determination of dissociation constants and thermodynamic parameters of pyrimidine derivatives in...

Journal of Molecular Liquids xxx (2014) xxx–xxx

MOLLIQ-04377; No of Pages 7

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Journal of Molecular Liquids

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Determination of dissociation constants and thermodynamic parametersof pyrimidine derivatives in organic-water mixed solvents atdifferent temperatures

Kapil Bhesaniya, Shipra Baluja ⁎Physical Chemistry Laboratory, Department of Chemistry, Saurashtra University, Rajkot 360005, Gujarat, India

⁎ Corresponding author.E-mail addresses: [email protected] (K. Bh

[email protected] (S. Baluja).

http://dx.doi.org/10.1016/j.molliq.2014.07.0420167-7322/© 2014 Published by Elsevier B.V.

Please cite this article as: K. Bhesaniya, S.derivatives in organic-water mixed solvents

a b s t r a c t
a r t i c l e i n f o
Article history:Received 30 April 2014Received in revised form 27 July 2014Accepted 28 July 2014Available online xxxx

Keywords:Pyrimidine derivativeDissociation constantAverage methodThermodynamic parameterGibb's free energyEntropy change

Twonovel pyrimidine derivatives have been synthesized and their structures havebeen confirmedby IR, 1HNMRandmass spectral data. The dissociation constants of these two compounds; 4-(naphthalen-1-yloxy)-N-(p-tolyl)pyrimidin-2-amine (TP-1) and N-(4-fluoro phenyl)-4-(naphthalen-1-yloxy) pyrimidin-2-amine (FP-1) werestudied in methanol/DMF - water (60:40 v/v) solvent systems at different temperatures ranging from 25 °C to45 °C at a 10 °C interval using the Calvin–Bjerrum pH titration method. The results are interpreted in terms ofsubstituent present in the compounds. Furthermore, some thermodynamic parameters such as enthalpy(ΔH°), Gibb's free energy (ΔG°) and entropy (ΔS°) of these dissociations have also been evaluated at differenttemperatures from dissociation constant data.

© 2014 Published by Elsevier B.V.

1. Introduction

Dissociation constant is the core physicochemical parameter of anyelectrolyte, which helps in understanding various processes of industri-al applications. In biological terms, the dissociation constant i.e., pKvalue gives an idea about the presence of a compound in the polar ornon-polar phase (partition). From a computational chemistry point ofview, pK calculations are a benchmark for quantum mechanical andfree solvation energy calculations. Further, the accurate determinationof pK value is important for the analysis of drugs to understand theirdistribution, transport behavior, binding to receptors and mechanismof action in various chemical and biochemical processes [1,2].

Heterocyclic compounds play an important role in industry as wellas in our life. Various heterocyclic compounds are known to act as ther-apeutic agents. Among these compounds, pyrimidines are one of themost important heterocycles exhibiting remarkable pharmacologicalactivities. It is an essential constituent of all cells of living matter as apyrimidine ring is a building unit of DNA and RNA [3]. These pyrimidinecompounds exhibit a range of biological activities such as anti-microbial[4], analgesic [5,6], anti-inflammatory [7,8], anti-tubercular [9], anti-tumor [10], anti-cancer [11], etc.

esaniya),

Baluja, Determination of disat ..., Journal of Molecular Liq

The applications of these compounds attracted us to determine theirdissociation constants and thermodynamic parameters in differentsolvents.

A literature survey shows that very little work has been done forthe determination of the dissociation constant of synthetic organiccompounds [12,13]. In continuation of our research work [14,15], thepresent work deals with the synthesis, characterization and determina-tion of the dissociation constant of synthesized pyrimidine derivativesin a methanol/DMF–water system by the Calvin–Bjerrum pH titrationmethod at different temperatures (298.15 to 318.15 K).

