Thermal Stability of Particle-Phase Monoethanolamine Salts

Article pubs.acs.org/est

Cite This: Environ. Sci. Technol. 2018, 52, 2409−2417

Thermal Stability of Particle-Phase Monoethanolamine Salts Xiaolong Fan,†,‡ Joseph Dawson,§,# Mindong Chen,† Chong Qiu,*,∥ and Alexei Khalizov*,‡,⊥ †

School of Environmental Science and Engineering, Nanjing University of Information Science and Technology, Nanjing, 210044, China ‡ Department of Chemistry and Environmental Science, New Jersey Institute of Technology, Newark, New Jersey 07102, United States § Department of Chemistry and Industrial Hygiene, University of North Alabama, Florence, Alabama 35632, United States ∥ Department of Chemistry and Chemical Engineering, University of New Haven, New Haven, Connecticut 06516, United States ⊥ Department of Chemical, Biological, and Pharmaceutical Engineering, New Jersey Institute of Technology, Newark, New Jersey 07102, United States S Supporting Information *

ABSTRACT: The use of monoethanolamine (MEA, 2hydroxyethanamine) for scrubbing of carbon dioxide from combustion flue gases may become the dominant technology for carbon capture in the near future. The widespread implementation of this technology will result in elevated emissions of MEA to the environment that may increase the loading and modify the properties of atmospheric aerosols. We have utilized experimental measurements together with aerosol microphysics calculations to derive thermodynamic properties of several MEA salts, potentially the dominant forms of MEA in atmospheric particles. The stability of the salts was found to depend strongly on the chemical nature of the acid counterpart. The saturation vapor pressures and vaporization enthalpies obtained in this study can be used to evaluate the role of MEA in the aerosol and haze formation, helping to assess impacts of the MEA-based carbon capture technology on air quality and climate change.

■

INTRODUCTION

Over the past decade, amines have been recognized as a major player in the formation and evolution of atmospheric particles.15−19 Amines may enhance the nucleation of aerosol and can be incorporated in atmospheric condensed phases through direct dissolution and also through acid-neutralizing reactions, as moderately strong bases.17,20−22 The salts formed in reactions of amines with atmospheric acids (i.e., aminium salts) often have lower volatility than parent chemicals, promoting the formation of secondary organic aerosols (SOA).23−26 Some aminium salts showed enhanced hygroscopicity at low relative humidity (RH), promoting growth of particles and exacerbating their environmental impacts.23,27 Previous studies have shown that thermodynamic stability of aminium salts depends on the chemical nature of both the amine and acid counterparts.28 Among aminium salts formed by different acids, nitrates are the least stable, with vaporization enthalpy ΔHvap ∼ 54−74 kJ mol−1 and saturation vapor pressure psat (298 K) ∼ 10−4 Pa.29 The most stable salts are sulfates (ΔHvap ∼ 114−168 kJ mol−1 and psat (298 K) ∼ 10−9− 10−12 Pa),30 with volatilities comparable to or lower than that of ammonium sulfate, a common constituent of atmospheric aerosols. The stability of monocarboxylates is comparable to

The combustion of fossil fuels dominates global energy generation, serving as a major anthropogenic emission source of CO2. Rapidly increasing concentration of CO2, which is a major driver of climate change, has become a prominent global environmental challenge.1,2 carbon capture and storage (CCS) technology is one way to reduce global warming driven by CO2 emitted from coal-fired power plants.3−6 According to the worldwide least-cost assumptions from International Panel on Climate Change,7 it is possible that CCS will remove 60−600 Gt-C of cumulative anthropogenic CO2 emissions by 2100 and the concentration of this atmospheric greenhouse gas may be stabilized between 450 and 750 ppm. Among various CCS technologies, chemical absorption by amine solution is a well-established method.6,8−10 Currently, monoethanolamine (MEA, 2-hydroxyethanamine) is widely utilized to extract CO2 from combustion flue gas.11−13 In this approach, the gas is exchanged with chilled aqueous MEA solution in an absorber tower, whereby CO2 reacts with MEA to form carbamate.14 The resulting solution is pumped into a stripper tower where the temperature is increased and the carbamate product is decomposed back to CO2 and MEA.1 During the CO2 capture process, MEA is gradually lost via thermal and chemical degradation and evaporation.4 Evaporative losses from widespread use of MEA may result in significant emissions of this amine into the environment. © 2018 American Chemical Society

