Some new cyano pyridine derivatives have been synthesized and their characterization was done by IR, 1H NMR and mass spectral data. The solubility of these synthesized compounds has been studied in dimethyl formamide and dimethyl sulphoxide at different temperatures at atmospheric pressure.

Keywords: cyano pyridine derivative, solubility, DMF, DMSO, thermodynamic parameter

DMF, dimethyl formamide; DMSO, dimethyl sulfoxide; RMSD, root-mean-square deviations; RD, relative deviations; RAD, relative average deviations

Pyridine compounds exist in nature in various forms and are integral part of various natural products.^1,2 The pyridine ring plays a key role in catalyzing both biological and chemical reactions.³ Various substituted pyridines demonstrate a wide range of applications. Among various substituted pyridines, cyano pyridine derivatives have been found to be an important sub class. Various substituted cyano pyridine derivatives are known to act as intermediates in the pharmaceutical, dye, photo and agrochemical industries.^4‒6 Further, various cyano pyridines have drawn attention due to their wide spectrum biological activities.^7‒11 Therefore, due to their applications in biological and chemical fields, it would be interesting to determine the solubility of some novel cyano pyridine derivatives in different solvents at various temperatures. The data may be useful for their application in other fields also.

Thus, in the present work, solubility of some newly synthesized cyano pyridine derivatives is determined in dimethyl formamide and dimethyl sulfoxide at different temperatures. Further, some thermodynamic parameters such as enthalpy, Gibb’s free energy and entropy of dissolution for these synthesized compounds have also been evaluated.

The solvents dimethyl formamide (DMF) and dimethyl sulfoxide (DMSO) were used for the present study were purified by standard methods.¹² All the synthesized compounds were crystallized and Figure 1 shows the general structure of these derivatives.

Figure 1 Asymmetric oxidation of the thioethers.

Solubility

The gravimetric method was used to study the solubility. An excess mass of compound was added to a known mass of solvent. The solution was heated to a constant temperature with continuous stirring. After, at least 3hrs the stirring was stopped and the solution was kept at a constant temperature for 2hrs. A portion of this solution was filtered and by a preheated injector, 5ml of this clear solution was taken to pre weighted measuring vial (m₀). The vial was quickly and tightly closed and weighted (m₁) to determine the mass of the sample (m₁- m₀). To prevent dust contamination, the vial was covered with a piece of filter paper. After completely dryness of vial mass, the vial was reweighed (m₂) to determine the mass of the constant residue solid (m₂- m₀). All the weights taken using Mettler Toledo AB204-S, Switzerland electronic balance with uncertainty of ±0.0001g. Thus the concentration of solid sample in the solution, mole fraction x, could be determined from equation

$x=\frac{(m2-m0)/M1}{(m2-m0)/M1+(m1-m2)/M2}$ …………….. (1)

Where M1 and M2 is the molar mass of solvent and compound respectively. At each temperature, the measurement was repeated three times and an average value is taken.

The molecular formula, molecular weight, melting point, % yield and Rf values along with the solvent systems of all the compounds are given in Table 1.

Figure 2 The variation of experimental mole fraction solubility (x) of compounds with temperature in DMF.

(♦); CP-1, (■); CP-2, (▲); CP-4, (♦); CP-4, (●); CP-5, (♦); CP-6, (■); CP-7, (▲); CP-8, (●); CP-9, (●); CP-10.

Figure 3 The variation of experimental mole fraction solubility (x) of compounds with temperature in DMSO.
(♦); CP-1, (■); CP-2, (▲); CP-4, (♦); CP-4, (●); CP-5, (♦); CP-6, (■); CP-7, (▲); CP-8, (●); CP-9, (●); CP-10.

