Kinetics and Mechanism of the Oxidation of Sulfide by Oxygen: Catalysis by Homogeneous Metal-Phthalocyanine Complexes Michael R. Hoffmann” and Brian C. Lim Environmental Engineering Program, Department of Civil & Mineral Engineering, University of Minnesota, Minneapolis, Minn. 55455
The kinetics of the catalyzed autoxidation of hydrogen sulfide to sulfur, thiosulfate, and sulfate in aqueous solution has been investigated potentiometrically over the p H range 5-12. The catalytic effects of cobalt(II), nickel(II), and copper(I1) 4,4’,4”,4”’-tetrasulfophthalocyaninehave been determined, and differences in catalytic ability have been explained in terms of relative capacity of the octahedral complexes to reversibly bind oxygen. Kinetic data have been analyzed in terms of a bisubstrate Michaelis-Menton rate law. T h e rate law and other data indicate t h a t the catalyzed reaction proceeds via the formation of a tertiary-activated complex in which HS- and 0 2 reversibly bind with CoTSP2-. p H variations have been explained in terms of the acid dissociation of H2S, general base catalysis, and deprotonation of pyrrole nitrogens within the phthalocyanine ring. Numerous attempts have been made in recent years to determine the kinetics and possible mechanisms for the autoxidation of hydrogen sulfide (I+), but unfortunately, results of these investigations show considerable disagreement. T h e rate of oxidation of sulfide in both fresh and seawater systems (7,9,2(1)has been shown to be sensitive to trace metal catalysis by Mn(II), Cu(II), Fe(II), Ni(II), and Co(I1). From these results, it is clear that trace metal catalysis is a dominant factor controlling the rate of autoxidation in natural systems. Sulfide oxidation is an important link in the natural sulfur cycle (22)
From a commercial standpoint, the oxidation of hydrogen sulfide and other reduced sulfur compounds in the presence of homogeneous (12, 13) or heterogeneous (14) organo-metallic complexes may provide convenient and economical methods for reduced-sulfur pollution control. The principal product of the catalyzed autoxidation of hydrogen sulfide and mercaptans in sour refinery distillates was reported to be colloidal sulfur with concomitant formation of polysulfides, polythionates, and sulfate (15). Little is known about the oxidation kinetics or the mechanistic pathways by which catalysis takes place. Phthalocyanines are macrocyclic tetrapyrrole compounds, which readily form square planar complexes in which the metal atom is located in the plane of the ring and is bonded by four pyrrole nitrogens as depicted in Figure 1. Phthalocyanine complexes have been shown to be effective catalysts for the autoxidation of hydroxylamine (16) and hydrazine (17), the heterogeneous dehydrogenation of cyclohexadiene (18), and the autoxidation of benzaldehyde (19). Catalytic properties of cobalt(I1) 4,4’,4”,4”’-tetrasulfophthalocyanine (CoTSP2-) have been compared to those of a n oxidase (26, 19) Because of their relationship to naturally occurring porphyrin pigments (ZO), coupled with their high stability, specificity, and strong catalytic activity, metal-phthalocyanine complexes are suitable models for studying catalytic effects of trace metals in aquatic systems. A review by Almgren and Hagstrom ( 1 )of the literature on the autoxidation of sulfide in natural waters reveals the extent of disagreement and lack of reproducibility in previous studies. Observed reaction orders and half-lives were inconsistent. Half-lives varied from 24 to 10 000 min for similar concentration ranges. Some investigators ( 2 , 6 )have reported 1406
Environmental Science & Technology
complex kinetics, whereas others ( 2 , 3, 5, 7 ) have reported simple kinetic observations. Numerous mechanisms have been proposed to account for observed rates and product distributions. Abel(21) originally proposed that the rate-controlling step involved the formation of HS02- as an intermediate. Subsequently, a radical mechanism (2) and a mechanism involving a series of parallel reactions ( 6 ) ,in which the major reaction products are formed, have been proposed. In order to further the understanding of the catalyzed autoxidation of sulfide, the stoichiometry, kinetics, and mechanisms of the oxidation of H2S by 0 2 have been investigated and reported here. Results of this study should help to elucidate the role of trace metal catalysis as an important pathway for sulfide removal by autoxidation in aquatic systems and suggest an alternative method for reduced sulfur pollution control.
