Calculation of the enthalpies of formation and proton affinities of some isoquinoline derivatives

J. Chem. Thermodynamics 40 (2008) 1627–1631

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Calculation of the enthalpies of formation and proton affinities of some isoquinoline derivatives Mansoor Namazian *, Michelle L. Coote * ARC Centre of Excellence for Free-Radical Chemistry and Biotechnology, Research School of Chemistry, Australian National University, Canberra ACT 0200, Australia

a r t i c l e

i n f o

Article history: Received 4 April 2008 Received in revised form 5 July 2008 Accepted 10 July 2008 Available online 18 July 2008 Keywords: Enthalpy of formation Isoquinoline G3 Proton affinity

a b s t r a c t Ab initio molecular orbital theory has been used to calculate enthalpies of formation of isoquinoline, 1hydroxyisoquinoline, 5-hydroxyisoquinoline, and 1,5-dihydroxyisoquinoline as well as some pyridine and quinoline derivatives. The proton affinities of the four isoquinoline derivatives were also obtained. The high-level composite methods G3(MP2), G3(MP2)//B3LYP, G3//B3LYP, and CBS-QB3 have been used for this study, and the results have been compared with available experimental values. For six of the eight studied compounds, the theoretical enthalpies of formation were very close to the experimental values (to within 4.3 kJ  mol1); where comparison was possible, the theoretical and experimental proton affinities were also in excellent agreement with one another. However, there is an extraordinary discrepancy between theory and experiment for the enthalpies of formation of 1-hydroxyisoquinoline and 1,5-dihydroxyisoquinoline, suggesting that the experimental values for these two compounds should perhaps be reexamined. We also show that popular low cost computational methods such as B3LYP and MP2 show very large deviations from the benchmark values. Ó 2008 Elsevier Ltd. All rights reserved.

1. Introduction Nitrogen heterocycles are an important feature in the chemical structures of many natural products, and often impart specific biological properties. An important class of such compounds are the isoquinolines, many of which have varying degrees of neurological effects, ranging from relaxation and euphoria to seizures and vasodilatation [1,2]. Their derivatives have potential therapeutic application in several diseases including cancer, diabetes, stroke, rheumatoid arthritis, and hemorrhagic shock [3–5]. Isoquinolines can be found in varying quantities in the prickly poppy, bloodroot, and celandine poppy [1,2]. The enthalpy of formation for a compound from its elements in their reference states is an important concept in chemical thermodynamics, and can be used to obtain the change of enthalpy for a chemical reaction [6]. The enthalpy of formation of isoquinoline (1) was first reported by Steele et al. [7] (see figure 1 for the structure). Recently, Riberio da Silva and his co-workers reported enthalpies of formation of 1-hydroxyisoquinoline (2), 5-hydroxyisoquinoline (3), and 1,5-dihydroxyisoquinoline (4), as measured by static bomb-combustion calorimetry [1,2]. While the enthalpies of formation of gaseous isoquinoline and 5-hydroxyisoquinoline were reported to be both positive (200.54 ± 0.94 kJ  mol1 and

22.8 ± 3.0 kJ  mol1, respectively), the enthalpies of formation of gaseous 1-hydroxyisoquinoline and 1,5-dihydroxyisoquinoline were reported as (33.2 ± 2.9 and 224.3 ± 3.3) kJ  mol1. It is of interest to examine theoretically why 1-hydroxyisoquinoline and 1,5-dihydroxyisoquinoline have such negative values for enthalpies of formation. Moreover, through the use of high-level ab initio calculations, it is also possible to confirm the experimental values of the enthalpies of formation, since some of these appear to be lower when compared with similar substituted pyridines [1,2]. In the present work, we have studied four isoquinoline derivatives (isoquinoline, 1-hydroxyisoquinoline, 5-hydroxyisoquinoline, and 1,5-dihydroxyisoquinoline) as well as four pyridine and quinoline derivatives (pyridine, 2-hydroxypyridine, quinoline, and 2hydroxyquinoline; figure 1), theoretically using high-level ab initio molecular orbital calculations. We show that, whilst the theoretical methods predict the enthalpies of formation for most of the studied compounds to within chemical accuracy, there is a significant discrepancy between theory and experiment for the cases of 1hydroxyisoquinoline and 1,5-dihydroxyisoquinoline. For further verification of the theoretical procedures, the proton affinities of these species have been also calculated and compared with available experimental values. 2. Computational methods

