DOCSLIB.ORG

Chapter 4 Calculation of Standard Thermodynamic Properties of Aqueous Electrolytes and Non-Electrolytes

Total Page：16

File Type：pdf, Size：1020Kb

Download full-text PDF Read full-text
Chapter 4 Calculation of Standard Thermodynamic Properties of Aqueous Electrolytes and Non-Electrolytes Chapter 4 Calculation of Standard Thermodynamic Properties of Aqueous Electrolytes and Non-Electrolytes Vladimir Majer Laboratoire de Thermodynamique des Solutions et des Polymères Université Blaise Pascal Clermont II / CNRS 63177 Aubière, France Josef Sedlbauer Department of Chemistry Technical University Liberec 46117 Liberec, Czech Republic Robert H. Wood Department of Chemistry and Biochemistry University of Delaware Newark, DE 19716, USA 4.1 Introduction Thermodynamic modeling is important for understanding and predicting phase and chemical equilibria in industrial and natural aqueous systems at elevated temperatures and pressures. Such systems contain a variety of organic and inorganic solutes ranging from apolar nonelectrolytes to strong electrolytes; temperature and pressure strongly affect speciation of solutes that are encountered in molecular or ionic forms, or as ion pairs or complexes. Properties related to the Gibbs energy, such as thermodynamic equilibrium constants of hydrothermal reactions and activity coefficients of aqueous species, are required for practical use by geologists, power-cycle chemists and process engineers. Derivative properties (enthalpy, heat capacity and volume), which can be obtained from calorimetric and volumetric experiments, are useful in extrapolations when calculating the Gibbs energy at conditions remote from ambient. They also sensitively indicate evolution in molecular interactions with changing temperature and pressure. In this context, models with a sound theoretical basis are indispensable, describing with a limited number of adjustable parameters all thermodynamic functions of an aqueous system over a wide range of temperature and pressure. In thermodynamics of hydrothermal solutions, the unsymmetric standard-state convention is generally used; in this case, the standard thermodynamic properties (STP) of a solute reflect its interaction with the solvent (water), and the excess properties, related to activity coefficients, correspond to solute-solute interactions. For dilute and moderately concentrated solutions, the standard-state functions have a dominant role and can be used as a reasonable approximation for semiquantitative modeling. The solute- solvent interactions particularly prevail at near-critical conditions, where all the STP of solutes undergo rapid variations. As shown in chapter 2, the standard derivative properties of a solute scale with the thermal expansivity and isothermal compressibility of the solvent and diverge at the solvent critical point. The direction of this divergence cannot be unambiguously predicted for certain classes of solutes without experimental evidence. In addition, this extreme behavior strongly affects the properties of solutions in a relatively wide range of conditions below and above the critical point. Thus, the modeling approaches used and the experimental data available at near-ambient conditions, while important in calculations, are 4-1 Calculation of Standard Thermodynamic Properties of Aqueous Electrolytes and Non-Electrolytes not sufficient for extrapolation of STP towards high temperatures and pressures, particularly up to the supercritical region. This chapter reviews the status of approaches to the determination, correlation and prediction of STP of aqueous solutes at superambient conditions. Its organization is in two parts. The thermodynamic background is presented in the first part, which introduces the essential terms regarding STP of aqueous species, shows their relationship to quantities accessible from experiments, and outlines the strategy for building up prediction schemes at superambient conditions. The second part discusses the physico- chemical concepts that have been used for description of STP over the last thirty years. It focuses particularly on practical models that allow calculation of STP for a variety of solutes and that are in some cases available as a software package. Main sources tabulating STP of aqueous electrolytes and nonelectrolytes are listed, and examples of calculations for selected systems are given. The focus of this chapter is on the transfer of a solute from its pure or ideal-gas state to the standard state in aqueous solution. The determination of standard thermodynamic properties of pure substances is beyond the scope of this chapter and is not addressed. Approaches to the thermodynamic description of aqueous systems at different conditions ranging from ambient to supercritical were developed by authors belonging to different communities: physical chemists, chemical engineers, geologists and environmental chemists. Each discipline addresses the issue of aqueous systems from its own perspective, using specific terminology and concepts regarding model formulation, standard-state conventions and concentration scales. The underlying objective of this chapter is to present a synthetic view of the topic with an effort to identify common denominators of various approaches and to unify the description of STP of aqueous solutes. 