Synthesis, Characterization, and Thermochemical Study on

Article pubs.acs.org/jced

Synthesis, Characterization, and Thermochemical Study on Complexes of Praseodymium(III) with Niacin or 8‑Hydroxyquinoline Sheng-Xiong Xiao,†,‡ Jian-Hong Jiang,†,‡ Xu Li,†,‡ Xiao-Fang Zheng,‡ Hang-Ying Xiao,‡ Li-Juan Ye,†,‡ and Qiang-Guo Li*,†,‡ †

Hunan Provincial Key Laboratory of Rare-Precious Metals Compounds and Applications, Chenzhou 423000, Hunan Province, P. R. China ‡ Department of Chemistry and Life Science, Xiangnan University, Chenzhou 423000, Hunan Province, P. R. China ABSTRACT: This paper reports the synthesis and thermodynamic properties of three rare earth complexes, which were synthesized from praseodymium nitrate hexahydrate (Pr(NO3)3·6H2O) and heterocyclic ligands of niacin (vitamin PP, C6H5NO2) or 8-hydroxylquinoline (C9H7NO). Their compositions and structures were characterized by elemental analysis, molar conductance, thermogravimetric analysis, UV−visible spectroscopy, and IR spectroscopy. The ligand C6H5NO2 was bidentate-coordinated with Pr3+ ions through an oxygen atom of its carboxylic group which was formed by removing the proton from the hydroxyl group: the heterocyclic nitrogen atom from C9H6NO− formed a chelating ring around with Pr3+ ion. At a constant temperature of 298.15 K, the dissolution enthalpies of the reactants and products of the coordination reactions in the optimized calorimetric solvent were determined by an advanced solution−reaction isoperibol microcalorimeter. The standard molar enthalpies of the coordination reactions were determined. By combination of the experimental values of enthalpies of dissolution with some other auxiliary thermodynamic data through a designed thermochemical cycle on the basis of a supposed chemical reaction, the standard molar enthalpies of formation of the synthetic coordination complexes were estimated to be ΔfHΘm[Pr(C6H4NO2)2C9H6NO(s), 298.15 K] = −(1448.3 ± 2.6) kJ·mol−1, ΔfHΘm[Pr(C6H4NO3)3·2H2O(s), 298.15 K] = −(2256.1 ± 2.9) kJ·mol−1, and ΔfHΘm[Pr(C9H6NO)3· 4H2O(s), 298.15 K] = −(1975.3 ± 4.6) kJ·mol−1.

1. INTRODUCTION It is known that rare earth ions possess the properties of antibacterial, antitumor, and antivirus agents when coordinated to some small organic molecules ligands.1,2 Due to their strong affinity to many biological molecules, rare earth ions can effectively participate in many important life processes and activities or inhibit a variety of enzymes or pro-enzymes. Many recent researches showed that vitamin PP has obvious effects on dilating vessels, decreasing blood lipid, inhibiting synthesis of cholesterol, dissolving fibrin, and restraining the formation of thrombus.3 Interestingly, complexing ligands such as 8hydroxylquinoline and its derivatives were also found to have potential bioactivities, including anticancer, antibacterial, and antioxidative.4−6 Accordingly, rare earth coordination complexes have been widely and deeply researched by many researchers all over the world in recent years, especially in China. It is well-known that the data of the standard molar enthalpy of formation plays a very vital role in theoretical studies, using in the development of industrial compounds on basis of theoretical analysis. The standard enthalpy of formation, together with the standard entropy, are equally important data in determining any chemical equilibria.7 As what we have mentioned above, researches on the thermo-kinetics and biological properties of rare earth coordination complexes are quite active. However, © 2013 American Chemical Society

only a few studies have been devoted to the determination of standard enthalpies of formation of rare earth coordination complexes especially in the situation that the thermodynamic fundamental database of rare earth coordination complexes are still insufficient. Solution-reaction calorimetry is a broad and versatile technique and has the advantages of being rapid, accurate, economic, and convenient. It is widely applied to the measurements of enthalpies of reaction, dissolution, dilution, mixing, adsorption, formation, and excess enthalpies.8−10 To further study the thermodynamic properties of rare earth coordination complexes, this article reports the synthesis and characterization of coordination complexes of [Pr(C6H4NO2)2C9H6NO], [Pr(C6H4NO3)3·2H2O], and [Pr(C9H6NO)3·4H2O]. In particular, its thermodynamic properties were investigated by an advanced solution−reaction isoperibol microcalorimeter11 which greatly improved the accuracy of the dissolution enthalpies of samples, with an accuracy of 99.5%. Obviously, these values are indispensable for the strict modern experimental technology. We believe that this Received: July 4, 2013 Accepted: September 9, 2013 Published: September 23, 2013 2868

dx.doi.org/10.1021/je4006294 | J. Chem. Eng. Data 2013, 58, 2868−2874

Journal of Chemical & Engineering Data

Article

Figure 1. Thermochemical cycle of the coordination reactions.