Received: Revised: Accepted: Published: 2409

December 10, 2017 January 23, 2018 January 25, 2018 January 25, 2018 DOI: 10.1021/acs.est.7b06367 Environ. Sci. Technol. 2018, 52, 2409−2417

Article

Environmental Science & Technology Table 1. Physical Properties of MEA Salts

a

MEA salt

abbr.

molar mass, g mol−1

appearance

acetate hydrogen oxalate oxalate succinate glutarate adipate dihydrogen citrate hydrogen citrate citrate sulfate benzoate

EaAc EaHOx Ea2Ox Ea2Su Ea2Gl Ea2Ad EaH2Ct Ea2HCt Ea3Ct Ea2SO4 EaBz

121 151 212 240 254 268 253 314 375 220 183

oily solid powder powder crystalline oil oil solid oil oil crystalline powder

surface energy,a J m−2 0.16 ± 0.07; 0.08 ± 0.05; 0.10c; 0.14 0.03 ± 0.01; 0.03 ± 0.01;

density, g cm−3

diffusion coeff. (298 K),b cm2 s−1

1.23 1.47 1.34 1.31 1.05 1.02 1.42 1.42 1.35 1.38 1.31

0.07 0.10c 0.05c 0.05c

0.108 0.108 0.095 0.091 0.088

The first value is calculated using Model I and the second value using Model II. bDetermined based on ref 68. cEstimated based on ref 29.

diffusion charger (Po-210, 400 μCi, NRD Staticmaster). The DMAs were operated in recirculating sheath flow mode with a sample flow of 1 L per minute (lpm) and a sheath-to-sample flow ratio of 10. Typical initial particle mobility diameters were 75, 100, and 150 nm. The size-classified aerosol was sent to the DMA2, either directly or through a TD, to measure initial and final particle mobility diameters, Dp,i and Dp,f, respectively, which were used to calculate the particle growth factor, Gfd = Dp,f/Dp,i. The TD was a 100 cm long, 0.5 in. (1.27 cm) outer diameter (O.D.) stainless steel tube, maintained at a preset temperature (298−498 K) using a heating tape connected to a PID temperature controller. The temperature was measured in situ inside the TD, using a 0.125 in. O.D. sheathed thermocouple. The accuracies in Gfd and temperature measurements were 0.3% and 1.5 K, respectively. Data Processing. For aerosol in the transition region (0.02−3 μm at atmospheric pressure), the particle evaporation rate is limited by vapor diffusion and can be described by the Maxwell equation,

that of ammonium sulfate, with an estimated vapor pressure psat (298 K) ∼ 10−9 Pa.16,31 The stability of dicarboxylates falls in between nitrates and sulfates (ΔHvap 73−134 kJ mol−1 and psat (298 K) ∼ 10−6 Pa).31 Among different alkyl-substituted amines, mono- and di- methyl/ethyl amines form most stable salts because of an interplay between basicity and alkyl substituent steric effects.23 The presence of the hydroxyl group in MEA makes it a weaker base (pKb is 4.5) than the commonly found atmospheric amines (pKb 3.3−4.2). However, despite the lower basicity of MEA, the stability of its nitrate is comparable to that of alkylamine nitrates.29 The stability of MEA salts with other atmospherically relevant acids, including sulfuric, monocarboxylic, and dicarboxylic acids, is still unexplored, hindering evaluation of the contribution of MEA to the formation of atmospheric aerosols. We have synthesized several representative salts of MEA with mono-, di-, and tricarboxylic acids and also with sulfuric acid (Table 1). A combination of experimental and computational approaches based on aerosol microphysics was then utilized to measure the evaporation rates of these salts and calculate their thermodynamic properties, followed by discussion of findings and atmospheric implications.

dDp

■

dt

EXPERIMENTAL SECTION Evaporation Rate Measurements. The synthesis of the MEA salts (Table 1) is described in the Supplementary, with purity of the products generally greater than 98% based on NMR spectroscopic measurements. To minimize decomposition, the synthesized salt samples were stored refrigerated. Particle evaporation rates were measured using the tandem differential mobility analyzer (TDMA) technique.32,33 In a TDMA, size-classified aerosol particles emerging from the first DMA1 (TSI 3081) pass through a thermal denuder (TD) maintained at a preset temperature and then their final size is measured by a second DMA in combination with a condensation particle counter (CPC, TSI 3772) (Supporting Information (SI) Figure S1).34 Polydisperse aerosol was generated by atomizing aqueous solutions of MEA salts (1.7−2.3 g/L), using a constant output atomizer (TSI 3079) Particle-free dry air was used to dilute the aerosol flow to control the aerosol number density and relative humidity. The aerosol flow was further dried using a Nafion dryer (Perma Pure, PD-07018T-24MSS) and a silica gel diffusion dryer. The RH, as measured using a Vaisala HUMICAP humidity module (HMM100), was below 7% in most experiments and below 10% in all experiments. Before entering the DMA1, the aerosol was brought to an equilibrium charge distribution with a

