Charge distribution in the carbonate ion

Notes

Department of Chemistry Imperial College London SW7 2A Y

P. J. MAYNE G. F. KIRKBRIGHT

REFERENCES 1. C.W. Sill and H. E. Peterson, Analyt. Chem. 21, 1266 (1949). 2. G. F. Kirkbright, C. G. Saw and T. S. West, Talanta 16, 65 (1969). 3. R. Bock and E. Zimmer, Z. analyt. Chem. 198, 170 (1%3). 4. G. F. Kirkbright, T. S. West and C. Woodward, Talanta 12, 517 (1%5). 5. H. Fromherz and Kun-Hu-Lih, Z phys. Chem. 153A, 321 (1931). 6. A. B. Scott and Kuo-Hao Hu, J. chem. Phys. 23, 1830 (1955). 7, R. C. Woodford, J. chem. Soc.(A), 651 (1970). 8. P. Brauer and D. Pelte, Z. Naturf. 17a, 875 (1962). 9. R. E. Curtice and A. B. Scott, lnorg. Chem. 3, 1383 (1%4). 10. R. O. Nilsson, Ark. Kemi. 10, 363 (1957). 11. D. Peschanski and S. Valladas-Dubois, Bull. Soc. chim. Ft. 1170 (1956).

Z

1529

12. R. Freeman, R. P. H. Gasser, R. E. Richards and D. H. Wheeler, Molec. Phys. 2, 75 (1959). 13. R. P. H. Gasser and R. E. Richards, Molec. Phys. 2, 357 (1959). 14. P. Debye, Trans. electrochem. Soc. 82, 265 (1942). 15. G. F. Kirkbright, P. J. Mayne and T. S. West, J. chem. Soc. Dalton 1918 (1972). 16. A. H. Mehler, B. Bloom and M. E. A. DeWitt Stetten, Science 126, 1285 (1957). 17. A. N. Fletcher, Photochem. Photobiol. 9, 439 (1%9). 18. J. G. Calvert and J. N. Pitts, Photochemistry, p. 800. Wiley, New York (1966). 19. J. N. Demas and G. A. Crosby, J. phys. Chem. 75, 1008, 1009 (1971). 20. C. A. Parker and W. T. Rees, Analyst 87, 83 (1%2). 21. G. Steffen and K. Sommermeyer, Biophys. 5, 192 (1968). 22. J. G. Calvert and J. N. Pitts, Photochemistry, p. 627. Wiley, New York (1%6). 23. E. A. Moelwyn-Hughes, The kinetics of reactions in solution, p. 179. Oxford University Press, Oxford (1933). 24. R. P. Bell and J. H. B. George, Trans. Faraday Soc. 49, 619 (1953).

inorg,nucl.Chem..1975,Vol.37. pp. 1529-1530. PergamonPress. Printedin GreatBritain

Charge distribution in the carbonate ion (Received 25 April 1974) AN ATI'EMPT has been made to determine the distribution of charge on the carbonate ion by both a cohesive energy study and theoretical calculations. A preliminary communication has been presented [l], and challenged by Jenkins and Waddington[2] on grounds which are questionable. These authors have reported[3] a value of -0"20 for the charge on oxygen in the carbonate ion, a value which, from simple considerations, is too small numerically. A model ion-gas (Ca ~+, C ~c, 0 % where zc + zo = - 2 ) has been used to define a reference level. It is satisfactory for the present calculations, as it has not been strictly related to any thermodynamic parameters. The usual model for the ion-gas (Ca 2÷, CO32-) differs from the above model by a quantity called[3] the electrostatic self-energy of the carbonate ion. This quantity represents the difference in energies of the process C zc + 3 0 ~o ~ C O 3 ~ in both the ion-gas and crystal situations. This difference is small, and permits an estimation of zo through the usual Madetung electrostatic energy term, A ( L ) / L , referred to the standard distance L. This term is related to z by a quadratic function, the nature of which has been demonstrated [4] both analytically and graphically. The calculations on calcite produced a value of Zo = - 0-77; with aragonite a value of -0.75 was obtained. These results reflect errors of about 2 per cent in the cohesive energy values for these two forms of CaCO3. In the calculations, a repulsion potential of the form exp ( - L / p ) was used. The values of p/L of 0.10 and 0-11 for calcite and aragonite respectively, obtained from compressibility data, were so close to the value of 0' 1 found so frequently for ionic crystals [5], that the model was regarded as satisfactory. The more complicated repulsion potential of Huggins [6] was not employed because it is not possible to make an unequivocal assignment of 'basic radii' for the compounds studied here, for reasons given in an earlier critical study[7]. The use of basic radii obtained from other studies will often give apparently satisfactory repulsion parameters. However, in predominantly ionic compounds, any repulsion potential which is equivalent to an inverse high power o f interatomic d i s t a n c e will give reasonable results. In ionic crystals containing polyatomic