3.2. Methods. Equipment used was as follows: Fourier IR spectrometer (Avatar 360, Nicolet, USA, KBr S.P., using KBr pressed pellet method); UV−visible spectrophotometer (U3010, Hitachi, Japan, using the conjugate principle of molecules); thermogravimetry analysis instrument (STA449C, Netzsch, Germany, weight loss rate of high temperature); elemental analyzer (Perkin-Elmer 2400 CHN, USA); digital Abbe refractometer (WAY-IS, Shanghai Precision & Scientific Instrument Co., Ltd., China). The dissolution enthalpies were measured by a solution− reaction isoperibol microcalorimeter (SRC-100, constructed by the Thermochemical Laboratory of Wuhan University, China).11 In the experimental process, the temperature was 298.15 K, the electric current has the intensity was 11.714 mA, and the resistance of heater was 1245.2 Ω. The calibration of the calorimeter was carried out by measuring the dissolution enthalpies of KCl (calorimetric primary standard) in three times distilled water and trihydroxymethyl aminomethane (THAM, NBS 742a, USA) in 0.0001 mol·cm−3 HCl at 298.15 K. The mean dissolution enthalpies were (17597 ± 17) J·mol−1 for KCl and −(29776 ± 16) J·mol−1 for THAM, which match well with the published data [(17536 ± 9) J·mol−1 for KCl11 and −(29766 ± 31.5) J·mol−1 for THAM12]. The eventual errors of the experimental results were within ± 0.5 % compared with the recommended reference data, demonstrating that the microcalorimeter was practicable. 3.3. Synthesis and Characterization of the Title Complexes. Pr(NO 3 ) 3 ·6H 2 O, NaOH, C 6 H 5 NO 2 , and C9H7NO were dissolved in appropriate amounts of anhydrous ethanol, respectively. Quantitative NaOH solution (20 mmoL of NaOH was dissolved in 40 mL of three times distilled water) was put into C6H5NO2 solution (20 mmoL of C6H5NO2 was dissolved in 20 mL of anhydrous ethanol) and placed into a three-necked round flask. C9H7NO solution (10 mmoL of C6H5NO2 was dissolved in 20 mL of anhydrous ethanol) was added dropwise with electromagnetic stirring. At a constant temperature of 333.15 K, Pr(NO3)3·6H2O solution (10 mmol of powdered Pr(NO3)3·6H2O(s) was dissolved in 50 mL of three times distilled water) was added dropwise to the previous mixture. Then, about 250 mL of the reaction solution was stirred for 6 h (the value of pH was adjusted to 6.5 to 7.0 by adding NaOH solution). After overnight deposition and air pump filtration, a solid complex was acquired. This product was washed alternatively with ethanol and water at the temperature of 353.15 K until no NO3− was detected in the filtrate. After that, the product was dried to a constant mass in a vacuum desiccator at 333.15 K. Subsequently, the contents of C, H, and N were determined by elemental analysis individually. The

type of study may potentially benefit researches and contribute to the development of the rare earth coordination chemistry.

2. THEORY According to Hess’s law, the standard molar enthalpy change of a chemical reaction can be deduced from the dissolution enthalpies which could be measured when the reactants and products were dissolved in an optimal calorimetric solvent, that is: Θ Δr HmΘ = −∑ νBΔsHm, T ,B B

(a)

Also, it can be calculated from the standard molar enthalpies of formation of the relevant reactants and products in a chemical reaction, that is: Δr HmΘ =

∑ νBΔf Hm,Θ T ,B B

(b)

In the above presented calculation formulas, ΔrHΘm represents the standard molar enthalpy change of a chemical reaction; ΔrHΘm,T,B stands for the dissolution enthalpy of a substance B when it was dissolved in the optimal calorimetric solvent at temperature T; Δ3HΘm,T,B denotes the standard molar enthalpy of formation of a substance B at temperature T; and vB is the stoichiometric number of a corresponding substance B in a chemical reaction equation, having negative values for reactants and positive values for products.

