Reaction Kinetics and Mechanism of Iron(II) - American Chemical

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Reaction Kinetics and Mechanism of Iron(II)-Induced Catalytic Oxidation of Sulfur(IV) during Wet Desulfurization Yu Zhang,*,† Jiti Zhou,† Chengyu Li,‡ Shiyuan Guo,‡ and Guodong Wang‡ †

School of Environmental Science and Technology and Key Laboratory of Industrial Ecology and Environmental Engineering, MOE, and ‡Environmental Engineering Design and Research Institute, Dalian University of Technology, Dalian 116024, People's Republic of China ABSTRACT: The reaction kinetics of the Fe(II)-induced catalytic oxidation of S(IV) in a concentration range of 106 M e [Fe(II)] e 1 M was studied. The S(IV) catalytic oxidation reaction order for O2 and HSO3 was found to be 0 and 1 in this iron concentration range. The reaction order toward Fe(II) was 1 for 106 M e [Fe(II)] e 104 M, 0 for 104 M < [Fe(II)] e 102 M, and 0.2 for 102 M e [Fe(II)] e 1 M. At 3050 °C the activation energy values were low but increased after an increase in the ionic strength. The ionic strength had a negative influence on the reaction rate constant in the iron concentration range. By analyzing the radical mechanisms and reaction kinetic results, a simplified model for the Fe(II)-induced catalytic oxidation of S(IV), which enables a better analysis and elucidation of the experimental results, is proposed.

1. INTRODUCTION In aqueous solutions, SO2 causes oxidation reactions upon catalysis by Fe, Mn, and other transition metal ions in the presence of O2. Many researchers have conducted studies on this reaction in different fields such as atmospheric chemistry, flue gas desulfurization, and hydrometallurgy, and they have covered reaction kinetics and mechanisms. In their mechanism studies, Siskos et al.1 proposed a nonradical mechanism for the Mn ion-induced catalytic oxidation of S(IV). Kraft and van Eldik2 also proposed a similar mechanism for the Fe(III)-induced catalytic oxidation of S(IV). However, most researchers have focused on the radical mechanism of transition metal ions in the liquid-phase catalytic oxidation of S(IV).3 For the reaction in the initial stages of the radical mechanism, many researchers believe that a compound composed of metal ions and sulfur is formed.46 However, PasiukBronikowska and Bronikowski7 contend that no complexation occurs between the metal ions and S(IV). Regarding the reoxidation of the reduced metal ions in the radical mechanism, numerous researchers4,6,8,9 have argued that free radicals oxidize the reduced metal ions to high valence metal ions. Therefore, metal ions can participate in the reaction. Regarding reaction termination in the radical mechanism, Huss et al.10 found that inert products form by the reaction between the radicals and free radical scavengers such as those from organic substances. The chain reaction is terminated to stop the aforementioned reaction. Other scholars4,6,8,11 have determined that the reaction termination is caused by the combination of free radicals. Podkrajsek et al.12 verified that carboxylic acids in the atmosphere inhibit the Mn-induced catalytic oxidation of S(IV). This inhibition is caused by the capture of sulfur radicals. Zuo et al.13 found that monochromatic UVvisible light and natural sunlight promote the Fe(III) catalytic oxidation of liquid-phase SO2. They attribute this result to the formation of hydroxyl radicals. Mudgal et al.14 attributed the inhibition of ammonia and ammonium ions to the S(IV) catalytic oxidation reaction, and this was determined r 2011 American Chemical Society

by capturing sulfur radicals and the formation of a Co(II)NH3 compound. The above-mentioned studies further confirmed the radical mechanism of transition metal ions in the liquid-phase catalytic oxidation of S(IV). In addition, Huss et al. proposed a mechanism for the combination of radicals and nonradicals. Their results show that the contribution of the nonradical chain transfer to the overall reaction in Fe catalytic reactions is higher than that found in the Mn catalytic reaction. Therefore, they believe that Fe(III) is more active and applicable than Mn(III).10 Researchers have obtained different results in kinetic studies. Noticeable differences have been observed in terms of reaction order, reaction rate constant, pH effect, activation energy, and the reaction induction period. Scholars have found that the order of the reactions for metal ions, S(IV) and O2, differ when the reaction is terminated differently.15,16 The results of Brandt et al.6 showed that the order of the reaction changes along with variations in the Fe(III), Fe(II), and S(IV) concentrations. Ermakov and Purmal17 concluded that the pH and the [Mn(II)]/[S(IV)] ratio affect the order of the reaction during the Mn-induced catalytic oxidation of S(IV). In different systems the complexity of transition metal ion reactions during the liquid-phase catalytic oxidation of S(IV) leads to different results. Even under the same experimental conditions, different results have been obtained when different metal ions were used as catalysts. Therefore, targeted study should be done in transition metal ions liquid-phase catalytic oxidation of S(IV) on the basis of these informed research results, ideas, and methods for different systems. Zhang et al.,18 Jiang et al.,19 and Zhang et al.20 proposed the use of flue gas desulfurization technology using iron scrap. With regard to the Fe ion-induced catalytic oxidation of SO2 in flue gas, they added iron scrap to their solutions. Iron scrap can consume Received: July 5, 2011 Accepted: December 14, 2011 Revised: December 6, 2011 Published: December 14, 2011 1158