=−

⎛ 4γM ⎞ i ⎟⎟f (Kni , α) psat, i exp⎜⎜ i ρi DpRT ⎝ ρi DpRT ⎠

4Di ,air Mi

(1)

where R is the gas constant, T is the temperature, t is the evaporation time, Dp is the particle diameter and Di,air is the diffusivity of the molecule i in air, psat,i is the saturation vapor pressure, Mi, ρi, and γi are the molar mass, density, and surface free energy, respectively. Several assumptions are made when using this equation, including that the particles are spherical, the surface energy is isotropic, and the contributions from the latent heat of evaporation and the far-field vapor pressure are negligible.35 Many of the synthesized MEA salts were ionic liquids that form spherical particles (Table 1), making the first assumption accurate. The latent heat effects are small for particles in the studied size range.36 During evaporation measurements, the particle concentration was kept at a sufficiently low level, 100−1000 cm−3, depending on the particle size, so that the far-field partial pressure never exceeded 1−2% of the determined saturation vapor pressure. The transition correction f(Kni, α) for particle diameters was estimated according to Fuchs and Sutugin,37 f (Kni , α) = 2410

1 + Kni 1 + 0.3773Kni + 1.33Kni

1 + Kni α

(2)

DOI: 10.1021/acs.est.7b06367 Environ. Sci. Technol. 2018, 52, 2409−2417

Article

Environmental Science & Technology

Figure 1. Measured growth factor (Gfd) for size-classified aerosol particles composed of ethanolammonium acetate (EaAc, a), sulfate (Ea2SO4, b), benzoate (EaBz, c), hydrogen oxalate (EaHOx, d), oxalate (Ea2Ox, e), succinate (Ea2Su, f), glutarate (EaGl, g), adipate (Ea2Ad, h), dihydrogen citrate (EaH2Ct, i), hydrogen citrate (Ea2HCt, j), and citrate (Ea3Ct, k). Black squares, red circles, and blue triangles correspond to particles with initial mobility diameters of 75, 100, and 150 nm, respectively. Thermal profiles of 100 nm particles for all MEA salts investigated here are aggregated in panel (l). Thermal profiles for several aminium salts in ref 31 are plotted for reference, including monomethylammonium acetate (MMA-AC, a), trimethylammonium sulfate (TMAS, b), and ammonium sulfate (AS, b).

where α is the accommodation coefficient that was set to unity, similarly to previous work,38,39 and Kni is the Knudsen number of molecule i,

Kni =

2λi Dp

proton transfer reaction mass spectrometry analysis to show that a significant fraction of particle-phase alkylamine dicarboxylates is transferred to the gas-phase in molecular form. Accordingly, we calculated the diffusion coefficient assuming that the MEA salts vaporized without thermal breakdown. Even when the salt was allowed to dissociate in the model into acidic salt and free MEA, the resulting error in the saturation pressure was significantly lower than the error associated with the uncertainty in the evaporation time.

(3)

In eq 3, λi is the mean free path of molecule i in air. The diffusivity of species i in air Di,air is estimated by an empirical correlation,40 which yields values similar to that based on the method described in Rader et al.,33 but permits an easier computer coding. Recently, Lavi et al.31 have utilized the 2411


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Environmental Science & Technology

calculated and measured growth factors for individual experiments (χ2 in eq 6) was minimized by varying ΔHvap, C, and γ.

If the temperature is constant along the TD, eq 1 can be rearranged into eq 4 and integrated over time to obtain the saturation vapor pressure, psat, at each set temperature. psat (T ) = −

ρi RT 4Di ,air

∫ M Δt D i

Dp, f

p, i

χ2 =

⎛ 4γM ⎞ Dp i ⎟dDp exp⎜⎜ i f (Kni , α) ⎝ ρi DpRT ⎟⎠

D,T

(6)

Initial guess values of ΔHvap, C, and γ were obtained from previous studies or using Model I. Multiple Model II runs were performed, using a range of guess values to ensure that the model converged to a unique solution. In the cases when the model failed to converge, the γ was fixed at a value estimated from Model I, and only ΔHvap and C were allowed to vary.

