Kinetics and Mechanism of Se-Catalyzed Disproportionation of Bisulfite

Article pubs.acs.org/IECR

Kinetics and Mechanism of Se-Catalyzed Disproportionation of Bisulfite: The Critical Role of Selenosulfate Bentao Yang,† Liyuan Chai,†,‡ Fangfang Zhu,† Xu Yan,† Kaisong Xiang,† and Hui Liu*,†,‡ †

School of Metallurgy and Environment, Central South University, Changsha 410083, China Chinese National Engineering Research Centre for Control & Treatment of Heavy Metal Pollution, Changsha 410083, China

‡

S Supporting Information *

ABSTRACT: Se-catalyzed disproportionation of bisulfite (HSO3−) is a facile and potential wet flue gas desulfurization. In this paper, the kinetics and mechanism of Se-catalyzed disproportionation of HSO3− were systematically investigated. The Se-catalyzed disproportionation of HSO3− is found to be pseudo-first-order with respect to the HSO3− and Se concentrations, and is zeroth-order with respect to the initial pH value. More importantly, the kinetics of Se-catalyzed disproportionation of HSO3− can be divided into two stages, i.e., a rate-controlling stage and an acceleration stage. In the rate-controlling stage, the reaction rate is slow and the intermediate HSeSO3− forms. This intermediate is decisive to accelerate the HSO3− disproportionation, and thus, it plays an important role in the Se-catalyzed process. Furthermore, a kinetic equation is determined according to the results: r0 = k1[HSO3−]0.94[Se]1.13, where k1 = (10.88 ± 0.04) × 10−2 M−1.07·min−1. This kinetics prediction model is fitted well with the experimental data, indicating that this model can satisfactorily describe the experimental kinetic data of Se-catalyzed disproportionation of HSO3−.

the importance of the HSO3− solution in potential applications, especially in environmental treatment, the study of its reaction and transformation in the environment is necessary. As is well-known, HSO3− ions in solution are unstable and easy to convert to other sulfur species via some irreversible side reactions. 12 One of the side reactions is the HSO 3 − disproportionation that produces sulfur or thiosulfate:13,14

1. INTRODUCTION A large amount of SO2 emission from the combustion process of fossil fuels and minerals causes serious problems to the environment and human health.1,2 Various flue gas desulfurization (FGD) processes are used to reduce the SO2 emission.3 Among these processes, the wet scrubbing technology is the most widely used (>87%) because of its high desulfurization efficiency and low capital cost.4 In addition, most of the elements (e.g., Se, As, and Hg) in the flue gas can be simultaneously removed.5 In a wet scrubbing process, absorption of SO2 using liquid is an essential stage. Common adsorption liquids for SO2 include sodium hydroxide, ammonia, organic solvents, and ionic liquids. No matter which liquid is used, the absorption process will produce a solution containing HSO3− or SO32− ions via the following reactions:6 SO2(g) ⇌ SO2(aq)

(1)

SO2(aq) + H 2O(l) ⇌ H+(aq) + HSO3−(aq)

(2)

HSO3−(aq) ⇌ H+(aq) + SO32 −(aq)

(3)

3HSO3−(aq) → S(s) + 2SO4 2 −(aq) + H+(aq) + H 2O(l)

or 4HSO3−(aq) → S2O32 −(aq) + 2SO4 2 −(aq) + 2H+(aq) + H 2O(l)

(5)

HSO3−

These disproportionation reactions happen in many processes, e.g., the Wellman−Lord process, the NH3 steam stripping process, and the regenerable organic amine process.13 Therefore, the studies of eqs 4 and 5 are widely reported. For example, Zhou et al. investigated the effect of SO2 on the thermal degradation of CO2 adsorption liquid in amine scrubbing technology. Their study confirmed that S2O32− could be formed via HSO3− disproportionation (eq 5).15,16 In addition, Talonen et al. found that S(s) could not be formed via HSO3− disproportionation when the temperature was lower than 433 K due to the high activation energy of eq 4.17 Huang et al.14 reported the photodisproportionation of HSO3− using a

