Solubility Measurement and Thermodynamic Study of Pimelic Acid in Water, 1-propanol, and Their Binary Mixtures at 293.15–303.15 k

 

Sandip B. Nahire*

Department of Chemistry, M.S.G. College, Malegaon (M.S.) India.

*Corresponding Author E-mail: nahiresandip@gmail.com

 

 

INTRODUCTION:

Pimelic acid (PA) [1,7-heptanedioic acid, HOOC–(CH₂)₅–COOH; molar mass: 160.17 g·mol⁻¹; CAS No. 111-16-0] is a linear aliphatic dicarboxylic acid that serves as a valuable intermediate in the chemical and pharmaceutical industries. It is widely used in the synthesis of polyamides, plasticizers, and resins, and is a precursor for 1,7-heptanediol—a compound employed in the manufacture of pharmaceuticals, surfactants, flavors, and cosmetic products^1-3. Additionally, PA is utilized in biochemical studies and synthetic applications due to its molecular flexibility and bifunctional carboxylic groups². Accurate solubility data and solid–liquid phase equilibrium (SLE) information is essential for the rational design of crystallization and separation processes in industrial and research settings⁴. In particular, the solubility behavior of active pharmaceutical ingredients and organic acids in mixed solvents is critical for optimizing yield, purity, and process efficiency.

 

However, comprehensive solubility data for pimelic acid in mixed aqueous–organic systems are limited in the literature⁵. In this study, the solubility of PA was experimentally determined in pure water, pure 1-propanol, and their binary mixtures over the entire composition range (0 to 1 weight fraction of 1-propanol) at five temperatures: 293.15, 295.15, 298.15, 300.15, and 303.15 K. Furthermore, thermodynamic parameters—standard enthalpy of solution (ΔH⁰ₛₒₗₙ), standard entropy (ΔS⁰ₛₒₗₙ), standard Gibbs free energy (ΔG⁰ₛₒₗₙ), and the relative contributions of enthalpy (%ζH) and entropy (%ζTS)—were evaluated using the van’t Hoff approach to elucidate the dissolution mechanism and solute–solvent interactions.

 

EXPERIMENTAL:

Materials and Apparatus:

Triple-distilled water was used for all solution preparations. Pimelic acid (PA) [CAS No. 111-16-0, purity: 99.5%] was purchased from Merck, and 1-propanol (purity: 99.9%) was obtained from Jiangyin Huaxi International Trade Co., China. All reagents were of analytical grade and used without further purification. The apparatus and methodology employed for solubility measurements were based on previously reported procedures with minor modifications^6-8. In brief, binary mixtures of water and 1-propanol were prepared gravimetrically using an analytical balance (Shimadzu AUXZZO) with a weighing accuracy of ±0.1 mg. An excess amount of pimelic acid was added to 100mL of the binary solvent mixture in a specially designed double-jacketed glass flask. The flask was connected to a thermostatic water circulator, which maintained the temperature within ±0.1 K by circulating water between the inner and outer jackets. The mixture was stirred continuously using a magnetic stirrer for approximately 1 hour to ensure saturation equilibrium, after which it was allowed to stand undisturbed for another hour to facilitate phase separation. After equilibration, a fixed volume of the supernatant was withdrawn using a pipette preheated to a temperature slightly above the solution temperature to avoid premature precipitation. The withdrawn solution was transferred to a pre-weighed glass weighing bottle, and the solvent was evaporated completely by heating the bottle in an oven maintained at 343 K. Complete evaporation was confirmed by repeated weighing until a constant mass was achieved. The solubility of pimelic acid was determined from the mass of the residue (solute) and the initial mass of the withdrawn solution. All solubility values reported represent the average of at least three independent measurements, ensuring reproducibility and consistency. The saturated mole fraction solubility of pimelic acid (X_b), the initial mole fraction of 1-propanol (), and the initial mole fraction of water () were calculated using the standard expressions given in Eqs. (1) and (2).

