NO2-Catalyzed Sulfite Oxidation - Industrial & Engineering Chemistry

Subscriber access provided by NEW YORK UNIV

Article

NO2 Catalyzed Sulfite Oxidation Virbin Nath A Sapkota, Nathan A Fine, and Gary T Rochelle Ind. Eng. Chem. Res., Just Accepted Manuscript • DOI: 10.1021/ie504767w • Publication Date (Web): 14 Apr 2015 Downloaded from http://pubs.acs.org on April 20, 2015

Just Accepted “Just Accepted” manuscripts have been peer-reviewed and accepted for publication. They are posted online prior to technical editing, formatting for publication and author proofing. The American Chemical Society provides “Just Accepted” as a free service to the research community to expedite the dissemination of scientific material as soon as possible after acceptance. “Just Accepted” manuscripts appear in full in PDF format accompanied by an HTML abstract. “Just Accepted” manuscripts have been fully peer reviewed, but should not be considered the official version of record. They are accessible to all readers and citable by the Digital Object Identifier (DOI®). “Just Accepted” is an optional service offered to authors. Therefore, the “Just Accepted” Web site may not include all articles that will be published in the journal. After a manuscript is technically edited and formatted, it will be removed from the “Just Accepted” Web site and published as an ASAP article. Note that technical editing may introduce minor changes to the manuscript text and/or graphics which could affect content, and all legal disclaimers and ethical guidelines that apply to the journal pertain. ACS cannot be held responsible for errors or consequences arising from the use of information contained in these “Just Accepted” manuscripts.

Industrial & Engineering Chemistry Research is published by the American Chemical Society. 1155 Sixteenth Street N.W., Washington, DC 20036 Published by American Chemical Society. Copyright © American Chemical Society. However, no copyright claim is made to original U.S. Government works, or works produced by employees of any Commonwealth realm Crown government in the course of their duties.

Page 1 of 27

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Industrial & Engineering Chemistry Research

1

NO2 Catalyzed Sulfite Oxidation

2

By Virbin Nath A. Sapkota, Nathan A. Fine, & Gary T. Rochelle*

3 4

The University of Texas at Austin, McKetta Department of Chemical Engineering, 200 E Dean Keeton St. Stop C0400, Austin, TX 78712-1589

5

* [email protected]

6 7

Keywords: NO2 absorption, sulfite oxidation, SO2 scrubbing

8

Abstract: Rates of sulfite (SO32-) oxidation during nitrogen dioxide (NO2) absorption into sulfite

9

solutions were measured using a high gas flow sparging apparatus. SO32- absorbs NO2 via a free

10

radical mechanism that produces sulfite radical (SO3 − ∙) and nitrite. The radical then catalyzes

11

SO32- oxidation to form sulfate. Free radical scavengers such as thiosulfate can be added to

12

suppress radical concentrations and inhibit sulfite oxidation. The effects of thiosulfate, sulfite,

13

NO2 absorption, temperature, and O2 partial pressure on sulfite oxidation were investigated under

14

NaOH scrubbing conditions. Oxidation is inverse half-order in thiosulfate, first order in sulfite,

15

and half order in NO2 absorbed.

16

pressures above 5 kPa but has a strong dependence on O2 at lower partial pressures. Oxidation

17

from 20 °C to 65 °C was fit using the Arrhenius equation with an activation energy of 24.1

18

kJ/mol. The addition of 0.01 mM Fe increased oxidation rates by a factor of 3 compared to

19

solutions with 0.01 mM EDTA added to chelate trace metals. A standard NaOH scrubber would

Sulfite oxidation shows little dependence on O2 partial

1

ACS Paragon Plus Environment


1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 2 of 27

20

need approximately 50 mM thiosulfate in the circulating solvent to maintain 10 mM SO32- while

21

removing 100 ppm SO2 and 2.5 ppm NO2.

22

Nitrosamines in Amine Scrubbing

23

Amine scrubbing is a mature and industrially proven technology for post-combustion carbon

24

capture. However, amine solvents can react with NO2 to form carcinogenic nitrosamines. While

25

these nitrosamines are unstable compared to the parent amine, they will still accumulate in the

26

scrubber to a steady state concentration. Assuming amine oxidation will be kept to a minimum,

27

the maximum concentration of accumulated nitrosamines in an amine scrubber is proportional to

28

the inlet NO2 concentration and inversely proportional to its decomposition rate in the desorber

29

(Equation 1).1 𝑁𝑁𝑂𝑆𝑐𝑟 =

𝑦𝑁𝑂2 𝑖𝑛𝑙𝑒𝑡 𝐺 ∗ 𝑘𝐷𝑒𝑐𝑜𝑚𝑝 𝜏𝐷𝑒𝑠𝑜𝑟𝑏𝑒𝑟 𝐿 𝑆𝑐𝑟

(1)

30

Nitrosamines are a possible safety risk in the case of accidental spills and a possible

31

environmental hazard from flue gas emissions or reclaiming waste streams. One possible way to

32

limit nitrosamine formation is to remove the NO2 from the flue gas before it reaches the amine

33

scrubber.

34

Sulfite Oxidation in Limestone and Sodium Hydroxide Scrubbing

35

Limestone slurry scrubbing is a common method for flue gas desulfurization. The process also

36

removes some NO2, although it does not remove NO because of the limited solubility and

37

reactivity of NO in aqueous SO32-. In the presence of gas phase O2, the absorption of NO2 is

38

accepted as follows: 𝑁𝑂2 + 𝑆𝑂3 2− → 𝑁𝑂2 − + 𝑆𝑂3 − ∙

(2)

2


Page 3 of 27

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


𝑆𝑂3 − ∙ +𝑂2 → 𝑆𝑂5 − ∙

(3)

𝑆𝑂5 − ∙ +𝑆𝑂3 2− → 𝑆𝑂4 − ∙ +𝑆𝑂4 2−

(4)

𝑆𝑂4 − ∙ +𝑆𝑂3 2− → 𝑆𝑂3 − ∙ +𝑆𝑂4 2−

(5)

2𝑆𝑂3 − ∙→ 𝑆2 𝑂6 2−

(6)

39

Equation 2 was first outlined by Nash2 and Equations 3–5 were proposed by Huie and Neta.3

40

These occur in the liquid mass transfer boundary layer. They imply that for every mole of NO2

41

absorbed, several moles of SO32- can be consumed due to free-radical propagation. Because NO2

42

absorption is first order in SO32-, NO2 absorption is strongly affected by sulfite oxidation.4 In a

43

limestone slurry scrubber, SO2 gas is absorbed into solution as SO32-, which balances the SO32-

44

loss from oxidation and leads to a steady state SO32- concentration.

