalcd k31
+
while the second pathway (38-40), and the more likely pathway, would be k , = 2.6
lo6 s-l) to give disulfide as the product.
(28)
-
(HOC2H4S--Co11TSP-HOCzH4S--Co*1TSPK2l
8-I
'RSSRO-
observed rate constant calculated rate constant from eq 3.27 electron-transfer rate constant (s-') of (HOCzH4S--CO"TSP-HOC~H~S--CO~~TSP-O )-6
k32
O:-P eq&brium constant (M-') of (H20-Co"TSP-HOC~H~S--CO''TSP-O~)-~ HOC2H4S- * (HOCZH ~ S - - C O ~ T S P - H O C ~ H ~ S - - C ~+ ~T H2O SP-~~)~ electron-transfer rate constant (d)of (HOC2H4S--CO"TSP-HOC~H~S--CO~~TSP-O )-7 ---*
+
(HOCzH4S~-Co11TSP-HOCzH4S--Co*1TSP03-7
Kzz In the former case, the ascorbate radical is formed directly, while in the latter case, it is formed indirectly from RS'. Nonetheless, the thioethanol radical is undoubtedly an intermediate of the catalyzed autoxidation. Several pathways may yield RSSR from RS'. The first pathway is
+ RS'
RS'
-
RSSR
(31)
k33
.-+
+ 2H+
HzO2
+ 02
(32)
K23
K,'
K2'
The third pathway (38, 39, 41, 42) is
RSOO' RSOO'
+ RS'
--
+ RSOO'
(33)
+0 2 RSSR + 202
RSSR
(34)
Ka1 Kaz
(35)
Given the above rate constants appropriate to the three pathways and the experimental conditions used in this study, we can estimate that eq 28-30 control the fate of RS'. With a small steady-state concentration of RS', the rate of reaction 31 should be very slow, and it can be ignored as the primary route for the disappearance of RS'. Also, the subsequent reactions in the third sequence are expected to be relatively slow because of the low concentrations of the various radical intermediates. Thus, RS' is more likely to react via eq 28, 29, and 32 (kifsd = 1.3 X
(HOC~H~S--CO'*TSP-HOC~H~S--CO~~'TSP-
0 2 7 8
The second pathway is a sequence that includes eq 28 and 29 as the initial reactions followed by 02'- + 0 2 ' -
equ&brium constant (M-') of (H20-Co"TSP-HOC~H~S--CO"TSP-O~)~ + HOCzH4S- + (HOCZH~S--CO~TSP-HOC~H~S--CO~TSP-O~)-~ + H2O electron-transfer rate constant (s-') of (HOC2H4S--CO~TSP-CO~TSP-HOC~H,S--CO~TSP-O~)~
Kll K3 K4 K5
equilibrium constant (M-I) of (H20-ConTSP-HOC~H~S--CO~TSP-O~)-~ + HOC2H4S- + (HOC2H4S--ConTSP-H0C2H4S--ConTSP-02)-s + H2O apparent acid dissociation constant (M) of the acid-base equilibria (HZ0-Co"TSP-HOCzH4S--Co"TSP-H20)" + (H~O-COI'TSP-HOCZH~S--CO"TSP-H~O)~ + H+ apparent acid dissociation constant (M) of the acid-base equilibria (Hz0-Co"TSP-HOC2H4S--CO"TSP-H~O)~+ (H,O-CO"TSP-HOC~H~S--CO''TSP-H~O)-~ + H+ apparent acid dissociation constant (M) of the acid-base equilibria HOC2H4SH+ HOCzH4S+ H+ apparent acid dissociation constant (M) of the acid-base equilibria HOCzH4S-+ -OC2H4S-+ H+ equilibrium constant (M-l) of (HzO-ConTSP-HOCzH4S--Co"TSP-H 0)-5+ O2 == (HzO-Col*TSP-HOC2H4S--C~rTSP-02)-6 + H20 equilibrium constant (M-') of CohTSP-OZ + HOC2H4S- + HOC2H4S--Co1'TSP-0 equilibrium constant of 2(HOC2H S--Con2TSP-Oz) + HOC2H4S--ConTSP-02-CodTSP--SH4C20H +0 2 apparent acid dissociation constant (M) of the
acid-base equilibria (H20-Co"TSP-HOC2H4S--CO"TSP-H,O)-~ * (HZ0-C~"TSP-HOCzH~S--CO"TSP-OH)~ + H+ Environ. Sci. Technol., Vol. 22, No. 3, 1988
281
Envlron. Scl. Technol. 1988, 22, 282-286
hydrogen ion activity ionic strength (M) total Co"TSP concentration (M)
aH* P
Redpath, J. L.; Willson, R. L. Int. J. Radiat. Biol. Relat. Stud. Phys. Chem. Med. 1973,23, 51. Schuler, R. H. Radiat. Res. 1977,69, 417. Leung, P. K.; Hoffmann, M. R. J. Phys. Chem. 1985,89,
[Co"%SPIT [Co total ith ConTSP catalytic center concentration (M) TSPlTt Registry No. Co(II)TSP, 29012-54-2; HSC2H,0H, 60-24-2; (HOCzH$)2, 1892-29-1.
5267.
Gruen, L. C.; Blagrove,R. J. Aust. J. Chem. 1973,26,319. Bevington, P. R. Data Reduction and Error Analysis for the Physical Sciences; McGraw-Hill: New York, 1969; p 225.
Beelen, T. P. M.; da Costa Gomez, C. 0.;Kuijer, M. Recl.
Literature Cited Sheldon, R. A.; Kochi, J. K. Metal-Catalyzed Oxidation of Organic Compounds; Academic: New York, 1981. Hoffmann, M. R. Environ. Sci. Technol. 1980, 14, 1061. Boyce, S. D.; Hoffmann, M. R.; Hong, P. A.; Moberly, L. M. Environ. Sci. Technol. 1983, 17, 602. Hoffmann, M. R.; Lim, B. C. Environ. Sci. Technol. 1979,
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Received for review April 20, 1987. Accepted August 20, 1987. Support for this research was provided by grants from the US. Environmental Protection Agency (R809198-01 and R811612-
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Effect of Cosolvents on the Solubility of Hydrocarbons in Water Frank R. Groves, Jr. Chemical Engineering Department, Louisiana State University, Baton Rouge, Louisiana 70803
rn New experimental data are reported on the aqueous solubility of n-hexane in the presence of cosolvents methanol and methyl tert-butyl ether (MTBE) and of benzene with MTBE. These data along with literature information on the benzene-ethanol, benzene-methanol) and hexane-ethanol systems are correlated with the aid of Margules and UNIQUAC activity coefficient equations. A rapid approximate method for predicting hydrocarbon solubility in the presence of a cosolvent is reported. Introduction The transport of spilled organic materials through the environment is a major concern for environmental protection. A current problem is the leakage of gasoline from 282
Environ. Scl. Technol., Vol. 22, No. 3, 1988
underground storage tanks. Leaked organics come into contact with groundwater, gradually dissolve, and are transported through the environment. Modeling of these transport processes requires a knowledge of basic thermodynamic data, including aqueous solubility. Solubility of most pure hydrocarbons is known. However, unleaded gasoline may contain organic octane enhancers, some of which are completely miscible with water. The effect of these cosolvents is to increase the hydrocarbon solubility. The objective of this research is to study the effects of cosolvents on hydrocarbon solubility with special emphasis on gasoline-range hydrocarbons. Specifically, we have investigated the aqueous solubility of benzene and nhexane in the presence of methanol, ethanol, and methyl
00 13-936X/88/0922-0282$0 1.50/0
0 1988 American Chemical Society
