Theinternal energy of athermodynamic system is theenergy of the system as astate function, measured as the quantity of energy necessary to bring the system from itsstandard internal state to its present internal state of interest, accounting for the gains and losses of energy due to changes in its internal state, including such quantities asmagnetization.[1][2] It excludes thekinetic energy of motion of the system as a whole and thepotential energy of position of the system as a whole, with respect to its surroundings and external force fields. The notion of internal energy was introduced by Clausius as part of the formulation of thefirst law of thermodynamics.
Without athermodynamic process, the internal energy of anisolated system does not change, as expressed in the law ofconservation of energy, a foundation of the first law of thermodynamics.[3] Without transfer of matter, internal energy changes equal the algebraic sum of the heat transferred and the thermodynamic work done.
The internal energy cannot be measured absolutely. Thermodynamics concernschanges in the internal energy, not its absolute value. The processes that change the internal energy are transfers, into or out of the system, of matter, or of energy, asheat, or bythermodynamic work.[4] These processes are measured by changes in the system's properties, such astemperature,entropy,volume,electric polarization, andmolar constitution. The internal energy depends only on the internal state of the system and not on the particular choice from many possible processes by which energy may pass into or out of the system. It is astate variable, athermodynamic potential, and anextensive property.[5]
Thermodynamics defines internal energy macroscopically, for the body as a whole. Instatistical mechanics, the internal energy of a body can be analyzed microscopically in terms of the kinetic energies of microscopic motion of the system's particles fromtranslations,rotations, andvibrations, and of the potential energies associated with microscopic forces, includingchemical bonds.
The internal energy of a system depends on itsentropy S, its volume V and its number of massive particles:U(S,V,{Nj}). It expresses the thermodynamics of a system in theenergy representation. As afunction of state, its arguments are exclusively extensive variables of state, which is why it is sometimes called acardinal function of state. Alongside the internal energy, the other cardinal function of state of a thermodynamic system is its entropy, as a function,S(U,V,{Nj}), of the same list of extensive variables of state, except that the entropy,S, is replaced in the list by the internal energy,U. It expresses theentropy representation.[7][8][9]
Each cardinal function is a monotonic function of each of itsnatural orcanonical variables. Each provides itscharacteristic orfundamental equation, for exampleU =U(S,V,{Nj}), that by itself contains all thermodynamic information about the system. The fundamental equations for the two cardinal functions can in principle be interconverted by solving, for example,U =U(S,V,{Nj}) forS, to getS =S(U,V,{Nj}).
In contrast,Legendre transformations are necessary to derive fundamental equations for other thermodynamic potentials andMassieu functions. The entropy as a function only of extensive state variables is the one and onlycardinal function of state for the generation of Massieu functions. It is not itself customarily designated a 'Massieu function', though rationally it might be thought of as such, corresponding to the term 'thermodynamic potential', which includes the internal energy.[8][10][11]
For real and practical systems, explicit expressions of the fundamental equations are almost always unavailable, but the functional relations exist in principle. Formal, in principle, manipulations of them are valuable for the understanding of thermodynamics.
The internal energy of a given state of the system is determined relative to that of a standard state of the system, by adding up the macroscopic transfers of energy that accompany a change of state from the reference state to the given state:
where denotes the difference between the internal energy of the given state and that of the reference state,and the are the various energies transferred to the system in the steps from the reference state to the given state.It is the energy needed to create the given state of the system from the reference state. From a non-relativistic microscopic point of view, it may be divided into microscopic potential energy,, and microscopic kinetic energy,, components:
The microscopic kinetic energy of a system arises as the sum of the motions of all the system's particles with respect to the center-of-mass frame, whether it be the motion ofatoms,molecules,atomic nuclei,electrons, or other particles (often expressed as thermal engergy/ temperature). The microscopic potential energy algebraic summative components are those of thechemical andnuclear particle bonds (e.g.intermolecular forces), and the physical force fields within the system, such as due to internalinducedelectric ormagneticdipolemoment (e.g.intramolecular forces), as well as the energy ofdeformation of solids (stress-strain). Usually, the split into microscopic kinetic and potential energies is outside the scope of macroscopic thermodynamics.
Internal energy does not include the energy due to motion or location of a system as a whole. That is to say, it excludes any kinetic or potential energy the body may have because of its motion or location in externalgravitational,electrostatic, orelectromagneticfields. It does, however, include the contribution of such a field to the energy due to the coupling of the internal degrees of freedom of the system with the field. In such a case, the field is included in the thermodynamic description of the object in the form of an additional external parameter.
For practical considerations in thermodynamics or engineering, it is rarely necessary, convenient, nor even possible, to consider all energies belonging to the total intrinsic energy of a sample system, such as the energy given by the equivalence of mass. Typically, descriptions only include components relevant to the system under study. Indeed, in most systems under consideration, especially through thermodynamics, it is impossible to calculate the total internal energy.[12] Therefore, a convenient null reference point may be chosen for the internal energy.
