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Inthermodynamics, thethermodynamic free energy is one of thestate functions of athermodynamic system. The change in the free energy is the maximum amount ofwork that the system can perform in aprocess atconstant temperature, and its sign indicates whether the process is thermodynamically favorable or forbidden. Since free energy usually containspotential energy, it is not absolute but depends on the choice of a zero point. Therefore, only relative free energy values, or changes in free energy, are physically meaningful.
The free energy is the portion of anyfirst-law energy that isavailable to perform thermodynamic work at constant temperature,i.e., work mediated bythermal energy. Free energy is subject toirreversible loss in the course of such work.[1] Since first-law energy is always conserved, it is evident that free energy is an expendable,second-law kind of energy. Several free energy functions may be formulated based on system criteria. Free energyfunctions areLegendre transforms of theinternal energy.
TheGibbs free energy is given byG =H −TS, whereH is theenthalpy,T is theabsolute temperature, andS is theentropy.H =U +pV, whereU is the internal energy,p is thepressure, andV is the volume.G is the most useful forprocesses involving a system atconstantpressurep and temperatureT, because, in addition to subsuming any entropy change due merely to heat, a change inG also excludes thep dV work needed to "make space for additional molecules" produced by various processes. Gibbs free energy change therefore equals work not associated with system expansion or compression, at constant temperature and pressure, hence its utility tosolution-phase chemists, including biochemists.
The historically earlierHelmholtz free energy is defined in contrast asA =U −TS. Its change is equal to the amount ofreversible work done on, or obtainable from, a system at constantT. Thus its appellation "work content", and the designationA (from German Arbeit 'work'). Since it makes no reference to any quantities involved in work (such asp andV), the Helmholtz function is completely general: its decrease is the maximum amount of work which can be doneby a system at constant temperature, and it can increase at most by the amount of work doneon a system isothermally. The Helmholtz free energy has a special theoretical importance since it is proportional to thelogarithm of thepartition function for thecanonical ensemble instatistical mechanics. (Hence its utility tophysicists; and to gas-phase chemists and engineers, who do not want to ignorep dV work.)
Historically, the term 'free energy' has been used for either quantity. Inphysics,free energy most often refers to the Helmholtz free energy, denoted byA (orF), while inchemistry,free energy most often refers to the Gibbs free energy. The values of the two free energies are usually quite similar and the intended free energy function is often implicit in manuscripts and presentations.
The basic definition of "energy" is a measure of a body's (in thermodynamics, the system's) ability to cause change. For example, when a person pushes a heavy box a few metres forward, that person exerts mechanical energy, also known as work, on the box over a distance of a few meters forward. The mathematical definition of this form of energy is the product of the force exerted on the object and the distance by which the box moved (Work = Force × Distance). Because the person changed the stationary position of the box, that person exerted energy on that box. The work exerted can also be called "useful energy", because energy was converted from one form into the intended purpose, i.e. mechanical use. For the case of the person pushing the box, the energy in the form of internal (or potential) energy obtained through metabolism was converted into work to push the box. This energy conversion, however, was not straightforward: while some internal energy went into pushing the box, some was diverted away (lost) in the form of heat (transferred thermal energy).
For a reversible process, heat is the product of the absolute temperature and the change in entropy of a body (entropy is a measure of disorder in a system). The difference between the change in internal energy, which is, and the energy lost in the form of heat is what is called the "useful energy" of the body, or the work of the body performed on an object. In thermodynamics, this is what is known as "free energy". In other words, free energy is a measure of work (useful energy) a system can perform at constant temperature.
