Chemistry
Free Energy
Free energy in chemistry refers to the energy available to do work in a system. It is the energy that is available to drive chemical reactions and processes. Free energy is a measure of the spontaneity of a reaction, with a negative change in free energy indicating a spontaneous process.
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11 Key excerpts on "Free Energy"
No longer available |Learn more- (Author)
- 2014(Publication Date)
- Library Press(Publisher)
• Entropy ( S ) • Gibbs Free Energy ( G ) Most identities in chemical thermodynamics arise from application of the first and second laws of thermodynamics, particularly the law of conservation of energy, to these state functions. Chemical energy Chemical energy is the potential of a chemical substance to undergo a transformation through a chemical reaction or to transform other chemical substances. Breaking or ________________________ WORLD TECHNOLOGIES ________________________ making of chemical bonds involves energy, which may be either absorbed or evolved from a chemical system. Energy that can be released (or absorbed) because of a reaction between a set of chemical substances is equal to the difference between the energy content of the products and the reactants. This change in energy is called the change in internal energy of a chemical reaction. Where is the internal energy of formation of the reactant molecules that can be calculated from the bond energies of the various chemical bonds of the molecules under consideration and is the internal energy of formation of the product molecules. The internal energy change of a process is equal to the heat change if it is measured under conditions of constant volume, as in a closed rigid con-tainer such as a bomb calorimeter. However, under conditions of constant pressure, as in reactions in vessels open to the atmosphere, the measured heat change is not always equal to the internal energy change, because pressure-volume work also releases or absorbs energy. (The heat change at constant pressure is called the enthalpy change; in this case the enthalpy of formation). Another useful term is the heat of combustion, which is the energy released due to a combustion reaction and often applied in the study of fuels. Food is similar to hydrocarbon fuel and carbohydrate fuels, and when it is oxidized, its caloric content is similar (though not assessed in the same way as a hydrocarbon fuel).
eBook - PDF- Allan Blackman, Steven E. Bottle, Siegbert Schmid, Mauro Mocerino, Uta Wille(Authors)
- 2022(Publication Date)
- Wiley(Publisher)
G = H - TS We will look at these thermodynamic functions in greater detail, and provide exact definitions, in the appropriate sections of this chapter. For now, it is sufficient to state that the value of Gibbs energy, G, allows us to predict spontaneity. We can liken a spontaneous change to a ball at the top of a hill — once pushed, it will roll spontaneously down the hill until it reaches the bottom, the point of minimum energy. CHAPTER 8 Chemical thermodynamics 349 In the same way, we will see that chemical reactions proceed in the direction that leads to a decrease in the Gibbs energy of the system. We will also see that overall chemical and physical change will cease once the Gibbs energy of the system is minimised, and that, under these conditions, the system is at equilibrium. 8.2 Thermodynamic concepts LEARNING OBJECTIVE 8.2 Use the terminology and units of chemical thermodynamics. Before we begin our study of chemical thermodynamics, we need to discuss some important thermody- namic concepts. Heat and temperature Arguably the most important and indeed most familiar thermodynamic term, heat (q), is perhaps the most conceptually difficult. Heat is the energy that is transferred as the result of a temperature difference. If two bodies having different temperatures are brought into direct contact, there will be heat flow from the hotter to the colder body, until both are at the same temperature. As we will see, we cannot actually measure heat directly. The definition of temperature also follows from the above; two bodies have the same temperature if they are in thermal equilibrium; that is, there is no heat flow between them when they are in direct contact. It is important to note that the thermodynamic temperature (sometimes called the absolute temperature) scale is used in nearly all thermodynamic calculations.