dx.doi.org/10.1021/ie2014372 | Ind. Eng. Chem. Res. 2012, 51, 1158–1165

Industrial & Engineering Chemistry Research the dilute sulfuric acid produced by SO2 removal, and this leads to reactions in a high-pH solution while a high SO2 removal efficiency is obtained. Conversely, SO2 can be transformed into a FeSO4 solution, which is more valuable, has more applications, and can even be used to produce polymeric ferric sulfate (PFS) during desulfurization for sulfur recovery.21,22 Iron-based flue gas desulfurization technology continuously consumes iron scrap by desulfurization to increase the Fe(II) concentration in the absorption solution. An absorption solution with higher concentrations of Fe(II) is thus obtained from the desulfurization system. For desulfurization the mechanism of S(IV) oxidation as a result of SO2 dissolution may change because of changes in the Fe(II) concentrations. However, few studies exist about the reaction kinetics of S(IV) oxidation by a wide range of Fe(II) concentrations.3 Therefore, a better understanding of iron-based desulfurization requires a study into the reaction kinetics of the Fe(II)-induced catalytic oxidation of S(IV) within the Fe(II) concentration ranges encountered in wet desulfurization technology. We used NaHSO3 as a reaction material to study the solution reaction kinetics of the Fe(II)-induced catalytic oxidation of S(IV) over an Fe(II) concentration range of 106 M e [Fe(II)] e 1 M. Our experimental results confirm the reaction orders of the reaction between HSO3 and Fe(II). The concentration of Fe(II) was divided into two ranges. The influence of ionic strength and temperature on the reaction rate constant was studied. Using a reaction kinetics analysis, we examined the reaction mechanism by combining the radical mechanisms and the reaction processes that have been proposed by other researchers.

2. EXPERIMENTAL SECTION 2.1. Materials and Methods. The reagents used in these experiments were of analytical grade and were used without further purification. 2.1.1. Na2S2O3 (0.1 M) Solution Preparation. Na2S2O3 3 5H2O (25 g) was weighed and placed into a 500 mL beaker. Cooled distilled water (300 mL) was then added. After complete dissolution, 0.2 g of Na2CO3 was added to the mixture. This mixture was then diluted to 1 L using distilled water. Finally, the solution was stored in a brown bottle. It was calibrated with K2Cr2O7 after 714 days. 2.1.2. I2 (0.01 M) Solution Preparation. I2 (2.6 g) and KI (8 g) were weighed and placed into a small mortar or beaker. A small amount of water was added and then the mixture was ground or stirred to completely dissolve the I2. Subsequently, the mixture was diluted to 1 L using distilled water and stored in a brown bottle. The bottle was tightly sealed and shaken overnight. Finally, it was calibrated using a Na2S2O3 solution. 2.1.3. FeSO4 (0.001 M) Solution Preparation. Solid FeSO4 (2.7801 g) was weighed and placed into a beaker for dissolution in water. After dilution to 100 mL, 10 mL was pipetted to prepare 100 mL of solution. This process was repeated after which a 0.001 M FeSO4 solution was obtained. 2.1.4. NaHSO3 (0.4 M) Solution Preparation. Solid NaHSO3 (16.6496 g) was weighed and placed into a beaker for dissolution in deionized water. After dilution to 100 mL, 25 mL was pipetted to prepare 100 mL of solution in a 100 mL volumetric flask. A NaHSO3 (0.4 M) solution was then obtained. 2.2. Experimental Setup and Procedures. The reactor used was a 250 mL four-necked round-bottomed flask, which was placed in a super thermostatic water bath, and the temperature was controlled to within (1 °C. The outlets of the flask were

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used to adjust the stirrer and to insert a pH electrode, and for ventilation. A NaOH solution was added dropwise for pH adjustment and sampling. The stirrer was used to keep the solution well-mixed. Air was supplied by a magnetic air pump, and a rotameter was used for quantity measurements. A polyethylene pipe (4 mm diameter) was used as a gas line. The porous end seals were baked so that they served as vent holes. In the experiment, the thermostatic water bath was opened first. A Beekmann thermometer was then adjusted to the experimental temperature. On the basis of the experimental requirements, a solution of a certain volume and ionic strength was prepared using deionized water, a Na2SO4 solution, and a catalyst solution and it was then mixed using a magnetic stirrer. The solution was adjusted to near-experimental pH with H2SO4 or NaOH. When the temperature of the water bath stabilized, the reaction base solution was added to the reaction flask. The magnetic air pump was used to supply air to the flask. Under constant temperature and stirring, ventilation was allowed for 15 min to enable the reaction base solution to reach the required temperature and to saturate it with dissolved oxygen. A certain amount of NaHSO3 solution was added to the reactor, and the response time was monitored. The start time of the reaction (t = 0) was the time at which the NaHSO3 solution was added. During the reaction process air was continuously pumped into the flask reactor to maintain oxygen saturation. Solutions of different NaOH concentrations were added dropwise to maintain a constant solution pH. The samples obtained at different time intervals were rapidly added into a conical beaker containing excess standard iodine solution. After mixing, the solution was placed into a beaker before measuring the S(IV) concentration. A sample with an initial S(IV) concentration (t = 0) was obtained from another set of experiments under similar conditions but without catalysis. 2.3. Analysis. Solutions with high Fe concentrations were measured using the PFS method (GB14591-2006).23 For total iron concentrations, stannous chloride was used to reduce ferric iron to ferrous iron in the acid solutions. After removing the excess stannous chloride with mercury chloride, a standard volumetric solution of potassium dichromate was used for the titration. A standard volumetric solution of potassium permanganate was used for the Fe2+ concentration titration in the acid solution. For low Fe concentrations the total iron and the Fe2+ concentrations were determined using the 1,10-phenanthroline spectrophotometric method (721-type spectrophotometer, Shanghai Analytical Instrument Factory). The solution pH was measured using an acidometer (pHS-25-type acidometer, Shanghai Weiye Instrument Factory). The concentration of O2 in solution was measured using a dissolved oxygen meter (JPB-607 portable dissolved oxygen meter, Shanghai Rex Instruments Factory) which has a DO-952-type oxygen electrode. The HSO3 concentration was measured using the iodometric method, and starch was used as the indicator.24

