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This chapter analyzes thermodynamic parameters (ΔG°, ΔH°, ΔS°, and ΔC_p°) of the H-bond association equilibrium in gas-phase, crystalline solids, and non-polar and aqueous solvents. The key feature is the extrathermodynamic enthalpy-entropy compensation relationship, ΔH° = βS°, induced by the loss of degrees of freedom during H-bond formation and inducing a considerable reduction of the association constants (e.g., ΔG° ≅ 2/3 Δ/H° in gas-phase and ≅ 1/3 Δ/H° in non-polar solvents). H-bonds in aqueous solutions are mostly studied through drug-receptor binding thermodynamic data. A distinction is made between hydrophobic binding to cytoplasmic and hydrophilic binding to membrane receptors, respectively characterized by negative or nearly zero ΔC_p° values. The unifying aspect of all drug-receptor phenomena remains enthalpy-entropy compensation. This is interpreted in terms of solvent reorganization by the Grunwald and Steel (1995) theoretical model showing that solvent effects do not significantly affect the intrinsic ΔG° of the H-bonds which directly connect the drug to the receptor binding site.

The internal energy U of a system is a function of state, which means that a system undergoes the same change in U when we move it from one equilibrium state to another, irrespective of which route we take through parameter space. This makes U a very useful quantity, though not a uniquely useful quantity. In fact, we can make a number of other functions of state, simply by adding to U various other combinations of the functions of state p, V , T, and S in such a way as to give the resulting quantity the dimensions of energy. These new functions of state are called thermodynamic potentials. Most thermodynamic potentials that one could pick are really not very useful but three of them are extremely useful and are given special symbols: H = U + pV, F = U - TS, and G = U + pV - TS. This chapter explores why these three quantities are so useful.

The relationships between atoms and molecules are mediated by the theory of chemical bonding. Atoms were forged in the extreme conditions that occurred at the cores of stars in the early stages of the development of the universe, but they did not survive as isolated systems in the much milder conditions of mature planets, where more stable structures are formed by highly stabilising interactions between electrons of different atoms, in terms of lowering of the potential energy and of spin pairing. This chapter discusses molecular structure, size and shape, along with classification concepts in many particle systems, symmetry and order, enthalpy and entropy, mass and length dimensions of a molecule, atomic radii, molecular volume and surface, and molecular size in terms of electron density.

In the gaseous state, molecules are to a good approximation isolated entities traveling in space at high speed with sparse and near elastic collisions. At the other extreme, a perfect crystal has a periodic and symmetric intermolecular structure. The structure is dictated by intermolecular forces, and molecules can only perform small oscillations around their equilibrium positions. In between these two extremes, matter has many more ways of aggregation. This chapter deals with proper liquids and the liquid state. Molecular diffusion in liquids occurs approximately on the timescale of nanoseconds, to be compared with the timescale of molecular or lattice vibrations. This chapter also discusses molecular dynamics (MD), the Monte Carlo (MC) method, structural and dynamic descriptors for liquids, physicochemical properties of liquids from MD or MC simulations, simulations of enthalpy, heat capacity and density, crystal and liquid equations of state, polarisability and dielectric constants, free energy simulations, and simulation of water.

Definition and mathematics of enthalpy. Definition of heat capacity at constant pressure as C_P = (∂H/∂T)_V. Endothermic and exothermic reactions. Role of the change in enthalpy as regards the direction and spontaneity of a change in state. Enthalpy changes and phase changes. Measuring enthalpy changes by calorimetry. Hess’s law of constant heat formation. Chemical standards and standard states. Standard enthalpies of formation, ionic enthalpies and bond energies. How the change in enthalpy varies with temperature. Kirchhoff’s equations. Applications of thermochemistry to a variety of worked examples, including flames and explosions.

Thermochemistry explores the basic principles of energy changes in chemical reactions. Calorimetry is described as a tool to measure the quantity of heat involved in a chemical or physical change. Quantitative overviews of enthalpy and the stoichiometry of thermochemical equations are provided, including the concepts of endothermic and exothermic reactions. Standard conditions are defined to allow comparison of enthalpies of reactions and determine how the enthalpy change for any reaction can be obtained. Hess"s Law, which allows the enthalpy change of any reaction to be calculated, is discussed with illustrative examples. A presentation of bond energies and bond dissociation enthalpies is offered along with the use of bond enthalpy to estimate heats of reactions.

