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Thermophysical Properties of Carboxylic and Amino Acid Buffers at Subzero Temperatures: Relevance to Frozen State Stabilization Prakash Sundaramurthi†,‡ and Raj Suryanarayanan*,† † ‡
Department of Pharmaceutics, University of Minnesota, Minneapolis, Minnesota 55455, United States Scientific Affairs, Teva Parenteral Medicines Inc., 11 Hughes, Irvine, California 92618, United States ABSTRACT: Macromolecules and other thermolabile biologicals are often buffered and stored in frozen or dried (freeze-dried) state. Crystallization of buffer components in frozen aqueous solutions and the consequent pH shifts were studied in carboxylic (succinic, malic, citric, tartaric acid) and amino acid (glycine, histidine) buffers. Aqueous buffer solutions were cooled from room temperature (RT) to 25 °C and the pH of the solution was measured as a function of temperature. The thermal behavior of frozen solutions was investigated by differential scanning calorimetry (DSC), and the crystallized phases were identified by X-ray diffractometry (XRD). Based on the solubility of the neutral species of each buffer system over a range of temperatures, it was possible to estimate its degree of supersaturation at the subambient temperature of interest. This enabled us to predict its crystallization propensity in frozen systems. The experimental and the predicted rank orderings were in excellent agreement. The malate buffer system was robust with no evidence of buffer component crystallization and hence negligible pH shift. In the citrate and tartrate systems, at initial pH < pKa2, only the most acidic buffer component (neutral form) crystallized on cooling, causing an increase in the freeze-concentrate pH. In glycine buffer solutions, when the initial pH was ∼3 units < isoelectric pH (pI = 5.9), β-glycine crystallization caused a small decrease in pH, while a similar effect but in the opposite direction was observed when the initial pH was ∼3 units > pI. In the histidine buffer system, depending on the initial pH, either histidine or histidine HCl crystallized.
’ INTRODUCTION The importance of the concept of pH is now widely recognized, since its introduction in 1909 by Sorensen.1,2 The reaction rate in numerous biological and chemical systems is dependent on solution pH. Many compounds of biological and pharmaceutical interest are stable only over a narrow pH range.35 For example, penicillin G has optimal stability in the pH range of 5.57.5.6,7 A small shift outside this limit significantly reduces its activity thereby necessitating the use of a buffer system. Most weak acids and bases are characterized by pKa values in the range of 410. By partially neutralizing the ionizable group in aqueous solution, they could be used as buffers. Such solutions, containing a weak acid/base and its conjugated base/acid, in concentration ratios of 1:10 to 10:1 could be used as buffers covering the pH range of pKa (1.8 The selection of a buffer system is based on its pKa, buffer capacity, and possibility of buffer specific catalysis.9,10 For compounds which degrade readily in solutions, freezedrying is a common method of stabilization.11 This method finds wide application in both food and pharmaceutical industries. The process consists of freezing, primary drying (ice sublimation), and secondary drying (removal of unfrozen water). In the selection of buffer system for freeze-dried formulations, propensity for crystallization and consequent pH shift in frozen solutions should be considered.12 In frozen phosphate buffer solutions, crystallization of disodium hydrogen phosphate dodecahydrate (Na2HPO4 3 12H2O) caused a pronounced pH shift in the freezeconcentrate.1315 Similarly, in frozen succinate buffer solutions, the “pH swing” was attributed to sequential crystallization of buffer components.16 r 2011 American Chemical Society
Several experimental methods have been used to study crystallization behavior and consequent pH changes in frozen buffer solutions. In earlier investigations, the buffer solutions were cooled to a desired temperature and the frozen and unfrozen fractions were physically separated.13,17,18 The pH and composition of the unfrozen fraction was determined at room temperature (RT). Thermal analysis was used to monitor the crystallization or eutectic melting of the buffer components.19,20 Frozen solution pH measurements were used to directly monitor the pH change during freezing.2123 Recently, low-temperature X-ray diffractometry was used to characterize crystallizing phases in frozen systems.2426 In our preliminary investigation, based on the reported solubility values at and above RT, we had predicted the crystallization propensity of a series of carboxylic acid buffer components in frozen systems.27 The experimental and the predicted rank orderings were in excellent agreement. In the current investigation, we have provided comprehensive experimental evidence, both calorimetric and diffractometric, for the buffer crystallization in frozen systems. We have also expanded the investigation to amino acid buffer systems. Finally, based on crystallization propensity, we have developed a rank ordering system encompassing both the carboxylic and amino acid systems to aid in buffer selection for frozen and freezedried systems. We hypothesize that the temperature-dependent solubility coupled with the predicted degree of supersaturation in frozen solution would aid in rank ordering the buffer systems. Received: March 7, 2011 Revised: April 21, 2011 Published: May 11, 2011 7154