* Corresponding authors. Tel.: +61 2 6125 5411; fax: +61 2 6125 0750. E-mail addresses: [email protected] (M. Namazian), [email protected]. edu.au (M.L. Coote). 0021-9614/$ - see front matter Ó 2008 Elsevier Ltd. All rights reserved. doi:10.1016/j.jct.2008.07.007

Standard ab initio molecular orbital theory [8,9] and density functional theory calculations [10,11] were carried out using the

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a

were also performed using the lower levels of theory, B3LYP and MP2, with the 6-311+G(3df,2p) basis set, in order to establish whether these lower-cost procedures could deliver acceptable accuracy for the present systems.

OH N

N

1

3. Enthalpies of formation

2

The standard enthalpy of formation of a gaseous compound is the change of enthalpy of formation of 1 mol of the compound from its elements in their reference states at the pressure of one atmosphere and a temperature of 298 K [6]. In the present work the enthalpies of formation were calculated following the methodology described by Nicolaides et al. [6] and were thus obtained from the calculated atomization energies of the molecules at 298 K and the experimental heats of formation of the gaseous atoms at 298 K. These latter values were in turn obtained from the experimental heats of formation of the gaseous atoms at 0 K and the relevant experimental thermal corrections, as tabulated in the original work of Nicolaides et al. [6]. The G3//B3LYP calculated enthalpies of formation of compounds 1, 3, 5, 6, 7, and 8, as well as the corresponding experimental values are presented in table 1. The deviations of the theoretical values of the enthalpies of formation for these compounds from the experimental values are in the range of 0.6–4.3 kJ  mol1. These deviations are within the typical level of error of G3// B3LYP calculations; therefore, the good agreement between the calculated values and experimental values verifies the validity of employed theoretical method for the selected molecules. We have employed the same level of theory (G3//B3LYP) for calculation of enthalpies of formation of species 2 and 4, and the results are shown in table 2. The theoretical enthalpies of formation for 2 and 4 have been obtained as (17.9 and 195.9) kJ  mol1, respectively, which can be compared with the experimental values of (33.2 ± 2.9 and 224.3 ± 3.3) kJ  mol1. For these two species, the deviations of theory from experiment (15.3 kJ  mol1 and 24.2 kJ  mol1) are much larger than the typical level of error of the G3//B3LYP method. To test the G3//B3LYP method further, we then considered alternative high levels of theory, including CBS-QB3 [15] and other members of the G3 family, G3(MP2) [14], and G3(MP2)//B3LYP [13]. The results are also shown in table 2. In each case, we find that the various high-level ab initio procedures show good agreement with one another, further supporting the G3//B3LYP results for these two species. We also conducted a small assessment study in which we examined the effect of the geometry optimization on the calculated heats of formation. table 3 shows the heats of formation of species 1 and 2 as obtained from single point energy calculations at a consistent level of theory but using the various geometries employed by the G3//B3, CBS-QB3, and G3(MP2) procedures. Results obtained using the high-level procedure QCISD/6311+G(d,p) are provided as a benchmark. From table 3, it is seen

OH N N

OH

OH

3

4

b

OH N N

5

6 N

7

N

OH

8

FIGURE 1. (a) Isoquinoline derivatives: Isoquinoline (1), 1-hydroxyisoquinoline (2), 5-hydroxyisoquinoline (3), 1,5-dihydroxyisoquinoline (4). (b) Pyridine and quinoline derivatives: pyridine (5), 2-hydroxypyridine (6), quinoline (7), and 2-hydroxyquinoline (8).

