One-Electron Oxidation of Hydrogen Sulfide by a Stable Oxidant

Article pubs.acs.org/IC

One-Electron Oxidation of Hydrogen Sulfide by a Stable Oxidant: Hexachloroiridate(IV) Ying Hu†,‡ and David M. Stanbury*,‡ †

College of Chemical Engineering, China University of Mining and Technology, Xuzhou 221116, People’s Republic of China Department of Chemistry and Biochemistry, Auburn University, Auburn, Alabama 36849, United States

‡

S Supporting Information *

ABSTRACT: Detailed reports on the oxidation of aqueous H2S by mild one-electron oxidants are lacking, presumably because of the susceptibility of these reactions to trace metal-ion catalysis and the formation of turbid sulfur sols. Here we report on the reaction of [IrCl6]2− with H2S in acetate buffers. Dipicolinic acid (dipic) is shown to be effective in suppressing metal-ion catalysis. In the presence of dipic the reaction produces [IrCl6]3− and polysulfides; turbidity develops primarily after the IrIV oxidant is consumed. Water-soluble phosphines are shown to prevent the development of turbidity; in the case of tris-hydroxymethylphosphine (THMP) the product is the corresponding sulfide, THMPS. THMP diminishes the rates of reduction of IrIV, and the rate law with sufficient THMP is first order in [IrIV] and first order in [HS−]. The rate-limiting step is inferred to be electron transfer from HS− to IrIV with ket = 2.9 × 104 M−1 s−1 at 25.0 °C and μ = 0.1 M. The kinetic inhibition by THMP is attributed to its interception of a polysulfide chain elongation process.

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idase.25,26 Rate constants for oxidation of H2S by one-electron oxidants are limited to highly reactive species such as MPO compounds I and II26 and radicals generated by pulse radiolysis, including SO4•−, I2•−, NO2•, OH•, and CO3•−.21,27 Efforts to study less reactive oxidants such as [IrCl6]2− and various Ru(III) ammines only revealed that the reactions are catalyzed by trace levels of copper ions.28 Studies with [NaP5W30O110]14− and [Nb3P2W15O62]9− did not yield to a full mechanistic analysis.15,29 Here we report the first kinetic study of oxidation of H2S by a mild one-electron oxidant in which the stoichiometry and kinetics of the noncatalyzed reaction are determined and a suitable mechanism is proposed. Crucial to this advance has been the development of methods that fully inhibit trace metalion catalysis and prevent the formation of sulfur sols. Our selection of [IrCl6]2− as the oxidant was based on its high stability and simplicity and the anticipation that Ir−S bonding would not be a complicating factor. We find that the ratelimiting step is the one-electron oxidation of HS− to HS•; polysulfides generated during the reaction are easily oxidized, which increases the overall rate of consumption of the oxidant and leads to qualitative changes in the empirical rate law.

INTRODUCTION Oxidation of aqueous hydrogen sulfide occurs readily and is of wide importance in the environment.1 Interest in H2S oxidation has increased recently because of the discovery of its role in antibiotic resistance,2 its various functions in eukaryotic organisms,3−7 and its involvement in unique aspects of nonlinear reaction dynamics.8−13 Despite this importance, there are remarkably few reports on the kinetics and mechanisms of these reactions. This scarcity may be attributed to the complexity of the reactions, where the kinetics are often affected by trace metal catalysis and the products are often heterogeneous mixtures of sulfur sols. Sulfur sols complicate data acquisition because their light scattering can influence the absorbance measurements; they also introduce uncertainty in interpretation of the results because of their indefinite composition14 and the potential for surface catalysis. The majority of reports on the kinetics of H2S oxidation deal with 2- (or 4)-electron oxidants. Autoxidation of H2S has been studied widely and is invariably metal-ion catalyzed;15−17 in fact, there is reason to doubt that the noncatalyzed reaction has ever been observed. Oxidation by H2O2 is less susceptible to catalysis, obeys a well-defined rate law, but produces a mixture of sulfate and sulfur sol.11,13,18 Peroxynitrite reacts with a rate law analogous to that of H2O2 but produces species believed to have S−N bonds.19−21 Oxidation by ferrate(VI) yields a variety of products depending on conditions and is proposed to have a radical chain mechanism.22 Oxidation by HOCl is very rapid, produces polysulfides and sulfur sol, and has a relatively simple rate law.23,24 One-electron oxidation of H2S to S•− has been implicated in “cross-talk” between the NO and the H2S biological signaling pathways and in the reactions of H2S with myeloperox© XXXX American Chemical Society