Mathematically, free energy is expressed as
This expression has commonly been interpreted to mean that work is extracted from the internal energy while represents energy not available to perform work. However, this is incorrect. For instance, in an isothermal expansion of an ideal gas, the internal energy change is and the expansion work is derived exclusively from the term supposedly not available to perform work. But it is noteworthy that the derivative form of the free energy: (for Helmholtz free energy) does indeed indicate that a spontaneous change in a non-reactive system's free energy (NOT the internal energy) comprises the available energy to do work (compression in this case) and the unavailable energy.[2][3][4] Similar expression can be written for the Gibbs free energy change.[5][3][4]
In the 18th and 19th centuries, thetheory of heat, i.e., that heat is a form of energy having relation to vibratory motion, was beginning to supplant both thecaloric theory, i.e., that heat is a fluid, and thefour element theory, in which heat was the lightest of the four elements. In a similar manner, during these years, heat was beginning to be distinguished into different classification categories, such as "free heat", "combined heat", "radiant heat",specific heat,heat capacity, "absolute heat", "latent caloric", "free" or "perceptible" caloric (calorique sensible), among others.
In 1780, for example,Laplace andLavoisier stated: “In general, one can change the first hypothesis into the second by changing the words ‘free heat, combined heat, and heat released’ into ‘vis viva, loss of vis viva, and increase of vis viva.’" In this manner, the total mass of caloric in a body, calledabsolute heat, was regarded as a mixture of two components; the free or perceptible caloric could affect a thermometer, whereas the other component, the latent caloric, could not.[6] The use of the words "latent heat" implied a similarity to latent heat in the more usual sense; it was regarded as chemicallybound to the molecules of the body. In theadiabaticcompression of a gas, the absolute heat remained constant but the observed rise in temperature implied that some latent caloric had become "free" or perceptible.
During the early 19th century, the concept of perceptible or free caloric began to be referred to as "free heat" or "heat set free". In 1824, for example, the French physicistSadi Carnot, in his famous "Reflections on the Motive Power of Fire", speaks of quantities of heat ‘absorbed or set free’ in different transformations. In 1882, the German physicist and physiologistHermann von Helmholtz coined the phrase ‘free energy’ for the expression, in which the change inA (orG) determines the amount of energy ‘free’ forwork under the given conditions, specifically constant temperature.[7]: 235
Thus, in traditional use, the term "free" was attached to Gibbs free energy for systems at constant pressure and temperature, or to Helmholtz free energy for systems at constant temperature, to mean ‘available in the form of useful work.’[8] With reference to the Gibbs free energy, we need to add the qualification that it is the energyfree for non-volume work and compositional changes.[9]: 77–79
An increasing number of books and journal articles do not include the attachment "free", referring toG as simply Gibbs energy (and likewise for theHelmholtz energy). This is the result of a 1988IUPAC meeting to set unified terminologies for the international scientific community, in which the adjective ‘free’ was supposedly banished.[10][11][12] This standard, however, has not yet been universally adopted, and many published articles and books still include the descriptive ‘free’.[citation needed]
Just like the general concept of energy, free energy has a few definitions suitable for different conditions. In physics, chemistry, and biology, these conditions are thermodynamic parameters (temperature, volume, pressure, etc.). Scientists have come up with several ways to define free energy. The mathematical expression of Helmholtz free energy is:
This definition of free energy is useful for gas-phase reactions or in physics when modeling the behavior of isolated systems kept at a constant volume. For example, if a researcher wanted to perform a combustion reaction in a bomb calorimeter, the volume is kept constant throughout the course of a reaction. Therefore, the heat of the reaction is a direct measure of the free energy change,. In solution chemistry, on the other hand, most chemical reactions are kept at constant pressure. Under this condition, the heat of the reaction is equal to the enthalpy change of the system. Under constant pressure and temperature, the free energy in a reaction is known as Gibbs free energy.
These functions have a minimum in chemical equilibrium, as long as certain variables (, and or) are held constant. In addition, they also have theoretical importance in derivingMaxwell relations. Work other thanp dV may be added, e.g., forelectrochemical cells, orf dx work inelastic materials and inmuscle contraction. Other forms of work which must sometimes be considered arestress-strain,magnetic, as inadiabatic demagnetization used in the approach toabsolute zero, and work due to electricpolarization. These are described bytensors.
