Ammonium thermodynamic data

The thermodynamic data of Tables 2 and 3 are remarkably selfconsistent, and correlate well with data for the various phenyltrimethyl ammonium perchlorates in 1,2-dichloroethane originally studied by Denison and Ramsey (44). Clearly dissociation is nominally exothermic though the actual enthalpy changes are quite small. Solvation of the free ions, therefore, is also small (in these solvents) but exceeds that of the ion pair species. The overall entropy changes confirm some ordering of solvent molecules on dissociation, the loss in entropy for this process being more than sufficient to balance the entropy increase associated with... 

Physical chemical studies of dilute alkali metal-ammonia solutions indicate the principal solution species as the ammoniated metal cation M+, the ammoniated electron e , the "monomer M, the "dimer" M2 and the "metal anion" M. Most data suggest that M, M2, and M are simple electrostatic assemblies of ammoniated cations and ammoniated electrons The reaction, e + NH3 - lf 2 H2 + NH2 is reversible, and the directly measured equilibrium constant agrees fairly well with that estimated from other thermodynamic data. Kinetic data for the reaction of ethanol with sodium and for various metal-ammonia-alcohol reductions of aromatic compounds suggest that steady-state concentrations of ammonium ion are established. Ethanol-sodium reaction data allow estimation of an upper limit for the rate constant of e + NH4+ 7, H2 + NH3. 

A listing of thermodynamic data for a number of substituted triaryloxy-boranes (XC6H40)3B, where X = w- orp-C, -Me, -OMe, or -CN, has been made. Similar data are reported for several quaternary ammonium salts of the tetramethoxyborate anion, e.g. l Hf(a) = —1310.91 1.72kJmol for the piperidinium, —1266.5 1.5 kJmoL for the 2-aminoethylam-monium salts, and —2486.0 2,0kJmol for ethylenediammoniurn di-tetramethoxyborate. 

A Ho(PDC)3(o-phen) complex has been obtained and studied by Xie et al. [184] from the reaction of hydrated holmium chloride, ammonium pyrrolidinedithiocarbamate (APDC) and 1,10-phenanthroline (o-phen-H20) in absolute ethanol. The enthalpy of complex formation from a solution of the reagents and the molar heat capacity of fre complex were determined by using a heat conduction microcalorimeter. The enthalpy of complex formation from the reaction of the reagents in the solid phase, was calculated on the basis of an appropriate thermochemical cycle and other auxiliary thermodynamic data. The thermodynamics of formation of the complex were investigated by the reaction in solution at the temperature range of 292.15 - 301.15 K. The constant-volume combustion energy of the complex, Act/, was also determined. 

Using the thermodynamic data of Ananthapadmanabhan, species distribution diagram of dodecylamine was constructed (see Figure 3). The maximum in amine-ammonium complex formation at 1.6 X 10 kmol/m concentration occurs at pH 10. Ananthapadmanabhan and Pugh have shown the significance of such complexes in flotation, even if present in small amounts. 

The a-form, which crystallines as the kinetic product, is the commercial form of ammonium pentaborate tetrahydrate and the P-form is the thermodynamic product but is slow to crystalline. Its heat capacity has been measured over a broad temperature range (85). Solubihty data are given ia Table 9 and pH data ia Table 10. 

In principle, a thermodynamic value for the ion-dipole interaction could be acquired as well by the incorporation of several alkanediamines into our data set in the foregoing analysis. However, a second ammonium ion has in fact a disproportionate effect, and it cannot be meaningfully placed on the same scale... 

Considering the thermodynamic parameters of the reaction, it has been determined that the silicic acid with a hexacoordinated atom of silicon H2[Si(OH)6] formed in die process of reaction displaces ammonium ion from phenolates, so that this acid is stronger than catechol. This is in agreement with our experimental data at the beginning of the reaction process the decrease in the pH of the reaction mixture from 7.16 to 5.85 was observed before the pH became constant. 

Figure 3. Pressure-temperature-reaction phase diagram for ammonium dinitramide (ADN) showing the estimated thermodynamic stability fields for the a and P poljnnorphs and the liquidus curve. The monoclinic a phase is stable up to about 2.0 GPa between -75 and 120 °C. The high-pressure P phase, which is also monoclinic, is stable above 2.1 GPa between -75 and 120 °C. Above 140 °C between 1.0 and 10.0 GPa, ADN undergoes a molecular rearrangement to form ammonium nitrate (AN) and N2O. The a-P transition pressure is estimated to be 2.0 0.2 GPa and is the result of a least squares fit of the data points. Reproduced with permission from Ref. [39].

However, there had not been a spectroscopic study that confirmed this HM. Moreover, this model is unable to explain certain experimental data and does not account for the thermodynamical requirement that the solution temperature be below the Krafft point for micellization to take place. In order to understand these discrepancies, adsorption of long-chain amines was studied spectroscopically on quartz and albite, and the results were compared with the indirect data [325, 372]. DRIFTS, IRRAS, and XPS of quartz and albite conditioned in solutions of dodecyl (C12) and hexadecyl (Cie) ammonium chloride (ACl) and acetate (AAc) of different initial concentrations (Cj,) were measured. 

Booth and coworkers also studied the surface tension of binary systems (dicar-boxylic acids (oxalic, malonic, succinic, glutaric, and adipic acids) and ammonium sulfate) both experimentally (at 21 °C) and using modeling to determine whether additive or thermodynamic surface tension models agreed better with experimental data [190]. Following the methodology of Topping et al. [143] they found that the LiLu model best predicted binary systems. 

In the following article we will compare the relative protonation energies of a number of amines at 25°C in three media the gas phase, fluorosulfuric acid (HSO3F), and water. By means of a simple thermodynamic cycle we will analyse the liquid-phase data to derive the relative solvation energies of the ammonium ions from the gas phase to HSO3F and to water. We will discuss these results in terms of the interplay between the internal and external factors which stabilize ions by charge distribution. 

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