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EXPERIMENTAL SECTION

Reagents and Solutions. Ammonium hexachloroiridate(IV) (41% Ir) purchased from Alfa Aesar was used after recrystallization as described in the literature.30 Analytical and higher grade reagents including Na2S·9H2O, 2,6-pyridinedicarboxylic acid (dipic), NH4Cl, NaOH, acetic acid, anhydrous sodium acetate (all from SigmaReceived: May 27, 2016

A

DOI: 10.1021/acs.inorgchem.6b01289 Inorg. Chem. XXXX, XXX, XXX−XXX

Article

Inorganic Chemistry Aldrich), tris(hydroxymethyl)phosphine (THMP, 95%, from ACROS), tris(2-carboxyethyl)phosphine hydrochloride (TCEP, 98%, from Beantown Chemical), and sodium diphenylphosphinobenzene-3sulfonate (TPPMS, >90%, from TCI) were used without further purification. Our sample of THMP was a white odorless crystalline solid, although it has been reported as a malodorous liquid previously.31 All solutions were prepared with deionized water with a conductivity of 18.2 MΩ−1 cm−1 at 25 °C from a Barnstead Nanopure water purification system. Stock solutions of Na2S (pH > 12) were freshly prepared daily from sodium sulfide nonahydrate crystals after rinsing their surfaces to remove colored polysulfide impurities. These solutions were standardized with a UV−vis spectrometer at 230 nm (ε = 7910 M−1 cm−1)32 near pH 9.5 and protected from oxygen by bubbling with argon gas. Working solutions of Na2S were prepared by dilution of the stock solutions with deoxygenated water. In a few cases, some control experiments in the presence of different phosphine compounds were also carried out by adding the appropriate phosphines into the Na2S working solutions. Solutions of Ir(IV) were prepared either with or without dipic by mixing both (NH4)2IrCl6 and dipic solids with buffer solutions in a volumetric flask and purged with argon for 20−30 min before all kinetic experiments. Buffers were prepared with 0.2 M sodium acetate and the appropriate acetic acid concentrations to maintain the pH. In the stopped-flow experiments equal volumes of buffered Ir(IV) and unbuffered Na2S solutions were mixed, yielding reaction mixtures with a 0.1 M ionic strength. Since H2S has pKa = 6.82 at μ = 0.1 M,33 Na2S is converted primarily to H2S immediately upon mixing the Na2S and IrIV/buffer solutions. For all solutions, only glass and Teflon materials were used for weighing and transferring samples because of both transition metal catalysis28 and reactivity of unknown materials. Product chlorination experiments were performed by generating Cl2(g) from the reaction of aqueous KMnO4 with HCl and passing the Cl2 into the reaction mixture with an N2 stream. After chlorination the residual Cl2 was removed by further N2 sparging. Instruments and Methods. Kinetic experiments were performed at 25 ± 0.1 °C on a Hi-Tech SF-51 stopped-flow spectrophotometer in the 1.0 cm optical path length configuration and equipped with a thermostat apparatus; reactions were monitored with Olis 4300 data acquisition and analysis software at 488 nm, which is a characteristic absorption wavelength of [IrCl6]2− (ε = 3980 M−1 cm−1).30 Some additional kinetic data were collected at 370 nm in order to monitor the formation of sulfur-containing products.18 Reactions were always investigated under pseudo-first-order conditions with sulfide in excess and with pH values between 3.6 and 5.2. Six repetitions were performed for every reaction, and a mean kobs was obtained by fitting the kinetic curves of 488 nm with an exponential decay function. Twodimensional nonlinear fitting for dependence equations kobs = f([H2S]tot) and kobs = f([H+]) were accomplished with the Prism 5 software package, while surface fitting for the two-variable function kobs = f([sulfide],[H+]) was conducted with the Datafit 8 software package. All fits were performed with weighting as the inverse square of kobs. Cyclic voltammetry measurements of the products was performed on a BASi 100B electrochemical analyzer with a BASi C3 cell stand by using a glassy carbon disk electrode (d = 3.2 mm, BASi MF-2012) as the working electrode, a Ag/AgCl (3 M NaCl) reference electrode (E° = 0.205 V vs NHE), and a Pt wire counter electrode. 31 P NMR spectra were obtained on a Bruker AVANCEII 250 MHz instrument with samples in D2O and chemical shifts reported relative to 5 mM H3PO4 as an internal standard.

