Enthalpies of complexation

The enthalpies of complexation of 3.8c to the copper(lf) - amino acid ligand complexes have been calculated from the values of at 20 C, 25 1C, 30 1C, 40 1C and 50 1C using the van t Hoff equation. Complexation entropies have been calculated from the corresponding Gibbs energies and enhalpies. 

Fig. 2. Components of Li enthalpies of complexation with methylamines. Successive steps indicate the effect on energy of interaction between Li and the amine of inclusion of additional components of the binding energy. The diagram shows that the permanent dipoles on amines (the charge on the nitrogen of the isolated amine) favor ammonia over trimethylamine complexation, but that polarizability and inductive effects (shift of negative charge onto the nitrogen in the complex) cause a massive turnaround in favor of complexation with trimethylamine rather than ammonia. Of particular importance is the near inversion of order caused by the addition of repulsive van der Waals terms. Modified after Ref. (9).

Fig. 4. Plot of enthalpies of complex formation in the gas phase of CpNi complexes (Cp = cyclopentadienyl) vs those for the corresponding Mn complexes. The ligands are segregated into correlations for soft (O) (N and S donors) and hard (oxygen donors) ( ). Energies are in kcal mol-1. Redrawn after Ref. (14). 

The EA/CA ratio was proposed as a measure of hardness of the Lewis acid, and EB/CB as hardness of the Lewis base in aqueous solution (17). It now seems that the E/C ratio is not a measure of hardness in the sense in which Pearson (5,5a) defined hardness. Rather, the E/C ratio for a Lewis acid or base is a measure of the tendency to ionicity in the M-L bonds formed. The EAICA ratio should rather be called IA, and the EbICb ratio IB, the tendency to ionic bonding in forming the M-L bond. Acids and bases in Tables I and II are placed in order of increasing tendency towards ionicity in the M-L bond, according to the E/C ratios IA and 7b. A justification for this interpretation is that the order of IA values for metal ions in aqueous solution strongly resembles the order of hardness derived by Pearson (19) from enthalpies of complex forma-... 

Calorimetric studies indicate that the enthalpies of complexation tend to show related trends to the observed stability constants and display selectivity peaks, although there is not necessarily a coincidence between the two sets of peaks. Complexation is characterized by the entropy becoming progressively less positive (less favourable) as the cation size decreases. This is illustrated in Figure 6.5 for the complexation of 2.2.1 with the alkali metals. 

Information on the thermodynamic properties (complexation constants, enthalpies of complexation, Gibbs energy of formation, and their relationships with structural and spectroscopic parameters) can be found in refs. 12, 23, and 24. 

The second-order rate-constants kp and kA for polymerisations in solution which we consider reliable are summarised in Table 12. The initiators used by the various investigators have not been listed, because by definition kp and Ep must be independent of these and there are insufficient data to permit any firm conclusions about the effects of the nature of the anion on and E. When considering the rate-constants in this Table it must be remembered that all of them, except those for isobutene, probably comprise a contribution from the polymer-complexed cation, p+p, greater or smaller according to circumstances (see Section 2.3), and correspondingly the activation energies would contain a term Ep+P and an enthalpy of complexation further, for the reason explained in Section 4.1.9, the kp from ionising radiation experiments are minimum values. 

Such equilibria are known to consist of a number of consecutive complex equilibria. The formation constant / represents the free enthalpy of complex formation in the gas phase. This quantity can not be determined by experiment. 

The enthalpy of complexation can be measured directly by reacting the metal and ligand in a calorimeter. It can also be determined indirectly by measuring log at different temperatures and applying the equation... 

The molar enthalpies of complexation of pyridine with iodine are - 36 kJ mol-1 (CC14) and - 34 kJ mol 1 (C6H12) by calorimetry, with corresponding values of - 33.1 and - 34.6, respectively, using the UV absorbance of the pyridine/iodine complex (84JCS(P2)731, 87JCS(P2)1713). For 2,6-dimethylpyr-idine the complex formation constant is reduced considerably from that for pyridine itself (from 106 to 46 in cyclohexane), but the enthalpy of formation remains sensibly constant, and thus independent of steric considerations. 

Pyridines are also well known as ligands in transition metal complexes, and if the equilibrium constants for the formation of such complexes can be related to base strength, it is expected that such constants would follow the Hammett equation. The problem has been reviewed,140 and a parameter S, formulated which is a measure of the contribution of the additional stabilization produced by bond formation to the stabilization constants of complexes expressed in terms of a.141 The Hammett equation has also been applied to pyridine 1 1 complexation with Zn(II), Cd(II), and Hg(II) a,/3,y,<5-tetraphenylporphins,142 143 the a values being taken as measures of cation polarizing ability. Variation of the enthalpy of complexation for adducts of bis(2,4-pentanediono)-Cu(II) with pyridines plotted against a, however, exhibited a curved relationship.144... 