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Table 1. Thermophysical Characterization of Frozen Carboxylica and Amino Acid Buffer Systems buffer succinic acid/Na succinate
malic acid/Na malate
tartaric acid/Na tartrate
citric acid/Na citrate
glycine/glycine HCl or Na
histidine/histidine HCl
initial pH (at RT) ΔpHUF (pH0 °C pH25 °C) ΔpHF (pH25 °C pHint)
XRD (frozen solution)
DSC (heating curve)
4.0
0.0
4.0b
SA, NaHSA
5.0
0.0
2.1b
multiple Te
6.0
0.1
NaHSA, Na2SA 3 6H2O SA, NaHSA, Na2SA 3 6H2O
multiple Te
1.7
multiple Te
4.0
0.0
0.1
amorphous
no transition
5.0
0.0
0.3
amorphous
no transition
6.0
0.1
0.4
amorphous
no transition
3.0
0.0
1.8
DL-
4.0 5.0
0.0 0.0
0.8 0.30
D-tartaric
acid amorphous
4.0
0.1
1.7
citric acid monohydratec
no transition
5.0
0.1
0.1
amorphous
no transition
6.0
0.1
0.1
amorphous
no transition
3.0
0.0
0.5
β-glycine, diglycine HCl
Tg, Te
9.0
0.3
2.5
β-glycine
Tg, Te
9.5
0.4
1.5
β-glycine
Tg, Te
5.5 6.0
0.3 0.3
1.9 0.9
L-histidine
7.0
0.3
0.7
L-histidine
and D-tartaric acid
Tg, Tc, Te1, and Te2 Te no transition
HCl no transition L-histidine HCl and L-histidine no transition Tc
Selected data reproduced from ref 27 with permission of the copyright owner. ΔpHUF: pH shift in unfrozen solution. ΔpHF: pH of frozen solution. SA: β-succinic acid. NaHSA: monosodium succinate. Na2SA 3 6H2O: disodium succinate hexahydrate. Te, eutectic melting; Tg, glass transition; and Tc, crystallization. b “pH swing” was observed (an increase in pH followed by a decrease; only the initial increase was considered). c Low-intensity XRD peaks. a
’ MATERIALS AND METHODS Materials. Citric acid, DL-malic acid, succinic acid, DL-tartaric acid, L-histidine, and glycine with purity >98% were purchased from Sigma. Aqueous sodium hydroxide (Mallinckrodt) solution and hydrochloric acid (Sigma) were used to partially neutralize the acid and base respectively to prepare buffer systems. All these chemicals were used as received without further purification. A pH meter (Oakton), calibrated with standard buffer solutions (Oakton standard buffers; pH 2.00, 4.01, 7.00, and 10.00; certified by NIST) was used. The systems investigated are listed in Table 1. Preparation of Buffer Solutions. Buffer solutions (200 mM) were prepared by dissolving the appropriate amount of acid or base in degassed water and adjusting to desired pH with either 2 M NaOH or 1 M HCl at RT. The solutions were membrane filtered (0.45 μm PTFE; Fisher, USA) and stored in tightly closed glass vials at RT. While HCl was used to adjust the pH of L-histidine, the pH of glycine solution was adjusted either with HCl or with NaOH. Temperature and pH Measurements during Freezing. The buffer solution (25 mL) was placed in a jacketed beaker (100 mL) connected to a water bath with an external temperature controller unit (Neslab RTE 740, Thermo Electron, NH). The bath fluid (Dynalene HC-50, Dynalene Heat Transfer Fluids, PA), with a working temperature range of 80 to 40 °C, was used. A lowtemperature pH electrode (Inlabcool, Mettler Toledo, Switzerland) was placed in the center of the sample and connected to a pH meter (pH 500 series, Oakton, Singapore) to monitor the electromotive force (EMF). The measured EMF was then used to calculate the solution pH. The reference electrolyte containing glycerol and formaldehyde (Friscolyte-B, Mettler Toledo, Switzerland) allowed pH measurements down to 30 °C. A copper-constantan thermocouple (0.05 in. diameter, Omega, Stanford, CT) with Teflon
insulation connected to a digital benchtop read-out device ((0.2 °C; Omega MDSi8 Series, Stanford, CT) was used to monitor the temperature changes in the sample upon cooling. In all the experiments, the thermocouple was placed in the middle of the sample close to the electrode bulb. More details of the experimental setup and calibration procedure are provided in our earlier publications.16,28 X-ray Diffractometry (XRD). A powder X-ray diffractometer (Model XDS 2000, Scintag; BraggBrentano focusing geometry) with a variable-temperature stage (High-Tran Cooling System, Micristar, Model 828D, R.G. Hansen & Associates; working temperature range 190 to 300 °C) and a solid-state detector was utilized for low-temperature XRD studies. The buffer solution (1 mL) was placed in an aluminum sample holder and