Ward N. Hubbard

The enthalpies of solution and neutralization of HF(I); enthalpies of dilution and derived thermodynamic properties of HF(aq)

Enthalpies of solution, ΔHsoln, of HF(I) in water and the enthalpy of neutralization, ΔHN, of HF(I) in dilute NaOH, as well as some enthalpies of dilution, ΔHdlln, of HF(aq) were measured in a reaction calorimeter. These results were combined with some enthalpies of dilution of HF(aq) reported in the literature to obtain the enthalpy of solution of liquid HF as a function of composition between HF·∞H2O and HF·H2O. The enthalpies of formation of all aqueous HF solutions in this concentration range were obtained by combination of the enthalpy of formation of HF(I) and the enthalpies of solution of HF(I). In addition, the relative apparent molar enthalpies, Lo, of HF(aq) and the relative partial molar enthalpies, L2 and L1, of HF and H2O, respectively, were calculated. Some of the key measured and derived results at 298.15 K obtained in this work are ΔHoN(HF, 1) = - (21017 ± 29) calth mol−1, ΔHosoln(HF, 1) = − (7672 ± 38) calth mol−1, and ΔHot,(HF, aq) = ΔHof(F−, aq) = − (80.22 ± 0.07) kcalth mol−1.

Fluorine bomb calorimetry. Part 14.—Enthalpies of formation of the hexafluorides of sulphur, selenium, and tellurium and their thermodynamic properties to 1500°K

The energies of formation of the hexafluorides of sulphur, selenium, and tellurium were measured by direct combination of the elements in a bomb calorimeter. From these measurements the standard enthalpies of formation ΔH°f298·15(g), were calculated (in kcal mole–1) : SF6, –291.77±0.24; SeF6, –266.95±0.14; TeF6, –327.20±0.56. These values are substantially more negative then those reported by Yost and Claussen in 1933. The thermodynamic properties of these compounds have been computed and tabulated from 0 to 1500°K.

The standard enthalpy of formation of LaNi5 The enthalpies of hydriding of LaNi5−xAlx☆

Abstract The enthalpies of reaction in HCl(aq) of LaNi 5 (cr), La(cr), and Ni(cr) were measured. From these measurements and appropriate auxiliary quantities, the standard enthalpy of formation, Δ f H m o , of LaNi 5 (cr) was determined. The enthalpies of reaction between H 2 and LaNi 5 , as well as LaNi 4.5 Al 0.5 , and LaNi 4 Al, have been measured by hydrogen titration calorimetry. The enthalpy of solution of H 2 in the α-phase of (LaNi 5 +H 2 ) is −28 kJ · mol −1 , and in the two-phase plateau region the enthalpy of hydriding by H 2 is −(31.6±0.2) kJ · mol −1 . In the two-phase plateau region of (LaNi 5− x Al x +H 2 ) the enthalpy of hydriding by H 2 is −38 kJ · mol −1 for LaNi 4.5 Al 0.5 and −46 kJ · mol −1 for LaNi 4 Al. The relation between the enthalpy of hydriding of the alloy by H 2 and its aluminum content was found to be Δ hyd H m o (LaNi 5− x Al x , cr, 298.15 K) = −(31.6 + 14.4 x ) kJ · mol −1 . The following thermodynamic quantities at T = 298.15 K are reported: Δ f H m o /(kJ · mol −1 ) Δ f S m o /(J · K −1 · mol −1 ) Δ f G m o /(kJ · mol −1 ) LaNi 5 −159.1±8.3 3.5±0.8 −160.1±8.3 LaNi 5 H 6.2 −258.6±8.3 −339.5±0.9 −157.4±8.3

Thermodynamics of the lanthanide trifluorides. I. The heat capacity of lanthanum trifluoride, LaF3 from 5 to 350°K and enthalpies from 298 to 1477°K

