The influence of fluoride ions on the equilibrium between titanium ions and titanium metal in fused alkali chloride melts

2016 ◽  
Vol 190 ◽  
pp. 421-432 ◽  
Author(s):  
Jianxun Song ◽  
Qiuyu Wang ◽  
Jinyu Wu ◽  
Shuqiang Jiao ◽  
Hongmin Zhu
Keyword(s):  
Equilibrium Constant ◽  
Coordination Compounds ◽  
Least Squares Method ◽  
Equilibrium Constants ◽  
Free Energies ◽  
Stable States ◽  
Alkali Chloride ◽  
Fluoride Ions ◽  
Formation Reaction ◽  
The Right

KF is employed as a source of fluoride ions added to the melt to disclose the influence of fluoride on the disproportionation reactions of titanium ions, 3Ti2+ = 2Ti3+ + Ti, and 4Ti3+ = 3Ti4+ + Ti. The results reveal that the equilibrium transferred to the right direction for the first reaction and the apparent equilibrium constant increased sharply, mainly because of the formation of coordination compounds: TiFi3−i. The accurate values of the equilibrium constants referring to the formation reactions of Ti3+ + iF− = TiFi3−i (i = 1–6) in NaCl–KCl melt at 1023 K were evaluated with a best fit least squares method. It is also revealed that the stable states of the coordination compounds are TiF2+, TiF2+, TiF4− and TiF63−. Moreover, the Gibbs free energies for complex formation were estimated. Ti2+ was undetectable when the concentration of fluoride ion was high enough. The equilibrium constant for the formation reaction, Ti4− + 6F− = TiF62−, was evaluated. The equilibrium constant, Kc2, for the disproportionation reaction 4Ti3+ = 3Ti4+ + Ti, in chloride melt was determined as 0.015.

The retroaldol reaction of cinnamaldehyde

1984 ◽  
Vol 62 (8) ◽  
pp. 1441-1445 ◽  
Author(s):  
J. Peter Guthrie ◽  
Kevin J. Cooper ◽  
John Cossar ◽  
Brian A. Dawson ◽  
Kathleen F. Taylor
Keyword(s):  
Equilibrium Constant ◽  
Rate Constants ◽  
Equilibrium Constants ◽  
Alkaline Solutions ◽  
Formation Reaction

Rate and equilibrium constants have been measured for the hydration and retroaldol reactions of cinnamaldehyde. The equilibrium constant for the 1,4-addition of water to cinnamaldehyde is 4.42 × 10−3. The rate constants for hydroxide catalyzed reaction, extrapolated to zero hydroxide concentration (to correct for the addition of hydroxide to the aldol carbonyl), are: [Formula: see text];[Formula: see text]; and [Formula: see text]. The rate of the formation reaction was measured by adding small amounts of acetaldehyde to alkaline solutions of benzaldehyde: [Formula: see text] and Koverall = 1480 M−1. The course of the synthetically useful reaction of acetaldehyde with benzaldehyde is discussed in the light of these results.

Automatic Reaction Balancing

1996 ◽  
Author(s):  
Craig M. Bethke
Keyword(s):  
Species Composition ◽  
Equilibrium Constant ◽  
Chemical Reactions ◽  
Equilibrium Constants ◽  
Transformation Matrix ◽  
Small Computer ◽  
Free Energies ◽  
Equilibrium Equations ◽  
Charged Species ◽  
The One

Conveniently, perhaps even miraculously, the equations developed in Chapter 4 to accomplish basis swaps can be used to balance chemical reactions automatically. Once the equations have been coded into a computer program, there is no need to balance reactions, compute equilibrium constants, or even determine equilibrium equations by hand. Instead, these procedures can be performed quickly and reliably on a small computer. To balance a reaction, we first choose a species to appear on the reaction’ s left side, and express that species’ composition in terms of a basis B. The basis might be a list of the elements in the species’ stoichiometry, or an arbitrary list of species that combine to form the left-side species. Then we form a second basis B´ composed of species that we want to appear on the reaction’ s right side. To balance the reaction, we calculate the transformation matrix relating basis B´ to B, following the procedures in Chapter 4. The transformation matrix, in turn, gives the balanced reaction and its equilibrium constant. Two methods of balancing reactions are of interest. We can balance reactions in terms of the stoichiometries of the species considered. In this case, the existing basis B is a list of elements and, if charged species are involved, the electron e–. Alternatively, we may use a dataset of balanced reactions, such as the LLNL database. Basis B, in this case, is the one used in the database to write reactions. We will consider each possibility in turn. A straightforward way to balance reactions is to use as the initial basis the stoichiometries of the species involved. If the species’ free energies of formation are known, the reaction’ s equilibrium constant can be determined as well. In the stoichiometric approach, basis B is the list of elements that will appear in the reaction, plus the electron if needed. We write swap reactions and calculate a transformation matrix as described in Section 3.1. The equations in Sections 3.2 and 3.3 give the balanced reaction and associated equilibrium constant.

