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Statistical mechanics
Bose Einstein condensate.png
Statistical thermodynamics
Kinetic theory

In thermodynamics, the internal energy of a thermodynamic system, or a body with well-defined boundaries, denoted by U, or sometimes E, is the total of the kinetic energy due to the motion of particles (translational, rotational, vibrational) and the potential energy associated with the vibrational and electric energy of atoms within molecules or crystals. It includes the energy in all of the chemical bonds, and the energy of the free, conduction electrons in metals.

One can also calculate the internal energy of electromagnetic or black body radiation. It is a state function of a system, and is an extensive quantity. The SI unit of energy is the joule although other historical, conventional units are still in use, such as the (small and large) calorie for heat.

One can have a corresponding intensive thermodynamic property called specific internal energy, commonly symbolized by the lower-case letter u, which is internal energy per mass of the substance in question. As such, the SI unit of specific internal energy would be the J/kg. If intensive internal energy is expressed on an amount of substance basis, then it could be referred to as molar internal energy and the unit would be the J/mol.



Internal energy does not include the translational or rotational kinetic energy of a body as a whole. It excludes any potential energy a body may have because of its location in external gravitational or electrostatic field, although the potential energy it has in a field due to an induced electric or magnetic dipole moment does count, as does the energy of deformation of solids (stress-strain).

The principle of equipartition of energy in classical statistical mechanics states that each molecular quadratic degree of freedom receives 1/2 kT of energy, [1] a result which had to be modified when quantum mechanics explained certain anomalies, such as discrepancies in the observed specific heats of crystals when the expected thermal energy per degree of freedom is less than the energy necessary to move that degree of freedom up one quantum energy level.

For monoatomic helium and other noble gases, the internal energy consists only of the translational kinetic energy of the individual atoms. Monoatomic particles, of course, do not (sensibly) rotate or vibrate, and are not electronically excited to higher energies except at very high temperatures.

From the standpoint of statistical mechanics, the internal energy is equal to the ensemble average of the total energy of the system. It is also called Intrinsic energy.


Internal energy is the sum of all forms of energy of a system. It is related to the molecular structure and the degree of molecular activity and may be viewed as the sum of kinetic and potential energies of the molecules; it is composed of the following types of energies:[2]

Type Composition of Internal Energy (U)
Sensible energy the portion of the internal energy of a system associated with kinetic energies (molecular translation, rotation, and vibration; electron translation and spin; and nuclear spin) of the molecules.
Latent energy the internal energy associated with the phase of a system.
Chemical energy the internal energy associated with the chemical bonds in a molecule.
Nuclear energy the very large amount of energy associated with the bonds within the nucleus of the atom itself.
Energy interactions those types of energies not stored in the system (e.g. heat transfer, mass transfer, and work), but which are recognized at the system boundary as they cross it, which represent gains or losses by a system during a process.

Sensible energy and latent energy together can be referred to as thermal energy.

The first law of thermodynamics

The internal energy is essentially defined by the first law of thermodynamics which states that energy is conserved:

 \Delta U = Q + W + W' \,


ΔU is the change in internal energy of a system during a process.
Q is heat added to a system (measured in joules in SI); that is, a positive value for Q represents heat flow into a system while a negative value denotes heat flow out of a system.
W is the mechanical work done on a system (measured in joules in SI)
W' is energy added by all other processes

The first law may be stated equivalently in infinitesimal terms as:

 \mathrm{d}U = \delta Q + \delta W + \delta W'\,

where the terms now represent infinitesimal amounts of the respective quantities. The d before the internal energy function indicates that it is an exact differential. In other words it is a state function or a value which can be assigned to the system. On the other hand, the δ's before the other terms indicate that they describe increments of energy which are not state functions but rather they are processes by which the internal energy is changed. (See the discussion in the first law article.)

From a microscopic point of view, the internal energy may be found in many different forms. For a gas it may consist almost entirely of the kinetic energy of the gas molecules. It may also consist of the potential energy of these molecules in a gravitational, electric, or magnetic field. For any material, solid, liquid or gaseous, it may also consist of the potential energy of attraction or repulsion between the individual molecules of the material.

Expressions for the internal energy

The internal energy may be expressed in terms of other thermodynamic parameters. Each term is composed of an intensive variable (a generalized force) and its conjugate infinitesimal extensive variable (a generalized displacement).

