The parametrization and testing of the OPLS all-atom force field for organic molecules and peptides are described. Parameters for both torsional and nonbonded energetics have been derived, while the bond stretching and angle bending parameters have been adopted mostly from the AMBER all-atom force field. The torsional parameters were determined by fitting to rotational energy profiles obtained from ab initio molecular orbital calculations at the RHF/6-31G*//RHF/6-31G* level for more than 50 organic molecules and ions. The quality of the fits was high with average errors for conformational energies of less than 0.2 kcal/mol. The force-field results for molecular structures are also demonstrated to closely match the ab initio predictions. The nonbonded parameters were developed in conjunction with Monte Carlo statistical mechanics simulations by computing thermodynamic and structural properties for 34 pure organic liquids including alkanes, alkenes, alcohols, ethers, acetals, thiols, sulfides, disulfides, a...

https://sakai.rutgers.edu/access/content/user/kparis/biomaps_513_references/02_B_01_JACS118_11225_OPLS.pdf

Development and Testing of the OPLS All-Atom Force Field on Conformational Energetics and Properties of Organic Liquids

Successive parameterizations of the GROMOS force field have been used successfully to simulate biomolecular systems over a long period of time. The continuing expansion of computational power with time makes it possible to compute ever more properties for an increasing variety of molecular systems with greater precision. This has led to recurrent parameterizations of the GROMOS force field all aimed at achieving better agreement with experimental data. Here we report the results of the latest, extensive reparameterization of the GROMOS force field. In contrast to the parameterization of other biomolecular force fields, this parameterization of the GROMOS force field is based primarily on reproducing the free enthalpies of hydration and apolar solvation for a range of compounds. This approach was chosen because the relative free enthalpy of solvation between polar and apolar environments is a key property in many biomolecular processes of interest, such as protein folding, biomolecular association, membrane formation, and transport over membranes. The newest parameter sets, 53A5 and 53A6, were optimized by first fitting to reproduce the thermodynamic properties of pure liquids of a range of small polar molecules and the solvation free enthalpies of amino acid analogs in cyclohexane (53A5). The partial charges were then adjusted to reproduce the hydration free enthalpies in water (53A6). Both parameter sets are fully documented, and the differences between these and previous parameter sets are discussed.

/pdf/a-biomolecular-force-field-based-on-the-free-enthalpy-of-1chefjehkc.pdf

A biomolecular force field based on the free enthalpy of hydration and solvation: the GROMOS force-field parameter sets 53A5 and 53A6.

In a preceding project, functional forms for “short” Helmholtz energy equations of state for typical nonpolar and weakly polar fluids and for typical polar fluids were developed using simultaneous optimization. In this work, the coefficients of these short forms for the equations of state have been fitted for the fluids acetone, carbon monoxide, carbonyl sulfide, decane, hydrogen sulfide, 2-methylbutane (isopentane), 2,2-dimethylpropane (neopentane), 2-methylpentane (isohexane), krypton, nitrous oxide, nonane, sulfur dioxide, toluene, xenon, hexafluoroethane (R-116), 1,1-dichloro-1-fluoroethane (R-141b), 1-chloro-1,1-difluoroethane (R-142b), octafluoropropane (R-218), 1,1,1,3,3-pentafluoropropane (R-245fa), and fluoromethane (R-41). The 12 coefficients of the equations of state were fitted to substance specific data sets. The results show that simultaneously optimized functional forms can be applied to other fluids out of the same class of fluids for which they were optimized without significant loss of a...

Short Fundamental Equations of State for 20 Industrial Fluids

A valence force field for saturated hydrocarbons

This is a compilation of ionization and appearance potentials of positive ions published from 1955 through June 1966. The compilation lists the ion formed, the parent species from which it was formed, the other products of the process, the threshold energy for the formation of this ion, and the method by which this data was obtained. Where feasible, the heat of formation at 298 K of the positive ion has been computed for each entry using auxiliary thermochemical data. From these computed values “best” values have been chosen.