45

Several factors affect the oxidation rate. Free radical scavengers, such as thiosulfate, can

46

provide an alternative termination step to the free radical process, drastically reducing the

47

amount of sulfite oxidized:4 𝑁𝑂2 + 𝑆𝑂3 2− → 𝑁𝑂2 − + 𝑆𝑂3 − ∙

(2)

𝑆𝑂3 − ∙ +𝑂2 → 𝑆𝑂5 − ∙

(3)

𝑆𝑂5 − ∙ +𝑆2 𝑂3 2− → 𝑆𝑂5 2− + 𝑆2 𝑂3 − ∙

(7)

𝑆𝑂5 2− + 𝑆𝑂3 2− → 2𝑆𝑂4 2−

(8)

2𝑆2 𝑂3 − ∙→ 𝑆4 𝑂6 2−

(9)

48

Takeuchi5 studied the effects of other antioxidants, such as hydroquinone, phenol,

49

ethanolamines, ethylene glycol monoethyl ether, glycine, ethylenediaminetetraacetic acid

50

(EDTA), and acetic acid. All of these antioxidants slowed sulfite oxidation.5 Changes in other

51

process variables such as NO2 flow, O2 flow, initial sulfite concentration, metals, and

52

temperature could affect the sulfite oxidation rate. Shen studied the absorption of NO2 into 3



1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 4 of 27

53

sulfite and thiosulfate inhibited solutions, investigated the effects of adding Fe2+, chloride, and

54

EDTA, and quantified the effects of sulfite, oxygen, and gas-phase SO2.6 However, all previous

55

experimental work has focused on limestone slurry scrubbing conditions with relatively high

56

inlet SO2 partial pressure, at pH 4–6, and a high feed of metals from both the fly ash and

57

limestone. This experimental effort addresses NaOH scrubbing with lower SO2 partial pressure,

58

at pH 7–10, and with lower dissolved metals.

59

Methods

60

All of the experiments were run in the high gas flow apparatus (HGF) previously described

61

(Figure 1).7

62 63

Figure 1. High gas flow apparatus 4


Page 5 of 27

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


64

The mixture of the sparging gas was controlled by mass flow controllers. Dry air exited the

65

mass flow controller and was passed through a temperature controlled water saturator to

66

maintain water balance in the HGF. CO2 and dilute NO2 were then directly mixed with the

67

hydrated air stream. The resulting gas stream was either sparged directly into the bottom of the

68

HGF or put through a bypass.

69

Before solutions were loaded into the HGF, it was triply rinsed with deionized distilled water

70

to remove any contaminants. The temperatures of the water saturator and HGF water baths were

71

adjusted and allowed to reach a steady temperature. To evaluate the amount of absorbed NO2,

72

the exit gas streams were diluted by a dry air stream and fed into a Thermo Scientific Model 42i

73

(NO-NO2-NOx) Analyzer chemiluminescent trace level NOx analyzer.8

74

The absorbing solution (approximately 360 g) included aqueous sulfite and thiosulfate in a

75

0.5 M NaHCO3 buffer with EDTA added to chelate any trace metals. The initial solution was

76

adjusted to pH 9.2 with NaOH. During startup, the gas stream was set to HGF bypass mode

77

while pouring the solution into the HGF. The gas stream was then sparged into the HGF for a

78

short period in order to coat the walls of the HGF and ensure the solution was well mixed. The

79

gas was set back to bypass mode and 1 ml of initial sample was taken from a sample port at the

80

bottom of the HGF. The sample was immediately injected into 0.1 g of 35 wt % formaldehyde

81

to react any free sulfite remaining to methylsulfonic acid (MSA). 𝐻𝐶𝑂𝐻 + 𝑆𝑂3 2− + 𝐻2 𝑂 → 𝐻𝐶𝐻𝑂𝐻𝑆𝑂3 − + 𝑂𝐻 −

(10)

82

MSA is oxidatively stable at room temperature, allowing indirect sulfite analysis using anion

83

chromatography (Supporting Information A).

84

Once the sample was taken, NO2 concentration values from the trace level NOx analyzer were

85

recorded while in bypass mode. The gas was set back to sparge into the HGF, and NO2 values 5



1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 6 of 27

86

were recorded as soon as they stabilized at a new value. At predetermined intervals, the liquid

87

sampling procedure was repeated to get sulfite oxidation kinetics in the semi-batch process. A

88

50x dilution with water was performed on each sample generated, and the diluted samples were

89

analyzed on the anion chromatograph.

90

Safety

91

All of the chemicals used in this experiment were safe to handle with standard laboratory

92

practices with the notable exception of formaldehyde.

Work with an open formaldehyde

93

container was always done under a fume hood with the exit gas vented to the atmosphere. The

94

sample vials were closed with slitted caps such that the formaldehyde could not escape, yet

95

sample could be injected into the formaldehyde solution. When the formaldehyde container was

96

not in use, it was always kept in a flammables cabinet.

97

Nitrogen oxides were the main hazardous gases. Stainless steel tubing was used with gases

98

containing nitrogen oxides. All joints were checked for leaks before usage. Lines containing

99

pressurized nitrogen oxides were depressurized nightly. The tanks containing nitrogen oxides

100

are stored in a vented cabinet next to low-level NOx sensors.

101

Results and Discussion

102

Table 2 gives overall results for the entire experimental set. A gas mixture of hydrated N2, Air,

103

CO2, and NO2 gases was sparged through aqueous solutions containing compositions of sodium

104

sulfite (NaSO3), sodium thiosulfate (NaS2O3), sodium bicarbonate (NaHCO3), and EDTA in the

105

High Gas Flow apparatus (HGF). The sulfite and sulfate concentrations were measured at

106

various times in the semi-batch process using anion chromatography. The decrease in the sulfite

107

concentration over time was attributed to sulfite oxidation, and the time series was used to

6


Page 7 of 27

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


108

regress sulfite oxidation rate constants. The standard error for each rate constant was on average

109

4% of the regressed constant.

