Temperature and Pressure Effects on the Separation Efficiency and

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Cite This: Ind. Eng. Chem. Res. 2019, 58, 12247−12252

Temperature and Pressure Effects on the Separation Efficiency and Desorption Kinetics in the NH3−CO2−H2O System Georgios Kolliopoulos and Vladimiros G. Papangelakis* Department of Chemical Engineering and Applied Chemistry, University of Toronto, 200 College Street, Toronto, Ontario, Canada M5S 3E5

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S Supporting Information *

ABSTRACT: The present work is a continuation of a previous study on the separation efficiency and desorption kinetics of thermolytic draw solutions in forward osmosis (FO). The aim of the present work was to study the combined effects of temperature (30, 40, and 50 °C) and pressure (0.75, 0.55, and 0.35 bar) on the separation efficiency of the draw solutes from an aqueous carbonated ammonia (NH3−CO2− H 2 O) draw solution and compare them to aqueous carbonated trimethylamine (TMA−CO2−H2O). In the present work, CO2 proved to desorb faster from solution compared to NH3 with an observed NH3:CO2 average molar ratio in solution of 1.71, ranging from 1.07 to 2.64. The separation process resulted in almost 90% draw solute removal after 90 min at 50 °C, 0.35 bar. The draw solute desorption kinetics were studied based on the overall concentration of total dissolved NH3 and CO2 in the draw solution. According to the proposed mechanism, the decomposition of NH4[NH2CO2] in the aqueous phase proved to be the rate-limiting step of the separation process. The chemistry of the NH3−CO2−H2O system is welldescribed in the literature.22−28 Liu et al.24 suggest that the overall CO2 absorption in aqueous NH3 proceeds as

1. INTRODUCTION Forward osmosis (FO) is a spontaneous process that recovers water indirectly. First, a concentrated draw solution (CDS) spontaneously draws water from a feed solution that is more dilute.1−4 The resulting diluted draw solution (DDS) is subsequently treated to recover the draw solutes, leaving behind fresh water. An ideal draw solution generates high osmotic pressure values to maximize the water recovery yield. Also, it is easily separable and recoverable from the resulting diluted draw solution (DDS). The main source of energy consumption in FO is during the separation and the regeneration of the draw solution.5 Thermolytic salts undergo a phase change from aqueous to gaseous upon mild heating and are promising draw solute candidates for FO.2,6 Aqueous carbonated ammonia (NH3− CO2−H2O)2,3,7,8 and aqueous carbonated trimethylamine (TMA−CO2−H2O) have been studied as draw solutes for FO.6,9−14 The TMA−CO2−H2O system results in similar water flux and lower reverse flux compared to the NH3−CO2− H2O system, while the separation of TMA and CO2 has been estimated to be 36% more energy efficient compared to the separation of NH3 and CO2.6 However, the high volatility and the low odor threshold of TMA (0.11−0.58 ppb)6,15−18 may limit its applications in FO, despite the reported nontoxic effects for chronic exposure to 0.1−8 ppm TMA.6,19 The NH3−CO2−H2O system produces stable solutions that do not decompose at low to moderate temperatures easily,20,21 which may be attributed to the direct reaction between NH3 and CO2 that results in the formation of stable carbamates. © 2019 American Chemical Society

NH3 + CO2 + H 2O F NH4HCO3

(R1)

The steps involved in this process include the formation of ammonium carbamate (NH4[NH2CO2]) from NH3 and CO2 according to reaction R2:24 2NH3 + CO2 → NH 2CO2− + NH4 +

(R2)

followed by an instantaneous irreversible reaction between NH2COO− and NH4+ (reaction R3).24 NH 2CO2− + NH4 + + 2H 2O → NH4HCO3 + NH4OH (R3)

Simultaneously, reactions R4−R8 reach equilibrium quickly in the aqueous solution.24 NH3 + H 2O F NH4 + + OH−

(R4)

NH4HCO3 F NH4 + + HCO3−

(R5)

(NH4)2 CO3 F 2NH4 + + CO32 −

(R6)

Received: Revised: Accepted: Published: 12247

March 27, 2019 June 18, 2019 June 21, 2019 June 21, 2019 DOI: 10.1021/acs.iecr.9b01699 Ind. Eng. Chem. Res. 2019, 58, 12247−12252

Article

Industrial & Engineering Chemistry Research

Figure 1. Experimental methodology.

Figure 2. Removal (%) of NH3 and CO2 vs time at 0.75 bar.