covered with a stainless steel dome with a beryllium window. Periodically, the system was exposed to Cu KR radiation (45 kV 40 mA), and the XRD patterns were obtained by scanning over an angular range of 535° 2θ with a step size of 0.05° and a dwell time of 1 s. The samples were cooled from RT to 10 at 1 °C/min and the first XRD pattern was recorded. The chamber temperature was then decreased to 24 °C at a step size of 2 °C and at each step an XRD pattern was collected. During each scan (about 10 min) the sample was held under isothermal conditions. The underlying cooling rate, from 10 to 24 °C, was 50% of the tartaric acid (H2A) was nonionized, and in light of its low aqueous solubility, a fraction crystallized. XRD provided evidence of the crystallization of both the D and DL isomers of tartaric acid in the annealed frozen solutions (Figure 4, bottom pattern). While the buffer solution was prepared using the DL-isomer, crystallization of the D-isomer was observed in the frozen system. According to Wallach’s rule the racemic crystals tend to be denser, (i.e., more stable) than their chiral counterparts. Thus, the formation of the D-isomer, the “metastable form”, in the frozen solution is not surprising.35 The DSC heating curve of tartrate buffer (pH 3.0) showed a glass transition (35.4 °C), recrystallization (23.7 °C), and three overlapping endothermic events at 2.7, 0.9 °C, and þ1.7 °C (Figure 2, panel b; see also inset). While these events appear as a shoulder, they were discernible in a derivative plot (not shown). We hypothesize that, upon cooling the buffer solution (initial pH 3.0), a fraction of tartaric acid crystallized, first as DL-isomer (least soluble form). The remaining solute in the amorphous freeze-concentrate exhibited a glass transition followed by recrystallization as D-isomer during warming in the DSC (Figure 2b; inset). The DSC heating curve of tartrate buffer (pH 4.0) showed overlapping eutectic (2.9 °C) and ice melting (Figure 2, panel b). XRD patterns indicated the crystallization of DL-tartaric acid in frozen solution (Figure 4). Both DSC and XRD suggested that only ice had crystallized in the frozen solution when it was buffered to pH 5.0. 7157
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Figure 3. XRD pattern of frozen malate buffer solutions. The solutions were initially cooled from RT to 10 °C, held for 10 min, and then cooled to 25 °C while collecting XRD patterns at 2 °C intervals. The frozen solution was annealed at 25 °C for 60 min. The cooling and heating rate was 1 °C/min. Only the XRD patterns of the annealed samples are shown.
Figure 4. XRD pattern of frozen tartrate buffer solutions. The solutions were initially cooled from RT to 10 °C, held for 10 min, and then cooled to 25 °C while collecting XRD patterns at 2 °C intervals. The frozen solution was annealed at 25 °C for 60 min. The cooling and heating rate was 1 °C/min. Only the XRD patterns of the annealed samples are shown. The characteristic peaks of DL- and D-tartaric acid are pointed out.
Citrate Buffer. The pH change was negligible upon freezing and annealing the citrate buffer solutions with initial pH values of 5.0 or 6.0. However, when the initial pH was 4.0, there was a pronounced increase in the freeze-concentrate pH, upon cooling and annealing (∼1.7 pH units). This is attributed to the crystallization of citric acid. While the DSC heating curve only indicated ice melting, XRD provided evidence for the crystallization of citric acid monohydrate (Figure 2, panel c). A characteristic peak, attributable to both citric acid and citric acid monohydrate (2θ 23.2°), was barely discernible (not shown). Interestingly, after the low-temperature pH measurements, when the solution was thawed, a white precipitate was observed. These particles
dissolved as the sample temperature increased to RT. We hypothesize that a citric acid phase precipitated in the frozen solution and since the monohydrate will be the least soluble species at this temperature, its appearance is expected. There have been reports suggesting that citric acid monohydrate crystallizes poorly if no seed crystals are available.3638 Ueda et al., using scanning electron microscopy, probed the crystallization mechanism of citric acid monohydrate from aqueous solutions of varying degrees of supersaturation. They reported the formation of aggregates (∼16 nm) of citric acid which phase separate as “clustercysts” (∼60 nm).39,40 Finally, when seeded, the clustercysts gather around the seeds, and then the crystal 7158
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Figure 5. Low-temperature pH measurement of amino acid buffer solutions as a function of time. The solutions were buffered to the desired pH at RT, cooled to 0 °C, equilibrated for 15 min, cooled further to 25 °C, and annealed for 70 min. The cooling rate was at 0.5 °C/min. Both the bath (9) and sample (b) temperatures were recorded.