3. EXPERIMENTAL DETAILS 3.1. Chemicals and Reagents. Pr2O3 (> 99.99 %, produced by Tianjin Guangfu Fine Chemical Research Institute); Pr(NO3)3·6H2O (synthesized from the reaction of Pr2O3 with 6.0 mol·L−1 HNO3 followed by crystallization); C6H5NO2 (> 99.5 %, A.R., produced by the Shanghai Reagent Company, recrystallized from anhydrous ethanol); C9H7NO (> 99.5 %, A.R., obtained from Tianjin Guangfu Fine Chemical Research Institute, recrystallized with anhydrous ethanol) were used. The above presented substances were dried in a vacuum desiccator containing P4O10 to a constant mass at room temperature. KCl (calorimetric primary standard purity is greater than 99.99 %) was dried in a vacuum oven for 6 h at 408.15 K. Al2O3 (spectral grade with the purity > 99.99 %, purchased from Shanghai No. 1 Reagent Company, can be used after dryness under 2173.15 K for 2 h, making the entire sample transformed into α-phase). HNO3 (A.R.); NaOH (A.R.); dimethylformamide (DMF, A.R.); ethanol (EtOH, A.R.); three times distilled water were also used. 2869



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Table 1. Dissolution Enthalpies of [2C6H5NO2(s)], [C9H7NO(s)], [Pr(NO3)3·6H2O(s)], [Pr(C6H4NO2)2C9H6NO(s)], and Solution E in the Calorimetric Solvent S at 298.15 K ΔsHΘm/(kJ·mol−1)

a

sample

2C6H5NO2(s)

C9H7NO(s)

Pr(NO3)3·6H2O(s)

[Pr(C6H4NO2)2C9H6NO](s)

solution E

1 2 3 4 5 6 7 avg.

10.477 10.298 10.395 10.304 9.949 10.751 10.245 10.346 ± 0.243a

−5.816 −5.470 −5.433 −5.125 −5.207 −5.501 −5.440 −5.427 ± 0.223

19.193 18.791 19.701 19.650 18.670 19.038 19.005 19.150 ± 0.398

−44.871 −45.449 −45.170 −46.842 −46.008 −46.519 −44.951 −45.687 ± 0.781

−27.920 −27.564 −27.341 −27.352 −27.262 −28.183 −27.415 −27.577 ± 0.346

The uncertainty was estimated as the standard deviation by the Bessel equation.

Table 2. Dissolution Enthalpies of [3C6H5NO2(s)], [Pr(NO3)3·6H2O(s)], [Pr(C6H4NO3)3·2H2O(s)], and Solution E in the Calorimetric Solvent S at 298.15 K ΔsHΘm/(kJ·mol−1) sample

3C6H5NO2(s)

Pr(NO3)3·6H2O(s)

[Pr(C6H4NO3)3·2H2O](s)

solution E

1 2 3 4 5 6 7 avg.

13.143 13.076 12.861 13.247 13.139 12.925 13.128 13.074 ± 0.135

18.290 18.801 17.897 18.375 18.456 18.688 18.182 18.384 ± 0.305

−52.433 −52.471 −50.731 −52.433 −50.611 −50.640 −51.646 −51.566 ± 0.894

−27.608 −27.257 −26.598 −27.050 −26.867 −27.121 −26.795 −27.042 ± 0.332

content of Pr3+ was determined by the EDTA titration after the complex was decomposed by heating together with concentrated hydrochloric acid, and the content of crystal water was determined by thermogravimetric analysis (TG-DSC). The analytical results showed that the general formula of the complex was [Pr(C6H4NO2)2C9H6NO] and its purity was more than 99.0 %. By the similar method, we synthesized orther two binary complexes: [Pr(C6H4 NO3) 3·2H 2O] and [Pr(C9H6NO) 3· 4H2O]. 3.4. Determination of the Molar Conductance of the Complexes. The complexes were prepared to solutions with concentrations of 0.01 mol/L by dissolving them in pure dimethyl sulfoxide (DMSO). Measuring the electric conductance of 0.02 mol/L KCl solution when the thermostatic bath temperature was adjusted to 25 ± 0.1 °C, the constant of a conductivity cell was calculated. The same steps were then followed to measure the electric conductivity of the complexes and deduct the conductivity of the solvent of DMSO, and the molar conductivity of complexes were determined; concrete data are listed in Table 4. 3.5. Determination of Dissolution Enthalpies. 3.5.1. Thermochemical Cycle of the Coordination Reaction. As the thermal effect of solid state coordination reaction was difficult to be determined, it was feasible to deduce the enthalpies of formation from the dissolution enthalpies measured when the samples were dissolved in the calorimetric solvent. In this paper, a convincing thermochemical cycle based on Hess’s law was designed and is shown in Figure 1. In Figure 1, C6H5NO2 is niacin, and C9H7NO is 8hydroxylquinoline. The equations for the coordination reactions are as follows:

Pr(NO3)3 · 6H 2O(s) + 2C6H5NO2 (s) + C9H 7NO(s) → [Pr(C6H4NO2 )2 C9H6NO](s) + 3HNO3(g) + 6H 2O(l)

(1)

Pr(NO3)3 ·6H 2O(s) + 3C6H5NO2 (s) → [Pr(C6H4NO3)3 ·2H 2O](s) + 3HNO3(g) + 4H 2O(l)

(2)

