Molar Heat Capacities Cp Research Articles

Molar heat capacity and thermodynamic properties of 1,2-cyclohexane dicarboxylic anhydride [C8H10O3

The molar heat capacity C p,m of 1,2-cyclohexane dicarboxylic anhydride was measured in the temperature range from T=80 to 390 K with a small sample automated adiabatic calorimeter. The melting point T m, the molar enthalpy Δfus H m and the entropy Δfus S m of fusion for the compound were determined to be 303.80 K, 14.71 kJ mol−1 and 48.43 J K−1 mol−1, respectively. The thermodynamic functions [H T-H 273.15] and [S T-S 273.15] were derived in the temperature range from T=80 to 385 K with temperature interval of 5 K. The thermal stability of the compound was investigated by differential scanning calorimeter (DSC) and thermogravimetry (TG), when the process of the mass-loss was due to the evaporation, instead of its thermal decomposition.

Apparent molar volumes and apparent molar heat capacities of aqueous tetrahydrofuran, dimethyl sulfoxide, 1,4-dioxane, and 1,2-dimethoxyethane at temperatures from 278.15 K to 393.15 K and at the pressure 0.35 MPa

We determined apparent molar volumes Vϕ at 278.15⩽(T/K)⩽368.15 and apparent molar heat capacities Cp,ϕ at 278.15⩽(T/K)⩽393.15 at p=0.35MPa for aqueous solutions of tetrahydrofuran at m from (0.016 to 2.5)mol·kg−1, dimethyl sulfoxide at m from (0.02 to 3.0)mol·kg−1, 1,4-dioxane at m from (0.015 to 2.0)mol·kg−1, and 1,2-dimethoxyethane at m from (0.01 to 2.0)mol·kg−1. Values of Vϕ were determined from densities measured with a vibrating-tube densimeter, and values of Cp,ϕ were determined with a twin fixed-cell, differential, temperature-scanning calorimeter. Empirical functions of m and T for each compound were fitted to our Vϕ and Cp,ϕ results.

Molar Heat Capacities of Aminouracils by Differential Scanning Calorimetry

The molar heat capacities Cp,m of four crystalline compounds (5-aminouracil, 6-aminouracil, 6-amino-1-methyluracil, and 6-amino-1,3-dimethyluracil) were determined using a differential scanning calorimeter (SETARAM TG-DSC 111). The measurements were made at constant pressure between (298.15 and 343.15) K. Then the molar heat capacities were correlated with temperature. The contributions of C-CH3, N-CH3, and C-NH2 groups to the standard molar heat capacity values were evaluated. The experimentally determined values were compared with those calculated by the Chickos method.

Heat Capacities of Uracil, Thymine, and Its Alkylated, Cyclooligomethylenated, and Halogenated Derivatives by Differential Calorimetry

The molar heat capacity (Cp) of solid uracil, its alkylated and halogenated derivatives, and cyclooligomethyl-enouracils in the temperature range of (298.15 to 343.15) K by a differential scanning calorimeter (SETARAM TG-DSC 111) were determined. It was demonstrated that the Cp value increases by increasing the number of methylene groups attached to the diketopyrimidine ring. The correlations Cp = f(T) are given. The contributions of C-CH 3 , N-CH 3 , and C-NO 2 groups as well as F, Cl, Br, and I atoms in the value of Cp for the temperature of 298.15 K are presented. The Chickos method for calculation of Cp values is discussed.

Thermodynamic studies of (R)-BINOL-menthyl dicarbonates

The (R)-BINOL-menthyl dicarbonates, one of the most important compounds in catalytic asymmetric synthesis, was synthesized by a convenient method. The molar heat capacities Cp,m of the compound were measured over the temperature range from 80 to 378 K with a small sample automated adiabatic calorimeter. Thermodynamic functions [HT–H298.15] and [ST–S298.15] were derived in the above temperature range with a temperature interval of 5 K. The thermal stability of the substance was investigated by differential scanning calorimeter (DSC) and a thermogravimetric (TG) technique.