3. RESULTS AND DISCUSSION The range of Fe(II) concentrations used in previous research varies widely, and the reaction mechanisms considered differ. Therefore, this study was performed under a wide Fe(II) concentration range. The experiments were conducted at 106 M e [Fe(II)] e 1 M and pH = 3. The primary form of S(IV) was HSO3, and, therefore, NaHSO3 was used as the reactive substance. Except for the experiment of the effect of ionic strength on the reaction rate constant, we used excess Na2SO4 electrolyte to adjust the ionic strength to 4.01 M to avoid the change in ionic strength 1159

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Figure 1. Relationship between ln{[HSO3]/[HSO3]0} and time for the Fe(II)-induced catalytic oxidation of HSO3 (T = 30 °C; I = 4.01 M; pH = 3; [HSO3]0 = 0.01 M): [, [Fe(II)] = 106 M; 9, [Fe(II)] = 5 106 M; 2. [Fe(II)] = 105 M; ], [Fe(II)] = 5 105 M; 0, [Fe(II)] = 104 M; Δ, [Fe(II)] = 5 104 M; , [Fe(II)] = 103 M; +, [Fe(II)] = 102 M.

during the reaction processes, which enabled us to obtain more accurate results. With use of a normal power exponent the relevant kinetic equation can be expressed as d½HSO3 ¼ k½FeðIIÞα ½HSO3 β ½O2 γ ð1Þ dt At a specific pH, the reaction rate constant is defined as a function of the ionic strength of the solution and the temperature is expressed as k ¼ f ðI, TÞ

ð2Þ

It is generally accepted that the reaction order toward oxygen has a magnitude of 0,10,2529 which indicates that SO3• + O2 f SO5• is a rapid reaction and is not limited in terms of the total reaction rate.30 Huss et al.10 showed that at O2 concentrations of less than 0.001 M the reaction is unaffected by the O2 concentration. On the basis of Henry’s law, at 25 °C the concentration of dissolved oxygen is around 2 104 M when oxygen is pumped into pure water.28 In sulfate media the concentration of dissolved oxygen is31 "

CO2 =ðM=atmÞ ¼ 5:909 10

6

#

1602:1 0:9407CNa2 SO4 exp T 1 þ 0:1933CNa2 SO4

ð3Þ A higher reaction temperature and stronger ionic strength were used in this research. At 30 °C and at an ionic strength of 4.01 M, the dissolved oxygen concentration was estimated using eq 3 and was found to be 1.11 104 M. The value obtained using a dissolved oxygen analyzer was 1.0 104 M. From preliminary experiments, we found that when air and pure oxygen are adopted as oxidants, the reaction rate is unchanged. The reaction rate is also unchanged when the experiments mixing speed doubles. The results show a reaction order of 0 for oxygen. The reaction order of the reaction between HSO3 and Fe(II) was confirmed in the next experiments. 3.1. Reaction Order with Respect to HSO3 and Iron(II). 3.1.1. Low Iron(II) Concentration. The relationship between ln{[HSO3]/[HSO3]0} and time t under Fe(II) concentrations from 106 to 102 M is shown in Figure 1.

Figure 2. Relationship between kobs and Fe(II) for the Fe(II)-induced catalytic oxidation of HSO3 (T = 30 °C; I = 4.01 M; pH = 3; [HSO3]0 = 0.01 M). (Notation 1E-6 represents, for example, 1 106.)

Figure 1 shows that no induction period is apparent in this reaction. The HSO3 oxidation rate was initially high but gradually slowed as the reaction progressed. We determined that the Fe(II)-catalyzed HSO3 oxidation reaction order with respect to HSO3 was 1, which is similar to the results obtained by many other researchers.4,28,29,32,33 When the Fe(II) concentration was changed from 106 to 104 M, the apparent firstorder reaction rate constant kobs increased with an increase in the Fe(II) concentration. The reaction gave a first-order reaction to Fe(II) in agreement with the results obtained under atmospheric conditions by other researchers. However, the apparent reaction rate constant kobs did not change with a change in the Fe(II) concentration when the Fe(II) concentration was higher than 104 M or even when it reached 0.01 M. Figure 2 shows the effect of Fe(II) concentration on the apparent first-order reaction rate constant kobs. Many other researchers obtained similar results. Barron and O’Hern34 studied the Cu(II)-induced catalytic oxidation reaction of sodium sulfite and found that the reaction rate increased 618 times for Cu(II) concentrations within 107106 M. However, the change on the reaction rate was negligible between 106 and 104 M. Mishra and Srivastava30 studied the Co(II)induced catalytic oxidation of sodium sulfite and found that the reaction order toward Co(II) in the reaction was 0.5 from 2 107 to 2 105 M and 0 from 2 105 to 104 M. On the basis of thermodynamic calculations they found that the Co(II) concentration has no effect on the reaction rate because of limitations in Co(II) solubility. Bal Reddy and van Eldik35 studied Fe(II)-induced S(IV) catalytic oxidation kinetics and mechanisms in acid solutions and found that Fe(II) concentrations of (222) 104 M did not affect the reaction. Berglund et al.5 studied the Mn(II)-induced catalytic oxidation of S(IV) at pH 2.4 and found that the reaction order of the Mn(II) concentration for this reaction was 1 from 0 to 104 M and 0 at concentrations higher than 104 M. Zhang et al.4 found that the reaction order of Fe(III) was 1 from 0 to 0.02 M and that Fe(III) at greater than 0.02 M had no effect on the reaction. Karatza et al.36 found that Fe(II) at greater than 0.03 M had no effect on the reaction. Sato et al.25 also found that the concentration of Fe(II) from 0.08 to 0.44 M did not affect the reaction. The concentration of Fe(II) affected the apparent first-order reaction rate constant, and this was dependent on the different Fe(II) concentration ranges. The reaction rate equation can be 1160



