Table of thermodynamic equations

For list of mathematical notation used in these equations, see mathematical notation.

Main article: List of thermodynamic properties

This article is a summary of common equations and quantities in thermodynamics (see thermodynamic equations for more elaboration). SI units are used for absolute temperature, not celsius or fahrenheit.

Definitions

Main articles: List of thermodynamic properties, Thermodynamic potential, Free entropy, Defining equation (physical chemistry)

Many of the definitions below are also used in the thermodynamics of chemical reactions.

General basic quantities

Quantity (Common Name/s)	(Common) Symbol/s	SI Units	Dimension
Number of molecules	N	dimensionless	dimensionless
Number of moles	n	mol	[N]
Temperature	T	K	[Θ]
Heat Energy	Q, q	J	[M][L]2[T]−2
Latent Heat	QL	J	[M][L]2[T]−2

General derived quantities

Quantity (Common Name/s)	(Common) Symbol/s	Defining Equation	SI Units	Dimension
Thermodynamic beta, Inverse temperature	β	$\beta = 1/k_B T \,\!$	J−1	[T]2[M]−1[L]−2
Thermodynamic Temperature	τ	$\tau = k_B T \,\!$ $\tau = k_B \left (\partial U/\partial S \right )_{N} \,\!$ $1/\tau = 1/k_B \left (\partial S/\partial U \right )_{N} \,\!$	J	[M] [L]2 [T]-2
Entropy	S	$S = -k_B\sum_i p_i\ln p_i$ $S = \left (\partial F/\partial T \right )_{V} \,\!$ , $S = \left (\partial G/\partial T \right )_{N,P} \,\!$	J K−1	[M][L]2[T]−2 [Θ]−1
Pressure	P	$P = - \left (\partial F/\partial V \right )_{T,N} \,\!$ $P = - \left (\partial U/\partial V \right )_{S,N} \,\!$	Pa	M L−1T−2
Internal Energy	U	$U = \sum_i E_i \!$	J	[M][L]2[T]−2
Enthalpy	H	$H = U+pV\,\!$	J	[M][L]2[T]−2
Partition Function	Z		dimensionless	dimensionless
Gibbs free energy	G	$G = H - TS \,\!$	J	[M][L]2[T]−2
Chemical potential (of component i in a mixture)	μi	$\mu_i = \left (\partial U/\partial N_i \right )_{N_{j \neq i}, S, V } \,\!$ $\mu_i = \left (\partial F/\partial N_i \right )_{T, V } \,\!$, where F is not proportional to N because μi depends on pressure. $\mu_i = \left (\partial G/\partial N_i \right )_{T, P } \,\!$, where G is proportional to N (as long as the molar ratio composition of the system remains the same) because μi depends only on temperature and pressure and composition. $\mu_i/\tau = -1/k_B \left (\partial S/\partial N_i \right )_{U,V} \,\!$	J	[M][L]2[T]−2
Helmholtz free energy	A, F	$F = U - TS \,\!$	J	[M][L]2[T]−2
Landau potential, Landau Free Energy, Grand potential	Ω, ΦG	$\Omega = U - TS - \mu N\,\!$	J	[M][L]2[T]−2
Massieu Potential, Helmholtz free entropy	Φ	$\Phi = S - U/T \,\!$	J K−1	[M][L]2[T]−2 [Θ]−1
Planck potential, Gibbs free entropy	Ξ	$\Xi = \Phi - pV/T \,\!$	J K−1	[M][L]2[T]−2 [Θ]−1