Molar Heat Capacity and Thermodynamic Properties of 4-Methyl-4-cyclohexene-1,2-dicarboxylic Anhydride [C9H10O3

The molar heat capacity Cp,m of 4-methyl-4-cyclohexene-1,2-dicarboxylic anhydride was measured over the temperature range from T = 80 K to T = 361 K with a small-sample automated adiabatic calorimeter. The melting point Tm and the molar enthalpy ΔfusHm and entropy ΔfusSm of fusion for the compound were determined to be (335.54±0.41) K, (17.67±0.04) kJ·mol-1, and (52.65±0.04) J·K-1·mol-1, respectively. Thermodynamic functions [HT − H298.15] and [ST − S298.15] were derived in the temperature range from T = 80 K to T = 361 K with a temperature interval of 5 K. The purity of the sample in mole fraction used in the adiabatic calorimetry study was determined to be 0.9940 by using the fractional melting technique. The thermal stability of the compound was investigated by differential scanning calorimetry (DSC) and a thermogravimetric (TG) technique, and the process of the mass loss of the sample was due to evaporation instead of thermal decomposition.

Thermodynamic Properties of Sr2RuO4(s) and Sr3Ru2O7(s) by Using Solid‐State Electrochemical Cells.

AbstractFor Abstract see ChemInform Abstract in Full Text.

Low-temperature heat capacities and standard molar enthalpy of formation of aspirin

Molar heat capacities (Cp,m) of aspirin were precisely measured with a small sample precision automated adiabatic calorimeter over the temperature range from 78 to 383 K. No phase transition was observed in this temperature region. The polynomial function of Cp,mvs. T was established in the light of the low-temperature heat capacity measurements and least square fitting method. The corresponding function is as follows: for 78 K≤T≤383 K, Cp,m/J mol-1 K-1=19.086X4+15.951X3-5.2548X2+90.192X+176.65, [X=(T-230.50/152.5)]. The thermodynamic functions on the base of the reference temperature of 298.15 K, {ΔHT -ΔH298.15} and {ST-S298.15}, were derived. Combustion energy of aspirin (ΔcUm) was determined by static bomb combustion calorimeter. Enthalpy of combustion (ΔcHom) and enthalpy of formation (ΔfHom) were derived through ΔcUm as - (3945.26±2.63) kJ mol-1 and - (736.41±1.30) kJ mol-1, respectively.

Calorimetric investigation of excess molar heat capacities for water + ethylene glycol from T = 273.15 to T = 373.15 K

The molar excess heat capacity Cp, mE of(1 −x)H2O +xHOCH2CH2OH has been measured in the temperature range from 273.15 K to 373.15 K at 5 K intervals. The Cp, mE values are negative and positive at all compositions between temperatures of (273.15, 283.15, 333.15 and 373.15) K, respectively. BetweenT= 288.15 K and T= 328.15 K, Cp,mE values are positive for the mixtures rich in water and negative for the mixtures rich in ethylene glycol with maxima at values of x about 0.07, 0.3 and with minima at values of x about 0.6. The curve ofCp, mE against x at T= 303.15 K is unusual in showing three maximum values at values of x at about 0.07, 0.3, and 0.4, respectively. The curves ofCp, mE against T show that the effects of T onCp, mE at x= 0.03124, 0.5372, and 0.7232 are more complex. The shapes of the curves of Cp, mE against x and their changes with temperature can be accounted for by the structure of the binary system.

Thermodynamics of aqueous amines: excess molar heat capacities, volumes, and expansibilities of {water + methyldiethanolamine (MDEA)} and {water + 2-amino-2-methyl-1-propanol (AMP)}

The heat capacities of aqueous solutions of methyldiethanolamine (MDEA) and 2-amino-2-methyl-1-propanol (AMP) were measured in a heat-flux differential scanning calorimeter at temperatures 278.15K⩽T⩽ 368.15 K, to yield excess molar heat capacities Cp,mEover the whole mole fraction range. Densities of aqueous AMP were determined with a vibrating tube densimeter over the temperature range 293.15 K⩽T⩽ 353.15 K, and used to derive excess molar volumesVmE and excess molar thermal expansibilitiesEmE= (∂VmE/∂T)p. A modified Redlich–Kister treatment was employed to describe the excess functions. The standard partial molar propertiesCp,2o and V2oatT= 298.15 K for MDEA, as determined by extrapolating the Redlich–Kister expressions to infinite dilution, are consistent with results from a Picker calorimeter and densimeter at mole fractionsx2 (MDEA) ⩽ 0.03. Values ofCp,2o and V2ofor AMP from measurements at low mole fractions { x2(AMP) ⩽ 0.03} differ from the Redlich–Kister extrapolations, suggesting that specific interactions take place in dilute solutions. The results are consistent with Lumry’s model (Faraday Symp. Chem. Soc. 1982, 17, 93–108), with MDEA being less hydrophobic than AMP. The mole-fraction dependence of the reduced excess heat capacities is most pronounced at T< 298.15 K, with clear differences in the behaviour of Cp,mE(water+MDEA) and Cp,mE(water+AMP), while the behaviour ofVmE for the two aqueous alkanolamines is similar.