Pr(NO3)3 · 6H 2O(s) + 3C9H 7NO(s) → [Pr(C9H6NO)3 ·4H 2O](s) + 3HNO3(g) + 2H 2O(l)

(3)

The UV spectra and refractive indexes of the final solution of reactants and products can be used to determine whether they have the same thermodynamic state. In the present experiment, the UV spectra and refractive indexes of solution C and solution F were determined. The experimental results show that both solutions have similar UV spectrum curves and equal refractive indexes, which proves that they have the same thermodynamic state and the thermochemical cycles of the coordination reactions designed were reliable. 3.5.2. Choice of Calorimetric Solvent. It is very important to choose a solvent which could dissolve all of the reactants and products rapidly and completely. Research indicated that the relevant substances in the coordination reaction are highly soluble in the mixed solvent of 2 mol·L−1 HCl, DMF, and EtOH. By mixing these three solvents uniformly at different ratios and testing the solubility of samples continually. The best solubility appeared when the volume ratio of the three solvents was VHCl/VDMF/VEtOH = 1:1:1. Consequently, it was chosen as the optimized calorimetric solvent S. 3.5.3. Determination of Dissolution Enthalpies of Reactants and Products. Thoroughly dried samples were ground 2870



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Table 3. Dissolution Enthalpies of [3C9H7NO(s)], [Pr(NO3)3·6H2O(s)], [Pr(C9H6NO)3·4H2O(s)], and Solution E in the Calorimetric Solvent S at 298.15 K ΔsHΘm/(kJ·mol−1) sample

3C9H7NO (s)

Pr(NO3)3·6H2O(s)

[Pr(C9H6NO)3·4H2O](s)

solution E

1 2 3 4 5 6 7 avg.

−5.845 ν-6.363 −6.096 −6.163 −6.144 −6.088 −6.032 −6.105 ± 0.155

19.153 18.426 19.076 18.292 18.944 18.903 18.017 18.687 ± 0.439

−128.360 −128.086 −127.661 −128.005 −127.096 −127.275 −127.011 −127.642 ± 0.529

−37.156 −36.616 −36.541 −37.564 −36.833 −36.446 −36.699 −36.837 ± 0.396

Table 4. Results of Elemental Analysis and Molar Conductance of Complexes results of elemental analysis observed/% (calculated/%)

molar conductivity

sample

C%

H%

N%

Pr %

product’s formula

S·cm2·mol−1

1 2 3

39.69(39.8) 50.36(50.25) 47.66(47.89)

3.05(2.97) 4.15(4.06) 2.67(2.76)

7.63(7.74) 6.43(6.51) 7.94(8.01)

25.78(25.94) 21.68(21.83) 26.62(26.91)

[Pr(C6H4NO3)3·2H2O] [Pr(C9H6NO)3·4H2O] [Pr(C6H4NO2)2C9H6NO]

17.44 20.90 10.35

completely in an agate mortar; then 0.5 mmol of samples were placed exactly into the sample container of the microcalorimeter. The calorimetric solvent S (100.00 mL) was poured into the reaction vessel in advance. When the calorimeter was adjusted to a constant temperature of (298.150 ± 0.001) K, the samples were added into the reaction vessel and their dissolution enthalpies measured. After seven parallel measurements, the results were listed in Tables 1 to 3.

4. RESULTS AND DISCUSSION 4.1. Elemental Analysis and General Properties of the Complexes. The theoretical values and found values of the elemental analysis for [Pr(C 6 H 4 NO 2 ) 2 C 9 H 6 NO], [Pr(C6H4NO3)3·2H2O] and [Pr(C9H6NO)3·4H2O] are as follows: [Pr(C6H4NO2)2C9H6NO] and [Pr(C9H6NO)3·4H2O] were obtained as yellow powders, and [Pr(C6H4NO3)3·2H2O] is a light green powder. These compounds are very stable in the atmosphere. It should be noted that the compound of [Pr(C9H6NO)3·4H2O] can be synthetized in a wide range of pH. Even in a high value of pH, a light green product of Pr(OH)3 appeared first and then gradually changed to yellow. These all show the high theromdynamic stabilities of the compound. These compounds are soluble in dimethyl sulfoxide (DMSO) but cannot be soluble in water, chloroform, methanol, ethanol, dimethylformamide, acetone, petroleum ether, or tetrahydrofuran. The molar conductance of the complexes in DMSO were determined, and corresponding results are listed in Table 4, indicating that these complexes are nonelectrolyte and exist as neutral molecules in DMSO. 4.2. Mechanism of Thermal Decomposition. TG-DSC curves (β = 10 °C·min−1, argon atmosphere) of one complex are shown in Figure 2. No obvious mass loss until 473.15 K indicates that no crystal water exists in the molecular complex. The thermal decomposition process can be divided into three steps. First stage, 1 mol of C9H6NO− (actual value: 25.76 %, theoretical value: 25.27 %) was eliminated from (660.49 to 738.23) K. The next two stages, 1 mol of (C6H4NO2)− decomposed; the mass losses were 20.45 % and 20.07 %, respectively, at the temperature goes from (738.23 to 821.95) K