2. Experimental

2.1. Materials

2,4-Dichloropyrimidine (DCP) (CASNo.: 3934-20-1) and 1-naphthol(NTL) (CAS No.: 90-15-3) used in the synthesis were purchased fromSigma-Aldrich. Potassium carbonate (K2CO3) (CAS No.: 584-08-7) waspurchased from Sisco Chem. Pvt. Ltd. (Mumbai, India). Sodium nitrate(CAS No.: 7631-99-4), nitric acid (CAS No.: 7697-37-2) and sodium hy-droxide (CAS No.: 1310-73-2) were purchased from SD Fine Chem. Ltd.(Vadodara, India). The solvents methanol and dimethylformamide(DMF) used in the present work were of AR grade supplied bySpectrochem Pvt. Ltd. (Mumbai, India) and were purified according tothe standard procedure [16] and were kept over molecular sieves. The

sociation constants and thermodynamic parameters of pyrimidineuids (2014), http://dx.doi.org/10.1016/j.molliq.2014.07.042

http://dx.doi.org/10.1016/j.molliq.2014.07.042

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2 K. Bhesaniya, S. Baluja / Journal of Molecular Liquids xxx (2014) xxx–xxx

purity of solvents was checked by GC–MS (SHIMADZU Model No. QP-2010) and found to be greater than 99.7%.

2.2. Spectroscopy study

Spectroscopic study of both the synthesized compounds was doneby IR, 1H NMR and mass spectroscopy. IR spectra were recorded onKBr discs, using FT-IR (Shimadzu spectrophotometer Model no. 8400).1H-NMR spectra were taken on a Bruker Avance II 400. In all the cases,NMR spectra were obtained in DMSO-d6 using tetramethylsilane(TMS) as an internal standard. The NMR signals are reported in δppm.Mass spectra were determined using a direct inlet probe on a GCMS-QP-2010 mass spectrometer. Melting points of compounds were deter-mined by a differential scanning calorimeter (Shimadzu-DSC-60).

2.3. Synthesis of 4-(naphthalen-1-yloxy)-N-(p-tolyl) pyrimidin-2-amine(TP-1) and N-(4-fluorophenyl)-4-(naphthalen-1-yloxy)pyrimidin-2-amine (FP-1)

A mixture of 2,4-dichloropyrimidine (DCP) (0.1 mmol), 1-naphthol(NTL) (0.1 mmol) and potassium carbonate (K2CO3) (0.15 mmol) inDMF was refluxed for 4 h. The completion of the reaction was con-firmed by analytical thin layer chromatography (TLC) using hexane:ethyl acetate (7:3) as mobile phase. After the completion of the reac-tion, the reactionmass was cooled and the resulting solid was filtered,washed with cold water and dried under vacuum to get a crudeproduct.

This resulting product (0.1 mmol) was refluxed for 4–5 h withethanolic solution of appropriate aromatic amines (0.12 mmol) usinghydrochloric acid as catalyst. The completion of the reaction was con-firmed by TLC using hexane:ethyl acetate (7.5:2.5) mobile phase. Afterthe completion of the reaction, the reaction mass was cooled and theresulting solid was filtered, washed with cold ethanol and dried undervacuum. The obtained crude product was purified by adding a suitablesolvent (diethyl ether) to remove colored, nonpolar impurity byscratching/stirring. The product was then allowed to stabilize and theabove solution was decanted. The procedure was repeated 3–4 timesto free the product from impurities (trituration) and the purity of allthese synthesized compounds was ascertained by TLC (performed onaluminum coated plates Gel 60 F254 (E. Merck)). The reaction schemeis given in Fig. 1.