eBook - PDFChemistry
Structure and Dynamics
- James N. Spencer, George M. Bodner, Lyman H. Rickard(Authors)
- 2011(Publication Date)
- Wiley(Publisher)
The name of this quantity recognizes the contributions of J. Willard Gibbs, a professor of mathematical physics at Yale from 1871 until the early 1900s, who some consider to be the greatest scientist produced by the United States. The Gibbs Free Energy of a system at any moment in time is defined as the enthalpy of the system minus the product of the temperature times the entropy of the system. The Gibbs Free Energy is therefore a state function because it is defined in terms of thermodynamic properties (enthalpy, entropy, and temperature) that are state functions. At a given temperature, the change in the Gibbs Free Energy of the sys- tem that occurs during a reaction is therefore equal to the change in the enthalpy of the system minus the product of the temperature times the change in entropy of the system. The beauty of the equation defining changes in the Free Energy of a system is its ability to determine the relative importance of the enthalpy and entropy terms for a particular reaction at a given temperature. The change in the Free Energy of the system that occurs during a reaction measures the balance between the two driving forces that determine whether a reaction is spontaneous. As we have seen, the enthalpy and entropy terms have different sign con- ventions. The enthalpy term is favorable when it is negative, whereas the entropy term is favorable when positive. ¢G = ¢H - T¢S G = H - TS The overall entropy of reaction is negative because the reaction transforms 4 moles of reactants into 2 moles of products. According to these calculations, ¢H° for the reaction is 92.2 kJ/mol rxn and ¢S° for the reaction is 198.8 J/mol rxn K. Enthalpy (¢H° 6 0) therefore drives the reaction toward the products, but entropy (¢S° 6 0) drives the reac- tion toward the reactants.
eBook - PDF- Donald T. Haynie(Author)
- 2008(Publication Date)
- Cambridge University Press(Publisher)
Chapter 4 Gibbs Free Energy – theory A. Introduction This chapter discusses a thermodynamic relationship that provides a basis for explaining spontaneous chemical reactivity, chemical equilibrium, and the phase behavior of chemical compounds. The relationship involves a thermodynamic state function that enables prediction of the direction of a chemical reaction at constant tem-perature and pressure . The constraints of fixed T and p might seem annoyingly restrictive, because they are less general than the requirements of the Second Law, but in fact the gains made on imposing the constraints will outweigh the losses. How is that? One reason is at any given time an individual organism is practically at uniform pressure and temperature (but be sure to see the Exercises at the end of the chapter). Another is that constant temperature and pressure are the very conditions under which nearly all bench-top biochemistry experiments are done. Yet another is that, although the total entropy of the universe must increase in order for a process to be spontaneous, evaluation of the total entropy change requires mea-surement of both the entropy change of the system and the entropy change of the surroundings. Whereas 1 S system can often be found without too much difficulty, albeit only indirectly, 1 S surroundings can be hard to measure! How could one measure the entropy change of the rest of the universe? The subject of the present chapter provides a way around the difficulty. A particularly clear example of the inadequacy of 1 S system to predict the direction of spontaneous change is given by the behavior of water at its freezing point. Table 4.1 shows the thermo-dynamic properties of water for the liquid ! solid phase transition. The decrease in internal energy (which is practically identical to the enthalpy as long as the number of moles of gas doesn’t change; see Chapter 2 ) would suggest that water freezes spontaneously in the range 263–283 K.