(4)

By fitting the pressure and temperature data to the Clausius− Clapeyron eq 5, enthalpy of vaporization ΔHvap can be determined. ⎞ ⎛ ΔH vap psat (T ) = exp⎜ − + C⎟ ⎠ ⎝ RT

exp 2 ∑ (Gf dcalc D , T − Gf d D , T )

■

RESULTS AND DISCUSSION Aerosol Evaporation Profiles. The composition of the investigated MEA salts is listed in Table 1. In the evaporation experiments, aerosol particles of several initial diameters were passed through the TD maintained at a predefined exit temperature, and the change in the particle size (Gfd) was measured. The temperature dependencies of Gfd for these salts are summarized in Figure 1, along with the literature data for several similar salts.31 As shown previously,44 for singlecomponent particles, the rate of evaporation is controlled by the gas-phase diffusion of the vaporized particle material. The particle size remained nearly constant at lower temperatures, and only above the critical temperature, the particles began to shrink at a progressively faster rate, showing a distinct size dependence (SI Figure S5). In our study, MEA salts of dicarboxylic acids, including oxalate, hydrogen oxalate, succinate, glutarate, and adipiate, showed this type of simple vaporization behavior. Other MEA salts, including sulfate, acetate, benzoate, dihydrogen citrate, hydrogen citrate, and citrate, displayed more complex evaporation profiles. For instance, the MEA acetate particles shrank considerably already at lower temperatures, switching sharply to a less steep temperature dependence above ∼370 K. For all other salts, the first region was nearly linear, lacking a pronounced size dependence; in the second region, the particles shrank at a progressively faster rate with a clear size dependence. The transition between the two regions was at ∼370 K for citrate and ∼430 K for benzoate and sulfate. Evaporation profiles that fail to follow the simple diffusion-limited mechanism have been reported previously, but the cause of this behavior has remained poorly understood.31 The lack of a clear size dependence may indicate that the rate of the particle size change is controlled by in-particle processes rather than by the vapor diffusion. One such process could be the evaporation of water trapped inside highly viscous or amorphous solid particles or the partial evaporation of MEA to form acidic salts, as reported previously for alkylaminium sulfates.45 For instance, at lower temperatures, the hydrogen citrate and citrate may lose MEA, forming dihydrogen citrate, which at higher temperatures decomposes into MEA and citric acid. Finally, in addition to the reversible thermal dissociation into the MEA and acid (or acidic salt) counterparts, various irreversible intramolecular reactions may also take place inside the particles. For monocarboxylic acids, the pyrolysis of aminium salts at temperatures above ∼370 K causes reversible retrogradation, which is followed by nucleophilic attack of the amine on the carbonyl group of the acid, resulting in the water elimination and amide formation.46,47 Amine salts with 2,4dichlorophenoxyacetic acid are a good example,48 where amides are formed already at 353 K, peaking at 413−443 K, and

(5)

To integrate eq 4, the initial and final diameters (Dp,i and Dp,f) were taken directly from TDMA measurements. The material density of MEA salts, ρi, was measured using the aerosol particle mass analyzer against polystyrene latex nanospheres, as described previously.41,42 The only parameter not readily available was the surface free energy. Some studies relied on an estimated γ,29,31,43 while others35 derived γ along with psat using size-dependent evaporation experiments. When possible, we followed the latter approach, introduced by Tao and McMurry,35 which involved experiments on particles of several initial diameters (75, 100, and 150 nm) to solve eq 4 for a range of surface energies (0.005−0.200 J m−2), as shown in SI Figure S2. The vapor pressure versus surface energy curves must cross at a single point (or within a certain region due to experimental errors), corresponding to the true vapor pressure and surface energy at a chosen temperature. This point was found by the least-squares minimization of the standard deviation of the saturation vapor pressure between particles of different initial sizes, using an automated procedure implemented in MATLAB (Mathworks, Inc.).The procedure derived the saturation vapor pressure and surface tension from the experimental data by numerical integration of eq 4. When no suitable minimum satisfying all sizes was found for a particular salt, an estimated surface energy was used, as in previous studies.29 The procedure described above will be referred to as Model I. Solving eq 4 requires the knowledge of the residence time of aerosol, Δt, which can be calculated from the flow rate measured at the aerosol inlet and further corrected by the TD temperature. Although it is often assumed that the gas flow attained the set temperature immediately after entering the TD, our measurements indicate that the gas temperature inside the TD increased gradually, following a sigmoidal profile (SI Figure S3). To account for this nonconstant temperature, we developed a more advanced numerical model that treated the axial temperature profile in the TD explicitly, following the approach of Earl et al.36 The model (referred in the following as Model II) consisted of two components to perform aerosol microphysics and minimization calculations sequentially. The microphysics component was represented by eq 1, where the temperature-dependent psat was defined by eq 5 of the two parameters, ΔHvap and C. Beginning with the initial particle diameter Dp,i, the microphysics component solved the timedependent differential eq 1, along with associated algebraic equations for all temperature-dependent parameters, subject to the selected axial temperature profile, producing Dp,f. The calculation was conducted simultaneously for all experimentally assessed initial particle diameters and temperature profiles. The normalized sum of squares of differences between the 2412