Ionization Potentials, Appearance Potentials, and Heats of Formation of Gaseous Positive Ions

A sample of furan was purified and used for calorimetric studies and for measurement of vapor pressure, density, refractive index, infrared spectrum (2 to 15..mu..) and mass spectrum. The heat of formation of liquid furan (-..delta..H/sub t/ = 14.90 kcal mole/sup -1/) was derived from heat of combustion measurements. Low temperature thermal studies included measurement of the heat capacity of the solid, transition temperature (150.0/sup 0/K), heat of transition (489.2 cal. mole/sup -1/), triple point (187.55/sup 0/K), heat of fusion (908.8 cal. mole/sup -1/) and heat capacity of the liquid. The entropy of the liquid at 298.16/sup 0/K, calculated from these data is 42.22 cal. deg./sup -1/ mole/sup -1/. The observed vapor pressures between 2 and 62/sup 0/ are represented by the equation log/sub 10/p (mm) = 6.97523 - 1060.851 (t + 227.740). The heat of vaporization at the normal boiling point (31.36/sup 0/) is 6474 cal. mole/sup -1/. The vapor heat capacity, measured at five temperatures between 317 and 488/sup 0/K, is given by the equation C/sup 0/psub p/ = -7.55 + 9.352 x 9.352 x 10/sup -2/T - 5.287 x 10/sup -6/ T/sup 2/ cal. deg./sup -1/ mole/sup -1/. The equation B = -279 - 22.6 exp(950/T) ccmore » mole/sup -1/ was derived for the second virial coefficient from calorimetric data. A vibrational assignment which was partly empirical was made and used with molecular structure data to compute thermodynamic functions of the ideal gas to 1500/sup 0/K. The heat, free energy and equilibrium constant of formation at temperatures to 1500/sup 0/ were calculated using appropriate calorimetric data and thermodynamic functions.« less

Thermodynamic Properties of Furan

The vapor heat capacity of 2-methylbutane was measured at five temperatures between 317 and 487/sup 0/K. The experimental values of C/sub p//sup 0/, the heat capacity in the ideal gas state, may be represented by the equation C/sub p//sup 0/ = -0.01 + 0.10608T - 3.676 x 10/sup -5/T/sup 2/. The heat of vaporization at 279.48, 298.16 and 301.01/sup 0/K. was found to be 6181, 5937 and 5901 cal./mole, respectively. The following empirical equation for the second virial coefficient, B, in the equation of state PV = RT + BP, was obtained from thermal data: B = -577 - 12.01e/sup 1150/T/ cc./mole. Calorimetric entropy and vapor heat capacity data for both 2-methylbutane and 2,3-dimethylbutane were utilized to obtain information about the energy differences between the rotational isomers of these compounds. It was necessary to interpret the spectra and make vibrational assignments for both molecules. It was found that the C/sub 8/ form of 2-methylbutane is at least several thousand cal./mole less stable than the C/sub 1/ form. The energy difference between the rotational isomers of 2,3-dimethylbutane, on the other hand, was found to be very small; spectroscopic data set an upper limit of about 100 cal./mole. Thermodynamic functions were computedmore » by the methods of statistical mechanics for 2-methylbutane and 2,3-dimethylbutane for selected temperatures up to 1500/sup 0/K.« less

Rotational isomerism and thermodynamic functions of 2-methylbutane and 2,3-dimethylbutane. Vapor heat capacity and heat of vaporization of 2-methylbutane

3-Thiapentane: Heat Capacity, Heats of Fusion and Vaporization, Vapor Pressure, Entropy, Heat of Formation and Thermodynamic Functions1