Experimental
Synthesis of Catalysts. The monosodium salt of 4-sulfophthalic acid was obtained by neutralization of a 50% aqueous solution of 4-sulfophthalic acid (Eastman Kodak, practical grade) with 15 M NaOH. T h e resulting brown solution was evaporated until precipitation of the yellow sodium salt. This precipitate was recrystallized with absolute ethanol. The tetrasodium salts of cobalt(II), copper(II), and nickel(II), 4,4’,4”,4~~’-tetrasulfophthalocyanine 2-hydrate were synthesized using procedures outlined by Weber and Busch (22). Sodium .l-sulfophthalate, ammonium chloride. urea, ammonium molybdate, and the appropriate metal salt were ground together and refluxed in nitrobenzene. The dark blue products were washed in methanol, 1 N HC1 saturated with NaCl, 0.1 N NaOH, and 80%ethanol. The final products were purified by Soxhlet extraction with absolute ethanol. All chemicals used in synthetic procedures were reagent grade. Reagents. Sulfide solutions were prepared from reagent grade Na2S.SH20 (Mallinckrodt) crystals which were washed with distilled water to remove any oxidized surface layers. Deoxygenated doubly distilled water was used to prepare all solutions. Sulfide solutions were analyzed by potentiometric titration with PbC104 (Orion). Buffers were prepared using reagent grade tris(hydroxymethy1)aminomethane (Sigma), sodium phosphate-monobasic (Mallinckrodt), sodium phosphate-dibasic (Mallinckrodt), sodium borate (BA), potassium chloride (BA), sodium hydroxide (BA), and hydrochloric acid (BA). Sodium perchlorate (Fischer) was used to maintain ionic strength constant a t p = 0.4 M. Kinetic Data. Kinetic data were obtained by two continuous measurement procedures. Sulfide activity was monitored as a function of time using a n Orion (Model 94-16) specific sulfide ion electrode and an Orion (Model 801-A) digital ionalyzer linked to a n Orion (Model 751-02) digital printer. Dissolved oxygen concentrations were followed with a YSI (Model 5739) DO probe, a YSI (Model 54A) oxygen meter coupled to a Hewlett-Packard (Model 410C) voltmeter. The response of the DO probe was calibrated against the standardized Winkler method in the absence of HPS and HS-. Potential interference from H2S was minimized because the sensing element of the YSI probe is protected by a n On-permeable plastic membrane that serves as a diffusion barrier
0013-936X/79/0913-1406$01.00/0
@ 1979 American Chemical Society
H,S,
HS-
catolyst 0 2
Ss , S,Of, S O ; , S,O,',
S4=, S i , SO4=
catalyst = M (E), 4, 4: 4', 4"'- tetrosulfcphtholocyonlne
Table 1. Effect of pH on kobsd PH
5.52 6.46 6.70 7.01
M=Fe,Co, N i , C u
0
s 0;
Figure 1, Generalized stoichiometric expression for the aqueous autoxidation of sulfide catalyzed by homogeneous metal phthalocyanine
catalysts
8.29 9.23 9.83 10.47
against HYS. Kinetic intermediates and final products were determined by standard analytical procedures previously described in detail by O'Brien and Birkner ( 6 ) and Chen and Morris ( 2 ) .Aliquots were taken during the course of a reaction for sulfide, polysulfide, sulfur, thiosulfate, sulfite, and sulfate analyses. Spectrophotometric measurements were made with a Beckrnan 26 UV-vis spectrophotometer. For kinetic determinations, reactions were run in a waterjacketed, glass and Teflon reactor with a total volume of 2.0 I,. The reactor profile and schematic diagram for the complete reactor system are depicted in Figure 2. Five electrodes, a precision thermometer, a sampling port, and a pressure release valve were adapted with precision-designed screw-type Teflon plugs, which were sealed with O-rings to provide an air-tight seal. T o minimize the effect of trace metal catalysis in noncatalyzed experiments, careful washing procedures were used. The reactor was cleaned successively with alcoholic KOH, 10% "03, and doubly distilled water. EDTA salts were added to noncatalyzed kinetic runs to sequester any residual trace metals and to minimize trace metal catalysis. In a typical kinetic run, a 2.0-L stock buffer solution was transferred to the reactor cell and high-purity oxygen (Matheson), nitrogen (Airco),or a controlled 02-Nz mixture was purged through the cell at 20 mL/min with constant stirring for 2 h. After saturation, the cell was sealed from the atmosphere and any free head space was displaced with a movable Teflon piston. Constant temperature was maintained a t 25.0 f 0.2 "C with a Haake FK-2 circulator system and temperature controller. When the p H was 38.0, some reactions were run with a stream of oxygen purged constantly through the reactor. For these situations, the concentration of oxygen was maintained a t a constant value. T o distinguish between the will be used to refer to a fixed initial 0 2 two conditions, [0]0 concentration and [O,], will be used to refer to a constant 0 2 concentration. T o initiate the reaction, known volumes of a stock sulfide solution and a concentrated catalyst solution, which were calculated to give [HzS]o and [CATIo, were injected into the reactor by syringe with rapid stirring. Complete mixing and dynamic equilibrium were obtained within 2 min.
