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Enthalpies of solution, limiting solubilities, and partial molar heat capacities of n-alcohols in water and in trehalose crowded media

  • Guangyue Bai , Sandra C.C. Nunes , Marisa A.A. Rocha , Luís M.N.B.F. Santos , M. Ermelinda S. Eusébio , M. João Moreno and Margarida Bastos EMAIL logo
Published/Copyright: January 17, 2014
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Pure and Applied Chemistry
From the journal Pure and Applied Chemistry Volume 86 Issue 2

Abstract

The enthalpies of solution of alcohols were determined by calorimetry in HEPES and (HEPES + trehalose) at 298.15 K. The used methodology and experiment’s design allowed us to extract from a single titration experiment the enthalpy of solution (ΔsolHm), the limiting solubility of the alcohol in each aqueous media, and an estimation of the enthalpy of solution of water in the alcohol phase. From these values the changes in Gibbs energy (ΔsolGm) and in entropy (ΔsolSm) of solution were derived. A decrease in solubility for 1-butanol and 1-pentanol in the crowded media (HEPES + trehalose) was observed which is driven by a significant decrease in the favorable enthalpy of solution. The partial molar heat capacity, in each media was determined in our heat capacity drop calorimeter, also at 298.15 K. A significant decrease of the partial molar heat capacity was observed for both alcohols in (HEPES + trehalose), which together with the obtained decrease in favorable ΔsolHm, is consistent with a decrease in hydrophobic solvation, as a result of a decrease in free solvent availability induced by the trehalose. Finally, we tentatively predict that in the aqueous media of the crowded solutions that characterize cells and biological fluids, solutes with low aqueous solubility will be more soluble, whereas the solubility of highly polar solutes will be reduced.

Keywords: alcohol; calorimery; crowded media; enthalpy of solution; isothermal titration calorimetry (ITC); solubility

Introduction

Most of the thermodynamic data in solution for simple compounds is obtained in the high dilution region, in many cases extrapolated to infinite dilution [1–11]. This choice was for many years related to their use, as well as to accepted standard states’ definitions. Nevertheless, in the biological domain the situation is quite different. The cytoplasm is an aqueous solution with a very high concentration of polar solutes, small molecules, macromolecules, and supramolecular assemblies. At these high solute concentrations only about half the total water molecules behave as bulk water, whereas the remaining are somehow involved in the solvation of the solute molecules. This medium is usually referred to as a solution with “molecular crowding”. Under these conditions, the structure and dynamics of solutes and supramolecular assemblies is significantly different from that found in dilute aqueous solutions. Interactions resulting from electrostatics and the hydrophobic effect are dependent on the properties of the solvent (water) and thus in these circumstances will be strongly affected by the reduced availability of free water due to molecular crowding. Due to lack of appropriate data, most available biochemical and biophysical characterizations of the solute behavior inside a cell are derived from information taken from studies performed in the dilute regime.

In recent years, some efforts have been made towards the study of the effects of crowded media on the structure and dynamics of macromolecules. It has been shown that the presence of small solutes that interact stronger with water than with the protein surface (salting out agents) increases the stability of the protein folded state due to a tendency to reduce its surface area exposed to the solvent [12]. The diffusion-dependent dynamics of species from small solutes to proteins and DNA was shown to be affected by molecular crowding [13]. The binding to proteins and the kinetics of enzyme-catalyzed reactions is also predicted to be strongly altered due both to changes in the protein and substrate structure and to anomalous diffusion [12, 14, 15].

Under conditions of stress (heat and desiccation) cells accumulate small molecules, to reduce the water availability even further. An example of this is the up to 100-fold increase in trehalose concentration in yeasts as a response to heat shock [16]. Under those conditions, the concentration of trehalose may reach 1 g per g of protein. Aqueous solutions with high concentrations of this disaccharide have been used as a model system for solutions with “molecular crowding”.

In order to assess the effect of “molecular crowding” on small solutes’ thermodynamic properties we used HEPES buffer alone or with a high trehalose concentration as our solvent media to characterize the solution behavior of n-alcohols at 298.15 K. The enthalpies of solution (ΔsolH) were determined by solution calorimetry, and for 1-butanol and 1-pentanol we could also estimate their solubility in each medium through isothermal titration calorimetry (ITC) measurements. From these the thermodynamic properties, ΔsolG and ΔsolS were calculated for each alcohol. Finally, the partial molar heat capacities of these alcohols were directly measured in each media by a high-precision heat capacity drop calorimeter.