ions, independent confirmation of a repulsion energy model is difficult to obtain. In constructing and interpreting a curve of the electrostatic energy A ( L ) / L as a function of zo for CaCO3, it is necessary to bear in mind the features illustrated by Table l, for aragonite. In this Table, (a) represents a point-charge model, in which the carbonate ion is represented as a point charge (-2) centred on the position of the carbon atom. Models (b)-(e) are different distributed-charge situations, subject to the condition that the sum of the charges on carbon and oxygen, IEz~l, is,

0.78

0.77

0"76

0"75

120

1.22

124

126

128

130

------ d (c-o),

Fig. 1. The variation of zowith d (C-O) in the CO3z ion.

1530

Notes Table 1. Calculations on aragonite.

(a) (b) (c) (d) (e)

zo

zo

[zc+3zol

-2.0 -1.1 -0.5 -0.4 +1.0

0.0 -0.3 -0.5 -0.8 -1.0

2.0 2-0 2.0 2.0 2.0

[zol+3lZol A ( L ) / L 2.0 2-0 2-0 2.8 4.0

2.107 1.298 1.301 2,119 3.207

numerically, 2.0. Situations like (b) and (c) lead to values of A ( L ) / L much lower than that for the point charge model, whereas (d) and (e) are higher than (a), (b) or (c), This effect arises mainly from a Y.z~2 term in the Madelung constant calculation. The condition that IXz, I = 2 is, therefore, a necessary but not sufficient criterion of a feasible model; the individual values of z, must be inspected carefully for reasonableness. The value of Zo obtained[l] by the CNDO/2 method[8] was -0.77, in good agreement with the results from the energetics study. The variation of zo with C - O distance in the carbonate ion is shown in Fig. 1. It is to be expected that the value of zo will be different in the condensed state. An attempt has been made to extend the quantum mechanical calculation to take account of the nearest-neighbour interactions. In order to

keep the basis set small, it was necessary to work with MgCOn instead of CaCOn. The result obtained for Zo is -0.80. Bearing in mind the limitations of the procedures used, the overall agreement in Zo seems s a t i s f a c t o r y . .

Department of C_hemical Physics University of Surrey Guildyord Surrey

M.F.C.

LADD

REFERENCES 1. M. F. C. Ladd, Nature, Lond. (phys. Sci.), 238, 125 (1972). 2. H. D. B. Jenkins and T. C. Waddington, Nature, Lond. (phys. Sci.) 238, 126 (1972).

3. H. D. B. Jenkins and T. C. Waddington, Nature, Lond, (phys. Sci.) 232, 5 (1971). 4. M. F. C. Ladd, Theoret. Chim. Acta 25, 400 (1972). 5. M. F. C. Ladd and W. H. Lee, in Progress in Solid State Chemistry, (Edited by H. Reiss) Vol. 1. Pergamon Press, Oxford (1964). 6. M. L. Huggins, J. chem. Phys. 5, 143 (1937). 7. M. F. C. Ladd and W. H. Lee, Jr. inorg, nucl. Chem. 11, 264 (1959). 8. J. A. Pople and D. L. Beveridge, Approximate Molecular Orbital Theory. McGraw-Hill, N e w York (197o).

