Evaluation and Modeling of VaporâLiquid Equilibrium and CO2

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Evaluation and Modelling of Vapour-Liquid-Equilibrium and CO2 Absorption Enthalpies of Aqueous Designer Di-amines for Post Combustion Capture Processes Weiliang Luo, Qi Yang, William Owen Conway, Graeme Puxty, Paul H. M. Feron, and Jian Chen Environ. Sci. Technol., Just Accepted Manuscript • Publication Date (Web): 31 May 2017 Downloaded from http://pubs.acs.org on June 2, 2017

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Environmental Science & Technology

1

Evaluation and Modelling of Vapour-Liquid-Equilibrium and CO2

2

Absorption

3

Combustion Capture Processes

Enthalpies

of

Aqueous

Designer

Di-amines

for

Post

4 5

Weiliang Luo†, Qi Yang‡, William Conway§,*, Graeme Puxty§, Paul Feron§, and Jian Chen†,*

6

†

State Key laboratory of Chemical Engineering, Tsinghua University, Beijing 100084, China

7

‡

CSIRO Manufacturing, Clayton, VIC 3168, Australia

8

§

CSIRO Energy, Mayfield West, NSW 2304, Australia

9 10 11

* Corresponding authors:

12

William Conway, e-mail: [email protected]; phone: +61 2 4960 6098. ORCID: 0000-0002-7958-3872

Jian Chen, e-mail: [email protected]; phone: +86 10 6279 8627. ORCID: 0000-0003-1695-7790

13 14

Abstract

15

Novel absorbents with improved characteristics are required to reduce the existing cost

16

and environmental barriers to deployment of large scale CO2 capture. Recently, bespoke

17

absorbent molecules have been specifically designed for CO2 capture applications, and their

18

fundamental properties and suitability for CO2 capture processes evaluated. From the study,

19

two unique di-amine molecules, 4-(2-hydroxyethylamino)piperidine (A4) and 1-(2-

20

hydroxyethyl)-4-aminopiperidine (C4), were selected for further evaluation including

21

thermodynamic characterisation. The solubilities of CO2 in two di-amine solutions with a

22

mass fraction of 15% and 30% were measured at different temperatures (313.15 - 393.15 K)

23

and CO2 partial pressures (up to 400 kPa) by thermostatic VLE stirred cell. The absorption

24

enthalpies of reactions between di-amines and CO2 were evaluated at different temperatures

25

(313.15 and 333.15 K) using a CPA201 reaction calorimeter. The amine protonation

26

constants and associated protonation enthalpies were determined by potentiometric titration.

27

The interaction of CO2 with the di-amine solutions was summarized and a simple

28

mathematical model established that could make a preliminary but good prediction of the

29

VLE and thermodynamic properties. Based on the analyses in this work, the two designer di-

30

amines A4 and C4 showed superior performance compared to amines typically used for CO2

31

capture and further research will be completed at larger scale.

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ACS Paragon Plus Environment


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TOC/Abstract art

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1. Introduction

37

Extreme weather events most likely related to the enhanced greenhouse effect have

38

attracted widespread attention worldwide. The emission of CO2, the major components

39

driving rapid enhancement of the greenhouse effect, has long been suspected as a major

40

cause of this issue. Carbon capture and storage (CCS), specifically post-combustion capture

41

(PCC) technologies, have been identified as a crucial and key strategy in the development

42

and implementation of short to intermediate term carbon reduction strategies from coal-fired

43

power generation1, 2. Currently, commercial PCC processes employ amine-based chemical

44

absorption to selectively bind CO2 present in industrial process flue gases, and it is this

45

method, often referred to as amine scrubbing, that is becoming one of the dominant

46

technology options for CO2 capture from coal-fired power plants by 20303. Aqueous

47

alkanolamine solutions have been well researched for use in chemical absorption processes.

48

Among the vast list of absorbents, monoethanolamine (MEA) is considered to be the most

49

mature reference solvent for its high capacity and fast absorption rate with CO2. However,

50

several shortcomings with MEA including high thermal and oxidative degradation rates, high

51

corrosion, and notably high energy consumption for solvent regeneration, limit the large scale

52

deployment of MEA processes. New and improved solvents are needed to overcome these

53

issues and improve confidence in PCC technology.

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Traditionally, large amounts of experimental data are required to characterize a new

55

solvent and the measurement process is very time-consuming. As a result, different criteria

56

have been presented in the literature for pre-selection of molecules preceding full

57

characterization and, if successful, subsequent pilot plant experimentation. Rapid CO2

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absorption rates as well as high cyclic capacity have been a critical criterion for absorbents; a

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promising amine ideally possessing both characteristics among other properties4. 2


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Furthermore, energy conservation in solvent regeneration is considered to be a brand new

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criterion in recent research5. Previous studies proved that heterocyclic amines show better

62

performance in almost every field of CO2 capture6, 7, and piperazine (PZ) was even suggested

63

as a substitute for MEA as the reference solvent for its superior capabilities over MEA

64

including higher CO2 capture capacity, faster CO2 absorption rate, lower degradation, lower

65

corrosion, and above all, lower regeneration energy consumption8. However, precipitation

66

issues stemming from high concentrations of PZ solution and high CO2 loadings restrict its

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application window9. PZ derivatives with different substituents, such as 1-methylpiperazine

68

(1-MPZ), 1-ethylpiperazine (1-EPZ), 1,4-dimethylpiperazine (DMPZ), have been proven to

69

have better performance for CO2 capture than PZ10, 11. Based on the above, we attempted to

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utilise the robust cyclic structure of PZ or similar while minimising or eliminating

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precipitation generated in the presence of CO2 at high concentrations.