OH− + HCO3− F CO32 − + H 2O

(R7)

CO32 − + H 2O + CO2 F 2HCO3−

(R8)

dihydrate (99+%, ACROS ORGANICS, CAS No. 10035-04-8, Canada). 2.2. Chemical Analysis. Total dissolved NH3 (wt %), i.e., the sum of NH3(aq) and NH4+, and total dissolved CO2 (wt %), i.e., the sum of HCO3−, CO32−, CO2(aq), and carbamate ions, were determined following the analytical procedures developed by Kolliopoulos et al.9 The analytical error was 4% for NH3 and 5% for CO2. Nuclear magnetic resonance (NMR) spectroscopy was used to obtain the CO2 speciation at 25 °C using procedures presented previously.11 2.3. Experimental Design. The draw solute initial concentration was 3.7 ± 0.2 wt % total dissolved nitrogen (NH3) and 9.4 ± 0.5 wt % total dissolved CO2 (equivalent to 2.11 ± 0.05 M NH4HCO3). We followed the experimental methodology shown in Figure 1. The experiments were conducted in an agitated (1200 rpm) 1 L glass separation reactor (Ace Glass Inc.) for t = 0, 15, 30, 45, and 60 min at each temperature−pressure combination. Temperature and pressure were monitored using a J-KEM scientific temperature controller (Model 210, VWR) and an Ashcroft digital pressure gauge (0−15 psi, Cole Palmer Canada), respectively. A Welch Dry Fast Ultra (Model 2032B-01, Fisher Scientific) vacuum pump was used to control the pressure in our setup. A known DDS mass (493.3 ± 3.6 g) was initially introduced in the separation reactor, and temperature and pressure were set. We assumed t = 0 as the time when both temperature and pressure reached the target values, which resulted in nonzero percent removal values at t = 0. All experiments were carried out sequentially and on a mass basis to account for any solution concentration changes due to water evaporation. The evolved NH3(g) and CO2(g) were collected and measured in a H2SO4 acid trap (within 3% error) and a KOH base trap

In a previous work, we studied the combined effects of temperature (30, 40, and 50 °C) and pressure (0.75, 0.55, and 0.35 bar) on the separation efficiency of the draw solutes from a TMA−CO2−H2O diluted draw solution (DDS).11 In this work, we investigate the combined effects of temperature and pressure on the draw solute separation efficiency in an NH3− CO2−H2O DDS. The aim is to study the NH3 and CO2 desorption kinetics. The above is necessary in the design of the draw solute separation process that follows the osmotic recovery of water in an FO process operating with ammonium bicarbonate as the draw solute. Finally, we compare the aqueous carbonated NH3 and TMA systems in terms of their draw solute separation efficiencies from water as well as their respective draw solution decomposition mechanisms.

2. MATERIALS AND METHODS 2.1. Materials. The NH3−CO2−H2O DDS was prepared by dissolving NH4HCO3 certified powder (Fisher Scientific, CAS No. 1066-33-7, Canada) in deionized (DI) water. A 2 M HCl standard solution (Fisher Scientific, CAS No. 7647-01-0, Canada), a 2 M NaOH standard solution (Fisher Scientific, CAS No. 1310-73-2, Canada), a 2 M CH3COONa solution, and a 4 M CaCl2 solution were used to determine total dissolved NH3 and total dissolved CO2. The solutions of CH3COONa and the CaCl2 were prepared using sodium acetate trihydrate certified ACS crystalline powder (Fisher Scientific, CAS No. 6131-90-4, Canada) and calcium chloride 12248

DOI: 10.1021/acs.iecr.9b01699 Ind. Eng. Chem. Res. 2019, 58, 12247−12252

Article

Industrial & Engineering Chemistry Research



Figure 5. Molar ratio of NH3:CO2 in solution during separation.

3. RESULTS AND DISCUSSION

(within 6% error), following the methodology described by Kolliopoulos et al.11 Samples were collected and stored in 50 mL Falcon tubes (Fisher Scientific, Canada) with minimum free headspace at 3.5 °C. They were analyzed after 24 h for total dissolved NH3, total dissolved CO2, speciation, and pH at 25 °C (1.01 bar). The reproducibility of the experimental results was evaluated in selected triplicate experiments; similar errors were assumed in all other experimental runs thereafter.