Materials and methods

Chemicals

The alcohols used, 1-propanol (ACS reagent, 99.4 % pure, St. Louis, USA), 1-butanol (anhydrous, 99.8 %, St. Louis, USA), and 1-pentanol (puriss. pro analysi, >99.0 % pure determined by GC, St. Louis, USA) were purchased from Sigma-Aldrich and used without further purification. Trehalose was from starch, Sigma (≥99 % pure, St. Louis, USA), and HEPES, Sigma (≥99.5 % pure). Buffer was prepared at a total concentration of 10 mmol·dm–3 and pH = 7.4, 0.02 % NaN3, 1 mM EDTA, and 150 mM NaCl. The aqueous solutions containing trehalose at a concentration of 1 mol·dm–3 (HEPES + trehalose) were prepared via the dissolution of the necessary amount of all solid components in the amount of water needed to perform the total required volume. The solutions were filtered through a 1-μm nylon filter and stored at 277 K.

KCl Panreac (x > 0.995, Barcelona, Spain) was purified according to the procedure recommended by Goffredi et al. [17] for this standard. NaCl Riedel deHaen (x > 99.8 %, Seelze, Germany) was dried at 773 K for 24 h, and used thereafter.

The solutions to be used in the heat capacity drop calorimeter [1-butanol and 1-pentanol, in HEPES buffer and in (HEPES + trehalose)] were prepared by weight at room temperature, 1 day before the measurements and were kept at constant stirring overnight. An analytical balance Mettler AT220 with mass resolution of (1 × 10–5) g and weighting uncertainty of ±(2 × 10–5) g was used.

All the uncertainties are given as twice of the standard deviation of the mean. The relative atomic masses used were those recommended by the IUPAC Commission in 2009 [18].

Determination of enthalpies of solution and limiting solubilities

Enthalpies of solution of 1-propanol in HEPES and in (HEPES + trehalose) were determined at T = 298.15 K, using a Calvet-type calorimeter Setaram, model C80, and the cell assembly described by Tomé et al. [19]. The amounts of solute and solvent in the calorimeter were determined by weight, using an analytical balance Mettler AE 163 with mass resolution of (1 × 10–5) g and weighting uncertainty of ±(2 × 10–5) g.

The calorimeter calibration was carried out using the enthalpy of solution of KCl in water [20, 21], (500 H2O) = (17.584 ± 0.017) kJ·mol−1. The calibration procedure was checked by determining the enthalpy for the dissolution of NaCl in water; five experimental determinations were made for solutions with concentrations ranging from (0.13 to 0.15) mol·kg–1. The result obtained, ∆solHm = (4.15 ± 0.13) kJ·mol–1 is in excellent accordance with literature values [22].

The enthalpies of solution of 1-butanol and 1-pentanol in HEPES buffer and in (HEPES + trehalose) were determined by ITC, due to the lower solubility of these alcohols in both media. The microcalorimeter unit used in this work consists of a twin heat conduction calorimeter equipped with a 1- or 3-mL titration cell (ThermoMetric AB, Järfalla, Sweden), a high-precision temperature controller, developed and constructed at the Lund University, Sweden. The calorimetric signal was measured by means of a 7½ digit nanovoltmeter, Hewlett-Packard model 34420A, and the data was recorded and visualized in real time using a modified version of the software application LABTERMO. This calorimeter unit, instrumentation, as well as the experimental methodology is already described in detail in the literature [23].Briefly, the volume of solvent in the calorimetric vessel was either 0.9 or 2.6 mL. The calorimetric titration experiments consisted of a series of consecutive additions of the alcohol into the solvent media contained in the calorimetric vessel. The titrating solution (pure alcohol) was added to the vessel automatically, in aliquots of volumes between 1.6 and 4 μL, from a modified gas-tight Hamilton syringe (100 or 500 μL), through a thin stainless-steel capillary, until the desired range of concentrations had been covered. A gold propeller was used throughout, with stirring rates between 120 and 140 rpm.The experiments were performed either in “fast mode” (injections every 7 min) [24] or in normal, slow mode (injections every 40 min). When working in “fast mode”, the curves were dynamically corrected [24] and integrated thereafter using the Origin 7 software. The heat evolved per injection was calculated using the electrically determined calibration constant [23]. All experiments were carried out at 298.15 ± 0.01 K.

