Standard solution enthalpy

If A is a gas at ambient temperature and pressure, AsoiVH° can be determined experimentally by calorimetric methods or from measurements of the solubility change with temperature [42-44], When A is a liquid or a solid, its solvation enthalpy in a given solvent is usually calculated from its standard solution enthalpy (Asin//°) and its standard vaporization or sublimation enthalpy ... 

STANDARD SOLUTION ENTHALPY The measured standard solution enthalpy... 

STANDARD SOLUTION ENTHALPY The measured standard solution enthalpy of a-quartz and of the as-synthesized and the calcined MFI-type samples are as follows ... 

The difficulty in this cycle is the nonexistence of the undissociated hydrochloric acid in water. Hence, the measurement of the standard solution enthalpy... 

It is seen from equation (22) that there will, indeed, be a temperature at which the separation ratio of the two solutes will be independent of the solvent composition. The temperature is determined by the relative values of the standard free enthalpies of the two solutes between each solvent and the stationary phase, together with their standard free entropies. If the separation ratio is very large, there will be a considerable difference between the respective standard enthalpies and entropies of the two solutes. As a consequence, the temperature at which the separation ratio becomes independent of solvent composition may well be outside the practical chromatography range. However, if the solutes are similar in nature and are eluted with relatively small separation ratios (for example in the separation of enantiomers) then the standard enthalpies and entropies will be comparable, and the temperature/solvent-composition independence is likely be in a range that can be experimentally observed. 

For a Raoulf s law standard state, H° = Hf and L, = //, — Hf. These are the differences described in Chapter 5. For a Henry s law standard state, H° is the enthalpy in a hypothetical m = 1 (or X2 — 1 or c = 1) solution that obeys Henry s law. To help in understanding the nature of these standard state enthalpies, we will show that... 

If a2/ax = gim2lgxmx m2/mx (a = activity, g = molal activity coefficient, m = molality) if Henry s law is obeyed and a = gm, mh and m2 are the molal solubilities of the polymorphs, and if the standard molar enthalpies and standard molar entropies of solution are, respectively,... 

Having thus settled on Pedley s tables for the pure organic compounds, we have then decided to use NBS Tables to derive the solution enthalpies in figure 2.1. The values can be easily evaluated from the differences between the standard enthalpies of formation of the compounds in solution and the standard enthalpies of formation of pure substances, viz. 

In summary, we selected one database (Pedley s) to quote the standard enthalpies of formation of the pure organic compounds and another database (NBS) to derive the solution enthalpies. Although these databases are not mutually consistent, that did not affect our final result because the experimental enthalpies of solution were calculated with NBS data only. The exercise illustrates the sort of caution one should keep in mind whenever two or more nonconsistent databases are used. 

A similar exercise can be made with other anions and cations, producing a list of relative values of standard enthalpies of formation, anchored on Af77°(H+, ao) = 0. This database is rather useful, because it allows the enthalpies of formation (equation 2.53) and the lattice enthalpies (equation 2.47) of many crystalline ionic salts to be predicted, since their solution enthalpies are usually easy to measure. 

Reza, J., Trejo, A., and Vera-Avila, L.E. Water solubility and solution enthalpy of polycyclic aromatic hydrocarbons, in Fourteenth Symposium on Thermophysical Properties, (Boulder, CO National Institute of Standards and Technology. 2000). 

When a solute is dissolved in a solvent, heat may either be evolved (as with sulfuric acid in water, where strong heating is observed) or absorbed (as with ammonium nitrate in water, where strong cooling is observed). Thus the dissolution may be either exothermic or endothermic, and the standard molar enthalpy of solution, is then either negative or positive. The sign of AsaJl°... 

Table 9.8 Standard molar enthalpies (Ai7°) and standard entropies (A5 ) of solution of noble gases in silicate melts, after Lux (1987). (1) leucite-basanite (2) tholeiite (3) alkali-olivine basalt.

