Solubility of Fumaric Acid and Its Monosodium Salt - Industrial

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Solubility of Fumaric Acid and Its Monosodium Salt Carol A. Roa Engel,†,§ Joop H. ter Horst,‡ Mervin Pieterse,† Luuk A. M. van der Wielen,† and Adrie J. J. Straathof*,† †

Department of Biotechnology, Delft University of Technology, Julianalaan 67, 2628 BC Delft, The Netherlands Intensified Reaction and Separation Systems, Process & Energy Department, Delft University of Technology, Leeghwaterstraat 44, 2628 CA Delft, The Netherlands

‡

ABSTRACT: Fumaric acid is a dicarboxylic acid applied in food industry and in some polymers. Currently, its fermentative production from renewable resources is receiving much attention, and crystallization is used to recover it. To determine the window of operation for crystallization from multicomponent fermentation mixtures, the aqueous solubilities of fumaric acid and its sodium salts were investigated. For fumaric acid, sodium hydrogen fumarate, and sodium fumarate, solubilities and pH increased in this order because of increasing polarity and dissociation. A mathematical model was developed to predict crystal type and amount as function of temperature and pH. The effect of glucose (up to 3.0 mmol·mol−1) on the solubility can be neglected, but ethanol (1.0 mmol·mol−1) slightly increased the solubility of fumaric acid and significantly decreased the solubility of the sodium salts, because the aqueous solution becomes less polar upon ethanol addition but not upon glucose addition.

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study,8 we use saturation temperature measurements of mixed composition to obtain the necessary thermodynamic data.

INTRODUCTION Fumaric acid (trans-1,2-ethenedicarboxylic acid) is an important starting material for the production of food additives and has great potential in the production of some polymers.1 The production of fumaric acid (H2FA) by fermentation of carbohydrates is an attractive alternative to the current process, because it avoids the use of petroleum-derived maleic anhydride. In addition, the fermentation process involves carbon dioxide fixation,2 increasing the sustainability of the fumaric acid production. As with other carboxylic acids produced by fermentation (e.g., lactic acid, citric acid), the highest cell-specific production rates are found around neutral pH values.3 In the case of fumaric acid, to achieve the optimal pH (∼5), sodium hydroxide (NaOH) has been identified as a very effective neutralizing agent.4 pH shift crystallization, upon acidification with sulfuric acid, is the most common method used to recover carboxylic acid fermentation products from fermentation broths.5,6 However, this acidification leads to production of waste salts. In the case of fumaric acid, the acidification step leads to the production of sodium sulfate (Na2SO4) if NaOH is used as neutralizing agent. Recovery from the fermentation broth by cooling crystallization is a very interesting option if the temperature dependence of the solubility is sufficiently high and if the aqueous solubility at low temperature is sufficiently low.5 However, the speciation of fumaric acid might be affected by the multicomponent fermentation broth composition and pH.7 This can result in the formation of different fumarate salts upon crystallization. We aim to control the fumaric acid cooling crystallization in the pH range from 3.0 to 5.0. Therefore, the solubilities of fumaric acid and its salts are determined as a function of temperature, pH and concentration of fermentation cosolutes, focusing on glucose, which is the fermentation feed material, and ethanol, which is the main byproduct. In addition, speciation of fumarate salts is addressed, to predict which solid state will be obtained from the fermentation broth under different pH and temperature conditions. Similar to a previous © XXXX American Chemical Society

