Enthalpy sulfate

Almost all of the directly measured thermochemical data for the sulfoxides, sulfones, sulfites and sulfates are due to the work of Busfield and Mackle and their coworkers at the University of Leeds and The Queens University, Belfast1-14. This work involved measurement of enthalpies of combustion, fusion and vaporization. It is the basis of the subsequent compilations of Benson and coworkers15, Cox and Pilcher16 and Pedley, Naylor and Kirby11. The data given by the latter are used as the basic data set in the present work. Corrections and omissions are noted in the next section. Data on additional compounds were sought by searching the IUPAC Bulletin of Thermochemistry and Thermodynamics for the years 1980 198318, and by searches of Chemical Abstracts. 

Stability Constants, Enthalpies, and Entropies of Plutonium(III) and Plutonium(IV) Sulfate Complexes... 

DeCarvalho and Choppin (10, 11) previously have reported the stability constants, complexation enthalpies, and entropies for a series of trivalent lanthanide and actinide sulfates. As their work was conducted a pH 3, the dominant sulfate species was S0 and the measured reaction was as in equation 12. 

Stability constants, enthalpies, and entropies have been published for both the 1 1 and 1 2 sulfate complexes of the Th1 "1"... 

In contrast to the situation observed in the trivalent lanthanide and actinide sulfates, the enthalpies and entropies of complexation for the 1 1 complexes are not constant across this series of tetravalent actinide sulfates. In order to compare these results, the thermodynamic parameters for the reaction between the tetravalent actinide ions and HSOIJ were corrected for the ionization of HSOi as was done above in the discussion of the trivalent complexes. The corrected results are tabulated in Table V. The enthalpies are found to vary from +9.8 to+41.7 kj/m and the entropies from +101 to +213 J/m°K. Both the enthalpy and entropy increase from ll1 "1" to Pu1 with the ThSOfj parameters being similar to those of NpS0 +. Complex stability is derived from a very favorable entropy contribution implying (not surprisingly) that these complexes are inner sphere in nature. 

I.ithium sulfate dissolves exothermically in water, (a) Is the enthalpy of solution for Ei,S04 positive or negative ... 

Khim 10(1), 454-7(1965) Sc CA 62, 12515 (l9659(Enthalpy of formation of guanidine perchlorate, nitrate, and sulfate) (Give the value —287kcaljn)... 

Estimate the standard enthalpy of hydration of the hydrogen sulfate ion. given that the lattice enthalpy of sodium hydrogen... 

Where this factor plays a role, the hydrophobic interaction between the hydrocarbon chains of the surfactant and the non-polar parts of protein functional groups are predominant. An example of this effect is the marked endothermic character of the interactions between the anionic CITREM and sodium caseinate at pH = 7.2 (Semenova et al., 2006), and also between sodium dodecyl sulfate (SDS) and soy protein at pH values of 7.0 and 8.2 (Nakai et al., 1980). It is important here to note that, when the character of the protein-surfactant interactions is endothermic (/.< ., involving a positive contribution from the enthalpy to the change in the overall free energy of the system), the main thermodynamic driving force is considered to be an increase in the entropy of the system due to release into bulk solution of a great number of water molecules. This entropy... 

The reaction of Problem 10 was studied at two different temperatures, and, from the temperature dependence of the rate constants, the authors determined AH% and ASJ, the enthalpy and entropy of activation, respectively. The following values of these parameters were obtained in pure water and in 0.01 M sodium dodecyl sulfate (NaLS) and 0.01 M hexadecyl trimethyl ammonium bromide (CTABr) ... 

Figure 16.3. Enthalpy-composition diagrams of some salt solutions. Several other diagrams are in the compilation of Landolt-Bomstein, IV 4b, 1972, pp. 188-224. (a) sodium sulfate/water (b) magnesium sulfate/water (after Chemical Engineers Handbook, 1963 edition, McGraw-Hill, New York) (c) sodium carbonate/water.

Figure 15.9 Enthalpy of formation of a series of silicates, tungstates, carbonates and sulfates from the oxides.

The type of the oxidation product on galena is independent of the chemical environment during preparation. Rao152) measured the adsorption heat of K amyl xanthate (KAX) on unactivated and Cu2+-activated pyrrhotite (FeS) and compared his results with heats of the reaction between KAX and Fe2+ or Cu2+ salts. With the unactivated mineral, the interaction involves a chemical reaction of xanthate with Fe2+ salts present at the interface (i.e. not bound to the crystal surface). The adsorption enthalpy is identical with the formation of Fe2+ amyl xanthate FeS04 + 2 KAX —> FeX2 + K2S04, and -AH = 97.45 kJ/mol Fe2+). As revealed from the enthalpy values and the analysis of anions released into the solution, the interaction of xanthate with Cu2+-activated pyrrhotite consists of xanthate adsorption by exchange for sulfate ions (formed by an oxidation of sulfides) at isolated patches (active spots), and by further multilayer formation of xanthate. The adsorption heat of KAX on pyrrhotite at the initial pH 4.5 was - AH (FeS unactivated) = 93.55 kJ/mol Fe2+ and - AH (FeS activated) = 70.03 kJ/mol Cu2+. 

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