Results The sensitivity, selectivity, and dynamic response characteristics of a silver sulfide membrane electrode make it ideally suited for continuously monitoring changes in sulfide activity ( 2 3 ) .Ostlund and Alexander ( 3 ) ,Demirjian ( 7 ) ,Almgren and Hagstrom ( I ) , and Hoffmann ( 2 4 ) have used a Ag/AgpS electrode successfully to follow sulfide oxidations in both natural and artificial systems. Other ion-selective membrane electrodes have been used for reaction rate measurements and
11.06 11.68 12.12
buffer systems
0.1 M NaH2P04 0.1 M NaOH 0.1 M NaH2P04 0.1 M NaOH 0.1 M NaH2P04 0.1 M Na2HP04 0.1 M NaH2P04 0.1 M NaOH 0.1 M Tris 0.1 M HCI 0.1 M borax 0.1 M HCI 0.1 M borax 0.1 M NaOH 0.1 M NaHC03 0.1 M NaOH 0.1 M Na2HP04 0.1 M NaOH 0.1 M Na2HP04 0.1 M NaOH 0.1 M KCI 0.1 M NaOH
kobsd, min-'
2.71 X lo-' 5.82 X lo-' 3.58 X lo-' 6.74 X lo-' 1.12 x lo-' 1.42 X lo-' 5.51 X lo-' 2.86 X lo-' 1.12 1.32 no observable
reaction
M, [H2SIT = 1.0 X IOm3 M, [Oplc = 1.0 a Where [CoTSPIo = 2.5 X X M,I./ = 0.4 M,and T = 25.0 O C for pH > 8.0 and [COTSP]~= 2.5 X lo-' M for pH < 8.0. M, and [O2lO= 1.0 X M, [ H P S ] = ~ 1.0 X
mechanistic investigations (25).For a pseudo-first-order reaction in sulfide, the observed rate constant can be obtained from the slope of a linear EMF vs. time relationship according to Equation 1 ( 2 4 ) a t constant ionic strength: -2F d(EMF) (1) RT dt where 2FIRT is the reciprocal of the Nernst coefficient. For example, the slope of an EMF vs. time plot a t pH 8.8 was 0.14, the y intercept was -674.9 mV, and x2 = 0.17, for which [HS-]o = M, [ 0 2 ] 0 = M, [CoTSP]o = 5 X M, p = 0.4 M, and T = 25.0 O C . Pseudo-first-order relationships of this type were linear after an initial mixing period for 90% of the reaction. Linearity of E M F vs. time relationships over a broad range of conditions suggests that the autoxidation of sulfide is first order in total sulfide concentrations when [ 0 2 ] 0 >> [HS-Io. The amperometric response of the dissolved oxygen electrode is a linear function of the aqueous oxygen concentration ( 2 6 ) . Amperometric response (i, 0: [Oz]) against time was linear for reactions in which [02]0 >> [HS-]o. For example, a t M, [CoTSPIo = p H 8.3, where [HS-Io = M, [ 0 2 ] 0 = 2.5 X M, the slope of a linear [Oz]vs. time relationship was -0.19, t h e y intercept was 1.1 X 10-3 M, and x2 = 0.17. This result indicates that the apparent reaction order in oxygen was zero. However, a nonintegral dependence of the reaction rate on [O,] was determined by the van't Hoff method (35).The apparent reaction order in oxygen was found to be 0.65. Similarly, the apparent reaction order for [CoTSPIo was determined to be 1.0. A summary of kinetic data presented in Table I shows a complex p H dependency over the p H range 5-12. In the p H range 5-8 with increasing p H , [HzS] decreases while [HS-] increases due to the acid dissociation of H2S (pK,1 = 7.0). In the p H range 10-12, the CoTSP catalyst undergoes an acid dissociation of a proton bound to a coordinating pyrrole nitrogen (pK, 1: 11). Deprotonation tends to stabilize a monomeric form of the catalyst (27), which apparently enkobsd
= -___
Volume 13, Number 11, November 1979
1407
a.
Metal Strip
Teflon Sleeve\
Gummed Rubber
Gas Dispersion
REACTOR CELL
b.