shown that dipic has several advantages over the ligands surveyed in that report, and it is a very effective reagent for suppressing copper catalysis in oxidation of sulfur-containing compounds by [IrCl6]2−.30 Here we use dipic to suppress copper catalysis of the [IrCl6]2−−sulfide reaction. In a preliminary experiment at pH 4.5, UV spectra showed that dipic does not react with H2S. However, the half-life (t1/2, s) for the reaction of IrIV with H2S increased from 0.013 and 0.458 s upon adding 1 mM dipic at pH 4.64 with [IrIV]0 = 0.056 mM, [H2S]tot = 1 mM; moreover, the loss of IrIV obeyed excellent pseudo-first-order kinetics in the presence of dipic. As shown in Figure S1, the values of kobs become independent of [dipic] above 6 mM. Thus, all further experiments reported below were performed with 6 mM dipic. Kinetics. The kinetics of the [IrCl6]2−−sulfide reaction with 6 mM dipic was examined as a function of sulfide concentration and pH by monitoring the consumption of [IrCl6]2− at 488 nm. Experiments were conducted under pseudo-first-order conditions with sulfide in excess and with [IrIV]0 = 0.05 mM. Figure 1 (data in Table S1) displays the results as a series of plots of

Figure 1. Sulfide concentration dependence of kobs in the [IrCl6]2−− sulfide reaction at different pH. Solid lines are the results of fitting with eq 1. [IrIV]0 = 0.05 mM, [dipic] = 6 mM. Acetate buffers, μ = 0.1 M, and 25.0 °C.

kobs vs [H2S] at pH values ranging from 3.71 to 4.76 with [H2S] = 0.25 to 2.5 mM. These plots clearly demonstrate that the dependence on sulfide concentration is nonlinear and that there is a pronounced pH dependence. At each pH value the data are fit quite well by eq 1 kobs = ka[H 2S]tot + k b[H 2S]tot 2

(1)

Fitted values of ka and kb at each pH are collected in Table S2. Both ka and kb increase with increasing pH, and satisfactory fits are obtained with eqs 2 as shown in Figure S2 ka =

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RESULTS Effect of dipic. Strong catalysis of the [IrCl6]2−−sulfide reaction by transition metal ions, specifically submicromolar levels of copper, has been reported previously, and it leads to nonpseudo-first-order kinetic traces for Ir(IV) consumption.28 Despite efforts with various chelating ligands, a tractable rate law was not attained in that study.28 It has subsequently been

A1(1/[H+])2 B1(1/[H+])2 and k = b 1 + A 2 (1/[H+]) 1 + B2 (1/[H+])

(2)

These results lead to an overall expression for kobs given by eq 3 kobs =

k1[H 2S] tot K ′[H+] + [H+]2

+

k 2[H 2S] tot 2 K ″[H+] + [H+]2

(3)

The corresponding overall empirical rate law is eq 4 B


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Inorganic Chemistry −

⎧ k1[H 2S] tot d[Ir IV] = [Ir IV]⎨ + + 2 dt ⎩ K ′[H ] + [H ] +

⎫ ⎬ K ″[H ] + [H ] ⎭

result was reported by Novoselov and Muzykantova, although obtained in the absence of chelating agents such as dipic.35 The sulfur-containing products were detected by UV−vis spectroscopy. Figure 3 displays the absorbance variation at 370

k 2[H 2S] tot 2 +

+ 2

(4)