Determination of partial molar heat capacity of the solute

The heat capacities at 298.15 K for HEPES buffer, (HEPES + trehalose) and solutions of 1-butanol and 1-pentanol in both media were determined using a heat capacity drop calorimeter, described in detail in the literature [25, 26] and recently rebuilt in our lab [27]. The receiving calorimeter consists of a twin heat conduction calorimeter and was operated between 303.15 K (temperature of the furnace) and 293.15 K (temperature of the receiving calorimeter). Therefore, the twin calorimeter measures the heat exchange resulting from the ampoule cooling from the initial temperature (Ti = 303.15 K) to the final temperature (Tf = 293.15 K), and the measured heat capacity can be assigned to the mean temperature, 298.15 K. In the present work, the apparatus was used in single-drop mode (no reference ampoule was used) with blank correction that were measured independently using empty ampoules. The calorimeter was calibrated with water and sapphire (α-Al2O3), using the respective standard molar heat capacities at 298.15 K reported in the literature [21], (H2O) = (75.32 ± 0.01) J·K–1·mol–1 and (α-aluminum oxide) = (79.03 ± 0.08) J·K–1·mol–1. The calibration constant was found to be ε = (6.6143 ± 0.0048) W·V–1.

Results and discussion

Enthalpies of solution and limiting solubilities

1-Propanol was soluble in both media at all concentrations. The values obtained at 298.15 K for the enthalpy of solution, at infinite dilution, of 1-propanol in HEPES and in (HEPES + trehalose) using the Calvet-type calorimeter, Setaram C80, are presented in Table 1. The enthalpy of solution obtained when using HEPES as the solvent media is in very good agreement with the one determined in water, –(10.16 ± 0.02) kJ·mol–1 [3]. This indicates that the buffer itself does not affect the enthalpy of solution of the alcohol.

Table 1

Enthalpies of solution, at infinite dilution, ΔsolHm∞, for the three alcohols in both media, at 298.15 K.

Alcohol HEPES
 HEPES + trehalose
ΔsolHm∞/kJ·mol–1 Solubility/mol·dm–3 ΔsolHm∞/kJ·mol–1 Solubility/mol·dm–3
1-Propanol –10.0 ± 0.2 – –6.0 ± 0.3 –
1-Butanol –9.2 ± 0.5 0.78 –4.7 ± 0.6 0.13
1-Pentanol –8.1 ± 0.8 0.18 –2.8 ± 0.9 0.07

Similar values for 1-butanol and 1-pentanol were obtained by ITC as described above. For both alcohols and in the two media, each titration experiment showed initially a number of exothermic peaks, followed by a “transition range” where the peaks had exothermic and endothermic components, and ended in a series of endothermic peaks. The number of exothermic peaks before the transition range depended on alcohol, solvent media, and injected volume. An example of a typical titration profile is depicted in Fig. 1.

Fig. 1 Typical titration of aliquots of 1-butanol into (HEPES + trehalose) at 298.15 K, obtained by ITC.
Fig. 1

Typical titration of aliquots of 1-butanol into (HEPES + trehalose) at 298.15 K, obtained by ITC.

The first part of the titration curve reflects the enthalpy of solution of the alcohol into the aqueous media, until saturation occurs, giving rise to the “transition region”, where a mixture of endothermic and exothermic components can be seen in each peak. The decrease in enthalpy per injection is expected, due to the onset of solute–solute interactions upon increase in solute concentration [2, 3]. Therefore, the enthalpy of solution at infinite dilution, ΔsolHm∞, for 1-butanol and 1-pentanol in HEPES and (HEPES + trehalose) were taken as the ΔsolHm values at zero alcohol concentration, using the results before saturation. The solubilities for 1-butanol and 1-pentanol in each media were taken as the values of alcohol concentration at the beginning of the saturation zone. After saturation, the endothermic heat effect observed could be ascertained to the enthalpy of solution of water into the alcohol phase.