At any temperature greater thanabsolute zero, microscopic potential energy and kinetic energy are constantly converted into one another, but the sum remains constant in anisolated system (cf. table). In the classical picture of thermodynamics, kinetic energy vanishes at zero temperature and the internal energy is purely potential energy. However, quantum mechanics has demonstrated that even at zero temperature particles maintain a residual energy of motion, thezero point energy. A system at absolute zero is merely in its quantum-mechanical ground state, the lowest energy state available. At absolute zero a system of given composition has attained its minimum attainableentropy.
The microscopic kinetic energy portion of the internal energy gives rise to the temperature of the system.Statistical mechanics relates the pseudo-random kinetic energy of individual particles to the mean kinetic energy of the entire ensemble of particles comprising a system. Furthermore, it relates the mean microscopic kinetic energy to the macroscopically observed empirical property that is expressed as temperature of the system. While temperature is an intensive measure, this energy expresses the concept as an extensive property of the system, often referred to as thethermal energy,[13][14] The scaling property between temperature and thermal energy is the entropy change of the system.
Statistical mechanics considers any system to be statistically distributed across an ensemble ofmicrostates. In a system that is in thermodynamic contact equilibrium with a heat reservoir, each microstate has an energy and is associated with a probability. The internal energy is themean value of the system's total energy, i.e., the sum of all microstate energies, each weighted by its probability of occurrence:
Thermodynamics is chiefly concerned with the changes in internal energy.
For a closed system, with mass transfer excluded, the changes in internal energy are due to heat transfer and due tothermodynamic work doneby the system on its surroundings.[note 1] Accordingly, the internal energy change for a process may be written
When a closed system receives energy as heat, this energy increases the internal energy. It is distributed between microscopic kinetic and microscopic potential energies. In general, thermodynamics does not trace this distribution. In an ideal gas all of the extra energy results in a temperature increase, as it is stored solely as microscopic kinetic energy; such heating is said to besensible.
A second kind of mechanism of change in the internal energy of a closed system changed is in its doing ofwork on its surroundings. Such work may be simply mechanical, as when the system expands to drive a piston, or, for example, when the system changes its electric polarization so as to drive a change in the electric field in the surroundings.
If the system is not closed, the third mechanism that can increase the internal energy is transfer of substance into the system. This increase, cannot be split into heat and work components.[4] If the system is so set up physically that heat transfer and work that it does are by pathways separate from and independent of matter transfer, then the transfers of energy add to change the internal energy:
If a system undergoes certain phase transformations while being heated, such as melting and vaporization, it may be observed that the temperature of the system does not change until the entire sample has completed the transformation. The energy introduced into the system while the temperature does not change is calledlatent energy orlatent heat, in contrast to sensible heat, which is associated with temperature change.
Thermodynamics often uses the concept of theideal gas for teaching purposes, and as an approximation for working systems. The ideal gas consists of particles considered as point objects that interact only by elastic collisions and fill a volume such that theirmean free path between collisions is much larger than their diameter. Such systems approximatemonatomic gases such ashelium and othernoble gases. For an ideal gas the kinetic energy consists only of thetranslational energy of the individual atoms. Monatomic particles do not possess rotational or vibrational degrees of freedom, and are notelectronically excited to higher energies except at very hightemperatures.
Therefore, the internal energy of an ideal gas depends solely on its temperature (and the number of gas particles):. It is not dependent on other thermodynamic quantities such as pressure or density.
The internal energy of an ideal gas is proportional to itsamount of substance (number of moles) and to its temperature
where is the isochoric (at constant volume)molar heat capacity of the gas; is constant for an ideal gas. The internal energy of any gas (ideal or not) may be written as a function of the three extensive properties,, (entropy, volume,number of moles). In case of the ideal gas it is in the following way[15]
where is an arbitrary positive constant and where is theuniversal gas constant. It is easily seen that is a linearlyhomogeneous function of the three variables (that is, it isextensive in these variables), and that it is weaklyconvex. Knowing temperature and pressure to be the derivatives theideal gas law immediately follows as below:
The above summation of all components of change in internal energy assumes that a positive energy denotes heat added to the system or the negative of work done by the system on its surroundings.[note 1]
This relationship may be expressed ininfinitesimal terms using the differentials of each term, though only the internal energy is anexact differential.[16]: 33 For a closed system, with transfers only as heat and work, the change in the internal energy is
For example, the mechanical work done by the system may be related to thepressure andvolume change. The pressure is the intensive generalized force, while the volume change is the extensive generalized displacement:
This defines the direction of work,, to be energy transfer from the working system to the surroundings, indicated by a positive term.[note 1] Taking the direction of heat transfer to be into the working fluid and assuming areversible process, the heat is
In case of an ideal gas, we can derive that, i.e. the internal energy of an ideal gas can be written as a function that depends only on the temperature.