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The broader concept of di-amines specifically designed to perform well in CO2 capture

73

applications, were initially investigated in our previous work12, 13. The two parent skeleton

74

amines share structural similarities to the sub-family of piperidine derivatives that, as a group

75

of amines, were found to display excellent CO2 loadings in our previous study14. Selected

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functional groups were added for structural modification of the skeleton molecules, including

77

carbon chains to provide connection sites for additional groups, hydroxyl groups to increase

78

the boiling point and reduce volatilization, and additional amine groups to improve the

79

solubility of CO2. Through a series of preliminary molecule screening tests, the research

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scope has been narrowed down to four promising candidates. Due to limited availability of

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material and costs, two molecules, designated herein as A4 and C4, were selected for further

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research.

83

In this work, designer di-amines A4 and C4 were evaluated for their CO2 solubilities and

84

absorption enthalpies alongside MEA for validation and comparison. Cyclic capacities were

85

calculated from the solubility data as a function of temperature and CO2 partial pressure.

86

Fundamental reaction mechanisms between two di-amines and CO2 were established

87

according to which, a simple thermodynamic model of the di-amine-CO2-H2O vapour-liquid

88

equilibrium was employed for regression of the equilibrium constants and enthalpies for the

89

carbamate formation reactions. Using this model and the equilibrium constants, CO2

90

solubilities and absorption enthalpies of designer di-amines were recalculated and then

91

compared with the experimental data to determine the adequacy of the model for the

92

prediction of CO2 equilibrium pressure.

93 3



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2. Experimental

95

2.1 Chemicals

96

Carbon dioxide (CO2, 99.99%) and nitrogen (N2, 99.99%) were purchased from Beijing

97

Qianxi Gas Chemical Industry Co., Ltd., monoethanolamine (MEA, 99%) was purchased

98

from Aladdin Industrial Corporation; they were all used without further purification. 4-(2-

99

hydroxyethylamino)piperidine (A4, 99%) and 1-(2-hydroxyethyl)-4-aminopiperidine (C4,

100

99%) were synthesised using the published methods13, 15. All solutions were prepared with

101

deionized water. The structures of the di-amines investigated in this work are shown in Table

102

1. The concentrations of solutions are expressed herein in mass fraction (% w/w).

103

2.2 Vapour liquid equilibrium

104

Solubility measurements were performed in a thermostatic VLE reactor shown in Figure

105

1.The CO2 solubility data represents the relationship between the CO2 partial pressure and the

106

total CO2 loading in the solution. The CO2 loading was determined using the difference

107

between the total amount of CO2 introduced into the cell and that of unreacted CO2 in the gas

108

phase using the Peng-Robinson cubic equation16. A detailed description of the stirred cell,

109

calculations, and estimated uncertainty can be found in our previous work17.

110

2.3 CO2 Absorption enthalpies

111

Enthalpy measurements were performed using the CPA201 reaction calorimeter from

112

ChemiSens AB shown in Figure 1. The true heat flow was measured by a heat transducer at

113

the bottom of the 250 mL sealed reactor, and a high precision Peltier element was used to

114

guarantee one-directional flow of heat. 100 mL of amine solution was placed in the reactor

115

and degassed under vacuum. The experiment was performed at a constant temperature until

116

equilibrium was reached. The vapour pressure of amine and water in the gas phase of the

117

reactor was assumed to be constant and equal to the total equilibrium pressure before the first

118

addition of CO2. CO2 was introduced batch-wise, and approximately same amount of CO2

119

was added into the reactor when equilibrium was achieved after the previous addition. An

120

automation script, which sought stability in pressure and true heat flow between each addition

121

of CO2, was used to control the experiments. The maximum deviation in the pressure and true

122

heat flow were required to be below ±0.01 bar and ±0.02 W respectively within a period of

123

10 minutes, so that the equilibrium for the system was considered to be reached.

124

The main source of uncertainty in the experimental data stems from the amount of CO2

125

added to the reactor. The amount of CO2 added to the reactor was recorded using a BIOS 4


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flow meter with an accuracy of ±0.8%. The accuracies of the temperature sensor and the

127

pressure transducer were ±2.5 K and ±1.0%, respectively. To obtain the total amount of heat

128

released during the absorption of CO2, the heat flow curve was integrated over the duration of

129

the loading interval. The uncertainty in the amount of heat released was estimated to be

130

±2.5%.

