Enthalpy aqueous species

References (20, 22, 23, 24, 29, and 74) comprise the series of Technical Notes 270 from the Chemical Thermodynamics Data Center at the National Bureau of Standards. These give selected values of enthalpies and Gibbs energies of formation and of entropies and heat capacities of pure compounds and of aqueous species in their standard states at 25 °C. They include all inorganic compounds of one and two carbon atoms per molecule. 

Much the same information can be displayed on an enthalpy diagram, in which we plot the enthalpies of formation (relative to the proton) of aqueous species instead of free energies. This facilitates the inclusion on the diagram of unknown and presumably unstable species, making reasonable assumptions concerning their hydration enthalpies. As an example, Fig. 5.5 presents enthalpy diagrams for M"+(aq) (M = Al, Tl) at a pH of zero. The enthalpy changes are AH° values for the processes ... 

TABLE A13.1 Gibbs free energy and enthalpy of formation from the elements at 25 C and 1 bar total pressure of some uranium aqueous species and solids of geochemical interest... 

TABLE A13.8 Gibbs free energies and enthalpies of formation from the elements of some geochemically important aqueous species and solids of plutonium at 25T and 1 bar total pressure... 

Just as we could not determine the absolute value of G ", we also cannot measure H/". As with G/" we circumvent this problem by assigningH/" a value of zero to all elements in their most stable form at 25"C and 1 atm pressure. In aqueous solution 1 mole/liter of the hydrogen ion, H", in ideal solution (y = 1) also is assigned an H° value of zero. We can determine values of enthalpy of specie based on these assignments and call t se the enthalpy of formation, A/. Similarly to the computations for AG/" values, we can compute the AH/° values of various compounds from th assigned H values of their component elements. A selection of these AH/° values is given in Table 3-1. 

Every enthalpy calculation requires a basis. For purposes of this development, we will take, for aqueous species, the pure component heat of formation at 25°C and infinite dilution as the basis. Electrolyte solutions, as we have seen, contain either molecular or ionic species. Thus, in order to compute H (T) we will need formulations for both molecular and ionic species. 

These thermochemical tables consist of enthalpies of formation at 298.15 K calculated from thermochemica data networks. Included are appropriate references to the literature and estimated errors in the enthalpies of formation. An interesting and Important feature of this scheme is that the tables can be readily updated by computer. These tables contain a substantial amount of data for aqueous species. The following tables have been published ... 

The thermodynamic values determined for the phases and species of iron(III) are listed in Table 11.31. In general, there is excellent agreement with the Gibbs energy data derived in the present review with that available in the literature (Ziemniak, Jones and Combs, 1995 Lemire et al., 2013). In all cases, the data derived herein and the respective uncertainty limits overlap the respective data available in the literature. There is also excellent agreement between the enthalpy of formation for solid phases derived in the present review with those available in the literature. There is poorer agreement with the enthalpy of formation data of aqueous species, typically due to differences in the enthalpy of reaction... 

Fluoroacetic acid [144-49-OJ, FCH2COOH, is noted for its high, toxicity to animals, including humans. It is sold in the form of its sodium salt as a rodenticide and general mammalian pest control agent. The acid has mp, 33°C bp, 165°C heat of combustion, —715.8 kJ/mol( —171.08 kcal/mol) (1) enthalpy of vaporization, 83.89 kJ /mol (20.05 kcal/mol) (2). Some thermodynamic and transport properties of its aqueous solutions have been pubHshed (3), as has the molecular stmcture of the acid as deterrnined by microwave spectroscopy (4). Although first prepared in 1896 (5), its unusual toxicity was not pubhshed until 50 years later (6). The acid is the toxic constituent of a South African plant Dichapetalum i mosum better known as gifirlaar (7). At least 24 other poisonous plant species are known to contain it (8). 