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Figure 3. Relationship between ln{[HSO3]/[HSO3]0} and time for the Fe(II)-induced catalytic oxidation of HSO3 (T = 30 °C; I = 4.01 M; pH = 3; [HSO3]0 = 0.01 M): [, [Fe(II)] = 0.01 M; 9, [Fe(II)] = 0.05 M; 2, [Fe(II)] = 0.1 M; ], [Fe(II)] = 0.4 M; 0, [Fe(II)] = 0.6 M; Δ, [Fe(II)] = 0.8 M; , [Fe(II)] = 1 M.

expressed as

Figure 4. Effect of temperature on the first-order reaction rate constant of the Fe(II)-induced catalytic oxidation of HSO3 (pH = 3; [Fe(II)] = 0.001 M; [HSO3]0 = 0.01 M): [,I = 0.71 M; 9, I = 2.51 M; 2, I = 4.01 M.

was 0.2. This result indicates that the reaction rate of S(IV) oxidation can be expressed as

106 M e [Fe(II)] e 104M: d½HSO3 ¼ k½FeðIIÞ½HSO3 dt

d½HSO3 ¼ k½FeðIIÞ0:2 ½HSO3 dt ð4Þ

104 M < [Fe(II)] e 102 M: d½HSO3 ¼ k½HSO 3 dt

ð5Þ

3.1.2. High Iron(II) Concentration. The relationship between ln{[HSO3]/[HSO3]0} and time t under Fe(II) concentrations from 102 to 1 M is shown in Figure 3. Figure 3 shows that the reaction order toward HSO3 remained 1 as it was free of the influence of an increased Fe(II) concentration. However, Figure 3 also shows that the reaction rate decreased as the Fe(II) concentration increased. The reaction order toward Fe(II) was negative. In experiment, an obvious induction period (300 s long) occurred at an Fe(II) concentration of 1 M. The abovementioned studies showed that increasing the concentration of Fe(II) does not obviously affect the reaction when the Fe(II) concentration is more than 104 M. If the concentration of Fe(II) is increased, the association between the Fe(II) and SO42 ions that form a passive ion pair is enhanced. This affects the dielectric constant of the solution as well as other characteristics, and the reaction rate decreases.37 Conversely, Fe(III) and SO42 are generated by Fe(II) oxidation and form a passive ion pair.5 Some Fe(III) in the solution can generate a precipitate,38 which also affects the reaction rate of S(IV) oxidation. Siskos et al.1 found that adding excess Mn(II) decreases the reaction rate of S(IV). Exponential regression was performed for the apparent firstorder reaction rate constants at different concentrations of Fe(II). The following equation was obtained: kobs ¼ 6:56 104 ½FeðIIÞ0:1859

3.2. Effect of Temperature on the Reaction Rate Constant and the Activation Energy. Many scholars have studied the

effect of temperature on the reaction rate constant of the transition metal ion-induced catalytic oxidation of S(IV). Different experimental conditions have been adopted, and hence the reaction activation energies that have been obtained differ. Some researchers obtained high activation energies between 50 and 130 kJ/mol.10,27,29,32,39 Studies into this reaction have generally been under atmospheric conditions, and, therefore, the selected temperatures have generally been below 30 °C. Some researchers have obtained far lower activation energies. Zhang et al.4 obtained a 22 kJ/mol activation energy. The activation energies of the Fe and Mn catalytic reactions obtained by Lu et al.33 were 10.54 and 9.91 kJ/mol, respectively. Copson and Payne40 found that temperature had no obvious effect on the reaction. All of the studied S(IV) catalytic oxidation reactions were carried out under industrial conditions. Pasiuk-Bronikowska and Sokolowski41 explained that this was caused by differences in the temperature ranges used. The reaction activation energy was found to be higher when a lower temperature range was used to study the effect of temperature on the reaction. The activation energy was lower when a higher temperature range was chosen. Generally, the temperature of absorption solution in the wet flue gas desulfurization process is higher than that of atmospheric conditions. Therefore, the temperature chosen for this study was 3050 °C. Different ionic strengths were used to measure the first-order reaction rate constant k1 at different temperatures. On the basis of the activation energy equation, ln k ¼

ð6Þ

In eq 6, the related coefficient is 0.931. Therefore, we can approximately determine that the reaction order toward Fe(II)

ð7Þ

Ea þ ln A RT

ð8Þ

We found a linear relationship between ln k1 and 1/T, as shown in Figure 4. The activation energy Ea under different ionic 1161



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Table 1. Activation Energy and Preexponential Factor I = 0.71M

I = 2.51M

I = 4.01M

A

2.2445

79.23

134.53

Ea, kJ/mol

16.38

26.44

28.59

strengths was determined using the preexponential factor A, and the values are listed in Table 1. As shown in Table 1, activation energies were 16.38, 26.44, and 28.59 kJ/mol at 3050 °C, with which the ionic strengths were 0.71, 2.51, and 4.01 M. The activation energies in the scope of experimental temperature were lower than those under atmospheric conditions. The activation energy increased as the ionic strength increased. This result is similar to that obtained by Zhang et al.4 and Lu et al.33 Pasiuk-Bronikowska and Sokolowski41 obtained activation energies of 4.3, 11.5, 18.1, 25.5, 36.9, and 44.7 kJ/mol at 2050 °C using acid concentrations of 0, 0.204, 0.408, 0.612, 0.816, and 1.02 M, respectively. Higher activation energies were obtained for lower reaction temperatures. The mechanism of the effect of ionic strength on the activation energy that was obtained in this study is the same as that obtained by Lagrange et al.42 3.3. Effect of Ionic Strength on the Reaction Rate Constant. Changes occur in ionic strength, and this affects the reaction rate to a certain extent because HSO3 is continuously transferred to SO42 during catalytic oxidation. Therefore, studying the effect of ionic strength on the reaction rate is necessary.37 Huss et al.43 and Berglund et al.5 found that the ionic strength has a negative effect on the reaction rate constant. Martin and Hill44 determined that the differences found in the kinetics literature are caused by the different ionic strengths of the solutions used. On the basis of the kinetics equation of Br€onsted,37 k γ γ log ¼ log A B ð9Þ k0 γAB The reaction rate constant k in an electrolyte solution can thus be calculated. In eq 9, k0 represents the reaction rate constant while the ionic strength is 0; A and B stand for Fe(II) and HSO3, respectively. The activity coefficients γA, γB, and γAB can be calculated using the DebyeH€uckel activity coefficient formula 10, which can be expressed as37 log γi ¼