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MATERIALS AND METHODS

Chemicals. Fumaric acid (99%), sodium fumarate (98%), glucose (99%), and ethanol (98%) (Sigma−Aldrich products) were used without further purification. Distilled water was used in all preparations. Saturation Temperature Determination. Saturation temperatures were determined by use of Crystal16 equipment of Avantium Technologies (Amsterdam, The Netherlands). In the Crystal16, cloud points and clear points of 16 1-mL solution aliquots were measured in parallel by automatic turbidity detection. The temperature at which the suspension became a clear solution upon heating at 0.5 °C·min−1 was taken as the saturation temperature of the measured sample. After recrystallization at 5 °C, the saturation temperature was measured again. By use of preliminary solubility values, the van ’t Hoff equation was used to calculate the fumaric acid mole fractions at saturation, xH2FA*, of 2.33, 5.35, and 7.55 mmol·(mol of solvent)−1 at T = 48, 70, and 80 °C, respectively. A saturated sodium fumarate mole fraction (xNa2FA*) of 24.15 mmol·(mol of solvent)−1 was assumed at all temperatures. Saturation temperatures of mixtures were determined in this way for a number of preweighed samples containing fumaric acid, sodium fumarate, or a mixture of both in water, water/ ethanol mixtures, or water/ethanol/glucose mixtures. The mole fractions x used for samples containing fumaric acid, sodium fumarate, and water were determined according to eq 1, and eq 2 was used to express solvent-free compositions.8 Figure 1 Received: March 12, 2013 Revised: June 12, 2013 Accepted: June 12, 2013

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dx.doi.org/10.1021/ie400794r | Ind. Eng. Chem. Res. XXXX, XXX, XXX−XXX

Industrial & Engineering Chemistry Research

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mmol·L−1 in Milli-Q water) at a column temperature of 59 °C and a flow rate of 0.6 mL·min−1.

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RESULTS AND DISCUSSION Three solid states were expected to occur in the system consisting of water, sodium ions, and fumaric acid: fumaric acid, disodium fumarate, and the less known sodium hydrogen fumarate.9,10 Dissociation and solubility equilibria of these species in aqueous solution are presented in Figure 2.

Figure 2. Fumarate equilibria in aqueous solution for Na+ as the counterion. Figure 1. Mole fractions of fumaric acid (xH2FA) and sodium fumarate (xNa2FA) used to generate aqueous solutions of these two compounds.

At low pH values, dissolved fumaric acid [H2FA (aq)] is the prevailing form in solution, and if the aqueous solution is saturated, fumaric acid can crystallize. On the other hand, at high pH, where fumaric acid is completely dissociated, sodium fumarate (Na2FA) is prevailing. At intermediate pH, sodium hydrogen fumarate (NaHFA) could be expected. The following sections deal with identification of the pH and temperature ranges for the formation of these three fumarate compounds.

shows the compositions of the mixtures prepared. The component pairs were dissolved at T = 80 °C before being placed in the Crystal16. x Na 2FA x H2FA =1− x Na 2FA* x H2FA* (1) yH FA = 2

x H2FA x H2FA + x Na 2FA

(2)

Solubility as a Function of pH. In a 100 mL jacketed glass reactor, a certain excess amount of dry fumaric acid crystals was added to 50 mL of a 140 g·L−1 solution of sodium fumarate in water. The resulting suspension was stirred magnetically for 24 h at constant temperature (thermostat bath, Lauda RMG). The pH was measured on a Metrohm 691-pH meter. After the crystals had settled for 1 h, solution samples were taken via pipet and immediately diluted with a known amount of water to avoid recrystallization. The total concentration of fumaric acid species in the diluted sample was determined by HPLC analysis. Analytical Methods. X-ray powder diffraction (XRPD) patterns were recorded in a Bragg−Brentano geometry on a Bruker D5005 diffractometer equipped with an incident beam monochromator and a Braun PSD detector. Data collection was carried out at room temperature with monochromatic Cu Kα1 radiation (λ = 1.540 562 Å) in the 2θ region between 5° and 90°. Diffraction patterns were compared by use of PowderCell 2.3 (Berlin, Germany). Inductively coupled plasma optical emission spectrometry (ICP-OES) was used to determine the sodium contents of crystalline material. Perkin-Elmer Optima 3000dv equipment was used at 589.592 and 588.995 nm wavelengths. Measured concentrations at the selected wavelengths were averaged. Fumaric acid, glucose, and ethanol concentrations were quantified by HPLC with a Bio-Rad Aminex HPX-87H ionexclusion column (300 × 7.8 mm) with a refractive index detector (Waters 2414) and UV detector at 210 nm (Waters 2489). The column was eluted with dilute phosphoric acid (1.5

Figure 3. Saturation temperature as a function of the solvent-excluded mole fraction yH2FA of H2FA in the presence of Na2FA. (□) xH2FA* = 2.33 mmol·mol−1; (○) xH2FA* = 5.35 mmol·mol−1; (◇) xH2FA* = 7.55 mmol·mol−1. Lines are model lines.