The heat effect per peak, i.e., the enthalpy of solution (in J), together with the saturation range can be seen in Fig. 2 for the titration shown above, 1-butanol into (HEPES + trehalose). Similar plots were obtained for 1-butanol in HEPES and for 1-pentanol in both solvent systems.

Fig. 2 Enthalpy as a function of number of moles of alcohol added for the titration of aliquots of 1-butanol into (HEPES + trehalose) at 298.15 K.
Fig. 2

Enthalpy as a function of number of moles of alcohol added for the titration of aliquots of 1-butanol into (HEPES + trehalose) at 298.15 K.

The obtained values of the enthalpy of solution, at infinite dilution, ΔsolHm∞, as well as the solubilities for these two alcohols in both solvent media are presented in Table 1. Similarly to what was found for 1-propanol, the results for ΔsolHm∞ in HEPES for 1-butanol and 1-pentanol are identical to those previously determined in water [3], –9.27 ± 0.02 and –7.99 ± 0.05 kJ·mol–1, respectively.

Fig. 3 Results for the enthalpy of solution, at infinite dilution, ΔsolHm∞, for the three studied alcohols in HEPES () and in (HEPES + trehalose) (), together with literature values [3] for ΔsolHm∞ of n-alcohols (from ethanol to 1-nonanol) in water (). All results are at 298.15 K.
Fig. 3

Results for the enthalpy of solution, at infinite dilution, ΔsolHm∞, for the three studied alcohols in HEPES () and in (HEPES + trehalose) (), together with literature values [3] for ΔsolHm∞ of n-alcohols (from ethanol to 1-nonanol) in water (). All results are at 298.15 K.

The results obtained here for the enthalpy of solution, at infinite dilution, ΔsolHm∞ for the three alcohols in both solvent media are plotted together with literature data in Fig. 3.

The solubility results derived from the ITC measurements in HEPES are in reasonable agreement with solubility data available in the literature for these alcohols in water [28], 0.93 and 0.24 mol·dm–3 for 1-butanol and 1-pentanol, respectively. This shows that this method can be used to extract solubility values from an ITC run where at the same time the enthalpy of solution is determined, and substantiates the values determined for (HEPES + trehalose), for which there is no literature data.

Finally, the solubilities were used to obtain the change in Gibbs energy of solution, ΔsolGm∞, which together with ΔsolHm∞ allowed the calculation of the change in entropy of solution, ΔsolSm∞, for 1-butanol and 1-pentanol in both media. The values of these thermodynamic properties can be found in Table 2.

Table 2

Thermodynamic properties of solution ΔsolHm, ΔsolGm and TΔsolSm for the 1-butanol and 1-pentanol in the two solvent media used, at 298.15 K.

Alcohol HEPES
 HEPES + trehalose
ΔsolHm
 ΔsolGm
 TΔsolSm
 ΔsolHm
 ΔsolGm
 TΔsolSm
kJ·mol–1 kJ·mol–1 kJ·mol–1 kJ·mol–1 kJ·mol–1 kJ·mol–1
1-Butanol –9.2 ± 0.5 0.59 –9.8 –4.7 ± 0.6 5.1 –9.8
1-Pentanol –8.1 ± 0.8 4.2 –12.4 –2.8 ± 0.9 6.6 –9.4

As the solubility is estimated, we do not assign uncertainty to ΔsolGm, and consequently neither to ΔsolSm.