Effect of the acyl substituent on the equilibrium constant for hydration of esters

1980 ◽  
Vol 58 (13) ◽  
pp. 1281-1294 ◽  
Author(s):  
J. Peter Guthrie ◽  
Patricia A. Cullimore
Keyword(s):  
Aqueous Solution ◽  
Equilibrium Constant ◽  
Rate Constant ◽  
Rate Constants ◽  
Acid Catalysis ◽  
Methyl Esters ◽  
Equilibrium Constants ◽  
Free Energies ◽  
General Acid ◽  
Aldehydes And Ketones

Heats of hydrolysis have been measured for the trimethyl orthoesters of isobutyric, propionic, benzoic, methoxyacetic, chloroacetic, and cyanoacetic acids using aqueous acid with an organic cosolvent where necessary, and of the corresponding esters in alkaline solution. Solubilities or free energies of transfer from gas to aqueous solution have been measured, permitting calculation of the free energies of formation of the aqueous orthoesters, and by methods which we have published previously, calculation of the free energies of formation of the covalent hydrates of the esters, and the free energy changes for hydration of these esters.Using estimated pKa values equilibrium constants were calculated for the addition of hydroxide to the esters. The data are in good agreement with the appropriate Marcus equation relating rate and equilibrium constants with a value for b of 8.99 ± 0.17. This line was used to estimate the equilibrium constant for addition of hydroxide, and thence of water, to some additional esters where only the rate constant was available. Rate constants for hydrolysis of methyl esters in aqueous solution at 25 °C were calculated from literature data, correcting for the effect of other conditions as necessary. From the equilibrium constants for addition of water we could estimate the rate constants for uncatalyzed hydrolysis; for the cases where this rate constant has been measured, the agreement was satisfactory. For acid catalyzed hydrolysis the data permit a test of the two alternative mechanisms considered previously, namely specific acid catalysis and general acid catalysis with hydronium ion acting as a general acid. For esters the mechanism is clearly specific acid catalysis, but for aldehydes and ketones it appears very likely that the mechanism is general acid catalysis.

Chemical equilibrium and chemical kinetics

2018 ◽  
Author(s):  
Dennis Sherwood ◽  
Paul Dalby
Keyword(s):  
Free Energy ◽  
Equilibrium Constant ◽  
Gibbs Free Energy ◽  
Chemical Kinetics ◽  
Chemical Equilibrium ◽  
First Principles ◽  
Effect Of Temperature ◽  
Total System ◽  
Free Energies

Building on the previous chapter, this chapter examines gas phase chemical equilibrium, and the equilibrium constant. This chapter takes a rigorous, yet very clear, ‘first principles’ approach, expressing the total Gibbs free energy of a reaction mixture at any time as the sum of the instantaneous Gibbs free energies of each component, as expressed in terms of the extent-of-reaction. The equilibrium reaction mixture is then defined as the point at which the total system Gibbs free energy is a minimum, from which concepts such as the equilibrium constant emerge. The chapter also explores the temperature dependence of equilibrium, this being one example of Le Chatelier’s principle. Finally, the chapter links thermodynamics to chemical kinetics by showing how the equilibrium constant is the ratio of the forward and backward rate constants. We also introduce the Arrhenius equation, closing with a discussion of the overall effect of temperature on chemical equilibrium.