Thermal properties of matter

Main Articles: Heat capacity, Thermal expansion

Quantity (common name/s)	(Common) symbol/s	Defining equation	SI units	Dimension
General heat/thermal capacity	C	$C = \partial Q/\partial T\,\!$	J K −1	[M][L]2[T]−2 [Θ]−1
Heat capacity (isobaric)	Cp	$C_{p} = \partial H/\partial T\,\!$	J K −1	[M][L]2[T]−2 [Θ]−1
Specific heat capacity (isobaric)	Cmp	$C_{mp} = \partial^2 Q/\partial m \partial T \,\!$	J kg−1 K−1	[L]2[T]−2 [Θ]−1
Molar specific heat capacity (isobaric)	Cnp	$C_{np} = \partial^2 Q/\partial n \partial T \,\!$	J K −1 mol−1	[M][L]2[T]−2 [Θ]−1 [N]−1
Heat capacity (isochoric/volumetric)	CV	$C_{V} = \partial Q/\partial T \,\!$	J K −1	[M][L]2[T]−2 [Θ]−1
Specific heat capacity (isochoric)	CmV	$C_{mV} = \partial^2 Q/\partial m \partial T \,\!$	J kg−1 K−1	[L]2[T]−2 [Θ]−1
Molar specific heat capacity (isochoric)	CnV	$C_{nV} = \partial^2 Q/\partial n \partial T \,\!$	J K −1 mol−1	[M][L]2[T]−2 [Θ]−1 [N]−1
Specific latent heat	L	$L = \partial Q/ \partial m \,\!$	J kg−1	[L]2[T]−2
Ratio of isobaric to isochoric heat capacity, heat capacity ratio, adiabatic index	γ	$\gamma = C_p/C_V = c_p/c_V = C_{mp}/C_{mV} \,\!$	dimensionless	dimensionless

Thermal transfer

Main article: Thermal conductivity

Quantity (common name/s)	(Common) symbol/s	Defining equation	SI units	Dimension
Temperature gradient	No standard symbol	$\nabla T \,\!$	K m−1	[Θ][L]−1
Thermal conduction rate, thermal current, thermal/heat flux, thermal power transfer	P	$P = \mathrm{d} Q/\mathrm{d} t \,\!$	W = J s−1	[M] [L]2 [T]−3
Thermal intensity	I	$I = \mathrm{d} P/\mathrm{d} A$	W m−2	[M] [T]−3
Thermal/heat flux density (vector analogue of thermal intensity above)	q	$Q = \iint \mathbf{q} \cdot \mathrm{d}\mathbf{S}\mathrm{d} t \,\!$	W m−2	[M] [T]−3

Equations

The equations in this article are classified by subject.

Phase transitions

Physical situation	Equations
Adiabatic transition	$\Delta Q = 0, \quad \Delta U = -\Delta W\,\!$
Isothermal transition	$\Delta U = 0, \quad \Delta W = \Delta H \,\!$ For an ideal gas $W=kTN \ln(V_2/V_1)\,\!$
Isobaric transition	p1 = p2, p = constant $\Delta W = p \Delta V, \quad \Delta Q = \Delta U + p \delta V\,\!$
Isochoric transition	V1 = V2, V = constant $\Delta W = 0, \quad \Delta Q = \Delta U\,\!$
Adiabatic expansion	$p_1 V_1^{\gamma} = p_2 V_2^{\gamma}\,\!$ $T_1 V_1^{\gamma - 1} = T_2 V_2^{\gamma - 1} \,\!$
Free expansion	$\Delta U = 0\,\!$
Work done by an expanding gas	Process $\Delta W = \int_{V_1}^{V_2} p \mathrm{d}V \,\!$ Net Work Done in Cyclic Processes $\Delta W = \oint_\mathrm{cycle} p \mathrm{d}V \,\!$

Kinetic theory

Ideal gas equations

Physical situation	Nomenclature	Equations
Ideal gas law	p = pressure V = volume of container T = temperature n = number of moles R = Gas constant N = number of molecules k = Boltzmann's constant	$pV = nRT = kTN\,\!$ $\frac{p_1 V_1}{p_2 V_2} = \frac{n_1 T_1}{n_2 T_2} = \frac{N_1 T_1}{N_2 T_2} \,\!$
Pressure of an ideal gas	m = mass of one molecule Mm = molar mass	$p = \frac{Nm \langle v^2 \rangle}{3V} = \frac{nM_m \langle v^2 \rangle}{3V} = \frac{1}{3}\rho \langle v^2 \rangle \,\!$