The physical parameters of these compounds are summarizedbelow:

TP-1: yield-71%; m.p. 294.6 °C; molecular formula—C21H17N3O.IR (cm−1, KBr): 3267.52 (\NH (sec.) str.), 3188.44 (aryl ether C\Hstr.), 3057.27 (Ar\H str.), 2872 (Ar\CH3 str.) 1585.54-1454.38(C_C str. phenyl nucleus), 1421–1327 (C\H in plane bending),1255–1217 (diarylethers str.), 1153.47-1076 (C\O\C sym. str.),

Fig. 1. Synthesis scheme of

Please cite this article as: K. Bhesaniya, S. Baluja, Determination of disderivatives in organic-water mixed solvents at ..., Journal of Molecular Liq

981 (pyrimidine ring breathing), 808 (C\H in plane bending),628.81-723.33 (C\C out of plane bending).1H NMR (DMSO-d6) δ(ppm): 2.107 (s, 3H), 6.628-6.744 (d, 1H, J =5.6), 6.941-6.960 (d, 2H, J = 7.6), 7.284 (s, 2H), 7.475-7.493 (d, 1H,J = 7.2), 7.527-7.642 (m, 3H),7.808-7.829 (m, 1H) 7.968-7.988 (d,1H, J = 8), 8.062-8.082 (d, 1H, J = 8), 8.429-8.443 (d, 1H, J = 5.6),9.71 (s, 1H); MS: (m/z) = 327.38.FP-1: yield-75%; m.p. 292.18 °C; molecular formula—C20H14FN3O.IR (cm−1, KBr): 3262 (\NH (sec.) str.), 3084.28, 3061.13 (Ar-H str.),1656.91-1664.02 (C_C str. phenyl nucleus), 1398.44 (C\H in planebending), 1244.13-1103.32 (diarylethers str.), 1035 (C\F str.), 831(C\H in plane bending), 794.7-619.17 (C\C out of plane bending).1H NMR (DMSO-d6) δ(ppm): 6.610-6.624 (d, 1H, J = 5.6), 6.762(s, 2H), 7.616-7.643 (m, 5H), 7.946-7.967 (d, 1H, J = 8.4), 8.069-8.089 (d, 1H, J = 8), 8.168-8.183 (d, 1H, J = 6), 8.400-8.414 (d, 1H,J = 5.6), 9.651 (s, 1H); MS: (m/z) = 331.11.

2.4. Dissociation constant measurement

0.1 M solutions of studied compounds were prepared in methanoland DMF. These solutions were retained at the desired temperature.The stock solutions of desired concentrations of nitric acid (HNO3),sodium hydroxide (NaOH) and sodium nitrate (NaNO3) required for ti-trations were prepared in Milli-Q water (Millipore Pvt. Ltd. Bangalore,India). An electrical balance (Mettler Toledo AB204-S) with an accuracyof ±0.1 mg was used for solution preparation.

The Calvin–Bjerrum pH titration method [17,18] was used to findout dissociation constants of both compounds. For this, two sets ofsolutions were prepared.

(i) 2.0ml HNO3 (0.1M)+ 4.0mlwater + 30.0 ml (methanol/DMF)+ 4.0 ml NaNO3 (1.0 M),

(ii) 2.0ml HNO3 (0.1M)+ 4.0mlwater + 28.0 ml (methanol/DMF)+ 2.0 ml compound solution (0.1 M) + 4.0 ml NaNO3 (1.0 M).

For different temperatures, both sets of solutions were titratedagainst 0.25 M NaOH and the corresponding pH was recorded by aSystronic pH meter (Model No. EQ-664). The accuracy of the pH meterwas±0.01 pH unit. The Systronic glass electrode and saturated calomelelectrode were used as indicator and reference electrodes respectively.Beforemeasurement, thepHmeterwas calibratedwith a buffer solutionof the known pH 4.0 (0.05Mpotassiumhydrogen phthalate buffer) and9.18 (0.01 M sodium borate decahydrate buffer) for aqueous media.However, in the present study, methanol–water (60:40 v/v) andDMF–water (60:40 v/v) solvent systems are used, so the followingVan Uitert and Haas relation [19] was used for pH correction.

− log Hþh i¼ pHþ log f þ logU0

H ð1Þ

pyrimidine derivatives.



3K. Bhesaniya, S. Baluja / Journal of Molecular Liquids xxx (2014) xxx–xxx

where f is the activity coefficient of the hydrogen ions in the solventmixtures under consideration at the same temperature and ionicstrength and UH

0 is a correction factor at zero ionic strength, whichdepends only on the solvent composition and temperature.

