Showing papers on "Standard molar entropy published in 1973"
TL;DR: In this article, the temperature, enthalpy, entropy of melting and crystal transitions of 21 polycyclic aromatic hydrocarbons, containing from 2 to 6 unsubstituted condensed rings, were determined by differential scanning calorimetry.
55 citations
TL;DR: In this paper, the van't Hoff analysis of the data for the reaction CO2H++CH4⇄ CH5++CO2 results in a value for the standard enthalpy change, ΔH° (196 to 553°K), of −0.151±0.004 eV and a value of −1.8±1.1 eV for ΔS° (277 to 505°K).
Abstract: The flowing afterglow measurements of the equilibrium constants for selected proton transfer reactions have been extended to temperatures other than room temperature in order to completely specify the thermodynamics of such reactions. The van't Hoff analysis of the data for the reaction CO2H++CH4⇄ CH5++CO2 results in a value for the standard enthalpy change, ΔH° (196 to 553°K), of −0.064±0.004 eV and for the standard entropy change, ΔS° (196 to 553°K), of +1.4±0.6 e.u. Studies of the reaction N2OH++CO⇄ COH++N2O led to a value of −0.151±0.009 eV for ΔH° (277 to 505°K) and a value of −1.8±1.1 e.u. for ΔS° (277 to 505°K). ΔH° is a direct measure of the difference between the proton affinities of CO2 and CH4 in the former case and of N2O and CO in the latter. The measured values for ΔS° yield differences in the standard entropies of the protonated species: S°298(CH5+)−S°298(CO2H+)= −5.2± 0.6 e.u and S°298(COH+)−S°298(N2OH+)= −7.1± 1.1 e.u. Reference to calculated values for the standard entropy of CH5+ an...
38 citations
TL;DR: The proton transfer reaction NH2−+H2⇄ H−+NH3 has been investigated in a flowing afterglow system independently in both directions at 297°K and the ratio of rate constants was found to be in agreement with the equilibrium constant, K, determined from equilibrium concentrations.
Abstract: The proton transfer reaction NH2−+H2⇄ H−+NH3 has been investigated in a flowing afterglow system independently in both directions at 297°K. The rate constants were found to be kforward= (2.3± 0.5)× 10−11 cm3 molecule−1· sec−1 and kreverse= (9.2± 1.8)× 10−13 cm3 molecule−1· sec−1. The ratio of rate constants, kf/kr= 26−5+6 was found to be in agreement with the equilibrium constant, K= 27± 9, determined from equilibrium concentrations. The value for K corresponds to a standard free energy change, Δ G°297, of− 1.9± 0.2 kcal mole−1. A calculated value for the standard entropy change, ΔS°298, of −4.3± 0.5 cal deg−1· mole−1 leads to a value for the standard enthalpy change, Δ H°298, of−3.2± 0.3 kcal mole−1. Adoption of a recent photodetachment value for the electron affinity of NH2 leads to D°0(NH2–H)= 106.0± 1.1 kcal mole−1, D°298(NH2–H)= 107.4± 1.1 kcal mole−1, and Δ H°f,298(NH2)= 44.3± 1.1 kcal mole−1.
31 citations
TL;DR: In this article, a simple-mixture model of non-ideal mixing is combined with equations of state for oxygen buffers to calculate the chemical potentials of the two components in a binary gas mixture.
Abstract: Reversed univariant hydrothermal phase-equilibrium reactions, in which a redox reaction occurs and is controlled by oxygen buffers, can be used to extract thermochemical data on minerals. The dominant gaseous species present, even for relatively oxidizing buffers such as the QFM buffer, are H2O and H2; the main problem is to calculate the chemical potentials of these components in a binary mixture. The mixing of these two species in the gas phase was assumed by Eugster and Wones (1962) to be ideal; this assumption allows calculation of the chemical potentials of the two components in a binary gas mixture, using data in the literature. A simple-mixture model of nonideal mixing, such as that proposed by Shaw (1967), can also be combined with the equations of state for oxygen buffers to permit derivation of the chemical potentials of the two components. The two mixing models yield closely comparable results for the more oxidizing buffers such as the QFM buffer. For reducing buffers such as IQF, the nonideal-mixing correction can be significant and the Shaw model is better. The procedure of calculation of mineralogical thermochemical data, in reactions where hydrogen and H2O simultaneously appear, is applied to the experimental data on annite, given by Wones et al. (1971), and on almandine, given by Hsu (1968). For annite the results are: Standard entropy of formation from the elements, S
f
0
(298, 1)=−283.35±2.2 gb/gf, S
0 (298, 1) =+92.5 gb/gf. G
f
0
(298, 1)=−1148.2±6 kcal, and H
f
0
(298, 1)=−1232.7±7 kcal. For almandine, the calculation takes into account the mutual solution of FeAl2O4 (Hc) in magnetite and of Fe3O4 (Mt) in hercynite and the temperature dependence of this solid solution, as given by Turnock and Eugster (1962); the calculations assume a regular-solution model for this binary spinel system. The standard entropy of formation of almandine, S
f,A
0
(298, 1) is −272.33±3 gb/gf. The third law entropy, S
0 (298, 1) is +68.3±3 gb/gf, a value much less than the oxide-sum estimate but the deviation is nearly the same as that of grossularite, referring to a comparable set of oxide standard states. The Gibbs free energy G
f,A
0
(298, 1) is −1192.36±4 kcal, and the enthalpy H
f,A
0
(298, 1) is −1273.56±5 kcal.