growth proceeds through the steps of transfer, aggregation, and regular rearrangement of clustercysts.38 While the phase separation of citric acid resulted in an increase in the freeze-concentrate pH, we did not have an unambiguous evidence of solute crystallization. Amino Acid Buffers: Characterization of Frozen Solutions. Both histidine and glycine are amphoteric molecules. The pKa values of histidine are 1.8 (carboxylic acid), 6.0 (imidazole side chain), and 9.2 (amino group) and those of glycine are 2.3 (carboxylic acid) and 9.6 (amino group). The isoelectric points of histidine and glycine are 7.6 and 6.0, respectively. When the carboxylic acid buffer solutions were cooled, between 25 and 0 °C, they exhibited only a small change in pH of the unfrozen solution. This is expressed as ΔpHUF (pH25 °C pH0 °C) in Table 1. On the other hand, the corresponding pH change in the amino acid buffers was much more pronounced. As pointed out earlier, the effect of temperature on pKa is pronounced for nitrogenous bases but not for carboxylic acids. Interestingly, ΔpHUF = 0 for glycine solution buffered to pH 3.0, but when buffered to pH values of 9.0 and 9.5, the ΔpHUF were 0.3 and 0.4 units, respectively. Figure 5 contains the temperature and pH of the amino acid buffer solutions during the cooling and annealing temperature program. When the histidine solution buffered to pH 5.5 was
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cooled, there was first a modest decrease in pH (0.4 units) followed by a pronounced increase (1.8 units). The initial decrease can be attributed to the effect of the increase in temperature brought about by ice crystallization and the consequent increase in solute concentration (Figure 5, panel a). This was followed by a pronounced pH increase, brought about by the crystallization of acidic buffer component (L-histidine HCl). The DSC heating curves and the XRD patterns of the frozen amino acid buffer solutions are given in Figures 68). XRD patterns (Figure 7) showed the characteristic peaks of L-histidine HCl. However, the DSC heating curve (Figure 6, panel a) did not show any thermal event except ice melting. The possibility of the overlapping of the L-histidine HClice eutectic melting with that of ice cannot be ruled out. Similarly, when the solution buffered to pH 6.0 was cooled, there was a marked increase in the freezeconcentrate pH, indicating crystallization of histidine HCl. Again, while the characteristic peaks of L-histidine HCl and Lhistidine were evident in XRD, no thermal event indicative of solute crystallization was observed in the DSC heating curve. In the case of the system buffered to pH 7.0 (RT), there was a sharp decrease in the pH of the freeze-concentrate (∼0.7 units), indicating crystallization of the basic (L-histidine) buffer component. This was followed by a gradual continuous but small decrease in pH. In the DSC heating curve, there was an exotherm (13.2 °C) likely due to crystallization, followed by ice melting endotherm. In spite of the recrystallization, no discernible eutectic melting was observed. This indicated that the crystallized solute melted along with ice. DSC may not be a suitable characterization technique for such systems. In the case of both succinic acid and L-histidine, in spite of solute crystallization, we do not see any evidence of eutectic melting in the DSC. Eutectic melting close to ice melting is possible. Again, in these cases, DSC may not be a suitable characterization technique. However, XRD provided direct evidence of L-histidine crystallization in frozen solutions. In addition, when the frozen histidine buffer (pH 7.0) solution was thawed, there were visible particulates, suggesting precipitation. Chang and Randall, while studying the thermal behavior of frozen L-histidine (not a buffer system) solution reported that it remained amorphous with Tg00 and Tg0 values of 42 and 33 °C, respectively.41 However, in addition to the Tg0 , a recrystallization exotherm at 10 °C was also reported by Osterberg and Wadsten.42 They observed large crystals in the open DSC sample pan