Figure 2. TG-DSC curves (β = 10 °C·min−1) of the complex [Pr(C6H4NO2)2C9H6NO].

and (821.95 to 1011.15) K. which roughly coincides with the value of 21.19 %. The characteristic absorption bands of the IR spectra of the residue are similar to those of Pr2O3. So we concluded that the complex was decomposed to Pr2O3 completely (actual value: 33.72 %, theoretical value: 32.66 %). The high decomposition temperature of the complex suggested that the complex has a good thermal stabilization. On the basis of experimental and calculated results, the thermal decomposition of [Pr(C6H4NO2)2C9H6NO] can be postulated as follows:

By contrast, the thermal decomposition of the complex [Pr(C9H6NO)3·4H2O] could be divided into four stages. First, 2 mol of crystal water and coordination water were lost from (293.15 to 381.28) K and (381.28 to 475.79) K, respectively. Second, the loss of mass is 10.75 % from (475.79 to 696.49) K; this showed that 0.5 mol of C9H5NO− was eliminated, and the result was very close to the theory values of 10.55 %. Finally, with a mass loss of 21.10 % from (696.49 to 1108.15) K, this 2871



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value does not fix the theoretical value of 27.32 %, indicating that 1 mol of C9H5NO− was lost incompletely. The thermal decomposition of the complex [Pr(C6H4NO3)3· 2H2O] could be divided into four steps. At first, from (408.37 to 451.77) K and (451.77 to 493.28) K, 1.5 mol and 0.5 mol of water was lost, respectively. According to its high dehydration temperature, this was recognized as the coordination water. Next, from (649.19 to 761.75) K with mass loss of 10.22 % (theoretical loss: 10.50 %), indicating that 0.5 mol of C6H4NO3− was eliminated. Then, with mass loss of 23.72 % (theoretical loss: 21.10 %) from (761.75 to 851.54) K, 1 mol of C6H4NO3− was lost. Finally, with mass loss of 25.59 % (theoretical loss: 21.01 %) from (851.54 to 1111.55) K attributing to the loss of 1 mol of C6H4NO3− and leaving a residue of 32.72 % (calculated 30.36 %). In all cases, the high decomposition temperature of the complexes indicated that the complexes have a good thermal stabilization. 4.3. UV Spectra of the Complexes. The 0.1 mmol/L solution of the synthetic complexes [Pr(C6H4NO2)2C9H6NO], [Pr(C9H6NO)3·4H2O], and [Pr(C6H4NO3)3·2H2O] was prepared with free ligands C6H5NO2 and C9H7NO by using the mixed solvent (VDMSO/VEtOH = 1:9). The UV spectra of the complexes and free ligands are shown in Figure 3. C6H5NO2

n−π* transition of the phenolic hydroxyl oxygen and cyclobenzene. Comparing with the UV spectra of the synthetic complexes and free ligands, we can drew three conclusions. From the picture, it is obvious that the absorption bands of the complexes are harmony with the free ligands, it was not only prove that Pr3+ has a small effect on the conjugate of the aromatic ring, but also prove that the coordination of rare earth is priority to the nature of typical ionic bond . By contrast, the complex [Pr(C9H6NO)3·4H2O] has a strongest and sharpest absorption peak at 246 nm, and the intensity of absorption is almost 3 times of 8-hydroxyquinoline. In addition, the complex [Pr(C6H4NO2)2C9H6NO] has a wide absorption band from (240 to 278) nm, and it totally covers the absorption bands of the two free ligands. But the intensity of its absorption band is slightly weak than the complex [Pr(C9H6NO)3·4H2O], so this is closely associated with the number of the quinoline ring. 4.4. IR Spectra of the Complexes. The main IR bands of the complexes are presented in Table 5. It is observed that the characteristic absorption bands of the complexes are different of the absorption bands of the free ligands as shown in Table 5 and Figure 4.