Gaussian 03 [12] software. Geometries of all species were optimized at the B3LYP/6-31G(d) level of theory [10,11], and where necessary, extra care was taken to select the minimum-energy conformation via systematic conformational searching at this level. The nature of each stationary point was established via B3LYP/631G(d) frequency calculations. The optimized geometries have been used for further calculations. For calculation of energies and enthalpies, a variety of high-level ab initio methods from the G3 family [13], including G3//B3LYP [13], G3(MP2) [14], G3(MP2)// B3LYP [13], and also the Complete Basis Set (CBS) method, CBSQB3 [15] have been used. These methods approximate CCSD(T) or QCISD(T) calculations with a large triple zeta basis set via additivity and/or extrapolation procedures at the MP2 and/or MP4 levels of theory, and have been demonstrated to provide ‘‘chemical accuracy” (ca. 4.2 kJ  mol1) when assessed against large test sets of thermochemical data [13–15]. Full descriptions of these methods can be found in their original references, and a brief summary is provided in a recent publication in this journal [16]. Calculations

TABLE 1 G3//B3LYP results for calculation of energies and enthalpies of the studied isoquinolines, quinolines, and pyridines: isoquinoline (1), 5-hydroxyisoquinoline (3), pyridine (5), 2hydroxypyridine (6), quinoline (7), and 2-hydroxyquinoline (8) Entry

1

3

5

6

7

8

Energy (0 K) Energy (298 K) Enthalpy (298 K) DHf (g, 0 K) DHf (g, 298 K)

1054495.7 1054477.4 1054474.9 225.6 202.3 (2.3) 204.6 ± 0.9

1251917.0 1251895.6 1251893.1 48.4 23.7 (0.9) 22.8 ± 3.0

651382.0 651370.4 651367.9 156.9 139.8 (0.6) 140.4 ± 3.0

848839.4 848825.2 848822.8 56.4 75.4 (4.3) 79.7 ± 1.5

1054500.3 1054482.1 1054479.6 221.1 197.6 (2.9) 200.5 ± 0.9

1251962.9 1251942.0 1251939.5 2.5 22.7 (2.8) 25.5 ± 2.4

DHf;exp: (g, 298 K)a

All values are in kJ  mol1. a Experimental values are taken from reference [1,2].

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M. Namazian, M.L. Coote / J. Chem. Thermodynamics 40 (2008) 1627–1631 TABLE 2 G3(MP2)//B3LYP, G3(MP2) and CBS-QB3 results for 1-hydroxyisoquinolines (2) and 1,5-dihydroxyisoquinolines (4) G3(MP2)//B3LYP

G3(MP2)

G3//B3LYP

CBS-QB3

Species 2 Energy (0 K) Energy (298 K) Enthalpy (298 K) DHf (g, 0 K) DHf (g, 298 K) Deviation

1250888.0 1250867.1 1250864.6 0.8 24.4 (8.8)

1250833.9 1250813.0 1250810.6 4.5 19.9 (13.3)

1251958.1 1251937.2 1251934.7 7.3 17.9 (15.3)

1250678.1 1250657.7 1250655.2 11.7 13.1 (20.1)

Species 4 Energy (0 K) Energy (298 K) Enthalpy (298 K) DHf (g, 0 K) DHf (g, 298 K) Deviation

1448201.0 1448176.8 1448174.4 173.8 200.1 (24.2)

1448140.0 1448115.8 1448113.3 169.1 194.5 (29.8)

1449379.1 1449354.8 1449352.4 169.6 195.9 (28.4)

1447985.2 1447961.6 1447959.1 168.6 194.4 (29.9)

All values are in kJ  mol1. The experimental value of enthalpy of formation of 2 and 4, DHf;exp: (g, 298 K), has been reported to be (33.2 ± 2.9 and 224.3 ± 3.3) kJ  mol1, respectively [1,2]. The deviations of experimental values from the theoretical heats of formation are in parentheses.