Apparent molar volumes and apparent molar heat capacities of aqueous ethane-1,2-diol, propane-1,2-diol, and propane-1,3-diol at temperatures from 278.15 K to 393.15 K and at the pressure 0.35 MPa

Apparent molar heat capacities Cp, φand apparent molar volumesVφ were determined for aqueous solutions of ethane-1,2-diol, propane-1,2-diol, and propane-1,3-diol at temperatures T from 278.15 K to 393.15 K and at the pressure 0.35 MPa. The molalities m investigated ranged from 1.0 mol · kg −1to about 0.02 mol · kg −1. We used a vibrating tube densimeter (DMA 512P, Anton PAAR, Austria) to determine the densities and volumetric properties. Heat capacities were obtained using a twin fixed-cell, power-compensation, differential-output, temperature-scanning calorimeter (NanoDSC 6100, Calorimetry Sciences Corporation, American Fork, UT, U.S.A.). The results were fit by regression to equations that describe the surfaces ( Vφ, T, m) and (Cp, φ , T, m). Infinite dilution partial molar volumesV2o and heat capacities Cp,2owere obtained over the range of temperatures by extrapolation of these surfaces to m= 0 mol · kg−1 .

Thermodynamic properties of the natural zeolite gmelinite at low temperatures

The heat capacity of the natural zeolite gmelinite Na1.39K0.01Ca0.21Fe0.01Al1.83Si4.17O12· 5.87H2O was investigated from (5.8 to 302.2) K. The measurements were carried out using a small calorimeter cell of internal volume of 2 cm 3sealed with a threaded lid. To test the accuracy of the calorimetric system, the heat capacity of benzoic acid was measured from (5.6 to 312) K and compared with reliable reference data. Experimental data on the molar heat capacity Cp, mof gmelinite are listed, together with the tabulated thermodynamic functions. Anomalous behavior of the Cp, mfunction was found at T> 160 K. It is suggested that the additional contribution to the heat capacity is a result of the disordering in the (water, cations) subsystem of gmelinite.

Apparent and partial molar heat capacities and volumes of the amino acids [formula omitted] -lysine monohydrochloride and[formula omitted] -arginine monohydrochloride in aqueous solution at temperatures from T = 288.15 K to T = 328.15 K

Relative densities and massic heat capacities have been determined for the amino acidsl -lysine monohydrochloride and l-arginine monohydrochloride in aqueous solution at the temperatures T= (288.15, 298.15, 313.15, and 328.15) K. These data have been used to calculate apparent molar properties which, in turn, were analyzed using extended Debye–Hückel equations to yield partial molar heat capacities Cp, m, φ∞and partial molar volumes Vm, φ∞at infinite dilution. The partial molar properties have been used to calculate amino acid side-chain contributions which have been compared to those values previously reported in the literature.

Apparent molar volumes and apparent molar heat capacities of aqueous solutions ofα- andβ-cyclodextrins at temperatures from 278.15 K to 393.15 K and at the pressure 0.35 MPa

Apparent molar heat capacities Cp, φand apparent molar volumesVφ were determined for aqueous solutions of α - and β -cyclodextrins at temperatures from 278.15 K to 393.15 K and at the pressure 0.35 MPa. The molalities investigated ranged from 0.008 mol · kg−1to 0.12 mol · kg−1forα -cyclodextrin and from 0.004 mol · kg−1to 0.014 mol · kg−1for β -cyclodextrin. We used a vibrating-tube densimeter (DMA 512P, Anton PAAR, Austria) to determine the densities and volumetric properties. Heat capacities were obtained using a twin fixed-cell, power-compensation, differential-output, temperature-scanning calorimeter (NanoDSC 6100, Calorimetry Sciences Corporation, Spanish Fork, UT, USA). Equations were fit by regression to our experimental (Vφ, T, m) and (Cp, φ,T , m) results. Infinite dilution partial molar volumes V2oand heat capacities Cp,2owere obtained over the range of temperatures by extrapolation of these surfaces to m= 0.