The experiment was repeated at different temperatures. Theconstant temperature was adjusted to ±0.05 K by circulating thethermostated water through the outer jacket of the vessel (NOVA NV-8550 E).

3. Results and discussion

3.1. IR spectra

The IR spectra of TP-1 and FP-1 are depicted in Figs. 2 and3 respective-ly. The IR data of TP-1 clearly shows a\NH (secondary) stretching bandat 3267.62 cm−1. The IR stretching bands at 3180.44–3057.27 cm−1

and 1585–1454.38 cm−1 confirm the presence of aromatic C\H andC_C in the basic skeleton. The absorption bands at 2872.10 cm−1 and1250.70–1076.32 cm−1 indicate the presence of a CH3 groupand ether linkage respectively in synthesized compounds. The bendingvibration of pyrimidine ring C\H was observed at 625.81 cm−1.

For FP-1, the IR stretching bands at 3304–3230 cm−1 show the\NH(secondary). The absorbance peaks at 3084–3061 and 1656.51–1464.02 cm−1 confirmed the presence of aromatic C\H and C_C inthe basic skeleton. The absorption band at 1153.47–1076.32 indicatesthe presence of C\F stretching in the compound. IR stretching band at808 cm−1 indicates the presence of a 1,4 disubstituted phenyl ring inthe synthesized compound.

All the IR group frequencies suggest that both pyrimidine derivativeswere prepared successfully.

3.2. 1H-NMR spectra

The proton NMR spectra of TP-1 and FP-1 are given in Figs. 4 and 5respectively. The peak due to residual DMSO was observed at about2.51 δppm. It can be seen from the chemical structure of compounds

Fig. 2. IR spectrum of


TP-1 and FP-1 that protons of the phenyl ring attached to the carbons,appeared as an appropriate multiplicity in the aromatic region at6.614–8.414 δppm. The protons, which were present in amine linkage(\NH) (attached to phenyl ring and pyrimidine ring), appeared as asinglet at 9.649 and 9.651 δppm in TP-1 and FP-1 respectively. Theother characteristic \CH3 group appeared as a singlet at 2.107 δppmin TP-1. All the 1H NMR splitting of peaks suggests that both pyrimidinederivatives were prepared successfully.

3.3. Dissociation constant and thermodynamic study

The dissociation constants were evaluated at different temperaturesby using Bjerrum's average and half integral methods. Figs. 6 and 7show typical titration curves of the acid in the absence and presenceof compounds TP-1 and FP-1 respectively in methanol + water [A]and in DMF+ water [B] systems respectively at 298.15 K. It is observedfrom these curves that for the same volume of NaOH added, the com-pound titration curves show a lower pH value than the titration curveof free acid. From these titration curves, the average number of protonsassociated with the compound (nH) was calculated by using the Irvingand Rossotti equation [20].

nH ¼ Y− V″−V ′ð Þ N0 þ E0� �n o.

V0 þ V′� �

T0L

n oð2Þ

where V′ and V″ are the volume of alkali required at the same pH forboth acid and compound titration curves respectively. V0 is the initialvolume of the test solution. N0, E0 and TL

0 are the initial concentrationsof the alkali, acid and compound respectively. Y is number of replaceableprotons. For both the compounds, value of nH is found between 0 and 1,which suggests that there is only one replaceable proton present in thestudied compounds.

For the average method, the pK value for each pHwas calculated bythe following equation:

log pK ¼ pHþ log nH= nH−1Þ�:ð½ ð3Þ

TP-1 compound.



Fig. 3. IR spectrum of FP-1 compound.


For the half-integral method, i.e., at nH ¼ 0:5 value, the pK valuewas evaluated from the plot of nH versus pH. These evaluated valuesby both methods are given in Table 1. It is observed that pK values

Fig. 4. 1H NMR spectrum


decrease with increasing temperature i.e., the dissociation of the com-pounds increases with temperature. Similar results were observed byMubarak et al. while studying the dissociation constant of metal

of TP-1 compound.