Using equation (4), Torocheshnikov et al. (50) have calculated the values of <5C and K for Me3SnCl and Et3SnCl in acetone, acetonitrile, and dioxan. The molar enthalpy of complex formation, AH, can be obtained from a In K vs. l/T plot (Fig. 5) using the Van t Hoff relationship ... 

These data provide a useful mean of checking the reliability of the thermodynamic parameters associated with the solution and complexation processes. Indeed, coordination data should be the same (within the experimental error) independently of the solvent used in the solution process. This statement is now corroborated by representative examples involving the fluoride anion and 1. Thus in fulfilling the requirements of equation (6), the enthalpies of solution of the reactants (Bu4NF and 1) and product (Bu4N1F) are combined with the enthalpies of complexation of the fluoride anion and 1 in acetonitrile (MeCN) (equation (7)) and N,N- dimethylformamide (DMF) (equation (8)) to derive the enthalpy of coordination for this system in the solid state. 

Similar results have been obtained [23] in determining spectroscopic equilibrium constants and enthalpy of complexing 1,3,5-trinitrobenzene (as an acceptor) with aromatic hydrocarbons (as electron donors). Based on the results, it has been also inferred that donor force of alkyl-benzenes increases with increasing number of methyl substituents. 

For review, see also MacDiarmid (269). S(29Si) and 5(31P), Chemical shift in ppm AH, enthalpy of complex formation inkcal/mol A5, entropy of complex formation in cal/mol degree Py, pyridine NMI, [Pg.283]

For the quantitative evaluation of the inductive effect of R3 and R4, the pKBH+ values of the conjugate acid of the enaminone, or the enthalpy of complex formation, AHBFi, of boron trifluoride with the carbonyl oxygen for different NR3R4 groups were proposed. The variation of <5H(2) was considered to be a measure of the steric interaction. The following correlations were found ... 

Table 2-4. Molar enthalpies of complex formation between boron trifluoride and several non-HBD solvents, determined in dichloromethane at 25 °C, according to Eq. (2-lOa) [211, 212].

A selection of A//p gp values has already been given in Table 2-4 in Section 2.2.6. This new Lewis basicity scale is more comprehensive and seems to be more reliable than the donor number scale. Analogously, a Lewis basicity scale for 88 carbonyl compounds (esters, carbonates, aldehydes, ketones, amides, ureas, carbamates) has been derived from their standard molar enthalpies of complexation with gaseous boron trifluoride in dichloromethane solution [143]. The corresponding Aff Q gp values range from 33 kJ mol for di-t-butyl ketone to 135 kJ mol for 3-diethylamino-5,5-dimethyl-cyclohexen-2-one. 

Contrary to the classical approach the recently applied functional anion concept considers the substitution on H-sites by anions. First results indicate that this approach again is very promising for altering reaction enthalpies of complex hydrides. 

Two typical sets of the calorimetric data are presented in Table IV. The enthalpies of complexation were calculated as described previously (6). The heat of dilution of the lanthanide perchlorate was checked by titrating two of the lanthanides (/ = IM, pH = 3.0) with NaC104 (jn = IM, pH = 3.0). No heat was observed. The heat of dilution of the sodium fluoride was checked by titrating IM NaC104 (pH = 3.0) with NaF fx = IM, pH = 3.0). A small heat of dilution was observed for the first two points in the titration after this it was negligible. The heats of dilution of the sodium fluoride were accounted for in the calculations. 

For solute HBD acidity different scales were proposed mainly based on complexa-tion constants and enthalpies of complexation. TTie most important are reported here. 

Terdentate 4-amino-2,6-bis(pyridin-2-yl)-l,3,5-triazine (ADPTZ) can coordinate to Am and Ln with the formation of the 1 1 complexes. In the work [33] the thermodynamic characteristics for the complex formation of Am with ADPTZ were calculated AG=-32.9 0.6 kJ/mole, A//=-28.9 3.0 kJ/mole, A5 =14.0 10.0 J/Kmole. The stability constant equal to logy0 =5.8 O.l for [Am(ADPTZ)] complex was defined using spectrophotometry. The thermodynamic data show that the observed selectivity of the ligand arises from a difference in the enthalpies of complexation for Am and lanthanides. The geometry and electronic structure of the [M(ADPTZ)(H20)6] complexes (for M = Am, Cm, Pu) were calculated using DFT theory with no symmetry constraint. Selected distances are presented in Table 4. 

Analysis of enthalpies of complex formation between functional monomer and template... 

TABLE 18. Enthalpies of complexation (kJ mol ) of 4-fluorophenol with various bases measured by the pure base method ... 

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