High-Pressure Solubility of Carbon Dioxide (CO2) in Aqueous 1

Article pubs.acs.org/jced

High-Pressure Solubility of Carbon Dioxide (CO2) in Aqueous 1‑Methyl Piperazine Solution Aravind V. Rayer,† Yamuna Armugam,†,‡ Amr Henni,*,† and Paitoon Tontiwachwuthikul† †

International Test Center for Carbon Dioxide Capture (ITC), Faculty of Engineering, and ‡Department of Chemistry and Biochemistry, University of Regina, Regina, Saskatchewan, Canada S4S 0A2 ABSTRACT: The solubility of CO2 in aqueous solutions of 1-methyl piperazine (1-MPZ) was measured at (313 and 353) K for (15 and 30) wt % and up to a pressure of 7815 kPa, reaching a gas loading of 0.67 molCO2/molalkalinity. The electrolyte−NRTL (nonrandom two-liquid) model was used to correlate the data. Five adjustable ionic pair interaction parameters and six molecular interaction parameters were regressed, which allowed for the correlation of the experimental solubility data within an average deviation of 2 % in pressure and temperature and 7 % in experimental loading. The model was used to predict the speciation, heat of absorption, and pH of the loaded solutions. The predicted properties were compared with PZ to show the advantages in using piperazine derivatives in CO2 capture.

1. INTRODUCTION The absorption of CO2 and H2S in aqueous solutions of alkanolamines such as monoethanolamine (MEA), diethanolamine (DEA), and n-methyldiethanolamine (MDEA) is considered to be an essential step in the processing of sour natural gas streams, refinery off-gases, and synthesis gas for ammonia production.1 A better absorption capacity, a higher absorption rate, and a reduced solvent regeneration energy can be achieved by blending primary or secondary amines with tertiary amines.1 Piperazine (PZ) has been used as an activator in the BASFactivated MDEA technology. It is considered to be more effective than conventional accelerators and was suggested to be used in industrial processes.2 As a cyclic diamine with a sixmembered saturated ring, piperazine has a capacity to absorb theoretically more CO2 moles than all other conventional amines including MEA. Kadiwala et al.3 showed that at high partial pressures, a CO2 loading higher than 2 can be achieved. Due to their minimized angle or torsional strain and six-member heterocyclic molecule, six-membered cyclic alkanes were identified to be stable in terms of their thermal degradation.4 Cyclic diamines (piperazine, PZ; 1-methylpiperazine, 1-MPZ; 2-methyl piperazine, 2-MPZ and 1-ethyl piperazine, 1-EPZ) have higher reaction rates compared to MEA (commonly used a benchmark solvent in CO2 capture technology).5 The absence of a hydroxyl group in piperazine makes it less soluble in water, and as consequence, these solvents are, for now, used as additives to other slower reacting alkanolamines such as methyl diethanolamine (MDEA).2 The reaction rate and solubility of CO2 in aqueous PZ have been studied extensively by several researchers.3,5−16 Khalili et al.17 measured the pKa values of a series of six piperazine derivatives in water from T = (298.15 K to 323.15) K and found that 2-MPZ, PZ, and 1-MPZ have higher pKa values than primary, secondary, or tertiary amines. © 2014 American Chemical Society

Their higher basicity at lower temperature and their lower basicity at higher temperature make them good candidates for efficient CO2 absorption and regeneration. One of the advantages for suggesting concentrated PZ for amine-based absorption/ stripping process is its exceptional resistance to thermal degradation. PZ has a resistance to degradation up to 423.15 K, well above the standard stripper operating conditions.18 The disadvantage of 2-MPZ and PZ is their solid state at room temperature, in addition to a tendency to precipitate at lower temperature and at high concentrations.19 Freeman et al.20 give an explanation to the thermal degradation of structural analogs of PZ. Not having an alcohol group in their molecular structure enhances the reaction pathways leading to thermal degradation.21−24 Freeman et al.20 concluded that the addition of a methyl group on the amino group (2-MPZ) accelerated the degradation, and the addition of a methyl group to the amino group (1-MPZ) accelerated the degradation further. On the other hand, 1-MPZ is liquid at room temperature and has a lower and similar trend in pKa as PZ and 2-MPZ. Due to its solubility in water and its high reaction rate, we suggested that 1-MPZ might be a potential solvent for CO2 capture compared to conventional amines.5 Physical properties of 1-MPZ (densities, viscosities, refractive indices, and surface tensions) were measured in order to study the structural, physical, and transport behavior of the molecule and were reported by us elsewhere.25 Chen26 reported the solubility and kinetic data of 1-MPZ at (313.15 to 373.15) K at low pressures up to 100 kPa using a wetted wall column. Xu and Rochelle27 measured the solubility in 1-MPZ solution at high temperatures of (354 to 464) K and Received: June 10, 2014 Accepted: October 21, 2014 Published: November 4, 2014 3610

dx.doi.org/10.1021/je500526m | J. Chem. Eng. Data 2014, 59, 3610−3623

Journal of Chemical & Engineering Data

Article

for the solubility was found to be less than 5 % from our previous work3, by comparing the experimental data for the CO2 solubility in 3.29 M aqueous MEA solution at 313 K with data published by Jou and Mather.28 Jou and Mather.28 reported a general uncertainty in solubility of 3 % when BaCl2 is used and 2 % in the case where a gas chromatograph is used.

low pressures of (115 of 2819) kPa using an autoclave and a calorimeter. In this work, high-pressure solubilities of aqueous 1-MPZ (15 wt % and 30 wt %) were measured in a Jerguson cell at (313.15 and 353.15) K. The electrolyte-nonrandom two-liquid (E-NRTL) model was used to regress the binary interaction parameters and correlate the experimental data within a percentile absolute average deviation of 2 % (%AAD). The model was used to predict the pH of loaded solutions, speciation, activity coefficients, and heats of absorption, and all of the results were compared to those of aqueous PZ.

3. MODELING 3.1. Liquid Phase Equilibria. When CO2 diffuses into the liquid phase of aqueous 1-MPZ through the gas liquid interface, the reactions mentioned below are considered to happen. CO2 is dissolved in the chemical solvent primarily in nonvolatile ionic form. Reaction 1: Autoprotolysis of water