131 132

2.4 Amine protonation constants

133

The protonation constants of A4 and C4 (Eq. (4) and (5) below) was determined

134

between 288.15 and 318.15 K by potentiometric titration adapted from the method of

135

Fernandes, et al18. The amine solution was initially acidified by addition of standardised

136

hydrochloric acid (HCl) and back titrated to high pH with sodium hydroxide (NaOH).

137

Titrations were performed using a 665 Dosimat automated burette system in a jacketed

138

titration vessel connected to a Julabo ED-5 heating circulator. The pH of the solution was

139

measured with a Metrohm combined micro-pH glass electrode. The electrode signal was

140

acquired and amplified by a National Instruments NI-DAQ 7 board running with Matlab

141

software on a personal computer. The electrode signal, initially recorded in mV, was

142

transformed into the pH scale using the Hyperquad2008 software19. At each temperature, the

143

electrode was calibrated by titrating HCl solutions of known concentration with standardised

144

NaOH solution.

145 146

3. Modelling

147

Absorption of CO2 into aqueous di-amine solutions is described by the series of

148

equilibrium reactions in Eq. (1) - (9). CO2 reacts with water to form bicarbonate/carbonate

149

according to the reactions of Eq. (1) - (3). Eq. (1) is bicarbonate/carbonate protonation

150

equilibria, Eq. (2) is the reaction of CO2 and water to form bicarbonate and Eq. (3) represents

151

water autoprotolysis. The equations describing the equilibrium constants and their

152

temperature dependence for these reaction were taken from previous work20.

153

K1 H+ +CO32- ← →HCO3∆H1

154

K2 H+ +HCO3- ← →H2O+CO2 ∆H2

155

K3 H+ +OH- ← →H2O ∆H3

(1) (2) (3)

5



156

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Incorporating the di-amine into the model, the first and second protonation of the di-

157

amine, K4 and K5, are described by Eq. (4) and (5).

158

K4 H+ +Am← →AmH+ ∆H4

159

K5 H+ +AmH+ ← →AmH22+ ∆H5

(4) (5)

160

Formation of the mono and di-carbamate, K6 and K7, together with the corresponding

161

protonation of the mono and di-carbamate, K8 and K9, are described in Eq. (6) - (9).

162

Carbamate protonation has a negligible impact on the prediction of chemical speciation for

163

typical mono-amines due to the weakly basic nature of the carbamate carboxyl group.

164

However, in the case of di-amines the often strongly basic nature of the unreacted amine

165

group plays a more significant role in the chemistry and must be included in the chemical

166

model. Formation of the doubly protonated carbamate and doubly protonated di-carbamate is

167

unlikely under the conditions here, hence these reactions have been omitted from the model.

168

K6 HCO3- +Am← →AmCOO− +H2O ∆H6

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K7 HCO3- +AmCOO- ← → Am ( COO ) 2 +H2O ∆H7

170

K8 H+ +AmCOO- ← →HAmCOO ∆H8

171

K9 H + +Am ( COO )2 ← → HAm ( COO ) 2 ∆H 9

2-

2-

172 173

(6) (7) (8) -

(9)

Since tertiary amines are unable to react directly with CO2, as is the case for the second amine group in C4, Eq. (7) and (9) are not required for C4.

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To fully determine the speciation in the system of equilibria, equations defining the

175

equilibrium constants and mass balance of the system are required. The equilibrium constants

176

are defined according to Eq. (10), or Eq. (11) if water is involved, and the mass balance by

177

Eq. (12) - (15).

178

Ki =

179 180

cP γP γP ⋅ = K i0 ⋅ γ R1 ⋅ γ R 2 cR 1 ⋅ c R 2 γ R 1 ⋅ γ R 2

(10)

Ki K0 cP γP γP = ⋅ = i ⋅ cH 2 O cR1 ⋅ c R 2 γ R 1 ⋅ γ R 2 cH 2 O γ R 1 ⋅ γ R 2

(11)

(

total c AM = cAM + c AMH + + cAMCOO - + cHAMCOO + cAMH 2+ + cAM ( COO )2- + cHAM ( COO )2

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(

total cCO = cCO 2 + cCO 2- + cHCO - + cAMCOO - + cHAMCOO +2 cAM 2 3

3

2

( COO )2-2

2

+ 2 cHAM

6


( COO )-2

) )

(12) (13)

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c Htotal = c H + + c H C O - + 2 c C O 2 + c AM H + + 2 c AM H 2 + + c H AM C O O + c H A M + 3

2

183

( C O O )-2

+ cH 2O

(14) 184

total cOH + cH2O − = c OH−

(15)

cP and cR represent the concentration of the product species and each reactant

185

where

186

respectively for each equilibrium reaction, and the concentrations in parentheses were

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included in the case of A4, but not in C4. Activity coefficients were estimated using the

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simple Specific Interaction Theory (SIT) as described from Puxty and Maeder20. The

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regression of interaction parameters between ion pairs was found to not improve the model

190

fitting results and as such were not included in the SIT based calculation of activity

191

coefficients.