An excellent fit of all the kobs values (Table S1) is obtained with eq 3, providing a low average deviation of 4.4% for kobs and the following parameters: k1 = (2.3 ± 0.9) × 10−5 M s−1, k2 = (1.08 ± 0.15) × 10−3 s−1, K′ = (1.2 ± 0.5) × 10−3 M, K″ = (7.2 ± 1.5) × 10−5 M. Product Identification. Cyclic voltammetry (CV) was used to determine the iridium-containing reaction product. These experiments consist of a comparison of the CVs of the IrIV reactant solution and the product mixture as shown in Figure S3. The IrIV reactant solution contained 1 mM IrIV, acetate buffer, and dipic, was purged with nitrogen, and yielded a reversible CV with E1/2 = 0.677 V vs Ag/AgCl and ΔEp/p = 0.072 V. The product mixture was prepared similarly except with the addition of 0.5 mM H2S; under these conditions of relatively high [IrIV]0 the IrIV is consumed completely and the product mixture becomes turbid due to formation of sulfur. The product mixture CV consisted of a reversible wave with E1/2 = 0.676 V (0.882 vs NHE) and ΔEp/p = 0.072 V. The close similarity of the two CVs provides strong evidence that the iridium-containing product is [IrCl6]3−. The presence of [IrCl5(H2O)]2− would have been indicated by a pair of waves about 140 mV higher than for [IrCl6]3−.34 Further evidence for the iridium-containing products is provided by UV−vis spectrophotometry. Using excess sulfide over a low concentration of [IrCl6]2−, the characteristic spectrum of [IrCl6]2− in the range of 400−650 nm provides a specific and sensitive signal for the qualitative and quantitative analysis of the oxidant. Three spectra were collected, corresponding to 0.1 mM [IrCl6]2− solutions with and without 1 mM sulfide and the product solution after chlorination, respectively. The other conditions are [dipic]0 = 6 mM, pH 4.64. The spectra, shown in Figure 2, demonstrate a 100% recovery of the initial [IrCl6]2− after the chlorination of product solution. This further evidence confirms that the reduction of [IrCl6]2− by sulfide proceeds with the formation of [IrCl6]3− as the unique chloride/iridium-containing product. A similar

Figure 3. Kinetic traces at different wavelengths for the [IrCl6]2−− sulfide reaction. [IrIV]0 = 0.5 mM, [H2S]tot = 2.5 mM, [dipic] = 6 mM, pH 3.7.

and 488 nm with time upon mixing a solution of 0.5 mM IrIV and 2.5 mM H2S at pH 3.7. The 370 nm trace is characterized by an initial absorbance increase that occurs on the same time scale as the absorbance decrease at 488 nm. This is followed by a rapid decrease at 370 nm and then a much slower apparent increase. The slow increase is also detected at 488 nm and is accompanied by formation of a white turbid suspension. We interpret these results as indicating that [IrCl6]2− is consumed with the immediate formation of polysulfides such as S42− and S52−, which have an absorbance maximum at 370 nm.18 Under the acidic conditions of this experiment the polysulfides disproportionate quickly to give clear colloidal sulfur solutions and at longer times to give turbid colloidal sulfur; the latter leads to an apparent absorbance increase at both wavelengths. We also observed that the turbid product solution is stable for several days before the irreversible transformation into orthorhombic sulfur S8, a water-insoluble precipitate. The above results can be summarized by the overall reaction 5 2[IrCl 6]2 − + H 2S → 2[IrCl 6]3 − + S + 2H+

(5)

Effects of Phosphines. From the results reported above it is evident that light scattering from the sulfur products can become significant, raising the possibility that the rate law in eq 4 is influenced by this effect. A further concern is that the polysulfides produced initially could be susceptible to oxidation by IrIV, thus opening the possibility for a chain reaction mechanism. Because of these two concerns, results are reported here on the effects of water-soluble phosphines, which are known to be highly thiophilic.36 The phosphines selected for study are THMP (tris-hydroxymethylphosphine), TPPMS (triphenylphosphine monosulfonate), and TCEP (tris-carboxyethylphosphine) TCEP and THMP have been used previously as sulfur-atom acceptors but not to our knowledge in the context of oneelectron oxidation reactions.31,37−39 By comparing the UV spectra of mixtures with spectra calculated from the individual components it was found that neither dipic nor H2S reacts directly with the three phosphines. On the other hand, these three phosphines do react with [IrCl6]2− as shown in Figure S4; analogous one-electron oxidation of phosphines has been reported previously for