2. EXPERIMENTAL SECTION 2.1. Materials. 1-Methyl piperazine (1-MPZ, 99 mass %), sodium hydroxide (NaOH, 10 mol), barium chloride-60 mesh (BaCl2, 99.9%), and methyl orange-0.1 wt % were purchased from Sigma-Aldrich and were used without further purification. Hydrochloric acid (HCl, 0.1 N) was obtained from VWR scientific products. Carbon dioxide (CO2, 99.99%) used was purchased from Praxair. 2.2. Apparatus and Procedure. The experimental setup consisted of a high-pressure Jerguson cell with a design pressure of 34 474 kPa and a magnetic pump with a working pressure of 82 700 kPa and has a capability of bubbling 100 mL·min−1 of vapor through the liquid phase. A 150 cc vapor phase reservoir was connected to the top of the vapor−liquid equilibrium cell. The entire experimental setup was enclosed in an air bath (Z16, Cincinnati, OH). The temperature of the equilibrium cell was held constant within ± 0.1 K. The temperature and pressure inside the cell were measured by a digital temperature indicator from Omega (DP97) and a high accuracy 1000 psi (6895 kPa) digital Heise pressure gauge with an accuracy of ± 0.1 % of the full scale (± 6.9 kPa). The experimental procedure was explained in our previous work.3 A liquid phase sample was taken from the side of the cell into a 40 cc sampling bomb containing 20−25 mL of 2 M NaOH. The vessel was weighed before and after the addition of the caustic solution. Approximately 2 g of the cell liquid was withdrawn from the cell to purge the sampling line. The sampling bomb valve was attached to the sampling valve, and the two valves were opened to allow the cell liquid into the steel bomb. The valves were then closed and the bomb disconnected and agitated to mix the sample with the caustic. The sampling bomb was then reweighed to determine the mass of the sample collected. A 250 cm3 Erlenmeyer flask was prepared for CO2 analysis with 50 mL of distilled water and 25 mL of BaCl2 (0.5 M) solution. An aliquot of sample was added and allowed to stand for (48 to 72) h. The barium carbonate precipitate was filtered using Whatman filter paper #5. During the filtration, the funnel was covered with a watch glass to minimize exposure to atmospheric CO2. The precipitate was washed with distilled water until Dutoest pH paper indicated a filtrate pH of 5.9 (≈ 300 mL of water). The filter paper and precipitate were then shifted to a 250 cm3 Erlenmeyer flask and topped up with distilled water to the 100 mL mark. A magnetic stirring bar was used to shred the filter paper. Two to three drops of methyl orange were added, and the mixture was titrated to a reddishorange end point using a standard 0.1 N HCl. A sample calculation for the above procedure is detailed in the Appendix. In the analysis, a correction is made for the presence of carbonate in the caustic used; a small correction for the change in pH of pure water during the titration was also made, generally amounting to the subtraction of (0.2 to 0.3) cm3 from the volume of 0.1 N HCl used. The average absolute deviation

K1

2H 2O ⇔ H3O+ + OH−

(1)

Reaction 2: Formation of bicarbonate ions K2

CO2 + 2H 2O ⇔ H3O+ + HCO3−

(2)

Reaction 3: Dissociation of bicarbonate ions K3

HCO3− + H 2O ⇔ H3O+ + CO32 −

(3)

Reaction 4: First protonation of 1-methyl piperazine K4

MPZH+ + H 2O ⇔ MPZ + H3O+

(4)

Reaction 5: Formation of 1-methyl piperazine carbamate K5

MPZ + HCO3− ⇔ MPZCOO− + H 2O

(5)

Reaction 6: Second protonation of 1-methyl piperazine +

K4

HMPZH+ + H 2O ⇔ MPZH+ + H3O+

(6)

Reaction 7: Formation of protonated carbamate K6

H+MPZCOO− + H 2O ⇔ MPZCOO− + H3O+

(7)

The chemical equilibrium described above can be expressed in terms of the activity of reactants and products by the following expression:29 ln Kj =

∑ νi ,j ln ai = ∑ νi ,j ln(xiγi) i

i

(8)

where νi, j is the reaction stoichiometric coefficient of component i in jth reaction; xi is the mole fraction of the component i and γi is the activity coefficient of the component i. The temperature dependence of the equilibrium constants is represented as ln K = a +

b + c ln T + dT T

(9)

The coefficients a, b, and c are given in Table 1, for all reactions. In this work, water and 1-MPZ are considered as solvents, and the solution is considered as a mixed solvent system. The reference state of water is described by a symmetric reference state. γw → 1 as xi → 1

(10)

1-MPZ, CO2, and ionic activity coefficients are referred to pure water as asymmetric reference state conventions: γi* → 1 as xi → 0 3611

(11)



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Table 1. Chemical Equilibrium Constants Used in the E-NRTL Model K1 K2 K3 K4 K5 K6 K7 a

A

B

132.899 231.465 216.049 −10.61 7.153 −1.403 −79.658

−13445.9 −12092.1 −12431.7 −4289.2 −3616.136 −3595.893 −819.700

C

D

−22.477 0 −36.782 0 −35.482 0 0 0 −0.522 0.00149 0.149 −0.00609 10.185 0.00227 ln Kx = a + (b/T) + c ln T + dT, where x = 1 to 6

T/K

reference

273−498 273−498 273−498 298−323 313−353 313−353 313−353

Posey and Rochelle29 Posey and Rochelle29 Posey and Rochelle29 this worka this work this work this work

Regressed from Khalili et al.17

3.2. Vapor Phase Equilibria. The fugacity of each component in the liquid and vapor phase are equal at equilibrium. CO2 molecules in the equilibrium condition swap themselves between the liquid and vapor phase as per the activity coefficient approach as follows:30 ⎛ v∞(P − P 0 ) ⎞ solv * ⎟ ⎜ i exp x H φivyP = γ i i i i RT ⎠ ⎝

The long-range molecular interactions are given as a function of the Debye−Hückel parameter (Aϕ), ionic strength (Ix: mole fraction scale), the “closest approach” parameter (ρ), and the molecular weight of the solution (Ms): ⎛ 1000 ⎞1/2 ⎡⎛ 2Zi2 ⎞ Zi2Ix1/2 ⎤ ⎥ ⎟ln(1 + ρIx1/2) + ln γi PDH = ⎜ ⎟ Aϕ⎢⎜ ⎢⎣⎝ ρ ⎠ ⎝ Ms ⎠ 1 + ρIx1/2 ⎥⎦

(12)

(17)

(13)

Equation 15, mentioned above, is the Pitzer−Debye−Hückel formula, normalized to mole fractions of unity for solvent and zero for electrolytes. Debye−Hückel parameter (Aϕ) can be calculated as follows:

Solvent species like water and 1-MPZ follow: ⎛ v (P − P 0) ⎞ i 0 ⎜ i ⎟ φivyP = x γ P exp i i i i RT ⎝ ⎠

2 3/2 1/2 1 ⎛ 2πNAds ⎞ ⎛ Q e ⎞ ⎜ ⎟ ⎟ Aϕ = ⎜ 3 ⎝ 1000 ⎠ ⎜⎝ εskT ⎟⎠

In this work, the Redlich−Kwong−Soave (RKS) equation of state31 was used to calculate φvi . The RKS equation is given as P=

a α(T , ω) RT − 1 Vm − b Vm(Vm + b1)

(14)

Ix =

where a1 = 0.42748

R2Tc2 ; Pc

b1 = 0.08664

Tr =

T = 0.7 Tc

∑ ∑ yyi j [(aiαiajαj)0.5 (1 − kij)]; i

j

b=

i

where kij(= kji) is the binary interaction parameter. 3.3. Electrolyte Nonrandom Two-Liquid (E-NRTL) Activity Coefficient Model. Excess Gibbs energy for molecular and ionic interactions was calculated using the E-NRTL equation30 as the sum of the short-range interactions that exist between the immediate neighborhood of any species, and the Born correction was included to adjust the activity coefficients of the ions for the effect of the dielectric constant and the long-range interactions that exist beyond the immediate neighborhood of a central ionic species: (15)

ln γi = ln γiPDH + ln γi Born + ln γi E,lc

(16)