eBook - PDF- Young, William Vining, Roberta Day, Beatrice Botch(Authors)
- 2017(Publication Date)
- Cengage Learning EMEA(Publisher)
Because it can be confusing to know when to use the different energies and to what situa-tions they apply, their definitions are summarized here. 1. Internal energy, E , is the sum of all the submicroscopic kinetic and potential energies of all the particles that make up a system. The change in internal energy, uni2206 E , is calcu-lated from the first law of thermodynamics, uni2206 E 5 q 1 w . 2. Enthalpy, H , is a defined energy that is based on the internal energy, H 5 E 1 PV . It is convenient to use enthalpy in the constant-pressure processes that are common in chemistry because uni2206 H 5 q , the heat absorbed at constant pressure. You can think about enthalpy as the internal energy measured at constant pressure. 3. Gibbs Free Energy, G , like enthalpy is a defined energy ( G 5 H 2 TS ) . Free Energy is related to reaction spontaneity by the second law, uni2206 S universe 7 0 . At constant temperature and pressure, the change in the entropy of the universe is equal to 2 uni2206 G > T , and under these conditions, uni2206 G points the direction of chemical change. Internal energy, E , and enthalpy, H , are first law energies. They are associated with the amount of energy exchanged. Gibbs Free Energy is a second law energy. It has to do with the direction of change and the quality of the energy, the maximum amount of energy that is available to do work. Interactive Table 20.3.1 Relationship between uni2206 G and Reaction Spontaneity uni2206 G 6 0 Spontaneous in the forward direction uni2206 G 5 0 At equilibrium; no net change will occur uni2206 G 7 0 Nonspontaneous in the forward direction; spontaneous in the reverse direction Copyright 2018 Cengage Learning. All Rights Reserved. May not be copied, scanned, or duplicated, in whole or in part. WCN 02-300
eBook - PDFBiosimulation
Simulation of Living Systems
- Daniel A. Beard(Author)
- 2012(Publication Date)
- Cambridge University Press(Publisher)
5 Chemical reaction systems: thermodynamics and chemical equilibrium Overview This and the following two chapters are focused on analyzing and simulating chem- ical systems. These chapters will introduce basic concepts of thermodynamics and kinetics for application to biochemical systems, such as biochemical synthesis, cel- lular metabolism and signaling processes, and gene regulatory networks. Although we have seen examples of chemical kinetics in previous chapters, notably in Sections 2.3 and 2.4, in those examples we developed the expressions governing the chemistry more from intuition than from a physical theory. One of the primary goals here will be to develop a formal physical/chemical foundation for analyzing and simulating complex biochemical systems. As is our practice throughout this book, these concepts will be applied to analyze real data (and understand the behavior of real systems) later in this chapter and elsewhere. Yet, because the rules governing the behavior of biochemical systems are grounded in thermodynamics, we must begin our investigation into chemical systems by establishing some fundamental concepts in chemical thermo- dynamics. The concept of Free Energy is particularly crucial to understanding thermodynamic driving forces in chemistry. We will see that both a physical definition and an intuitive understanding of Free Energy require physical definitions and intuitive understandings of temperature and entropy. All of this means that this chapter will begin with some abstract thought experiments and derivations of physical concepts. 5.1 Temperature, pressure, and entropy 5.1.1 Microstates and macrostates All thermodynamic theory arises from the fact that physical systems composed of many atoms and/or molecules attain a large number (often a practically infinite 146 Chemical reaction systems: thermodynamics and chemical equilibrium number) of microstates under defined macroscopic conditions, such as temper- ature, pressure, and volume.
- Frank R. Foulkes(Author)
- 2012(Publication Date)
- CRC Press(Publisher)
CHAPTE R TWELV E THERMODYNAMICS (VII) 12.1 GIBBS Free Energy ( G ) When the process taking place in the system is spontaneous , in principle it can be harnessed to de-liver work to the surroundings. Conversely, when the process is non-spontaneous , we must do work on the system in order to make the process occur. According to the Second Law of Thermo-dynamics, the process taking place in our system is spontaneous if 6 S univ = 6 S syst + 6 S surr > 0 . . . [1] Calculating 6 S univ is cumbersome because, although we usually are concerned primarily with what is taking place in the system––where the process of interest occurs––we also must calculate 6 S surr to find out if our process is spontaneous. For this purpose a more useful function than the entropy is the Gibbs Free Energy G , defined as G > H – TS Gibbs Free Energy . . . [2] In Eqn [2] H is the enthalpy of the system, T is its temperature, and S is its entropy. G is a state function because all its components are state functions. Furthermore, G is a property of the sys-tem , not of the surroundings. The Gibbs Free Energy is named after J. Willard Gibbs, a famous nineteenth century American thermodynamicist who is considered to be the “father” of chemical thermodynamics. 12.2 GIBBS Free Energy CHANGES, “ OTHER” WORK, AND SPONTANEITY By definition, G H TS > < . . . [3] But, H in turn is defined as H U PV > + . . . [4] Substituting [4] into [3] gives: G = U + PV – TS . . . [5] Differentiating: dG = dU + PdV + VdP – TdS – SdT But dU = b Q + b W 12-2 THERMODYNAMICS (VII) Therefore, dG = ( b Q + b W) + PdV + VdP – TdS – SdT . . . [6] Now consider any kind of work other than PV -work (e.g., electrical work, shaft work, etc.): If we let W total = ( PV -work) + (“other” work) . . . [7] then b W total = –P ext dV + b W° . . . [8] where W° is “other” work; i.e., any work other than PV-work. 1 Substituting Eqn [8] into Eqn [6] gives: dG = ( b Q – P ext dV + b W°) + PdV + VdP – TdS – SdT .