3.1. Draw Solute Separation Efficiency. The separation efficiency of NH3 and CO2 is presented in Figures 2−4; a second-degree polynomial function was used to fit the experimental data points (average R2 > 0.95 ± 0.04). Under all conditions, the separation efficiency increases with increasing time, temperature, and vacuum. A vacuum of 0.35 bar is needed to ensure almost 90% separation at 50 °C and 1200 rpm within 90 min. It is interesting that, under the same 12249


Article

Industrial & Engineering Chemistry Research Table 1. Reactions at Equilibrium Occurring during Separation of NH3 and CO2 from NH3−CO2−H2O DDS Keq

reaction 30 °C

no. R9 R10 R11 R12 R13 R14 R15 R16 R17

−

NH4 + OH ⇌ NH3(aq) + H2O NH4OH(aq) ⇌ NH4+ + OH− NH3(aq) ⇌ NH3(g) CO32− + H3O+ ⇌ HCO3− + H2O HCO3− + H3O+ ⇌ CO2(aq) + 2H2O CO2(aq) ⇌ CO2(g) NH3(aq) + HCO3− ⇌ NH2CO2− + H2O 2H2O ⇌ H3O+ + OH− H2O ⇌ H2O(g) +

5.54 × 5.28 × 0.02 1.94 × 2.09 × 33.59 1.95 1.48 × 0.04

40 °C 5.40 × 5.57 × 0.03 1.66 × 1.92 × 42.37 1.54 2.89 × 0.07

4

10 10−5 1010 106

10−14

50 °C 5.40 × 5.77 × 0.05 1.47 × 1.85 × 51.57 1.24 5.37 × 0.12

4

10 10−5 1010 106

10−14

104 10−5 1010 106

10−14

kinetics, using the results collected from t = 15 min onward. The latter was due to the initial spike in the pH value (t = 0 to t = 15 min), which is attributed to the nonstoichiometric NH3−CO2 separation. The rate equation for the desorption of a gas from a liquid is given by eq 1:

experimental conditions, only 60 min was required to remove almost 100% of the draw solutes from the TMA−CO2−H2O DDS.11 The draw solute separation process is faster and more efficient in the TMA−CO2−H2O DDS. It is also evident that the separation process does not proceed stoichiometrically in the case of the NH3−CO2−H2O DDS. As shown in Figure 5, the NH3:CO2 average molar ratio measured in solution was found to be 1.71, ranging from 1.07 at the beginning of the experiments to 2.64. In addition, the solution pH increased from 7.5−8 (t = 0) to around 9−9.5 (t = 15 min); similar behavior was observed in the concentration levels of NH2CO2− and CO32− during the first 15 min (Figure S1 in the Supporting Information). Thereafter, NH2CO2− and CO32− concentrations and pH remain approximately constant, while the concentration of HCO3− (dominant species) keeps decreasing with time. The above indicate that CO2 escapes faster than NH3 from the DDS, which contradicts the observations for the TMA−CO2−H2O system, where the separation proved stoichiometric with a TMA:CO2 average molar ratio in solution of 1.14 ± 0.04.11 Therefore, it appears that the reverse process, i.e., the regeneration of the CDS in an absorption column, would be more efficient in the TMA− CO2−H2O system, especially considering that the alkalinity resulting from the presence of NH3 or TMA in solution is necessary to capture CO2 during the CDS regeneration. 3.2. Draw Solute Desorption Kinetics. Based on all measurements, it appears that total dissolved NH3 and total dissolved CO2 are temperature independent, since there are no subsequent losses of NH3 or CO2 to the vapor phase from the samples collected once cooled to 3.5 °C. However, speciation and pH are temperature dependent and affected by the various equilibria in solution. The chemical reactions that occur during draw solute separation and their equilibrium constant values (Keq’s) at each temperature are obtained from the OLI-MSE model and are shown in Table 1. To obtain the draw solute desorption kinetics, speciation data at high temperature are required. Despite the excellent agreement between the experimental data and the predicted by OLI-MSE model pH values (error < 5.59%), the procedure we developed to reconstruct solution speciation at high temperatures11 could not be performed for the NH3−CO2−H2O system. Our methodology assumes quasi-equilibrium systems, which include reactions involving single proton exchanges.11 The carbamate ions form by the direct reaction of NH3 with CO2, which cannot be assumed at equilibrium at any given time, because based on reaction R15 (Table 1) it does not involve a single proton exchange. Therefore, the overall total dissolved NH3 and CO2 rates of disappearance in the DDS were estimated, assuming they follow pseudo-first-order