2 Q e2 ⎛ 1 1 ⎞ Zi ⎜ − ⎟ ·10−2 2kT ⎝ εs εw ⎠ ri

(20)

The local interaction contribution is accounted for by the nonrandom two-liquid theory. The basic assumption of the NRTL model is that the nonideal entropy of mixing is negligible compared to the heat of mixing; this is indeed the case for electrolyte systems. The short-range interactions are expressed in the E-NRTL model for molecular components as

∑ yb i i

GE GE,PDH GE,Born GE,lc = + + RT RT RT RT

(19)

i

ln γi Born =

where Tc is the critical temperature and Pc is the critical pressure. For a vapor phase with a multicomponent, the following mixing rules were used: (aα)m =

∑ xizi2

where NA is the Avogadro’s number; ds is the mass density of solvent; Qe is the electron charge; εs is the dielectric constant of the solvent; T is the temperature (K); k is the Boltzmann constant; and Zi is the charge number of ion i. The Born equation is used for the Gibbs energy of transfer of ionic species from the infinite dilution state in a mixed-solvent to the infinite dilution state in an aqueous phase:

RTc Pc

α(T , ω) = [1 + (0.48508 + 1.55171ω − 0.15613ω2) ·(1 − Tr0.5)]2

ω = −log10(Prsat) − 1;

1 2

(18)

ln γBk = +

∑j XjGjBτjB ∑k XkGk B

∑∑ c

a′

+

∑ B

∑ XG τ ⎤ XBG BB ⎡ ⎢τBB − k k k B k B ⎥ ∑k XkGk B ⎢⎣ ∑k XkGk B ⎥⎦

XcG Bc,a ′ c Xa ′ · ∑a ″ Xa ″ ∑k XkGkc,a ′ c

⎡ ∑ XkGkc,a ′ cτkc,a ′ c ⎤ ⎥ × ⎢τBc,a ′ c − k ⎢⎣ ∑k XkGkc,a ′ c ⎥⎦ +

∑∑ a

c

⎤ ∑ XG τ XaG Ba,c ′ a ⎡ Xc ⎢τBc,c ′ a − k k ka,c ′ a ka,c ′ a ⎥ · ∑c ″ Xc ″ XkGka,c ′ a ⎢⎣ ∑k XkGka,c ′ a ⎥⎦ (21)

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For cations: 1 ln γ k = Zc c +

⎡

Xa ′ ⎤ ∑k XkGk c,a ′ cτk c,a ′ c ⎥ ⎢⎣ ∑a ″ Xa ″ ⎥⎦ ∑k XkGk c,a ′ c

∑⎢ a′

∑ XG τ ⎤ XBGcB ⎡ ⎢τcB − k k k B k B ⎥ ∑k XkGk B ⎢⎣ ∑k XkGk B ⎥⎦

∑ B

+

τca,B = Cca,B +

∑∑ c′

a

τB,ca = C B,ca +

Dca,B T DB,ca T

⎡ (T ref − T ) ⎛ T ⎞⎤ + Eca,B⎢ + ln⎜ ref ⎟⎥ ⎝ T ⎠⎦ T ⎣

(34)

⎡ (T ref − T ) ⎛ T ⎞⎤ + E B,ca⎢ + ln⎜ ref ⎟⎥ ⎝ T ⎠⎦ T ⎣

(35)

c. For electrolyte−electrolyte pair parameters:

∑ XkGk a,c ′ aτk a,c ′ a ⎤ XaGca,c ′ a ⎡ Xc ′ ⎥ · ·⎢τca,c ′ a − k ∑c ″ Xc ″ ∑k XkGk a,c ′ a ⎢⎣ ∑k XkGk a,c ′ a ⎥⎦

τc ′ a,c ″ a = Cc ′ a,c ″ a +

Dc ′ a,c ″ a T

(22)

For anions: 1 ln γ k = Za a +

τca ′ ,ca ″ = Cca ′ ,ca ″ +

∑

∑∑ a′

c

(24)

where T = 298.15 K. 3.3.1. Model Parameters and Regression. The E-NRTL model involves several pure component parameters such as critical constants, acentric factor, compressibility factor, Brelvi− O’Connell parameter, and the Antoine equation constants for the vapor pressure of molecular species and the coefficients for Henry’s constant for CO2 in water. Molecular parameters were assumed to be similar to PZ, and they were taken from the DIPPR (Design Institute for Physical Properties) database. All required physical property constants for the VLE model are given in Table 2. The standard free energy of formation and the

(25)

Table 2. Physical Properties of Pure Component for the VLE Model32

∑ XkGk c,a ′ cτk c,a ′ c ⎤ XcGac,a ′ c ⎡ Xa ′ ⎥ · ·⎢τac,a ′ c − k ∑a ″ Xa ″ ∑k XkGk c,a ′ c ⎢⎣ ∑k XkGk c,a ′ c ⎥⎦ (23)

Ga,B =

∑a XaGca,B ∑a ′ Xa ′

∑c XcGca,B ∑c ′ Xc ′

αBc = αcB =

αBa = αaB =

∑a XaαB,ca ∑a ′ Xa ′

(26)

∑c XcαB,ca ∑c ′ Xc ′

T

⎡ (T ref − T ) ⎛ T ⎞⎤ + Eca ′ ,ca ″⎢ + ln⎜ ref ⎟⎥ ⎝ T ⎠⎦ T ⎣

ref

where GcB =

Dca ′ , ca ″

(37)

∑ XG τ ⎤ XBGaB ⎡ ⎢τaB − k k k B k B ⎥ ∑k XkGk B ⎢⎣ ∑k XkGk B ⎥⎦

B

+

⎡ X ⎤∑ X G τ ∑ ⎢ c ′ ⎥ k k ka,ca ka,ca ⎢ ⎥ c ′ ⎣ ∑c ″ Xc ″ ⎦ ∑k Xk Gk a,ca

⎡ (T ref − T ) ⎛ T ⎞⎤ + Ec ′ a,c ″ a⎢ + ln⎜ ref ⎟⎥ ⎝ T ⎠⎦ T ⎣ (36)

(27)

properties

H2O

CO2

1-MPZ

Tc/K Pc/kPa Vc/(m3·kmol−1) acentric factor (ω) racket (ZRA) Brelvi−O’Connell parameter

647.3 22048 0.0559 0.344 0.2432 0.0464

304.2 7376 0.0939 0.225 0.2736 0.0939

607.6 4135 0.38847 0.4138 0.323

τcB = −

ln GcB αcB

(28)

standard enthalpy of formation for 1-MPZ were assumed to be the same values as PZ (Table 3). The dielectric constant for

τaB = −

ln GaB αaB

(29)