Phase Diagram and Fumarate Speciation. Figure 3 depicts the saturation temperature (TS) as a function of yH2FA. When yH2FA = 1, fumaric acid is the only solute and the points plotted there are the TS values of mixtures with different amounts of fumaric acid. When a higher proportion of fumaric acid is added to the solution, a higher temperature is required B



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for dissolution in the range ∼0.4 < yH2FA < 1. However, Figure 3 shows a local maximum when 0.05 < yH2FA < 0.4. Therefore, another stable salt is formed in this region. When yH2FA = 0, the solutions contain only sodium fumarate and water, but no saturation temperature was observed for the conditions studied, which was attributed to the very high solubility of sodium fumarate. The phase transition points of the two systems for xH2FA* = 7.55 and xH2FA* = 5.33 are close to each other, around yH2FA = 0.394 (29 °C) and 0.366 (23 °C), respectively. The phase transition point for the system where xH2FA* = 2.33 seems to be located at temperatures below 0 °C and it was not determined experimentally. The stable additional phase that was observed in the phase diagram system in the region of the local maximum is expected to be crystalline sodium hydrogen fumarate (NaHFA), by analogy to systems containing potassium hydrogen fumarate11 and sodium hydrogen glutamate.7 The X-ray powder diffraction pattern of crystals at yH2FA ≈ 0.2 was compared not only with the theoretical sodium hydrogen fumarate pattern but also with experimental patterns of fumaric acid and sodium fumarate (Figure 4). The XRPD

convention was adopted that activity coefficient values tend to 1.0 as the solution tends to infinite dilution. In contrast, for water and solids, if the mole fraction approaches 1.0, the activity coefficient tends to 1.0 according to the convention used.12 AESolve (www.halotec.com), a simulator for performing complex equilibrium calculations on aqueous electrolyte systems, was used to calculate the equilibrium compositions. The model simultaneously solved the equilibria shown in Figure 2 and the conservation balances for the 10 species involved. The aqueous species were H2O, H+, OH−, Na+, H2FA, HFA−, and FA2− and the solid species were H2FA, NaHFA, and Na2FA. Four independent conservation balances were used: for total sodium, total fumarate, H2O + OH−, and charge. Values of six thermodynamic equilibrium constants were required, and these are treated subsequently. The first and second dissociation constant of fumaric acid (Ka,1 and Ka,2) can be expressed as K a,1 =

aq aq −)(γ +)(m +) (γHFA−)(m HFA H H

(γH FA )(m Haq2FA )

(3)

aq aq (γFA2−)(m FA 2 −)(γ +)(m +) H H aq −) (γHFA−)(m HFA

(4)

2

K a,2 =

where m is the molality (moles per kilogram of solvent) and γ is the activity coefficient. Values of Ka,1 and Ka,2 were obtained with the correlation pKa,i = Ai/T + Bi + CiT, where A1 = −205.2277 K, A2 = −87.5800 K, B1 = 6.8883, B2 = 4.0771, C1 = −0.010 390 0 K−1, and C2 = 0.002 871 5 K−1.13 Activity coefficients for ions with charge z were calculated by using an extended Debye−Hückel expression,12 which is valid up to ionic strength I = 0.5 mol·(kg of solvent)−1: −log γi =

AEDHz 2I 0.5 1 + I 0.5

− bI

(5) −3/2

The constants are taken as AEDH = 1.8249 × 10 ρwater(εT) , b = 0.2 with I in moles per kilogram and defined by 0.5∑(zi2mi) for a solution with ions with charges zi. Temperature correlations were used for the density ρwater and the dielectric permittivity ε. For undissociated fumaric acid, a simple relationship was used to obtain the activity coefficient:12 3

log γi = 0.1I

(6)

Water dissociation into H+ and OH− is an equilibrium reaction not shown in Figure 2, but to be able to properly calculate the amounts of all species involved, a literature correlation for the water dissociation constant was used.14 The equilibrium constant for the conversion H2FA(s) ↔ H2FA(aq) (i.e., the solubility constant) equals