during thawing. In spite of the unambiguous XRD evidence of L-histidine crystallization in frozen solution, DSC did not provide any evidence of solute crystallization. When the glycine solution buffered to pH 3.0 was cooled, there was first a decrease (0.3 units) and then an increase (0.5 units) in the pH of the freeze-concentrate (Figure 5, panel b). This can be explained by the sequential crystallization of the basic (β-glycine) and acidic (diglycine HCl) buffer components. In the DSC, a glass transition at 24.8 °C was followed by two endotherms due to eutectic melting at 7.7 °C and ice melting at 1.8 °C (Figure 6, panel b). XRD pattern recorded at 16 °C showed peaks of both β-glycine (2θ 18.1°, 19.2°, and 20.8°) and diglycine HCl (2θ 22.1°, 26.7°, and 27.5°; Figure 8). When buffer solution with an initial pH value of 9.0 was cooled, there was a pronounced increase in the freeze-concentrate pH (∼2.5 units) indicating crystallization of the acidic buffer component (glycine). The DSC heating curve showed a glass transition at 23.6 °C, followed by eutectic and ice melting at 6.2 and 1.7 °C, respectively. XRD revealed the crystallization of β-glycine (data not shown). The glass transition indicated that only a fraction of glycine had crystallized during cooling. Similar, but less 7159
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Figure 6. DSC heating curves of frozen amino acid buffer solutions. The solutions were buffered to the desired pH at RT, cooled to 50 °C, held for 15 min, and warmed back to RT. The heating and cooling rate was 1 °C/min. Only the final heating curve is shown.
Figure 7. XRD pattern of frozen histidine buffer solutions. The solutions were initially cooled from RT to 10 °C, held for 10 min, and then cooled to 25 °C while collecting XRD patterns at 2 °C intervals. The frozen solution was annealed at 25 °C for 60 min. The cooling and heating rate was 1 °C/min. Only the XRD patterns of the annealed samples are shown. The characteristic peaks of L-histidine HCl and L-histidine are pointed out.
pronounced, pH shift (∼1.5 units) was observed when the glycine buffered to pH 9.5 was cooled. When this frozen solution was warmed in the DSC, a glass transition at 22.2 °C and a weak eutectic melting at 7.4 °C were evident. The less pronounced pH shift and a weak eutectic melting, coupled with the poorly crystalline nature of the solute observed in the XRD (not shown), indicated that a significant fraction of glycine had remained amorphous. Prediction of Crystallization Propensity. The intrinsic solubility (neutral form, C0) of the carboxylic and amino acids at various temperatures, reported in the literature, was used to
construct the van’t Hoff plot (Figure 9).4345 Over the temperature range of 400 °C, there was an excellent linear relationship between solubility (logarithmic scale) and reciprocal temperature. The solubility at 25 °C for each compound was predicted by extrapolating this linear relationship (Table 2). While we have assumed linearity of the van’t Hoff plot over a wide temperature range, in several organic compounds, this is not the case.46 Thus, the predicted solubility values, while only approximate, will be a reasonable estimate. The calculated enthalpy of solution, reported in Table 2, was obtained from the slope of the van’t Hoff 7160
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Figure 8. XRD pattern of frozen glycine buffer solutions (pH 3.0 at RT) recorded during cooling and annealing. The solution was first cooled to 10 °C, held for 10 min, and then cooled to 25 °C while collecting XRD patterns at 2 °C intervals. The system was annealed at 25 °C for 45 min. The cooling and heating rate was 1 °C/min. The characteristic peaks of diglycine HCl and β-glycine are pointed out.
Figure 9. van’t Hoff plots of carboxylic and amino acids. The solubility values were obtained from the literature.4345,48,49 The plots of the carboxylic acids were published earlier and are reproduced with permission of the copyright owner.27 SA, succinic acid; TA, DL-tartaric acid; CA, citric acid anhydrate; MA, DL-malic acid.