Table 3. Reference State Parameters Used in This Work (kJ·kmol−1)26

τBa,ca = τaB − τca,B + τB,ca

(30)

τBc,ac = τcB − τca,B + τB,ca

(31)

MPZ MPZCOO− H2O CO2 HCO3− CO32− OH−

The pure component dielectric constant coefficients of nonaqueous solvents and Born radius of ionic species are required for a mixed-solvent electrolyte system. The temperature dependency of the dielectric constant of solvent B is ⎡1 1 ⎤ εB(T ) = AB + BB⎢ − ⎥ CB ⎦ ⎣T

(32)

BBB + FBB ln(T ) + G BBT T

ΔfHigi

170113.9 −229100 −228743 −394648

16410.96 −564400 −241976 −393773

ΔfG∞,aq i

ΔfH∞,aq i

−586770 −527810 −157244

−587333 −528336 −157403

1-MPZ (C5H12N2) is assumed to be the same as that of piperidine (C5H11N) for which temperature-dependent correlations are available and shown in Table 4. The mixed solvent dielectric constant εs is calculated by a simple mass fraction average.32 The temperaturedependent dielectric constant and Henry’s law are given by

The temperature dependency relations of the electrolyte NRTL parameters are: a. For molecule−molecule binary parameters: τBB = ABB +

ΔfGigi

(33)

ε=a+

b. For electrolyte−molecule pair parameters: 3613

b ⎡ 1 1 ⎤ − ⎢ ⎥ T (K) ⎣ T (K) 298.15 ⎦

(38)



Article

Table 4. Dielectric Constant of 1-MPZ and Water15 a1

species

Table 7. Pair Parameters Used as the Default in the E-NRTL Model15

b1

H2O 78.65 31989 1-MPZ 4.253 1532.2 ε = a1 + [b1/(T/K)]·{[1/(T/K)] − (1/298.15)}

b + c ln(T /K) + d T /K

ln(HCO2 /kPa) = a +

(39)

Table 5. Coefficients Regressed for Henry’s Law Constant in H2O (Pa) B

C

D

source

HCO2−H2O

170.7

−8478

−21.96

5.78e-3

Aspen Databank30

HHMPZCOO−H2O

−10

HMPZ−H2O

33.1

0

0

0

this work

−9180

0

0

this work

ln(HCO2/Pa) = a + [b1/(T/K)] + c ln(T/K) + d(T/K)

+ FT (K)G

a

72.55 −7207 0 0 −7.14 4.05·10−6 2 A + [B/(T/K

ions

32.73 72.83 −1.00·10−20 −9929 −3403 0 0 0 0 0 9.49·10−3 0 0 −8.56 0 0 2.91·10−16 0 0 6 0 + C)] + D(T/K) + E ln (T/K) + F(T/K)G

τca,m = Aca,m +

)⎤ ⎥ ⎥⎦

T , p , nj ≠ i

(43)

id nGmE,l = ΔGmixing − ΔGmixing

(44)

Hml(T ) =

(45)

∫T

T +ΔT

cpl ,m dT

∑ xiHi + ∑ xkHk∞ + HmE i

For solvents:

(46) (47)

k

35

Hi(T ) = ΔHfi g(T ref ) +

∫T

T ref

cpig dT

+ [Hi(T , p) − Hii g(T , p)]

ion pair and ion-pair are expressed as a function of temperature in the following forms:

(48) 36

For molecular solutes, cations, or anions:

Bm,ca T

1 ⎡ ∂(nGm ⎢ RT ⎢⎣ ∂ni

Hml(T + ΔT ) − Hml(T ) =

Regressed in this work using the data available in literature.18

τm,ca = A m,ca +

0.2 0.2 0.2

where n is the total mole number of the mixture; ni is the mole number of component in the mixture. ΔGmixing is the liquid Gibbs free energy of mixing; it is defined as the difference between the Gibbs free energy of the mixture and that of the pure components. ΔGidmixing is the ideal Gibbs free energy of mixing. Once the excess liquid functions are known, the thermodynamic properties of liquid mixtures can be computed as follows:35

Table 6. Coefficients Used for Calculating the Vapor Pressure for Molecular Species from the DIPPR Database15 (eq 40) A B C D E F G ln Pi0(N/m2) =

αij

0 0 0

⎛ ∂ ln γi ⎞ HmE,l = −RT 2 ∑ xi⎜ ⎟ ⎝ ∂T ⎠ i

(40)

CO2

Bca,m

−4.072 −8 −2

The liquid mixture enthalpy is:

Parameters A, B, C, D, E, and F are given in Table 6. Binary interaction parameters for molecule−molecule and molecule−

MPZa

Aca,m

0 0 0

E,l

⎛N⎞ B ln Pi0⎜ 2 ⎟ = A + + DT (K) + E ln T(K) ⎝m ⎠ T (K) + C

H2O

Bm,ca

8.045 15 10

ln γi =

zwitterion and considered to be not volatile; therefore, an extremely small Henry’s law constant is assigned to it.26 The vapor pressure of molecular species was predicted from the DIPPR database using the following equation:

components

Am,ca

pressure data of Xu and Rochelle27 for aqueous 1-MPZ solutions were entered into the data regression system (DRS) along with the liquid excess enthalpy and heat capacity data.36 Heat of mixing was calculated in Aspen Plus from temperature derivatives of activity coefficients. The heat capacity is calculated from the secondary temperature derivative of the activity coefficient. As a result, the temperature-dependent parameters are critical for modeling correctly the enthalpy, and it is recommended that enthalpy and heat capacity data be used to obtain the temperature-dependent interaction parameters. The related properties to the activity coefficient model are:

Henry’s law constants for CO2 in H2O, 1-MPZ, and H+MPZCOO− are listed in Table 5. H+MPZCOO− is a

A

parameter τ (H2O−ion pair) τ (CO2−ion pair) τ (molecule−ion pair)

Hk∞(T ) = ΔHf,∞k(T ref ) +

(41)

∫T

T ref

cp∞, k dT

(49)

where ΔT is the perturbation in temperature from T; HEm is the excess enthalpy of the mixture; ΔHigf (Tref) is the standard enthalpy of formation of component i at Tref; Tref is the reference temperature (298.15 K); cigp is the ideal gas heat capacity of component i; Higi is the ideal gas enthalpy of component i; H∞ k (T) is the infinite dilution aqueous enthalpy of ref component k; ΔH∞ f,k(T ) is the infinite dilution aqueous phase standard enthalpy of formation of component k at Tref; and c∞ p,k

Bca,m

(42) T The ion pair−ion pair parameters are considered insignificant and assigned a value of zero.34 The default values used for molecule−ion pair and ion pair−molecule interactions are listed in Table 7.35 The experimental total pressure data of this work, low pressure data of Chen and Rochelle,26 and total 3614



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Table 8. Coefficients for the Ideal Gas Heat Capacity (Cigp for Molecules, kJ·kmol−1·K−1) and Aqueous Infinite Dilution Heat Capacity (C∞,aq for Ions, kJ·kmol−1·K−1)26 p components C0 C1 C2 C3 C4 C5

a

MPZa

MPZH+

MPZCOO−

98.49 0.73 0 0 0 0

221 0 0 0 0 0

107 0.604 0 0 0 0

H+MPZCOO−

H3O+

OH−

HCO3−

CO3

−255 19.8 31.492 75.3 0 0.0734 0.063 0 0 5.6·10−5 2.73·10−5 0 0 0 −1.67·10−8 0 0 0 4.3·10−12 0 0 0 −4.17·10−16 0 Cpig = C0 + C1T + C2T2 + C3T3 + C4T4 + C5T5 Cp∞,aq = C0 + C1T + C2T2 + (C3/T) + (C4/T2) + (C5/T1/2)

−149 0 0 0 0 0

211 −0.882 8.75·10−4 −18800 0 0

1330 −5.56 5.19·10−3 −119000 0 0

H2Oa

CO2

Regressed in this work.