XZi 2 I 1=2 1 þ YaI 1=2

ð10Þ

In eq 10, a represents the minimum distance of the ion, X is the coefficient, Y is the DebyeH€uckel coefficient, Zi represents the ionic valence of ion i, and I refers to the ionic strength of the solution, which can be expressed as37

∑

1 bi Zi 2 ð11Þ 2 In eq 11, bi is the mass molarity of the ion i. Equation 10 can only be applied to a 1:1 electrolyte solution with an ionic strength of less than 0.01 M. This increased ionic strength may offset eq 10. The corrected DebyeH€ukel equation, that is, the Guggenheim activity coefficient equation, has, therefore, been adopted.37 In eq 12, C is the specific interaction parameter. I ¼

XZi 2 I 1=2 þ CI log γi ¼ 1 þ I 1=2

ð12Þ

Figure 5. Effect of ionic strength on the reaction rate of the Fe(II)induced catalytic oxidation of HSO3 (T = 30 °C; pH = 3; [HSO3]0 = 0.01 M): [, [Fe(II)] = 0.001 M; 9, [Fe(II)] = 0.1 M.

Equation 9 can be rearranged to pffiffi ZA 2 þ ZB 2 ZAB 2 pffiffi log k ¼ log k0 X I 1 þ I

! þ ðCA þ CB CAB ÞI

ð13Þ For a solution with a stronger ionic strength, log k is generally linear with (I1/2 + PI). P is an adjustable constant, and we can thus establish that ð14Þ

m ¼ log k0 n ¼ Að

ZA 2 þ ZB 2 ZAB 2 pffiffi Þ 1 þ I

C ¼ CA þ CB CAB Equation 12 can be simplified to pffiffi log k ¼ m n I þ CI

ð15Þ ð16Þ ð17Þ

Researchers have studied the Fe(II)-induced catalytic oxidation reaction of S(IV) under atmospheric conditions. Generally, the Fe(II) concentrations studied have ranged from 107 to 105 M. This paper, however, focuses on the industrial application of flue gas desulfurization. The concentrations of Fe(II) were 0.001 and 0.1 M in studying the effect of ionic strength on the reaction rate constant. The effect of ionic strength on the reaction rate constant is shown in Figure 5. From Figure 5, eq 17 can be used at an ionic strength of 4.01 M. The reaction order toward Fe(II) (Fe(II) = 0.001 M) had a magnitude of 0 according to eq 5, and the obtained reaction rate constant was of the first-order type. After fitting the reaction rate constant from this study, the following equation is obtained when Fe(II) was 0.001 M: pffiffi ð18Þ log k1 ¼ 2:2104 0:3354 I þ 0:0159I In eq 18, the related coefficient R is 0.9918, which confirms that this equation can represent the effect of ionic strength on the reaction rate constant. When Fe(II) was 0.1 M, the following relevant expression was obtained: pffiffi ð19Þ log k ¼ 3:0752 þ 0:0847 I 0:0744I In eq 19 R = 0.9954. At a high Fe(II) concentration the use of eq 19 is sufficient to indicate the effect of ionic strength on the reaction rate constant. 1162



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Millero and Izaguirre45 adopted a similar expression in their study on the effect of ionic strength and the interaction among ions during Fe(II) oxidation. The effect of temperature was included in their formulation, which enabled them to obtain a better result. Gu and Han46 used the same equation for the effect of ionic strength on the reaction rate constant.

4. ANALYSIS OF THE REACTION MECHANISM Many researchers have suggested the mechanism of Fe(II)induced catalytic oxidation of S(IV) based on radical mechanisms. The process is mainly divided into a chain initiation reaction, a chaintransfer reaction, and product formation as well as the termination of the chain reaction. This can be summarized as follows (Fe2+ and Fe3+ are represented by the Fe(II) and Fe(III) hydrolyzed species and the generation of OH• and O2• is disregarded):36,8,10,11,25,30,35,47 chain initiation: FeðIIÞ þ HSO3 h FeHSO3 þ

ðrapid equilibriumÞ

1 1 FeHSO3 þ þ O2 f FeSO3 þ þ H2 O 4 2 þ

FeSO3 f FeðIIÞ þ SO3

•

SO4

þ HSO3 f SO3

K1

FeðIIÞ þ HSO3 R s FFeHSO3 þ s

•

þ HSO4


k2 1 1 FeHSO3 þ þ O2 sf FeSO3 þ þ H2 O 4 2

ð23Þ

SO5 • þ HSO3 f SO3 • þ HSO5 SO5 • þ HSO3 f SO4 • þ HSO4

ð21Þ ð22Þ

chain transfer: SO3 • þ O2 f SO5 •

•

ð20Þ

(1) After generating SO5• radicals, many reactions are required to obtain SO5•. The reaction rate constants of reactions 27, 35, and 36 are around 108 M1 s1. Reactions 24, 25, and 2830 have rate constants around 104107M1 s1. Therefore, reactions 24, 25, and 2830 can be disregarded. (2) Because reaction 25 is insignificant, no additional SO4• radicals are generated. Therefore, reactions 26, 34, and 37 can be disregarded. (3) Plenty of dissolved oxygen is present in the solution, and the generated SO3• radicals result in a reaction based on reaction 23. In addition, the reaction rate of reaction 23 is an order of magnitude higher than that of reactions 32 and 36. Therefore, reactions 32 and 36 can be disregarded. In our work the concentration of the selected HSO3 was high so that abundant free HSO3 was present in the solution. Reaction 35 seldom occurs and can also be disregarded. (4) In the chain-transfer step, reactions 31 and 22 are added to reflect the reduction of Fe(III), in which the order of reaction 22 follows reaction 31 from the reaction sequence. The rate of reaction 22 is slower than that of reactions 2330, so reaction 22 can be disregarded in the chain transfer.5,35 After simplification the reaction mechanism is as follows:

k3

FeSO3 þ sf FeðIIÞ þ SO3 •

ð24Þ ð25Þ

k4

ð41Þ k5

FeðIIÞ þ SO5 • þ Hþ f FeðIIIÞ þ HSO5

ð27Þ

FeðIIÞ þ SO5 • þ Hþ sf FeðIIIÞ þ HSO5

FeðIIÞ þ HSO5 f FeðIIIÞ þ SO4 • þ OH

ð28Þ

s FeðIIIÞ þ SO3 2R FFeSO3 þ s

FeðIIÞ þ SO4 • f FeðIIIÞ þ SO4 2

ð29Þ

FeðIIÞ þ SO3

•

f FeðIIIÞ þ SO3

FeðIIIÞ þ SO3 2 h FeSO3 þ

2


FeSO3 þ f FeðIIÞ þ SO3 •


K2

k6

HSO5 þ HSO3 sf 2SO4 2 þ 2Hþ

ð32Þ

SO3 • þ SO3 • f S2 O6 2 SO4 • þ SO4 • f S2 O8 2

ð33Þ ð34Þ ð35Þ ð36Þ ð37Þ

SO5 • þ SO3 • f S2 O6 2 þ O2 SO4 þ HSO3 f HSO4 þ SO3 •

On the basis of the reaction rate constants of the reactions reported by researchers, the above-mentioned mechanism can be simplified as follows:46


On the basis of the equilibrium approximation principle, reaction 40 is the rate-controlling step because the SO3• radical is generated in this reaction. Subsequently, the reactions are rapid and do not affect the observed reaction kinetics.35 Therefore, the reaction kinetics equation can be represented as

product formation/chain termination:

SO5 • þ SO5 • f S2 O6 2 þ 2O2 f S2 O8 2 þ O2


d½HSO3 ¼ k2 ½FeHSO3 þ dt

HSO5 þ HSO3 f 2SO4 2 þ 2Hþ

ð39Þ ð40Þ

SO3 • þ O2 sf SO5 •

ð26Þ

ð38Þ

ð45Þ

From eq 38, ½FeHSO3 þ ¼ K1 ½FeðIIÞ½HSO3 F ½HSO3 ¼ ½HSO3 F þ ½FeHSO3 þ

ð46Þ ð47Þ

[HSO3]F means the concentration of free HSO3.5 Reactions 46 and 47 can be integrated into eq 45, and the reaction kinetics is represented as d½HSO3 ½FeðIIÞ ¼ k2 1 ½HSO3 dt K1 þ ½FeðIIÞ

ð48Þ

According to literature K1 is large. Berglund et al.5 obtained K1 ≈ 3 104 M1 for Mn(II). The experimental conditions used 1163



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in this study were close to theirs, and, thus, we determined that the K1 of Fe(II) was around 104 M1 leading to 106 M e [Fe(II)] e 104 M:

d½HSO3 ¼ k2 K1 ½FeðIIÞ½HSO3 dt

ð49Þ

Corresponding Author

*Tel.: +86-411-84706240. Fax: +86-411-84706252. E-mail: [email protected].

’ REFERENCES

104 M < [Fe(II)] e 102 M: d½HSO3 ¼ k2 ½HSO3 dt

’ AUTHOR INFORMATION

ð50Þ

As shown in eqs 49 and 50, the S(IV) catalytic oxidation reaction order for HSO3 was found to be 1 in the iron concentration range of 106102 M. The reaction order toward Fe(II) was 1 for 106 M e [Fe(II)] e 104 M and 0 for 104 M < [Fe(II)] e 102 M. Equations 49 and 50 gave the same reaction order toward HSO3 and Fe(II) as the kinetics shown in eqs 4 and 5 during the experiment. This result indicates that the above reaction mechanism given for Fe(II) during the liquid-phase catalytic oxidation of S(IV) is reasonable.

5. CONCLUSIONS (1) The Fe(II)-induced catalytic oxidation of S(IV) reaction order for O2 and HSO3 was found to be 0 and 1. For 106 M e [Fe(II)] e 104 M the reaction order toward Fe(II) was found to be 1. d½HSO3 ¼ k½FeðIIÞ½HSO3 dt For 104 M < [Fe(II)] e 102 M the reaction order toward Fe(II) was found to be 0. d½HSO3 ¼ k½HSO 3 dt For 102 M e [Fe(II)] e 1 M the reaction rate expression was found to be 0.2. d½HSO3 ¼ k½FeðIIÞ0:2 ½HSO3 dt (2) From 30 to 50 °C the activation energies were 16.38, 26.44, and 28.59 kJ/mol when the ionic strengths were at 0.71, 2.51, and 4.01 M, respectively. The activation energy was low but increased with an increase in the ionic strength. (3) The ionic strength has a negative influence on the reaction rate constant. When Fe(II) was 0.001 M, the relationship between these two factors is pffiffi log k1 ¼ 2:2104 0:3354 I þ 0:0159I The relevant coefficient R = 0.9918. When Fe(II) was 0.1 M, the relationship between these two factors is pffiffi log k ¼ 3:0752 þ 0:0847 I 0:0744I R = 0.9954. (4) We found that a simplified model for the Fe(II)-induced catalytic oxidation reaction with S(IV) using radical reaction mechanisms, as proposed by other researchers, combined with experimental conditions enables a better analysis and explanation of the experimental results.