TABLE 3 Calculated enthalpies of formation (kJ  mol1) of species 1 and 2 calculated at B3LYP/ 6-311+G(3df,2p) level of theory using various optimised geometries

DHf (g, 298 K)

Geometries

Species 1 B3LYP/6-31G(d) B3LYP/CBSB7a MP2(Full)/6-31G(d) QCISD/6-311+G(d,p)

244.2 243.5 244.5 246.0

Species 2 (1.8) (2.6) (1.5) (0.0)

34.1 33.2 34.7 35.7

(1.6) (2.5) (1.0) (0.0)

Values in parentheses are the deviation of each calculated enthalpy of formation from the benchmark level of theory. a CBSB7 basis set which has been used at CBS-QB3 level of theory stands for 6311G(2d,d,p).

that the B3LYP/6-311+G(3df,2p) single point energies are relatively unaffected by the level of theory used in the geometry optimization, and it is thus clear that this is not a source of error in the calculations. It therefore appears that the extraordinary discrepancies between theory and experiment for the enthalpies of formation of 2 and 4, may be due to experimental error. Therefore, we suggest the value of 18 ± 4 kJ  mol1 for heat of formation of 2 and 196 ± 4 kJ  mol1 for enthalpy of formation of 4 based on G3// B3LYP method [17]. Interestingly, as reported by Ribeiro da Silva et al. [1], considering the enthalpic increments for the introduction of the hydroxy group into positions 1 and 5 of the isoquinoline, the enthalpy of formation of 4 was estimated as 215.0 ± 1.4 kJ  mol1. This value is 10 kJ  mol1 closer to the theoretical value of 196 kJ  mol1 presented here than the experimental value and further supports the present theoretical calculations.

In order to investigate whether lower-cost computational methods provide an accurate alternative to G3//B3LYP, we have also calculated enthalpies of formation of species 1–8, using the popular B3LYP and MP2 methods with the large 6-311+G(3df,2p) basis set (table 4). As can be seen from table 4, the B3LYP/6311+G(3df,2p) enthalpies of formation deviate significantly from the G3//B3LYP results (the mean absolute deviation, MAD, is 40 kJ  mol1). For the ab initio method, MP2, the MAD is even larger at 65 kJ  mol1. It would thus appear that, even when a large triple zeta basis set is used, popular low-cost methods such as B3LYP and MP2 are not suitable for the calculation of accurate heats of formation; instead, high levels of theory, such as the composite methods studied in the present work, should be used. However, of these methods, the lower cost procedure, G3(MP2)//B3LYP does appear to reproduce the G3//B3LYP results with reasonable accuracy (the MAD is just 6 kJ  mol1). Natural bond orbital (NBO) [18] charge calculations at the B3LYP/6-311+G(3df,2p) level were performed in order to examine why 5-hydroxyisoquinoline, 3, is much less stable than the other mono-substituted isoquinoline, 1-hydroxyisoquinoline (2). figure 2 shows the charge distribution for the parent compound (1), those

-0.17

0.10

-0.21

N

-0.19

-0.46 0.03

-0.19

-0.22

FIGURE 2. Natural charges on the carbon and nitrogen atoms of isoquinoline.

TABLE 4 Heats of formation of studied species (1–8) using different levels of theory Entry

B3LYP/6-311+G(3df,2p)

1 2 3 4 5 6 7 8 MADa

244 38 76 134 147 61 239 28

MP2/6-311+G(3df,2p) (42) (52) (52) (62) (7) (15) (42) (51) 40

140 102 63 306 126 111 135 107

G3(MP2)//B3LYP (63) (85) (87) (110) (14) (36) (63) (84) 65

193 24 17 200 135 78 188 29

The deviations of calculated heats of formation from the G3//B3LYP values are in parentheses. All values are in kJ  mol1. a Mean absolute deviations of calculated heats of formation from G3//B3LYP results.