In most cases of interest there are internaldegrees of freedom and processes, such aschemical reactions andphase transitions, which create entropy. Even for homogeneous "bulk" materials, the free energy functions depend on the (often suppressed)composition, as do all properthermodynamic potentials (extensive functions), including the internal energy.
| Name | Symbol | Formula | Natural variables |
|---|---|---|---|
| Internal energy | |||
| Helmholtz free energy | |||
| Enthalpy | |||
| Gibbs free energy | |||
| Landau potential, or grand potential | , |
is the number of molecules (alternatively,moles) of type in the system. If these quantities do not appear, it is impossible to describe compositional changes. Thedifferentials for processes at uniform pressure and temperature are (assuming only work):
whereμi is thechemical potential for theithcomponent in the system. The second relation is especially useful at constant and, conditions which are easy to achieve experimentally, and which approximately characterizeliving creatures. Under these conditions, it simplifies to
Any decrease in the Gibbs function of a system is the upper limit for anyisothermal,isobaric work that can be captured in the surroundings, or it may simply bedissipated, appearing as times a corresponding increase in the entropy of the system and/or its surrounding.
An example issurface free energy, the amount of increase of free energy when the area of surface increases by every unit area.
Thepath integral Monte Carlo method is a numerical approach for determining the values of free energies, based on quantum dynamical principles.
For a reversible isothermal process, ΔS =qrev/T and therefore the definition ofA results in
This tells us that the change in free energy equals the reversible or maximum work for a process performed at constant temperature. Under other conditions, free-energy change is not equal to work; for instance, for a reversible adiabatic expansion of an ideal gas,. Importantly, for a heat engine, including theCarnot cycle, the free-energy change after a full cycle is zero,, while the engine produces nonzero work. It is important to note that for heat engines and other thermal systems, the free energies do not offer convenient characterizations; internal energy and enthalpy are the preferred potentials for characterizing thermal systems.
According to thesecond law of thermodynamics, for any process that occurs in a closed system, theinequality of Clausius, ΔS >q/Tsurr, applies. For a process at constant temperature and pressure without non-PV work, this inequality transforms into. Similarly, for a process at constant temperature and volume,. Thus, a negative value of the change in free energy is a necessary condition for a process to be spontaneous; this is the most useful form of the second law of thermodynamics in chemistry. In chemical equilibrium at constantT andp without electrical work, dG = 0.
The quantity called "free energy" is a more advanced and accurate replacement for the outdated termaffinity, which was used by chemists in previous years to describe theforce that causedchemical reactions. The term affinity, as used in chemical relation, dates back to at least the time ofAlbertus Magnus.[13]
From the 1998 textbookModern Thermodynamics[14] by Nobel Laureate and chemistry professorIlya Prigogine we find: "As motion was explained by the Newtonian concept of force, chemists wanted a similar concept of ‘driving force’ for chemical change. Why do chemical reactions occur, and why do they stop at certain points? Chemists called the ‘force’ that caused chemical reactions affinity, but it lacked a clear definition."
During the entire 18th century, the dominant view with regard to heat and light was that put forth byIsaac Newton, called theNewtonian hypothesis, which states that light and heat are forms of matter attracted or repelled by other forms of matter, with forces analogous to gravitation or to chemical affinity.
In the 19th century, the French chemistMarcellin Berthelot and the Danish chemistJulius Thomsen had attempted to quantify affinity usingheats of reaction. In 1875, after quantifying the heats of reaction for a large number of compounds, Berthelot proposed theprinciple of maximum work, in which all chemical changes occurring without intervention of outside energy tend toward the production of bodies or of a system of bodies which liberate heat.
In addition to this, in 1780Antoine Lavoisier andPierre-Simon Laplace laid the foundations ofthermochemistry by showing that the heat given out in a reaction is equal to the heat absorbed in the reverse reaction. They also investigated thespecific heat andlatent heat of a number of substances, and amounts of heat given out in combustion. In a similar manner, in 1840 Swiss chemistGermain Hess formulated the principle that the evolution of heat in a reaction is the same whether the process is accomplished in one-step process or in a number of stages. This is known asHess' law. With the advent of themechanical theory of heat in the early 19th century, Hess's law came to be viewed as a consequence of the law ofconservation of energy.