We can see that there is a significant decrease in solubility and in the absolute value of ΔsolHm when changing from HEPES to (HEPES + trehalose), whereas the corresponding entropic contribution to the transfer Δ(TΔsolSm) is about zero for 1-butanol and positive for 1-pentanol (+3 kJ·mol–1). Although we did not assign uncertainties to these values, for the reasons described above, the values for the entropic contribution are probably equal within uncertainty. The interpretation of the entropy of transfer is complex, as together with the solvent structuring effect one also needs to consider the difference in cavity formation in each media, whereas in terms of cp this last effect is not so problematic, and we will see below that the effect is clearer for this later thermodynamic property. Balancing the relative contributions of the ΔsolHm and TΔsolSm to the Gibbs energy change we can conclude that it is the decrease in favorable enthalpy of solution that is mainly responsible for the observed decrease in solubility in the crowded media. It is well known that alcohols can interact with water through hydrogen bonds involving the OH group, and for higher alcohols (i.e., long alkyl chain) hydrophobic interactions also play a role [3, 29]. The observed decrease in absolute value of ΔsolHm as the alkyl chain increases for n(CH2) > 3 (see Fig. 3) reflects the decrease in relative importance of the exothermic contribution of hydrogen bond formation to the enthalpy of solution. The observed decrease in absolute value of ΔsolHm in the crowded media could thus reflect a decrease in water availability for hydrogen bond formation. Further, it is well established that the hydration of hydrocarbon moieties is characterized by large and negative entropy changes and large and positive heat capacity changes [29–31]. Accordingly, we did obtain a negative entropy change for the dissolution of the alcohols in HEPES, and a decrease for 1-pentanol in the negative magnitude of this property when we change the media to (HEPES + trehalose), whereas the value for 1-butanol is equal within the quoted uncertainty. This reflects the known fact that the importance of hydrophobic hydration is larger for higher alcohols.

From the average values of the heat absorbed in the positive peaks (after saturation) we could get an estimate of the value for the enthalpy of solution of water into the alcohol, provided the amount of water transferred to the organic phase was known. As the small amounts we worked with in the calorimetric cell precluded an easy separation of the two phases at the end of the experiment, we simulated outside the calorimetric cell its final conditions. Thus, we prepared an alcohol/solvent media mixture with the same proportion as found in the calorimetric cell at the end of a titration run, but in a larger total amount, where the phase separation was clearly seen. The water content in the organic phase was thereafter determined by Karl Fischer coulometry. The enthalpy values retrieved were 1.9 and 3.5 kJ·mol–1 for water into 1-butanol and 1-pentanol, respectively. Although these values represent a crude estimation, they are in reasonable agreement with the enthalpy of solution of water into 1-butanol and 1-pentanol, 1.68 ± 0.04 and 2.66 ± 0.04 kJ·mol–1, respectively, that were determined by ITC [32], in experiments where water was titrated into pure alcohol phase contained in the calorimetric cell at 298.15 K.

Partial molar heat capacity of the solute

The apparent molar heat capacities for the alcohol, Cp,φ, were calculated taking into account the specific heat capacities obtained for the solutions, according to the equation

where and are the specific heat capacities of the solution and pure solvent, respectively, M is the molar mass of the solute and m is the molality of the solution. The concentration range used in each case was chosen to span the concentration as much as possible while keeping it below the solubility limit in each case. The results for the heat capacity obtained for the solvents HEPES and (HEPES + trehalose) as well as for the alcohols solutions in each media are presented in Table 3.

Table 3

Specific heat capacities of the pure solvent, cp, and specific heat capacities solutions of 1-butanol and 1-pentanol in HEPES buffer and (HEPES + trehalose) at 298.15 K.

 cp/J·K–1·g–1a
Pure 1-Butanol 1-Pentanol
HEPES buffer 4.101 ± 0.005 4.124 ± 0.008 4.116 ± 0.002
(HEPES + trehalose) 3.433 ± 0.003 3.442 ± 0.004 3.444 ± 0.003

acpwas calculated as the mean value of the obtained results for each solution in the respective media, as no concentration dependence was found.

It can be seen that there is a very significant decrease in the heat capacity of the solvent media when changing from HEPES to (HEPES + trehalose). Since it is known that sugars have a high heat capacity both in water and in nonaqueous solvents [33], the observed decrease in heat capacity of the solvent media cannot be attributed to the presence of trehalose, but must reflect a decrease in hydrogen bond network of the solvent water due to the high trehalose concentration. From these values, the apparent molar heat capacities of the solutes were calculated according to eq. 1. Since no concentration dependence was observed, these values represent the partial molar heat capacity of the solute, at infinite dilution, The obtained values for 1-butanol and 1-pentanol in each solvent media are presented in Table 4.