By applying equation 9.97 to plots of the type shown in figure 9.15B, Lux (1987) obtained standard molar enthalpies AH° and standard molar entropies ASf of solution of noble gases in melts, as listed in table 9.8. Enthalpies of solution are positive but close to zero, within error approximation. 

When strong bases neutralize strong acids in solutions that have molar concentrations of 1 mol dm-3, the enthalpy of the reaction is observed to be -55.83 kJ mol - irrespective of the counter ions (e.g. the chloride ion derivable from HC1 and the sodium ion contained in NaOH) present. For example, when a standard solution (1 mol dm-3) of hydrochloric acid is neutralized by a standard solution (1 mol dm-3) of sodium hydroxide, the change in enthalpy of the reaction is -55.83 kJ mol-1. Because the strong acid HCI and the strong base NaOH are 100% dissociated in aqueous solution, theucutruli/atiun reaction may be written as ... 

The enthalpies of formation of aqueous ions may be estimated in the manner described, but they are all dependent on the assumption of the reference zero that the enthalpy of formation of the hydrated proton is zero. In order to study the effects of the interactions between water and ions, it is helpful to estimate values for the enthalpies of hydration of individual ions, and to compare the results with ionic radii and ionic charges. The standard molar enthalpy of hydration of an ion is defined as the enthalpy change occurring when one mole of the gaseous ion at 100 kPa (1 bar) pressure is hydrated and forms a standard 1 mol dm-3 aqueous solution, i.e. the enthalpy changes for the reactions Mr + (g) — M + (aq) for cations, X (g) — Xr-(aq) for monatomic anions, and XOj (g) —< XO (aq) for oxoanions. M represents an atom of an electropositive element, e.g. Cs or Ca, and X represents an atom of an electronegative element, e.g. Cl or S. 

Solution The standard reaction enthalpy for the forward reaction in C is... 

Another remarkable Lewis basicity scale for 75 non-HBD solvents has been established by Gal and Maria [211, 212]. This involved very precise calorimetric measurements of the standard molar enthalpies of 1 1 adduct formation of EPD solvents with gaseous boron trifluoride, A//p gp, in dilute dichloromethane solution at 25 °C, according to Eq. (2-10a). 

Fig. 5-4. Schematic one-dimensional enthalpy diagram for the exothermic bimolecular Finkelstein reaction Cl -I- CFI3—Br Cl—CH3 -I- Br in the gas phase and in aqueous solution [469, 474, 476]. Ordinate standard molar enthalpies oi (a) the reactants, (b, d) loose ion-molecule clusters held together by ion-dipole and ion-induced dipole forces, (c) the activated complex, and (e) the products. Abscissa not defined, expresses only the sequence of (a). ..(e) as they occur in the chemical reaction.

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. 

The standard transformed enthalpies of formation of the three species are given by the following six equations. dHlzero, dH2zero, and dH3zero are also related by the equations for the enthalpy of dissociation. ) solution s Solve [ (dHlexpt == dHlzero +1.4775 (zi[[l]] 2 - nHi[ [1] ]) ionstr 0.5 / (1 + 1.6 ionstr 0.5), dH2expt == dH2zero +... 

From a thermodynamic point of view, the variation of standard free enthalpy associated to the electron transfer process represented by Equation (2.1), tAG °, can be related with the variation of such tliermodynamic quantity for the electron transfer process for species in solution phase, lAG °. and for the transfer of the oxidized, IaG, and reduced, forms of the electroactive species and the electrolyte cations, from the solution phase to the porous solid. The corresponding Bom-Haber-type cycle is shown in... 

The calorimetric measurements comprised determinations of the enthalpies of dissolution in 70 mass-% H2SO4, H2S04(aq, 1 2.33), as shown in Table A-40. This solution will be denoted by sin. The partial standard molar enthalpies of formation of H2O and H2SO4 were estimated from data in [82WAG/EVA]. 

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