Thermodynamics for ionization of water at temperatures from 278.15 K to 393.15 K and at the pressure 0.35 MPa: apparent molar volumes of aqueous KCl, KOH, and NaOH and apparent molar heat capacities of aqueous HCl, KCl, KOH, and NaOH

Apparent molar volumes Vφof aqueous KCl, KOH, and NaOH and apparent molar heat capacities Cp, φof aqueous HCl, KCl, KOH, and NaOH have been determined at the pressure p= 0.35 MPa, and at molalities 0.015 ⩽m/mol · kg−1⩽ 0.5. Densities were measured using a vibrating-tube densimeter (DMA 512, Anton Paar, Austria) at temperatures 278.15 ⩽T/K⩽ 368.15. These values were used to calculate the apparent molar volumes. A fixed-cell, differential-output, power-compensating, temperature-scanning calorimeter (NanoDSC model 6100, Calorimetry Sciences Corporation, Spanish Fork, UT, U.S.A.) was used to measure the heat capacities of the same solutions at temperatures 278.15 ⩽T/K⩽ 393.15. Results were fitted by using equations that describe the surfaces (m, T, Vφ) and (m, T, Cp, φ). Using these equations, we have calculated the surfaces (m, T, ΔrVm), (m, T, ΔrCp, m), (m, T, ΔrHm), (m,T , p Qa), and (m, T,ΔrSm ) for the ionization of water in the presence of combinations of the above electrolytes. The last three surfaces were calculated by integration using our (m,T , ΔrCp, m) surface and literature values for the molality dependence of ΔrHmand pQa at T= 298.15 K.

Apparent molar volumes and apparent molar heat capacities of aqueous solutions ofN,N- dimethylformamide andN,N- dimethylacetamide at temperatures from 278.15 to 393.15 K and at the pressure 0.35 MPa

Apparent molar heat capacities Cp, φand apparent molar volumesVφ were determined for aqueous solutions of N, N - dimethylformamide andN , N - dimethylacetamide at temperatures from 278.15 to 393.15 K and at the pressure 0.35 MPa. The molalities investigated ranged from 0.015 mol ·kg−1to 1.0 mol · kg−1. We used a vibrating-tube densimeter (DMA 512P, Anton PAAR, Austria) to determine the densities and volumetric properties. Heat capacities were obtained using a twin fixed-cell, power-compensation, differential-output, temperature-scanning calorimeter (NanoDSC 6100, Calorimetry Sciences Corporation, Spanish Fork, UT, U.S.A.). The results were fit by regression to equations that describe the surfaces (Vφ,T , m) and (Cp, φ, T, m). Infinite dilution partial molar volumes V2oand heat capacitiesCp,2o were obtained over the range of temperatures by extrapolation of these surfaces to m= 0.

(Solid + solute) phase equilibria in aqueous solution. XIII. Thermodynamic properties of hydrozincite and predominance diagrams for (Zn2 + + H2O + CO2)

The thermodynamic properties ofZn5(OH)6(CO3)2 , hydrozincite, have been determined by performing solubility and d.s.c. measurements. The solubility constant in aqueous NaClO4media has been measured at temperatures ranging from 288.15 K to 338.15 K at constant ionic strength (I= 1.00 mol · kg−1). Additionally, the dependence of the solubility constant on the ionic strength has been investigated up to I= 3.00 mol · kg−1NaClO4at T= 298.15 K. The standard molar heat capacity Cp, mofunction fromT= 318.15 K to T= 418.15 K, as well as the heat of decomposition of hydrozincite, have been obtained from d.s.c. measurements. All experimental results have been simultaneously evaluated by means of the optimization routine of ChemSage yielding an internally consistent set of thermodynamic data (T= 298.15 K): solubility constant log*Kps 00= (9.0 ± 0.1), standard molar Gibbs energy of formationΔfGmo {Zn5(OH)6(CO3)2 } = (−3164.6 ± 3.0)kJ · mol−1, standard molar enthalpy of formation ΔfHmo{Zn5(OH)6(CO3)2 } = (−3584 ± 15)kJ · mol−1, standard molar entropy Smo{Zn5(OH)6(CO3)2 } = (436 ± 50)J · mol−1· K−1and Cp,mo/(J · mol−1· K−1) = (119 ± 11) + (0.834 ± 0.033)T/K. A three-dimensional predominance diagram is introduced which allows a comprehensive thermodynamic interpretation of phase relations in(Zn2++ H2O + CO2) . The axes of this phase diagram correspond to the potential quantities: temperature, partial pressure of carbon dioxide and pH of the aqueous solution. Moreover, it is shown how the stoichiometric composition{n(CO3)/n(Zn)} of the solid compoundsZnCO3 and Zn5(OH)6(CO3)2can be checked by thermodynamically analysing the measured solubility data.