110

Sulfite and sulfate concentrations were also used to check for mass balance closure in sulfite

111

and sulfate. Almost all sulfite should oxidize to sulfate (Equation 4), so total S (sulfite + sulfate)

112

should remain constant throughout the run. In Experiment 14, the sum of sulfite and sulfate held

113

at an average 43.3 mmol/kg with a standard deviation of 0.6 mmol/kg (Table 1). With the

114

exception of the first sample, there was a small but clear loss of total S over the run, possibly due

115

to dithionate formation (Equation 6). All experiments showed near constant total S as sulfite

116

oxidized, proving that sulfate is the dominant product of sulfite oxidation.

117

Table 1. Experiment 14 total S balance Time (min:sec)

Sulfite (mM)

Sulfate (mM)

Total S (mM)

0

41.56

1.10

42.66

5:00

40.67

3.94

44.61

13:00

37.82

5.74

43.57

20:50

35.68

7.72

43.40

28:00

33.60

9.63

43.22

37:40

31.91

10.99

42.90

45:35

30.07

12.73

42.80

Average

43.31

Standard Deviation

0.61

118

NO2 flux was initially determined by using a chemiluminescent trace level NOx analyzer to

119

measure the difference between the concentration of NO2 when the gas passed through the HGF

120

and when it bypassed the HGF. Assuming no leaks, this difference in NO2 concentration was the

121

amount of NO2 absorbed into solution.

However, this method was not reliable at higher 7



1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 8 of 27

122

temperature, when the sparging gas was over-saturated at room temperature. Having significant

123

amounts of water vapor present in the exit stream presented a risk to the analyzer used in the

124

experiment, which typically operates under ambient, unsaturated conditions.

125

empirical model regressed from NO2 flux measurements at 20 °C was used to estimate NO2

126

absorbed: 𝑁𝑂2 𝑎𝑏𝑠𝑜𝑟𝑏𝑒𝑑[=]

𝑚𝑜𝑙 𝑘𝑔𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛 ∙ 𝑚𝑖𝑛

=

(1 − 𝑒 −𝑁𝑜𝑔 )[𝑁𝑂2 ]𝐵𝑦𝑝𝑎𝑠𝑠 𝑉̇ 1 𝑚𝑜𝑙 ∗[ ] 𝑚𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛 22.4 𝐿

Instead, an

(11)

2− ]

[𝑆𝑂 𝐿 𝑎 𝑉̇ = 7.5 𝑚𝑖𝑛 , 𝑚𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛 = 0.36 𝑘𝑔 , 𝑁𝑜𝑔 = 𝑘𝑔 ′ ∗ 𝐺 𝑒 = 7.37√ 1 3𝑀 𝐻𝐺𝐹

127

This model assumes that the HGF operates in a semi-batch mode and that NO2 flux is first

128

order in NO2 partial pressure.

The mass-transfer kinetics for determining the number of

129

theoretical transfer units (Nog) is assumed to be controlled by mass transfer with fast reaction in

130

the liquid boundary layer with minimal impact from the diffusion of reactants and products.9

131

The model fit all the NO2 flux data measured at 20 °C for inhibited solutions. However, NO2

132

flux into uninhibited solutions was noticeably slower due to significant sulfite depletion at the

133

gas-liquid interface, so NO2 flux measured by the NOx analyzer was used in place of the model.

134

Similar slow rates of aerobic NO2 absorption into uninhibited sulfite solutions have been

135

reported in the literature.4,5 The amount of SO32- oxidized per mole of NO2 absorbed (f) was

136

calculated at a normalized SO32- concentration of 0.040 mol/kg. High values imply that a single

137

NO2 molecule catalyzes large amounts of sulfite oxidation. In Experiments 6–8 when the effect

138

of SO32- was examined, the f value is shown for the initial concentration of SO32- instead of the

139

normalized 0.040 mol SO32-/kg. The calculated f was determined using Equation 11 and a

140

simple least-squares regression with the prefactor and activation energy as free parameters; the

141

effect of iron and O2 partial pressure were not regressed. 8


Page 9 of 27

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


5 ∗ 10−6 𝑓 = 𝑓𝑜 √ ∗ (1 − 𝑒−𝑁𝑜𝑔 )𝑦𝑁𝑂 2

𝑓𝑜 [=]

[𝑆𝑂3 2− ]

∗ exp [

√[𝑆2 𝑂3 2− ] ∗ √1.0 M

𝑚𝑜𝑙 𝑆𝑂3 2− 𝑜𝑥 𝑚𝑜𝑙 𝑁𝑂2 𝐴𝑏𝑠

= 233

−𝐸𝑎 1 1 ( − )] 𝑅 293 𝐾 𝑇

(12)

𝐸𝑎 = 24.1 𝑘𝐽/𝑚𝑜𝑙

142

Many of the various species were changed independently during separate experiments, but in

143

all of the following: the total gas flow rate was kept at 7.5 SLPM; the volume of solution

144

introduced to the HGF was constant; 0.5 M NaHCO3 was added to maintain a buffered solution;

145

NaOH was added to buffer to pH = 9.2; and 0.02 mM EDTA was added to chelate any metal

146

impurities. Additionally, a base case was picked to compare all of the variations in the different

147

independent conditions. The base case solution started with 0.040 mol/kg Na2SO3 and 0.025

148

mol/kg Na2S2O3. The temperature of the HGF was held at 20 °C, and the gas contained 5 ppm

149

NO2 in hydrated air diluent.