plots. These were in reasonably good agreement with the reported experimental values (at 25 °C).43 Under our experimental conditions, in all the buffer systems, the neutral form is the least soluble species and is therefore expected to crystallize first. Since crystallization from the frozen system was monitored at 25 °C, it was also of interest to estimate the degree of supersaturation of the neutral species at this temperature. When the initial solution pH values were pKa2 ; at lower pH values; the crystallization propensity is higherÞ
From Figure 9 (data in Table 2), it is evident that the solubility of succinic and tartaric acids at RT is substantially lower than that of citric and malic acids. More importantly, the decrease in
solubility as a function of temperature for succinic and tartaric acids (manifested by the heat of solution values; Table 2) is much more pronounced than that of the two acids. Thus, the lower solubility coupled with the pronounced temperature effect on solubility makes these systems vulnerable to crystallization at subzero temperatures. In the case of glycine and histidine, the solubility is approximately constant in the pH range studied. Thus, both of them crystallized at all the pH values. Significance and Practical Implications. Solutions are buffered to maintain them at the desired pH value. When buffered solutions are frozen, either as a long-term stabilization strategy or as the first step in the freeze-drying process, any undesirable pH shift brought about by the buffer component crystallization defeats the very purpose of buffering the solution. Effective prediction of buffer component crystallization propensity at subzero temperatures was possible from the van’t Hoff plots coupled with the calculated degree of supersaturation. This work will facilitate the selection of a suitable buffer while avoiding the problem of buffer crystallization and the consequent pH shift in frozen solutions. Freezing and freeze-drying of buffered solutions is of interest in many fields including food and pharmaceutical sciences, biotechnology, and cryobiology. Our approach can be extended to evaluate the crystallization propensity of other buffer systems.
’ CONCLUSIONS Crystallization of buffer components in frozen aqueous solutions and the consequent pH shifts were studied in carboxylic and amino acid buffers. Solute crystallization in frozen systems and identification of the crystallizing phase was accomplished using DSC and XRD, respectively. Based on the estimated degree of supersaturation at the subzero temperature of interest, the crystallization propensity in frozen systems was predicted and rank ordered. The experimental and the predicted rank orderings were in excellent agreement. ’ AUTHOR INFORMATION Corresponding Author
*Phone: 612-624-9626. Fax: 612-626-2125. E-mail: surya001@ umn.edu. 7162
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’ ACKNOWLEDGMENT The XRD studies were carried out in the College of Science and Engineering Characterization Facility, University of Minnesota, which receives partial support from NSF through the MRSEC program. Support from the William and Mildred Peters Endowment Fund is gratefully acknowledged. ’ REFERENCES (1) Bates, R. G. Determination of pH: Theory and Practice; John Wiley & Sons, Inc.: New York, 1973. (2) Galster, H. pH Measurement: Fundamentals, Methods, Applications, Instrumentation; VCH Publishers, Inc.: New York, 1991. (3) Trissel, L. A. Handbook on Injectable Drugs, 9th ed.; American Society of Health-System Pharmacists: Bethesda, MD, 1996. (4) Waterman, K. C.; Adami, R. C.; Alsante, K. M.; Antipas, A. S.; Arenson, D. R.; Carrier, R.; Hong, J.; Landis, M. S.; Lombardo, F.; Shah, J. C.; Shalaev, E.; Smith, S. W.; Wang, H. Hydrolysis in Pharmaceutical Formulations. Pharm. Dev. Technol. 2002, 7, 113–146. (5) Xu, Q.; Trissel, L. Stability-Indicating HPLC Methods for Drug Analysis, 2nd ed.; Pharmaceutical Press: London, 2003. (6) Benedict, R. G.; Schmidt, W. H.; Coghill, R. D.; Oleson, A. P. Penicillin III. The Stability of Penicillin in Aqueous Solution. J. Bacteriol. 