Table 9. Binary NRTL Parameters Regressed from the E-NRTL Model parameter

Aij

Aji

Bij

Bji

Cij

H2O−MPZ H2O−CO2

18.033 10.064

−2725.569 10.064

−2.431 −3268.135

−953.515 −3268.135

0.112 0.2

Table 10. Experimental Data for 15 wt % (1.76 mol·kgH2O−1) and 30 wt % 1-MPZ (4.28 mol·kgH2O−1) at 313 Ka 15 wt % 1-MPZ (1.76 mol·kgH2O−1)

30 wt % 1-MPZ (4.28 mol·kgH2O−1)

P/kPa

loading/αCO2

mCO2 mol·kgH2O−1

PCO2/kPa

P/kPa

loading/αCO2

mCO2 mol·kgH2O−1

PCO2/kPa

6805 6403 5525 4883 4067 3486 2900 2236 1642 1163 753 443 271

0.392 0.386 0.385 0.371 0.366 0.334 0.327 0.319 0.293 0.288 0.287 0.260 0.245

1.38 1.36 1.35 1.31 1.29 1.18 1.15 1.12 1.03 1.01 1.01 0.91 0.86

6798 6396 5518 4876 4060 3479 2893 2229 1635 1156 746 436 265

7154 6491 6019 5725 5540 4648 4365 3558 3036 2702 2132 1579 1158 805 580

0.670 0.644 0.656 0.642 0.639 0.616 0.600 0.611 0.589 0.558 0.560 0.554 0.544 0.524 0.519

5.73 5.51 5.61 5.50 5.47 5.27 5.14 5.22 5.04 4.77 4.79 4.74 4.65 4.49 4.44

7148 6485 6013 5719 5534 4642 4359 3552 3030 2696 2126 1573 1152 799 574

a Standard uncertainties u are u(P) = 6.9 kPa, ur(mCO2) = 2%, u(αCO2) = 0.005 molCO2·molalkalinity−1 (level of confidence = 0.95). PCO2/kPa is a calculated value from the total pressure using the equation: (PCO2 = P − PH0 2O − P01‑MPZ). PH0 2O and P01‑MPZ are vapor pressures of water and 1-MPZ calculated using eq 40.

Hilliard35 suggested a back algorithm procedure for selecting the interaction parameters for the E-NRTL model. Kadiwala et al.3 and Dash et al.15 employed few pair parameters for regressing the total pressure data of PZ. The same procedure and parameters were regressed in DRS by considering the liquid phase equilibrium. The maximum likelihood principal and Deming search method30 were used to minimize the following objective function:

is the infinite dilution of aqueous heat capacity polynomial of component k: Cpig = C0 + C1T + C2T 2 + C3T 3 + C4T 4 + C5T 5 C p∞ ,aq = C0 + C1T + C2T 2 +

C3 C C + 42 + 5 T T T

(50)

(51)

N

The values of heat capacity parameters for the 1-MPZ− CO2−H2O system are assumed to have the same values as similar species in the Frailie’s PZ model.33 The values for the heat capacity model parameters are reported in Table 8. (Cigp ) parameters for 1-MPZ and H2O and parameters were regressed in this work using experimental heat capacities measured in this laboratory.36 Experimental excess liquid molar enthalpies36 were used to obtain the molecular NRTL parameters for the 1-MPZ−H2O system, and the regressed parameters are given in Table 9.

F=

⎡ (P Est − P exp)2 i i

∑ ⎢⎢ i=1

⎣

σP2i

+

(TiEst − Tiexp)2 σT2i

+

(xiEst − xiexp)2 ⎤⎥ σx2i ⎦⎥

(52) Est Est where N is the experimental data points and PEst i , Ti , and xi are the estimated values of the corresponding measured values Exp of pressure (PExp i ), temperature (Ti ), and mole fraction of 2 2 2 Exp CO2 (xi ). σPi, σTi, and σxi are the standard deviations of the corresponding measured data and were maintained within the values reported by Austgen et al.34

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Table 11. Experimental Data for 15 wt % (1.76 mol·kgH2O−1) and 30 wt % 1-MPZ (4.28 mol·kgH2O−1) at 353 Ka 15 wt % 1-MPZ P/kPa

loading/αCO2

7815 6935 6773 5967 5628 5483 4908 4648 4145 3669 3173 2635 2149 1631 1132 726 406 253

0.376 0.357 0.357 0.357 0.356 0.353 0.341 0.332 0.320 0.306 0.288 0.274 0.269 0.239 0.236 0.227 0.206 0.193

mCO2 mol·kgH2O 1.32 1.26 1.26 1.26 1.25 1.24 1.20 1.17 1.13 1.08 1.01 0.96 0.95 0.84 0.83 0.80 0.72 0.68

30 wt % 1-MPZ −1

PCO2/kPa

P/kPa

loading/αCO2

mCO2 mol·kgH2O−1

PCO2/kPa

7769 6890 6728 5922 5583 5438 4863 4603 4101 3625 3137 2592 2106 1590 1092 686 368 216

7817 5526 5048 4320 3775 3497 2728 2477 2323 1920 1468 909 615 425 215

0.560 0.560 0.560 0.548 0.532 0.500 0.501 0.467 0.457 0.450 0.439 0.434 0.432 0.424 0.408

4.79 4.79 4.79 4.69 4.55 4.28 4.29 4.00 3.91 3.85 3.75 3.71 3.70 3.63 3.49

7780 5488 5014 4284 3740 3462 2693 2443 2280 1885 1435 876 582 392 182

Standard uncertainties u are u(P) = 6.9 kPa, ur(mCO2) = 2 %, u(αCO2) = 0.005 molCO2·molalkalinity−1 (level of confidence = 0.95). PCO2/kPa is a calculated value from the total pressure using the equation: (PCO2 = P − PH0 2O − P01‑MPZ). PH0 2O and P01‑MPZ are vapor pressures of water and 1-MPZ calculated using eq 40.

a

Figure 1. Comparison of the solubility data of CO2 in aqueous 1-MPZ (30 wt %), PZ (15 wt %3), and MEA (30 wt %28) solutions at 313.15 K.