£ inorg,nucl.Chem.,1975,Vol.37, pp. 1530--1532. PergamonPress. Printedin Grea!Britain P a l l a d i u m ( l l ) c o m p l e x e s of Schiff bases

(First received l April 1974; infinalform 24 June 1974) SCI-nFFbase complexes of palladium(II) and other members of the second and third row transition metal ions have not been widely studied[I-3]. Recently[4] we have reported the preparation and properties of some palladium(II) chelates of dibasic quadridentate Schiff bases. We have further isolated some palladium(II) complexes by reacting bis(salicylaldehydato) palladium(II) (=Pd(salh) with different amines, aminoalcohols, and aminoacids in rettuxing chloroform. This note describes the preparation and properties of these newly isolated palladium(II) chelates. While our work was in progress Patel and Bailar[5] described the preparation of a palladium(II) complex of N,N'-bis(salicylidene)1,1-(dimetbyl) ethylenediamine (= saln-i-bn) by reacting Pd(salh with isobutylenediamine in chloroform. It has also been observed in the present investigation that the reaction of tetrachloropalladate(II) with the appropriate Schiff base in dry ethanol and in the presence of a suitable base can yield the corresponding palladium(II) Schiff base chelate. This is a slightly different method than the one described recently by Yamada and Yamanouchi[6] for the synthesis of palladium(II) complexes of N-substituted salicylideneimines, The following abbreviations have been used for the Schiff base ligands used in the present study: (i) BSTNOL--H2 = N,N'-(2-hydroxytrimethylene) bis(salicylideneimine), (ii) BSTN-H2 = N,N'-trimethylenebis(salicylideneimine), (iii) BSOP-H2 = N'N'-°rth°phenylenebis(salicylideneimine)' (iv) SOAP-H2 = N-(2-hydroxyphenyl) salicylideneimine, (v) SAA-Hz = N-(2-carboxyphenyl) salicylideneimine, (vi) SETOL-H2 = N-(2 hydroxyethyl) salicylideneimine, and (vii) SAN-H = N-phenylsalicylideneimine.

EXPERIMENTAL Elemental analyses were done by conventional methods. Magnetic susceptibilities were measured in a Gouy balance at 24.5°C. Electronic absorption spectra were recorded in a Beckman DU-2 spectrophotometer. I.R. spectra were recorded

(KBr phase) by C.D.R.I., Lucknow. Reagents and solvents were purified by the usual techniques.

Preparation of palladium(H) complexes of Schiff bases These chelates were prepared by the following general method: Equimolar quantities of amines, aminoalcohols or aminoacids and Pd(salh in chloroform were refluxed on the waterbath for one hour and filtered while hot. The colored filtrate was concentrated, and ultimately yielded colored crystalline compounds. In many cases addition of ether was necessary to get the crystals. The compounds were filtered, washed with an ether--chloroform mixture and finally with chloroform and dried in a vacuum desiccator. RESULTS AND DISCUSSION The palladium(II) Schiff base complexes along with their elemental analyses and other physical data are given in Table 1. The complexes are all found to be diamagnetic at room temperature, which supports square-planar, four coordinate configurations for them. It is well known that the palladium(II) ion having a d 8 configuration, favours the formation of complexes with square-planar geometry [2]. The electronic spectra of the present palladium(II) Schiff base complexes, as recorded in Table 2, are quite similar to each other, showing that their configurations are all alike. The absorption peak at about 2 5 , 0 0 0 c m - t may be due to an electronic transition within the ligand bound to the palladium(II) ion. The ligand field bands, mainly due to palladium(II) ion, are considered to be hidden by the strong absorption bands of other origins. It is to be noted that Beers' law holds for these palladium(II) chelates, indicating that no association of the square-planar complexes occurs in solution. This finding is in agreement with the current view that the tendency of palladium(II) to take coordination numbers larger than four is extremely low. This is further substantiated by the fact that the spectra of all these palladium(II) complexes in a donor solvent like pyridine are nearly the same as those of the same complexes in n0d-donor solvents, This is not the case with