 

                m_b⁄M_b

  X_b= ––––––––––––––––––––                                     (1)

    m_a⁄M_a +m_b⁄M_b +m_c⁄M_c

 

                   m_c⁄M_c                                                  m_b⁄M_a

 = ––––––––––––––––     = ––––––––––––––––  (2)

             m_a⁄M_a +m_c⁄M_c                               m_a⁄M_a +m_c⁄M_c

 

Where m_b, m_a, and m_c are the mass of PA, water, 1-propanol respectively, and M_b, M_a, and M_c are the molecular weight of the solute, water, and 1-propanol, respectively.

 

RESULTS AND DISCUSSION:

Verification of the experimental methods:

To ensure the reliability and accuracy of the experimental apparatus and procedures, the solubility of PA in pure water was measured and compared with values reported in the literature^9–11. The experimental results obtained in this study show good agreement with previously published data, confirming that the experimental setup and methodology are both accurate and dependable for solubility determination.

 

Solubility Data:

The experimental mole fraction solubility data (X_b) of PA in pure water, pure 1-propanol, and binary mixtures of water + 1-propanol at 293.15, 296.15, 298.15, 300.15, and 303.15 K are presented in Table 1 and illustrated graphically in Figure 2.

 

 

Fig 1. Comparison of Experimental Solubility of PA in water as a function of temperature: (■) this work, (◊)⁹ (∆)¹⁰ (×)¹¹.

 

At a constant temperature, the solubility of PA increases progressively with increasing 1-propanol content in the binary solvent mixture. The data show that solubility rises continuously with increasing mole fraction of 1-propanol () up to approximately = 0.6100, beyond which the solubility (X_b) approaches a saturation plateau. As observed from Table 2 and Figure 2, in the case of pure 1-propanol, the solubility of PA increases and reaches its maximum within the mole fraction range  = 0.5453 to 0.7296, indicating a solubility maximum before leveling off. The most pronounced solubility enhancement is observed in the mixed solvent system water + 1-propanol, highlighting the synergistic solvation effect of the binary mixture. Furthermore, Figure 3 shows that the solubility of PA in water + 1-propanol mixtures increase linearly with temperature, confirming the temperature dependence of the dissolution process.

 

Table 1 Experimental (X_b) and calculated mole fraction solubility of PA in various initial mole fraction  of 1-propanol at T= 293.15 to 303.15 K

PA+ Water + 1-propanol
T/K
293.15	0.0000	0.0044	0.0043	0.0045
	0.0322	0.0091	0.0086	0.0080
	0.0697	0.0203	0.0208	0.0223
	0.1139	0.0248	0.0258	0.0293
	0.1666	0.0572	0.0565	0.0559
	0.2307	0.0748	0.0740	0.0737
	0.3102	0.0820	0.0819	0.0829
	0.4116	0.1068	0.1061	0.1077
	0.5453	0.1215	0.1211	0.1186
	0.7296	0.1171	0.1162	0.1166
	1.0000	0.1006	0.0996	0.1001
296.15	0.0000	0.0053	0.0053	0.0053
	0.0322	0.0100	0.0108	0.0107
	0.0697	0.0285	0.0281	0.0283
	0.1139	0.0354	0.0366	0.0372
	0.1666	0.0632	0.0642	0.0641
	0.2307	0.0815	0.0829	0.0828
	0.3102	0.0916	0.0927	0.0928
	0.4116	0.1180	0.1181	0.1183
	0.5453	0.1297	0.1292	0.1288
	0.7296	0.1272	0.1272	0.1273
	1.0000	0.1080	0.1093	0.1093
298.15	0.0000	0.0058	0.0059	0.0059
	0.0322	0.0128	0.0127	0.0130
	0.0697	0.0355	0.0338	0.0332
	0.1139	0.0510	0.0449	0.0434
	0.1666	0.0695	0.0699	0.0701
	0.2307	0.0891	0.0894	0.0895
	0.3102	0.1015	0.1003	0.0999
	0.4116	0.1254	0.1263	0.1258
	0.5453	0.1345	0.1353	0.1360
	0.7296	0.1329	0.1349	0.1348
	1.0000	0.1157	0.1160	0.1159
300.15	0.0000	0.0066	0.0067	0.0065
	0.0322	0.0145	0.0150	0.0156
	0.0697	0.0393	0.0401	0.0387
	0.1139	0.0555	0.0541	0.0506
	0.1666	0.0765	0.0762	0.0766
	0.2307	0.0968	0.0963	0.0965
	0.3102	0.1089	0.1082	0.1075
	0.4116	0.1342	0.1348	0.1337
	0.5453	0.1411	0.1419	0.1435
	0.7296	0.1435	0.1430	0.1427
	1.0000	0.1235	0.1230	0.1227
303.15	0.0000	0.0077	0.0078	0.0076
	0.0322	0.0210	0.0196	0.0206
	0.0697	0.0504	0.0509	0.0487
	0.1139	0.0661	0.0690	0.0635
	0.1666	0.0872	0.0867	0.0873
	0.2307	0.1077	0.1076	0.1079
	0.3102	0.1198	0.1207	0.1197
	0.4116	0.1488	0.1477	0.1463
	0.5453	0.1512	0.1532	0.1552
	0.7296	0.1538	0.1555	0.1552
	1.0000	0.1337	0.1339	0.1335