G3//B3LYP (9) (7) (7) (4) (5) (2) (9) (6) 6

202 18 24 196 140 75 198 23

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M. Namazian, M.L. Coote / J. Chem. Thermodynamics 40 (2008) 1627–1631

TABLE 5 G3(MP2) results for calculation of energies and enthalpies of studied isoquinolines and their protonated form Entry

1

Zero-point energy Energy (0 K) Energy (298 K) Enthalpy (298 K) PA (calc.) PA (exp.)

2

3

4

Neutral

Protonated

Neutral

Protonated

Neutral

Protonated

Neutral

Protonated

342.1 1053486.7 1053468.6 1053466.1 953.4 951.7

377.0 1054434.4 1054415.8 1054413.3

354.9 1250833.9 1250813.0 1250810.6 921.8

387.3 1251750.4 1251728.6 1251726.1

353.3 1250793.4 1250772.0 1250769.5 958.9

387.8 1251746.5 1251724.7 1251722.2

365.9 1448140.0 1448115.8 1448113.3 927.2

398.0 1449061.9 1449036.7 1449034.2

All values are in kJ  mol1.

TABLE 6 Calculated proton affinities of studied isoquinones (1–4) using different levels of theory Entry

B3LYP//6-311+G(3df,2p)

MP2//6-311+G(3df,2p)

G3(MP2)//B3LYP

1 2 3 4 MAD

965.5 934.1 971.5 940.1

938.9 904.6 945.0 909.3

952.9 921.6 958.2 926.5

(12.1) (12.3) (12.6) (12.9) 12.5

(14.5) (17.2) (13.9) (17.9) 15.9

G3//B3LYP B3LYP (0.5) (0.2) (0.7) (0.6) 0.5

954.1 922.0 959.2 927.0

G3MP2 (0.7) (0.2) (0.3) (0.2) 0.3

953.4 921.8 958.9 927.2

The deviations of calculated proton affinities from the G3(MP2) values are in parentheses. All values are in kJ  mol1.

for the other species are provided in the Supporting Information. As shown in figure 2, the natural charge on the nitrogen is negative, whilst those on the carbon atoms adjacent to the nitrogen atom are positive. One would thus expect that lone-pair donating substituents attached to those electron-deficient carbons could help to stabilize the molecule via resonance. In 2 (and also 4), the hydroxyl group can play this role and thereby stabilize the molecule compared with 1 and 3 (which have no hydroxyl group in either of these positions). The electron donating effect of the hydroxyl group in 2 and 4 is clearly evident in the increased negative charge on nitrogen in these two species, which increases from 0.46 in 1 and 0.45 in 3 to 0.55 in 2 and 0.54 in 4.

4. Proton affinity The proton affinity, PA, of a species M is the change of enthalpy (with the opposite sign) of the following gaseous reaction in which 1 mol of H+(g) is consumed to form one mole of species MH+(g) [19]:

MðgÞ þ Hþ ðgÞ ! MHþ ðgÞ PA ¼ DH

ð1Þ

Of the G3 methods considered in this study, G3(MP2) has been reported to perform better for the calculation of the proton affinity, having an average absolute error of 4.3 kJ  mol1 [14]. The corresponding reported errors were (5.6 and 5.1) kJ  mol1 for the G3 and G3//B3LYP methods, respectively [13]. Our previous calculation of proton affinity of dimethyl methylphosphonate using different G3 methods also showed that G3(MP2) presented a slightly better calculated PA compared with the experimental value [16]. Therefore, we have employed G3(MP2) for the calculation of the proton affinity of studied isoquinoline derivatives, and results are shown in table 5. We have used the value of 6.2 kJ  mol1 for enthalpy (H) of proton which has been achieved by using the available values of the entropy (S = 108.95 J  mol1  K1) [20] and Gibbs free energy of proton (G = 26.28 kJ  mol1) [21] and considering the usual, H = G + TS relation. This value is in a perfect agreement with the value of 5/2RT at 298.15 K based on equi-partition principle [16,19]. Using enthalpies of isoquinoline (1) and its protonated form, as calculated at G3(MP2) level of theory, the proton affinity of