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In addition to the chemistry model, the Henry’s coefficient of CO2 in solution is required

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to relate the free aqueous CO2 concentration to CO2 partial pressure. In lieu of experimental

194

data for these specific aqueous amines, the value for water was used which has previously

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been shown to be a reasonable approach20.

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With established knowledge of the equilibrium constants for the series of CO2 hydration

197

reactions available from the literature21, and the corresponding amine reactions measured or

198

estimated in this work, the chemical speciation can be determined by solving the system of

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nonlinear simultaneous equations consisting of the equilibrium constants and mass balance

200

equations for the concentrations of all species using the Newton-Raphson method22. From the

201

calculated speciation, the free aqueous CO2 concentration in solution was determined, and

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then the CO2 partial pressure above the solution can be calculated using the Henry’s

203

coefficient. Conversely, equilibrium constants for the di-amine reactions in Eq. (4) - (9) can

204

be determined from regression of the experimental CO2 partial pressures and known total

205

concentrations from the VLE measurements. The Newton-Gauss-Levenberg/Marquardt

206

method was used as described from Puxty et al.23. The objective function to be minimised in

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the regression was the sum of the squared relative error (SSQRE) between calculated and

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experimental CO2 partial pressure:

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exp cal  pCO -pCO 2 ,i 2 ,i SSQRE = ∑  exp  pCO 2 ,i i =1  n

n

  

2

(16)

exp is the total number of data points, pCO2 ,i is the experimental CO2 partial pressure

210

where

211

cal and pCO2 ,i is the calculated CO2 partial pressure from the model. The relative error, instead of

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the absolute error, was used since the data varies over orders of magnitude. The agreement

213

between calculated and experimental data has been expressed as the average absolute relative

214

deviation (AARD):

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exp cal 1 n pCO2 ,i -pCO2 ,i AARD = ∑ × 100% exp n i =1 pCO 2 ,i

(17)

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The temperature dependence of the amine reaction equilibrium constants was described

217

using the form of the Van’t Hoff equation of Eq. (18). The regressed parameters for each

218

equilibrium reaction were the zero ionic strength reference state equilibrium constant at

219

0 313.15 K, lg K j ,313.15K , and the reaction enthalpy, ∆Hj .

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lg K j = lg K 0j ,313.15K −

∆H j

1  1 ⋅ −  2.3026 R  T 313.15 

(18)

221 222

4. Results and discussion

223

4.1 Vapour-Liquid-Equilibrium (VLE)

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CO2 solubilities in 30% MEA solutions from 313.15 to 393.15 K were initially

225

evaluated to validate the experimental VLE setup for measurements of CO2 solubilities.

226

Experimental CO2 partial pressure data as a function of CO2 loading is available in Table S1

227

in Supporting Information. The experimental VLE data together with that from literature24, 25

228

for comparison are shown in Figure 2. Good agreement was observed with the data from

229

literature at different temperatures, hence confirming the validity and suitability of the VLE

230

setup for measurements of CO2 solubilities.

231

CO2 solubilities in 15% and 30% aqueous solutions of the designer di-amines A4 and C4

232

were measured at temperatures of 313.15, 333.15, 363.15, and 393.15 K. CO2 partial

233

pressures as a function of CO2 loading in solutions of A4 and C4 in various concentrations

234

are available in Tables S2 - S5 in Supporting Information. Graphical representations of CO2

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solubilities of A4 and C4 solutions are shown in Figure 3 with fitted trends from the

236

modelling results.

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Cyclic loading and cyclic capacity were evaluated to assess the capability for CO2

238

capture between different solutions. Cyclic loading, in terms of mole CO2 per mole amine,

239

represents the difference between CO2 loadings after absorption and after desorption. Cyclic

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capacity, in terms of mole CO2 per kilogram solution, describes the total amount of CO2

241

removed by regenerating one kilogram of solution in desorption process for one cycle. 8


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According to the analysis results in our previous work10, the cyclic loading and cyclic

243

capacity of different solutions in an absorber-stripper cycle can be approximately evaluated

244

by the equilibrium solubility of CO2 corresponding with the simulation conditions

245

(absorption at 15 kPa, 313.15 K and desorption at 15 kPa, 393.15 K). A comparison of the

246

results is shown in Figure 4.

247

For the designer di-amines, the CO2 loadings in absorption and desorption of 15%

248

solution are marginally higher than those of 30% solution, demonstrating that a lower amine

249

concentration could provide a better opportunity for the reactive amine groups to contact with

250

free CO2 in dilute solution. Cyclic loading of the 15% solution is similar to that of 30%

251

solution, while the cyclic capacity of 15% solution is only half that of 30% solution as a

252

result of lower molality. Overall, higher amine concentrations would be favoured.