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Environmental Science & Technology Table 2. Enthalpy of Vaporization of MEA Salts Model IIb

Model I chemical

a

EaAc EaBz oxalic acid EaHOx Ea2Ox succinic acid

−1

ΔHvap, kJ mol

e

T1/2, K

f

Psat , Pa

−1

ΔHvap, kJ mol

Psatf, Pa

372 407 343 357

Ea2Su glutaric acid Ea2Gl adipic acid

315

Ea2Ad EaH2Ct Ea2HCt Ea3Ct Ea2SO4 EaNt MMA-ACi DMA-ACi (NH4)2SO4 MEA

313 392 387 386 443 345 432 459 393

312

10−2 (s); 2.9 × 10−2 (l)g 10−5 10−6 10−5 (s); 1.3 × 10−3 (l)g 10−5h 10−4i 10−4c; 1.9 × 10−3d 10−4 10−4 (s), 1.0 × 10−3 (l)g 10−4 10−5 (s), 6.1 × 10−5 (l) 10−5h 10−4i 10−5c; 6.9 × 10−4d 10−4

91 ± 9 (s); 79 ± 15 (l)g 88 ± 3 96 ± 6 115 ± 15 (s); 105 ± 1 (l)g 112 ± 12h 88 ± 3i 82 ± 3c; 81 ± 4d 68 ± 5 130 ± 11 (s); 100 ± 5 (l)g 69 ± 4 131 ± 18 (s); 94 ± 5 (l)g 97 ± 8h 135 ± 13i 92 ± 4c; 87 ± 4d 88 ± 1

1.4 2.0 3.2 7.7 6.4 3.7 2.1 2.7 1.7 4.3 1.9 5.8 3.7 6.4 2.3

74.0h

8.9 × 10−5h

100i 88 ± 3l

1.0 × 10−7i 35.3l

× × × × × × × × × × × × × × ×

83.5 100.0

2.9 × 10−6 2.3 × 10−7

88c; 87d 67

2.4 × 10−5c; 2.1 × 10−4d 3.2 × 10−5

73

2.7 × 10−5

96c; 92d 91

8.2 × 10−6c; 8.2 × 10−5d 1.6 × 10−5

Unless specified otherwise, see Table 1 for the definitions of abbreviations. bRecommended values. cUsing accommodation α = 1. dUsing accommodation α = 0.1. eEstimated temperature when the particle loses half of its initial volume. fAt 298 K. gRecommended values from ref 28. h Values from ref 29. iValues from ref 57. jBased on the results in ref 29.: EaNt, monoethanolamine nitrate. kBased on the results in ref 31.: MMA, monomethylamine; DMA, dimethylamine. lBased on ref 69.; MEA, monoethanolamine. a

and comparable to that of the previously studied MEA nitrate.29 The MEA sulfate had a T1/2 of 443 K, higher than that of ammonium sulfate (T1/2 = 393 K) and comparable to that of methylaminium acetates (T1/2 = 432−459 K).31 The three MEA citrates showed comparable T1/2 in the range 386−392 K, On the contrary, there was a noticeable difference (∼14 K) in T1/2 between the oxalate and hydrogen oxalate. Saturation Vapor Pressure and Vaporization Enthalpy. To test the performance of our models, we performed calculations using Gfd measured for pure succinic and adipic acids (SI Figure S5). These chemicals have well-documented vaporization enthalpies and their particles showed simple evaporation profiles. Only the experimental data in the range of 0.3 < VFR < 0.9 were processed, following previously reported approach.29,31 The upper bound was set to exclude the possible contribution from the evaporation of residual water. The lower bound was set to discard the heavily processed particles, which could be enriched in low-volatility products generated by the irreversible decomposition or from residual contaminants present in water.44 In Model I calculations, two different assumptions were made regarding the residence time of the aerosol. Under the first assumption, the aerosol attained the set temperature immediately and the particles were evaporating throughout the entire length of the TD. Under the second assumption, the evaporation of particles began at the point corresponding to the e-folding time of the saturation ratio, p(T)/p(Tmax), as shown in SI Figure S3. The p(T) and p(Tmax) were calculated by eq 5, using available from literature or