Ideal gas

Quantity	General Equation	Isobaric Δp = 0	Isochoric ΔV = 0	Isothermal ΔT = 0	Adiabatic $Q=0$
Work W	$\delta W = -p dV\;$	$-p\Delta V\;$	$0\;$	$-nRT\ln\frac{V_2}{V_1}\;$ $nRT\ln\frac{P_2}{P_1}\;$	$\frac{PV^\gamma (V_f^{1-\gamma} - V_i^{1-\gamma}) } {1-\gamma} = C_V \left(T_2 - T_1 \right)$
Heat Capacity C	(as for real gas)	$C_p = \frac{5}{2}nR\;$ (for monatomic ideal gas) $C_p = \frac{7}{2}nR \;$ (for diatomic ideal gas)	$C_V = \frac{3}{2}nR \;$ (for monatomic ideal gas) $C_V = \frac{5}{2}nR \;$ (for diatomic ideal gas)
Internal Energy ΔU	$\Delta U = C_v \Delta T\;$	$Q + W\;$ $Q_p - p\Delta V\;$	$Q\;$ $C_V\left ( T_2-T_1 \right )\;$	$0\;$ $Q=-W\;$	$W\;$ $C_V\left ( T_2-T_1 \right )\;$
Enthalpy ΔH	$H=U+pV\;$	$C_p\left ( T_2-T_1 \right )\;$	$Q_V+V\Delta p\;$	$0\;$	$C_p\left ( T_2-T_1 \right )\;$
Entropy ΔS	$\Delta S = C_v \ln{T_2 \over T_1} + nR \ln{V_2 \over V_1}$ $\Delta S = C_p \ln{T_2 \over T_1} - nR \ln{p_2 \over p_1}$	$C_p\ln\frac{T_2}{T_1}\;$	$C_V\ln\frac{T_2}{T_1}\;$	$nR\ln\frac{V_2}{V_1}\;$ $\frac{Q}{T}\;$	$C_p\ln\frac{V_2}{V_1}+C_V\ln\frac{p_2}{p_1}=0\;$
Constant	$\;$	$\frac{V}{T}\;$	$\frac{p}{T}\;$	$p V\;$	$p V^\gamma\;$

Entropy

• $S = k_B (\ln \Omega)$, where k_B is the Boltzmann constant, and Ω denotes the volume of macrostate in the phase space or otherwise called thermodynamic probability.

• $dS = \frac{\delta Q}{T}$, for reversible processes only

Statistical physics

Below are useful results from the Maxwell–Boltzmann distribution for an ideal gas, and the implications of the Entropy quantity. The distribution is valid for atoms or molecules constituting ideal gases.

Physical situation	Nomenclature	Equations
Maxwell–Boltzmann distribution	v = velocity of atom/molecule, m = mass of each molecule (all molecules are identical in kinetic theory), γ(p) = Lorentz factor as function of momentum (see below) Ratio of thermal to rest mass-energy of each molecule:$\theta = k_B T/mc^2 \,\!$ K2 is the Modified Bessel function of the second kind.	Non-relativistic speeds $P\left ( v \right )=4\pi\left ( \frac{m}{2\pi k_B T} \right )^{3/2} v^2 e^{-mv^2/2 k_B T} \,\!$ Relativistic speeds (Maxwell-Jüttner distribution) $f(p) = \frac{1}{4 \pi m^3 c^3 \theta K_2(1/\theta)} e^{-\gamma(p)/\theta}$
Entropy Logarithm of the density of states	Pi = probability of system in microstate i Ω = total number of microstates	$S = - k_B\sum_i P_i \ln P_i = k_\mathrm{B}\ln \Omega\,\!$ where: $P_i = 1/\Omega\,\!$
Entropy change		$\Delta S = \int_{Q_1}^{Q_2} \frac{\mathrm{d}Q}{T} \,\!$ $\Delta S = k_B N \ln\frac{V_2}{V_1} + N C_V \ln\frac{T_2}{T_1} \,\!$
Entropic force		$\mathbf{F}_\mathrm{S} = -T \nabla S \,\!$
Equipartition theorem	df = degree of freedom	Average kinetic energy per degree of freedom $\langle E_\mathrm{k} \rangle = \frac{1}{2}kT\,\!$ Internal energy $U = d_f \langle E_\mathrm{k} \rangle = \frac{d_f}{2}kT\,\!$