4. RESULTS AND DISCUSSION

Table 12. Pair Parameters Regressed Using the E-NRTL Model

Experimental data obtained in this work for CO2 solubility at (313 and 353) K for (15 and 30) wt % 1-MPZ solutions are listed in Tables 10 and 11. The solutions were prepared at room temperature. A comparison of the solubility data of aqueous MEA (30 wt %),28 aqueous 1-MPZ (30 wt %), and aqueous PZ (15 wt %)3 is shown in Figure 1. At high pressure (≈ 7000 kPa) and at 313.15 K, PZ had a maximum capacity of 1.01 molCO2/molalkalinity, while 1-MPZ had a capacity of 0.67 molCO2/molalkalinity, whereas MEA reached 1.04 molCO2/ molalkalinity. The electrolyte NRTL model correlated the total pressure data and solution loading within 2 % and 7 % AAD,

parameter

Am,ca

Aca,m

τH+MPZCOO−[MPZH+,MPZCOO−] τH+MPZCOO−[MPZH+,HCO3−]

76.043 2.648

91.385 −15.646

respectively. The regressed NRTL binary parameters and E-NRTL pair interaction parameters in 1-MPZ are given in Table 12. 4.1. CO2 Equilibrium Solubility. Figure 2 shows the experimental PCO2 and the correlated PCO2 in aqueous 1-MPZ (15 wt % and 30 wt %) at (313.15 and 353.15) K. Solution 3616



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Figure 2. Equilibrium partial pressure of CO2 in aqueous solutions of 1-MPZ (15 wt % and 30 wt %).

loading (molCO2·mol1‑MPZ−1) increased with decreasing temperature and increasing pressure. The partial pressure of CO2 decreased more rapidly at the same loading for 15 wt % 1-MPZ than for 30 wt % 1-MPZ. When compared to the same loading and temperature, PCO2 of 1-MPZ was sometimes about an order magnitude higher than that of PZ. This observation is comparable to the work of Chen26 done at lower partial pressure solubility for 1-MPZ and PZ. The reason for this may be that the methylated tertiary amino group is much less reactive with CO2, and the conversion to carbamate group is not possible. Cullinane et al.37 studied the aqueous PZ system at low pressures and suggested that the equilibrium constant of transformation of CO2 to bicarbonate is less than the formation of PZ monocarbamate. Using the regressed parameters in this work, the experimental partial pressure data of CO2 in aqueous 1-MPZ at different temperatures26 were predicted and shown in Figure 3. The partial pressure was predicted within 1 % AAD and the loading within 3 % AAD. Xu and Rochelle27 reported the solubility data in terms of total pressure data at higher temperatures, which were predicted within 2 % AAD as shown in Figure 4. 4.2. Speciation of the CO2-Loaded 1-MPZ. Equilibrium concentrations of various species were predicted using the E-NRTL model for 30 wt % 1-MPZ. It was observed that 1-MPZ is steadily consumed with an increase in CO2 loading. The maximum concentration of [MPZCOO−] ions is 2.5 % and occurred at a loading of 0.1 molCO2·molalkalinity−1 and then became negligible with a further increase in loading. The maximum concentration of PZ carbamate is 2 % and occurred at loading 0.21 molCO2·molalkalinity−1.15 [MPZH+] and [HCO3−] ions showed an increasing trend with the increase in CO2 loading. These observations indicate that the formation of bicarbonate ions (reaction 2) and the protonation of 1-MPZ (reaction 4) control the absorption process. Protonation reactions are the main contributors in the case of solubility of CO2 in 1-MPZ, and the trend here is similar to PZ as observed by Dash et al.,15 Ermatchkov et al.,11 and Derks et al.8 This may be the reason for lower heat of absorption values for PZ and 1-MPZ.26 Cullinane and Rochelle37 suggested that, at high PCO2, the relative amount of bicarbonate (HCO3−) in the

loaded solution will be higher. Similar observations can be made in terms of the concentration of HCO3− ions at high pressure and at higher CO2 loading for the aqueous 1-MPZ solution. Molar heat capacity of aqueous 1-MPZ, liquid excess enthalpy, and the liquid phase speciation are predicted using the coefficients regressed in this work. The predicted heat capacity data and liquid phase excess enthalpy are shown in Figures 5 and 6 at various temperatures. The predicted molar heat capacity was found to deviate up to ± 1 kJ·kmol−1·K−1. The prediction of liquid excess enthalpy was found to deviate up to ± 20 kJ·kmol−1. These deviations mainly result from the assumption of reference state properties of PZ for 1-MPZ in Table 3. The model can be improved by providing the corresponding reference state parameters. As is well-known also, the model parameters can be regressed to better predict the heat of mixing values, but it will then not do such a good job with the solubility data. Using 1H and 13C1 NMR spectroscopy, the formation of 1methylpiperazine (1-MPZ) carbamate was monitored. Both NMR spectra were obtained with a Varian INOVA LC-500 MHz spectrometer operating at 499.59 MHz for 1H and 125.63 MHz for 13C1 NMR. D2O was used as an external standard to find the relative chemical shifts of the observed species. The 1H NMR provided detailed information about the formation of 1-MPZ carbamate and the unreacted 1-MPZ, whereas the 13C1 NMR spectrum not only provided information about the formation of 1-methylpiperazine carbamate and the unreacted 1-MPZ, but also information about other carbonated species such as carbonates and bicarbonates. Overall, both 1H and 13C NMR spectra showed that the reaction products between 1-MPZ and CO2 contained the expected 1-MPZ carbamate/ protonated 1-MPZ carbamate and the unreacted 1-MPZ/ protonated 1-MPZ. Additionally, the presence of a signal at δ161.8 ppm in the 13C NMR spectrum is attributed to the carbon atoms of the carbonate and bicarbonate ions, which are rapidly exchanging on the NMR time scale through proton scrambling. Table 13 and Figure 7 provide 1H NMR details about the proton environments of 1-MPZ/protonated 1-MPZ and 1-MPZ carbamate/protonated 1-MPZ carbamate, including the chemical shifts in ppm and the coupling constants J in Hz. 3617



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Figure 3. Predicted solubility data for aqueous solutions of 1-MPZ (45 wt %26) at different temperatures.

