Ziley Singh

Determination of standard molar Gibbs' energy of formation of Sr3MoO6(s)

An electromotive oxygen concentration cell using calcia-stabilized zirconia (CSZ) electrolyte has been constructed and used to determine ΔfGmo(SrMoO4, s, T). The e.m.f. of cell (II), Pt/SrMoO3(s) + SrMoO4(s)/CSZ/air/Pt, could be represented by E (V) ± 0.006 = 1.515-4.299 × 10−4T (K), while the e.m.f. of cell (III), Pt/SrMoO3(s) + SrMoO4(s)/ CSZ/FexO(s) + Fe(s)/Pt, could be represented by E (V) ± 0.002 = 0.1459 - 5.859 × 10−5T (K). The e.m.f. values were combined with ΔfGmo(T) of SrMoO3(s) and FexO(s) from our earlier studies to get ΔfGmo(SrMoO4, s, T) values as ΔfGmo(SrMoO4, s, T) (kJmol−1) ± 1.2 = −1582.6 + 0.3692T (K) (1046.0 ⩽ T ⩽ 1255.5 K) for cell (II) and ΔfGmo(SrMoO4, s, T) (kJmol−1)± 0.4 = −1583.0 + 0.3699T (K) (1037.5 ⩽ T ⩽ 1278.0 K) for cell (III). From the variation in ΔfGmo(SrMoO4, s, T) (kJmol−1) with temperature and the relevant heat capacity values from the literature, ΔfHmo(298.15 K) of SrMoO4(s) has been obtained as −(1582.9±0.5) and −(1583.3±0.1) kJmol−1 respectively and the average value is −(1583.1±0.6) kJmol−1.

Calorimetric investigations of SrMoO3 and BaMoO3 compounds

Abstract Enthalpy increment (HT°−H298.15 K°) values of the compounds SrMoO3(s) and BaMoO3(s) were measured in the temperature range 426.2 to 1010.8 K (±0.1 K) and 477.8 to 1010.8 K (±0.1 K) respectively, using a high temperature Calvet calorimeter. The data was least square fitted with respect to temperature and is given by the following polynomials: SrMoO 3 (s) : H ° T −H ° 298.15 K ( J/mol )±1143=−29291(±2963)+91.942(±9.034)×T( K )+0.0264(±0.0066)×T 2 ( K ), BaMoO 3 (s) : H ° T −H ° 298.15 K ( J/mol )±812=−31427(±2279)+97.977(±6.952)×T( K )+0.0275(±0.0050)×T 2 ( K ). Heat capacities of the compounds calculated from the above equations were compared with estimated heat capacity values obtained by Neumann Kopps rule. Standard entropy (Sm°(298.15 K)), standard enthalpy of formation (ΔfH°(298.15 K)) and free energy functions of the compounds were calculated using heat capacities obtained from present experiments, Gibbs free energy of formation of the compounds reported in the literature and elemental thermodynamic data.

Standard molar Gibbs free energy of formation of NaCrO2 by e.m.f. measurements

Abstract The standard molar Gibbs free energy of formation of sodium chromite ΔfGmo(NaCrO2, s, T) has been determined in the temperature range 820 to 1006 K using a solid-electrolyte galvanic cell of the type: Pt|NaCrO 2 +Na 2 CrO 4 +Cr 2 O 3 |(1−x) ZrO 2 +x CaO|air|Pt where x = 0.15 and p(O2) in air is taken to be 21.21 kPa. The measured e.m.f.s. could be represented by (E/ mV )±2.8=605.8−0.4696(T/ K ) Using this equation and the reported standard molar Gibbs free energies of vormation of Na2CrO4 and Cr2O3, the standard molar Gibbs free energy of formation of NaCrO2 was calculated and is given by Δ f G° m ( NaCrO 2 , s,T ) ( kJ · mol −1 )±2.5=−856.1+0.1693(T/K)

Calorimetric studies on uranium molybdate

Abstract Enthalpy increment measurements on UMoO6(s) have been carried out using a high-temperature Calvet micro-calorimeter in the temperature range 299 to 1000 K. The enthalpy increments were least squares analyzed using Shomate’s method. The complete thermodynamic information for UMoO6(s) has been computed. The enthalpy increment expression for UMoO6(s) as a function of temperature is given by H o (T)−H o (298.15 K )( J mol −1 )=−53928.8+158.65T( K )+21.443×10 −3 T 2 ( K )+14.077×10 5 /T( K ).