24 citations
TL;DR: In this article, the α-phase of the palladium/hydrogen system in the temperature range 372-724 K, by direct equilibration with gaseous hydrogen, was determined.
Abstract: Pressure against composition isotherms have been determined for the α-phase of the palladium/hydrogen system in the temperature range 372–724 K, by direct equilibration with gaseous hydrogen. The partial molar enthalpy and partial molar entropy of solution of hydrogen in palladium have been derived as a function of composition. The values obtained at 600 K are, respectively, –15.56 kJ (mol H2)–1, and –96.5 JK–1(mol H2)–1.
22 citations
TL;DR: In this paper, the Gibbs adsorption equation is rearranged in a suitable form, where the coefficient of dT does not depend on the choice of the standard entropy, and the heat of reversible adorption at constant temperature and surfactant mole fraction in the bulk can be calculated on a thermodynamic basis.
19 citations
TL;DR: In this paper, the authors studied the monomer formation from the dimer by n.m.r. spectroscopy at various temperatures and showed that the resulting gem-chloronitroso derivatives obtained isomerize into hydroximoyl chlorides.
16 citations
TL;DR: In this article, the Gibbs energy of formation of IrO2 has been determined by means of electromotive force measurements using the cell Pt/Fe, FeO/ZrO2, Y2O3/Ir, YO2/Pt.
10 citations
TL;DR: In this article, the standard enthalpy and standard entropy changes at 25°C for the trimerisation of liquid acetaldehyde to form liquid paraldehyde are shown to be thermodynamically unpolymerisable.
Abstract: Equilibrium constants for reaction (1) 3CH3CHO(g)⇌(CH3CHO)3(g)(1) have been measured in the temperature range 19 to 40°C. The results can be expressed by eqn (2). log(Kp/m4 kN–2)=(6.97 ± 0.09)103K/T–(27.90 ± 0.30)(2) The standard enthalpy and standard entropy changes at 25°C for the trimerisation of liquid acetaldehyde to form liquid paraldehyde are ΔH°298=–98.1 kJ (mol paraldehyde)–1 and ΔS°298=–299.5 J K–1(mol paraldehyde)–1. Liquid paraldehyde is shown to be thermodynamically unpolymerisable.
10 citations
TL;DR: In this paper, a calorimetric enthalpy of dilution technique has been applied to the investigation of the self-association of several chloro-substituted carboxylic acids in dilute solutions in anhydrous carbon tetrachloride at 25°C.
9 citations
TL;DR: In this paper, the authors measured the heat capacity of KFCT and the anhydrous salt from 10 K to 300 K. The results gave a value of 593.7 J K−1 mol−1 for the molar entropy derived from the equilibrium study, in almost exact agreement with the calorimetric entropy.
Abstract: The heat capacity of potassium ferrocyanide trihydrate (KFCT) and the anhydrous salt has been measured from 10 K to 300 K. The contribution made to the heat capacity of KFCT by the ferroelectric → paraelectric transition has been assessed (a) by using the heat capacity of the anhydrous salt as a reference, (b) by subtracting from the total heat capacity of KFCT the contribution from the internal vibrations of the ferrocyanide ion. The values so obtained for the entropy of transition are (a) 7.8 J K–1 mol–1, (b) 9.7 J K–1 mol–1. Measurements have also been made, at 25°C, of the heat of solution of the two salts and of the water vapour dissociation pressure of the KFCT + anhydrous salt system. The results give a value of 593.7 J K–1 mol–1 for the molar entropy of KFCT derived from the equilibrium study, in almost exact agreement with the calorimetric entropy of 593.8 J K–1 mol–1, showing that the water molecules in KFCT in the ferroelectric phase are uniquely ordered at 0 K.The standard entropy of the ferrocyanide ion in aqueous solution at 25°C, text-decoration:overlineS°(Fe(CN)4–6, 298.15 K), is estimated to be 76.6 J K–1 mol–1.
TL;DR: In this article, the acid and base hydrolysis of nitratopentaamminerhodium (III) and iridium(III) cations have been studied kinetically and the observed rate constants were found to follow the equation k obs = k aq + k OH [OH − ].
01 Sep 1973-Journal of Research of the National Bureau of Standards Section A: Physics and Chemistry
TL;DR: A reevaluation of the data for the enthalpy of solution of AgNO3(c) has resulted in the selected best value, Δ H ° ( ∞ ) ( 298.15 K) , which is found to be S ° [ Ag + ( aq ) ] = 73.42 ± 0.20 ’ J ⋅ mol - 1⋅ K - 1 = 17.55 ± 0%.