253

Considering the two designer di-amines, CO2 loadings after absorption and desorption in

254

A4 are both higher than those in C4 at same concentration, indicating that A4 possesses a

255

greater affinity for CO2 than C4 (assuming physical CO2 absorption does not play a

256

significant role in the absorption capacity). This result is consistent with the molecular

257

structure characteristics that there are two secondary amine groups in the A4 molecule, while

258

there is one primary and one tertiary amine group in the C4 molecule. Primary and secondary

259

amine groups are able to directly react with CO2 while tertiary amine cannot. Similarly, both

260

cyclic loading and cyclic capacity of A4 are marginally higher than C4, making A4 the more

261

promising molecule.

262

MEA was used as a comparative example to evaluate the capture performance of these

263

two designer di-amines. With the same mass fraction, it can be seen that cyclic loadings of

264

A4 and C4 are both at least twice as high as that of MEA, which is likely related to the high

265

CO2 loading after absorption (74% and 52% higher than that of MEA for A4 and C4

266

respectively). The CO2 loading after desorptionis a different situation with A4 17% higher

267

and C4 37% lower than that of MEA. However, the cyclic capacity of A4 and C4 are 14%

268

and 12% lower than that of MEA respectively, resulting from the higher molecular weight

269

and therefore lower molality of the designer di-amines when at equal mass fraction to MEA.

270

A straight chain di-amine solution, N-(2-aminoethyl)ethanolamine26 (AEEA, 30%) and a

271

straight chain tri-amine solution, diethylenetriamine17 (DETA, 20%) were used to investigate

272

how the molecular structure affects the CO2 capture performance of amines. The effective

273

average cyclic loadings assigned to each amine group of the designer di-amines (0.357 and

274

0.349 for 30% A4 and C4 respectively) are higher than those of AEEA and DETA (0.287 and

275

0.280 for 30% AEEA and 20% DETA respectively), indicating that the designer di-amines 9



276

utilize each reactive amine group more effectively. Among all these amines, both the

277

designer di-amines possess a lower CO2 loading at high temperature, enhancing the high

278

cyclic loading. It was postulated that this phenomenon could be due to the steric hindrance

279

and electron distribution introduced by the ring in the molecule similar to piperidine and

280

piperazine. It is understandable that the sterical hindrance of A4 makes its 4-carbamate

281

unstable, hence benefiting bicarbonate formation. Additionally, the tertiary amine group of

282

C4 also promotes bicarbonate formation. AEEA has a similar chemical structure to the

283

designer di-amines, however, its arrangement is one of a straight chain instead of the ring.

284

The CO2 loading after absorption in AEEA is approaching those of A4 and C4, however, the

285

CO2 loading after desorption is significantly higher. Thus, assumptions surrounding the

286

importance of the ring structure in improving the desorption performance of amines were

287

judged to be plausible. DETA has an extra primary amine group, replacing the hydroxyl

288

group in AEEA, and still has a high cyclic capacity despite the modification. The hydroxyl

289

group in the designer di-amines was initially designed to increase the boiling point of the

290

amine, however, the capture performance from DETA indicates that the hydroxyl group in

291

the designer di-amines could be replaced by another amine group to further improve the

292

capture performance.

293

4.2 CO2 absorption enthalpies

294

Differential and integral CO2 absorption enthalpies have been determined using a

295

CPA201 reaction calorimeter. The differential heat of CO2 absorption was determined by

296

dividing the heat released from each addition of CO2 by the amount of CO2 absorbed with

297

each addition. The integral heat of absorption was determined by dividing the total amount of

298

heat released by the total amount of CO2 absorbed up to the equilibrium point. With

299

increasing CO2 loading the amount of heat released following CO2 addition became smaller

300

which is consistent with the amines becoming saturated with CO2.

301

CO2 absorption enthalpies in 30% MEA solutions at 313.15 and 353.15 K were initially

302

determined to validate the experimental enthalpy setup. A summary of differential and

303

integral enthalpies is available in Table S6 in Supporting Information. The experimental

304

enthalpies comparing the data from literature27-29 in MEA solutions are shown in Figure 5.

305

High consistency for heat curves was observed, from the figures, thus confirming reliability

306

and validity for measurements of CO2 absorption enthalpies.

307

CO2 absorption enthalpies in 15% and 30% aqueous solutions of A4 and C4 were

308

measured at 313.15 K and 333.15 K. The measured differential enthalpies of CO2 absorption 10


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309

at various CO2 loading, as well as the integral values by integration, are given in Table S7

310

and S8 in Supporting Information. Graphical representations are shown in Figure 6. There

311

was obvious similarity with MEA in the CO2 absorption enthalpy for the designer di-amines,

312

where CO2 absorption enthalpies decreased with increasing CO2 loading. However, unlike

313

MEA there was an approximately linear decrease for the integral enthalpies at the full range

314

of CO2 loadings, while the integral enthalpies of MEA remained stable before reaching the

315

critical loading point of 0.5. This indicates the ongoing reactions in solutions of the designer

316

di-amines include more than one dominant reaction. According to the previous work in which

317

the reaction products in designer di-amine solutions with CO2 were determined13, carbamate

318

was the primary product at low CO2 loading, and carbonate/bicarbonate with increasing CO2

319

loading. As a result, the reaction enthalpies should be composed of the formation enthalpies

320

of carbamate and carbonate/bicarbonate, and the former is higher than the latter. This

321

phenomenon indicates the reaction mechanism of designer di-amines and CO2 at low loading

322

is similar to the mechanism of primary amines with CO2, however, the formation of

323

carbonate/bicarbonate should be taken into account as CO2 loading increases.