Table 4

Apparent molar heat capacities at infinite dilution of 1-butanol and 1-pentanol, in HEPES buffer and (HEPES + trehalose) at 298.15 K.
1-Butanol 1-Pentanol
HEPES buffer 425 ± 16 551 ± 8
HEPES containing trehalose 329 ± 24 457 ± 35

The value of obtained in HEPES are in reasonable agreement with the ones determined in water by Hallén et al. [3], 441 ± 3 J·K–1·mol–1 and 532 ± 6 J·K–1·mol–1, and Makhatadze and Privalov [5] 445.9 ± 6.0 J·K–1·mol–1 and 539.5 ± 7.1 J·K–1·mol–1, for 1-butanol and 1-pentanol, respectively. This shows again that the buffer does not significantly affect the thermodynamic properties of solution of the alcohols. At odds, the values obtained for the same parameter in (HEPES + trehalose) show a significant decrease. The partial molar heat capacity of hydrophobic solutes in aqueous solution has been considered to reflect hydrophobic solvation, interpreted as resulting in an increased network of hydrogen bonds induced in the solvent by the presence of the hydrophobic solute [34, 35]. From the values above, the change in heat capacity on transfer of the alcohol from aqueous HEPES solution to (HEPES + trehalose) can be calculated. The values obtained for (transfer) were –(96 ± 29) and –(94 ± 36) J·K–1·mol–1 for 1-butanol and 1-pentanol, respectively. We should recall here that the unfavorable entropy change associated with the dissolution of the alcohols in the crowded media (HEPES + trehalose) does not increase from 1-butanol to 1-pentanol at odds with what happens in the absence of trehalose This observation (despite the uncertainty in the entropy change values), together with the observed decrease in is consistent with a decrease in hydrophobic hydration and a smaller contribution from hydrophobic solvation in this media. This interpretation goes in line with a significant decrease in the availability of bulk water in this media for a complete hydrophobic hydration.

Finally, these observations and their interpretation allow us to predict some trends in behavior in solutions with molecular crowding. It is to be expected that hydrogen bond forming molecules will show a lower solubility in aqueous media with “molecular crowding” (due to the reduction in water available) and nonpolar molecules an increase in the solubility (due to the decrease in hydrophobic solvation). The existing data for linear alcohols as a function of the length of the hydrocarbon chain show a minimum, thus a transition from solubility reduction to solubility increase appears. We therefore predict that in the aqueous media of the crowded solutions that characterize cells and biological fluids solutes with low aqueous solubility will be more soluble, whereas the solubility of highly polar solutes will be reduced.

Conclusion

The methodology used allowed us to extract from a single ITC experiment all thermodynamic properties of solution (ΔsolHm, ΔsolGm, and ΔsolSm) of the studied alcohols in both aqueous media.

The present results show that there is a significant change in the values of thermodynamic properties of solution when we add a high concentration of trehalose to the HEPES aqueous solution. The main effect of the crowded media on the thermodynamic properties of solution for these alcohols is a reduction in solubility that is enthalpically driven, whereas the entropic contribution is less clearly defined.

The values of the partial molar heat capacity here directly obtained by heat capacity drop calorimetry showed a significant reduction in this property in the media with trehalose. This can be taken as a clear sign of a decrease in hydrophobic solvation that we propose to be a result of a decrease in free solvent availability.

Finally, we tentatively predict that in the aqueous media of the crowded solutions that characterize cells and biological fluids, solutes with low aqueous solubility will be more soluble, whereas the solubility of highly polar solutes will be reduced.

Corresponding author: Margarida Bastos, CIQ (UP), Faculty of Sciences, Department of Chemistry and Biochemistry, University of Porto, R. Campo Alegre, 687, P-4169-007 Porto, Portugal, Tel.: +351 22 0402511, Fax: +351 22 0402659, E-mail:

A collection of invited papers based on presentations at the 33rd International Conference on Solution Chemistry (ICSC-33), Kyoto, Japan, 7–12 July 2013.

This work received support from FCT, Lisbon, Portugal, and European Social Fund through strategic projects PEst-C/QUI/UI0081/2013 awarded to CIQUP and project PTDC/QUI/64565/2006 to MJM and ESE, a Post-Doc grant to GB (SFRH/BPD/41407/2007) and a PhD grant to MAAR (SFRH/BD/60513/2009). GB acknowledges the support from the National Natural Science Foundation of China (21273061).