4.2 Thermodynamic Background 4.2.1 Basic definitions Any thermodynamic function characterizing an aqueous system can be expressed as a linear combination of the partial molar property of water (solvent) X1 and the partial molar properties of dissolved ionic or molecular species (solutes) Xi (i > 1). In physical chemistry of solutions, it is usual to consider the solvent separately from solutes present in the system: n = + X n1 X 1 ∑ ni X i . (4.1) i=2 Each partial molar property can be divided for convenience into a standard-state term and an activity (a) term expressing its variation from the standard state due to changes in concentration and nonideality of the system. Since water is generally present in much greater quantity than other species, it is common practice to treat it differently from solutes. The unsymmetric standard state convention is adopted where the solvent is referenced to its pure state, complying with Raoult’s law, while the solutes are referenced to the state of “infinite dilution,” complying with Henry’s law.1 We will focus first on the partial molar Gibbs energy (chemical potential); all other partial molar thermodynamic properties can be obtained by derivations with respect to temperature and pressure. For water: 1 IUPAC recommends defining standard thermodynamic properties at the temperature of the system and standard pressure of 0.1 MPa. However, when describing aqueous systems over a wide range of conditions, it is more convenient to introduce standard thermodynamic properties that are both temperature and pressure dependent; this approach is adopted in this chapter. 4-2 Calculation of Standard Thermodynamic Properties of Aqueous Electrolytes and Non-Electrolytes = • + = • + γ R → γ R → G1 G1 RT ln a1 G1 RT ln x1 1 x1 1, 1 1, (4.2) • γ R where G1 is the molar Gibbs energy of pure water and x1 and 1 are its mole fraction and Raoult’s activity coefficient, respectively. In the literature, thermodynamic functions of solutes are expressed on three concentration scales. While mole fraction xi is mainly used by physical chemists and process engineers for nonelectrolytes, molality mi (moles of solute per kilogram of solvent) is popular with geochemists, and molarity ci (moles of solute per liter of solution) dominates as a concentration unit in environmental and analytical chemistry. The latter two scales are almost always used for electrolyte concentrations. Since often xi<<x1, and at ambient ≈ ≈ conditions mi ci 55xi , the molality and molarity scales are able to express concentration in magnitudes not far from 1. However, molarity changes with temperature and pressure, which makes this concentration scale impractical at elevated conditions. The partial molar Gibbs energy (chemical potential) of a solute is expressed as µ = = o + = o + γ H → γ H → i Gi Gix RT ln aix Gix RT ln(xi ix ) xi 0, ix 1 = o + = o + γ H → γ H → Gim RT ln aim Gim RT ln(mi / m0 im ) mi 0, im 1 (4.3) = o + = o + γ H → γ H → Gic RT ln aic Gic RT ln(ci / c0 ic ) ci 0, ic 1. o γ H The standard Gibbs energy (standard chemical potential) Gi and the Henry’s activity coefficient i of a solute are specific for a selected concentration scale. Since both activity and activity coefficient are by ⋅ −1 ⋅ −3 definition dimensionless, the introduction of constants m0 = 1 mol kg and c0 = 1 mol dm is necessary (although often neglected) for obtaining dimensionless concentration variables. It is apparent that ideal behavior of the system is attained at infinite dilution of solutes where activity coefficients of both solvent and solutes are unity. The relationship between the standard chemical potential of solute for different concentration scales can be derived by writing Eq. (4.3) in the limit of infinite dilution. After introducing for mi and ci the limiting conversion relations x x ρ = i = i 1 lim mi lim ci , (4.4) x →0 x →0 i M1 i M 1 ρ ⋅ −1 ⋅ −3 where molar mass M1 and density 1 of water are in kg mol and kg dm , respectively, it follows: o = o − = o − ρ Gix Gim RT ln(M 1m0 ) Gic RT ln(M 1c0 / 1 ) (4.5) o o Thus, the difference between Gix and Gim is an additive constant linearly dependent on temperature, o while for conversions where Gic is involved this constant also changes with pressure at conditions of high water compressibility. It should be stressed that the standard chemical potential of a solute must have a finite value and therefore it is not equivalent to the chemical potential of solute at infinite dilution (which is minus infinity as suggested by Eq. (4.3)). For any concentration scale, the standard chemical potential can be expressed generally as 4-3 Calculation of Standard Thermodynamic Properties of Aqueous Electrolytes and Non-Electrolytes G o = lim(G - RT ln (dcv ) ) , (4.6) i → i i x1 1 where dcv stands for a dimensionless concentration variable of a solute (xi, mi/m0 or ci/c0). It means that o Gi corresponds to the behavior of an infinitely dilute solution whose chemical potential is extrapolated to unit concentration (xi=1, mi=m0 or ci=c0). In other words, the standard chemical potential relates to a γ hypothetical solution of unit activity referenced to infinite dilution where i=1.