Figure 4. Predicted total pressure data for aqueous solutions of 1-MPZ at higher temperatures (393.15 K to 433.15 K).

and then the pH dropped to 3 at loading 1 molCO2·molalkalinity−1. This shows that 2 (NH) groups in PZ takes part in absorption and forms more carbamate. In 1-MPZ, due to the steric hindrance in the second (−N) group for 1-MPZ, which makes it less basic than PZ and forms less carbamate, Khalili et al.17 suggested that the addition of a methyl group in PZ molecule (1-MPZ) introduces a hindrance effect and lowers the basicity. At lower loadings, the increase in temperature decreased the pH and at higher loadings the increase in temperature increased the pH. This behavior indicates that the absorption capacity of the amine decreases as the CO2 loading increases in the solution with respect to the temperature. 4.4. Heat of Absorption of CO2 (ΔHabs) in 1-MPZ Solution. The heat released during the absorption process (heat of absorption) is an important process parameter in CO2 capture studies, because it provides information about the amount of heat produced in the absorber and the heat required

Table 14 and Figure 7 gives the 13C NMR details about the different carbon environments in 1-MPZ/protonated 1-MPZ and 1-MPZ carbamate/protonated 1-MPZ carbamate in ppm. 4.3. pH of the CO2-Loaded 1-MPZ. The pH values were calculated from the concentration of H3O+ ion using the following relation:

pH = − log10[H3O+]

(53)

In Figure 8, the predicted pH values of 15 wt % 1-MPZ loaded with CO2 are compared with the predicted pH values of 0.3 M (10 wt %) PZ loaded with CO2 using the regressed parameters of Kadiwala et al.3 It can be observed that the addition of CO2 decreased the pH of the aqueous 1-MPZ solution, and then the pH remained around 6 at loading up to 1 molCO2·molalkalinity−1, whereas the pH of the aqueous PZ solution decreased and dropped at loading around 0.5 molCO2·molalkalinity−1 3618



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Figure 5. Predicted heat capacity of aqueous 1-MPZ solutions.

Figure 6. Predicted liquid excess enthalpy of aqueous 1-MPZ solutions.

Table 13. 1H NMR Summary for 1-MPZ Solution in D2O δppm

splitting pattern

coupling constant J (Hz)

Table 14. 13C NMR Summary for 1-MPZ Solution in D2O

species

2.19 singlet 2.50 2.88 2.18 3.26

3.54

protons of methyl group of 1-MPZ/ protonated 1-MPZ broad singlet methylene protons adjacent to the N−H (expected as triplet) group of 1-MPZ/protonated 1-MPZ 3 J = 4.5 methylene protons adjacent to the N−CH3 triplet group of 1-MPZ/protonated 1-MPZ singlet protons of methyl group of 1-MPZ carbamate/protonated 1-MPZ carbamate 3 J = 6.2 methylene protons adjacent to the N−CH3 triplet group of 1-MPZ carbamate/protonated 1-MPZ carbamate 3 J = 6.2 methylene protons adjacent to the N-CO2− triplet group of 1-MPZ carbamate/protonated 1-MPZ carbamate

δppm

species

43.3 44.1 44.6 50.2

methyl carbon of 1-MPZ/protonated 1-MPZ methyl carbon of 1-MPZ carbamate/protonated 1-MPZ carbamate carbamate ring C’s methylene C’s adjacent to the N−H group of 1-MPZ/protonated 1-MPZ methylene C’s adjacent to the N−CH3 group of 1-MPZ/protonated 1-MPZ methylene C’s adjacent to the N-CO2− group of 1-MPZ carbamate/ protonated 1-MPZ carbamate methylene C’s adjacent to the N-CO2− group of 1-MPZ carbamate/ protonated 1-MPZ carbamate carbamate carbon of 1-MPZ carbamate/protonated 1-MPZ carbamate C’s of carbonate/bicarbonate species

52.1 53.9 60.9 164.4 161.8 3619



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Figure 7. 1H NMR and 13C NMR for 1-MPZ solution in D2O for aqueous 1-MPZ (30 wt %) at 298.15 K and 6394 kPa.

Rochelle27 for aqueous 45 wt % 1-MPZ solution as at 313.15 K. The experimental data reported by Hilliard35 for 2 M aqueous PZ were also plotted in Figure 9 to compare with other calorimetric values published in the literature. The heat of absorption values for PZ were compared with the predicted values for 30 wt % aqueous 1-MPZ in Figure 9. It can be observed that the heat of absorption values for 1-MPZ was less than PZ for all temperatures and loadings. This may be due to the fact that 1-MPZ is made of one secondary (−NH) group and one tertiary amine (−N) group. All calorimetric experiments reveal the fact that tertiary amino (−N) groups have less heat of absorption than the primary (−NH2) and secondary amino (−NH) groups.38 Primary amines and hindered amines have the highest heats of absorption ranging from (76 to 86) kJ·mol−1 and (68 to 77) kJ·mol−1. The average

for regenerating the solvent from the CO2 loaded solution. Using a Gibbs−Helmholtz equation, the heat of absorption can be estimated from VLE data as ⎡ d ln PCO ⎤ −ΔHabs 2 =⎢ ⎥ R d(1/ T ) ⎣ ⎦α

CO2

(53)

In this work, the E-NRTL model was used to predict the PCO2 for different temperature ranges of (313.15 to 393.15) K using the regressed parameters. Figure 9 shows the calculated ΔHabs for 30 wt % aqueous 1-MPZ as a function of loading. The average −ΔHabs was estimated to be nearly 52 kJ·molCO2−1 at 313.15 K. The average −ΔHabs value of 67 kJ·molCO2−1 was estimated by Chen26 for aqueous 45 wt % 1-MPZ solution as at 313.15 K, and a value of 69 kJ·molCO2−1 was reported by Xu and 3620



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Figure 8. Predicted pH of aqueous 1-MPZ (30 wt %) and PZ (15 wt %) loaded with CO2.

Figure 9. Predicted heat of absorption values (−ΔHabs) for CO2 in aqueous 1-MPZ (30 wt %) and PZ (17 wt %) solutions.

ΔHabs values for PZ derivatives were estimated to be around −70 kJ·mol−1.26 ΔHabs values were mainly the contribution of the heat of carbamate formation and the heat of protonation reactions depending upon the type of amines reacting with CO2. Concerns in using this Gibbs−Helmholtz equation were discussed in the literature (Lee et al.39 and Kim and Svendsen40). Although the solubility data can be accurate to 2−5%, the predicted heat of absorption values have errors of (20 to 30) %.40 Heats of absorption obtained in this work are differential in loading, but integral in temperature. Therefore, calorimetric experiments are needed to confirm the predicted heat of absorption from the VLE data obtained in this work. Recently, Freeman et al.41 concluded that the use of aqueous 1-MPZ in a blend with PZ demonstrated noticeable levels of overall thermal degradation but degrade to useful products that can react with CO2 and maintain alkalinity of the solution.