Fig. 5. 1H NMR spectrum of FP-1 compound.

0.0

4.0

8.0

12.0

0 5 10

pH

Volume of NaOH (ml )

[A]

III

0

3

6

9

12

0 5 10

pH


[B]

III

Fig. 6. Plot of pHof acid (I) and acid+ compound (II) against volume ofNaOH (V) for TP-1[A] in methanol + water and [B] in DMF + water at 298.15 K.



complexes [21]. Further, there is good agreement between pK valuesevaluated by the two methods.

It is evident from Table 1 that the pK value is higher in TP-1 than thatfor FP-1 in both the solvent systems. It means that maximum dissocia-tion takes place in FP-1 and minimum in TP-1. This may be due to thepresence of different substituents in these compounds [22]. Both thecompounds have the same central nucleus but TP-1 contains a methylgroup at the para position whereas in FP-1, at the para position a fluorogroup is present. Overall, the dominating effect of different groupsis: positive resonating effect (+R) N positive hyper conjugation effect(+H) N negative inductive (−I). TP-1 is more basic because of positivehyper conjugation effect of \CH3. The negative inductive effect of thefluoro group decreases the basic character of FP-1. Similar results havebeen reported by earlier studies on pyrimidine compounds where thepolarity of a substituent affects the basicity of the compounds [23].

Further, pK is higher in the methanol–water system as compared tothe DMF–water system, indicating thereby that the solvent also plays akey role in the dissociation of the compounds. The dissociation processof the subjected compounds ismore pronounced in themedia containingpoorer hydrogen bond donating solvent i.e., DMF as compared to meth-anol (Table 1). DMF has a high basic character which is reflected in itshigher tendency to accept hydrogen from the un-ionized solute [24].

Using these values of dissociation constant of compounds, somethermodynamic parameters such as enthalpy change (ΔH°), Gibb's en-ergy change (ΔG°) and entropy change (ΔS°) have also been evaluated.

The enthalpy change (ΔH°) of the dissociation process have beenevaluated by the Van't Hoff relation [25]

d lnKdT

¼ ΔHB

RT2 : ð4Þ

The Gibb's free energy (ΔG°) was calculated by following Eq. (5).

ΔGB ¼ 2:303RTpK: ð5Þ



0

4

8

12

0 5 10

pH


[A]

I

II

0

4

8

12

0 5 10

pH

Volume of NaOH (ml)

[B]

III

Fig. 7. Plot of pH of acid (I) and acid+ compound (II) against volume of NaOH (V) for FP-1[A] in methanol + water and [B] in DMF + water at 298.15 K.


Using these evaluated ΔH° and ΔG° values, entropy change (ΔS°)was determined by Eq. (6):

ΔSB¼ ΔHB−ΔGB� �

T: ð6Þ

Table 1The pK values and some thermodynamic parameter for TP-1 and FP-1 compounds in methano

Compd. T/K pKa ΔG°/(kJ/mol)b

ae be ae be

MethanolTP-1 298.15 9.50 ± 0.03 9.50 ± 0.04 54.23 ± 0.08 54.23 ±

308.15 9.49 ± 0.05 9.47 ± 0.07 55.99 ± 0.01 55.87 ±318.15 9.29 ± 0.09 9.29 ± 0.09 56.59 ± 0.06 56.59 ±

FP-1 298.15 9.39 ± 0.04 9.41 ± 0.02 53.60 ± 0.04 53.72 ±308.15 9.38 ± 0.07 9.33 ± 0.10 55.34 ± 0.09 55.05 ±318.15 9.21 ± 0.01 9.21 ± 0.04 56.10 ± 0.02 56.10 ±

DMFTP-1 298.15 9.41 ± 0.01 9.40 ± 0.06 53.72 ± 0.09 53.66 ±

308.15 9.38 ± 0.03 9.35 ± 0.08 55.34 ± 0.04 55.17 ±318.15 8.97 ± 0.04 8.97 ± 0.07 54.64 ± 0.05 54.64 ±