The heat capacity of a sample of LaF3 was determined in the temperature range 5–350°K by aneroid adiabatic calorimetry and the enthalpy from 298.15 to 1477°K by drop calorimetry. The heat capacity at constant pressure C°p(298.15°K), the entropy S° (298.15°K), the enthalpy [H° (298.15°K)−H° (0)] and the Planck function −[G° (298.15°K)−H° (0)]/298.15°K; were found to be (90.29±0.09) J °K−1⋅mole−1, (106.98±0.11) J °K−1⋅mole−1, (16717±17) J mole−1, and (50.91±0.05) J °K−1⋅mole−1. The thermal functions from the present research were extended up to the melting temperature (1766°K) by combination with previously published results. The anomalously high heat capacity from about 1100 to 1766°K is discussed.

The enthalpies of formation of XeF6(c), XeF4(c), XeF2(c), and PF3(g)☆☆☆

Abstract The energies of reaction of XeF6(c), XeF4(c), and XeF2(c) with PF3(g) were measured in a bomb calorimeter. These results were combined with the enthalpy of fluorination of PF3(g), which was redetermined to be −(151.98 ± 0.07) kcalth mol−1, to derive (at 298.15 K) ΔHfo(XeF6, c, I) = −(80.82 ± 0.53) kcalth mol−1, ΔHfo(XeF4, c) = −(63.84 ± 0.21) kcalth mol−1, and ΔHfo(XeF2, c) = −(38.90 ± 0.21) kcalth mol−1. The enthalpies of formation of the solid xenon fluorides were combined with reported enthalpies of sublimation to derive (at 298.15 K) ΔHfo(XeF6, g) = −(66.69 ± 0.61) kcalth mol−1, ΔHfo(XeF4, g) = −(49.28 ± 0.22) kcalth mol−1, and ΔHfo(XeF2, g) = −(25.58 ± 0.21) kcalth mol−1. The average bond dissociation enthalpies,〈Do〉(XeF, 298.15 K), are (29.94 ± 0.16), (31.15 ± 0.13), and (31.62 ± 0.16) kcalth mol−1 in XeF6(g), XeF4(g), and XeF2(g), respectively. The enthalpy of formation of PF3(g) was determined to be −(228.8 ± 0.3) kcalth mol−1.

The enthalpies of formation of plutonium dioxide and plutonium mononitride

Abstract The energies of combustion in oxygen of α-plutonium metal and plutonium mononitride were measured in a bomb calorimeter. The standard enthalpies of formation, ΔH f o (298.15 K), of PuO 2 (c) and PuN(c) were calculated to be −(1055.85±0.72) kJ mol −1 [−(252.35±0.17) kcal mol −1 ] and −(299.2±2.6) kJ mol −1 [−(71.51±0.62) kcal mol −1 ], respectively.

Calorimetric measurements of the low-temperature heat capacity, standard molar enthalpy of formation at 298.15 K, and high-temperature molar enthalpy increments relative to 298.15 K of tungsten disulfide (WS2), and the thermodynamic properties to 1500 K

Low-temperature (5 to 350 K) heat capacity, fluorine combustion, and high-temperature (350 to 1500 K) drop-calorimetric measurements have been performed on a pure synthetic specimen of tungsten disulfide, WS2. The following molar thermodynamic quantities are reported at To = 298.15 K: the standard enthalpy of formation, ΔfHmo(To), −(240.8±3.1) kJ·mol−1; the heat capacity, Cp,mo(To), (63.82±0.32) J·K−1·mol−1; the standard entropy, Smo(To), (67.78±0.34) J·K−1·mol−1; and the standard Gibbs energy of formation, ΔfGmo(To), −(232.1±3.1) kJ·mol−1. The thermodynamic quantities have been calculated to 1500 K. Standard enthalpies of formation deduced from high-temperature equilibrium and e.m.f. studies in the literature are, in general, not in good agreement with one another or the present result. There have been no previous measurements of the low-temperature heat capacity, but the high-temperature enthalpy increments are in fair agreement with results published for WS1.97. There is some indication of a γT (i.e. conduction electrontribution to the heat capacity at low temperatures, but the evidence for this is not strong. The present thermodynamic quantities are consistent with geochemical field observations that molybdenite (MoS2) and not tungstenite (WS2) but tungstates and not molybdates are formed in hydrothermal deposits.