Load more

Recommended publications
  • Notation for States and Processes, Significance of the Word Standard in Chemical Thermodynamics, and Remarks on Commonly Tabulated Forms of Thermodynamic Functions
    Pure &Appl.Chem.., Vol.54, No.6, pp.1239—1250,1982. 0033—4545/82/061239—12$03.00/0 Printed in Great Britain. Pergamon Press Ltd. ©1982IUPAC INTERNATIONAL UNION OF PURE AND APPLIED CHEMISTRY PHYSICAL CHEMISTRY DIVISION COMMISSION ON THERMODYNAMICS* NOTATION FOR STATES AND PROCESSES, SIGNIFICANCE OF THE WORD STANDARD IN CHEMICAL THERMODYNAMICS, AND REMARKS ON COMMONLY TABULATED FORMS OF THERMODYNAMIC FUNCTIONS (Recommendations 1981) (Appendix No. IV to Manual of Symbols and Terminology for Physicochemical Quantities and Units) Prepared for publication by J. D. COX, National Physical Laboratory, Teddington, UK with the assistance of a Task Group consisting of S. ANGUS, London; G. T. ARMSTRONG, Washington, DC; R. D. FREEMAN, Stiliwater, OK; M. LAFFITTE, Marseille; G. M. SCHNEIDER (Chairman), Bochum; G. SOMSEN, Amsterdam, (from Commission 1.2); and C. B. ALCOCK, Toronto; and P. W. GILLES, Lawrence, KS, (from Commission 11.3) *Membership of the Commission for 1979-81 was as follows: Chairman: M. LAFFITTE (France); Secretary: G. M. SCHNEIDER (FRG); Titular Members: S. ANGUS (UK); G. T. ARMSTRONG (USA); V. A. MEDVEDEV (USSR); Y. TAKAHASHI (Japan); I. WADSO (Sweden); W. ZIELENKIEWICZ (Poland); Associate Members: H. CHIHARA (Japan); J. F. COUNSELL (UK); M. DIAZ-PENA (Spain); P. FRANZOSINI (Italy); R. D. FREEMAN (USA); V. A. LEVITSKII (USSR); 0.SOMSEN (Netherlands); C. E. VANDERZEE (USA); National Representatives: J. PICK (Czechoslovakia); M. T. RATZSCH (GDR). NOTATION FOR STATES AND PROCESSES, SIGNIFICANCE OF THE WORD STANDARD IN CHEMICAL THERMODYNAMICS, AND REMARKS ON COMMONLY TABULATED FORMS OF THERMODYNAMIC FUNCTIONS SECTION 1. INTRODUCTION The main IUPAC Manual of symbols and terminology for physicochemical quantities and units (Ref.
    [Show full text]
  • Chapter 8 Thermodynamic Properties of Mixtures
    Chapter 8 Thermodynamic Properties of Mixtures 2012/3/29 1 Abstract The thermodynamic description of mixtures, extended from pure fluids. The equations of change, i.e., energy and entropy balance, for mixtures are developed. The criteria for phase and chemical equilibrium in mixtures 2012/3/29 2 8.1 THE THERMODYNAMIC DESCRIPTION OF MIXTURES Thermodynamic property for pure fluids, θθ=()TPN , , where N is the number of moles. θθ=()TP , where the number of mole equals to 1. Thermodynamic property for mixtures, θθ=()TPN , ,12 , N ,L , Nc where Ni is the number of moles of the ith component. θθ=()TPx , ,12 , x ,L , xci where x is the mole fraction of the ith component. For example UUTPNN=() , ,12 , ,LL , Ncc or UUTPxx=() , , 12 , , , x VVTPNN=() , ,12 , ,L , Nc or VVTPxx=() , ,12 , ,L , xc 2012/3/29 3 Summation of the properties of pure fluids (before mixing at TP and ) C UTPxx(), ,12 , ,L , xci− 1= ∑ xUTPi () , (8.1-1) i=1 where UU is the molar internal energy, i is the internal energy of the pure i-th component at TP and . C ˆˆ UTPww()(), ,12 , ,L , wcii− 1= ∑ wUTP , (8.1-2) i=1 where wi is the mass fraction of component i. 2012/3/29 4 At the same T and P 50 cc 25 cc + 25 cc H2O H2O or 52 cc 25 cc + 25 cc 48 cc + 2 cc A B -2 cc Attractive Repulsive 2012/3/29 5 Property change upon mixing (at constantTP and ) C Δ=mixθθ()TPx,,ii −∑ x θi () TP , i=1 Volume change upon mixing C Δ=mixVTP(),,,, VTPx()ii −∑ xVTPi () i=1 Enthalpy change upon mixing C Δ=mix HTP(),,,, HTPx()ii −∑ xHTPi () i=1 2012/3/29 6 Experimental data : properties changes upon mixing (H and V) Figure 8.1-1 Enthalpy-concentration diagram for aqueous sulfuric acid at 0.1 MPa.