Figure 2. UV−vis spectral changes in the reaction of sulfide with IrIV. Black = reactant IrIV, red = product, green = IrIV after chlorination. [IrIV]0 = 0.1 mM, [H2S]tot = 1 mM, [dipic] = 6 mM, pH 4.64. C


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Inorganic Chemistry

[Fe(phen)3]3+.40 The sequence of reaction rates is as follows: TCEP > TPPMS > THMP, but in all cases the rates are slower than the rates of oxidation of H2S. Nevertheless, in the presence of these phosphines the reaction of H2S with [IrCl6]2− is dramatically altered. As Figure 4a at 370 nm demonstrates, the

Figure 5. Dependence on two phosphines of the half-life of the [IrCl6]2−−sulfide reaction. (Top) R3P = THMP. (Bottom) R3P = TPPMS. [H2S]tot = 2.5 mM, [dipic] = 6 mM, pH 4.85.

where the inhibition of H2S oxidation has reached its limiting value. An analysis of the products of the reaction of H2S with IrIV in the presence of THMP by 31P NMR spectroscopy revealed the formation of a species characterized by a singlet with a chemical shift of 47.4 ppm. This result is identical to that reported for the corresponding phosphine sulfide (THMPS, δ = 47.8 ppm).41 Accordingly, we infer that the overall reaction under these conditions has the following stoichiometry

Figure 4. Effect of phosphines on the reaction of [IrCl6]2− with sulfide: [IrIV]0 = 0.1 mM, [H2S]tot = 2.5 mM, [R3P] = 0.25 mM, [dipic] = 6 mM, pH 4.85; (A) 370 and (B) 488 nm.

2[IrCl 6]2 − + H 2S + THMP → 2[IrCl 6]3 − + THMPS + 2H+

colloidal sulfur that forms slowly in the absence of phosphines is not formed in the presence of 0.25 mM phosphine. Moreover, Figure 4b at 488 nm shows that the consumption of IrIV in H2S solutions is not accelerated by the presence of the phosphines. In fact, THMP and TPPMS slow the reaction quite distinctly. The direct reactions of THMP and TPPMS with IrIV can be ignored under these conditions since they have half-lives that are greater than that for the H2S reaction by factors of >400 and 23, respectively. The kinetic inhibition by THMP and TPPMS becomes more pronounced with increasing phosphine concentration, and at high concentrations it reaches a limiting value that is the same for both phosphines; this effect is shown in Figure 5 by plots of t1/2 for IrIV loss as a function of phosphine concentration. This limiting half-life (∼0.4 s at 2.5 mM H2S) is attained at about 0.2 mM TPPMS, but 3 mM THMP is required to achieve the same result. For both phosphines, even at these limiting concentrations their direct reactions with IrIV are considerably slower than the reaction with H2S (Table S3). Overall, THMP is the phosphine that has the lowest rate of direct oxidation by IrIV at concentrations

(6)

Kinetics with THMP. A detailed study of the kinetics of the H2S/IrIV reaction was performed with 0.05 mM IrIV, 5 mM THMP, and 6 mM dipic, i.e., a limiting concentration of THMP. Under these conditions the loss of IrIV obeys pseudofirst-order kinetics, with the rate constant designated as kobs. The direct (slow) oxidation of THMP by IrIV occurs with a rate constant kobs,THMP, which is applied as a small correction to yield Δkobs (= kobs − kobs,THMP). The dependence of Δkobs on [H2S]tot was investigated by varying the sulfide concentration from 0.5 to 3 mM at pH 4.83. These results are displayed as a linear plot of Δkobs vs [H2S]tot in Figure 6 (data in Table S4); linear regression analysis indicates that the reaction is first order with respect to [H2S]tot. The pH dependence was studied in 2.5 mM H2S at six different pH values between 3.9 and 5.2. Figure 7 is the plot of log(Δkobs) versus pH (data in Table S5); it demonstrates an excellent linear relationship with a slope = 0.995, indicating an inverse dependence of Δkobs on hydrogen-ion concentration: Δkobs ∝ 1/[H+]. There is a small pH dependence for kobs,THMP, D