FP-1 298.15 9.37 ± 0.08 9.34 ± 0.09 53.49 ± 0.09 53.32 ±308.15 9.36 ± 0.09 9.32 ± 0.07 55.23 ± 0.07 54.99 ±318.15 8.95 ± 0.04 8.95 ± 0.01 54.52 ± 0.01 54.52 ±

The ± uncertainties are standard deviation of the fit defined as s ¼ffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiΣ X−Mð Þ2=n

n orwhere X

a Dissociation constant, calculated by Eq. (3).b Gibb's energy change (ΔG°), obtained by Eq. (5).c Enthalpy change (ΔH°), obtained by Eq. (4).d Entropy change (ΔS°), obtained by Eq. (6).e Methods: (a) average, (b) half-integral.


All these thermodynamic parameters are given in Table 1 for bothaverage and half-integral methods. It is evident from Table 1 that ΔH°values are positive for both the compounds, suggesting thereby that dis-sociation is accompanied by the absorption of heat i.e., the dissociationprocess is endothermic. Comparison of ΔH° values in both solventsshows that the dissociation is more endothermic in methanol than inDMF.

The Gibb's free energy is positive whereas entropy change ΔS° isnegative for both the compounds in both the solvents. The positivevalue of ΔG° indicates that the studied compounds prefer the non-dissociate state in solution rather than the dissociate one. However,the entropy change is more negative in methanol which may be dueto the formation of hydrogen bonds between solvent and compoundsi.e., solvation of compounds takes place due to ion–dipole interactions.This could generate more organized structures as salvation spheres.

4. Conclusion

Thus, it is concluded that the dissociation is maximum in FP-1 andminimum in TP-1. The more basic nature of TP-1 having a 4-CH3

group is due to a hyper conjugation effect whereas due to the negativeinductive effect of a 4-F group in FP-1 the basicity is decreased. Thus, adifferent substitution affects dissociation. Further, the dissociation ofthe studied compounds is higher in DMF and increases with tempera-ture. The dissociation is found to be endothermic.

NomenclaturenH Average number of proton associated with the compoundY Number of displaceable protons per moleculeV′ Volume of alkali required for acid titration curveV″ Volume of alkali required for compound titration curveV0 Initial volume of the test solutionN0 Initial concentration of the alkaliE0 Initial concentration of the acidTL0 Initial concentration of the compoundspK Dissociation constant+R Positive resonating effect+H Positive hyper conjugation effect−I Negative inductive effectΔH° Enthalpy change (kJ·mol−1)ΔG° Gibb's energy change (kJ·mol−1)

l/DMF–water system at different temperatures.

ΔH°/(kJ/mol)c −ΔS°/(J/mol)d

ae be ae be

0.06 18.88 ± 0.13 18.92 ± 0.16 118.58 ± 0.13 118.45 ± 0.150.05 120.45 ± 0.16 119.94 ± 0.170.00 118.54 ± 0.18 118.42 ± 0.150.05 16.18 ± 0.13 18.12 ± 0.17 125.51 ± 0.15 119.42 ± 0.190.04 127.08 ± 0.17 119.86 ± 0.200.01 125.48 ± 0.19 119.41 ± 0.24

0.05 39.56 ± 0.14 38.70 ± 0.10 47.50 ± 0.18 50.19 ± 0.170.03 51.23 ± 0.19 53.45 ± 0.140.02 47.41 ± 0.14 50.12 ± 0.190.01 37.72 ± 0.11 35.06 ± 0.10 52.90 ± 0.15 61.25 ± 0.170.05 56.81 ± 0.19 64.68 ± 0.240.09 52.81 ± 0.14 61.17 ± 0.21

is the individual score, M is the mean and n is 3 consecutive measurements.