These reactions will produce a desulfurization solution loaded with a high concentration of HSO3−. In practical application, the desulfurization solution is regenerated thermally to release a concentrated SO2 stream. Then the concentrated SO2 can be treated to produce sulfuric acid or liquid SO2 (>99.9%).7,8 As for other applications, the HSO3− solution also can be used to reduce Fe3+ to Fe2+ in simultaneous removal of SO2 and NOx from flue gas.9,10 Similarly, HSO3− solution is reported to have the potential to reduce Cr(VI) to Cr(III) to decrease toxicity.11 Considering © XXXX American Chemical Society

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Received: December 18, 2015 Revised: March 12, 2016 Accepted: March 30, 2016

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DOI: 10.1021/acs.iecr.5b04840 Ind. Eng. Chem. Res. XXXX, XXX, XXX−XXX

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Industrial & Engineering Chemistry Research

temperature was varied in the range of 353−373 K with steps of 5 K, and the temperature was maintained constant ±1 K. The flask was put into a water bath (DF-101S, Yu Hua Instrument) with a magnetic stirrer at a speed of 260 rpm. At this stirring speed, the process was chemical reaction controlled rather than diffusion controlled.19 All the experiments were conducted by monitoring the HSO3− concentration as a function of time. The solution pH was measured by a pH meter (S220 K, Mettler Toledo) with an InLab Routine Pro (Mettler Toledo) pH electrode after two-point calibration using NIST/DIN buffer of pH 4.006 and 6.865 (Mettler Toledo). The sulfite concentration in solution was determined by an ion chromatograph (883 basic IC plus, Metrohm Ltd.) with an anion separation column (Metrosep a Supp 4) and an ion chromatography (IC) conductivity detector. All instrument control and data collection were performed by Metrohm chromatography software Magic Net 3.0. Before IC analysis, the solution was sampled every 10−30 min, and was filtered using a 0.22 μm filter to obtain the filtrate without a solid impurity.

standard low-pressure Hg vapor UV lamp (λ = 254 nm) at a low temperature (323 K). The milky solution containing S(s) will be formed when the solution pH value is 0.04 M. This is because part of HSO3− will directly react with Se at a slow reaction rate due to the excess

Figure 2. (a) HSO3− concentration as a function of time. The plot of 1 M was enlarged in the inset. (b) Initial rates at various HSO3− concentrations ([HSO3−] = 1−4 M, [Se] = 0.04 M, and T = 373 K).

can also be divided into two stages, i.e., the rate-controlling stage and the acceleration stage. For the rate-controlling stage, the reaction rate increases from (28.5 ± 0.4) × 10−4 to (115.4 ± 3.0) × 10−4 M·min−1 as the HSO3− concentration increases from 1 to 4 M. Also, the initial rate method was used to study the kinetics with respect to the HSO3− concentration (Figure 2b). From the plot of ln(r0) vs ln [HSO3−], the reaction order with respect to the HSO3− concentration is 0.94 (R2 = 0.97). This value is close to unity, indicating that the reaction follows a pseudo-first-order kinetics with respect to the HSO3− concentration. In addition, we find that this reaction order is lower than that with respect to the Se concentration. This result indicates that the rate-controlling stage seems less dependent on the HSO3− concentration than the Se concentration. This further confirms that the rate-controlling stage is the process that produces some intermediate derived from Se. For the acceleration stage, the influence of the HSO3− concentration is much more significant. As listed in Table 2, the reaction rate of the acceleration stage is increased ∼16-fold as the HSO3− concentration increases from 1 to 4 M, whereas the reaction rate of the rate-controlling stage is only increased ∼4-fold. This C