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Inorganic Chemistry

having an apparent second-order dependence on H2S concentration, while no such term appears in the rate law when phosphines are present. An important set of clues to the role of the phosphines in the mechanism is that (1) the phosphines inhibit the rate of reduction of IrIV, (2) the degree of inhibition depends on the specific phosphine at low phosphine concentration, and (3) the degree of inhibition becomes independent of the phosphine identity or concentration at higher (limiting) phosphine concentrations. A mechanism that explains the results at limiting (high) phosphine concentrations is as follows H 2S ⇌ HS− + H+ Figure 6. Sulfide dependence of the kinetics of the [IrCl6] −sulfide reaction in the presence of THMP. [dipic] = 6 mM, [THMP] = 5 mM, [IrIV]0 = 0.048 mM, pH 4.83, μ = 0.1 M, 25.0 °C.

K aH2S

(9)

2−

[IrCl 6]2 − + HS− → [IrCl6]3 − + HS• HS• ⇌ S•− + H+

ket

(10)

K a,rad

S•− + HS− ⇌ HS2•2 −

(11)

K rad

(12)

[IrCl6]2 − + HS2•2 − → [IrCl6]3 − + HS−2

R3P + HS2− → R3PS + HS−

(fast)

kR3P(fast)

(13) (14)

This mechanism leads to the observed overall stoichiometry in eq 6 and the derived rate law −

Consequently, the simple rate eq 7 for the [IrCl6]2−−H2S reaction is established under weakly acidic conditions in the presence of THMP [Ir IV][H 2S] tot d[Ir IV] = Δkobs[Ir IV] = k dt [H+]

(7) −3

−1

The value of the rate constant k is (8.9 ± 0.3) × 10 s as derived from the slope of a plot of Δkobs vs [H2S]tot/[H+] (Figure S5). The kinetic experiments described above were performed between pH 3.9 and 5.2, while the pKa of H2S is 6.82. As a result, more than 97% of the [H2S]tot was present in the form of H2S rather than HS−. Thus, to a good approximation eq 7 can be written as eq 8 IV

■

[IrCl 6]2 − + HS− → [IrCl 6]3 − + S•− + H+

IV

[Ir ][H 2S] d[Ir ] k − =k = H S [Ir IV][HS−] dt [H+] Ka 2

(15)

A comparison of this derived rate law and the empirical rate law 8 identifies k/KaH2S as 2ket. Given that the pKa of H2S is 6.82, the measured value of k leads to ket = 2.9 × 104 M−1 s−1. The mechanism proposed above is supported by several other lines of evidence. First, the rapid equilibrium association of S•− with HS− (eq 12) has already been reported,27,42 with Krad = 9 × 103 M−1, and the acid dissociation of HS• (pKa = 3.4 ± 0.7) is strong enough to permit this process.43 Second, the oxidation of HS2•2− to HS2− is strongly driven with E°(HS2−/ HS2•2−) = −1.13 V.43 Third, the ability of phosphines to act as sulfur atom acceptors is well supported with SCN− and S2O32− as sulfur atom donors.44 Fourth, the rate constant for the initial electron-transfer step, ket, is compatible with the driving force for that step, since E°(IrIV/IrIII) = 0.88 V and E°(HS•/HS−) = 1.13 V;43 although this step is moderately uphill, detailed balancing yields an acceptable value of 5 × 108 M−1 s−1 for the reverse rate constant. A conceivable alternative to the rate-limiting step is the concerted PCET (CPET) oxidation of HS− as in

Figure 7. Plot of log(Δkobs) against pH for the [IrCl6]2−−sulfide reaction in the presence of THMP and dipic. [dipic] = 6 mM, [THMP] = 5 mM, [IrIV]0 = 0.046 mM, [H2S]tot = 2.5 mM, μ = 0.1 M, 25.0 °C.