(1) Siskos, P. A.; Peterson, N. C.; Huie, R. E. Kinetics of the Manganese(III)Sulfur(IV) Reaction in Aqueous Perchloric Acid Solutions. Inorg. Chem. 1984, 23, 1134. (2) Kraft, J.; van Eldik, R. Kinetics and Mechanism of the Iiron(III)Catalyzed Autoxidation of Sulfur(IV) Oxides in Aqueous Solution. 2. Decomposition of Transient Iron(III)Sulfur(IV) Complexes. Inorg. Chem. 1989, 28, 2306. (3) Brandt, C.; van Eldik, R. Transition Metal-Catalyzed Oxidation of Sulfur(IV) Oxides. Atmospheric-Relevant Processes and Mechanisms. Chem. Rev. 1995, 95, 119. (4) Zhang, W.; Singh, P.; Muir, D. Iron(II) Oxidation by SO2/O2 in Acidic Media: Part I. Kinetics and Mechanism. Hydrometallurgy 2000, 55, 229. (5) Berglund, J.; Fronaeus, S.; Elding, L. I. Kinetics and Mechanism for Manganese-Catalyzed Oxidation of Sulfur(IV) by Oxygen in Aqueous Solution. Inorg. Chem. 1993, 32, 4527. (6) Brandt, C.; Fabian, I.; van Eldik, R. Kinetics and Mechanism of the Iron(III)-Catalyzed Autoxidation of Sulfur(IV) Oxides in Aqueous Solution. Evidence for the Redox Cycling of Iron in the Presence of Oxygen and Modeling of the Overall Reaction Mechanism. Inorg. Chem. 1994, 33, 687. (7) Pasiuk-Bronikowska, W.; Bronikowski, T. The Rate Equation for SO2 Autoxidation in Aqueous MnSO4 Solutions Containing H2SO4. Chem. Eng. Sci. 1981, 36, 215. (8) Ziajka, J.; Beer, F.; Warneck, P. Iron-Catalysed Oxidation of Bisulphite Aqueous Solution: Evidence for a Free Radical Chain Mechanism. Atmos. Environ. 1994, 28, 2549. (9) Brandt, C.; van Eldik, R. Kinetics and Mechanism of the Iron(III)-Catalyzed Autoxidation of Sulfur(IV) Oxides in Aqueous Solution. The Influence of pH, Medium and Aging. Transition Met. Chem. (Dordrecht, Neth.) 1998, 23, 667. (10) Huss, A.; Lim, P. K.; Eckert, C. A. Oxidation of Aqueous Sulfur Dioxide. 2. High-Pressure Studies and Proposed Reaction Mechanisms. J. Phys. Chem. 1982, 86, 4229. (11) Connick, R. E.; Zhang, Y. X. Kinetics and Mechanism of the Oxidation of HSO3 by O2. 2. The Manganese(II)-Catalyzed Reaction. Inorg. Chem. 1996, 35, 4613. (12) Podkrajsek, B.; Grgic, I.; Tursic, J.; Bercic, G. Influence of Atmospheric Carboxylic Acids on Catalytic Oxidation of Sulfur(IV). J. Atmos. Chem. 2006, 54, 103. (13) Zuo, Y. G.; Zhan, J.; Wu, T. X. Effects of Monochromatic UVVisible Light and Sunlight on Fe(III)-Catalyzed Oxidation of Dissolved Sulfur Dioxide. J. Atmos. Chem. 2005, 50, 195. (14) Mudgal, P. K.; Sharm, A. K.; Mishra, C. D.; Bansal, S. P.; Gupta, K. S. Kinetics of Ammonia and Ammonium Ion Inhibition of the Atmospheric Oxidation of Aqueous Sulfur Dioxide by Oxygen. J. Atmos. Chem. 2009, 61, 31. (15) Mishra, G. C.; Srivastava, R. D. Homogeneous Kinetics of Potassium Sulfite Oxidation. Chem. Eng. Sci. 1976, 31, 969. (16) Chen, T. I.; Barron, C. H. Some Aspects of the Homogeneous Kinetics of Sulfite Oxidation. Ind. Eng. Chem. Fundam. 1972, 11, 466. (17) Ermakov, A. N.; Purmal, A. P. Catalysis of HSO3/SO32‑ Oxidation by Manganese Ions. Kinet. Katal. 2002, 43, 249. (18) Zhang, Y.; Zhou, J. T.; Wang, D.; Zhang, A. L.; Wang, Y. O. Research on Removal of Sulfur Dioxide from Flue Gas by FeSO4 Solution Catalytic Oxidation. J. Dalian Univ. Technol. 2004, 45, 11. (19) Jiang, J. H.; Li, Y. H.; Cai, W. M. Experimental and Mechanism Research of SO2 Removal by Cast Iron Scraps in a Magnetically Fixed Bed. J. Hazard. Mater. 2008, 153, 508. 1164