r=−

d(C − C*) = kLa(C − C*) dt

(1) −1

−1

where r is the desorption rate (mol min L ), C is the concentration of the gas in the liquid phase (mol L−1), C* is the concentration at equilibrium (mol L−1), kL is the liquid mass transfer coefficient (m min−1), and a is the volumespecific gas−liquid interfacial area (m−1). We assumed C* to be zero in our experiments since the vapor phase was continuously removed from the overhead. Further, we confirmed with OLI-MSE that the equilibrium concentration of CO2(aq) in solution is negligible compared to the total dissolved CO2. Therefore, the integrated form of eq 1 was used to determine the rate constants for the desorption of NH3 and CO2 (eq 2). We performed linear regression on the solute concentration (C) vs t using the experimentally measured total dissolved NH3 and CO2 concentration values. ln C = −kt + ln C0

(2)

−1

where k (min ), in this work, is the product of the liquid mass transfer coefficient kL (m min−1) and the volume-specific gas− liquid interfacial area a (m−1), t is the time (min), C0 is the initial solute concentration (mol L−1), and C is the concentration of the gas in the liquid phase (mol L−1). The regressed desorption rate constants, k, are presented in Table S1 in the Supporting Information. Given the R2 values, the hypothesis for first-order kinetics with respect to the total dissolved NH3 or CO2 (reactions R18 and R19, Supporting Information) was proven true. The data in Table S1 were assessed by analysis of variance (ANOVA), using STATISTICA 7 (Statsoft Inc., Tulsa, OK, USA), while significant differences of mean values were estimated at the probability level P < 0.05. Duncan’s multiple range test was used to separate means of data when significant differences were observed. ANOVA results for overall NH3 desorption showed that k values are significantly influenced by temperature. According to Duncan’s test k values are significantly higher at 50 °C than the respective values at 30 and 40 °C, which do not differ significantly. The findings for the overall desorption of CO2 were very similar. 3.3. Decomposition Mechanism. Based on the above findings, a possible draw solution decomposition mechanism was proposed by means of reaction R20: 12250


Article

Industrial & Engineering Chemistry Research Notes

2NH4 + + 2HCO3− F NH4[NH 2CO2 ] + CO2 ↑ F 2NH3 ↑ +CO2 ↑

The authors declare no competing financial interest.

■

(R20)

ACKNOWLEDGMENTS The authors would like to acknowledge the Natural Sciences and Engineering Research Council of Canada (NSERC) (NSERC Fund No. RPGIN 105655-2013) for financial support of this project.

According to reaction R20, CO2 desorbs from the initial NH4HCO3 DDS and NH3(aq) is generated. The latter reacts with HCO3− (dominant species in solution) to form carbamate NH4[NH2CO2]. Thus, only a small fraction of ammonia desorbs, at a significantly lower rate than CO2. To achieve higher NH3 desorption rates, NH4[NH2CO2] must decompose subsequently. To elucidate further the decomposition mechanism, a direct or indirect online carbamate species measurement would be required, which was not available in this work.