The heat capacity of 2-thiabutane in the solid and liquid states was measured over the temperature range 14 to 298/sup 0/K. The melting point (167.23 = 0.05/sup 0/K.) and heat of fusion (2,333 cal./mole) were determined. The vapor pressure was measured over the temperature range 23 to 101/sup 0/ and the following equation was selected to represent the data: log/sub 10/p = 6.93849 - 1182.562/(t + 224.784). The heat of vaporization was measured at 301.66, 319.76 and 339.81/sup 0/K., and the values found were 7563, 7329 and 7055 cal./mole, respectively. The heat capacity of the vapor was measured at five different temperatures in the range 327 to 487/sup 0/K. The experimental values of C/sub p//sup 0/, the heat capacity of the ideal gas state, may be represented by the equation: C/sub p//sup 0/ = 5.11 + 6.5448 x 10/sup -2/ T - 2.1735 x 10/sup -5/T/sup 2/. An equation for the second virial coefficient, B, in the equation of state PV = RT + BP was obtained from thermal data. This equation is B (cc.) = -257 - 68.13 exp (900/T). The entropy of liquid 2-thiabutane at 298.16/sup 0/K. is 57.14 = 0.10 cal./deg./mole, and the entropy of the vapormore » in the ideal gas state at the normal boiling point, 339.81/sup 0/K., is 82.70 = 0.15 cal./deg./mole. A vibrational assignment was made for 2-thiabutane. Interpretation of the spectroscopic data led to the conclusion, which was confirmed by the thermal data, that the two rotational isomers have nearly the same energy. The heights of the potential barriers hindering internal rotation which were selected to fit the calorimetric data are 1970 cal./mole for rotation about either C-S bond and 3820 cal./mole for rotation about the C-C bond. The thermodynamic functions (H/sub 0//sup 0/ - F/sup 0//sub T/)/T, H/sup 0//sub T/ - H/sub 0//sup 0/, S/sup 0/ and C/sup 0//sub p/ were computed for selected temperatures up to 1000/sup 0/K.« less

Thermodynamic Properties and Rotational Isomerism of 2-Thiabutane1

The thermodynamic properties of ethanethiol have been measured in the solid, liquid and vapor states over the temperature range 14 to 450/sup 0/K. The entropy of the saturated liquid at 298.16/sup 0/K was derived from measurements of the heat capacity of the solid and liquid and the heat of fusion (1189 cal. mole/sup -1/ at the triple point, 125.26 +- 0.05/sup 0/K). Results obtained for the vapor pressure, heat of vaporization (..delta..H/sub vap./), vapor heat capacity in the ideal gaseous state (C), and second virial coefficient (B, in the equation of state, PV = RT + BP) are accurately represented by the empirical equations (1) log P = 6.96206 - 1084.531/(t + 231.385), (P, mm.; t, /sup 0/C.); (2) ..delta..H/sub vap./ = 8151 + 1.369T - 0.02288T/sup 2/, cal. mole/sup -1/ (280 to 308/sup 0/K); (3) C/sup 0//sub p/ = 4.45 + 4.836 x 10/sup -2/ - 1.695 x 10/sup -5/T/sup 2/, cal. deg./sup -1/ mole/sup -1/ (315 to 450/sup 0/K); (4) B = -27 - 60.1 exp(800/T), cc. mole/sup -1/ (280 to 450/sup 0/K). The entropy of the ideal gas at one atmosphere pressure and 298.16/sup 0/K (70.77 +- 0.15 cal. deg./sup -1/ mole/sup -1/) was computed from thesemore » data. A vibrational assignment was made for ethanethiol, based upon information from similar molecules. The heights of the potential barriers to internal rotation selected to fit the experimental thermal data were 3310 cal. mole/sup -1/ for the methyl barrier and 1640 cal. mole/sup -1/ for the thiol barrier. Values of the free energy function, heat content, entropy and heat capacity were computed on the basis of an harmonic-oscillator, rigid-rotator approximation at selected temperatures up to 1000/sup 0/K.« less

K. D. Williamson

Papers

Thermodynamic Properties of Furan

Rotational isomerism and thermodynamic functions of 2-methylbutane and 2,3-dimethylbutane. Vapor heat capacity and heat of vaporization of 2-methylbutane

3-Thiapentane: Heat Capacity, Heats of Fusion and Vaporization, Vapor Pressure, Entropy, Heat of Formation and Thermodynamic Functions1

Thermodynamic Properties and Rotational Isomerism of 2-Thiabutane1

Ethanethiol (Ethyl Mercaptan): Thermodynamic Properties in the Solid, Liquid and Vapor States. Thermodynamic Functions to 1000°K.1