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using the initial rate method. Because the HSO3− concentration is highly dependent on the initial pH at an acidic condition, r0 should be adjusted according to the HSO3− concentration at different pH values as presented in section 1 of the Supporting Information. However, the results of the initial rate method show that the reaction order with respect to the pH is 0.08 (Figure 3b). This reaction order is close to zero. Therefore, the HSO3− disproportionation is zeroth-order with respect to the initial pH value. This also indicates that the initial pH value has an insignificant effect on the first stage.33 Different from the first stage, the pH value shows a notable impact on the acceleration stage. The reaction rate of the acceleration stage increases from (314.9 ± 22.0) × 10−4 to (457.0 ± 17.5) × 10−4 M·min−1 as the initial pH decreases from 4.48 to 3.87 (Table 3). This indicates that the increase of

Table 2. Effect of the HSO3− Concentration on the Reaction Rate of Se-Catalyzed Disproportionation of HSO3− reaction rate × 10−4/M·min−1 HSO3−

concn/M 1 2 3 4

rate-controlling stage 28.5 60.5 74.7 115.4

± ± ± ±

0.4 1.0 1.0 3.0

acceleration stage 65.9 279.8 457.0 716.1

± ± ± ±

0.5 23.5 17.5 37.0

implies that more HSO3− reacts with the intermediate than with Se directly. 3.1.3. Effect of the Initial pH. As presented by eqs 4 and 5, the solution pH will decrease as the reaction progresses. Therefore, pH is an important factor for this process, and thus, the effect of the initial pH was studied at a HSO 3 − concentration of 3 M (Figure 3a). Notably, the change of the

Table 3. Effect of the Initial pH on the Reaction Rate of SeCatalyzed Disproportionation of HSO3− reaction rate × 10−4/M·min−1 initial pH 1.51 2.52 3.87 4.48

rate-controlling stage 55.5 56.2 74.7 46.1

± ± ± ±

1.5 2.0 1.0 0.5

acceleration stage 196.1 294.1 457.0 314.9

± ± ± ±

8.5 15.5 17.5 22.0

the H+ concentration has a positive effect on the acceleration stage. Nevertheless, with a further decrease of the initial pH value to below 2.52, the reaction rate of the acceleration stage sharply decreases from (457.0 ± 17.5) × 10−4 to (196.1 ± 8.5) × 10−4 M·min−1. This phenomenon is due to the acidolysis of the intermediate.34 Therefore, there is an optimal initial pH value in this Se-catalyzed disproportionation of HSO3−. 3.1.4. Effect of Temperature. The effect of temperature was examined at 353 ± 1, 358 ± 1, 363 ± 1, 368 ± 1, and 373 ± 1 K at pH 3.87 (Figure 4a). As seen, the reaction rate of HSO3− increases as the temperature increases in both the ratecontrolling and acceleration stages. For example, as the temperature rises from 353 to 373 K, the reaction rate of the rate-controlling stage increases from (32.0 ± 2.0) × 10−4 to (74.7 ± 1.0) × 10−4 M·min−1, and the reaction rate of the acceleration stage increases from (165.2 ± 3.5) × 10−4 to (457.0 ± 17.5) × 10−4 M·min−1 (Table 4). Furthermore, the study of the effect of temperature can be used to provide an evaluation of the reaction activation of Se-catalyzed disproportionation of HSO3−. The apparent rate constant was calculated at different temperatures, and an Arrhenius plot of ln(apparent rate constant) against 1/T (K−1) was made. The apparent activation energy (Ea) for this reaction is calculated from the slope of the straight line as 44.49 kJ/mol (Figure 4b). This value of Ea is much lower than that without addition of a catalyst (306.98 kJ/mol). More importantly, this value is lower than that with addition of thiosulfate as the catalyst (188.36 kJ/ mol).13 Therefore, Se is more effective than thiosulfate, and thus is an alternative catalyst for HSO3− disproportionation. 3.2. Investigation of Intermediates. Particular attention needs to be paid to the acceleration stage of Se-catalyzed disproportionation of HSO3−. On the basis of the above results and our previous study, the acceleration stage is highly related to some intermediates. The formation of these intermediates is just the rate-controlling stage. In this system, the possible intermediates include SeSO32−, S2O32−, and Se(sol) as shown in our previous study.19 Here, we suppose SeSO32− rather than