Figure 3. UV spectra of the free ligands and the synthetic complexes. Figure 4. IR spectra of the free ligands and the synthetic complexes.

has the maximum absorption at 263 nm due to the π−π* transition of the carboxyl and aromatic heterocyclic. Free C9H7NO has a sharp and strong absorption peak at 246 nm ascribed to the absorption spectrum of the π−π* transition of the condensed nucleus. In addition, there is a weak broad peak at 316 nm, which belongs to the absorption spectrum of the

According to the IR spectra analysis, all of the characteristic bands of Pr(C6H4NO3)3·2H2O and Pr(C9H6NO)3·4H2O are present in the IR spectra of Pr(C 6 H 4 NO 2 ) 2 C 9 H 6NO, illustrating that both niacin and 8-hydroxyquinoline are coordinated to Pr3+. Then, Pr(C6H4NO2)2C9H6NO and Pr(C9H6NO)3·4H2O have an absorption band centers on

Table 5. IR Data of Ligands and the Complex [Pr(C6H4NO2)2C9H6NO] (cm−1)a sample C6H5N O2(s) C9H7NO(s) [Pr(C6H4NO2)3·2H2O] [Pr(C9H6NO)3·4H2O] [Pr(C6H4NO2)2C9H6NO]

νO−H 2200 2700 3136 3392 3418

(m) (w) (w) (w) (w)

δO−H

νCO (COO−)

νas (COO−)

νs (COO−)

951 (w)

1715 (s)

1596 (w)

1417 (m)

1598 (s)

1404 (vs)

1379 (s) 1380 (vs) 1381 (vs)

νCC+CN

νC−N

1576 (m)

1597 (m)

1405 (m)

νC−O

νPr−O

1322 (s) 1299 (s) 1284 (m) 424 (m)

1570 (m) 1569 (m)

1314 (vs) 1314 (vs)

1274 (m) 1274 (m)

420 (w)

ν stands for stretching vibration; νs stands for symmetrical stretching vibration; νas stands for asymmetrical stretching vibration; δ stands for deformation vibration.

a

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1314 cm−1(νC−N), but each ligand and Pr(C6H4NO3)3·2H2O do not have any absorption band in this place. This indicated that the nitrogen atoms of thick heterocyclic of 8hydroxyquinoline took part in coordination and the nitrogen atom of pyridine ring of nicotinic acid did not participate in coordination. It is known from Cu(C9H6NO)2 IR spectra that C9H7NO coordinates to Pr3+ by the hydroxyl oxygen atom and heterocyclic nitrogen atom forming a five-membered chelate ring, which increased the conjugation degree of quinoline ring even decreasing the bond strength of CC and CN. After coordination, three characteristic absorption bands of carboxylate group in ligand C6H5NO2 molecule vanished, while the symmetric stretching vibration absorption band and asymmetric stretching vibration absorption band of COO − remained. The fact proved that C6H5NO2 removed the proton of carboxylic acid and coordinates as a bidentate group to Pr3+. The single IR absorption band of the coordination complex presenting at 420 cm−1, which was attributed to the stretching vibration of the Pr−O bond. Based on the above IR spectra, the data from elemental analysis and thermogravimetric analysis, the chemical structure of the coordination complex is given in Figure 5.

Δr HmΘ(1) =

∑ vBΔsHm,Θ T ,B = ΔsHmΘ[C9H 7NO(s) B

, 298.15 K] + ΔsHmΘ[2C6H5NO2 (s), 298.15 K] + ΔsHmΘ[Pr(NO3)3 ·6H 2O(s), 298.15 K] − ΔsHmΘ[Pr(C6H4NO2 )2 C9H6NO(s), 298.15 K] − ΔsHmΘ[HNO3(l), 298.15 K] − ΔsHmΘ [solution E, 298.15 K] = 19.1497 + 10.3456 + ( −5.4274) − ( −45.6871) − ( −27.5767) − ( −61.38) 0.3982 + 0.2432 + 0.2232 + 0.7812 + 0.3462

±

= 158.71 ± 0.92 kJ ·mol−1

According to refs 14 to 17, Δf HmΘ[HNO3(l), 298.15 K] = −174.10 kJ ·mol−1 Δf HmΘ[H 2O(l), 298.15 K] = −(285.83 ± 0.04) kJ ·mol−1

Δf HmΘ[C6H5NO2 (s), 298.15 K] = −(344.81 ± 0.92) kJ ·mol−1 Δf HmΘ[C9H 7NO(s), 298.15 K] = −(83.0 ± 1.5) kJ ·mol−1 Δf HmΘ[Pr(NO3)3 ·6H 2O(s), 298.15 K] = −3071.7 kJ·mol−1

Also, through eq b, we can obtain: Δf HmΘ[Pr(C6H4NO2 )2 C9H6NO(s), 298.15 K] =Δr HmΘ(l) + 2Δf HmΘ[C6H5NO2 (s), 298.15 K] + Δf HmΘ[C9H 7NO(s), 298.15 K] + Δf HmΘ[Pr(NO3)3 ·6H 2O(s), 298.150] − 3Δf HmΘ[HNO3(l), 298.15 K] − 6Δf HmΘ[H 2O(l), 298.15 K] Figure 5. Chemical structure of the complex [Pr(C6H4NO2)2C9H6NO].