 

 

Fig.2 Mole fraction solubility (X_B) variation with initial mole fraction () of 1-propanol at various temperatures

 

Ideal Solubilities and activity coefficients for PA:

The ideal mole fraction solubility of pimelic acid (X^idl) was calculated using the following thermodynamic equation^12,13:

 

                 (3)

 

where:

·       R is the universal gas constant (8.314 J·mol⁻¹·K⁻¹),

·       T is the absolute temperature (K),

·       T_fus is the melting point of pimelic acid (K),

·       ΔH_fus is the molar enthalpy of fusion (J·mol⁻¹),

 

ΔC_p is the difference in molar heat capacities between the liquid and crystalline states (J·mol⁻¹·K⁻¹).

 

 

Fig.3 Mole fraction solubility (X_b) variation with temperatures

 

The values of T_fus, ΔH_fus, and ΔC_p for PA were taken from the literature¹⁴ and used to calculate the ideal solubility (X^idl) at various temperatures using Eq. (3). The calculated ideal solubility values are listed in Table 2. These ideal solubility values were subsequently used to determine the activity coefficients (γ) of PA in different solvent systems, providing insights into solute–solvent interactions.

 

Table 2: - The  PA in pure water and 1-propanol at T = (293.15 -303.15 K).

 

The activity coefficients (γ) of PA in each solvent system were calculated using the following relation:

 

                                                                          (4)

 

Table 3. Activity coefficient (γ) of PA in water + 1-propanol binary mixture at T= 293.15-303.15 K

The activity coefficients of PA in water, 1-propanol, and their binary mixtures at temperatures ranging from 293.15 to 303.15 K are listed in Table 3.

 

The activity coefficients (γ) of pimelic acid in water + 1-propanol mixtures over the temperature range 293.15 to 303.15 K are presented in Table 3. As observed from the data, γ values in pure water are significantly higher than those in pure 1-propanol at all temperatures. For both solvents, γ values are greater than unity, indicating positive deviations from ideality, which suggests weaker or repulsive solute–solvent interactions in these systems. Moreover, as the mole fraction of 1-propanol ()  increases in the binary mixtures, the γ values systematically decrease. This trend indicates that the introduction of 1-propanol enhances solute–solvent affinity, likely due to better solvation or more favorable interactions between PA and 1-propanol molecules compared to water. The decrease in γ with increasing alcohol content is consistent with the observed increase in solubility, confirming the correlation between solubility enhancement and reduced deviation from ideality.