−

d[IrCl 62 −] = 2ket[IrCl 62 −][HS−] dt

ket

(16)

Although this CPET mechanism is compatible with the results, we consider that it is unlikely to occur given that it is thermodynamically less favorable than the direct electrontransfer process in eq 10. In the absence of phosphines we propose a different fate for the HS2−, where it undergoes further oxidation by IrIV

(8)

DISCUSSION From the above results it is immediately evident that the oxidation of H2S by [IrCl6]2− is profoundly affected by the presence of water-soluble phosphines. Without phosphines present the reaction yields a turbid suspension of sulfur colloids, while the presence of low concentrations of phosphines leads to the production of clear solutions of phosphine sulfides. Another notable difference is in the empirical rate laws 4 (no phosphine) and 7 (limiting phosphine): in the absence of phosphines there is a term

[IrCl 6]2 − + HS2− → [IrCl6]3 − + HS2•

HS2• + HS− → HS3•2 − + H+

fast

further polysulfide chain lengthening

(17) (18)

[IrCl 6]2 − + HS3•2 − → [IrCl6]3 − + HS3−

E

fast

fast

fast

(19) (20)


Inorganic Chemistry polysulfides → sulfur colloid

slow

■

(21)

AUTHOR INFORMATION

Corresponding Author

This mechanism describes the accumulation of polysulfides during the reduction of IrIV and the conversion of the polysulfides to sulfur colloids after the IrIV is consumed. It also leads to a rate law analogous to eq 15 except that the stoichiometric factor is n, where n is equal to twice the polysulfide chain length −2. If the chain length increases with increasing concentration of H2S then the value of n would also increase; this dependence of the stoichiometric factor on [H2S] could be the origin of the nonlinear dependence of kobs on [H2S] shown in eq 1. The data in Figure 5 show that the reaction rates at 2.5 mM H2S are reduced by approximately a factor of 5 at limiting concentrations of phosphines. In terms of the above mechanism this result implies that the average polysulfide chain length is about 7 in the absence of phosphines. From the data in Figure 5 it is also clear that the limiting concentration of phosphine is lower for TPPMS than for THMP, and the data in Figure 4 imply that TCEP would have a much higher limiting concentration. This variation in limiting phosphine concentrations can be understood if the phosphines compete with [IrCl6]2− for the HS2−, and the values for kR3P in eq 14 vary in the sequence k(TCEP) < k(THMP) < k(TPPMS). It is of interest to compare the reactions of various oneelectron oxidants with HS−. The powerful oxidants OH• and SO4•− react at nearly diffusion-controlled rates, as expected.21,27 The milder oxidant CO3•− (E°(CO3•−/CO32−) = 1.57 V)43 reacts somewhat more slowly (k = 2 × 108 M−1 s−1),27 while NO2• (E°(NO2•/NO2−) = 1.04 V)43 reacts even slower (k = 1.2 × 107 M−1 s−1).21 [IrCl6]2− extends this trend into the realm of stable oxidants with ket = 2.9 × 104 M−1 s−1 and E°(IrIV/IrIII) = 0.88 V. A practical consequence of the greater rate constants for the radical oxidants is that the initial electrontransfer step occurs so rapidly that its study is not complicated by the concurrent oxidation of the sulfide radical and polysulfide products. An analysis of the electron-transfer rate constant in terms of the Marcus cross relationship (see SI) with [IrCl6]2− as the oxidant yields an effective self-exchange rate constant for the HS•/HS− couple of 7 × 106 M−1 s−1; the high value of this rate constant may be an indication that the reaction has some inner-sphere character. Despite the uncertainty regarding inner/outer-sphere character, the calculated self-exchange rate constant provides a basis for predicting the rates of reaction of HS− with other one-electron oxidants. An important implication of the present work relates to the behavior of H2S in biological systems. It is now recognized that H2S has essential functions in lower and higher organisms and is present in diverse natural circumstances.3,4,6,7 Our results now show that H2S is susceptible to oxidation at nearphysiological pH by relatively mild one-electron oxidants, many of which can be found in living organisms.

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Article

*E-mail: [email protected]. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS This research was supported by a grant from the ACS-PRF (USA) (#55078-ND3), Fundamental Research Funds for the Central Universities (Grant No. 2015XKZD09), and Priority Academic Program Development of Jiangsu Higher Education Institutions (PAPD).

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REFERENCES

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ASSOCIATED CONTENT

S Supporting Information *

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.inorgchem.6b01289. Tables S1−S5, Figures S1−S5, details of the Marcus calculation (PDF) F


Article

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One-Electron Oxidation of Hydrogen Sulfide by a Stable Oxidant

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