    [Show full text]
  • Thermodynamics the Study of the Transformations of Energy from One Form Into Another
    Thermodynamics the study of the transformations of energy from one form into another First Law: Heat and Work are both forms of Energy. in any process, Energy can be changed from one form to another (including heat and work), but it is never created or distroyed: Conservation of Energy Second Law: Entropy is a measure of disorder; Entropy of an isolated system Increases in any spontaneous process. OR This law also predicts that the entropy of an isolated system always increases with time. Third Law: The entropy of a perfect crystal approaches zero as temperature approaches absolute zero. ©2010, 2008, 2005, 2002 by P. W. Atkins and L. L. Jones ©2010, 2008, 2005, 2002 by P. W. Atkins and L. L. Jones A Molecular Interlude: Internal Energy, U, from translation, rotation, vibration •Utranslation = 3/2 × nRT •Urotation = nRT (for linear molecules) or •Urotation = 3/2 × nRT (for nonlinear molecules) •At room temperature, the vibrational contribution is small (it is of course zero for monatomic gas at any temperature). At some high temperature, it is (3N-5)nR for linear and (3N-6)nR for nolinear molecules (N = number of atoms in the molecule. Enthalpy H = U + PV Enthalpy is a state function and at constant pressure: ∆H = ∆U + P∆V and ∆H = q At constant pressure, the change in enthalpy is equal to the heat released or absorbed by the system. Exothermic: ∆H < 0 Endothermic: ∆H > 0 Thermoneutral: ∆H = 0 Enthalpy of Physical Changes For phase transfers at constant pressure Vaporization: ∆Hvap = Hvapor – Hliquid Melting (fusion): ∆Hfus = Hliquid –
    [Show full text]
  • Physical Chemistry
    Subject Chemistry Paper No and Title 10, Physical Chemistry- III (Classical Thermodynamics, Non-Equilibrium Thermodynamics, Surface chemistry, Fast kinetics) Module No and Title 10, Free energy functions and Partial molar properties Module Tag CHE_P10_M10 CHEMISTRY Paper No. 10: Physical Chemistry- III (Classical Thermodynamics, Non-Equilibrium Thermodynamics, Surface chemistry, Fast kinetics) Module No. 10: Free energy functions and Partial molar properties TABLE OF CONTENTS 1. Learning outcomes 2. Introduction 3. Free energy functions 4. The effect of temperature and pressure on free energy 5. Maxwell’s Relations 6. Gibbs-Helmholtz equation 7. Partial molar properties 7.1 Partial molar volume 7.2 Partial molar Gibb’s free energy 8. Question 9. Summary CHEMISTRY Paper No. 10: Physical Chemistry- III (Classical Thermodynamics, Non-Equilibrium Thermodynamics, Surface chemistry, Fast kinetics) Module No. 10: Free energy functions and Partial molar properties 1. Learning outcomes After studying this module you shall be able to: Know about free energy functions i.e. Gibb’s free energy and work function Know the dependence of Gibbs free energy on temperature and pressure Learn about Gibb’s Helmholtz equation Learn different Maxwell relations Derive Gibb’s Duhem equation Determine partial molar volume through intercept method 2. Introduction Thermodynamics is used to determine the feasibility of the reaction, that is , whether the process is spontaneous or not. It is used to predict the direction in which the process will be spontaneous. Sign of internal energy alone cannot determine the spontaneity of a reaction. The concept of entropy was introduced in second law of thermodynamics. Whenever a process occurs spontaneously, then it is considered as an irreversible process.