9



1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48

Page 10 of 27

Table 2. NO2 catalyzed sulfite oxidation

T (°C)

Oxygen (kPa)

NO2 (ppm)

k1 obs*103 (1/min)

k0 obs*103 (mol/kg/min)

NO2 Flux*106 (mol/kg/min)

25

20

21

5

6.5±0.3

-

3.6

73.4

67.2

44.3

0

20

21

2

70.1

97.5

1.2

1751.7

-

3*

39.2

0

20

21

5

110.8

154.1

3.0

1107.9

-

4*

33.8

0

20

21

10

156.7

218.0

6.0

783.4

-

5

40.0

100

20

21

5

3.4±0.4

-

3.6

37.9

33.6

6

7.2

25

20

21

5

4.0±0.1

-

2.1

13.1

15.6

7

67.0

25

20

21

5

7.0±0.2

-

3.9

119.4

107.0

8

135.6

25

20

21

5

6.4±0.3

-

4.3

201.7

210.9

9

45.0

25

20

21

2

3.7±0.1

-

1.4

102.9

106.2

10*+

37.2

25

20

21

5

18.4±0.2

-

3.6

207.0

67.2

11

45.2

25

35

21

5

10.3±0.6

-

3.6

115.9

108.9

12

7.4

25

53

21

5

15.0±0.7

-

2.2

175.5

183.3

13

33.9

25

64

21

5

22.1±0.7

-

3.6

248.6

245.3

14*

41.6

25

20

10

5

5.9±0.1

-

3.6

66.8

67.2

15*

48.2

25

20

5

5

5.6±0.1

-

3.6

63.0

67.2

16*

49.1

25

20

2

5

3.9±0.2

-

3.6

44.3

67.2

Exp. #

Sulfite (mM)

Thiosulfate (mM)

1

40.2

2*

𝑺𝑶 𝟐− 𝒐𝒙

fobs (𝑵𝑶𝟑

𝟐 𝑭𝒍𝒖𝒙

)

fcalc

10


Page 11 of 27

132

*Experiments 2–4, 10, and 14–16 were not used to regress fCalc.

133

+

134

For Experiments 1 and 5–16, the rate law was assumed

135

regressed using ln[𝑆𝑂3 2− ] = ln[𝑆𝑂3 2− ]𝑖 − 𝑘1,obs 𝑡

136

For Experiments 2–4, the rate law was assumed

137

using a least squares regression on the numerical solution of the ODE.

138

Uninhibited Aerobic NO2 Absorption with Varying NO2 Flow

139

The rates for uninhibited sulfite oxidation were quantified in solutions with no added

140

thiosulfate. The dilute NO2 flow rate was varied, but sulfite, temperature, and pH were kept

141

constant at base case conditions (Figure 2).

In Experiment 10, 0.1 mM Fe was added instead of 0.02 mM EDTA.

𝑑[𝑆𝑂3 2− ] 𝑑𝑥

𝑑[𝑆𝑂3 2− ] 𝑑𝑥

= 𝑘1,obs [𝑆𝑂3 2− ] and linearly

= 𝑘1,obs [𝑆𝑂3 2− ] + 𝑘0,obs and regressed

50

Sulfite (mmol/kg)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


2 ppm NO2 10 ppm NO2

5

5 ppm NO2

0.5 0

10

20

30

40

Time (min) 142 143

Figure 2: Uninhibited SO32- oxidation with varying NO2 flow to adjust concentration: 0.5 M

144

NaHCO3; pH = 9.2; 20 °C; 21 kPa O2 (Experiments 2–4)


11


1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

145 146

The sulfite concentrations were fit with the following rate law equation, graphed with solid lines above: 𝑑[𝑆𝑂3 2− ] = 𝑘1 obs [𝑆𝑂3 2− ] + 𝑘0 obs 𝑑𝑡

147 148

Page 12 of 27

(13)

The zeroth order term dominates at low concentrations of sulfite. The dashed line models pseudo-first-order oxidation: 𝑑[𝑆𝑂3 2− ] = 𝑘1 obs [𝑆𝑂3 2− ] 𝑑𝑡

(14)

149

The observed oxidation only deviates significantly from the pseudo-first-order rate when the

150

concentration of sulfite is below 5 mmol/kg. Thus, k0 obs is only significant below 5 mmol/kg,

151

and otherwise, oxidation is pseudo-first order in sulfite.

152

This also verifies the half-order relationship of NO2 on sulfite oxidation. In all runs, increases

153

in NO2 result in half that proportional increase in k1 obs and k0 obs. Additionally, increases in NO2

154

have a half-order effect on reducing the ratio of SO32- oxidized to NO2 absorbed (f). From the 2

155

to the 5 ppm cases, f dropped by a factor of 1.56, and from the 2 to 10 ppm cases, f dropped by a

156

factor of 2.22. However, the large f values imply very low sulfite concentrations when oxidation

157

is uninhibited, making these systems poor absorbers of NO2 unless the inlet SO2 is very high.

158

Inhibited Aerobic NO2 Absorption with Varying Thiosulfate Concentrations

159

The rates for inhibited sulfite oxidation were quantified in solutions with thiosulfate added at

160

0 mM, 25 mM, and 100 mM. Sulfite, temperature, pH, gas composition and gas flow were kept

161

constant at base case conditions. Data for the 0 mM run came from the 5 ppm uninhibited

162

experiment and data for the 25 mM run were treated as the base case for all experiments

163

(Figure 3).


12

Page 13 of 27

100 mM S2O32-

40

Sulfite (mmol/kg)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


30

25 mM S2O320 mM S2O32-

20

10

0 0

10

20

30

40

Time (min)

164 165

Figure 3: Sulfite oxidation with varying S2O32- concentration: 0.5 M NaHCO3, pH = 9.2, 20 °C,

166

21 kPa O2, 5 ppm NO2 (Experiments 1, 3, 5)

167

SO32- for the 25 mM S2O32- and 100 mM S2O32- runs was fit with Equation 13 because the

168

sulfite remains much higher than 5 mmol/kg. The results without S2O32- were modeled with

169

Equation 12. Thiosulfate inhibited sulfite oxidation by an order of magnitude. Increasing

170

thiosulfate from 25 mM to 100 mM decreased both sulfite oxidation and f by approximately a

171

factor of 2. This half-order correlation in oxidation rate is corroborated in previous work.4

172

Because sulfite oxidation inhibited by thiosulfate is an order of magnitude slower than

173

uninhibited oxidation, sulfite will accumulate to a higher concentration in the SO2 scrubber. A


13


174

higher sulfite concentration will lead to high levels of simultaneous absorption of SO 2 and NO2

175

in the scrubber.