The solvophobic model of Hquid-phase nonideaHty takes into account solute—solvent interactions on the molecular level. In this view, all dissolved molecules expose microsurface area to the surrounding solvent and are acted on by the so-called solvophobic forces (41). These forces, which involve both enthalpy and entropy effects, are described generally by a branch of solution thermodynamics known as solvophobic theory. This general solution interaction approach takes into account the effect of the solvent on partitioning by considering two hypothetical steps. Eirst, cavities in the solvent must be created to contain the partitioned species. Second, the partitioned species is placed in the cavities, where interactions can occur with the surrounding solvent. The idea of solvophobic forces has been used to estimate such diverse physical properties as absorbabiHty, Henry s constant, and aqueous solubiHty (41—44). A principal drawback is calculational complexity and difficulty of finding values for the model input parameters. 

In recent years, aqueous solutions of Xe03 have been used to oxidize a species in solution, from which A[H°m can be calculated when AH for the oxidation reaction is combined with AH for other reactions. The noble gas oxide Xe03 is used as an oxidant because of its stability and the fact that the final reaction product is Xe(g), which has a zero enthalpy of formation and is easily removed from the reaction mixture. As an example, O Hare4 has reported AfHcm for UCI4. We will not go through the details of his procedure, but the critical step involved measuring A TH for the reaction... 

If a substance is to be dissolved, its ions or molecules must first move apart and then force their way between the solvent molecules which interact with the solute particles. If an ionic crystal is dissolved, electrostatic interaction forces must be overcome between the ions. The higher the dielectric constant of the solvent, the more effective this process is. The solvent-solute interaction is termed ion solvation (ion hydration in aqueous solutions). The importance of this phenomenon follows from comparison of the energy changes accompanying solvation of ions and uncharged molecules for monovalent ions, the enthalpy of hydration is about 400 kJ mol-1, and equals about 12 kJ mol-1 for simple non-polar species such as argon or methane. 

Last but not least of the liquid calorimetric media are aqueous solutions used in the hydrolysis of simple and complex fluorides. Stepwise replacement of F by OH occurs, and mixed products are not unusual. Thus the BFj ion hydrolyzes to species BF (OH)l and one has to ensure that the same product composition is formed in the auxiliary heat experiments (99). The problem is accentuated when polynuclear species form, as the equilibration can be slow. The inconsistencies in the heats of alkaline hydrolysis of MoF6 and WFe found by various authors and of the enthalpy of SbF5—derived by assuming SbF5 and Sb205 dissolved in 10 M HF produced the same species in solution—illustrate the difficulties. It is as well to confirm enthalpies of higher valent fluorides obtained by hydrolysis by alternative nonaqueous methods, especially since uncertainty in the Afl (Fderived enthalpy. The advantage of hydrolysis methods, apart from the simplicity of technique, is that the heats are small and one can tolerate... 

We start with butane-2,3-dione dioxime, more commonly known as dimethylglyoxime (dmg). It is a classic reagent for the analysis of NP, the green aqueous solution of metal ions transforming into a vibrantly red precipitate of Ni(dmg)2 complex it is one of the stars of the show in Ponikvar and Liebman s analytical chemistry chapter in the current volume. Here the stereochemistry is well-established and well-known—both OH groups are found on the same side as their adjacent CH3 group on the butanedione backbone. There have been several measurements of the enthalpy of formation of this species for which we take the one associated with this inorganic analytical chemistry application, i.e. with diverse metal complexes and chelates . 

An indication of the degree of exothermicity of sulphide oxidation reactions can be gained by comparing the enthalpy of formation (A//f), that is, a measure of the energy locked up in each chemical species, relative to native elements. The difference in enthalpies of formation of all reactants and all products defines the enthalpy (heat released or absorbed) of the reaction. Thermodynamic data on sulphide minerals, such as pyrite, are notoriously varied and disputed, and the values in Table 4 must be treated with caution. Nevertheless, depending on whether one defines the reaction as ending in an aqueous solution (equation 5), an intermediate secondary sulphate (e.g., melanterite - equation 6) or in complete oxidation to an oxyhydroxide (equation 7), the calculated reaction enthalpy (AH°) released is of the order of at least 1000 kJ/mol. 

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