SOME EQUILIBRIUM CONSTANTS OF TRANSITION METAL HALIDES

1960 ◽  
Vol 38 (10) ◽  
pp. 1827-1836 ◽  
Author(s):  
M. W. Lister ◽  
P. Rosenblum
Keyword(s):  
Ionic Strength ◽  
Transition Metal ◽  
Equilibrium Constant ◽  
Spectrophotometric Method ◽  
Equilibrium Constants ◽  
Cell Voltage ◽  
Metal Halides ◽  
Complex Ions ◽  
Anomalous Variation ◽  
A Cell

Measurements are reported on the formation of complex ions in solutions containing cupric and chloride or bromide ions, and solutions of nickel or cobalt with chloride. In each case the halide was present in very low amount. With copper a spectrophotometric method was used, and a cell voltage method with nickel and cobalt. The ionic strength was kept constant, but the temperature was varied. The data show difficulties of interpretation if it is assumed that only MX+ ions (M is the metal, X is the halogen) are formed, the difficulties arising from the anomalous variation of the equilibrium constant with temperature, and from the general drift of the calculated constants from the e.m.f. measurements. Various explanations are considered and it is shown that postulation of M2X+3 ions is at least a possible explanation.

scholarly journals Binding of coenzymes to yeast alcohol dehydrogenase

2010 ◽  
Vol 75 (2) ◽  
pp. 185-194 ◽  
Author(s):  
Vladimir Leskovac ◽  
Svetlana Trivic ◽  
Draginja Pericin ◽  
Mira Popovic ◽  
Julijan Kandrac
Keyword(s):  
Alcohol Dehydrogenase ◽  
Equilibrium Constants ◽  
Point Mutations ◽  
Free Energies ◽  
Wild Type ◽  
Wild Type Enzyme ◽  
Mutant Type ◽  
Yeast Alcohol Dehydrogenase ◽  
Internal Equilibrium ◽  
Individual Reaction

In this work, the binding of coenzymes to yeast alcohol dehydrogenase (EC 1.1.1.1) were investigated. The main criterions were the change in the standard free energies for individual reaction steps, the internal equilibrium constants and the overall changes in the reaction free energies. The calculations were performed for the wild type enzyme at pH 6-9 and for 15 different mutant type enzymes, with single or double point mutations, at pH 7.3. The abundance of theoretical and experimental data enabled the binding of coenzymes to enzyme to be assessed in depth.

Aqueous nonelectrolyte solutions — Part XX: Formula of structure I methane hydrate, congruent dissociation melting point, and formula of the metastable hydrate

2003 ◽  
Vol 81 (12) ◽  
pp. 1443-1450 ◽  
Author(s):  
David N Glew
Keyword(s):  
Melting Point ◽  
Equilibrium Constant ◽  
Methane Hydrate ◽  
Equilibrium Constants ◽  
Stability Range ◽  
Congruent Melting Point ◽  
Quadruple Point ◽  
Dissociation Equilibrium ◽  
Metastable Structure ◽  
The Stability

Sixteen new measurements of high precision for structure I methane hydrate with water between 31.93 and 47.39 °C are shown to be metastable and exhibit higher methane pressures than found by earlier workers. Comparison of earlier measurements between 26.7 and 47.2 °C permit positive identification of the structure II and the structure I hydrates. Forty-nine equilibrium constants Kp(h1[Formula: see text]l1g) for dissociation of structure I methane hydrate into water and methane, 32 between –0.29 and 26.7 °C for the stable hydrate and 17 between 31.93 and 47.39 °C for the metastable hydrate, are best represented by a three-parameter thermodynamic equation, which indicates a standard error (SE) of 0.63% on a single Kp(h1[Formula: see text]l1g) determination. The congruent dissociation melting point C(h1l1gxm) of metastable structure I methane hydrate is at 47.41 °C with SE 0.02 °C and at pressure 505 MPa. The congruent equilibrium constant Kp(h1[Formula: see text]l1g) is 102.3 MPa with SE 0.2 MPa. ΔH°t(h1[Formula: see text]l1g) is 62 281 J mol–1 with SE 184 J mol–1, and the congruent formula is CH4·5.750H2O with SE 0.059H2O. At the congruent point, ΔV(h1[Formula: see text]l1g) is zero within experimental precision, and its estimate is 1.3 with SE 1.6 cm3 mol–1. The stability range of structure I methane hydrate with water extends from quadruple point Q(s1h1l1g) at –0.29 °C up to quadruple point Q(h1h2l1g) at 26.7 °C, and its metastability range with water extends from 26.7 °C up to the congruent dissociation melting point C(h1l1gxm) at 47.41 °C. Key words: methane hydrate, clathrate structure I, metastability range, dissociation equilibrium constant, formula, congruent melting point, metastability of structure I hydrate.