    [Show full text]
  • Lecture Outline
    LECTURE OUTLINE 1. Equilibrium of heterogeneous system 2. Phase transformations R&Y, Chapter 2 Salby, Chapter 4 C&W, Chapter 4 A Short Course in Cloud Physics, R.R. Rogers and M.K. Yau; R&Y Thermodynamics of Atmospheres Fundamentals of Atmospheric Physics, and Oceanes, M.L. Salby; Salby J.A. Curry and P.J. Webster; C&W 2 /25 LECTURE OUTLINE 1. Equilibrium of heterogeneous system 2. Phase transformations Equilibrium conditions for a homogeneous system: • thermal equilibrium • mechanical equilibrium (at most an infinitesimal pressure difference exists between the system and its environment). A heterogeneous system must also be in: • chemical equilibrium. No conversion of mass occurs from one phase to the other. Chemical equilibrium requires a certain state variables to have no difference between the phases present. 4 /30 For a homogeneous system, two intensive properties describe the thermodynamic state. Only two state variables may be varied independently, so a homogeneous system has two thermodynamic degrees of freedom. For a heterogeneous system, each phase may be regarded as a homogeneous sub-system, one that is ‚open’ due to exchanges with the other phases present. The number of intensive properties that describes the thermodynamic state is proportional to the number of phases present. However, thermodynamic equilibrium between phases introduces additional constraints that actually reduce the degrees of freedom of a heterogeneous system below those of a homogeneous system. The system we consider is a two-component mixture of: • dry
    [Show full text]
  • Energy and Enthalpy Thermodynamics
    Energy and Energy and Enthalpy Thermodynamics The internal energy (E) of a system consists of The energy change of a reaction the kinetic energy of all the particles (related to is measured at constant temperature) plus the potential energy of volume (in a bomb interaction between the particles and within the calorimeter). particles (eg bonding). We can only measure the change in energy of the system (units = J or Nm). More conveniently reactions are performed at constant Energy pressure which measures enthalpy change, ∆H. initial state final state ∆H ~ ∆E for most reactions we study. final state initial state ∆H < 0 exothermic reaction Energy "lost" to surroundings Energy "gained" from surroundings ∆H > 0 endothermic reaction < 0 > 0 2 o Enthalpy of formation, fH Hess’s Law o Hess's Law: The heat change in any reaction is the The standard enthalpy of formation, fH , is the change in enthalpy when one mole of a substance is formed from same whether the reaction takes place in one step or its elements under a standard pressure of 1 atm. several steps, i.e. the overall energy change of a reaction is independent of the route taken. The heat of formation of any element in its standard state is defined as zero. o The standard enthalpy of reaction, H , is the sum of the enthalpy of the products minus the sum of the enthalpy of the reactants. Start End o o o H = prod nfH - react nfH 3 4 Example Application – energy foods! Calculate Ho for CH (g) + 2O (g) CO (g) + 2H O(l) Do you get more energy from the metabolism of 1.0 g of sugar or
    [Show full text]
  • Biochemical Thermodynamics
    © Jones & Bartlett Learning, LLC. NOT FOR SALE OR DISTRIBUTION CHAPTER 1 Biochemical Thermodynamics Learning Objectives 1. Defi ne and use correctly the terms system, closed, open, surroundings, state, energy, temperature, thermal energy, irreversible process, entropy, free energy, electromotive force (emf), Faraday constant, equilibrium constant, acid dissociation constant, standard state, and biochemical standard state. 2. State and appropriately use equations relating the free energy change of reactions, the standard-state free energy change, the equilibrium constant, and the concentrations of reactants and products. 3. Explain qualitatively and quantitatively how unfavorable reactions may occur at the expense of a favorable reaction. 4. Apply the concept of coupled reactions and the thermodynamic additivity of free energy changes to calculate overall free energy changes and shifts in the concentrations of reactants and products. 5. Construct balanced reduction–oxidation reactions, using half-reactions, and calculate the resulting changes in free energy and emf. 6. Explain differences between the standard-state convention used by chemists and that used by biochemists, and give reasons for the differences. 7. Recognize and apply correctly common biochemical conventions in writing biochemical reactions. Basic Quantities and Concepts Thermodynamics is a system of thinking about interconnections of heat, work, and matter in natural processes like heating and cooling materials, mixing and separation of materials, and— of particular interest here—chemical reactions. Thermodynamic concepts are freely used throughout biochemistry to explain or rationalize chains of chemical transformations, as well as their connections to physical and biological processes such as locomotion or reproduction, the generation of fever, the effects of starvation or malnutrition, and more.