Corollaries of the non-relativistic Maxwell–Boltzmann distribution are below.

Physical situation	Nomenclature	Equations
Mean speed		$\langle v \rangle = \sqrt{\frac{8 k_B T}{\pi m}}\,\!$
Root mean square speed		$v_\mathrm{rms} = \sqrt{\langle v^2 \rangle} = \sqrt{\frac{3k_B T}{m}} \,\!$
Modal speed		$v_\mathrm{mode} = \sqrt{\frac{2k_B T}{m}}\,\!$
Mean free path	σ = Effective cross-section n = Volume density of number of target particles ℓ = Mean free path	$\ell = 1/\sqrt{2} n \sigma \,\!$

Quasi-static and reversible processes

For quasi-static and reversible processes, the first law of thermodynamics is:

$dU=\delta Q - \delta W$

where δQ is the heat supplied to the system and δW is the work done by the system.

Thermodynamic potentials

Main article: Thermodynamic potentials

See also: Maxwell relations

The following energies are called the thermodynamic potentials,

Name	Symbol	Formula	Natural variables
Internal energy	$U$	$\int ( T \text{d}S - p \text{d}V + \sum_i \mu_i \text{d}N_i )$	$S, V, \{N_i\}$
Helmholtz free energy	$F$	$U-TS$	$T, V, \{N_i\}$
Enthalpy	$H$	$U+pV$	$S, p, \{N_i\}$
Gibbs free energy	$G$	$U+pV-TS$	$T, p, \{N_i\}$
Landau Potential (Grand potential)	$\Omega$, $\Phi_\text{G}$	$U - T S -$$\sum_i\,$$\mu_i N_i$	$T, V, \{\mu_i\}$

and the corresponding fundamental thermodynamic relations or "master equations" are:

Potential	Differential
Internal energy	$dU\left(S,V,{N_{i}}\right) = TdS - pdV + \sum_{i} \mu_{i} dN_i$
Enthalpy	$dH\left(S,p,{N_{i}}\right) = TdS + Vdp + \sum_{i} \mu_{i} dN_{i}$
Helmholtz free energy	$dF\left(T,V,{N_{i}}\right) = -SdT - pdV + \sum_{i} \mu_{i} dN_{i}$
Gibbs free energy	$dG\left(T,p,{N_{i}}\right) = -SdT + Vdp + \sum_{i} \mu_{i} dN_{i}$

Maxwell's relations

The four most common Maxwell's relations are:

Physical situation	Nomenclature	Equations
Thermodynamic potentials as functions of their natural variables	$U(S,V)\,$ = Internal energy $H(S,P)\,$ = Enthalpy $F(T,V)\,$ = Helmholtz free energy $G(T,P)\,$ = Gibbs free energy	$\left(\frac{\partial T}{\partial V}\right)_S = -\left(\frac{\partial P}{\partial S}\right)_V = \frac{\partial^2 U }{\partial S \partial V}$ $\left(\frac{\partial T}{\partial P}\right)_S = +\left(\frac{\partial V}{\partial S}\right)_P = \frac{\partial^2 H }{\partial S \partial P}$ $+\left(\frac{\partial S}{\partial V}\right)_T = \left(\frac{\partial P}{\partial T}\right)_V = - \frac{\partial^2 F }{\partial T \partial V}$ $-\left(\frac{\partial S}{\partial P}\right)_T = \left(\frac{\partial V}{\partial T}\right)_P = \frac{\partial^2 G }{\partial T \partial P}$

More relations include the following.