Based on these and other ideas, Berthelot and Thomsen, as well as others, considered the heat given out in the formation of a compound as a measure of the affinity, or the work done by the chemical forces. This view, however, was not entirely correct. In 1847, the English physicistJames Joule showed that he could raise the temperature of water by turning a paddle wheel in it, thus showing that heat and mechanical work were equivalent or proportional to each other, i.e., approximately,dW ∝dQ. This statement came to be known as themechanical equivalent of heat and was a precursory form of thefirst law of thermodynamics.
By 1865, the German physicistRudolf Clausius had shown that thisequivalence principle needed amendment. That is, one can use the heat derived from acombustion reaction in a coal furnace to boil water, and use this heat to vaporize steam, and then use the enhanced high-pressure energy of the vaporized steam to push a piston. Thus, we might naively reason that one can entirely convert the initial combustion heat of the chemical reaction into the work of pushing the piston. Clausius showed, however, that we must take into account the work that the molecules of the working body, i.e., the water molecules in the cylinder, do on each other as they pass or transform from one step of orstate of theengine cycle to the next, e.g., from () to (). Clausius originally called this the "transformation content" of the body, and then later changed the name toentropy. Thus, the heat used to transform the working body of molecules from one state to the next cannot be used to do external work, e.g., to push the piston. Clausius defined thistransformation heat as.
In 1873,Willard Gibbs publishedA Method of Geometrical Representation of the Thermodynamic Properties of Substances by Means of Surfaces, in which he introduced the preliminary outline of the principles of his new equation able to predict or estimate the tendencies of various natural processes to ensue when bodies or systems are brought into contact. By studying the interactions of homogeneous substances in contact, i.e., bodies, being in composition part solid, part liquid, and part vapor, and by using a three-dimensionalvolume-entropy-internal energy graph, Gibbs was able to determine three states of equilibrium, i.e., "necessarily stable", "neutral", and "unstable", and whether or not changes will ensue. In 1876, Gibbs built on this framework by introducing the concept ofchemical potential so to take into account chemical reactions and states of bodies that are chemically different from each other. In his own words, to summarize his results in 1873, Gibbs states:
If we wish to express in a single equation the necessary and sufficient condition ofthermodynamic equilibrium for a substance when surrounded by a medium of constantpressurep and temperatureT, this equation may be written:
δ(ε −Tη +pν) = 0whenδ refers to the variation produced by any variations in thestate of the parts of the body, and (when different parts of the body are in different states) in the proportion in which the body is divided between the different states. The condition of stable equilibrium is that the value of the expression in the parenthesis shall be a minimum.
In this description, as used by Gibbs,ε refers to theinternal energy of the body,η refers to theentropy of the body, andν is thevolume of the body.
Hence, in 1882, after the introduction of these arguments by Clausius and Gibbs, the German scientistHermann von Helmholtz stated, in opposition to Berthelot and Thomas' hypothesis that chemical affinity is a measure of the heat of reaction of chemical reaction as based on the principle of maximal work, that affinity is not the heat given out in the formation of a compound but rather it is the largest quantity of work which can be gained when the reaction is carried out in a reversible manner, e.g., electrical work in a reversible cell. The maximum work is thus regarded as the diminution of the free, or available, energy of the system (Gibbs free energyG atT = constant,P = constant orHelmholtz free energyA atT = constant,V = constant), whilst the heat given out is usually a measure of the diminution of the total energy of the system (Internal energy). Thus,G orA is the amount of energy "free" for work under the given conditions.
Up until this point, the general view had been such that: “all chemical reactions drive the system to a state of equilibrium in which the affinities of the reactions vanish”. Over the next 60 years, the term affinity came to be replaced with the term free energy. According to chemistry historian Henry Leicester, the influential 1923 textbookThermodynamics and the Free Energy of Chemical Reactions byGilbert N. Lewis andMerle Randall led to the replacement of the term "affinity" by the term "free energy" in much of the English-speaking world.
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