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Published Online: 2014-01-17
Published in Print: 2014-02-01

©2014 IUPAC & De Gruyter Berlin Boston

Articles in the same Issue

  1. Masthead
  2. Masthead
  3. Preface
  4. International Union of Pure and Applied Chemistry
  5. Conference paper
  6. Optimization of superamphiphobic layers based on candle soot
  7. Probing “ambivalent” snug-fit sites in the KcsA potassium channel using three-dimensional reference interaction site model (3D-RISM) theory
  8. Perspectives for hybrid ab initio/molecular mechanical simulations of solutions: from complex chemistry to proton-transfer reactions and interfaces
  9. The complex structure of ionic liquids at an atomistic level: from “red-and-greens” to charge templates
  10. Amide I IR probing of core and shell hydrogen-bond structures in reverse micelles
  11. Kinetic studies on cyclopalladation in palladium(II) complexes containing an indole moiety
  12. Properties of the one-electron oxidized copper(II) salen-type complexes: relationship between electronic structures and reactivities
  13. Solvation of a sponge-like geometry
  14. What is “hypermobile” water?: detected in alkali halide, adenosine phosphate, and F-actin solutions by high-resolution microwave dielectric spectroscopy
  15. Reentrant condensation, liquid–liquid phase separation and crystallization in protein solutions induced by multivalent metal ions
  16. Emulsion-templated macroporous polymer/polymer composites with switchable stiffness
  17. Effective interaction between small unilamellar vesicles as probed by coarse-grained molecular dynamics simulations
  18. Enthalpies of solution, limiting solubilities, and partial molar heat capacities of n-alcohols in water and in trehalose crowded media
  19. Protonation of alkanolamines and cyclic amines in water at temperatures from 293.15 to 373.15 K
  20. IUPAC Technical Report
  21. Defining the transfer coefficient in electrochemistry: An assessment (IUPAC Technical Report)
  22. IUPAC Recommendations
  23. Definition of the transfer coefficient in electrochemistry (IUPAC Recommendations 2014)
https://doi.org/10.1515/pac-2014-5014
Keywords for this article
alcohol; calorimery; crowded media; enthalpy of solution; isothermal titration calorimetry (ITC); solubility

Articles in the same Issue

  1. Masthead
  2. Masthead
  3. Preface
  4. International Union of Pure and Applied Chemistry
  5. Conference paper
  6. Optimization of superamphiphobic layers based on candle soot
  7. Probing “ambivalent” snug-fit sites in the KcsA potassium channel using three-dimensional reference interaction site model (3D-RISM) theory
  8. Perspectives for hybrid ab initio/molecular mechanical simulations of solutions: from complex chemistry to proton-transfer reactions and interfaces
  9. The complex structure of ionic liquids at an atomistic level: from “red-and-greens” to charge templates
  10. Amide I IR probing of core and shell hydrogen-bond structures in reverse micelles
  11. Kinetic studies on cyclopalladation in palladium(II) complexes containing an indole moiety
  12. Properties of the one-electron oxidized copper(II) salen-type complexes: relationship between electronic structures and reactivities
  13. Solvation of a sponge-like geometry
  14. What is “hypermobile” water?: detected in alkali halide, adenosine phosphate, and F-actin solutions by high-resolution microwave dielectric spectroscopy
  15. Reentrant condensation, liquid–liquid phase separation and crystallization in protein solutions induced by multivalent metal ions
  16. Emulsion-templated macroporous polymer/polymer composites with switchable stiffness
  17. Effective interaction between small unilamellar vesicles as probed by coarse-grained molecular dynamics simulations
  18. Enthalpies of solution, limiting solubilities, and partial molar heat capacities of n-alcohols in water and in trehalose crowded media
  19. Protonation of alkanolamines and cyclic amines in water at temperatures from 293.15 to 373.15 K
  20. IUPAC Technical Report
  21. Defining the transfer coefficient in electrochemistry: An assessment (IUPAC Technical Report)
  22. IUPAC Recommendations
  23. Definition of the transfer coefficient in electrochemistry (IUPAC Recommendations 2014)
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