CONFIGURATION OF THE REACTOR TOP (TEFLON)
C. Printer 7
Orion 801 - - - - - - - _ I lonalyzer
'2
I
e
SCHEMATIC DIAGRAM FOR MONITORING SULFIDE OXIDATIONS Figure 2. Schematic presentation of batch reactor profile (a), of the Teflon top (b), and of the overall monitoring system (c)
1408
Environmental Science & Technology
must be one less than the number of species in the reaction sequence. Given these conditions, the expressions for the concentrations of the various forms of the catalyst can be written as:
Table II. Comparison of Catalytic Activity of Cobalt( ll), Copper( Il),’and Nickel( II) Phthalocyanine a catalyst
PH
CuTSP NiTSP CoTSP CuTSP NiTSP CoTSP CuTSP NiTSP CoTSP
8.3 8.3 8.3 6.7 6.7 6.7 8.8 8.8 8.8
[CAT] x
lo8
11/2, min
500 500 0.05 10
76.0 6.0 7.5 35.0 5.0 2.0 58 4.0 13.0
10 10
250 250 2.5
turnover no.
1.5 17 130 000 14 100 25 000 3.5 50 1500
a F o r p H 8 . 3 a n d 8 . 8 , [HS-Io= 1.OX 10-3M, [ O 2 l c = 1.OX l O P M , T = 25.0 O C , and p = 0.4 M, and, for pH 6.7, [HS-],, = 1.0 X M and [02],, = 10X M. t112 = In 2 / k o b s d . Number of sulfide molecules transformed by a single molecule of catalyst in 1 min.
[CoTSP] 0~ h - l h : j [CoTSP-Or]
[HS-CoTSP-021
CoTSP2-
+ O2 3CoTSP-022-
3:
(6)
(7)
hlh2[02] [HS-]
(8)
The mass balance equation for the total catalyst concentration is given by Equation 9: [COTSPIT= [CoTSP’-]
+ [COTSP-O~”]
+ [HS-CoTSP-On’’-]
(9)
From steady-state considerat ions: [HS-CoTSP-Oi‘’-] =
ii lk 2 [C oTSP] [ 0 2 1 [H S-1
(h-lh:’
hances the reversible activation of molecular oxygen. Catalytic activity of CoTSP is significantly greater than either NiTSP or CuTSP. According to the data listed in Table 11, the catalytic activity follows the apparent order of CoTSP > NiTSP > CuTSP. This order remains t h e same over t h e entire p H range. The catalytic activity of metal phthalocyanines in aqueous solution was documented initially by Cook for the decomposition of H2O2 ( 1 9 )and the oxidation of H I (28) and later by Wagnerova and co-workers (16, 17) for the autoxidation of hydroxylamine and hydrazine. Both investigators suggested that the behavior of the phthalocyanine catalysts was similar to such enzymes as catalase and peroxidase. Wagnerova et al. ( 17) suggested t h a t reactions catalyzed by CoTSP follow Michaelis-Menton enzyme kinetics. Kinetic evidence presented above is also suggestive of enzyme-like kinetic behavior f’or the catalyzed autoxidation of sulfide. A well-known characteristic (29) of enzyme reactions is the variability of reaction orders for catalyst and substrate. Under certain conditions, the order in substrate concentration can vary between zero and one, but most likely a nonintegral order will he ohserved. In order to interpret t h e observed kinetic behavior more thoroughly, a bisubstrate model for catalytic activity was developed and a rate law was derived using the method of King and Altman (301, which is based on a standard determinant procedure used for solving a system of nonhomogeneous linear equations obtained from steady-state considerations. T h e mechanism postulated to account for ohserved kinetic behavior between p H 8.3 and 10.0 is the f’ollowingordered ternary-complex mechanism (31):
a
+ h-lh-2 + h:jh2[HS-] h-2h1[02] + h:jk1[02]
+ k - l h - 2 + h?h:j[HS-]
-
(10)
+ hl(h-2 + h:J[Oa] + hlh?[O~]I[HS-])
Using this expression and the rate-determining step represented by Equation 4, the final rate law can be written as: u = -d[HS-]/dt
=
k:3[CoTSP]~[02] [HS-]
(11)
A variant of this mechanism occurs when HS- reacts with CoTSP-OL2- without the formation of a ternary complex of iciently long life to be kinetically significant. This type of mechanism was originally suggested by Theorell and Chance (52) Although an exact rate equation can be derived directly, it can be obtained indirectly from Equation 11 if h 3 >> h-2. After simplification, Equation 11 will yield Equation 12: huff