    [Show full text]
  • Thermodynamics, Experimental, and Modelling of Aqueous Electrolyte and Amino Acid Solutions
    Downloaded from orbit.dtu.dk on: Sep 26, 2021 Thermodynamics, Experimental, and Modelling of Aqueous Electrolyte and Amino Acid Solutions Breil, Martin Peter Publication date: 2001 Document Version Publisher's PDF, also known as Version of record Link back to DTU Orbit Citation (APA): Breil, M. P. (2001). Thermodynamics, Experimental, and Modelling of Aqueous Electrolyte and Amino Acid Solutions. General rights Copyright and moral rights for the publications made accessible in the public portal are retained by the authors and/or other copyright owners and it is a condition of accessing publications that users recognise and abide by the legal requirements associated with these rights. Users may download and print one copy of any publication from the public portal for the purpose of private study or research. You may not further distribute the material or use it for any profit-making activity or commercial gain You may freely distribute the URL identifying the publication in the public portal If you believe that this document breaches copyright please contact us providing details, and we will remove access to the work immediately and investigate your claim. Thermodynamics, Experimental, and Modelling of Aqueous Electrolyte and Amino Acid Solutions Martin P. Breil 2001 IVC-SEP Department of Chemical Engineering Technical University of Denmark DK-2800 Kongens Lyngby, Denmark Preface iii Preface This thesis is submitted as a partial fulfilment of the Ph.D. degree at the Technical University of Denmark. The project, granted by the IVC-SEP, has been carried out from October 1998 to September 2001 at the Department of Chemical Engineering, Technical University of Denmark under the supervision of Jørgen Mollerup.
    [Show full text]
  • Chapter 9 Ideal and Real Solutions
    2/26/2016 CHAPTER 9 IDEAL AND REAL SOLUTIONS • Raoult’s law: ideal solution • Henry’s law: real solution • Activity: correlation with chemical potential and chemical equilibrium Ideal Solution • Raoult’s law: The partial pressure (Pi) of each component in a solution is directly proportional to the vapor pressure of the corresponding pure substance (Pi*) and that the proportionality constant is the mole fraction (xi) of the component in the liquid • Ideal solution • any liquid that obeys Raoult’s law • In a binary liquid, A-A, A-B, and B-B interactions are equally strong 1 2/26/2016 Chemical Potential of a Component in the Gas and Solution Phases • If the liquid and vapor phases of a solution are in equilibrium • For a pure liquid, Ideal Solution • ∆ ∑ Similar to ideal gas mixing • ∆ ∑ 2 2/26/2016 Example 9.2 • An ideal solution is made from 5 mole of benzene and 3.25 mole of toluene. (a) Calculate Gmixing and Smixing at 298 K and 1 bar. (b) Is mixing a spontaneous process? ∆ ∆ Ideal Solution Model for Binary Solutions • Both components obey Rault’s law • Mole fractions in the vapor phase (yi) Benzene + DCE 3 2/26/2016 Ideal Solution Mole fraction in the vapor phase Variation of Total Pressure with x and y 4 2/26/2016 Average Composition (z) • , , , , ,, • In the liquid phase, • In the vapor phase, za x b yb x c yc Phase Rule • In a binary solution, F = C – p + 2 = 4 – p, as C = 2 5 2/26/2016 Example 9.3 • An ideal solution of 5 mole of benzene and 3.25 mole of toluene is placed in a piston and cylinder assembly.