eBook - PDFChemistry
An Atoms First Approach
- Steven Zumdahl, Susan Zumdahl, Donald J. DeCoste, , Steven Zumdahl, Steven Zumdahl, Susan Zumdahl, Donald J. DeCoste(Authors)
- 2020(Publication Date)
- Cengage Learning EMEA(Publisher)
We saw in Section 16.4 that for a process at constant temperature and pressure, we can use the change in free en- ergy of the system to predict the sign of DS univ and thus the direction in which it is spontaneous. So far we have applied these ideas only to physical processes, such as changes of state and the formation of solutions. However, the main business of chem- istry is studying chemical reactions, and, therefore, we want to apply the second law to reactions. First, we will consider the entropy changes accompanying chemical reactions that occur under conditions of constant temperature and pressure. As for the other types of processes we have considered, the entropy changes in the surroundings are determined by the heat flow that occurs as the reaction takes place. However, the entropy changes in the system (the reactants and products of the reaction) are determined by positional probability. For example, in the ammonia synthesis reaction N 2 s gd 1 3H 2 s gd ¡ 2NH 3 s gd four reactant molecules become two product molecules, lowering the number of inde- pendent units in the system, which leads to less positional disorder. See Exercises 16.37 through 16.39 656 CHAPTER 16 Spontaneity, Entropy, and Free Energy Copyright 2021 Cengage Learning. All Rights Reserved. May not be copied, scanned, or duplicated, in whole or in part. Due to electronic rights, some third party content may be suppressed from the eBook and/or eChapter(s). Editorial review has deemed that any suppressed content does not materially affect the overall learning experience. Cengage Learning reserves the right to remove additional content at any time if subsequent rights restrictions require it. N H Less entropy Greater entropy Fewer molecules mean fewer possible configurations. To help clarify this idea, con- sider a special container with a million compartments, each large enough to hold a hydrogen molecule. Thus there are a million ways one H 2 molecule can be placed in this container.
eBook - PDF- Don Shillady(Author)
- 2011(Publication Date)
- CRC Press(Publisher)
6 Gibbs ’ Free Energy and Equilibria INTRODUCTION In previous chapters, we have stressed that in nature energy tends to decrease while entropy tends to increase. A naive fi rst consideration of any machine or process is that energy is needed to continue operation and we often overlook energy expended on various repair activities that are a form of entropy management. It becomes more obvious that entropy is a factor when one studies chemical processes that ‘‘ should ’’ occur based on energy considerations but nevertheless require some sort of a catalyst or other special conditions, which imply geometric constraints that overcome the natural tendency of randomness to increase. The value of D S is a change in a state variable but the path can be modi fi ed by special conditions such as the introduction of a catalytic surface, which allows reactants to meet side-by-side compared to random collisions in the gas phase. Josiah Willard Gibbs (1839 – 1903) was a foremost U.S. scientist (Figure 6.1) who made important advances in thermo-dynamics applying the new idea of ‘‘ chemical potential ’’ ( D G = n ) as a Free Energy per mole of a substance in phase diagrams and applied to equilibria. At the time of his work, few people understood it but it was later developed into the idea of Free Energy and greatly affected thinking, teaching, and problem solving in chemical engineering. Gibbs ’ research used what was advanced mathematics in his time but remained at what we call ‘‘ classical physics ’’ today since he predated quantum mechanics. Gibbs is especially noteworthy in that he carried out research in the United States at a time when the turmoil of the U.S. Civil War and settling in the West were not as conducive to research as was the case in Europe in the late 1800s. However, Gibbs had spent a year each in Paris, Berlin, and Heidelberg and had written contact with foremost scientists in Europe.