For example, for a non-viscous fluid, the mechanical work done on the system may be related to the pressure p and volume V. The pressure is the intensive generalized force, while the volume is the extensive generalized displacement:

Taking the default direction of work, W, to be from the working fluid to the surroundings,

\delta W = p \mathrm{d}V\,.
p is the pressure
V is the volume

Taking the default direction of heat transfer, Q, to be into the working fluid and assuming a reversible process, we have

\delta Q = T \mathrm{d}S\,.
T is temperature
S is entropy

The above two equations in the first law of thermodynamics imply for a closed system:

\mathrm{d}U = \delta Q - \delta W = T\mathrm{d}S-p\mathrm{d}V\,

If we also include the dependence on the numbers of particles in the system, the internal energy function may be written as U(S,V,N_{1}, N_{2},\ldots) where the Nj are the numbers of particles of type j in the system. U is an extensive function, so when considered as a function of the extensive variables variables S, V, and the particle numbers N_{1}, N_{2},\ldots, we have:

U(\alpha S,\alpha V,\alpha N_{1},\alpha N_{2},\ldots )=\alpha U(S,V,N_{1},N_{2},\ldots)\,

From Euler's homogeneous function theorem we may now write the internal energy as:

U=TS-pV + \sum_{i}\mu_{i}N_{i}\,

where the μi are the chemical potentials for the particles of type i in the system. These are defined as:

\mu_i = \left( \frac{\partial U}{\partial N_i} \right)_{S,V, N_{j \ne i}}

For an elastic substance the mechanical term must be replaced by the more general expression involving the stress σij and strain \varepsilon_{ij}. The infinitesimal statement is:


where Einstein notation has been used for the tensors, in which there is a summation over all repeated indices in the product term. The Euler theorem yields for the internal energy (Landau & Lifshitz 1986):


For a linearly elastic material, the stress can be related to the strain by:

\sigma_{ij}=C_{ijkl} \varepsilon_{kl}

Change in internal energy due to change in temperature and volume or pressure

The expressions given above for the internal energy involves the entropy. In practice one often wants to know the change in internal energy of a substance as a function of the change in temperature and volume, or as a function of the change in temperature and pressure.

To express dU in terms of dT and dV, we substitute

dS = \left(\frac{\partial S}{\partial T}\right)_{V}dT + \left(\frac{\partial S}{\partial V}\right)_{T}dV \,

in the fundamental thermodynamic relation

dU = T dS - P dV\,

This gives:

dU = T\left(\frac{\partial S}{\partial T}\right)_{V}dT +\left[T\left(\frac{\partial S}{\partial V}\right)_{T} - P\right]dV\,

The term T\left(\frac{\partial S}{\partial T}\right)_{V} is the heat capacity at constant volume CV.

The partial derivative of S with respect to V can be evaluated if the equation of state is known. From the fundamental thermodynamic relation, it follows that the differential of the Helmholtz free energy A is given by:

dA = -S dT - P dV\,

The symmetry of second derivatives of A with respect to T and V yields the Maxwell relation:

\left(\frac{\partial S}{\partial V}\right)_{T} = \left(\frac{\partial P}{\partial T}\right)_{V} \,

This gives the expression:

dU =C_{V}dT +\left[T\left(\frac{\partial P}{\partial T}\right)_{V} - P\right]dV\,\,\text{ (1)}\,

This is useful if the equation of state is known. In case of an ideal gas, P = NkT / V which implies that dU = CvdT, i.e. the internal energy of an ideal gas can be written as a function that depends only on the temperature.

When dealing with fluids or solids, an expression in terms of the temperature and pressure is usually more useful. The partial derivative of the pressure with respect to temperature at constant volume can be expressed in terms of the coefficient of thermal expansion

\alpha \equiv \frac{1}{V}\left(\frac{\partial V}{\partial T}\right)_{P}\,

and the isothermal compressibility

\beta_{T} \equiv -\frac{1}{V}\left(\frac{\partial V}{\partial P}\right)_{T}\,

by writing:

dV = \left(\frac{\partial V}{\partial P}\right)_{T}dP + \left(\frac{\partial V}{\partial T}\right)_{P} dT = V\left(\alpha dT-\beta_{T}dP \right)\,\,\text{ (2)} \,

and equating dV to zero and solving for the ratio dP/dT. This gives:

\left(\frac{\partial P}{\partial T}\right)_{V}= -\frac{\left(\frac{\partial V}{\partial T}\right)_{P}}{\left(\frac{\partial V}{\partial P}\right)_{T}}= \frac{\alpha}{\beta_{T}}\,\,\text{ (3)}\,

Substituting (2) and (3) in (1) gives:

dU = \left(C_{P}-\alpha P V\right)dT +\left(\beta_{T}P-\alpha T\right)VdP\,

where we have used that the heat capacity at constant pressure is related to the heat capacity at constant volume according to:

C_{P} = C_{V} + V T\frac{\alpha^{2}}{\beta_{T}}\,

as shown here.