Figure 3. (a) HSO3− concentration as a function of time at different initial pH values. (b) Initial rates at various initial pH values ([HSO3−] = 3 M, [Se] = 0.04 M, and T = 373 K).

initial pH has an obvious impact on the initial concentration of HSO3−. Specifically, the initial HSO3− concentration increases with the pH decreasing from 4.48 to 3.87, while it decreases with further a decrease of pH from 3.87 to 1.51. This phenomenon is attributed to the ionization equilibrium between SO32−, HSO3−, H2SO3, and H+.32 However, the initial pH has an insignificant impact on the first stage. The first stage for all initial pH values is finished within about 100 min, and the values of the reaction rate are all in the range of (46.1 ± 0.5) × 10−4 to (74.7 ± 1.0) × 10−4 M·min−1. Furthermore, the kinetics with respect to the initial pH value was also studied D


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Figure 5. HSO3− concentration as a function of time with addition of SeSO32− or S2O32− ([HSO3−] = 3 M, [Se] = 0.04 M, initial pH 4.50, and T = 373).

(52.3 ± 0.9) × 10−4 M·min−1 and from (314.9 ± 22.0) × 10−4 to (426.4 ± 7.8) × 10−4 M·min−1, respectively. This means that S2O32− can accelerate this process, but the effect is not very significant, possibly due to the low temperature. By comparison, a faster reaction rate is attained over SeSO32−, especially for the first stage. The time for the first stage is greatly shortened to about 60 min, and the time for the complete conversion is shortened to less than 115 min. The reaction rates of the first stage and second stage are further increased to (74.8 ± 22.0) × 10−4 and (451.7 ± 4.2) × 10−4 M· min−1, respectively. Such a fast reaction rate derived from SeSO32− confirms that SeSO32− is the predominant substance to accelerate the Se-catalyzed disproportionation of HSO3−. To further verify the role of SeSO 3 2− in HSO 3 − disproportionation, it is necessary to consider the kinetics that only SeSO32− is presented in the system without addition of Se as the catalyst. For this purpose, a 0.05 M exotic SeSO32− solution was added to the HSO3− solution (3 M) at different pH values without addition of Se (Figure 6). Unexpectedly, the process for only adding SeSO32− still has a slow phase. More importantly, the initial pH value also has an obvious influence on this slow phase. Specifically, the time for this slow phase is 90, 60, and 30 min at initial pH values of 4.54, 3.98, and 3.71, respectively. This phenomenon indicates that an appropriate

Figure 4. (a) HSO3− concentration as a function of time at different temperatures. (b) Arrhenius plots for the reaction rate ([HSO3−] = 3 M and [Se] = 0.04 M).

Table 4. Effect of Temperature on the Reaction Rate of SeCatalyzed Disproportionation of HSO3− reaction rate × 10−4/M·min−1 T/K 353 358 363 368 373

rate-controlling stage

acceleration stage

± ± ± ± ±

165.2 ± 3.5 204.1 ± 6.0 356.0 ± 7.0 364.3 ± 12.5 457.0 ± 17.5

32.0 40.8 41.8 56.9 74.7

2.0 4.0 1.0 2.0 1.0

the other two is the dominant intermediate according to the following reasons: (i) Although Se(sol) can be formed in this process, its existence is very transient. In SO32− solution, Se(sol) can be easily dissolved to form SeSO32− at room temperature.25,35 (ii) The optimal temperature for S2O32−-catalyzed reaction is 403−443 K. When the temperature is 373 K, the HSO3− disproportionation rate will be unacceptable only if the S2O32− concentration is >0.5 M.13 However, these hypotheses should be verified by experiments. For this purpose, 0.05 M SeSO32− or S2O32− was added before the Se-catalyzed disproportionation of HSO3− was initiated. The results of HSO3− concentration vs time are displayed in Figure 5. As seen, without addition of SeSO32− or S2O32−, the time for the first stage is about 105 min, and the time for complete conversion is nearly 180 min. Adding S2O32− can slightly shorten the time for both the first stage (90 min) and the second stage (50 min). The reaction rates of the first stage and second stage increase from (46.1 ± 0.5) × 10−4 to