Proof of pressure independence for an ideal gas
The expression relating changes in internal energy to changes in temperature and volume is
The equation of state is the ideal gas law
Solve for pressure:
Substitute in to internal energy expression:
Take the derivative of pressure with respect to temperature:
The partial derivative of with respect to can be evaluated if the equation of state is known. From the fundamental thermodynamic relation, it follows that the differential of theHelmholtz free energy is given by
In addition to including the entropy and volume terms in the internal energy, a system is often described also in terms of the number of particles or chemical species it contains:
where are the molar amounts of constituents of type in the system. The internal energy is anextensive function of the extensive variables,, and the amounts, the internal energy may be written as a linearlyhomogeneous function of first degree:[17]
where is a factor describing the growth of the system. The differential internal energy may be written as
which shows (or defines) temperature to be the partial derivative of with respect to entropy and pressure to be the negative of the similar derivative with respect to volume,
and where the coefficients are thechemical potentials for the components of type in the system. The chemical potentials are defined as the partial derivatives of the internal energy with respect to the variations in composition:
As conjugate variables to the composition, the chemical potentials areintensive properties, intrinsically characteristic of the qualitative nature of the system, and not proportional to its extent. Under conditions of constant and, because of the extensive nature of and its independent variables, usingEuler's homogeneous function theorem, the differential may be integrated and yields an expression for the internal energy:
that arises from changing the composition of the system at constant temperature and pressure. For a single component system, the chemical potential equals the Gibbs energy per amount of substance, i.e. particles or moles according to the original definition of the unit for.
For anelastic medium the potential energy component of the internal energy has an elastic nature expressed in terms of thestress and strain involved in elastic processes. InEinstein notation for tensors, with summation over repeated indices, for unit volume, the infinitesimal statement is
For a linearly elastic material, the stress is related to the strain by
where the are the components of the 4th-rank elastic constant tensor of the medium.
Elastic deformations, such assound, passing through a body, or other forms of macroscopic internal agitation or turbulent motion create states when the system is not in thermodynamic equilibrium. While such energies of motion continue, they contribute to the total energy of the system; thermodynamic internal energy pertains only when such motions have ceased.
James Joule studied the relationship between heat, work, and temperature. He observed that friction in a liquid, such as caused by its agitation with work by a paddle wheel, caused an increase in its temperature, which he described as producing aquantity of heat. Expressed in modern units, he found that c. 4186 joules of energy were needed to raise the temperature of one kilogram of water by one degree Celsius.[19]
^abcThis article uses the sign convention of the mechanical work as often defined in engineering, which is different from the convention used in physics and chemistry; in engineering, work performed by the system against the environment, e.g., a system expansion, is taken to be positive, while in physics and chemistry, it is taken to be negative.
^Grubbström, Robert W. (1985). "Towards a Generalized Exergy Concept". In van Gool, W.; Bruggink, J.J.C. (eds.).Energy and time in the economic and physical sciences. North-Holland. pp. 41–56.ISBN978-0444877482.
^Adkins, C. J. (Clement John) (1983).Equilibrium thermodynamics (3rd ed.). Cambridge [Cambridgeshire]: Cambridge University Press.ISBN0-521-25445-0.OCLC9132054.
^Landau, Lev Davidovich; Lifshit︠s︡, Evgeniĭ Mikhaĭlovich; Pitaevskiĭ, Lev Petrovich; Sykes, John Bradbury; Kearsley, M. J. (1980).Statistical physics. Oxford. p. 70.ISBN0-08-023039-3.OCLC3932994.{{cite book}}: CS1 maint: location missing publisher (link)
Callen, H. B. (1960/1985), Thermodynamics and an Introduction to Thermostatistics, (first edition 1960), second edition 1985, John Wiley & Sons, New York,ISBN0-471-86256-8.
Crawford, F. H. (1963).Heat, Thermodynamics, and Statistical Physics, Rupert Hart-Davis, London, Harcourt, Brace & World, Inc.
Haase, R. (1971). Survey of Fundamental Laws, chapter 1 ofThermodynamics, pages 1–97 of volume 1, ed. W. Jost, ofPhysical Chemistry. An Advanced Treatise, ed. H. Eyring, D. Henderson, W. Jost, Academic Press, New York, lcn 73–117081.
Lewis, Gilbert Newton; Randall, Merle: Revised by Pitzer, Kenneth S. & Brewer, Leo (1961).Thermodynamics (2nd ed.). New York, NY USA: McGraw-Hill Book Co.ISBN978-0-07-113809-3.{{cite book}}:ISBN / Date incompatibility (help)CS1 maint: multiple names: authors list (link)