= 158.71 + (− 3071.7) + (− 83.0) + 2(− 344.81) − 3( −174.10) − 6(− 285.83)

4.5. Calorimetric Determination of Enthalpies of Formation of the Complexes. 3HNO3(1) + 6H 2O(l) → solution E

±

= −(1448.3 ± 2.6) kJ·mol−1

(6)

The molality of solution E is 27.78 mol·kg−1. According to ref 13 ΔsHmθ (HNO3(l),

0.922 + 1.52 + (0.92· 2)2 + (0.04· 6)2

In the similar way, we can obtain: Δr HmΘ(2) = (171.45 ± 0.96) kJ ·mol−1

−1

298.15 K) = −33.28 kJ ·mol

Δf HmΘ[Pr(C6H4NO3)3 · 2H 2O(s), 298.15 K]

ΔsHmθ (solution E(aq), 298.15 K) = −ΦL(27.78 → 0)

= −(2256.1 ± 2.9) kJ·mol−1

−1

= −12.82 kJ ·mol

Δr HmΘ(3) = (238.44 ± 0.68) kJ ·mol−1

So that

Δf HmΘ[Pr(C9H6NO)3 · 4H 2O(s), 298.15 K]

θ θ ΔsHmθ (6) = 3Δd Hm( ∞→ 27.78) = Δs Hm(HNO3(l)

= −(1975.3 ± 4.6) kJ·mol−1

, 298.15 K) − ΔsHmθ (m = 27.78 mol · kg −1)

4.6. Experimental Discussion. As was shown by the above calculated results, the standard molar enthalpy change of the coordination reaction, ΔrHΘm(1) = (158.71 ± 0.92) kJ·mol−1, is

= 3[−33.28 − ( −12.82)] kJ ·mol−1 = −61.38 kJ ·mol−1

According to eq a, we have 2873



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a high positive value. This indicates that the reactants cannot react spontaneously in the solid phase. Since the thermal effect of solid state coordination reaction is difficult to be determined, if we want to measure it, some other methods should be employed. In this work, we designed a convincing thermochemical cycle based on Hess’s law and deduced the standard molar enthalpies of formation from the dissolution enthalpies which were measured by an advanced solution−reaction isoperibol microcalorimeter when the samples were dissolved in the optimized calorimetric solvent.

Notes

The authors declare no competing financial interest.

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(1) Yang, L.; Tao, D.; Yang, X.; Li, Y.; Guo, Y. Synthesis, Characterization, and Antibacterial Activities of Some Rare Earth Metal Complexes of Pipemidic acid. Chem. Pharm. Bull. 2003, 51, 494−498. (2) Zhou, J.; Wang, L. F.; Wang, J. Y.; Tang, N. Synthesis, Characterization, Antioxidative and Antitumor Activities of Solid Quercetin Rare Earth (III) Complexes. J. Inorg. Biochem. 2001, 83, 41−49. (3) Di, Y. Y.; Kong, Y. X.; Zhang, S.; Yang, W. W.; Wu, E. S.; Shi, Q.; Tan, Z. C. Synthesis, Characterization and Thermochemistry of the Hydrated Barium Nicotinate. Acta Phys. Chim. Sin. 2008, 24, 1884− 1890. (4) Darby, C. M.; Nathan, C. F. Killing of Non-Replicating Mycobacterium Tuberculosis by 8-Hydroxyquinoline. J. Antimicrob. Chemother. 2010, 65, 1424−1427. (5) Sashidhara, K. V.; Kumar, A.; Bhatia, G.; Khan, M. M.; Khanna, A. K.; Saxena, J. K. Antidyslipidemic and Antioxidative Activities of 8Hydroxyquinoline Derived Novel Keto-Enamine Schiffs Bases. Eur. J. Med. Chem. 2009, 44, 1813−1818. (6) Liu, Y.; Zhao, Y.; Zhai, X.; Feng, X.; Wang, J.; Gong, P. Synthesis and Anti-hepatitis B Virus Evaluation of Novel Ethyl 6-Hydroxyquinoline-3-carboxylates in Vitro. Bioorg. Med. Chem. 2008, 16, 6522−6527. (7) Suga, H. Perspectives of Low Temperature Calorimetry. Thermochim. Acta 2000, 355, 69−82. (8) Ferguson, H. F.; Frurip, D. J.; Pastor, A. J.; Apeerey, L. M.; Whiting, L. F. A Review of Analytical Applications of Calorimetry. Thermochim. Acta 2000, 363, 1−21. (9) Xiao, S. X.; Gu, H. W.; Zhang, J. J.; Xiao, H.; Ding, J.; Lu, X. Synthesis, Characterization, and Thermochemical Study on a Ternary Complex [Sm(m-MOBA)3phen]2. Int. J. Thermophys. 2012, 33, 289− 299. (10) Ribeiro da Silva, M. A. V.; Ribeiro da Silva, M. D. M. C.; da Silva, L. C. M.; Dietze, F.; Hoyer, E. Thermochemical Studies of Two Copper (II) Complexes with N-benzoyl-N′,N′-dialkylurea Derivatives. Thermochim. Acta 2001, 378, 45−50. (11) Yu, H. G.; Liu, Y.; Tan, Z. C.; Dong, J. X.; Zou, T. J.; Huang, X. M.; Qu, S. S. A Solution-Reaction Isoperibol Calorimeter and Standard Molar Enthalpies of Formation of Ln(hq)2Ac (Ln = La, Pr). Thermochim. Acta 2003, 401, 217−224. (12) Rychly, R.; Pekarek, V. The Use of Potassium Chloride and Tris(hydroxymethyl) Aminomethane as Standard Substances for Solution Calorimetry. J. Chem. Thermodyn. 1977, 9, 391−396. (13) Lide, D. R. CRC Handbook of Chemistry and Physics, 84th ed.; CRC Press Inc.: Boca Raton, FL, 2003/2004; pp 5−102. (14) Cox, J. D.; Wagman, D. D.; Medvedev, V. A. CODATA Key Values for Thermodynamics; Hemisphere Publishing Corp.: New York, 1984; p 1. (15) Sato-Toshima, T.; Kamagughi, A.; Nishiyama, K.; Sakiyama, M.; Seki, S. Enthalpies of Combustion of Organic Compounds. IV. Acetaneilide and Nicotinic Acid. Bull. Chem. Soc. Jpn. 1983, 56, 51−54. (16) Ribeiro da Silva, M. A. V.; Monte, M. J. S.; Matos, M. A. R. Enthalpies of Combustion, Vapour Pressures, and Enthalpies of Sublimation of 8-Hydroxyquinoline, 5-Nitro-8-hydroxyquinoline, and 2-Methyl-8-hydroxyquinoline. J. Chem. Thermodyn. 1989, 21, 159− 166. (17) Weast, R. C. CRC Handbook of Chemistry and Physics, 69th ed.; CRC Press Inc.: Boca Raton, FL, 1988/1989; p D-121.