Standard molar Gibbs free energy of formation of Na2ZrO3(s)

Abstract The standard molar Gibbs free energy of formation of Na2ZrO3(s) has been obtained by measuring equilibrium pressures of CO2 for the reaction: Na2CO3(s) + ZrO2(s) = Na2ZrO3(s) + CO2(g) by a static manometric method in the temperature range 878 to 1107 K. The CO2 pressures were least-squares analysed and can be represented by: lg ( p kPa ) = 7.470 − 7385.2( K T )±0.038, (878 to 1107 K ) . ΔfGmoNa2ZrO3(s) has been evaluated using ΔfGmos of CO2(g), ZrO2(s), and Na2CO3(s) and can be represented by Δ f G m o (Na 2 ZrO 3 , s , T) ( kJ · mol −1 ) = −1676.27+0.348( T K )±1.1 . Second- and third-law analyses yielded ΔfHmo(Na2ZrO3, s, 298.15 K) as −(1662.9±7.6) and −(1654.9±8.02) kJ · mol−1, respectively.

Thermodynamic studies on SrThO3(s)

Abstract The Gibbs energy of formation of SrThO3(s) has been determined using e.m.f. and manometric techniques. In the e.m.f. method, two fluoride cells have been constructed to determine ΔfG0m(SrThO3,s,T) using CaF2(s) as a solid electrolyte. The cells used are: (−) O 2 ( g ), Pt / SrO ( s )+ SrF 2 ( s )// CaF 2 //( SrThO 3 ( s )+ ThO 2 ( s )+ SrF 2 ( s )/ Pt , O 2 ( g )(+), ( I ) (−) O 2 ( g ), Pt / SrThO 3 ( s )+ SrF 2 ( s )+ ThO 2 ( s )// CaF 2 ( s )// CaO ( s )+ CaF 2 ( s )/ Pt , O 2 ( g )(+). ( II ) The observed e.m.f. values are represented by following respective expressions: E ( V )±0.0001=0.0998+3.254×10 −5 T ( K ), ( Cell I ) E ( V )±0.0001=0.0285−6.37×10 −5 T ( K ). ( Cell II ) From the measured e.m.f. values of the cells and the ΔfG0m(T) values from the literature, ΔfG0m(SrThO3,s,T) have been calculated and are respectively given as Δ f G 0 m ( SrThO 3 , s ,T)±10 kJ mol −1 =−1829.2+0.2735T ( K ) (978⩽T ( K )⩽1154), ( Cell I ) Δ f G 0 ( SrThO 3 , s ,T)±20 kJ mol −1 =−1853.5+0.2867T ( K ) (1008⩽T ( K )⩽1168). ( Cell II ) In the manometric technique, equilibrium CO2(g) pressures are measured over the three phase mixture: {SrThO3(s)+SrCO3(s)+ThO2(s)} using a mercury manometer from 1075 to 1197 K. The corresponding Gibbs energy as a function of temperature is given by Δ f G 0 m ( SrThO 3 , s ,T)( kJ mol −1 )±14=−1865.4+0.3086T ( K ).

Enthalpy increment measurements of SrMoO4(s) and BaMoO4(s)

Abstract Enthalpy increment measurements on SrMoO 4 (s) and BaMoO 4 (s) were carried out using a Calvet micro-calorimeter. The enthalpy increment values were least squares analyzed with the constraints that H o ( T )− H o (298.15 K) at 298.15 K is equal to zero and C p o (298.15 K) is equal to the known value. The dependence of enthalpy increments with temperature can be given as: H o (T)−H o (298.15 K )( J mol −1 )=126.810 T( K )+24.571×10 −3 T 2 ( K )+21.782×10 5 /T( K )−47299. ( SrMoO 4 ( s ), 299.0≤T( K )≤1020.3) H o (T)−H o (298.15 K )( J mol −1 )=136.280T( K )+15.122×10 −3 T 2 ( K )+20.502×10 5 /T( K )−48852. ( BaMoO 4 ( s ), 299.0≤T( K )≤1020.3) Thermodynamic functions of SrMoO 4 (s) and BaMoO 4 (s) have been generated using the experimental Δ f H m o (298.15 K), Δ f G m o ( T ) and estimated S m o (298.15 K) values for SrMoO 4 (s) and BaMoO 4 (s) from the literature.