1945, 49, 85–95. (7) Kheirolomoom, A.; Kazemi-Vaysari, A.; Ardjmand, M.; BaradarKhoshfetrat, A. The Combined Effects of pH and Temperature on Penicillin G Decomposition and Its Stability Modeling. Process Biochem. (Oxford) 1999, 35, 205–211. (8) Beynon, R. J.; Easterby, J. S. Buffer Solutions: The Basics; Oxford University Press: New York, 1996. (9) Akers, M. J. Excipient - Drug Interactions in Parenteral Formulations. J. Pharm. Sci. 2002, 91, 2283–2300. (10) Shalaev, E. Y.; Wang, W.; Gatlin, L. A. Rational Choice of Excipients for Use in Lyophilized Formulations. Drugs Pharm. Sci. 2008, 175, 197–217. (11) Pikal, M. J. Freeze Drying. In Encyclopedia of Pharmaceutical Technology; Swarbrick, J., Boylan, J. C., Eds.; Marcel Dekker: New York, 2002; pp 12991326. (12) Larsen, S. S. Stability of Drugs in Frozen Systems. VI. Effect of Freezing Upon pH for Buffered Aqueous Solutions. Arch. Pharm. Chemi, Sci. Ed. 1973, 1, 41–53. (13) van den Berg, L. pH Changes in Buffers and Foods During Freezing and Subsequent Storage. Cryobiology 1966, 3, 236–242. (14) Gomez, G.; Pikal, M. J.; Rodriguez-Hornedo, N. Effect of Initial Buffer Composition on pH Changes During Far-from-Equilibrium Freezing of Sodium Phosphate Buffer Solutions. Pharm. Res. 2001, 18, 90–97. (15) Pikal-Cleland, K. A.; Cleland, J. L.; Anchordoquy, T. J.; Carpenter, J. F. Effect of Glycine on pH Changes and Protein Stability During Freeze-Thawing in Phosphate Buffer Systems. J. Pharm. Sci. 2002, 91, 1969–1979. (16) Sundaramurthi, P.; Shalaev, E.; Suryanarayanan, R. “pH Swing” In Frozen Solutions-Consequence of Sequential Crystallization of Buffer Components. J. Phys. Chem. Lett. 2010, 1, 265–268. (17) van den Berg, L. The Effect of Addition of Sodium and Potassium Chloride to the Reciprocal System: KH2PO4Na2HPO4 H2O on pH and Composition During Freezing. Arch. Biochem. Biophys. 1959, 84, 305–315. (18) van den Berg, L.; Rose, D. Effect of Freezing on the pH and Composition of Sodium and Potassium Phosphate Solutions: The Reciprocal System KH2PO4Na2HPO4H2O. Arch. Biochem. Biophys. 1959, 81, 319–329. (19) Murase, N.; Franks, F. Salt Precipitation During the FreezeConcentration of Phosphate Buffer Solutions. Biophys. Chem. 1989, 34, 293–300. (20) Shalaev, E. Y.; Johnson-Elton, T. D.; Chang, L.; Pikal, M. J. Thermophysical Properties of Pharmaceutically Compatible Buffers at
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Sub-Zero Temperatures: Implications for Freeze-Drying. Pharm. Res. 2002, 19, 195–201. (21) Gomez, G. Crystallization-Related pH Changes During Freezing of Sodium Phosphate Buffer Solutions. Ph.D. Thesis, University of Michigan, 1995. (22) Szkudlarek, B. A. Selective Crystallization of Phosphate Buffer Components and pH Changes During Freezing: Implications on Protein Stability. Ph.D. Thesis, University of Michigan, 1997. (23) Orii, Y.; Morita, M. Measurement of the pH of Frozen Buffer Solutions by Using pH Indicators. J. Biochem. 1977, 81, 163–168. (24) Cavatur, R. K.; Suryanarayanan, R. Characterization of Frozen Aqueous Solutions by Low Temperature X-Ray Powder Diffractometry. Pharm. Res. 1998, 15, 194–199. (25) Varshney, D. B.; Kumar, S.; Shalaev, E. Y.; Kang, S.-W.; Gatlin, L. A.; Suryanarayanan, R. Solute Crystallization in Frozen Systems-Use of Synchrotron Radiation to Improve Sensitivity. Pharm. Res. 2006, 23, 2368–2374. (26) Varshney, D. B.; Sundaramurthi, P.; Kumar, S.; Shalaev, E. Y.; Kang, S.