The enthalpy of formation of lithium oxide (Li2O)

Abstract The enthalpy of reaction with water of a carefully characterized specimen of lithium oxide, Li 2 O, was found to be −(31.488 ± 0.025) kcal th mol −1 . Combination of the measured result with auxiliary thermochemical data from the literature yielded a value of −(142.902 ± 0.066) kcal th mol −1 for the standard enthalpy of formation of lithium oxide, ΔH f o (Li 2 O, c, 298.15 K).

Enthalpy of formation of disodium acetylide and of monosodium acetylide at 298.15 K, heat capacity of disodium acetylide from 5 to 350 K, and some derived thermodynamic properties

The enthalpies of reaction of Na 2 C 2 and of NaHC 2 with water were measured, and the standard enthalpies of formation of Na 2 C 2 and NaHC 2 at 298.15 K were determined to be (4.80 ± 0.40) and (23.10 ± 0.27) kcal th mol −1 , respectively. The heat capacity of Na 2 C 2 was measured from 5 K to 350 K. The following values were obtained for Na 2 C 2 (c) at 298.15 K: heat capacity at constant pressure: C o p (298.15 K) = (90.98 ± 0.27) J K −1 mol −1 ; entropy: S o (298.15 K) = (111.24 ± 0.33) J K −1 mol −1 ; enthalpy: { H o (298.15 K) − H o (0)} = (17389 ± 52) J mol −1 ; Gibbs energy divided by temperature: { G o (298.15 K) − H o (0)}/298.15 K = −(52.92 ± 0.16) J K −1 mol −1 . The standard Gibbs energy of formation, Δ G o f , of Na 2 C 2 at 298.15 K is (5.02 ± 0.41) kcal th mol −1 . The standard Gibbs energy of formation at higher temperatures was estimated from an extrapolation of the heat capacity curve. The concentration of Na 2 C 2 in liquid sodium in equilibrium with graphite was calculated from Δ G o f and the solubility of Na 2 C 2 in liquid sodium, and a comparison was made with experimental values for the solubility of graphite in liquid sodium. The comparison indicates that dissolved Na 2 C 2 can be an important species, and possibly the only species, formed when graphite dissolves in liquid sodium. It concluded that models based on dissolved Na 2 C 2 for the transport of carbon in liquidsodium-cooled reactors are thermodynamically possible.

Uranium Diboride: Preparation, Enthalpy of Formation at 298.15°K, Heat Capacity from 1° to 350°K, and Some Derived Thermodynamic Properties

A sample of uranium diboride was prepared and characterized as UB1.979±0.006 with 0.06 ± 0.03 wt % of identified impurities. The standard enthalpy of combustion in fluorine was determined to be − 1021.2 ± 1.1 kcal mole−1. The heat capacity was measured from 0.84° to 350°K. At 298.15°K the heat capacity CP°, entropy S°, and enthalpy increment H° − H°0 are 13.23 ± 0.03 cal K−1·mole−1, 13.17 ± 0.03 cal °K−1·mole−, and 2108 ± 4 cal mole−1, respectively. The following values were obtained for the standard enthalpy, entropy, and Gibbs energy of formation of UB2 at 298.15°K: ΔHf° = − 39.3 ± 4.0 kcal mole−1, ΔSf° = − 1.54 ± 0.05 cal °K−1·mole−, and ΔGf° = − 38.8 ± 4.0 kcal mole−1. These agree within experimental error with values calculated from high‐temperature effusion measurements. The heat‐capacity results below 4.2°K follow the equation CP = (9.40 ± 0.01)T + (3.18 ± 0.14) × 10−2T3mJ °K−1·mole−1. The relatively high value for the coefficient of the linear term indicates that uranium diboride is a good electri...