eBook - PDFGeneral Chemistry I as a Second Language
Mastering the Fundamental Skills
- David R. Klein(Author)
- 2015(Publication Date)
- Wiley(Publisher)
133 CHAPTER 5 ENERGY AND ENTHALPY This chapter is the first part of a much larger topic called thermodynamics. You will revisit thermodynamics in more detail during the second semester of chemistry. The topics in this chapter will lay the foundation that you need for the second semester, so it is important to master the terms, concepts, and problem-solving techniques in this chapter. If you don’t get these topics down now, you will find yourself strug- gling with thermodynamics next semester. In one sentence, thermodynamics is the study of energy and its interconver- sions. Put more simply, thermodynamics is the study of how, why, and when en- ergy can be transferred from one place to another. In this chapter we focus on how energy is transferred. In the second semester of chemistry, you will learn about why and when energy is transferred (entropy and Free Energy). The first half of this chapter will focus on theory, terminology, and analogies. The second half of the chapter will focus on problem-solving techniques. 5.1 ENERGY We will start off our discussion with the different types of energy, but as we do so, keep in mind that we have still not defined what energy really is. We will get to the definition a bit later. Energy can be classified into the following categories: kinetic energy and po- tential energy. Kinetic energy is energy associated with motion (or velocity), and potential energy is energy associated with position. Let’s start with kinetic energy. When a soccer ball is in motion, it has kinetic energy (you might even remember the term 1 ⁄ 2 mv 2 from your high school physics class). When the soccer ball hits another ball, it will transfer some of its energy to the other ball. Molecules can do the same thing. A molecule in motion has kinetic energy that it can transfer when it collides with another molecule.
eBook - PDFStatistical Physics of Biomolecules
An Introduction
- Daniel M. Zuckerman(Author)
- 2010(Publication Date)
- CRC Press(Publisher)
Once the conditions are set, the pertinent Free Energy is all of the following: (1) It is the logarithm of that partition function obeying the conditions, multiplied by ( − k B T ) . (2) It quantifies the amount of work that can be done by the system under the condi-tions. (3) It tends to become minimized if the system is initially out of equilibrium, as we will see below. The notation for free energies uses the conditions (the parameters held constant) as “natural” variables. Therefore, each Free Energy should always be written in a way that make the corresponding conditions obvious. The free energies we have discussed are • F ( T , V , N ) = E − TS , the Helmholtz Free Energy • G ( T , P , N ) = F + P V , the Gibbs Free Energy • ( T , V , μ ) = F − N μ , the grand canonical potential Note that the total energy is the sum of kinetic and potential contributions—that is, E = KE + U . 174 Statistical Physics of Biomolecules: An Introduction 7.6.2 F REE E NERGIES A RE “S TATE F UNCTIONS ” As we have seen more than enough times, a Free Energy is defined from a partition function. In turn, the partition function simply integrates to some number once the natural variables—for example, T , V , N in Z ( T , V , N ) —are fixed. Thus, the natural variables define the Free Energy: once you specify the variables, the Free Energy is some number. A Free Energy is called a “state function” because it comes from a definite integral depending only on the equilibrium “state” of the system—that is, only on the natural variables. It does not depend on how that state was reached (maybe by a weird nonequilibrium process) because this doesn’t change the equilibrium integral. The notion of a state function is important in understanding Free Energy differ-ences generally, and especially in understanding “thermodynamic cycles,” which we shall study in the context of binding in Chapter 9.
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