5. CONCLUSIONS Three praseodymium coordination complexes [Pr(C6H4NO2)2C9H6NO], [Pr(C6H4NO3)3·2H2O], and [Pr(C9H6NO)3·4H2O] were synthesized by using three starting materials: Pr(NO3)3·6H2O, C9H7NO, and C6H5NO2. Their compositions and structures were characterized by elemental analysis, molar conductance, thermogravimetric analysis, UV spectroscopy, and IR spectroscopy. The analytical results showed that their purity were more than 99.0 %. Particularly, an advanced solution−reaction isoperibol microcalorimeter was employed to determine the standard molar enthalpies of formation of the synthetic coordination complexes. According to Hess’s law and thermodynamic principles, the reasonable thermochemical cycles were designed. At 298.15 K, the dissolution enthalpies were measured when relevant substances were dissolved in the optimized calorimetric solvent. The standard molar enthalpy changes of the reactions were determined based on the experimental data to be ΔrHΘm(1) = (158.71 ± 0.92) kJ·mol−1, ΔrHΘm(2) = (171.45 ± 0.96) kJ· mol−1, and ΔrHΘm(3) = (238.44 ± 0.68) kJ·mol−1. From the above molar enthalpy changes of the coordination reactions and other auxiliary thermodynamic quantities, the standard molar enthalpies of formation of the three synthetic coordination complexes were derived to be Δ fH mΘ [Pr6H4NO2)2C9H6NO(s), 298.15 K] = −(1448.3 ± 2.6) kJ· mol−1, ΔfHΘm[Pr(C6H4NO3)3·2H2O(s), 298.15 K] = −(2256.1 ± 2.9) kJ·mol−1, and ΔfHΘm[Pr(C9H6NO)3·4H2O(s), 298.15 K] = −(1975.3 ± 4.6) kJ·mol−1. The present work demonstrates that solution−reaction microcalorimetry is a very useful tool for the thermodynamic research, as it is capable of providing accurate thermodynamic quantities of many significant substances in industry and scientific research. This work provides sufficient thermal and thermodynamic data of rare earth coordination complexes, and this information will also enrich and develop the thermodynamic fundamental database of rare earth coordination complexes and consequently may have theoretical instructive significance and application values for the development of rare earth coordination complexes.

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REFERENCES

AUTHOR INFORMATION

Corresponding Author

*E-mail address: [email protected]. Tel.: +86 735 2653128. Fax: +86 735 2653053. Funding

This research was financially supported by the National Natural Science Foundation of China (nos. 21273190 and 20973145), Science and Technology Plan Projects of Hunan Province, China (nos. 2012TP4021-6 and 2010FJ4080), Key Project of Hunan Provincial Key Laboratory Opening Topic Fund 2874


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