176

Inhibited Aerobic NO2 Absorption with Varying Sulfite Concentrations

177

Sulfite oxidation rates were quantified with 7 to 135 mM sulfite. Thiosulfate, temperature, pH,

178

gas composition, and gas flow were kept constant at base case conditions. Data for the 40 mM

179

sulfite experiment came from the base case. The results for this group are summarized in

180

Figures 4 & 5:

100

Sulfite (mM)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 14 of 27

65 mM SO32-

135 mM SO32-

40 mM SO3210

7 mM SO32-

1 0

10

20

30

40

Time (min) 181 182

Figure 4. SO32- oxidation with varying SO32- concentration: 0.5 M NaHCO3, pH = 9.2, 20 °C,

183

21 kPa O2, 5 ppm NO2, 25 mM S2O32- (Experiments 1, 6–8)


14

Page 15 of 27

0.008

150 0.006 100 0.005 50

0.004

0.003

0 0

184

f = SO32- oxidized/NO2 absorbed

200

0.007

k1 obs (min-1)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


50

100

150

Sulfite (mM)

185

Figure 5. SO32- oxidation with varying SO32- concentration: 0.5 M NaHCO3, pH = 9.2, 20 °C,

186

21 kPa O2, 5 ppm NO2, 25 mM S2O32- (Experiments 1, 6–8)

187

The sulfite concentrations were fitted with Equation 13. Because Equation 13 assumes sulfite

188

oxidation is pseudo-first order in sulfite, k1 obs should be the same for all runs. However, the

189

k1 obs for the 7 mM SO32- case is significantly lower than the rest. This is due to low NO2

190

absorption in the low initial SO32- concentration. Once normalized for NO2 absorption, the ratio

191

of the amount of SO32- oxidized to NO2 absorbed is first order in sulfite.

192

Inhibited Aerobic NO2 Absorption with Varying NO2 Flow


15


193

Sulfite oxidation rates were quantified with NO2 concentrations of 2 ppm and 5 ppm (base

194

case). Thiosulfate, sulfite, temperature, pH, and total gas flow were kept constant at base case

195

conditions (Figure 6). 45

Sulfite (mmol/kg)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 16 of 27

2 ppm NO2 40

35

5 ppm NO2

30 0 196

10

20

30

40

50

60

Time (min)

197

Figure 6. Inhibited SO32- oxidation with varying NO2 concentration: 0.5 M NaHCO3, pH = 9.2,

198

20 °C, 21 kPa O2, 25 mM S2O32- (Experiments 1, 9)

199

Sulfite was fit by Equation 13. Between the 2 ppm and 5 ppm cases, the k1 obs increased by a

200

factor of 1.78 while NO2 increased by a factor of 2.5, indicating half-order behavior. This half-

201

order correlation is corroborated in previous work.4 However, f decreased by a factor of 1.63

202

when NO2 increased by a factor of 2.5, suggesting an inverse half-order relationship between

203

NO2 absorbed and f. Shen showed that NO2 absorption is first order in NO2, but that SO32-


16

Page 17 of 27

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


204

oxidation is half-order in NO2. Thus, increases in NO2 partial pressures increase NO2 absorption

205

faster than they increase SO32- oxidation. This has important implications for NO2 absorbing

206

applications; higher partial pressures of NO2 are more efficiently removed with less sulfite loss

207

per unit NO2 absorbed even though the absolute rate of sulfite oxidation increases.

208

Inhibited Aerobic NO2 Absorption with Added Metals

209

Work done by Ulrich et al. indicates that Fe2+ is a powerful oxidation catalyst that is effective

210

as low as 3 ∗ 10−3 mM.10 Sulfite oxidation rates were quantified with 0.01 mM Fe2+ and

211

compared to the base case, which had 0.02 mM EDTA to chelate any background Fe2+.

212

Thiosulfate, sulfite, temperature, pH, and gas flow and composition were kept constant at base

213

case conditions. The addition of 0.01 mM Fe2+ increased k1

214

contradicts earlier conclusions that metals did not have a significant effect on NO2-catalyzed

215

sulfite oxidation.4 Shen reported that metals such as Fe2+ are insignificant catalysts of sulfite

216

oxidation in the presence of NO2, however, his experiments were conducted with 200–1000 ppm

217

NO2, which catalyzes far more sulfite oxidation than added metals do. The addition of iron,

218

therefore, did not cause a significant increase in the oxidation rate in his experiments.

obs

by a factor of 2.83, which

219

In contrast, work done by Ulrich indicates that Fe2+ is a powerful catalyst effective at

220

background concentrations. However, those experiments were run without NO2 absorption. The

221

current results show the importance of adding chelating agents to the NaOH scrubber to limit

222

sulfite oxidation when the flue gas contains less than 10 ppm NO2.

223

economically viable in NaOH scrubbing since most of the fly ash, a large source of metal ions,

224

will effectively be captured upstream of the scrubber. Furthermore, the NaOH feed is expected

225

to have much lower levels of dissolved metal compared to the limestone feed for traditional flue

226

gas desulfurization.

Chelating Fe2+ is


17


227

Inhibited Aerobic NO2 Absorption with Varying Temperature

228

Sulfite oxidation rates were quantified at 20 °C to 65 °C (Figure 7). Thiosulfate, sulfite, pH,

229

and gas flow and composition were kept constant at base case conditions.

0.024

64 °C

k1 obs (min-1)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 18 of 27

53 °C 0.012

35 °C

20 °C 0.006 230

Inverse Temperature (K-1)

231

Figure 7. SO32- oxidation with varying temperature: 0.5 M NaHCO3, pH = 9.2, 21 kPa O2, 5 ppm

232

NO2, 25 mM S2O32-, 40 mM initial SO32- (Experiments 11–13)

233

The kg’ for NO2 absorption has almost no temperature dependence in this temperature range,4

234

so NO2 flux was assumed to be constant. Since all other dependent variables for f were held

235

constant, an apparent activation energy of 24.1 kJ/mol was regressed using Equation 15. 𝑓 = 𝐴 ∗ 𝑒𝑥𝑝 [

236

−𝐸𝑎 1 1 ( − )] 𝑅 293𝐾 𝑇

(15)

Inhibited Aerobic NO2 Absorption with Varying Oxygen Flow


18

Page 19 of 27

237

Sulfite oxidation rates were quantified with at 2 to 21% oxygen.