Spectrophotometric evidence for the formation of monomeric adducts of CoTSP2- with HS- and 0 2 has been presented by Wagnerova e t al. (331,and ESR spectroscopy has been used by Abel et al. ( 3 4 )to demonstrate the reversible formation of 1:l molecular oxygen adduct with CoTSPZ-. When [O,] >> [HS-] and [HS-] is sufficiently low such that ( k l h i h - l h - J / h l h r >> [HS-1, Equation 11 can be reduced
+
to:
(2)
k-i
C ~ T S P - O ~ Z+- HS-
3HS-TSP-O~:’-
(3)
h -2
k .i
HS-CoTSP-02”- -+ CoTSPZ-
+ HS02-
+ 0 2 +HS04-
(4)
I n this case, the reaction is apparently zero order in [O,] and first order in [HS-1, which is consistent with kinetic observations. Similarly, when [HS-] >> [O,] and h:jlh1 >> [O,]
kq
(rapid) HS02-
(5)
I n this mechanism, there are two intermediates, CoTSP-022and HS-CoTSP-O$, and two steady-state equations for them. According to the method of King and Altman (30),t h e concentrations of CoTSP2-, CoTSP-O&, a n d HSCoTSP-O$ can be shown to be proportional t o t h e sums of terms that are obtained from reactions steps which individually or in sequence lead to t h e formation of the particular species. In general, the number of rate constants in each term
The reaction will be apparently zero order in [HS-] and first order in [O,]. In general, with the proposed mechanism, nonintegral orders should be observed for both substrates and a f irst-order dependence for the catalyst concentratitm. In order to verify the applicability of Equation 11,experimental data were analyzed with the use of Lineweaver-Burk plots ( 3 2 ) of l / v against ~ 1/[02]g, where v g and [ O ~are O the Volume 13, Number 11, November 1979
1409
DOUBLE RECIPROCAL PLOT
initial rate and the initial oxygen concentrations, for a wide range of [02]0. Equation 11 can be rewritten in the following f'orm: 4
which can be rearranged to give: 1/1jo = (1 + K ~ / [ 0 2 + ] KB/[HS-] + Kc/[O~][HS-I)/VO (16) where Vo = h COTSP SPIT, K c = ( h - ~ h st h - l k 2 ) / h l h 2 , K A = k J k 2 , and K B = (h-2 h3)/h2. A plot of l / u o against 1/[0210, at constant [HS-1, should be linear with a slope of
+
( K A+ Kc/[HS-])/Vo
(17)
3
Q X
s , -
2
and an intercept on the l / u o axis of: (1-t KB/[HS-])/VO
(18)
Shown in Figure 3 are the linear double-reciprocal plots for the catalyzed autoxidation of sulfide a t p H 6.7 and 8.3. The linear nature of these plots lends strong support to the propoqed ternary complex mechanism. Support for a Theorell-Chance type mechanism is obtained from examining the effect of ionic strength on the rate of autoxidation. If h 3 >> k-2, the rate controlling step is given by Equation 3 in which a negatively charged HS- forms a reactive intermediate conlplex with the monomeric CoTSP-02*- adduct that has a net negative charge. The primary salt effect predicts that for ionic reactions between ions of similar charge, an increase in ionic strength will result in an increased reaction rate ( 3 5 ) From the Debye-Huckel theory, it can be shown that: log k = log k o
+ 1.02ZaZbp'/2
(19)
in which Z,(HS-) = -1 and Zb(CoTSP-02*-) = - 2 . Even though the range of ionic strength is greater than the upper limit for strict applicability of the Debye-Huckel theory, a 2 obtained with positive slope for a plot of -log h&d vs. ~ 1 1 was m = 0.38, b = 1.18, and x2 = 0.009. In general, the reaction rate increases with increasing p. Both O'Brien and Birkner (6) and Alferova and Titova ( 3 6 )have reported positive effects of ionic strength. Based on his observation of a primary salt effect, Abel ( 2 2 ) postulated that the rate-controlling step of the noncatalyzed autoxidation was the reaction of two HS02intermediates formed from the combination of HS- and 02.