    [Show full text]
  • Topic 4 Thermodynamics Thermodynamics
    Topic 4 Thermodynamics Thermodynamics • We need thermodynamic data to: – Determine the heat release in a combustion process (need enthalpies and heat capacities) – Calculate the equilibrium constant for a reaction – this allows us to relate the rate coefficients for forward and reverse reactions (need enthalpies, entropies (and hence Gibbs energies), and heat capacities. • This lecture considers: – Classical thermodynamics and statistical mechanics – relationships for thermodynamic quantities – Sources of thermodynamic data – Measurement of enthalpies of formation for radicals – Active Thermochemical Tables – Representation of thermodynamic data for combustion models Various thermodynamic relations are needed to determine heat release and the relations between forward and reverse rate coefficients A statement of Hess’s Law ni is the stoichiometric coefficient Hess’s Law: The standard enthalpy of an overall reaction is the sum of the standard enthalpies of the individual reactions into which the reaction may be divided Tabulated thermodynamic quantities. 1. Standard enthalpy of formation o Standard enthalpy change of formation, fH The standard enthalpy change when 1 mol of a substance is formed from its elements in their reference states, at a stated temperature (usually 298 K). The reference state is the most stable state at that temperature, and at a pressure of 1 bar. o e.g. C(s) + 2H2(g) CH4(g) fH = -74.8 kJ mol-1 The standard enthalpies of formation of C(s) and H2(g) are both zero The reference state for carbon is graphite. Standard
    [Show full text]
  • Chapter 20: Thermodynamics: Entropy, Free Energy, and the Direction of Chemical Reactions
    CHEM 1B: GENERAL CHEMISTRY Chapter 20: Thermodynamics: Entropy, Free Energy, and the Direction of Chemical Reactions Instructor: Dr. Orlando E. Raola 20-1 Santa Rosa Junior College Chapter 20 Thermodynamics: Entropy, Free Energy, and the Direction of Chemical Reactions 20-2 Thermodynamics: Entropy, Free Energy, and the Direction of Chemical Reactions 20.1 The Second Law of Thermodynamics: Predicting Spontaneous Change 20.2 Calculating Entropy Change of a Reaction 20.3 Entropy, Free Energy, and Work 20.4 Free Energy, Equilibrium, and Reaction Direction 20-3 Spontaneous Change A spontaneous change is one that occurs without a continuous input of energy from outside the system. All chemical processes require energy (activation energy) to take place, but once a spontaneous process has begun, no further input of energy is needed. A nonspontaneous change occurs only if the surroundings continuously supply energy to the system. If a change is spontaneous in one direction, it will be nonspontaneous in the reverse direction. 20-4 The First Law of Thermodynamics Does Not Predict Spontaneous Change Energy is conserved. It is neither created nor destroyed, but is transferred in the form of heat and/or work. DE = q + w The total energy of the universe is constant: DEsys = -DEsurr or DEsys + DEsurr = DEuniv = 0 The law of conservation of energy applies to all changes, and does not allow us to predict the direction of a spontaneous change. 20-5 DH Does Not Predict Spontaneous Change A spontaneous change may be exothermic or endothermic. Spontaneous exothermic processes include: • freezing and condensation at low temperatures, • combustion reactions, • oxidation of iron and other metals.
    [Show full text]
  • Proquest Disserststions
    THERMODYNAMIC PROPERTIES OF NUCLEIC ACID BASES AND NUCLEOSIDES UNDER HYDROTHERMAL CONDITIONS A Thesis Presented to The Faculty of Graduate Studies of The University of Guelph by VANESSA MANN In partial fulfilment of requirements for the degree of Master of Applied Science August, 2009 © Vanessa Mann, 2009 Library and Archives Bibliothèque et 1*1 Canada Archives Canada Published Heritage Direction du Branch Patrimoine de l'édition 395 Wellington Street 395, rue Wellington OttawaONK1A0N4 OttawaONK1A0N4 Canada Canada Your file Votre référence ISBN: 978-0-494-68602-7 Our file Notre référence ISBN: 978-0-494-68602-7 NOTICE: AVIS: The author has granted a non- L'auteur a accordé une licence non exclusive exclusive license allowing Library and permettant à la Bibliothèque et Archives Archives Canada to reproduce, Canada de reproduire, publier, archiver, publish, archive, preserve, conserve, sauvegarder, conserver, transmettre au public communicate to the public by par télécommunication ou par l'Internet, prêter, telecommunication or on the Internet, distribuer et vendre des thèses partout dans le loan, distribute and sell theses monde, à des fins commerciales ou autres, sur worldwide, for commercial or non- support microforme, papier, électronique et/ou commercial purposes, in microform, autres formats. paper, electronic and/or any other formats. The author retains copyright L'auteur conserve la propriété du droit d'auteur ownership and moral rights in this et des droits moraux qui protège cette thèse. Ni thesis. Neither the thesis nor la thèse ni des extraits substantiels de celle-ci substantial extracts from it may be ne doivent être imprimés ou autrement printed or otherwise reproduced reproduits sans son autorisation.
    [Show full text]