Ionic thermodynamic quantities

The individual ionic thermodynamic quantities for hydration, listed in Table 4.1, are based on two extra-thermodynamic assumptions for the enthalpies, they are based on the tetraphenylphosphonium tetraphenylborate (TPTB), A // (Ph4P ) = A //" BPh and for the entropies on the temperature derivative of the electromotive force of... 

The small difference between the successive pK values (cf. values below) of tungstic acid was previously explained in terms of an anomalously high value for the first protonation constant, assumed to be effected by an increase in the coordination number of tungsten in the first protonation step (2, 3, 55). As shown by the values of the thermodynamic parameters for the protonation of molybdate it is actually the second protonation constant which has an abnormally high value (54, 58). An equilibrium constant and thermodynamic quantities calculated for the first protonation of [WO, - pertaining to 25°C and zero ionic strength (based on measurements from 95° to 300°C), namely log K = 3.62 0.53, AH = 6 13 kJ/mol, and AS = 90 33 J, are also consistent with a normal first protonation (131) (cf. values for molydate, Table V). 

In the system Th(IV)-H20 three sets of thermodynamic quantities can be derived from experimental data (1) the hydrolysis constants l°gio j3° and log10 /34 of ThOH3+ and Th(OH)4(aq), respectively, have been determined potentiometrically by several authors over a wide range of ionic strength (for references see Hummel et al. 2002) (2) the thermodynamic properties of Th02(cr) have been determined by calorimetry, and thus a solubility product logio K°s 0 (cr) for... 

The ionic bond is often described as the metal wants to lose an electron and the non-metal wants to accept an electron, so the two react with each other. Criticize this statement quantitatively using appropriate thermodynamic quantities. 

Nucleic acids References Ionic conditions Thermodynamic quantities... 

This principle serves as the basis for a number of models of fused salt systems. Perhaps the best known of these is the Temkin model, which uses the properties of an ordered lattice to predict thermodynamic quantities for the liquid state [79]. However, certain other models that have been less successful in making quantitative predictions for fused salts may be of interest for their conceptual value in understanding room temperature ionic liquids. The interested reader can find a discussion of the early application of these models in a review by Bloom and Bockris [71], though we caution that with the exception of hole theory (discussed in Section II.C) these models are not currently in widespread use. The development of a general theoretical model accurately describing the full range of phenomena associated with molten salts remains a challenge for the field. 

To clarify the different roles played by the Helmholtz and Gibbs free energies of ionic solutions, it is relevant to reconsider the derivation of these thermodynamic quantities in the original Debye-Hiickel theory [1—4],... 

Alot of information about the free energies of transfer of single ions between pure solvents has been accumulated. Less numerous are determinations in mixed solvents, and the ionic enthalpies of transfer and entropies of transfer as function of mole fraction are known as an exception only. In Table 1 ions and solvent mixtures are listed for which free energies of transfer and some other thermodynamic quantities have been determined. 

Considerable attention has been devoted to the nature of the solvent effects (as determined in water and in various mixed solvents) on the ionic dissociations (and related thermodynamic quantities) and other acid-base properties of aliphatic zwitterionic compounds. Such investigations include studies of tricine in 50 mass % methanol-water (1), Bes in pure water and in 50 mass % methanol-water 2,3), glycine in 50 mass % monoglyme-water (4), and glycine in pure water and in 50 mass % methanol-water (5,6, 7). The numerous factors (8,9,10) which... 

Quite apart from its theoretical calculation, by the use of one of the expressions developed above, it is possible to relate the lattice energy of an ionic crystal to various measurable thermodynamic quantities by means of a simple Hess s law cycle. This cycle was first proposed and used by Bom 15) and represented in its familiar graphical form by Haber (45). It is now usually referred to as the Born-Haber cycle. The cycle is given below for a uni-univalent salt in terms of enthalpies. 

The standard heats of formation of ions, as recorded in Table XL, may be combined with the free energies of formation and the entropies, in the same table, to calculate thermodynamic quantities for a variety of ionic reactions. Further, these data may be utilized in conjunction with those given in Tables V, XV, XIX and XXIV to determine standard free energy, entropy and heat content changes for reactions involving both ions and neutral molecules. 

The electron in H O becomes fully hydrated in ps time to become a discrete chemical species with a known charge (—1), ionic conductivity (190 cm 0 and diffusion coefficient (4.9 X 10 cm s )- From estimates of its thermodynamic quantities, the standard redox potential of [e ] is ca. —2.87 V, making it a powerful reducing agent. Because of its intense and broad optical absorption spectrum, (A = 710 nm ma. = cm ) extending from the UV into the ir and its relatively long... 

To summarize, for ionic solutions, one can either adopt the view that the system is a two-component mixture and use the KB theory for a two non-ionic species, say water and salt [as has been done by Friedman and Ramanathan (1970) and by Chitra and Smith (2002)], where the system is open with respect to the water and the salt (as a single entity). In this view, there is no place for ion-ion correlations hence the electro-neutrality condition is irrelevant. In this view only correlations between neutral molecules are meaningful. Or we can view the system as water, cation, and anion - but in this case one must open the system with respect to each of the species individually (hence allow also charged fluctuations in the system). In this case, the KBI are meaningful, but the thermodynamic quantities such as V, or fi/(,/6Nj are not available. In the second view, the ion-ion correlations do enter into the KB theory, but again the electro-neutrality condition is irrelevant. 

In this section we turn to the question of evaluating the pertinent thermodynamic quantities from experimental data. We discuss in this section the case of a solvaton s which does not undergo dissociation in any of the phases. The more complicated case of dissociated solvatons (such as ionic solutes) will be discussed separately in section 7.9. 

If the equilibrium constant depends on ionic strength then so do the other thermodynamic quantities relating to the equilibrium process such as AH , AG , A5 , AV and AC . Hence these must be calculated from the ideal K, or, if found independently of K, they must be extrapolated to zero ionic strength. 

Modem digital voltmeters allow the measurement of highly accurate emfs. If the emfs are measured over a range of ionic strengths this will allow the effect of ionic strength on the equilibrium constants to be studied in detail. Furthermore if the dependence of emf on temperature is measured, the thermodynamic quantities. A//, A5 and AC can be found and interpretations of the magnitudes made. If the experimental set-up can be adapted to a determination of the emfs at various pressures to be made, then AV values can also be found. 

The reported thermodynamic quantities seem to have been re-determined in each study, since the three papers report somewhat different values for the parameters referring to identical conditions. To avoid the over-representation of the data from this laboratory during our SIT analysis, we used an average value for each ionic strength studied. 

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