The p H dependence of the autoxidation of sulfide is unusually complicated as shown previously by Chen and Morris (2) and O'Brien and Birkner ( 6 ) . Increases in reaction rate between p H 4 and 8 have been attributed to the acid dissociation of H2S to give HS-. As the rate of autoxidation increases, the fraction of total sulfide present as HS- also increases. This behavior indicates that HS- is the principal reactive sulfide species in solution (22, 24) The p H dependency is further complicated by the fact that the catalyzed autoxidation is sensitive also to general base catalysis as indicated by the increase in reaction rate with increase in the total Tris buffer concentration. A plot of h&sd vs. [Tris] was linear with a slope of 0.60, a 3. intercept of 0.04, and a x2 of 0.04. Mechanistically, this relationship suggests that either hydrogen ions are transferred to a Bronsted base (A-) from the catalyst-substrate complex in the transition state, or that there is an equilibrium deprotonation of the substrate or catalyst followed by deprotonation of HA. In such cases, any basic molecule or ion can act as a receptor for the transfer of a proton from substrate or catalyst; therefore, a multiterm rate law should be observed ( 3 7 ) With the above pH-dependent considerations in mind, the mechanism postulated in Equa1410
Environmental Science & Technology
3
2
I I /
[OJO
X
IO' M-'
Figure 3. Lineweaver-Burk plot for the autoxidation of sulfide as a function of [O2lOat p = 0.4 M, T = 25.0 O C , where [HS-]= 1.0 X M and [CoTSPIo = 2.5 X lo-$ M
tions 2 through 5 can be modified by the addition of Equation 20 as a preequilibrium step and Equation 21 as the rate-determining step.
3H+ + HS-
H~S
hr
H S - C O T S P - O ~ ~t- A- +CoTSP3-
(20)
+ HSO2-
t HA (21)
CoTSP3-
+ HA 3CoTSP2- t A-
(22)
The slow rate-determining step given in Equation 21 involves a proton transfer to A- in the transition state. This may result in a more facile rearrangement of the catalyst-substrate complex. Using mathematical procedures described above, the following rate law can be derived from the proposed mechanism: uo = Vo[Ozl [H2Slo[A-l [Kd[HII([H+l+ Ka1)1+ KB[~~]{[H+I/([H ++K,i)/ ] -I-KA[HzS]O -k [Oz][H~SIoI (23) where [ H ~ S I=O [HzS] t [HS-1, K,1= [Hf][HS-]/[H2S],and A- = HPOd-, Tris. T o be complete, a second term in which [OH-] replaces [A-] should be included in the overall rate law, but in the pH range of 5 to 9, this term vanishes. Evidence for general base catalysis by buffers other than Tris and HP04was not established in this study. Mader ( 4 6 )and Hoffmann and Edwards ( 4 7 ) have reported a general catalytic influence of phosphate, TES, citrate, pivalate, and acetate buffers on the oxidation of sulfite, although borate buffers (46) were shown to be catalytically inactive. Chen and Morris (2, 38) have reported t h a t the autoxidation of sulfide was sensitive to catalysis by phosphate (2) and other weak conjugate bases such as resorcinol and pyrocatechol (38).A positive secondary salt effect should be expected for a reaction sensitive to general base catalysis ( 3 1 ) .
14
Ii
I C W
( 1 1 NtTSP
a U m
(111 C u T S P
SE
(1 1 1 )
CoTSP
a
[CI
I C x IO~'M
4
i
C
AT
01
350
1
400
I
450
1
500
1
550
WAVELENGTH
:4
I
600
min
A I
650
700
750
nm
Figure 4. (a) Visible absorption spectra for CoTSP'-, NiTSP'-, and CuTSP2- in water at T = 25.0 OC and pH 8.3. (b) Rapid scan spectra for reaction mixture at pH 12.0 after 1 and 4 min
adduct (670 nm) (33, 3 4 ) . Other investigators (41-43) have attributed the spectrum to a monomer-dimer equilibrium where the band a t 636 nm has been associated with the dimer. However, thse investigators failed to exclude oxygen from their systems. In the case of CuTSP and NiTSP, a monomer-dimer equilibrium is more likely. T h e electronic configurations of Cu and Ni do not favor the addition of axial ligands. Consequently, adducts of Co(I1) complexes are well known, whereas synthetic Cu(I1) and Ni(I1) oxygen carriers are uncommon. Spectral evidence for an increase in concentration of monomeric CoTSP-02 with a n increase in p H was obtained. As the p H increases from 9 to 12, the absorbance of the monomeric CoTSP-02 a t 670 nm increases while the peak a t 636 nm decreases. Evidence for this shift is presented in Figure 41). Rapid scan spectra of the HS7-02-CoTSP reaction mixture were taken 1 and 4 min after the injection of HS-. Upon complete mixing, the characteristic double hump spectrum of CoTSP is changed to a single peak, which has been attributed exclusively t o the CoTSP-O& adduct. Observed changes in the spectrum are consistent with additional coordination of HS- followed by enhanced coordination of 0 2 , which