Figure 4. XRPD analysis of the obtained crystalline phase. Theoretical values were obtained for the reported crystal structures of H2FA, NaHFA, and Na2FA.11,19,20.

pattern of the obtained crystalline material of yH2FA = 0.2 (xH2FA* = 7.55 mmol·mol−1) indicates the presence of sodium hydrogen fumarate crystals. The molar concentration ratio between sodium and fumarate contents of these crystals was determined to be 1.09 ± 0.1. A ratio of 1 would be expected for sodium hydrogen fumarate, 0 for fumaric acid, and 2 for sodium fumarate. According to these results, the additional crystal type is indeed sodium hydrogen fumarate. Thermodynamic Model. The experimental data presented in Figure 3 were validated by use of a thermodynamic model, which is based on the equilibria presented in Figure 2. For dissolved species, molalities were used and the reference

Kd1 =

a Haq2FA a Hs 2FA

=

(γHaqFA )(m Haq2FA ) 2

1

(7)

determine mHaq2FA, 10,15

To measured solubilities of fumaric acid from and this paper were used. The fumaric acid the literature will be partly dissociated, so the solubility constant of the undissociated species was found by correcting the measured solubilities also for the fraction of undissociated species, by use of eqs 3−6 at the measured solubilities. Due to this correction, the ionic strength of the solution changes slightly and hence the calculated activity coefficient and the fraction of undissociated species, but it was checked that this effect was negligible. Figure C



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5 shows the resulting values, indicating that all data can be used for obtaining a good linear fit. A fit against 1/T was clearly worse and was not used.

Figure 6. Temperature dependence of the negative logarithm of Kd2. Symbols are values calculated from individual measurements of solubility of sodium hydrogen fumarate; the line is a fit. Figure 5. Temperature dependence of mHaq2FA in saturated aqueous solution. Symbols are solubility measurements of fumaric acid corrected for calculated fractions of undissociated fumaric acid species; the line is a fit.

pKd2, the model lines do not represent very well the experimental data in the range when 0 < yH2FA < 0.3 at xH2FA* = 2.33 (Figure 3). The model lines confirm that the three-phase transition points (local minima in the curves in Figure 3) occur at higher yH2FA and TS when xH2FA increases. This phenomenon may be explained by the fact that Ka,1 increases with temperature whereas Ka,2 decreases, which leads to a higher proportion of sodium hydrogen fumarate compared to fumaric acid in aqueous solutions with increasing temperature.16 Then, a relatively large amount of fumaric acid has to be added to reach saturation of both sodium hydrogen fumarate and fumaric acid, and this will occur at high values of yH2FA. Apparent Solubility of Total Fumarate as a Function of pH and Temperature. By use of the model, the apparent (total) solubility can be calculated for the equilibria in Figure 2 as function of temperature and pH. This solubility is the dissolved amount measured when several species of a dissociating compound coexist and is often required in practical applications. To validate the model, apparent solubility measurements were performed, and Figure 7 shows the effect of different pH and temperature values. As expected, the apparent solubility increases with pH and temperature.15 In Figure 7 the model predicts the experimental data quite well, especially at 35 and 15 °C. However, at 5 °C the model does not predict so well experimental data above pH 3.5. This could be due to inaccuracy of the pKa predictions at lower temperature. Moreover, Figure 7 shows a discontinuity in the solubility curves because not fumaric acid but sodium hydrogen fumarate is determining the apparent solubility at higher pH, for example, for 35 °C at pH > 3.65. In addition, during the performance of solubility experiments in the range pH 3.8−4.2 (depending on the temperature), the pH did not increase any further when more sodium fumarate was added to a saturated solution of fumaric acid. The model explains that the reason for this phenomenon was the phase transition to formation of sodium hydrogen fumarate crystals. At the intersection of the two parts of the curves, the two crystal types can coexist. The model can predict the pH and temperature values of

According to eq 7, the solubility correlation found needs to be corrected by the activity coefficient of fumaric acid (eq 6) to obtain the equilibrium constant for dissolution: pKd1 = −0.0168T + 6.321 − 0.1I

(8)