Thiosulfate, sulfite,

238

temperature, pH, and total gas flow were kept constant at base case conditions. Total gas flow

239

was kept constant by blending air and N2 to the desired O2 concentration (Figure 8).

0.008

0.006

kobs (min-1)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


0.004

0.002

Maximum Anaerobic Sulfite Loss

0.000 0

0.05

0.1

0.15

0.2

0.25

Oxygen (bar)

240 241

Figure 8. SO32- oxidation with varying oxygen: 0.5 M NaHCO3, pH = 9.2, 5 ppm NO2, 25 mM

242

S2O32-, 40 mM initial SO32- (Experiments 1, 14–16)

243

In disagreement with the strong dependence on O2 partial pressure that Shen reported,4 O2 does

244

not have a significant effect on k1 obs until O2 partial pressure decreases below 5 kPa. This is due

245

to the free-radical species in the rate-limiting step of the propagation.

246

pressure, SO3 − ∙ oxidizes instantaneously in the boundary layer to form SO5 − ∙ (Equation 3). The

247

SO5 − ∙ then catalyzes oxidation until thiosulfate terminates the mechanism. At low O2 partial

248

pressure, the rate-limiting step occurs when SO3 − ∙ reacts with dissolved oxygen. Since SO3 − ∙ is

249

present in the bulk solution, thiosulfate can directly react with SO3 − ∙ .

250

concentrations of thiosulfate and dissolved oxygen then become determining factors in the

At high O2 partial

The relative


19


1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 20 of 27

251

overall oxidation rate. The O2 partial pressure in power generation applications will not affect

252

sulfite oxidation since coal flue gas usually contains over 5% O2 due to the excess air used

253

during combustion. The stronger dependence on O2 partial pressure that Shen reported was most

254

likely due to the lack of thiosulfate and the relatively high NO2 partial pressures (3.5–8 Pa) used

255

in this set of oxidation experiments, which depleted sulfite concentration at the gas-liquid

256

interface.

257

The maximum anaerobic sulfite loss solely from the NO2 absorption rate was calculated using

258

the stoichiometry for NO2 absorption as nitrite followed by the reaction of nitrite with sulfite to

259

form hydroxylamine disulfonate (HADS) (Equations 2 & 16–18). Neither nitrite nor HADS

260

accumulated to quantifiable concentrations during these experiments, but they are expected

261

major products from anaerobic absorption of NO2 into sulfite.4,11,12 Even at 2 kPa oxygen, sulfite

262

loss is dominated by NO2-catalyzed oxidation with less than 10% of sulfite loss attributable to

263

NO2/NO2- reactions. 𝑁𝑂2 + 𝑆𝑂3 2− → 𝑁𝑂2 − + 𝑆𝑂3 − ∙

(2)

1 (2𝑆𝑂3 − ∙→ 𝑆2 𝑂6 2− ) 2

(16)

3𝐻 + +𝑁𝑂2 − + 2𝑆𝑂3 2− → 𝐻𝑂𝑁(𝑆𝑂3 )2 3𝐻 + + 𝑁𝑂2 + 2𝑆𝑂3 2− → 𝐻𝑂𝑁(𝑆𝑂3 )2 264

2−

2−

+ 𝐻2 𝑂

(17)

+ 0.5𝑆2 𝑂6 2− + 𝐻2 𝑂

(18)

Industrial Application for Oxidation Results

265

The f value can be used to determine the thiosulfate makeup rate necessary to maintain a given

266

sulfite concentration in the NaOH scrubber. The NaOH polishing scrubber can be modeled as a

267

semi-batch reactor with the flue gas countercurrently contacting the circulating solvent (Figure

268

9). NaOH is fed to the system to maintain a basic pH while NaS2O3 can be fed to inhibit sulfite

269

oxidation. The unpolished flue gas contains approximately 30–300 ppm SO2 and 0.5–5 ppm


20

Page 21 of 27

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


270

NO2 depending on the coal type, burner technology, and upstream treatment. Almost all of the

271

SO2 will be absorbed in the scrubber, while NO2 absorption depends on the sulfite concentration

272

in the solvent.

273

approximately 2 M solid solubility limit. Inside the polishing scrubber SO2 reacts with OH- to

274

form SO32-, NO2 reacts with SO32- to form SO3 − ∙, and SO3 − ∙ catalyzes SO32- oxidation until

275

terminated by reacting with thiosulfate or another radical (Equations 2–9).

The circulating solvent is bled to maintain sulfate concentration below its

12 % 𝐶𝑂2 𝑦𝑁𝑂2 𝑓 = 0.1𝑦𝑁𝑂2 𝑖 𝑦𝑆𝑂2 𝑓 ≈ 0

𝑆𝑂2 + 2𝑂𝐻− → 𝑆𝑂3 2− + 𝐻2 𝑂

𝑁𝑂2 + 𝑆𝑂3 2− → 𝑁𝑂2 − + 𝑆𝑂3 − ∙ 1 𝑆𝑂3 − ∙ +𝑓𝑆𝑂3 2− + 𝑓𝑂2 → 𝑓𝑆𝑂4 2− + 𝑆𝑂3 − ∙ 2 𝑆𝑂3 − ∙ +𝑆2 𝑂3 /𝑟𝑎𝑑𝑖𝑐𝑎𝑙 → 𝑡𝑒𝑟𝑚𝑖𝑛𝑎𝑡𝑖𝑜𝑛

Feed 𝑛̇ 𝑁𝑎𝑂𝐻 [=]𝑚𝑜𝑙/𝑠 𝑛̇ 𝑁𝑎2𝑆2 𝑂3

Bleed [𝑆𝑂4 2− ] ≈ 1.5 𝑀 [𝑆𝑂3 2− ]

12 % 𝐶𝑂2 𝑦𝑁𝑂2 𝑖 𝑦𝑆𝑂2 𝑖 G [=]𝑚𝑜𝑙/𝑠

[𝑆2 𝑂3 2− ] 𝐿[=]𝐿𝑖𝑡𝑒𝑟𝑠/𝑠

276 277 278

Figure 9. Schematic for simultaneous NO2 and SO2 absorption in a NaOH scrubber using

279

NaS2O3 inhibition.