ΔS° Entropy change (J·mol−1)R Universal gas constant (J·mol−1·K−1)T Temperature in Kelvinf Activity coefficient of the hydrogen ionsUH0 Correction factor at zero ionic strength

AbbreviationsTP-1 and FP-1 Compound codeDCP 2,4-DichloropyrimidineNTL NTL = 1-NaphtholDMF N,N-dimethylformamideAR Analytical reagent

Acknowledgement

The authors are thankful to Prof. P. H. Parsania, Head, Department ofChemistry, Saurashtra University, Rajkot for providing facilities.

References

[1] G.H. Rochester, Acidity Functions, Academic Press, London, 1970.[2] P.A. Frey, F.C. Kokesh, F.H. Westheimer, J. Am. Chem. Soc. 93 (1971) 7266–7269.


[3] T. Sasada, F. Kobayashi, N. Sakai, T. Konakahara, Org. Lett. 11 (2009) 2161–2164.[4] Y. Kumar, V. Gupta, S. Singh, J. Pharm. Res. 7 (2013) 491–495.[5] J.K. Gupta, P.K. Sharma, R. Dudhe, S.C. Mondal, A. Chaudhary, P.K. Verma, Acta Pol.

Pharm. 68 (2011) 785–793.[6] O. Bruno, C. Brullo, S. Schenone, A. Ranise, F. Bondavalli, E. Barocelli, M. Tognolini, F.

Magnanini, V. Ballabeniet, Farmaco 57 (2002) 753–758.[7] S.M. Sondhi, N. Singh, M. Johar, A. Kumar, Bioorg.Med. Chem. 13 (2005) 6158–6166.[8] S.M. Sondhi, M. Dinodia, R. Rani, R. Shukla, R. Raghubir, Indian J. Chem. 49B (2009)

273–281.[9] T.S. Chitre, M.K. Kathiravan, A.S. Chothe, V.K. Rakholiya, K.D. Asgaonkar, S.M. Patil, K.

G. Bothara, J. Pharm. Res. 4 (2011) 1882–1883.[10] N.R. Mohamed, M.M.T. El-Saidi, Y.M. Ali, M.H. Elnagdi, Bioorg. Med. Chem. 15 (2007)

6227–6235.[11] S. Sridhar, Y.R. Prasad, S.C. Dinda, Int. J. Pharm. Sci. Res. 2 (2011) 2562–2565.[12] A. Chylewska, D. Jacewicz, D. Zarzeczanska, L. Chmurzynski, J. Chem. Thermodyn. 40

(2008) 1290–1294.[13] M. Meloun, Z. Ferencikova, A. Vrana, J. Chem. Thermodyn. 43 (2011) 930–937.[14] K.D. Bhesaniya, S. Baluja, J. Mol. Liq. 190 (2014) 190–195.[15] S. Baluja, R. Bhalodia, P. Kasundra, Russ. J. Phys. Chem. A 84 (2010) 2268–2269.[16] J.A. Riddick, W.B. Bunger, T. Sakano, Organic Solvents—Physical Properties and

Methods of Purification, Techniques of Chemistry, New York, 1986.[17] M. Calvin, K.W. Wilson, J. Am. Chem. Soc. 67 (1945) 2003–2007.[18] J. Bjerrum, Metal Amine Formation in Aqueous Solution, P Hasse and Son, Copenha-

gen, 1941.[19] L.G. Van Uitert, G.G. Hass, J. Am. Chem. Soc. 75 (1953) 451–454.[20] H. Irving, H.S. Rossoti, J. Chem. Soc. 12 (1954) 2904–2910.[21] A.T. Mubarak, A.S. Al-Shihri, H.M. Nassef, A.A. El-Bindary, J. Chem. Eng. Data 55

(2010) 5539–5542.[22] Y.M. Issa, O.E. Sherif, S.M. Abbas, Monatsh. Chem. 129 (1998) 985–998.[23] M. Mazik, W. Zielinski, Monatsh. Chem. 127 (1996) 587–591.[24] A.A.A. Boraei, J. Chem. Eng. Data 46 (2001) 939–943.[25] J.H. Van't Hoff, Etudes de Dynamique Chimique, Muller, Amsterdam, 1884. 114–118.


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