decomposing completely above 473 K. In addition to amides, other compounds, including imines and lactones, are observed above ∼430 K.48 Relatively simple inorganic salts also can undergo irreversible decomposition on heating, e.g., ammonium sulfate in the first reversible decomposition step produces ammonium bisulfate, which subsequently decomposes irreversibly into ammonium pyrosulfate, sulfamic acid, sulfur dioxide, water, and molecular nitrogen.49−51 Even in the case of dicarboxylic acid salts an additional pathway leading to the formation of the cyclic anhydride is possible at high temperatures.52 The gas-phase and particle-phase decomposition products of MEA salts may be identified by collecting particles after VTDMA measurements and analyzing their chemical compositions with mass spectrometry. To compare thermal stabilities of different salts, we converted Gfd profiles to volume fraction remaining (VFR) profiles (SI Figure S4), using eq 7. VFR = Gf d3 = (Dp, f /Dp, i)3

(7)

From VFR profiles, we calculated, T1/2, the temperature at which the particles lost half of their initial volume (Table 2). The table also lists T1/2 estimated from the VFR values reported previously for MEA nitrate,29 ammonium sulfate,30 and alkylaminim acetates.31 Overall, MEA salts with complex evaporation profiles showed a higher T1/2 than the salts with simple evaporation profiles. Glutarate, succinate, and adipate were the least stable salts of MEA (T1/2 = 312−315 K). The stability of MEA oxalate was somewhat higher (T1/2 = 357 K), 2413


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Environmental Science & Technology estimated ΔHvap. In Model II calculations, the temperature variation along the TD was treated explicitly. It should be noted that after exiting the TD, the aerosol temperature and hence p(T)/p(Tmax) dropped rapidly because of a large temperature gradient, terminating the particle evaporation. Thus, explicit treatment of the cooling temperature branch was not required. Clausius−Clapeyron plots derived from the evaporation curves of succinic and adipic acids by Model I show an excellent linearity (SI Figure S6). The associated values of ΔHvap and psat(298 K) determined using two different residence time assumptions are provided in SI Table S1. It can be seen that psat(298 K) derived under the entire TD length assumption is a factor of 2 lower that the value obtained using the shorter residence time, in full agreement with eq 4, which states that the saturation vapor pressure scales inversely with the residence time inside the TD. However, the use of different residence times had little impact on the derived values of ΔHvap, which differed by 0.9 and 0.2 kJ mol−1 for succinic and adipic acids, respectively. Model II calculations produced a factor of 5−10 lower psat(298 K) and 1.7−5.5 kJ mol−1 higher ΔHvap than Model I with the shorter residence time. The discernible difference in psat(298 K) of the reference materials predicted by Model I and II is caused by the exponential amplification of small errors in ΔHvap and C when extrapolating the saturation vapor pressures from experimental temperatures down to 298 K. However, the values of ΔHvap from Model I and II are in agreement with each other, with relative differences of 2−9%. The values obtained for reference materials using both models are in agreement with the data from previous studies. For instance, psat(298 K) and ΔHvap of adipic acid from Model I (shorter residence time) are 6.4 × 10−5 Pa and 92 kJ mol−1, respectively, whereas the reported values are (1.9−76) × 10−5 Pa and (89−105) kJ mol−1.28 The significant scatter in the literature data can be traced down to the use of different experimental techniques, different temperature intervals, the presence of water, and the phase state of particles.28 For example, in most VTDMA-based measurements aerosol samples were dried at RH < 10%; based in thermodynamic considerations, this humidity is sufficiently low to solidify particles generated by atomizing organic acid solutions. However, in some cases the particles have insufficient time to solidify on time scale of experiment due to kinetic limitation.53 Another factor that could affect the calculated vapor pressure, and hence influence the determined ΔHvap and psat(298 K) is the accommodation coefficient α in eq 2.54−57 Currently, the information on accommodation coefficients for different chemicals is sparse. Results in Table 2 show that using α = 0.1 rather than α = 1 resulted in a 5% or lower decrease in ΔHvap, but up to ten times increase in psat (298 K). We thus conclude that future measurements of the accommodation coefficient may further improve the accuracy of the saturation vapor pressure determination, but will have little impact on the vaporization enthalpy. Overall, according to the sensitivity analysis performed for a similar experimental system by perturbing each independent variable and calculating the response29 and excluding the contributions from accommodation coefficient and residence time, the maximum relative error in the saturation vapor pressure is under 20%. We applied Model I (shorter residence time) and Model II to derive the psat and ΔHvap from the evaporation profiles of MEA hydrogen oxalate, oxalate, succinate, glutarate, and adipiate (Table 2 and Figure 2). The literature values for pure acids are provided for comparison. It can be seen that ΔHvap and psat(298