324

4.3 Amine protonation constants

325

As both amines are di-amines two protonation constants were determined for each amine

326

at each measurement temperature. The results from potentiometric titration and linear

327

regression are tabulated in Table 2.

328

4.4 Modelling regression

329

95 and 92 points of experimental VLE data, covering an expansive range of CO2

330

loadings and temperatures, were used to regress equilibrium constants for the chemical model

331

of A4 and C4 respectively. Thus any chemical model regressed by this data set can be

332

broadly considered to be representative of the chemical and physical behaviour relevant to

333

the CO2 absorption process with these di-amines.

334

As the CO2-H2O chemistry has been well characterised, the parameters for Eq. (1) to (3)

335

were fixed at the values used in literature21. The parameters for the protonation reactions of

336

Eq. (4) and (5) were fixed at the values given in Table 2. The regressed parameters were the

337

equilibrium constants at zero ionic strength and infinite dilution, and enthalpies for carbamate

338

formation, Eq. (6) and (7) and carbamate protonation, Eq. (8) and (9). It was found that the

339

parameters of the di-carbamate formation, Eq. (7) and di-carbamate protonation, Eq. (9),

340

were not defined by the data sets. This indicates that these reactions did not occur to a

341

significant extent and could be excluded from the model. This was expected for amine C4, as 11



342

one amine is tertiary, however for amine A4, it was not previously known if di-carbamate

343

formation would be significant. Thus, the resulting chemical models only included Eq. (6)

344

and (8) resulting in 4 regressed parameters. Inclusion of interaction parameters between ionic

345

species did not result in an improved fit (reduced AARD) and therefore, were not included in

346

the model (that is the SIT model was used without interaction parameters). The regressed

347

parameter values are also given in Table 2. The parameter of particular interest is the stability

348

0 constant of carbamate formation, lg K 6,313.15K . The value for A4 is smaller than C4,

349

indicating a less stable carbamate. This is consistent with the speciation determined by NMR

350

in Yang, et al.13 which showed significantly more carbamate formation for C4. Also the

351

enthalpy of protonation for the C4 carbamate is smaller than the A4 carbamate. This is

352

consistent with protonation of a tertiary amine group rather than a primary/secondary amine

353

group18.

354

A parity plot of CO2 partial pressures is shown in Figure 7 with a random error

355

distribution. The model calculated VLE values using the regressed parameters are shown in

356

Figure 3. As can be seen, the AARD of the model for C4 was 11.9%, which was good enough

357

to represent the behaviour in the absorption process and showed agreement across the full

358

dataset. Meanwhile, the AARD for A4 was 17.3%, which is larger but still acceptable for a

359

simple model.

360

The model calculated integral enthalpies are shown in Figure 6. The modelling results

361

for A4 show good agreement with the experimental results below a CO2 loading of 0.5 and

362

the agreement is good for C4 below a loading of 0.8. Above these loadings the model

363

systematically over estimates the absorption enthalpy with a largest deviation of ~20 kJ/mol

364

CO2. The model also predicts that the enthalpy profiles for A4 cross for both temperatures. It

365

is unclear why this is the case, however it may be related to the slow mass transfer at high

366

CO2 loading resulting in small heat flows over extended time periods that are difficult to

367

measure. Additionally, uncertainty in the Henry coefficient may also be a factor as it has a

368

larger influence upon the estimated enthalpy as the CO2 loading increases. In order to

369

improve the model, the measured enthalpy data could be used in the regression of the

370

equilibrium constants and reaction enthalpies. However, combining such disparate data sets

371

in a model regression is a challenging task requiring complex statistical analysis beyond the

372

scope of this work.

373

The calculated results of VLE data were in good agreement with the experimental data,

374

and the predicted absorption enthalpies were well match to measured data at low CO2

12


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375

loadings. These two designer di-amines were found to have beneficial performance

376

characteristics for CO2 capture from flue gases. Relevant research of their properties at high

377

CO2 loading is still in progress.

378 379 380 381

Supporting Information Nomenclature, CO2 solubility data in Table S1 - S5, CO2 reaction enthalpy data in Table S6 - S8.

382 383

Acknowledgements

384

Financial supports from National Science and Technology Support Program of China

385

(No.2015BAC04B01), the National Natural Science Foundation of China (No.51134017), are

386

greatly appreciated. Financial supports from Australia-China Joint Coordination Group (JCG)

387

Partnership Fund is also appreciated. Huaneng Clean Energy Research Institute is gratefully

388

acknowledged for providing access to their experimental instruments.