Absolute gas-phase acidities of CH3NH2, C2H5NH2, (CH3)2 NH, and (CH3)3N

1976 ◽  
Vol 54 (10) ◽  
pp. 1624-1642 ◽  
Author(s):  
Gervase I. Mackay ◽  
Ronald S. Hemsworth ◽  
Diethard K. Bohme
Keyword(s):  
Equilibrium Constant ◽  
Gas Phase ◽  
Equilibrium Constants ◽  
Heats Of Formation ◽  
Molecular Orbital Calculations ◽  
Electron Affinities ◽  
Flowing Afterglow ◽  
Experimental Assessment ◽  
Gas Phase Acidity ◽  
Conjugate Bases

The flowing afterglow technique has been employed in measurements of the rate and equilibrium constants at 296 ± 2 K for reactions of the type[Formula: see text]and[Formula: see text]where R1 and R2 may be H, CH3, or C2H5. The equilibrium constant measurements provided absolute values for the intrinsic (gas-phase) acidities of the Brønsted acids CH3NH2, C2H5NH2, (CH3)2NH, and (CH3)3N, the heats of formation of their conjugate bases, and the electron affinities of the corresponding radicals R1R2N. Proton removal energies, ΔG0298/(kcal mol−1), were determined to be 395.7 ± 0.7 for [Formula: see text] 391.7 ± 0.7 for [Formula: see text] 389.2 ± 0.6 for [Formula: see text] and > 396 for [Formula: see text] Heats of formation, ΔH0f.,298, were determined to be 30.5 ± 1.5 for CH3NH−, 21.2 ± 1.5 for C2H5NH−, and 24.7 ± 1.4 for (CH3)2N−. Electron affinities (in kcal mol−1) were determined to be 13.1 ± 3.5 for CH3NH, 17 ± 4 for C2H5NH, and 14.3 ± 3.4 for (CH3)2N. These results quantify earlier conclusions regarding the intrinsic effects of substituents on the gas-phase acidity of amines and provide an experimental assessment of recent molecular orbital calculations of proton removal energies for alkylamines.

Kinetic and equilibrium studies of fluoride sorption onto surfactant-modified smectites

Clay Minerals ◽  
2012 ◽  
Vol 47 (4) ◽  
pp. 429-440 ◽  
Author(s):  
S. Gamoudi ◽  
N. Frini-Srasra ◽  
E. Srasra
Keyword(s):  
Aqueous Solution ◽  
Thermodynamic Parameters ◽  
Adsorption Kinetics ◽  
Equilibrium Constants ◽  
Second Order ◽  
Operating Conditions ◽  
Polluted Water ◽  
Isotherm Models ◽  
Fluoride Ions ◽  
Pseudo Second Order

AbstractThe use of organoclays as adsorbents in the remediation of polluted water has been the subject of many recent studies. In the present work, a Tunisian smectite modified with two cationic surfactants was used as an adsorbent to examine the adsorption kinetics, isotherms and thermodynamic parameters of fluoride ions from aqueous solution. Various pH values, initial concentrations and temperatures have been tested. Two simplified kinetic models, first-order and pseudo-second-order, were used to predict the adsorption rate constants. It was found that the adsorption kinetics of fluoride onto modified smectites at different operating conditions can best be described by the pseudo-second-order model. Adsorption isotherms and equilibrium adsorption capacities were determined by the fitting of the experimental data to well known isotherm models including those of Langmuir and Freundlich. The results showed that the Langmuir model appears to fit the adsorption better than the Freundlich adsorption model for the adsorption of fluoride ions onto modified smectites. The equilibrium constants were used to calculate thermodynamic parameters, such as the change of free energy, enthalpy and entropy. Results of this study demonstrated the effectiveness and feasibility of organoclays for the removal of fluoride ions from aqueous solution.

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