Hydration enthalpies, lanthanide ions

It is quite difficult to measure an accurate enthalpy of solution A//( olutioni with a calorimeter, but we can measure it indirectly. Consider the example of sodium chloride, NaCl. The ions in solid NaCl are held together in a tight array by strong ionic bonds. While dissolving in water, the ionic bonds holding the constituent ions of Na+ and Cl- in place break, and new bonds form between the ions and molecules of water to yield hydrated species. Most simple ions are surrounded with six water molecules, like the [Na(H20)6]+ ion (VI). Exceptions include the proton with four water molecules (see p. 235) and lanthanide ions with eight. 

Patterns in Hydration Energies (Enthalpies) for the Lanthanide Ions... 

Table 2.4 shows the hydration energies (enthalpies) for all the 3+ lanthanide ions, and also values for the stablest ions in other oxidation states. Hydration energies fall into a pattern Ln + > Ln + > Ln +, which can simply be explained on the basis of electrostatic attraction. 

Table 2.4 Enthalpies of hydration of the lanthanide ions (values given as — AH hydr/kj mol ...

When is it most likely that Ln +(aq) ions will be stable For the first of the two reactions above to be favoured, the single factor that will help make A// positive is a high value of h. Less important would be the size of the ions, as this could affect the hydration enthalpies the difference between the hydration enthalpies will be less, the larger the lanthanide ions. Substituting into the above equation, we can investigate the relative stabilities of La +(aq) and Eu +(aq), making use of ionization energies from... 

Figure 5.2 The variation in hydration enthalpy of Ln ions across the lanthanide series...

In these studies the enthalpies and entropies of lanthanide complex formation had been measured and the enthalpies and entropies of hydration estimated from theoretical and semiempirical equations. Bertha and Choppin (44) examined the hydration thermodynamics of the lanthanide ions directly and obtained molar entropies of hydration of 338 4 J/mol-K and 401 4 J/mol-Kfor La-Pr and for Dy-Lu, respectively, and a range of 351-392 J/mol-K for Nd-Tb. These observations supported die existence of two dffferendy sized hydration spheres corresponding to the La-Pr and Dy-Lu groups, with Nd-Tb comprising the transition between them. Spedding et al (45) later confirmed die AShyd data of Bertha and Choppin. 

Fig. 5. Hydration enthalpies of lanthanide and actinide ions. Data from David (1986b).

Hydration enthalpies have also been derived from a Born-Haber cycle (e.g., Morss 1971) for ions and An if all other terms are known. Recently, this approach has been updated (Schoebrechts et al. 1989) with new lanthanide ionization potentials and has been extended to the actinide ions using experimental results for the... 

A similar trend is seen for the enthalpy of formation of Ln +(aq) which shows more or less constant values for all Ln + ions except Eu and Yb. The trend as a function of lanthanide is shown in Figure 14, which contrasts with the rather smoother trend seen in the enthalpy of hydration of the ions seen in Figure 10. 

Figure 30.3 Variation with atomic number of some properties of La and the lanthanides A, the third ionization energy (fa) B, the sum of the first three ionization energies ( /) C, the enthalpy of hydration of the gaseous trivalent ions (—A/Zhyd)- The irregular variations in I3 and /, which refer to redox processes, should be contrasted with the smooth variation in A/Zhyd, for which the 4f configuration of Ln is unaltered.

Hproblems associated with all the trihalides of this review of the presence of small amounts of hydrates or oxochlorides. While on the matter of possible impurities, it may be recalled that in Bommer and Hohmann s early work there is a discrepancy between enthalpies of solution of anhydrous trichlorides and of respective metals in hydrochloric acid. Here the more likely impurity to be responsible is unreacted potassium metal in the lanthanide metal used in the hydrochloric acid dissolution experiments. 

The enthalpies of solution and solubilities reviewed here provide much of the experimental information required in the derivation of single-ion hydration and solvation enthalpies, Gibbs free energies, and entropies for scandium, yttrium, and lanthanide 3+ cations. 

The enthalpies of hydration of the Ln ions were calculated and their relationship to the lanthanide contraction was discussed. 

An electrostatic hydration model has been applied to the trivalent lanthanide and actinide ions in order to predict the standard free energy (AG°) and enthalpy (AHt) of hydration for these series. Assuming crystallographic and gas-phase radii for Bk(III) to be 0.096 and 0.1534 nm, respectively, and using 6.1 as the primary hydration number, AG298 was calculated to be -3357 kJ/mol, and A/Z298 was calculated to be -3503 kJ/mol (187). 

An electrostatic hydration model, previously developed for ions of the noble gas structure, has been applied to the tervalent lanthanide and actinide ions. For lanthanides the application of a single primary hydration number resulted in a satisfactory fit of the model to the experimentally determined free energy and enthalpy data. The atomization enthalpies of lanthanide trihalide molecules have been calculated in terms of a covalent model of a polarized ion. Comparison with values obtained from a hard sphere modeP showed that a satisfactory description of the bonding in these molecules must ultimately be formulated from the covalent perspective. 

Semiempirical calculations of free energies and enthalpies of hydration derived from an electrostatic model of ions with a noble gas structure have been applied to the ter-valent actinide ions. A primary hydration number for the actinides was determined by correlating the experimental enthalpy data for plutonium(iii) with the model. The thermodynamic data for actinide metals and their oxides from thorium to curium has been assessed. The thermodynamic data for the substoicheiometric dioxides at high temperatures has been used to consider the relative stabilities of valence states lower than four and subsequently examine the stability requirements for the sesquioxides and monoxides. Sequential thermodynamic trends in the gaseous metals, monoxides, and dioxides were examined and compared with those of the lanthanides. A study of the rates of actinide oxidation-reduction reactions showed that, contrary to previous reports, the Marcus equation ... 

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