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(1) Cath, T. Y.; Childress, A. E.; Elimelech, M. Forward osmosis: Principles, applications, and recent developments. J. Membr. Sci. 2006, 281 (1−2), 70−87. (2) McCutcheon, J. R.; McGinnis, R. L.; Elimelech, M. A novel ammoniacarbon dioxide forward (direct) osmosis desalination process. Desalination 2005, 174 (1), 1−11. (3) McCutcheon, J. R.; McGinnis, R. L.; Elimelech, M. Desalination by ammonia−carbon dioxide forward osmosis: Influence of draw and feed solution concentrations on process performance. J. Membr. Sci. 2006, 278 (1−2), 114−123. (4) Akther, N.; et al. Recent advancements in forward osmosis desalination: A review. Chem. Eng. J. 2015, 281, 502−522. (5) Ge, Q.; Ling, M.; Chung, T.-S. Draw solutions for forward osmosis processes: Developments, challenges, and prospects for the future. J. Membr. Sci. 2013, 442 (0), 225−237. (6) Boo, C.; Khalil, Y. F.; Elimelech, M. Performance evaluation of trimethylamine−carbon dioxide thermolytic draw solution for engineered osmosis. J. Membr. Sci. 2015, 473 (0), 302−309. (7) McGinnis, R. L.; Elimelech, M. Energy requirements of ammonia−carbon dioxide forward osmosis desalination. Desalination 2007, 207 (1−3), 370−382. (8) Elimelech, M. Yale constructs forward osmosis desalination pilot plant. Membr. Technol. 2007, 2007 (1), 7−8. (9) Kolliopoulos, G.; Carlos, M.; Clark, T. J.; Holland, A. M.; Peng, D.-Y.; Papangelakis, V. G. Chemical Modeling of the TMA−CO2− H2O System: A Draw Solution in Forward Osmosis for Process Water Recovery. J. Chem. Eng. Data 2017, 62 (4), 1214−1222. (10) Kolliopoulos, G.; Holland, A. M.; Papangelakis, V. G. Modeling of density and electrical conductivity of aqueous carbonated trimethylamine (TMA−CO2−H2O) solutions at 20 °C. Monatsh. Chem. 2018, 149 (2), 453−460. (11) Kolliopoulos, G.; Papangelakis, V. G. Temperature and Pressure Effects on the Separation Efficiency and Desorption Kinetics in the TMA-CO2-H2O System. Ind. Eng. Chem. Res. 2018, 57 (43), 14767−14773. (12) Kolliopoulos, G.; Shum, E.; Papangelakis, V. G. Forward Osmosis and Freeze Crystallization as Low Energy Water Recovery Processes for a Water-Sustainable Industry. Environ. Processes 2018, 5 (S1), 59−75. (13) Kolliopoulos, G.; Papangelakis, V. G., Extraction of Water from Contaminated Effluents by Forward Osmosis. In Extraction 2018; Davis, B., et al., Eds.; The Minerals, Metals & Materials Series. Springer, Cham, Switzerland, 2018; pp 1893−1902.. (14) Kolliopoulos, G.; Martin, J. T.; Papangelakis, V. G. Energy requirements in the separation-regeneration step in forward osmosis using TMA−CO2−H2O as the draw solution. Chem. Eng. Res. Des. 2018, 140, 166−174. (15) Amoore, J. E.; Hautala, E. Odor as an aid to chemical safety: odor thresholds compared with threshold limit values and volatilities for 214 industrial chemicals in air and water dilution. J. Appl. Toxicol. 1983, 3 (6), 272−90. (16) Leonardos, G.; Kendall, D.; Barnard, N. Odor Threshold Determinations of 53 Odorant Chemicals. J. Air Pollut. Control Assoc. 1969, 19 (2), 91−95. (17) Ruth, J. H. Odor Thresholds and Irritation Levels of Several Chemical Substances: A Review. Am. Ind. Hyg. Assoc. J. 1986, 47 (3), A-142−A-151.

4. CONCLUSIONS The current study investigated the combined effects of temperature (30, 40, and 50 °C) and pressure (0.75, 0.55, and 0.35 bar) on the separation efficiency of the draw solutes and the decomposition kinetics of the NH3−CO2−H2O DDS. Experimental data for total dissolved NH3, total dissolved CO2, pH, and speciation were obtained at t = 0, 15, 30, 45, and 60 min. Draw solute separation was found to be slower and less efficient compared to TMA−CO2−H2O where a similar degree of separation (>90%) requires 30 min less retention time at the same experimental conditions (50 °C, 0.35 bar, and 1200 rpm). Also, the overall separation process is nonstoichiometric; the molar ratio of NH3:CO2 was found to range from 1.07 to 2.64, with an average value of 1.71. The latter is a clear indication that CO2 escapes faster than NH3 from solution. In order to increase the separation efficiency of the draw solutes from the NH3−CO2−H2O DDS and obtain similar separation results with the TMA−CO2−H2O system, higher operating temperatures or lower operating pressures are required. However, that would result in an increase in the operating cost of the FO process. Finally, a detailed desorption kinetics study similar to that for the TMA−CO2−H2O system could not be carried out for the NH3−CO2−H2O system, because of carbamate ion formation in solution. This ion forms by direct reaction of NH3 with CO2 and cannot be assumed to be at equilibrium at all times. A possible reaction mechanism was proposed, but to elucidate further, an online carbamate species measurement is required.

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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.iecr.9b01699. Comparison between experimental CO2 speciation and pH measurements (at 25 °C, 1.01 bar) with the OLIMSE predicted values (25 °C, 1.01 bar) in NH3−CO2− H2O system for all experimental runs (temperature, pressure) and regressed overall desorption rate constants k (min−1) for NH3 and CO2 from NH3−CO2−H2O DDS (PDF)

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REFERENCES

AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. ORCID

Vladimiros G. Papangelakis: 0000-0003-2551-1336 12251


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Temperature and Pressure Effects on the Separation Efficiency and

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