Thermodynamics describes the relationship between heat, work, energy and motion. The key concepts are the conservation of energy and the maximisation of entropy (or disorder) as given by the first and second laws of thermodynamics. Boltzmann’s equation explains how the entropy of a material is related to the disorder of its atoms or molecules, as measured by the probability or the number of equivalent atomic or molecular structures. This chapter examines thermodynamic properties such as internal energy, enthalpy and Gibbs and Helmholtz free energy; physical properties such as specific heat and thermal expansion coefficient; and the application of thermodynamics to chemical reactions, solid and liquid solutions, and phase separation. Ludwig Boltzmann’s early life as the son of a minor tax official in Austria is described, as are: his scientific career in a series of Austrian and German universities; his philosophical arguments with Ernst Mach and the phenomenalists about whether atoms do or do not exist; his increasing moodiness, paranoia and bipolar disorder; and his ultimate suicide while trying to recuperate from depression in Trieste.

The behavior of energy in bulk-matter systems is subtle. We observe that energy flows spontaneously from high to low temperature; we refer to this flowing energy as heat; and we distinguish heat from work, the transfer of energy through mechanical or other means unrelated to temperature. On the other hand, simple models of gases and solids strongly suggest that at the molecular level all energy is purely mechanical. This introductory chapter surveys these basic concepts of thermal physics, illustrates them with a wide variety of familiar examples, and sets the stage for developing a deeper understanding.

Chemical Thermodynamics discusses the fundamental laws of thermodynamics along with their relationships to heat, work, enthalpy, entropy, and temperature. Predicting the direction of a spontaneous change and calculating the change in entropy of a reaction are core concepts. The relationship between entropy, free energy and work is covered. The Gibbs free energy is used quantitatively to predict if reactions or processes are going to be exothermic and spontaneous or endothermic under the stated conditions. Also explored are the enthalpy and entropy changes during a phase change. Finally the Gibbs free energy of a chemical reaction is related to its equilibrium constant and the temperature.

Thermodynamics specifies the relation between an independent, predictor variable, and what is predicted. It is often the case that changing the variables regarded as independent can greatly simplify problem solving. The chapter shows how using an intensive variable (like temperature or pressure) as the predictor loses information that can be retained if it is expressed by a different function. It shows the importance of Legendre transforms, which contain the same information about the system as is available by using extensive variables. Legendre transforms exploiting the fundamental relation are shown to yield the Helmholtz free energy, the enthalpy, and the Gibbs free energy. Massieu functions are introduced as an alternative that is particularly important for models exhibiting negative temperatures.

This chapter derives the energy minimum principle from the entropy maximum principle. It postulates and consider the consequences of extensivity. From this are further derived minimum principles for the Helmholtz free energy, enthalpy, and Gibbs free energy. Because of its importance in engineering, exergy is also introduced, and the exergy minimum principle is justified. Analogously to these minimum principles, maximum principles can be derived for the Massieu functions from the entropy maximum principle. For the analysis of the entropy maximum principle, we isolated a composite system and released an internal constraint. Since the composite system was isolated, its total energy remained constant. The composite system went to the most probable macroscopic state after release of the internal constraint, and the total entropy went to its maximum.

Energy is the capacity to do work and heat is the spontaneous flow of energy from one body or system to another through the random movement of atoms or molecules. The entropy of a system determines how much of its internal energy is unavailable for work under isothermal conditions, and the Gibbs energy is the energy available for work under isothermal conditions and constant pressure. The Second Law of Thermodynamics states that for any reaction to proceed spontaneously the total entropy (system plus surroundings) must increase, which is why metabolic processes release heat. All organisms are thermodynamically open systems, exchanging both energy and matter with their surroundings. They can decrease their entropy in growth and development by ensuring a greater increase in the entropy of the environment. For an ideal gas in thermal equilibrium the distribution of energy across the component atoms or molecules is described by the Maxwell-Boltzmann equation. This distribution is fixed by the temperature of the system.