lead ultimately to a more rapid autoxidation. T h e dramatic discontinuity a t p H 12.3 is consistent with the formation of a dimeric adduct CoTSP-02-CoTSP reported by Wagnerova et al. (27) to occur only a t p H > 12. Maas et al. (14) have concluded t h a t in polymer-bonded CoTSP t h e reactive species for the autoxidation of thiols a t high p H was t h e monomeric adduct. Increased catalytic activity in the heterogeneous system over the aqueous system was attributed t o differences between the monomeric and dimeric adduct. Formation of CoTSP-02-CoTSP, which may be catalytically inactive, may account for the dramatic decrease in catalytic reactivity. In order to explain the apparent rate maximum near p H 12, the dissociation of the protons on the pyrrole nitrogens of the phthalocyanine ring and the concomitant enhancement of oxygen activation must be considered. The mechanistic sequence presented in Equations 2-5 can be modified to include the successive deprotonations of the pyrrole nitrogens. In the modified reaction mechanism, the principal reactive species is C(JTSP~I-.If the primary acid-base equilibria for the various forms of the catalyst, such as CoTSP2-, CoTSP4-, CoTSP-O2'-, C O T S P - O ~ ~ -HS-CoTSP-O2:j-, , 2nd HSC O T S P - O ~ ~are - , assumed to be important, the determinant techniques of Laidler and Bunting ( 3 1 ) and Laidler ( 4 4 ) can be used to obtain a general solution ( 4 5 ) for t h e p H dependency in the p H range of 9 to 12. A reduced form of this solution is presented in Equation 24: u=
T h e final region of p H dependency was observed between p H 9 and 12. Fallab (39) has shown t h a t the weakly acidic hydrogen atoms bound to the pyrrole nitrogen of the phthalocyanine ring dissociate in strongly basic solution. For uncomplexed phthalocyanine tetrasulfonic acid, P K N , = ~ 9.6 and P K N , >~' p K ~ , lAcceleration . of the autoxidation of sulfide may be due t o a n increase in catalytic efficiency of CoTSP0 2 : + . Deprotonation of a pyrrole nitrogen may lead to coordination of HS- or A- in t h e axial position opposite oxygen. Bonding of a strong u donor promotes nonsymmetric splitting of the t2# orbitals of Co(II), which in turn favors the formation of a bond with a TT acceptor ( 0 2 ) in t h e trans position ( 4 0 ) . Studies of the catalytic properties of heme have shown t h a t an axial ligand trans to 0 2 affects both the formation constant and its reversibility ( 4 0 ) . The visible spectrum for CoTSP is given in Figure 4a. The absorption bands have been attributed to an oxygen-free monomer (626 nm) and to a monomeric
where K1 is the equilibrium dissociation constant for the first protonated pyrrole nitrogen. According to Equation 24, v will be small when [H+] >> K1 and it will increase as [H+] K1. When p H > 12, the general solution shows that:
-
where K P is the equilibrium dissociation constant for the second protonated pyrrole nitrogen. Under these conditions, when [H+] > [HzSIO>> k,:
s Of ST.0
products
thiosulfate} HSO2-
= - '/Zd [ 0 2 1 l d t = d [ HS04-] l d t
under the conditions [ 0 2 1
mol % as Icatalystl, M
sulfate} HS02-
and from this stoichiometry:
-d [HS-] l d t
Table 111. Observed Production Distributions for Catalyzed and Noncatalyzed Reactions a
(35)
+ HzO
(381
Generally, no direct experimental evidence can be obtained that will elucidate the exact sequence of steps that will occur after the rate-determining one, although evidence for the existence of polysulfide intermediates, S d 2 - and Ss2-, is fairly convincing (24, 49). In Equation 33, HS- attacks HS02- t o give the postulated S202- intermediate, which reacts with an additional HS- ion to give S,2-, and water as a leaving group. Since the rate of this reaction is assumed to be rapid, water would be the preferred leaving group (51).In order to account for the appearance of colloidal sulfur, which is a complex
mixture of cyclohexasulfur, cyclooctasulfur, polycatenasulfur, and higher molecular weight sulfanes (52), a mechanism, in which the reaction proceeds through a series of nucleophilic displacements of polysulfides on t h e intermediate HS02until Ss,is produced in the final step by intramolecular displacement of S&. SScould form through an intramolecular displacement of ST^-; however, the yield of SSshould be low due to the instability of the Sg ring as compared to the s8 ring (55) Similar mechanisms have been proposed by Davies ( 5 4 ) to account for the acid-catalyzed conversion of S20z2- to Sa, and Hoffmann ( 2 4 ) to account for t h e formation of s8 in t h e oxidation of H2S by H202. Polysulfide intermediates are subsequently oxidized by oxygen to give the sulfur oxyanions (39).