Figure 6. HSO3− concentration as a function of time in the systems at various initial pH values with constant [HSO3−] (3 M) and [SeSO32−] (0.05 M) at 373 K. E


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where k1 = (10.88 ± 0.04) × 10−2 M−1.07·min−1, which is the average value of k1a and k1b. To further validate eq 11, a least-squares optimization of eq 10 was performed on the basis of the data in Tables 1 and 2. The least-squares optimization is performed by the minimization of the standard deviation with variation of the values of a, b, and k1. The calculation of the standard deviation by the leastsquares optimization method is presented in section 2 of the Supporting Information. According to the experimental results in Table 1, the model values of a, b, and k1 are calculated to be 0.92, 1.13, and 10.67 × 10−2 M−1.07·min−1, respectively, with a standard deviation of 0.15. When the experimental results in Table 2 are used, the model values of a, b, and k1 are calculated to be 0.96, 1.13, and 10.89 × 10−2 M−1.07·min−1, respectively, with a standard deviation of 0.05. These values are in good agreement with the experimental results, i.e., a = 0.94, b = 1.13, and k1 = (10.88 ± 0.04) × 10−2 M−1.07·min−1, indicating the good reliability of the kinetic model (eq 11). Furthermore, the comparison between the experimental data and the calculated model is presented in Figure 7. As seen in

H+ concentration is important for SeSO32−-catalyzed disproportionation of HSO3−, possibly because H+ is one of the reactants to form the active and operative intermediate. The operative intermediate will react with HSO3− rapidly.36,37 On the basis of the above results, we proposed a three-step mechanism for Se-catalyzed disproportionation of HSO3−. First, solid Se was dissolved in an aqueous solution to form nascent SeSO32− ions. The rate of this step is slow and highly related to the Se concentration. The second step is the activation of the intermediate; i.e., nascent SeSO32− ions combine with H+ ions at optimal pH to form HSeSO3−, which is the active and operative intermediate.38 The activation step is a key process in the catalytic reaction but is always too fast to be detected.39,40 Third, HSeSO3− reacts with HSO3− to produce the final products. Because the second step is very fast, only the first step and the third step are observed in the Secatalyzed disproportionation of HSO3−. Differently, an exotic SeSO32−-catalyzed HSO3− solution without Se has no need to experience the slow dissolution step. However, these exotic SeSO32− ions are more stable than the nascent SeSO32−. They are inactive, and thus, the activation step becomes the slow step. 3.3. Model Construction. The reactions for the three-step mechanism based on the above results are as follows: HSO3− + Se → SeSO32 −(active) + H+ +

2−

H + SeSO3 (active) → HSeSO3

2−

HSO3− + HSeSO32 − → product + Se

(k1)

(6)

(k 2 )

(7)

(k 3)

(8)

The complete rate equation is −

d[HSO−3 ] = k1[HSO3−]a [Se]b dt + k 2[H+]c [SeSO32 −(active)]d + k 3[HSO3−]a ′ [HSeSO3−]e

(9)

Figure 6 shows that direct addition of SeSO32− instead of Se to the HSO3− solution shortened the time of the first stage. This phenomenon indicates that the first stage in Figures 1−5 is the formation of SeSO32− from Se (eq 6), which is slow and is the rate-controlling step. Considering the fast reaction rate of eqs 7 and 8 as mentioned above, the complete rate equation is simplified and only the rate-controlling step (eq 6) is involved: r0 = −

d[HSO−3 ] = k1[HSO3−]a [Se]b dt

(10)