 

Thermodynamic Analysis:

Thermodynamic functions provide valuable insight into the molecular mechanisms governing the dissolution process of a solute in various solvents. In the present study, thermodynamic parameters associated with the dissolution of PA were calculated based on its solubility data. According to the van’t Hoff equation^15,16, the standard molar enthalpy of solution (ΔH⁰ₛₒₗₙ) can be estimated from the temperature dependence of solubility, specifically from the slope of the linear plot of ln X_b versus 1/T. The analysis was performed over a limited temperature range of 293.15 to 303.15 K, with the mean temperature (T_mean) taken as 298.15 K for calculation purposes. The values of ΔH⁰ₛₒₗₙ were obtained from the slope of the best-fit line for each solvent system.

 

=                        (5)

 

The slope and intercept of the plot of ln X_b versus (1/T − 1/T_mean) for each pure solvent and binary mixture are presented in Table 4.

 

Table 4. Thermodynamic Functions Relative to dissolution Process of glutaric acid at T_mean = 298.15 K

 

The standard molar Gibbs energy change for the solution process  , can be calculated by Eq. 6:       

 

                                        (6)

 

The standard molar entropy change is obtained from Eq.7:

 

                                                 (7)

Both the standard Gibbs free energy of solution (ΔG⁰ₛₒₗₙ) and the standard entropy of solution (ΔS⁰ₛₒₗₙ) were calculated at the mean temperature T_mean = 298.15 K. The corresponding values are summarized in Table 4, along with the relative enthalpic and entropic contributions, expressed as %ζH and %ζTS, respectively. These parameters provide a comparative evaluation of the roles of enthalpy and entropy in the dissolution process and were calculated using Eq. (8).

 

% and

 

                                    (8)

 

The positive values of the standard enthalpy of solution (ΔH⁰ₛₒₗₙ) indicate that the dissolution of pimelic acid (PA) in all solvent systems is an endothermic process. The order of ΔH⁰ₛₒₗₙ values in pure solvents is: 1-propanol (21.56 kJ·mol⁻¹) < water (41.49 kJ·mol⁻¹). This suggests that less energy is required for the dissolution of PA in 1-propanol compared to water, implying stronger solute–solvent interactions in the alcoholic medium. Additionally, both ΔG⁰ₛₒₗₙ and ΔS⁰ₛₒₗₙ values are positive across all binary solvent mixtures. The positive ΔG⁰ₛₒₗₙ values indicate that the dissolution process is non-spontaneous under standard conditions, while the positive entropy change (ΔS⁰ₛₒₗₙ) reflects an increase in molecular disorder upon dissolution. As seen from Table 4, the percentage contribution of enthalpy (%ζ_H) is greater than that of entropy (%ζ_TS) in all solvent systems, suggesting that the enthalpic contribution dominates the free energy of solution (ΔG⁰ₛₒₗₙ).

 

CONCLUSIONS:

In this study, the solubility of PA in water, 1-propanol, and their binary mixtures was experimentally determined over the temperature range of 293.15 to 303.15 K using a gravimetric method. The solubility of PA was found to increase with temperature in all solvent systems, exhibiting a nearly linear trend. Among the solvents tested, 1-propanol showed higher solubility for PA than water, and the maximum solubility enhancement was observed in the water+1-propanol binary mixtures. Ideal solubility analysis further confirmed that 1-propanol behaves as a more effective solvent than water for dissolving PA. Thermodynamic parameters, evaluated using the van’t Hoff approach, indicate that the dissolution process is endothermic, enthalpy-driven, and non-spontaneous under standard conditions. These findings contribute to a better understanding of solute–solvent interactions in mixed solvent systems and may aid in the design of solvent media for improved solubility of structurally similar compounds.

 

ACKNOWLEDGEMENTS:

The authors are thankful to Principal of MSG Arts, Science and Commerce College Malegaon for providing laboratory facilities. The authors also express their sincere thanks to Dr Apoorva Hiray (Co-ordinator M.G. Vidyamandir Malegaon).

 

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Received on 28.07.2025      Revised on 13.08.2025 Accepted on 30.08.2025      Published on 30.09.2025 Available online from October 07, 2025 Asian J. Research Chem.2025; 18(5):319-323. DOI: 10.52711/0974-4150.2025.00048 ©A and V Publications All Right Reserved
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