Present results do not necessarily preclude a radical mechanism; but if radical pathways were significantly involved, then considerable levels of dithionate and tetrathionate should be formed. Since these products were not detected and sulfate was the primary product, contribution of a free-radical pathway in t h e overall reaction seems unlikely. Furthermore, strong evidence for the rapid catalytic oxidation of sulfite to sulfate has been presented by numerous investigators (55, 56) Differences in catalytic activity exhibited by Co(11)-, Ni(I1)-, and Cu(II)TSP2- can be explained in terms of three different mechanisms. Catalytic activity ( 6 1 ) for CoTSP2can arise from the reversible addition of oxygen to form Co(II1) with the superoxide anion (Ol-) as a ligand or the Co(I1)-singlet oxygen adduct or t h e one-electron oxidation of Co(I1) to Co(II1) to give a reactive radical species ( 3 ) ,or the two-electron oxidation of the 1 8 ~ - e l e c t r o nphthalocyanine ring to give a 16~r-electronsystem. Catalytic activity for NiTSP2- and C U T S - appears to be limited to the phthalocyanine redox equilibrium (57):
TSP4-
2
TSP4-
+ 2e- + 2H+
1 8 electrons ~
16x electrons
hydroquinone
quinone
(39)
Manassen and Bar-Ilan ( 1 8 ) have correlated the catalytic activity of metal-pht halocyanine complexes for the dehydrogenation of cyclohexadiene by nitrobenzene with the decrease in first oxidation potentials of metal-phthalocyanine complexes. Rollman and Iwamoto (57) report oxidation half-wave potentials for M(I1)TSP complexes in Me2SO and subsequent catalytic activity t o decrease in the following order: Co (+1.09 V) > Ni (+0.98 V) > Cu (+0.872 V). For the autoxidation of sulfide, catalytic activity of NiTSP2- and CuTSPL- appears to rise solely from t h e reversible redox equilibria of the metal complexes. T h e reversible addition of 0 2 to CoTSP2- is well documented ( 3 3 , 3 4 , 5 8 , 5 9 )and appears to correlate well with the enhanced catalytic activity of CoTSP2- compared to NiTSP2- or CuTSP2-. Conclusions
CoTSP2- appears to be an effective homogeneous catalyst for the oxidative control of hydrogen sulfide in aqueous systems. Addition of the cobalt catalyst a t parts per billion levels can result in an increase in the rate of autoxidation by a factor of lo4. CuTSP2- and NiTSP2- are less effective catalysts. Trace catalysis by metals, anions, and organic molecules can significantly affect the rate of autoxidation of sulfide. The wide range of results observed previously can be attributed to the variation introduced by these factors. The relative catalytic effects of CoTSP2-, NiTSPZ-, and CuTSP2- have been determined, and differences in catalytic ability have been attributed to t h e relative capacity of the octahedral metal complexes to reversibly bind oxygen. The kinetics of autoxidation catalyzed by CoTSP2- have been
characterized in terms of a bisubstrate Michaelis-Menton rate law. T h e rate law and other data indicate t h a t t h e catalyzed reaction proceeds via the formation of a tertiary activated complex in which HS- and 0 2 reversibly bind with CoTSP. T h e p H dependency has been explained in terms of acid dissociation of H2S, general base catalysis, and deprotonation of pyrrole nitrogens of the phthalocyanine ring a t high pH. Improved catalytic ability may be attained by attachment of CoTSP t o solid supports such as cross-linked polyacrylamide ( 1 4 ) or activated carbon ( 6 0 ) . Supported CoTSP or FeTSP catalysts may have widespread applicability for pollution control processes in which oxidation by oxygen is an important step. For example, posttreatment of secondary effluents through polymer-supported CoTSP may enhance the oxidative degradation of refractory organics. Autoxidation of Klaus plant effluents in the presence of polymer-supported CoTSP or FeTSP may provide a convenient and economical method for residual H2S and SO2 control. These potential applications should be explored more thoroughly.
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(1976). I < c w i i v d /or rei’ieic,June 23, 1978. Accepted August 17, 1979. Funds /or /hi\ rcwnrch project were protided by the Graduate School and t h o .Sommc~rf’c~/d Endou,ment of the Department o/ Civil and Mineral /