389 390

Reference

391 392 393 394 395 396 397 398 399 400 401 402 403 404 405 406 407 408 409 410 411 412 413 414 415

1. Leung, D. Y. C.; Caramanna, G.; Maroto-Valer, M. M., An overview of current status of carbon dioxide capture and storage technologies. Renewable Sustainable Energy Rev. 2014, 39, 426-443. 2. MacDowell, N.; Florin, N.; Buchard, A.; Hallett, J.; Galindo, A.; Jackson, G.; Adjiman, C. S.; Williams, C. K.; Shah, N.; Fennell, P., An overview of CO2 capture technologies. Energy Environ. Sci. 2010, 3, (11), 1645-1669. 3. Rochelle, G. T., Amine scrubbing for CO2 capture. Science 2009, 325, (5948), 1652-1654. 4. Ma’mun, S.; Svendsen, H. F.; Hoff, K. A.; Juliussen, O., Selection of new absorbents for carbon dioxide capture. Energy Convers. Manage. 2007, 48, (1), 251-258. 5. Zhang, X.; Fu, K.; Liang, Z.; Rongwong, W.; Yang, Z.; Idem, R.; Tontiwachwuthikul, P., Experimental studies of regeneration heat duty for CO2 desorption from diethylenetriamine (DETA) solution in a stripper column packed with Dixon ring random packing. Fuel 2014, 136, 261-267. 6. Conway, W.; Wang, X.; Fernandes, D.; Burns, R.; Lawrance, G.; Puxty, G.; Maeder, M., Toward rational design of amine solutions for PCC applications: the kinetics of the reaction of CO2(aq) with cyclic and secondary amines in aqueous solution. Environ. Sci. Technol. 2012, 46, (13), 7422-7429. 7. Robinson, K.; McCluskey, A.; Attalla, M. I., An ATR-FTIR study on the effect of molecular structural variations on the CO2 absorption characteristics of heterocyclic amines, part II. Chemphyschem 2012, 13, (9), 2331-41. 8. Rochelle, G.; Chen, E.; Freeman, S.; Van Wagener, D.; Xu, Q.; Voice, A., Aqueous piperazine as the new standard for CO2 capture technology. Chem. Eng. J. 2011, 171, (3), 725-733. 9. Li, H.; Li, L.; Nguyen, T.; Rochelle, G. T.; Chen, J., Characterization of piperazine/2aminomethylpropanol for carbon dioxide capture. Energy Procedia 2013, 37, 340-352. 13



416 417 418 419 420 421 422 423 424 425 426 427 428 429 430 431 432 433 434 435 436 437 438 439 440 441 442 443 444 445 446 447 448 449 450 451 452 453 454 455 456 457 458 459 460 461 462

10. Li, H.; Moullec, Y. L.; Lu, J.; Chen, J.; Marcos, J. C. V.; Chen, G., Solubility and energy analysis for CO2 absorption in piperazine derivatives and their mixtures. Int. J. Greenh. Gas Con. 2014, 31, 25-32. 11. Li, H.; Le Moullec, Y.; Lu, J.; Chen, J.; Valle Marcos, J. C.; Chen, G.; Chopin, F., CO2 solubility measurement and thermodynamic modeling for 1-methylpiperazine/water/CO2. Fluid Phase Equilib. 2015, 394, 118-128. 12. Conway, W.; Yang, Q.; James, S.; Wei, C.-C.; Bown, M.; Feron, P.; Puxty, G., Designer amines for post combustion CO2 capture processes. Energy Procedia 2014, 63, 1827-1834. 13. Yang, Q.; Puxty, G.; James, S.; Bown, M.; Feron, P.; Conway, W., Toward intelligent CO2 capture solvent design through experimental solvent development and amine synthesis. Energy Fuels 2016, 30, (9), 7503-7510. 14. Puxty, G.; Rowland, R.; Allport, A.; Yang, Q.; Bown, M.; Burns, R.; Maeder, M.; Attalla, M., Carbon dioxide postcombustion capture: A novel screening study of the carbon dioxide absorption performance of 76 amines. Environ. Sci. Technol. 2009, 43, (16), 6427-6433. 15. Yang, Q.; James, S. N.; Ballard, M. J.; Bown, M.; Feron, P.; Puxty, G. D. Gas capture process. WO2012142668A1, 2012. 16. Peng, D.-Y.; Robinson, D. B., A new two-constant equation of state. Ind. Eng. Chem. Fundam. 1976, 15, (1), 59-64. 17. Luo, W.; Guo, D.; Zheng, J.; Gao, S.; Chen, J., CO2 absorption using biphasic solvent: Blends of diethylenetriamine, sulfolane, and water. Int. J. Greenh. Gas Con. 2016, 53, 141148. 18. Fernandes, D.; Conway, W.; Wang, X.; Burns, R.; Lawrance, G.; Maeder, M.; Puxty, G., Protonation constants and thermodynamic properties of amines for post combustion capture of CO2. J. Chem. Thermodyn. 2012, 51, 97-102. 19. Gans, P.; O'Sullivan, B., GLEE, a new computer program for glass electrode calibration. Talanta 2000, 51, (1), 33-37. 20. Puxty, G.; Maeder, M., A simple chemical model to represent CO2-amine-H2O vapourliquid-equilibria. Int. J. Greenh. Gas Con. 2013, 17, 215-224. 21. Edwards, T. J.; Maurer, G.; Newman, J.; Prausnitz, J. M., Vapor-liquid equilibria in multicomponent aqueous solutions of volatile weak electrolytes. AlChE J. 1978, 24, (6), 966976. 22. Maeder, M.; Neuhold, Y.-M., Practical data analysis in chemistry. Elsevier Science: 2007; Vol. 26, p 350. 23. Puxty, G.; Maeder, M.; Hungerbühler, K., Tutorial on the fitting of kinetics models to multivariate spectroscopic measurements with non-linear least-squares regression. Chemometrics Intellig. Lab. Syst. 2006, 81, (2), 149-164. 24. Lee, J. I.; Otto, F. D.; Mather, A. E., Equilibrium between carbon dioxide and aqueous monoethanolamine solutions. J. Appl. Chem. Biotechnol. 1976, 26, (1), 541-549. 25. Jou, F.-Y.; Mather, A. E.; Otto, F. D., The solubility of CO2 in a 30 mass percent monoethanolamine solution. Can. J. Chem. Eng. 1995, 73, (1), 140-147. 26. Ma'mun, S.; Jakobsen, J. P.; Svendsen, H. F.; Juliussen, O., Experimental and modeling study of the solubility of carbon dioxide in aqueous 30 Mass % 2-((2aminoethyl)amino)ethanol solution. Ind. Eng. Chem. Res. 2006, 45, (8), 2505-2512. 27. Kim, I.; Svendsen, H. F., Heat of absorption of carbon dioxide (CO2) in monoethanolamine (MEA) and 2-(aminoethyl)ethanolamine (AEEA) solutions. Ind. Eng. Chem. Res. 2007, 46, (17), 5803-5809. 28. Lee, J. I.; Otto, F. D.; Mather, A. E., The solubility of H2S and CO2 in aqueous 14