21


1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 22 of 27

280

The mass balances for sulfite, sulfate, and thiosulfate accumulation can be solved assuming

281

that sulfite loss is dominated by NO2-catalyzed oxidation and the NaOH scrubber is large enough

282

to absorb 90 % of the inlet NO2 and practically all of the SO2 (Equations 19–21). 𝑑[𝑆𝑂4 2− ] = 𝐺𝑦𝑆𝑂2 𝑖 − 𝐿[𝑆𝑂4 2− ] 𝑑𝑡

(19)

𝑑[𝑆𝑂3 2− ] = 𝐺𝑦𝑆𝑂2 𝑖 − 𝑓 ∗ 0.9𝐺𝑦𝑁𝑂2 𝑖 𝑑𝑡

(20)

𝑑[𝑆2 𝑂3 2− ] = 𝑛̇ 𝑁𝑎2 𝑆2 𝑂3 − 𝐿[𝑆2 𝑂3 2− ] 𝑑𝑡

283

(21)

284

At steady state and a conservative sulfate concentration of 1.5 M in the bleed, the f value from

285

Equation 12 and the molar feed rate for Na2S2O3 can be written in terms of inlet gas conditions

286

(Equations 22 & 23). 𝑦𝑆𝑂2𝑖 5 ∗ 10−6 𝑓= = 233√ ∗ 0.9𝑦𝑁𝑂2 𝑖 0.9𝑦𝑁𝑂 𝑖 2

[𝑆𝑂3 2− ] √[𝑆2 𝑂3 2− ] ∗ √1.0 M

𝑛̇ 𝑁𝑎2𝑆2𝑂3 = 𝐺𝑦𝑆𝑂2 𝑖 ∗ 287

∗ exp [

[𝑆2 𝑂3 2− ]

−24.1 1 1 ( − )] 𝑅 293 𝐾 𝑇 (22) (23)

1.5 𝑀

288

For an inlet concentration of 100 ppm SO2 and 2.5 ppm NO2 with a NaOH scrubber operating

289

at 55 °C, 0.050 M thiosulfate would be needed to keep SO32- concentration at 0.010 M in the

290

solvent (Equation 24). Assuming 90% CO2 capture from the flue gas, this corresponds to a feed

291

of 3.1*10-5 mole Na2S2O3 per mole of CO2 captured (Equation 25). The economics of the

292

process are a balance between the cost of Na2S2O3 feedstock and the size of the polishing

293

scrubber.


22

Page 23 of 27

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


𝑓 = 44.4 = 994 ∗

[𝑆𝑂3 2− ] √[𝑆2 𝑂3 2− ]

[𝑆2 𝑂3 ] = (

994 ∗ 0.01 M 2 ) = 0.050 𝑀 44.4

𝑛̇ 𝐶𝑂2 𝑐𝑎𝑝𝑡𝑢𝑟𝑒𝑑 = 0.9 ∗ 0.12 ∗ G 𝑛̇ 𝑁𝑎2𝑆2 𝑂3 294

𝑛̇ 𝐶𝑂2 𝑐𝑎𝑝𝑡𝑢𝑟𝑒𝑑

=

100 ∗ 10−6 0.9 ∗ 0.12

∗

0.050 𝑀 1.5 𝑀

=

3.1 ∗ 10−5 𝑚𝑜𝑙𝑁𝑎2 𝑆2 𝑂3

(24a)

(24b) (25a) (25b)

𝑚𝑜𝑙𝐶𝑂2 𝑐𝑎𝑝𝑡𝑢𝑟𝑒𝑑

295

Conclusions

296

The kinetics for NO2-catalyzed sulfite oxidation were measured at conditions relevant to NO2

297

pre-scrubbing with NaOH. An uninhibited sulfite system is not effective for NO2 absorption; the

298

ratio of SO32- oxidized/NO2 absorbed is on the order of 1000. As in limestone slurry scrubbing,

299

thiosulfate drastically reduces sulfite oxidation with a half-order dependence on thiosulfate

300

concentration. Increasing sulfite does not affect k1

301

between SO32- oxidized/NO2 absorbed and sulfite. Similar to previous results at higher partial

302

pressures of NO2, sulfite oxidation is half order in NO2 in the range of 0.2–1 Pa NO2. However,

303

NO2 is removed more efficiently at higher NO2 partial pressure (SO32- oxidized/NO2 absorbed is

304

inverse half order with NO2). At low mole fractions of NO2 (less than 10 ppm), EDTA or a

305

similar chelating agent should be added to chelate the metal ions. Sulfite oxidized with 0.01 mM

306

Fe2+ has a k1

307

sulfite oxidation rates were shown for the first time to follow an Arrhenius temperature

308

dependence with an apparent activation energy of 24.1 kJ/mol. Oxygen partial pressures do not

309

affect the oxidation rate until below roughly 5 kPa O2. For common industrial applications,

310

oxygen concentrations will not affect sulfite oxidation. A NaOH scrubber has been modeled to

obs

obs,

but produces a first-order correlation

2.8 times greater than when oxidized with 0.02 mM EDTA. NO2-catalyzed


23


1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 24 of 27

311

show how the f value can be used to calculate thiosulfate concentrations.

For an inlet

312

concentration of 100 ppm SO2 and 2.5 ppm NO2 with a NaOH scrubber operating at 55 °C,

313

0.050 M thiosulfate would be needed to keep SO32- concentration at 0.010 M. This corresponds

314

to a feed rate of 3.1*10-5 mole Na2S2O3 per mole of CO2 captured.

315

Acknowledgements

316

The authors acknowledge financial support from the Texas Carbon Management Program.

317

The authors declare the following competing financial interest(s): One author of this

318

publication consults for Southern Company and for Neumann Systems Group on the

319

development of amine scrubbing technology. The terms of this arrangement have been reviewed

320

and approved by the University of Texas at Austin in accordance with its policy on objectivity in

321

research. The authors have financial interests in intellectual property owned by the University of

322

Texas that includes ideas reported in this paper.