Apparent molar heat capacities and apparent molar volumes ofY2(SO4)3(aq),La2(SO4)3(aq),Pr2(SO4)3(aq),Nd2(SO4)3(aq),Eu2(SO4)3(aq),Dy2(SO4)3(aq),Ho2(SO4)3(aq), andLu2(SO4)3(aq)atT = 298.15 K andp = 0.1 MPa

The apparent molar heat capacities Cp, φ∘and apparent molar volumes Vφ∘of Y2(SO4)3(aq), La2(SO4)3(aq), Pr2(SO4)3(aq), Nd2(SO4)3(aq), Eu2(SO4)3(aq), Dy2(SO4)3(aq), Ho2(SO4)3(aq), and Lu2(SO4)3(aq) were measured at T= 298.15 K and p= 0.1 MPa with a Sodev (Picker) flow microcalorimeter and a Sodev vibrating-tube densimeter, respectively. These measurements extend from lower molalities of m= (0.005 to 0.018) mol ·kg−1to m= (0.025 to 0.434) mol ·kg−1, where the upper molality limits are slightly below those of the saturated solutions. There are no previously published apparent molar heat capacities for these systems, and only limited apparent molar volume information. Considerable amounts of the R SO4+(aq) and R(SO4)2−(aq) complexes are present, where R denotes a rare-earth, which complicates the interpretation of these thermodynamic quantities. Values of the ionic molar heat capacities and ionic molar volumes of these complexes at infinite dilution are derived from the experimental information, but the calculations are necessarily quite approximate because of the need to estimate ionic activity coefficients and other thermodynamic quantities. Nevertheless, the derived standard ionic molar properties for the various R SO4+(aq) and R(SO4)2−(aq) complexes are probably realistic approximations to the actual values. Comparisons indicate that Vφ∘{RSO4+, aq, 298.15K} =−(6 ± 4)cm3· mol−1and Vφ∘{R(SO4)2−, aq, 298.15K} = (35 ± 3)cm3· mol−1, with no significant variation with rare-earth. In contrast, values of Cp, φ∘{ RSO4+, aq, 298.15K } generally increase with the atomic number of the rare-earth, whereas Cp, φ∘{ R(SO4)2−, aq, 298.15K } shows a less regular trend, although its values are always positive and tend to be larger for the heavier than for the light rare earths.

Thermodynamic properties of tetraphenylphosphonium perchlorate and tetraphenylarsonium perchlorate fromT → 0 toT = 340 K

The temperature dependence of the standard molar heat capacity Cp, moof samples of crystalline tetraphenylphosphonium perchlorate and tetraphenylarsonium perchlorate was measured in an adiabatic low-pressure calorimeter between T= 4.8 K and T= 340 K and from T= 5.8 K to T= 340 K, respectively, mostly to within a precision of 0.2 per cent. For tetraphenylphosphonium perchlorate, an anomalous change of the heat capacity in the range T= 125 K to T= 185 K, probably arising from the excitation of hindered rotations of atomic groups, was found and its thermodynamic characteristics were determined. No such anomaly was observed for tetraphenylarsonium perchlorate. The data obtained were used to calculate the thermodynamic functions Cp, mo(T)/R, Δ0THmo/R·K, Δ0TSmo/R, and Φmo=Δ0TSmo−Δ0THmo/T(where R is the universal gas constant) of the compounds between T→ 0 and T= 340 K.

Apparent molar volumes and apparent molar heat capacities of aqueous solutions of isomeric butanols at temperatures from 278.15 K to 393.15 K and at the pressure 0.35 MPa

Apparent molar heat capacities Cp, φand apparent molar volumesVφ were determined for aqueous solutions of 1-butanol, 2-butanol (both R andS isomers), isobutanol (2-methyl-1-propanol), and t -butanol (2-methyl-2-propanol) at temperatures from 278.15 K to 393.15 K and at the pressure 0.35 MPa. The molalities investigated ranged from 0.02 mol · kg−1to 0.5 mol · kg−1. We used a vibrating-tube densimeter (DMA 512P, Anton Paar, Austria) to determine the densities and volumetric properties. Heat capacities were obtained using a twin fixed-cell, power-compensation, differential-output, temperature-scanning calorimeter (NanoDSC 6100, Calorimetry Sciences Corporation, Provo, UT, U.S.A.). The results were fit by regression to equations that describe the surfaces (Vφ, T, m) and (Cp,φ, T, m). Infinite dilution partial molar volumesV2o and heat capacities Cp,2owere obtained over the range of temperatures by extrapolation of these surfaces to m= 0.