This equation is valid for T in kelvins and I in moles per kilogram. For the reaction NaHFA(s) ↔ Na+(aq) + HFA−(aq), the equilibrium constant is Kd2 =

aq aq (a Na +)(a HFA−) s aNaHFA

=

aq aq aq aq (γNa )(γHFA −)(m HFA−) +)(m Na +

1 (9)

The experimental data from Figure 3 were taken where the system was supposed to be saturated by sodium hydrogen fumarate, that is, the points at low fractions of fumaric acid. For these data, at various temperatures, Kd2 was calculated by use of the model equations. The results (Figure 6) show an increasing K value with increasing temperature, which is usual for dissolution. The scatter is relatively large. This might be due to the relatively high ionic strengths in the experiments, which would lead to inaccurate calculations of the activity coefficients. The model equation obtained from Figure 6 is 1.2 × 103 − 3.2 (10) T 9 −1 For sodium fumarate, a solubility of 228 g·kg at 25 °C was included in the model, but the exact value of this high solubility was not important because sodium fumarate was completely dissolved in all cases evaluated by the model. The model was used to describe the experimental data in Figure 3. It behaves quite well for the systems with xH2FA* = 7.55 and xH2FA* = 5.35. Due to the inaccuracy in the fit for pKd2 =

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Figure 7. Solubility of total fumarate as a function of pH at different temperatures: experimental points at (□) 35, (○) 15, and (◇) 5 °C. Lines are model lines at different temperatures. The dashed line is a model line at standard conditions (25 °C). For each temperature, the data point at the left-hand side indicates the pH of a saturated solution of H2FA. For the curve at 35 °C, the solid at equilibrium with the saturated solution has been indicated.

Figure 8. Calculated pH of aqueous mixtures of fumaric acid and sodium fumarate at different temperatures.

the saturation temperature of sodium hydrogen fumarate and fumaric acid solutions are shown in Figure 9. Model lines and experimental points from Figure 3 have been added to Figure 9 to facilitate the comparison. No new solid states seem to appear when glucose and ethanol are present in the fumarate solution, as confirmed by XRPD. Figure 9a illustrates the phase diagram of the mixture of sodium fumarate where xH2FA* = 7.55 mmol·mol−1 plus the changes in the phase diagram when 1.5 and 3.0 mmol·(mol of glucose)−1 and 1.0 mmol·(mol of ethanol)−1 were part of the mixture. The same systems are shown in Figure 9b,c for xH2FA* = 5.35 and 2.33 mmol·mol−1, respectively. When only glucose up to 3.0 mmol·mol−1 is added to the sodium fumarate systems, the solubility of these mixtures is hardly affected. Ethanol, however, affects the solubility of sodium hydrogen fumarate and fumaric acid in different ways. Figure 9 shows that when ethanol is present, the saturation temperature left of the minimum increases. The minimum value of yH2FA leading to fumaric acid crystals was shifted from about 0.4 to about 0.5 for the conditions employed. This means that the solubility of sodium hydrogen fumarate decreases by adding ethanol. Because ethanol is a less polar solvent than water, it is expected that the solubility of the fumarate salt decreases when ethanol is present. On the other hand, right of the minimum, the presence of ethanol slightly decreases the saturation temperature, corresponding to an increase in the solubility of fumaric acid. This is occurring because fumaric acid is a more hydrophobic compound, which will dissolve better when the solution becomes less polar upon addition of ethanol (fumaric acid solubility in ethanol18 is 58 g·L−1 in 95% ethanol at 30 °C). This effect of ethanol is stronger at lower temperature.