All other things being equal, physiological reaction rate increases roughly exponentially with temperature. Organisms that have adapted over evolutionary time to live at different temperatures can have enzyme variants that exhibit similar kinetics at the temperatures to which they have adapted to operate. Within species whose distribution covers a range of temperatures, there may be differential expression of enzyme variants with different kinetics across the distribution. Enzymes adapted to different optimum temperatures differ in their amino acid sequence and thermal stability. The Gibbs energy of activation tends to be slightly lower in enzyme variants adapted to lower temperatures, but the big change is a decrease in the enthalpy of activation, with a corresponding change in the entropy of activation, both associated with a more open, flexible structure. Despite evolutionary adjustments to individual enzymes involved in intermediary metabolism (ATP regeneration), many whole-organism processes operate faster in tropical ectotherms compared with temperate or polar ectotherms. Examples include locomotion (muscle power output), ATP regeneration (mitochondrial function), nervous conduction and growth.

A critical chapter, explaining how the principles of thermodynamics can be applied to real systems. The central concept is the Gibbs free energy, which is explored in depth, with many examples. Specific topics addressed are: Spontaneous changes in closed systems. Definitions and mathematical properties of Gibbs free energy and Helmholtz free energy. Enthalpy- and entropy-driven reactions. Maximum available work. Coupled reactions, and how to make non-spontaneous changes happen, with examples such as tidying a room, life, and global warming. Standard Gibbs free energies. Mixtures, partial molar quantities and the chemical potential.

This chapter draws together all the main mathematical equations into a single, structured, sequence, so providing a source of reference, as well as enabling the student to appreciate how superficially different equations are, in fact, component parts of a ‘bigger picture’. The chapter also introduces some new material, such as the Maxwell relations, the chain rule, the thermodynamic equations-of-state, isenthalpic throttling processes, the Joule-Thomson coefficient and the compressibility factor – so setting the scene for the discussion of real systems in the following chapter.

Conformational isomers represent one form of stereoisomer in which the two structures can be interconverted by simple rotations about a single bond. Such conformational isomers are called rotomers and play a role in many areas of chemistry. In this experiment the 1,2 –dichloroethane rotomer molecule is studied using Raman spectroscopy at temperatures ranging from room temperature to that of liquid nitrogen. The anti and gauche forms are clearly differentiated in the Raman spectra. The enthalpy difference, Δ‎H_f, between the anti and gauche forms is then determined from a van’t Hoff plot. Calculations of the Raman spectra and the Δ‎H_f are carried out using the Gaussian program and excellent agreement between experiment and theory is seen.

This chapter discusses some of the more important thermodynamic properties of liquid metallic elements in the area of materials process science. These are evaporation enthalpy, vapour pressure, and heat capacity. Theoretical and empirical equations describing these thermodynamic properties are presented along with experimental data.

This chapter provides a compilation (in tabular form) of experimental data for the thermophysical properties of liquid metallic elements. The properties covered include melting and boiling points, atomic numbers and relative atomic masses, molar melting and evaporation enthalpies, densities and related data, vapour pressure equations, molar heat capacities at constant pressure, sound velocities, surface tension data, viscosities, self-diffusivity data, and electrical resistivities and thermal conductivities.

Chapter 6 introduces quantum-mechanical ensemble theory by proving the asymptotic equivalence of the quantum-mechanical, microcanonical ensemble average with the quantum grand canonical ensemble average for many-particle systems, based on the method of Darwin and Fowler. The procedures involved identify the grand partition function, entropy and other statistical thermodynamic variables, including the grand potential, Helmholtz free energy, thermodynamic potential, Gibbs free energy, Enthalpy and their relations in accordance with the fundamental laws of thermodynamics. Accompanying saddle-point integrations define temperature (inverse thermal energy) and chemical potential (Fermi energy). The concomitant emergence of quantum statistical mechanics and Bose–Einstein and Fermi–Dirac distribution functions are discussed in detail (including Bose condensation). The magnetic moment is derived from the Helmholtz free energy and is expressed in terms of a one-particle retarded Green’s function with an imaginary time argument related to inverse thermal energy. This is employed in a discussion of diamagnetism and the de Haas-van Alphen effect.