Abstract: The enthalpies of precipitation of silver halides and the enthalpies of solution of AgNO3, KCl, and KBr in H2O were measured in an adiabatic solution calorimeter. From the enthalpy measurements of KCl(c) and KBr(c) in AgNO3(aq), and of AgNO3(c) in KCl(aq), in KBr(aq), and in KI(aq), we calculated (in kJ · mol-1) -65.724, -84.826, and -111.124 for ΔH° pptn(298.15 K) for the averages of the chloride, bromide, and iodide reactions, respectively. A reevaluation of the data for the enthalpy of solution of AgNO3(c) has resulted in our selected best value, Δ H ° ( ∞ ) ( 298.15 K ) = 22.730 + 0.084 kJ ⋅ mol - 1 = 5.433 ± 0.020 kcal ⋅ mol - 1 A table of enthalpies of dilution of AgNO3(aq) is also given. The average standard entropy for the aqueous silver ion at 298.15 K is found to be S ° [ Ag + ( aq ) ] = 73.42 ± 0.20 J ⋅ mol - 1 ⋅ K - 1 = 17.55 ± 0.05 cal ⋅ mol - 1 ⋅ K - 1 . .
TL;DR: In this article, the formation of piperidinium tetramethoxyborate from piperidine, trimethoxy-borane, and methanol is discussed.
Abstract: Thermodynamic parameters for the formation of piperidinium tetramethoxyborate from piperidine, trimethoxyborane, and methanol are presented. The results show that the formation of the salt is due to a favourable enthalpy change, the standard entropy change being negative. The fact that similar salts with larger tetra-alkoxyborate anions are thermodynamically unstable is discussed. The derived enthalpy of formation of piperidinium tetramethoxyborate ΔHf° C5H12NB(OMe)4(s)=–1310·91 ± 1·72 kJ mol–1. The derived enthalpy of formation of 2-aminoethylammonium tetramethoxyborate ΔHf°enH B(OMe)4(s)=–1266·5 ± 1·5 kJ mol–1 and of ethylenediammonium ditetramethoxyborate ΔHf°enH2[B(OMe)4]2(s)=–2486·0 ± 2·0 kJ mol–1.
TL;DR: In this paper, the saturated molar heat capacity of oxalyl fluoride gas was measured from 13°K to its normal boiling point, 270.13°K, for a sample of 99.94 mole% purity.
Abstract: The saturated molar heat capacity has been measured from 13°K to its normal boiling point, 270.13°K, for a sample of 99.94 mole% purity. The molar heat of fusion was found to be 3204 cal/mole at the triple‐point temperature of 260.73°K. The molar heat of vaporization at the normal boiling point was found to be 6753 cal/mole. The vapor pressure of the solid in the range from 234°K to the triple point of 260.73°K is represented by the equation lnP(torr)=115.38986 − 8720.5964/T − 13.623524 lnT. The vapor pressure of the liquid from the triple point to the boiling point of 270.13°K is represented by the equation lnP(torr)=185.93601 − 10 317.218/T − 25.202996 lnT. The standard entropy of oxalyl fluoride gas at the normal boiling point was found to be 73.01± 0.16 cal/mole·° K. Neither of the two fundamental frequency assignments reported in the literature when used in the statistical entropy calculation gives satisfactory agreement with the calorimetric entropy. It is believed that the discrepancy results from ...
TL;DR: In this article, the vapor pressure of the initiation of decomposition of C 8 Br has practically the same value as that of initiation of bromination of C 12 Br, which is believed to give the equilibrium decomposition pressure of C8 Br.
TL;DR: In this paper, the authors investigated the complexation of cobalt(II) ion with azide ion, by spectroscopic methods, when water and absolute methanol are the solvents.
Abstract: Complexation of cobalt(II) ion with azide ion is investigated, by spectroscopic methods, when water and absolute methanol are the solvents. In water the octahedral monoazidopentaaquocobalt(II) ion was obtained, whereas in alcohol the tetrahedral tetraazidocobalt(II) was obtained. Stability constants were computed for both complexes at different temperatures. An estimate of the standard enthalpy as well as the standard entropy of complexation was obtained.
TL;DR: In this article, the formation of piperidinium tetramethoxyborate from piperidine, trimethoxy-borane, and methanol is discussed.
Abstract: Thermodynamic parameters for the formation of piperidinium tetramethoxyborate from piperidine, trimethoxyborane, and methanol are presented. The results show that the formation of the salt is due to a favourable enthalpy change, the standard entropy change being negative. The fact that similar salts with larger tetra-alkoxyborate anions are thermodynamically unstable is discussed. The derived enthalpy of formation of piperidinium tetramethoxyborate ΔHf° C5H12NB(OMe)4(s)=–1310·91 ± 1·72 kJ mol–1. The derived enthalpy of formation of 2-aminoethylammonium tetramethoxyborate ΔHf°enH B(OMe)4(s)=–1266·5 ± 1·5 kJ mol–1 and of ethylenediammonium ditetramethoxyborate ΔHf°enH2[B(OMe)4]2(s)=–2486·0 ± 2·0 kJ mol–1.