Calorimetric studies on Tl2Th(MoO4)3(s) and Tl2UO2(MoO4)2(s)

Abstract Enthalpy increment measurements on Tl2Th(MoO4)3(s) and Tl2UO2(MoO4)2(s) have been carried out using a high temperature Calvet microcalorimeter. Both compounds have shown transition in the experimental temperature range. The enthalpy increment data were least squares analyzed using Shomate’s method (Ttra>T) and origin programme version 6 (T>Ttra). In the Shomate method, the constraints used are Hom(T)−Hom (298.15 K) is zero at 298.15 K and Cop,m (298.15 K) equal to 364.6 and 269.44 J K−1 mol−1 for Tl2Th(MoO4)3(s) and Tl2UO2(MoO4)2(s), respectively. The resulting expressions are given below: H o m (T)−H o m (298.15 K) (J mol −1 )=−163 137+ 447.26T+19.286 ·10 −3 T 2 +83.669·10 5 /T (Tl 2 Th(MoO 4 ) 3 (s) , 299≤T≤500.6) H o m (T)−H o m (298.15 K) (J mol −1 ) =−86 844+237.13T+96.656·10 −3 T 2 +22.516 ·10 5 /T (Tl 2 UO 2 (MoO 4 ) 2 (s) , 299≤T≤496.5) Hom(T)−Hom (298.15 K) data, above the transition temperatures, were least squares analyzed using the origin program version 6 in which only the experimental Hom(T)−Hom (298.15 K) data points are taken. The corresponding expressions for Tl2Th(MoO4)3(s) and Tl2UO2(MoO4)2(s), respectively, are H o m (T)−H o m (298.15 K) (J mol −1 ) =−189 283+468.232T+19.350·10 −3 T 2 +18.878 ·10 6 /T (Tl 2 Th(MoO 4 ) 3 (s) , 775.5≤T≤828.0) H o m (T)−H o m (298.15 K) (J mol −1 ) =−78 811+235.391T+96.420·10 −3 T 2 +2.742·10 4 /T (Tl 2 UO 2 (MoO 4 ) 2 (s) , 626≤T≤835.4) The first differential of above equations with respect to temperature gives the molar heat capacities, Cop,m(T). The complete thermodynamic functions for the Tl2Th(MoO4)3(s) and Tl2UO2(MoO4)2(s) compounds have been computed for the first time.

Thermodynamics of the vaporisation of thorium tetrachloride

Abstract Vapour pressures of solid and liquid ThCl4 have been measured in the temperature range 880 to 1037 K and 1045 to 1160 K both by transpiration and evaporation-temperature techniques. The vapour pressures obtained from the two methods agree well and so were combined to get the vapour-pressure equations for solid and liquid ThCl4. Melting temperature, evaporation temperature, and enthalpy of fusion at 1043 K for ThCl4 were evaluated to be 1043 K, 1205 K, and (14.84±0.89) kcalth mol−1 respectively. The values of the enthalpy of vaporization ΔHo(298.15 K) calculated from the vapour pressures for solid ThCl4 by second-law and third-law methods are (57.20±0.38) kcalth mol−1 and (52.97±0.34) kcalth mol−1.

The vaporization thermodynamics of uranium tetrachloride

Vapour pressures of solid and of liquid uranium tetrachloride have been measured in the temperature ranges 763 to 862 K and 868 to 971 K respectively using a transpiration technique. The vapour pressures have been compared with data from the literature. The vapour pressure of solid and of liquid UCl4 can be represented by equations: log10(p/atm)=(10.427±0.101)−(10412±82)(K/T), (763 to 862 K) , and log10(p/atm)=(7.245±0.133)−(7649±121)(K/T), (868 to 971 K) , respectively. The standard enthalpies of vaporization ΔHo(298.15 K) obtained from vapour pressures of solid and of liquid UCl4 are (50.69 ± 0.36) and (51.01 ± 0.54) kcalth mol−1.