coexistence. It will also predict the crystal types at other conditions. For example, the literature mentions apparent solubilities of sodium hydrogen fumarate.9 According to our model at the temperatures used by the authors, the solutions would become saturated by fumaric acid crystals at 25 and 40 °C (at solubilities of 43 and 89 g·kg−1 and pH values of 3.55 and 3.48, respectively) but by sodium hydrogen fumarate crystals at 60 and 100 °C (at solubilities of 147 and 209 g·L−1 and pH values of 3.43 and 3.33, respectively). The explanation for this interesting switch is the same as the aforementioned explanation of the position of three-phase transition points in Figure 3. Fumarate fermentations using Rhizopus oryzae have been performed at 35 °C in a pH range from 3.0 to 6.5.2,17 We wondered which pH would be required at 5 °C to obtain fumaric acid crystals upon cooling. Figure 8 shows the calculated pH as a function of yH2FA for the conditions of Figure 3. Figure 8 illustrates, together with Figure 3, that to obtain fumaric acid crystals, the pH of fumaric acid solutions should be below 3.7 during fermentation. The lines of the three mixtures are crossing at yH2FA = 0.35. This could be occurring because at higher xH2FA* values more extreme pH values can be achieved because there is more acid or base added to the solutions. Above pH 3.7, sodium hydrogen fumarate crystallizes. However, cooling from 35 to 5 °C moves the transition point from fumaric acid to sodium hydrogen fumarate around pH 4, giving a larger pH range for fumaric acid production. Influence of Cosolutes on the Sodium Fumarate System. We further investigated the effect of the cosolutes glucose and ethanol on the solubility of sodium hydrogen fumarate and fumaric acid, since these are common compounds present in the production of fumaric acid by fermentation (glucose as a feed compound and ethanol as a byproduct). The mole fractions used were xGlu = 1.5 and 3 mmol·mol−1 and xEtOH = 1 mmol·mol−1. The effects of glucose and ethanol on

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CONCLUSIONS The aqueous solubilities of H2FA and NaHFA both increase with temperature and with pH. This solubility behavior was described by a mathematical model based on dissociation and solubility equilibria. The solubility is determined by H2FA in the region pH < ∼3.7 and by NaHFA in the region pH > ∼3.7. E



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Figure 9. continued fumarate solution + xGlu = 3.0 mmol·mol−1; (●) sodium fumarate solution + xGlu = 1.5 mmol·mol−1 + xEtOH = 1.0 mmol·mol−1; (◆) sodium fumarate solution + xGlu = 3.0 mmol·mol−1 + xEtOH = 1.0 mmol·mol−1. Lines are model lines from Figure 3.

As the three-phase equilibrium point shifts with temperature, solutions of NaHFA (pH ≈ 4) may, upon cooling, yield H2FA crystals. Glucose up to 3.0 mmol·mol−1 did not significantly influence the solubility of NaHFA and H2FA, but 1.0 mmol·(mol of ethanol)−1 decreased the solubility of NaHFA while it increased the solubility of H2FA. Moreover, the presence of ethanol increases the stability region for NaHFA crystals, and the minimum value of yH2FA leading to fumaric acid crystals was shifted from about 0.4 to about 0.5 for the conditions employed.

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AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. Phone: +31-15-278 2330. Fax: +31-15-278 2355. Present Address §

TNO, Leeghwaterstraat 46, NL-2628 CA Delft, The Netherlands. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS This work has been financially supported by The Netherlands Ministry of Economic Affairs and the B-Basic partner organizations (www.b-basic.nl) through B-Basic, a public private NWO-ACTS programme (ACTS, Advanced Chemical Technologies for Sustainability). Ruud Hendrikx and Joop Padmos are acknowledged for the X-ray diffraction and sodium analysis, respectively. Guido Breuer is acknowledged for the experimental study of fumarate solubility.

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NOMENCLATURE a = activity AEDH = Debye−Hückel parameter b = Debye−Hückel parameter I = ionic strength i = initial value Ka = dissociation equilibrium constant Kd = dissolution equilibrium constant m = molality T = absolute temperature x = mole fraction x* = mole fraction at saturation y = solvent-free mole fraction z = ion charge

Greek

γ = activity coefficient ε = dielectric permittivity ρ = density

Subscripts

Figure 9. Influence of fermentation cosolutes on saturation temperatures of mixtures of H2FA and Na2FA. (a) xH2FA* = 7.55 mmol·mol−1; (b) xH2FA* = 5.35 mmol·mol−1; (c) xH2FA* = 2.33 mmol·mol−1. (○) Sodium fumarate solution + xGlu = 1.5 mmol·mol−1; (◇) sodium

EtOH = ethanol Glu = glucose H2FA = fumaric acid Na2FA = sodium fumarate F



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NaHFA = sodium hydrogen fumarate Superscripts

aq = aqueous phase s = solid phase

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REFERENCES

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Solubility of Fumaric Acid and Its Monosodium Salt - Industrial

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