Industrial & Engineering Chemistry Research (20) Zhang, Q.; Gui, K. T.; Yao, G. H. Catalytic Oxidation of S(IV) in the Flue Gas Desulfurization Process Using Magnetically Fluidized Bed. J. Eng. Thermophys. 2008, 29, 987. (21) Butler, A. D.; Fan, M. H.; Brown, R. C.; van Leeuwen, J. H.; Sung, S. W.; Duff, B. Pilot-Scale Tests of Poly Ferric Sulfate Synthesized Using SO2 at Des Moines Water Works. Chem. Eng. Proc. 2005, 44, 413. (22) Zhang, Y.; Guo, S. Y.; Zhou, J. T.; Li, C. Y.; Wang, G. D. Flue Gas Desulfurization by FeSO4 Solutions and Coagulation Performance of the Polymeric Ferric Sulfate By-product. Chem. Eng. Proc. 2010, 49, 859. (23) General Administration of Quality Supervision; Inspection and Quarantine of the People’s Republic of China; Standardization Administration of the People’s Republic of China. Water Treatment Chemicals-Poly Ferric Sulfate (GB 14591-2006). National Standard of the People’s Republic of China; Standards Press of China: Beijing, 2006. (24) Wei, F. S. Standard Methods for Water and Wastewater Monitoring and Analysis, 3 rd ed.; China Environmental Science Press: Beijing, 1989. (25) Sato, T.; Goto, T.; Okabe, T.; Lawson, F. The Oxidation of Fe(II) Sulfate with Sulfur and Oxygen Mixtures. Bull. Chem. Soc. Jpn. 1984, 57, 2082. (26) Coughanowr, D. R.; Krause, F. E. The Reaction of SO2 and O2 in Aqueous Solutions of MnSO4. Ind. Eng. Chem. Fundam. 1965, 4, 61. (27) Grgic, I.; Hudnik, V.; Bizjak, M.; Levec, J. Aqueous S(IV) Oxidation—I. Catalytic Effects of Some Metal Ion. Atmos. Environ. 1991, 25A, 1591. (28) Zhang, H. T.; Peng, C. J.; He, J. Q.; Zhu, L. X.; Li, G. L. Study of Aqueous Phase SO2 Catalytic Oxidation in Test Solution. J. Wuhan Inst. Chem. Technol. 1993, 15, 55. (29) Kong, Q. C.; Shen, J. Kinetics of S(IV) Oxidation Catalyzed by Fe3+ and Mn2+ in Atmospheric Aqueous Phase. Acta Scientiae Circumstantiae 1997, 17, 289. (30) Mishra, G. C.; Srivastava, R. D. Kinetics of Oxidation of Ammonium Sulfite by Rapid-Mixing Method. Chem. Eng. Sci. 1975, 30, 1387. (31) Linek, V.; Vacek, V. Chemical Engineering Use of Catalyzed Sulfite Oxidation Kinetics for the Determination of Mass Transfer Characteristics of Gas-Liquid Contactors. Chem. Eng. Sci. 1981, 36, 1747. (32) Ibusuki, T.; Barnes, H. M. Manganese(II) Catalyzed Sulfur Dioxide Oxidation in Aqueous Solution at Environmental Concentrations. Atmos. Environ. 1984, 18, 145. (33) Lu, J. Y.; Cheng, T. Y.; Shen, Y. Study on the Catalytic Oxidation of Dynamics of SO32‑ in Aqueous Solution. Journal of Xinjiang Petroleum Institute 2000, 12, 35. (34) Barron, C. H.; O’Hern, H. A. Reaction Kinetics of Sodium Sulfite Oxidation by the Rapid-Mixing Method. Chem. Eng. Sci. 1966, 21, 397. (35) Bal Reddy, K.; van Eldik, R. Kinetics and Mechanism of the Sulfite-Induced Autoxidation of Fe(II) in Acidic Aqueous Solution. Atmos. Environ. 1992, 26A, 661. (36) Karatza, D.; Prisciandaro, M.; Lancia, A.; Musmarra, D. Calcium Bisulfite Oxidation in the Flue Gas Desulfurization Process Catalyzed by Iron and Manganese Ions. Ind. Eng. Chem. Res. 2004, 43, 4876. (37) Jin, J. J. Principles of Liquid Chemical Reaction Kinetics; Shanghai Scientific & Technical Publishers: Shanghai, 1984. (38) Mathews, C. T.; Robins, R. G. The Oxidation of Aqueous Ferrous Sulphate Solutions by Molecular Oxygen. Proceedings of the Australasian Institute of Mining and Metallurgy (ISSN: 0004-8364) 1972, 242, 47. (39) Ibusuki, T.; Takeuchi, K. Sulfur Dioxide Oxidation by Oxygen Catalyzed by Mixtures of Manganese(II) and Iron(III) in Aqueous Solutions at Environmental Reaction Conditions. Atmos. Environ. 1987, 21, 1555. (40) Copson, R. L.; Payne, J. W. Recovery of Sulfur Dioxide as Dilute Sulfuric Acid —Catalytic Oxidation in Water Solution. Ind. Eng. Chem. 1933, 25, 909. (41) Pasiuk-Bronikowska, W.; Sokolowski, A. Activation Energy Variation for Catalytic Oxidation of Aqueous SO2. Chem. Eng. Sci. 1983, 38, 121.

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(42) Lagrange, J.; Pallares, C.; Wenger, G.; Lagrange, P. Electrolyte Effects on Aqueous Atmospheric Oxidation of Sulphur Dioxide by Hydrogen Peroxide. Atmos. Environ. 1993, 27A, 129. (43) Huss, A.; Lim, P. K.; Eckert, C. A. Oxidation of Aqueous Sulfur Dioxide. 1. Homogeneous Manganese(II) and Iron(II) Catalysis at Low pH. J. Phys. Chem. 1982, 86, 4224. (44) Martin, L. R.; Hill, M. W. The Iron Catalyzed Oxidation of Sulfur: Reconciliation of the Literature Rates. Atmos. Environ. 1987, 21, 1487. (45) Millero, F. J.; Izaguirre, M. Effect of Ionic Strength and Ionic Interactions on the Oxidation of Fe(II). J. Solution Chem. 1989, 18, 585. (46) Gu, G. Y.; Han, S. J. Studies on Salt Effect to the Saponification Reaction Rate Constant of Ethyl Acetate. Acta Univ. Acta Univ. Med. Secondae Shanghai 1997, 17, 67. (47) Kraft, J.; van Eldik, R. The Possible Role of Iron(III)-Sulfur(IV) Complexes in the Catalyzed Autoxidation of Sulfur(IV)-Oxides. A Mechanistic Investigation. Atmos. Environ. 1989, 23, 2709.

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Reaction Kinetics and Mechanism of Iron(II) - American Chemical

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