Figure 2. Temperature dependence of the saturation vapor pressure of MEA salts. Red circles are for hydrogen oxalate (EaHOx), blue downpointing triangles are oxalate (Ea2Ox), orange up-pointing triangles are for succinate (Ea2Su), purple diamonds are for glutarate (EaGl), and brown squares are for adipate (Ea2Ad). The solid lines are linear fits to the experimental data.

K) generally agree within 3−9% and a factor of 5−16, respectively between the values from Model I and II. Since most of the reported literature values from TD aerosol measurements have been derived using an approach similar to Model I, we use the results from the latter in our discussion. Table 2 shows that ΔHvap spans from 69 kJ mol−1 for glutarate to 96 kJ mol−1 for oxalate. The associated psat(298 K) varies from 4.3 × 10−4 to 3.2 × 10−6 Pa. Most MEA dicarboxylates (except for oxalate) show a lower ΔH vap than the corresponding acids because of the disruption of the strong hydrogen-bonding network of dicarboxylic acids when organic salt is formed. In salts with strong bases, such as alkylamines, the disruption is compensated by the ion pair formation, resulting in a higher ΔHvap. For instance, ΔHvap of methylamine succinate and adipate are both 134 kJ mol−131 versus ∼70 kJ mol−1 for the corresponding MEA salts. The corresponding psat(298 K) are 4.2 × 10−5 and 5.5 × 10−6 Pa for methylamine salts31 and 2.7 × 10−4 and 2.3 × 10−4 Pa for MEA salts. It is notable that an opposite trend has been reported for nitric acid salts, with MEA nitrate (74 kJ mol−1 and 8.9 × 10−5 Pa) being somewhat more stable than small alkylamine nitrates, 54−66 kJ mol−1 and (2.7−5.4)×10−4 Pa.29 The reason for this trend is unclear at the present; the salt stability may depend on the phase change of particles with the change in RH.58 When applicable, surface energies γ of several MEA salts were also obtained from Model I and/or II. Solid MEA salts (e.g., oxalate) generally showed higher γ values (0.1−0.2 J m−2). Liquid MEA salts (e.g., adipate and glutarate) exhibited lower surface tension (0.03 J m−2), similar to those of long-chain aliphatic organic acids.35 We believe that Model II yields results that are more accurate because it explicitly considers the thermal profile. Since most of the ΔHvap and psat(298 K) reported in the literature have been derived assuming a unity mass accommodation coefficient, we recommend the values listed in Columns 5 and 6 of Table 2, obtained using α = 1, until better knowledge of α for MEA salts becomes available. The data in Table 2 can be used to inspect stability trends in MEA salts involving dicarboxylic acids of different carbon chain length and dicarboxylic acids of different degree of acidic hydrogen substitution. Among the dicarboxylic acid salts, oxalate shows the largest ΔHvap of 96 kJ mol−1 because oxalic acid is the strongest in the series (pKa1 1.23).59 The weaker 2414