5. CONCLUSION 1-Methyl piperazine (1-MPZ) is one of the structural analogues of piperazine (PZ) and has a higher reaction rate than conventional amines such as MEA. 1-MPZ has also the advantage of a lower heat of absorption (−55 kJ·mol−1 to −67 kJ·mol−1) compared to primary and secondary amines (−86 kJ·mol−1 to −120 kJ·mol−1). In this work, experimental data of VLE are reported for 1-MPZ at high pressures up to 7815 kPa. When compared to the same loading and temperature, PCO2 of 1-MPZ was sometimes about an order magnitude higher than that of PZ. The electrolyte nonrandom two-liquid (E-NRTL) model was used to correlate the experimental data within an acceptable average absolute deviations (for P and T with less than 2 % AAD and less than 7 % for the loading). The model was used to predict the speciation, pH, and heat of absorption of the solutions. The results were compared to experimental data and those for aqueous PZ. The model can be improved by 3621



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(2) Appl, M.; Wagner, U., Henrici, H. J.; Kuessnet, K.; Volkamer, F.; Ernst-Neust, N. Removal of CO2 and/or H2S and/or COS from gases containing these constituents. U.S. Patent 4336233, 1982. (3) Kadiwala, S.; Rayer, A. V.; Henni, A. High pressure solubility of carbon dioxide (CO2) in aqueous piperazine solutions. Fluid Phase Equilib. 2010, 292, 20−28. (4) McMurry, J. Organic Chemistry, 5th ed.; Brooks/Cole: Pacific Grove, 2000. (5) Rayer, A. V.; Sumon, K. Z.; Henni, A.; Tontiwachwuthikul, P. Kinetics of the reaction of carbon dioxide (CO2) with cyclic amines using the stopped-flow technique. Energy Proc. 2011, 4, 140−147. (6) Dugas, R.; Rochelle, G. T. Absorption and desorption rates of carbon dioxide with monoethanolamine and piperazine. Energy Proc. 2009, 1, 1163−1169. (7) Aroua, M. K.; Salleh, M. R. Solubility of CO2 in aqueous piperazine and its modeling using the Kent-Eisenberg approach. Chem. Eng. Technol. 2004, 27, 65−70. (8) Derks, P. W. J.; Dijkstra, B. S.; Hogendoorn, J. A.; Versteeg, G. F. Solubility of carbon dioxide in aqueous piperazine solutions. AIChE J. 2005, 8, 2311−2327. (9) Bishnoi, S.; Rochelle, G. T. Absorption of carbon dioxide into aqueous piperazine: reaction kinetics, mass transfer and solubility. Chem. Eng. Sci. 2000, 22, 5531−5543. (10) Pérez-Salado Kamps, A.; Xia, J.; Maurer, G. Solubility of CO2 in (H2O + piperazine) and in (H2O + MDEA + piperazine). AIChE J. 2003, 49, 2662−2670. (11) Ermatchkov, V.; Pérez-Salado Kamps, A.; Speyer, D.; Maurer, G. Solubility of carbon dioxide in aqueous solutions of piperazine in the low gas loading region. J. Chem. Eng. Data 2006, 51, 1788−1796. (12) Chakravarty, T.; Phukan, U. K.; Weiland, R. H. Reaction of acid gases with mixture of amines. Chem. Eng. Prog. 1985, 81, 32−36. (13) Nguyen, T.; Hilliard, M.; Rochelle, G. T. Amine volatility in CO2 capture. Int. J. Greenhouse Gas Control 2010, 4, 707−715. (14) Samanta, A.; Bandyopadhyay, S. S. Kinetics and modeling of carbon dioxide absorption into aqueous solutions of piperazine. Chem. Eng. Sci. 2007, 62, 7312−7319. (15) Dash, S. K.; Samanta, A.; Samanta, A. N.; Bandyopadhyay, S. S. Vapour liquid equilibria of carbon dioxide in dilute and concentrated aqueous solutions of piperazine at low to high pressure. Fluid Phase Equilib. 2011, 300, 145−154. (16) Bougie, F.; Iliuta, M. C. CO2 absorption in aqueous piperazine solutions: Experimental study and modeling. J. Chem. Eng. Data 2011, 56, 1547−1554. (17) Khalili, F.; Henni, A.; East, A. L. L. pKa values of some piperazines at (298, 303, 313 and 323) K. J. Chem. Eng. Data 2009, 4, 2914−2917. (18) Nguyen, T. Amine volatility in CO2 capture. Ph.D. Dissertation, University of Texas at Austin, Austin, TX, 2013. (19) Freeman, S. A.; Davis, J.; Rochelle, G. T. Degradation of aqueous piperazine in carbon dioxide capture. Int. J. Greenhouse Gas Control 2010, 4, 756−761. (20) Freeman, S. A.; Dugas, R.; Van Wagener, D. H.; Nguyen, T.; Rochelle, G. T. Carbon dioxide capture with concentrated aqueous piperazine. Int. J. Greenhouse Gas Control 2010, 4, 119−124. (21) Freeman, S. A.; Rochelle, G. T. Thermal degradation of piperazine and its structural analogs. Energy Proc. 2011, 4, 43−50. (22) Davis, J. Thermal degradation of aqueous amines used for carbon dioxide capture. Ph.D. Dissertation, University of Texas at Austin, Austin, TX, 2009. (23) Kennard, M. L.; Meisen, A. Mechanisms and kinetics of diethanolamine degradation. Ind. Eng. Chem. Fundam. 1985, 24, 129− 140. (24) Dawodu, O. F.; Meisen, A. Degradation of alkanolamine blends by carbon dioxide. Can. J. Chem. Eng. 1996, 74, 960−966. (25) Rayer, A. V.; Sumon, K. Z.; Henni, A.; Tontiwachwuthikul, P. Physicochemical properties of {1-methylpiperazine (1) + water (2)} system at T = (298.15 to 343.15) K and atmospheric pressure. J. Chem. Thermodyn. 2011, 43, 1897−1905.

providing better reference state parameters for aqueous 1-MPZ solutions. Due to its liquid state at ambient temperature, low heat of absorption, high reaction rate, and absorption capacity, compared to primary and secondary amines and other cyclic amines, 1-MPZ should be considered as a promising solvent for CO2 capture operations.

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AUTHOR INFORMATION

Corresponding Author

*Tel.: 306 585 4960. Fax: 306 585 4855. E-mail: amr.henni@ uregina.ca. Funding

The financial support of the Natural Sciences and Engineering Research Council of Canada (NSERC), Natural Resources Canada (NRCan), Canadian Foundation for Innovation (CFI), Petroleum Technology Research Centre (PTRC) and the International Test Center for Carbon Dioxide Capture (ITC, University of Regina) are gratefully acknowledged. Notes

The authors declare no competing financial interest.

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NOMENCLATURE A, B parameters in eqs 43 and 44 a, b, c, d coefficients for equilibrium constant in 8 ai activity of component of species i ds density of solvent (kg·m−3) E G excess Gibbs energy H Henry’s law constant (Pa) K equilibrium constant k Boltzmann constant (1.38065·10−23 (J·K−1) M molarity (kmol·m−3) NA Avogadro’s number (6.02205·1023) (1/mol) Exp Pi experimental pressure (kPa) PEst estimated pressure (kPa) i P0i vapor pressure (kPa) PCO2 equilibrium partial pressure of CO2 (kPa) Qe electron charge (1.60219·10−19) (C) R gas constant (8.314) (J·mol−1·K−1) ri Born radius (m) T temperature (K) vm molar volume (m3·mol−1) x mole fraction in liquid phase y mole fraction in vapor phase Z compressibility factor zi charge number of ion i αCO2 CO2 loading in liquid phase (molCO2·molamine−1) γ activity coefficient ρ closest approach parameter ε dielectric constant φ fugacity coefficient Subscripts and Superscripts

∞ infinite dilution * unsymmetric convention a, a′, a″ anion c, c′, c″ cation s solvent v vapor phase w water

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REFERENCES

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