323

Associated Content

324

Supporting Information.

325

Analysis of anion species by anion chromatography. This material is available free of charge

326 327

via the Internet at http://pubs.acs.org. B. List of Uncommon Abbreviations and Symbols Abbreviation

Description

Unit

𝑁𝑁𝑂𝑆𝑐𝑟

Steady state nitrosamine concentration in the amine scrubber

mmol/kg

𝑦𝑁𝑂2 𝑖𝑛𝑙𝑒𝑡

Inlet NO2 mole fraction to the amine scrubber

mol NO2/mol flue gas


24

Page 25 of 27

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


𝑘𝐷𝑒𝑐𝑜𝑚𝑝

Pseudo-first-order decomposition rate constant of the nitrosamine at desorber conditions.

s-1

𝜏𝐷𝑒𝑠𝑜𝑟𝑏𝑒𝑟

Residence time of the desorber

s

𝐺 𝐿 𝑆𝑐𝑟

Ratio of molar flow rate of flue gas to volumetric flow rate of solvent bleed

mol flue gas/L solvent

𝑁𝑜𝑔

Number of theoretical transfer units

-

𝑚𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛

Mass of solution in the HGF

kg

𝑉̇

Volumetric flow rate of inlet gas at STP

L/min

[NO2]Bypass

NO2 Concentration in Bypass mode outlet gas stream

mol CO2/mol gas stream

[NO2]HGF

NO2 Concentration in outlet HGF gas stream

mol CO2/mol gas stream

kg’

Overall liquid-side mass transfer coefficient in the HGF

mol/s m2 Pa

𝑎𝑒 𝐺𝐻𝐺𝐹

Specific gas-liquid interfacial area of the HGF over gas flow rate.

s m2 Pa/mol

𝑚𝑜𝑙 𝑆𝑂3 2− 𝑜𝑥

Number of moles of SO32- oxidized in HGF solution

mol

𝑚𝑜𝑙 𝑁𝑂2𝐴𝑏𝑠

Number of moles of NO2 absorbed into HGF solution

mol

𝑓

Amount of SO32- oxidized per mol of NO2 absorbed

mol SO32-/ mol NO2

𝑓𝑜

Regressed prefactor for calculating 𝑓

mol SO32-/ mol NO2

𝑦𝑁𝑂2

Inlet NO2 mole fraction

mol NO2/mol flue gas

𝑘1 obs

Sulfite oxidation first order rate constant

min-1

𝑘0 obs

Sulfite oxidation zeroth order rate constant

𝑚𝑜𝑙 𝑘𝑔 ∗ 𝑚𝑖𝑛


25


1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

328

Page 26 of 27

𝑓𝑜𝑏𝑠

Observed 𝑓 ratio for each experiment

mol SO32-/ mol NO2

𝑓𝑐𝑎𝑙𝑐

Regressed 𝑓 ratio normalized to 0.40 mol/kg SO32-

mol SO32-/ mol NO2

G

Molar flow rate of flue gas

mol/s of flue gas

L

Volumetric flow rate of solvent bleed

L/s of solvent

References

329 330

(1)

Fine, N. A.; Rochelle, G. T. Absorption of Nitrogen Oxides in Aqueous Amines. Energy Proc. 2014, 63, 830–847.

331 332

(2)

Nash, T. The Effect of Nitrogen Dioxide and of Some Transition Metals on the Oxidation of Dilute Bisulphite Solutions. Atmos. Environ. 1979, 13, 1149–1154.

333 334

(3)

Huie, R.; Neta, P. Chemical Behavior of Sulfur Trioxide (1-)(SO3-) and Sulfur Pentoxide (1-)(SO5-) Radicals in Aqueous Solutions. J. Phys. Chem. 1984, 88, 5665–5669.

335 336

(4)

Shen, C. Nitrogen Dioxide Absorption in Aqueous Sodium Sulfite, Ph.D. Dissertation, The University of Texas at Austin, Austin, TX, 1997.

337 338

(5)

Takeuchi, H.; Ando, M.; Kizawa, N. Absorption of Nitrogen Oxides in Aqueous Sodium Sulfite and Bisulfite Solutions. Ind. Eng. Chem. Process Des. Dev. 1977, 16, 303–308.

339 340

(6)

Shen, C.; Rochelle, G. T. Nitrogen Dioxide Absorption and Sulfite Oxidation in Aqueous Sulfite. Environ. Sci. Technol. 1998, 32, 1994–2003.

341 342

(7)

Sexton, A. J. Amine Oxidation in CO2 Capture Processes, Ph.D. Dissertation, The University of Texas at Austin, Austin, TX, 2008.

343 344 345

(8)

Fine, N. A.; Nielsen, P. T.; Rochelle, G. T. Decomposition of Nitrosamines in CO2 Capture by Aqueous Piperazine or Monoethanolamine. Environ. Sci. Technol. 2014, 48, 5996–6002.

346 347 348

(9)

Kameoka, Y.; Pigford, R. L. Absorption of Nitrogen Dioxide into Water, Sulfuric Acid, Sodium Hydroxide, and Alkaline Sodium Sulfite Aqueous Solutions. Ind. Eng. Chem. Fundam. 1977, 16, 163.

349 350 351

(10)

Ulrich, R. K.; Rochelle, G. T.; Prada, R. E. Enhanced Oxygen Absorption Into Bisulfphite Solutions Containing Transition Metal Ion Catalysts. Chem. Eng. Sci. 1986, 41, 2183– 2191.


26

Page 27 of 27

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


352 353 354

(11)

Oblath, S. B.; Markowitz, S. S.; Novakov, T.; Chang, S. G. Kinetics of the Formation of Hydroxylamine Disulfonate by Reaction of Nitrite with Sulfites. J. Phys. Chem. 1981, 85, 1017.

355 356 357

(12)

Susianto; Pétrissans, M.; Pétrissans, A.; Zoulalian, A. Experimental Study and Modelling of Mass Transfer during Simultaneous Absorption of SO2 and NO2 with Chemical Reaction. Chem. Eng. Process. Process Intensif. 2005, 44, 1075.


27

NO2-Catalyzed Sulfite Oxidation - Industrial & Engineering Chemistry

Recommend Documents