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monoethanolamine solutions. Can. J. Chem. Eng. 1974, 52, (6), 803-805. 29. Jou, F.-Y.; Otto, F. D.; Mather, A. E., Vapor-Liquid Equilibrium of Carbon Dioxide in Aqueous Mixtures of Monoethanolamine and Methyldiethanolamine. Ind. Eng. Chem. Res. 1994, 33, (8), 2002-2005.

15



468 469

Table 1. The structures of di-amines A4 and C4. Di-amines

Structure

MW (g.mol-1)

A4

144.22

C4

144.22

470

16


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Page 17 of 25


471

Table 2. Amine protonation constants at zero ionic strength as determined from potentiometric titration and

472

the values of

lg K 0j ,313.15K and ∆Hj from Eq. (4), (5), (6), and (8) determined by regression. Experimental

Regression

lg K 0j ,313.15K

∆H j (kJ/mol)

Eq.

288.15 K

298.15 K

308.15 K

318.15 K

(4)

10.56

10.39

10.11

9.95

10.04±0.02

37.0±2.9

(5)

7.42

7.24

7.04

6.86

6.95±0.01

33.0±0.9

(6)

-

-

-

-

-0.01±0.36

-16.2±12.4

(8)

-

-

-

-

9.95±0.32

-66.7±20.4

A4

AARD

17.3%

(4)

10.04

9.73

9.48

9.15

9.31±0.02

51.2±2.4

(5)

7.07

6.89

6.71

6.55

6.63±0.01

30.5±0.3

(6)

-

-

-

-

0.62±0.06

-14.8±2.7

(8)

-

-

-

-

8.86±0.05

-24.0±3.0

11.9%

C4

473

17



474

Figure Captions

475

Figure 1. Schematic diagram of (a) the thermostatic VLE stirred cell, (b) the CPA201 reaction calorimeter.

476

Figure 2. Comparison of CO2 solubilities in 30% MEA solutions from this work (Solid) and with literature

477

data (hollow) at temperatures of (a) 313.15 and 373.15 K, (b) 333.15 and 393.15 K.

478

Figure 3. Experimental data (points) and model calculated data (lines) of CO2 solubilities in A4 and C4

479

solutions from 313.15 to 393.15 K.

480

Figure 4. Equilibrium loadings, cyclic loadings, average cyclic loadings, cyclic capacities in various amine

481

solutions.

482

Figure 5. Differential and integral CO2 absorption enthalpies in 30% MEA solutions with comparison of

483

literature data.

484

Figure 6. Differential CO2 absorption enthalpies in 15% and 30% A4 and C4 solutions.

485

Figure 7. The parity plot of model calculated versus measured CO2 partial pressures.

18


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Page 19 of 25


486 487

Figure 1.

488

19



489 490

Figure 2.

491

20


Page 20 of 25

Page 21 of 25


492

493 494

Figure 3.

495

21



496 497

Figure 4.

498

22


Page 22 of 25

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499 500


Figure 5.

501

23



502

503 504

Figure 6.

505

24


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506 507


Figure 7.

508

25


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