isoquinoline has been calculated as 953.4 kJ  mol1. The experimental value for proton affinity of isoquinoline has been reported as 951.7 kJ  mol1. There is only a 1.7 kJ  mol1 difference between theory and experiment, which can be considered as an indication of validity of G3(MP2) enthalpies. Using the G3(MP2) method, the proton affinities of 2, 3, and 4 have been calculated as 921.8, 958.9, and 927.2 kJ  mol1, respectively. There are no available experimental values for these species for comparison; however, we expect an average error of 4.3 kJ  mol1 for the theoretical values, based on the performance of this level of theory over the G2/97 test set [14]. The proton affinities of both 2 and 4 are less than the proton affinity of isoquinoline, 1, further showing the stability of these species, in line with their negative heats of formation. In order to investigate the effect of level of theory on the accuracy of the calculated proton affinities of studied species, we have calculated proton affinity of species 1–4, using the B3LYP and MP2 methods with the 6-311+G(3df,2p) basis set and the results are compared our benchmark values (table 6). As can be seen from table 6, the B3LYP/6-311+G(3df,2p) proton affinities deviate significantly from the G3MP2 results (with an MAD of 12.5 kJ  mol1). The MP2/6-311+G(3df,2p) level of theory again results in even larger errors, having an MAD of 15.9 kJ  mol1. We therefore observe that the calculation of proton affinity also requires more accurate procedures than the popular methods of B3LYP and MP2, even when a large basis set such as 6-311+G(3df,2p) is used. Finally, the proton affinities were also calculated using the high-level composite methods G3(MP2)//B3LYP and G3//B3LYP for comparison with the G3MP2 results (see table 6). In those cases deviations from G3MP2 are less than 1.0 kJ  mol1, indicating all that all variants of G3 method are sufficiently accurate for these systems. 5. Conclusions The enthalpy of formation of four isoquinoline derivatives (isoquinoline 1, 1-hydroxyisoquinoline 2, 5-hydroxyisoquinoline 3, and 1,5-dihydroxyisoquinoline 4) as well as four pyridine and quinoline derivatives (pyridine 5, 2-hydroxypyridine 6, quinoline 7, and 2-hydroxyquinoline 8) have been calculated theoretically. In general there is excellent agreement between theory and experiment (within 4.3 kJ  mol1), except in the case of 1-hydroxyisoquinoline

M. Namazian, M.L. Coote / J. Chem. Thermodynamics 40 (2008) 1627–1631

(2) and 1,5-dihydroxyisoquinoline (4). Based on G3//B3LYP calculations, values of (196 and 18) kJ  mol1 have been suggested for the enthalpies of formation of these species, respectively. NBO calculations show that hydroxyl groups carbon adjacent to nitrogen help to stabilize the molecule via electron donation. The proton affinity of isoquinoline calculated using G3(MP2) method is in good agreement with the experimental value of 951.7 kJ  mol1 with a deviation of only 1.7 kJ  mol1. Proton affinities of the other isoquinoline derivatives have also been calculated and it is found that those of 2 and 4 are significantly less than that of the parent compound 1, further highlighting the stability of these species, in line with their negative heats of formation. Acknowledgements We gratefully acknowledge generous allocations of computing from the Australian Partnership for Advanced Computing and the Australian National University Supercomputing Facility, and financial support from the Australian Research Council under their Centres of Excellence program. Appendix A. Supplementary data B3LYP/6-31G(d) optimized geometries in the form of GAUSSIAN archive entries, corresponding total energies, and relevant NBO charges. Supplementary data associated with this article can be found, in the online version, at doi:10.1016/j.jct.2008.07.007. References [1] M.A.V. Ribeiro da Silva, M.A.R. Matos, L.M.P.F. Amaral, J. Chem. Thermodyn. 37 (2005) 1312–1317.

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JCT 08-131