Equipartition theorem

The equipartition theorem yields simple expressions for the internal energy. In case of an ideal gas, the internal energy is exactly given by the kinetic energy of the constituent particles. The internal energy per particle is equivalent to the average translational kinetic energy of each particle. Ignoring quantum effects, this is given by equipartition of energy.[3]

According to the equipartition theorem, the thermal energy of a molecule in a thermal bath is

U_{thermal} = f \cdot \tfrac{1}{2} kT

where f is the number of degrees of freedom, T is the temperature, and k is Boltzmann's constant. For example, for a monatomic ideal gas, each particle has three degrees of freedom, and thus

U_{thermal, monatomic} = \tfrac{3}{2} kT.

When the spacing between the energy levels of a particular degree of freedom becomes of the order of k T or less, the energy in that degree of freedom becomes less than given by the equipartition theorem and it vanishes exponentially as k T becomes much less than the energy difference. The system is then frozen in the ground state.

For a gas at room temperature at normal densities, the vibrational degrees of freedom are usually frozen, while the rotational and vibrational degrees of freedom can be treated classically. Quantum effects for the translational degrees of freedom become important when the specific volume per particle is of the same order or smaller than Λ3 where Λ is the thermal de Broglie wavelength. In this regime quantum statistical effects become important. Depending on whether the molecules are Fermions or Bosons, the gas will become a degenerate Fermi gas or a Bose-Einstein condensate, respectively. E.g., at room temperature, the electrons in a metal form a degenerate Fermi gas, the internal energy per electron is of the order of kTf with Tf the Fermi temperature which can be of the order of 80,000 K. So, in this case, the equipartition theorem underestimates the internal energy by many orders of magnitude.


James Joule studied the relationship between heat, work, and temperature. He observed that if he did mechanical work on a fluid, such as water, by agitating the fluid, its temperature increased. He proposed that the mechanical work he was doing on the system was converted to "thermal energy". Specifically, he found that 4200 joules of energy were needed to raise the temperature of a kilogram of water by one degree Celsius.


  • Lewis, Gilbert Newton; Randall, Merle: Revised by Pitzer, Kenneth S. & Brewer, Leo (1961). Thermodynamics (2nd Edition ed.). New York, NY USA: McGraw-Hill Book Co.. ISBN 0-07-113809-9. 


  1. ^ Reif, Frederick (1965). Statistical Physics. New York: McGraw-Hill Book Company. pp. 246–250. 
  2. ^ Cengel, Yungus, A.; Boles, Michael (2002). Thermodynamics - An Engineering Approach, 4th ed.. McGraw-Hill. pp. 17–18. ISBN 0-07-238332-1. 
  3. ^ Thermal energy – Hyperphysics

See also


Thermal energy is a form of energy that manifests itself as an increase of temperature. It is also the sum of sensible heat and latent heat.



The thermal energy of a single particle in a thermal bath is:

U_{thermal} = f \cdot \frac{1}{2} kT.

where f refers to the degrees of freedom, T refers to the temperature, and k to Boltzmann's constant. For example, a monatomic particle in an ideal gas has three degrees of freedom, and thus,

U_{thermal, monatomic} = \frac{3}{2} kT.

The total thermal energy is the sum of the thermal energies of all particles in the system. Thus, for a system of N particles,

U_{thermal} = N \cdot f \cdot \frac{1}{2} kT.

Note that Uthermal is rarely the total energy of a system; for instance, there can be static energy that doesn't change with temperature, such as potential energy, bond energy or rest energy (E=mc2).

History of the term

The term was first used explicitly by James Prescott Joule, who studied the relationship between heat, work, and temperature. He observed that if he did mechanical work on a fluid such as water, by agitating the fluid, its temperature increased. He proposed that the mechanical work he was doing on the system was converted to "thermal energy." Specifically, he found that 4200 joules of energy were needed to raise the temperature of a kilogram of water by one degree Celsius. It was discovered that the law may have had something to do with Christian Lausberg's Law of thermal heat.

Thermal energy in an ideal gas

Thermal energy is most easily defined in the context of an ideal gas. In a monatomic ideal gas, the thermal energy is exactly given by the kinetic energy of the constituent particles.[citation needed]

Other definitions

Thermal energy per particle is also called the average translational kinetic energy possessed by free particles given by equipartition of energy.[1]

Thermal energy is the difference between the internal energy of an object and the amount that it would have at absolute zero.[citation needed] It includes the quantity of kinetic energy due to the motion of the internal particles of an object, and is increased by heating and reduced by cooling.

See also


  1. ^ Thermal energy – Hyperphysics

Simple English

In thermal physics, thermal energy is the type of energy that has a system that increases with its temperature. In thermodynamics, thermal energy is the internal energy present in a system in a state of thermodynamic equilibrium because of its temperature.[1] That is, heat is defined as a spontaneous flow of energy (energy in transit) from one object to another, caused by a difference in temperature between two objects; so objects do not possess heat.[2]

Other pages


  1. Thermal energy - Britannica
  2. Schroeder, Daniel, R. (2000). Thermal Physics. New York: Addison Wesley Longman. ISBN 0201380277. 


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