Theoretical possibilities of determining energetic and thermodynamic characteristics of chemical entities in gaseous and condensed
(solid and liquid) phases are briefly reviewed. The considerations include quantum chemistry methods which enable evaluation
of energetic quantities and statistical thermodynamics dependencies necessary for determining other thermodynamic characteristics.
The possible applications of these methods are also discussed in brief.
In a recent publication, Drebushchak [ 1 ] has made the startling claim that “Thermodynamics is not a rigorous mathematical science.” He also stated that for a specific thermodynamic expression, “two different
Beside four approaches to the thermodynamics of GaSb-M(=S, Te) solid solutions the doping limits for extremely narrow concentration
regions are analysed and ranked in the Cu, Ge, Mn (p-dopants), S, Se, Te (n-dopants) and N, In (isoelectric) groups.
The amplitudes of the relaxation curves, as obtained by the Temperature-jump method have been used to measure simultaneously
equilibrium constant and enthalpy for the reaction of complex formation of Ni2+ ion by 2,6-dihydroxobenzoic acid in the presence of a buffer. The experiments have been performed by changing the concentration
of metal ion at constant ligand concentration andpH as in a complexometric titration. The points of such ‘dynamic titrations’ have been analysed by means of the concept of ‘normal
reactions’ which enabled us to transform a set of coupled individual steps into a set of kinetically independent reactions.
The potentialities of the dynamic titrations are discussed.
A method has been purposed to calculate some of the thermodynamic quantities for the thermal deformation of a smectite without
using any basic thermodynamic data. The Hançılı (Keskin, Ankara, Turkey) bentonite containing a smectite of 88% by volume
was taken as material. Thermogravimetric (TG) and differential thermal analysis (DTA) curves of the sample were obtained.
Bentonite samples were heated at various temperatures between 25–900°C for the sufficient time (2 h) until to establish the
thermal deformation equilibrium.
Cation-exchange capacity (CEC) of heated samples was determined by using the methylene blue standard method. The CEC was used
as a variable of the equilibrium. An arbitrary equilibrium constant (Ka) was defined similar to chemical equilibrium constant and calculated for each temperature by using the corresponding CEC-value.
The arbitrary changes in Gibbs energy (ΔGa0) were calculated from Ka-values. The real change in enthalpy (ΔH0) and entropy (ΔS0) was calculated from the slopes of the lnK vs. 1/T and ΔG vs. T plots, respectively. The real changes in Gibbs energy (ΔG0) and real equilibrium constant (K) were calculated by using the ΔH0 and ΔS0 values. The results at the two different temperature intervals are summarized as below: ΔG10=ΔH10−ΔS10T=−RTlnK1=47000−53t, (200–450°C), and ΔG20=ΔH20-ΔS20T=−RTlnK2=132000−164T, (500–800°C).
The role played by the metal ion in thermodynamics of azurin folding was addressed by studying the thermal denaturation of
the apo-form by differential scanning calorimetry (DSC), and by comparing the results with data concerning the holo protein.
The thermal unfolding experiments showed that at 25°C the presence of metal ion increases the thermodynamic stability of azurin
by 24 kJ mol−1. A comparison between the unfolding and the copper binding free energies allow us to assert that the unfolded polypeptide
chain binds copper and subsequently folds into native holo azurin, being this the thermodynamically most favourable process
in driving azurin folding.
A model was proposed to calculate some thermodynamic parameters for the acid dissolution process of a bentonite containing
a calcium-rich smectite as clay mineral along with quartz, opal and feldspar as impurities. The bentonite sample was treated
with H2SO4 by applying dry method in the temperature range 50–150°C for 24 h. The acid content in the dry bentonite-sulphuric acid mixture
was 45 mass%. The total content (x) of Al2O3, Fe2O3 and MgO remained in the undissolved sample after treatment was taken as an equilibrium parameter. An apparent equilibrium
constant, Ka, was calculated for each temperature by assuming Ka=(xm−x)/x where xm is the total oxide content of the natural bentonite. Also, an apparent change in Gibbs free energy, ΔGao, was calculated for each temperature by using the Ka value. The graphs of lnKavs. 1/T and ΔGaovs. T were drawn and then the real change in both the enthalpy, ΔHo and the entropy, ΔSo, values were calculated from the slopes of the straight lines, respectively. Inversely, real ΔGo and K values were calculated from the real ΔHo and ΔSo values through ΔGo = −RT ln K = ΔHo − TΔSo equation. The best ΔHo and ΔSo fittings to this relation were found to be 65687 J mol−1 and 164 J mol−1K−1, respectively.
The enthalpies of solution of α- and β-cyclodextrins is aqueous peptide solutions were determined experimentally at 298.15
K. The obtained results were used to calculate pair cross interaction parameters between solutes. The results are discussed
in terms of the likelysolute–solute interactions. For systems α-cyclodextrin+peptide and β-cyclodextrin+peptide the diametrically
opposite character of interaction defined by structure and solvation of the molecules is observed.
A novel system using a potassium aluminosilicate electrolyte under applied potential that is able to split H2O (or OH) into H2 and 1/2O2 (or O22-) with higher yields than the value deduced from Faraday"s law is presented. There were three steps by which H2 and O2 were generated stoichiometrically, and it was predicted that the high yields were due to the occurrence of chemically endothermic
reactions: dehydration of the catalytic cell at a temperature below 100C (step I), disproportionation of KOH (2KOH→H2+K2O2) at a temperature around 200C (step II), and disproportionation of K2O (2K2O→K2+K2O2) at a temperature above 500C (step III). So-called Nemca might be caused in the course of step III, since the rate of H2 was ca 102 times larger than the value deduced from Faraday"s law.