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Environmental Science & Technology succinic acid (pKa1 4.21)59 forms a salt with a lower ΔHvap of 68 kJ mol−1, which is comparable to the ΔHvap (69 kJ mol−1) of the MEA salt of glutaric acid with a similar pKa1 of 4.34. However, ΔHvap of MEA adipiate is 88 kJ mol−1, higher than those of succinate and adipate despite the lowest acidity of glutaric acid (pKa1 4.42),59 suggesting that increasing molecular weight of these acids outweighs the acidity contribution and enhances the thermal stability of MEA dicarboxylates. The MEA oxalate (ΔHvap 96 kJ mol−1) is more stable then MEA hydrogen oxalate (88 kJ mol−1) because of a significant difference in pKa1 and pKa2 (1.23 and 4.19).59 Atmospheric Implications. Organic acids enter the atmosphere from direct emissions and through photochemical oxidation of volatile organic compounds.60 According to their saturation vapor pressure,28 most organic acids can be considered semivolatile under typical atmospheric aerosol loadings.61 Indeed, the detection of 5−45% backup filter fractions for 0.3−15.4 ng m−3 filter loadings62 indicates that many organic acids can be present comparably in gas and particle phases. Important factors in the stabilization of organic acids in the particle phase include the mixing with other chemicals (including hydration and formation of nonideal mixtures) and particle-phase reactions to form less volatile or higher molecular weight accretion products.17,63 Among those reactions, neutralization of acids by ammonia and amines is a major process. These two chemicals, originating from a variety of anthropogenic and biogenic sources,64 can react with acids to form stable salts across a broad range of particle sizes.17,23 For most volatile acids, such as formic and acetic, neutralization is probably the most important process controlling their gas-toparticle transfer.16 A semiquantitative evaluation of the MEA contribution to the aerosol formation by acid neutralization can be performed by comparing the vapor pressure of MEA salts against that of a model compound, which exists in both gas and particle phases under typical atmospheric conditions. For instance, in the presence of an aerosol with a mass loading of 1 μg m−3, a compound with a saturation vapor pressure of 10−5 Pa (or a vapor mass concentration of ∼1 μg m −3 ) would be equipartitioned between gas and particle phases.63 Accordingly, if a substance has a vapor pressure of 10−4 and 10−6 Pa, its particle-phase fraction will be 10 and 90%, respectively. Based on these criteria and our results from Model I, only MEA oxalate is sufficiently stable to exist predominantly in the particle-phase at 298 K. The stability of MEA hydrogen oxalate, corresponds to MEA nearly equally partitioned between gas and particle phases. For succinic, glutaric, and adipic acids, only a small fraction of MEA will be present in the particles at room temperature, but the equilibrium may shift significantly toward higher MEA content in the particles at lower temperatures. An alternative way to evaluate the contribution of MEA to the aerosol loading is by considering whether the formation of salts would result in a relative stabilization of organic acids in the particle phase, i.e., whether the saturation vapor pressure of the salt is lower than that of the acid. It is clear that the formation of MEA oxalates will stabilize the particle-phase oxalic acid by a factor of 103−104, regardless if the oxalic acid is solid or subcooled liquid. This may be significant considering that oxalic acid is one of the most abundant organic acids in the atmosphere. Notably, MEA oxalate is thermally more stable than hydrogen oxalate, whereas particle-phase trimethylaminium sulfate is not thermodynamically stable and decomposes partially into hydrogen sulfate.58,65 Adipic acid is also stabilized,

regardless of its phase state. On the contrary, succinic and glutaric acids are only stabilized (by a factor of 10) relative to subcooled state, yet solid forms of these two acids will be slightly destabilized by the formation to MEA salts. Atmospheric aerosols are complex internal mixtures of inorganic and organic components. The presence of nonreactive organics, such as oleic acid66 and sucrose67 can enhance the uptake of amines by forming aminium salts. Further rigorous evaluation of the MEA contribution to the aerosol formation can be performed by explicit numerical modeling of the acid neutralization and water uptake processes in the presence of representative organic/inorganic aerosol constituents, using physical parameters derived in our work.

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ASSOCIATED CONTENT

S Supporting Information *

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.est.7b06367. Synthesis of the MEA salts, schematic of the TDMA system, original Gfd data, processing parameters, and plots of the volume fraction remaining of the reference materials and MEA salts (PDF)

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AUTHOR INFORMATION

Corresponding Authors

*(C.Q.) E-mail: [email protected]. *(A.K.) E-mail: [email protected]. ORCID

Alexei Khalizov: 0000-0003-3817-7568 Present Address #

(J.D.) Department of Chemistry, Pennsylvania State University, University Park, Pennsylvania 16802, United States. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS This work was supported by the National Science Foundation (AGS 1463702 (A.K.) and 1463703 (C.Q.)), startup funds from New Jersey Institute of Technology (A.K.) and Tagliatela College of Engineering at University of New Haven (C.Q.), the Faculty Research Grant (#42956, C.Q.) of the College of Arts and Sciences and Undergraduate Research Program (Award #45091, J.D.) at the University of North Alabama, and by Nanjing University of Information Science and Technology and China Scholarship Council (X.F.).

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REFERENCES

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