-W.; Gatlin, L. A.; Suryanarayanan, R. Phase Transitions in Frozen Systems and During Freeze-Drying: Quantification Using Synchrotron X-Ray Diffractometry. Pharm. Res. 2009, 26, 1596–1606. (27) Sundaramurthi, P.; Suryanarayanan, R. Predicting the Crystallization Propensity of Carboxylic Acid Buffers in Frozen SystemsRelevance to Freeze-Drying. J. Pharm. Sci. 2011, 100, 1288–1293. (28) Sundaramurthi, P.; Shalaev, E.; Suryanarayanan, R. Calorimetric and Diffractometric Evidence for the Sequential Crystallization of Buffer Components and Consequent pH Swing in Frozen Solutions J. Phys. Chem. B 2010, 114, 4915–4923. (29) Powder Diffraction File. Hexagonal ice, Card # 00-042-1142; L-histidine HCl, Card # 00-005-0459; DL-histidine HCl, Card # 00-0050538; D-histidine, Card # 00-031-1716; DL-histidine, Card # 00-0311717; histidine HCl monohydrate, Card # 00-054-1737; L-malic acid, Card # 00-023-1631; DL-tartaric acid, Card # 00-020-1901; D-tartaric acid, Card # 00-031-1912; citric acid, Card #00-016-1157. International Centre for Diffraction Data: Newtown Square, PA, 2004. (30) Goodno, C. C.; Swenson, C. A.; Bull, H. B. A Mathematical Method for Temperature Compensation of pH Measurements. Anal. Biochem. 1975, 67, 220–225. (31) Albert, A.; Serjeant, E. P. The Determination of Ionization Constants: A Laboratory Manual; Chapman and Hall: London. 1971. (32) Pinching, G. D.; Bates, R. G. First Dissociation Constant of Succinic Acid from 0° to 50°C and Related Thermodynamic Quantities. J. Res. NBS 1950, 45, 444–449. (33) Pinching, G. D.; Bates, R. G. Second Dissociation Constant of Succinic Acid from 0° to 50°C. J. Res. NBS 1950, 45, 322–328. (34) Sundaramurthi, P.; Suryanarayanan, R. Trehalose Crystallization During Freeze-Drying: Implications on Lyoprotection. J. Phys. Chem. Lett. 2010, 1, 510–514. (35) Brock, C. P.; Schweizer, W. B.; Dunitz, J. D. On the Validity of Wallach’s Rule: On the Density and Stability of Racemic Crystals Compared with Their Chiral Counterparts. J. Am. Chem. Soc. 1991, 113, 9811–9820. (36) Hoppu, P.; Hietala, S.; Schantz, S.; Juppo, A. M. Rheology and Molecular Mobility of Amorphous Blends of Citric Acid and Paracetamol. Eur. J. Pharm. Biopharm. 2009, 71, 55–63. (37) Hoppu, P.; Jouppila, K.; Rantanen, J.; Schantz, S.; Juppo, A. M. Characterization of Blends of Paracetamol and Citric Acid. J. Pharm. Pharmacol. 2007, 59, 373–381. (38) Ueda, M.; Hirokawa, N.; Harano, Y.; Moritoki, M.; Ohgaki, K. Change in Microstructure of an Aqueous Citric Acid Solution under Crystallization. J. Cryst. Growth 1995, 156, 261–266. (39) Ohgaki, K.; Hirokawa, N.; Ueda, M. Heterogeneity in Aqueous Solutions: Electron Microscopy of Citric Acid Solutions. Chem. Eng. Sci. 1992, 47, 1819–1823. (40) Ohgaki, K.; Makihara, Y.; Morishita, M.; Ueda, M.; Hirokawa, N. Solute Clusters in Aqueous Citric Acid Solutions. Chem. Eng. Sci. 1991, 46, 3283–3287. (41) Chang, B. S.; Randall, C. S. Use of Subambient Thermal Analysis to Optimize Protein Lyophilization. Cryobiology 1992, 29, 632–656. 7163
dx.doi.org/10.1021/jp202167p |J. Phys. Chem. B 2011, 115, 7154–7164
The Journal of Physical Chemistry B
ARTICLE
(42) Osterberg, T.; Wadsten, T. Physical State of L-Histidine after Freeze-Drying and Long-Term Storage. Eur. J. Pharm. Sci. 1999, 8, 301–308. (43) Apelblat, A.; Manzurola, E. Solubility of Oxalic, Malonic, Succinic, Adipic, Maleic, Malic, Citric, and Tartaric Acids in Water from 278.15 to 338.15 K. J. Chem. Thermodyn. 1987, 19, 317–320. (44) Bhagwat, W. V.; Doosaj, S. S. Limitations of Solubility Method for Determining Dissociation Constant. J. Indian Chem. Soc. 1933, 10, 477–490. (45) Doosaj, S. S.; Bhagwat, W. V. Solubilities of Weak Acids in Salts of Weak Acids at Very High Concentrations. J. Indian Chem. Soc. 1933, 10, 225–232. (46) Grant, D. J. W.; Mehdizadeh, M.; Chow, A. H.-L.; Faribrother, J. E. Non-Linear van’t Hoff Solubility-Temperature Plots and Their Pharmaceutical Interpretation. Int. J. Pharm. 1984, 18, 25–38. (47) Towler, C. S.; Davey, R. J.; Lancaster, R. W.; Price, C. J. Impact of Molecualar Speciation on Crystal Nucleation in Polymorphic Systems: The Conundrum of γ-glycine and Molecular “Self Poisoning. J. Am. Chem. Soc. 2004, 126, 13347–13353. (48) Dunn, M. S.; Ross, F. J.; Read, L. S. The Solubility of the Amino Acids in Water. J. Biol. Chem. 1933, 103, 579–595. (49) Mullin, J. W. Crystallization, 4th ed.; Butterworth-Heinemann: Oxford, UK, 2001.
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