Table 2 and Table 3 show the experimental solubility values of compounds at different temperatures in DMF and DMSO respectively. The variation of mole fraction solubility of compounds with temperature in DMF and DMSO is shown in Figure 2 and Figure 3 respectively. It is observed that in both the solvents, solubility increases with temperature. Further, comparison of solubility in both the solvents; DMF and DMSO shows that overall solubility is greater in DMSO than that in DMF. Thus, the solvent polarity plays an important role on the solubility of studied compounds. The dielectric constant of DMSO (46.6) is greater than that of DMF (36.71). However, there is very small variation is in their dipole moments (3.9 for DMSO and 3.86 for DMF). This suggests that dielectric constant of solvent plays an important role in dissolution for the studied compounds.
The temperature dependence of solubility was described by the modified Apelblat equation13,14

$\ln x=A+\frac{B}{T}+ C\ln T$ ….. (2)

Where T is the absolute temperature, and A, B, and C are empirical constants. The values of these parameters are listed in Table 4. The root-mean-square deviations (RMSD) are calculated using the following equation:

$\text{RMSD}=\sqrt{{\displaystyle \sum }_{i=1}^N\frac{{\left(x_i-x\right)}^{2 }}{N-1}}$ …….. (3)

Where N is the number of experimental points and x and xi represent the mole fraction solubility of the experiment and that calculated from eq 2, respectively. These values are given in Table 4. Further, relative deviations (RD) and relative average deviations (RAD) are calculated by eq (4) and (5).

$RD=\left(\frac{x-x_i}{x}\right)$

$\text{RAD}=\frac{1}{N}{\displaystyle \sum }_i^N\frac{(x-x_i)}{x}$ …… (5)

Where N is the number of experimental points and xi is the solubility calculated by eq 2. The values of relative deviation are listed in Table 2 and Table 3 for DMF and DMSO respectively and relative average deviation values are reported in Table 4. It is evident from Table 2 and Table 3 that relative deviation (RD) values are not more than 1.85% for DMF and 2.86% for DMSO. Thus, there is good agreement between experimental and calculated solubility values in both the solvents. Using experimental data of solubility in different solvents, some thermodynamic parameters such as dissolution enthalpy, Gibb’s energy of dissolution and entropy have also been evaluated. According to modified Van’t Hoff equation15,16 the dissolution enthalpy  were evaluated by following relation.

$\frac{\partial lnx}{\partial \left(\frac{1}{T}-\frac{1}{T_{hm}}\right)}=-\frac{\Delta H_{sol}}{R}$ ….. (6)

Where T is the experimental temperature and R is universal gas constant. T_hm represent the mean harmonic temperature which is given as

$T_{hm}=\frac{n}{{{\displaystyle \sum }}_i^n\left(\frac{1}{T}\right)}$ ….. (7

Where n is the number of experimental temperatures.¹⁷ In present case, the value of T_hm is obtained only 308K. The slope of the plot of ln x versus (1/T-1/308) gives the value of ∆H_sol. From the intercepts of these plots, Gibbs energy change (ΔG_sol) for dissolution process were calculated from the following relation¹⁵

$\Delta G_{sol}=-R{T}_{hm}.intercept$ ….. (8)

Using these evaluated ∆H_sol and ∆G_sol values, the entropies of solutions ∆S_sol were obtained from the following equation:

$\Delta S_{sol}=\frac{\Delta H_{sol}-\Delta G_{sol}}{T_{hm}}$ ….. (9)

All these thermodynamic parameters are listed in Table 5.

It is evident from Table 5, that for all the compounds, the evaluated thermodynamic parameters i.e., ∆H_sol, ∆G_sol and ∆S_sol values are positive for both the solvents. The positive ∆H_sol suggests endothermic dissolution of compounds in both the solvents. The endothermic effect may be due to strong interactions between compound and solvent molecules.^17,18 Whereas, positive ∆G_sol values indicate that the dissolution process is spontaneous. The positive entropy indicates that dissolution process increases the randomness in solution.¹⁸ However, for some compounds entropy is less than half value than those of other compounds. This depends on the functional groups present in the compound as well as on the solvent. Different functional groups interact differently with the solvent, so randomness will be different.

It is concluded that solubility increases with temperature in both the solvents. Overall, solubility is greater in DMSO than that in DMF for all the compounds. Further, the evaluated thermodynamic parameters i.e., enthalpy, Gibb’s free energy and entropy of dissolutions values are positive for both the solvents. The positive enthalpy suggests endothermic dissolution of compounds in both the solvents indicating thereby strong interactions between compound and solvent molecules. The positive Gibb’s free energy and entropy indicate that dissolution process is spontaneous and it increases the randomness in solution.

Authors are thankful to Head of Chemistry Department, Saurashtra University, Rajkot, India for providing necessary facilities.

There is no conflict of interest.

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