The reaction orders with respect to the HSO3− and Se concentrations were calculated by plotting ln(r0) with ln [C],41 where [C] is the concentration of the reactant. According to the results in sections 3.1.1 and 3.1.2, the values of a and b are determined to be 0.94 and 1.13, respectively. The apparent rate coefficient with respect to the HSO3− concentration (Figure 2) (kar = k1[Se]b) is 2.85 × 10−3 min−1, which yields a rate constant of k1a = 10.84 × 10−2 M−1.07·min−1. The rate constant with respect to the Se concentration is k1b = 10.92 × 10−2 M−1.07·min−1 (Figure 1). Therefore, the following expression is operational for HSO3− disproportionation by Se catalysis: −

d[HSO−3 ] = k1[HSO3−]0.94 [Se]1.13 dt

Figure 7. Comparison of the experimental data with the catalytic model (a) in the systems with excess [HSO3−] (3 M) over [Se] (0.01−0.10 M) at 373 K and (b) in the systems with excess [HSO3−] (1−4 M) over [Se] (0.04 M) at 373 K.

Figure 7a, the calculated model gives a good fit to the experimental results when the Se concentration is from 0.01 to 0.04 M. However, when the Se concentration is higher than 0.04 M, a slight deviation is observed. This is possibly caused by the aggregation of Se particles at high Se concentration, resulting in a decrease of the available surface of the Se catalyst. The aggregation of Se particles was ignored in the model

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Industrial & Engineering Chemistry Research calculation.42−44 The other reason for this deviation at high Se concentration might be derived from the simplification of eq 9 without consideration of eqs 7 and 8. Differently, Figure 7b indicates that the calculated model is fitted well with the experimental data in the test range of the HSO3− concentration.

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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.iecr.5b04840. Correction method to calculate the reaction with respect to the H + ion and calculation of least-squares optimization (PDF)

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REFERENCES

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4. CONCLUSIONS The process of Se-catalyzed disproportionation of HSO3− can be divided into two stages: a rate-controlling stage and an acceleration stage. The kinetics of Se-catalyzed disproportionation of HSO3− was investigated via the initial rate method with respect to the Se concentration, HSO3− concentration, and initial pH value. The results show that Se-catalyzed disproportionation of HSO3− is pseudo-first-order with respect to the HSO3− and Se concentrations, and is zeroth-order with respect to the initial pH value. In addition, the apparent activation energy of Se-catalyzed disproportionation of HSO3− is 44.49 kJ/mol. More importantly, the mechanism for the acceleration stage was explored. We find that SeSO32− is an important intermediate to accelerate this process. At optimal pH, SeSO32− will react with H+ to form HSeSO3−, which is the active and operative substance for the acceleration stage. Furthermore, the kinetic model was determined as r0 = k1[HSO3−]0.94[Se]1.13, where k1 = (10.88 ± 0.04) × 10−2 M−1.07·min−1. This kinetic model prediction fits the experimental data well, indicating the good reliability of the kinetic model.

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k1a, k1b = rate constant obtained under different HSO3− concentrations or different Se concentrations

AUTHOR INFORMATION

Corresponding Author

*Tel.: +86 731 88836921. Fax: +86 731 88710171. E-mail: [email protected]. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS This work was financially supported by the National Natural Science Foundation of China (Grants 51474246 and 51404306) and Major Science and Technology Project in Hunan Province (Grant 2013FJ1009).

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LIST OF SYMBOLS r0 = initial rate [HSO3−] = concentration of HSO3− [Se] = concentration of Se R2 = determination coefficient Ea = apparent activation energy k1 = reaction constant of eq 6 k2 = reaction constant of eq 7 k3 = reaction constant of eq 8 a, b, c, d, a′, e = reaction order [C] = concentration of the reactant kar = apparent rate coefficient G


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Kinetics and Mechanism of Se-Catalyzed Disproportionation of Bisulfite

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