$\left ( {\partial S\over \partial U} \right )_{V,N} = { 1\over T }$	$\left ( {\partial S\over \partial V} \right )_{N,U} = { p\over T }$	$\left ( {\partial S\over \partial N} \right )_{V,U} = - { \mu \over T }$
$\left ( {\partial T\over \partial S} \right )_V = { T \over C_V }$	$\left ( {\partial T\over \partial S} \right )_P = { T \over C_P }$
$-\left ( {\partial p\over \partial V} \right )_T = { 1 \over {VK_T} }$

Other differential equations are:

Name	H	U	G
Gibbs–Helmholtz equation	$H = -T^2\left(\frac{\partial \left(G/T\right)}{\partial T}\right)_p$	$U = -T^2\left(\frac{\partial \left(F/T\right)}{\partial T}\right)_V$	$G = -V^2\left(\frac{\partial \left(F/V\right)}{\partial V}\right)_T$
	$\left(\frac{\partial H}{\partial p}\right)_T = V - T\left(\frac{\partial V}{\partial T}\right)_P$	$\left(\frac{\partial U}{\partial V}\right)_T = T\left(\frac{\partial P}{\partial T}\right)_V - P$

Quantum properties

• $U = N k_B T^2 \left(\frac{\partial \ln Z}{\partial T}\right)_V ~$

• $S = \frac{U}{T} + N * ~ S = \frac{U}{T} + N k_B \ln Z - N k \ln N + Nk ~$ Indistinguishable Particles

where N is number of particles, h is Planck's constant, I is moment of inertia, and Z is the partition function, in various forms:

Degree of freedom	Partition function
Translation	$Z_t = \frac{(2 \pi m k_B T)^\frac{3}{2} V}{h^3}$
Vibration	$Z_v = \frac{1}{1 - e^\frac{-h \omega}{2 \pi k_B T}}$
Rotation	$Z_r = \frac{2 I k_B T}{\sigma (\frac{h}{2 \pi})^2}$ where: σ = 1 (heteronuclear molecules) σ = 2 (homonuclear)

Thermal properties of matter

Coefficients	Equation
Joule-Thomson coefficient	$\mu_{JT} = \left(\frac{\partial T}{\partial p}\right)_H$
Compressibility (constant temperature)	$K_T = -{ 1\over V } \left ( {\partial V\over \partial p} \right )_{T,N}$
Coefficient of thermal expansion (constant pressure)	$\alpha_{p} = \frac{1}{V}\left(\frac{\partial V}{\partial T}\right)_p$
Heat capacity (constant pressure)	$C_p = \left ( {\partial Q_{rev} \over \partial T} \right )_p = \left ( {\partial U \over \partial T} \right )_p + p \left ( {\partial V \over \partial T} \right )_p = \left ( {\partial H \over \partial T} \right )_p = T \left ( {\partial S \over \partial T} \right )_p$
Heat capacity (constant volume)	$C_V = \left ( {\partial Q_{rev} \over \partial T} \right )_V = \left ( {\partial U \over \partial T} \right )_V = T \left ( {\partial S \over \partial T} \right )_V$

Derivation of heat capacity (constant pressure)

Since $\left(\frac{\partial T}{\partial p}\right)_H \left(\frac{\partial p}{\partial H}\right)_T \left(\frac{\partial H}{\partial T}\right)_p = -1$ $\begin{align} \left(\frac{\partial T}{\partial p}\right)_H & = -\left(\frac{\partial H}{\partial p}\right)_T \left(\frac{\partial T}{\partial H}\right)_p \\ & = \frac{-1}{\left(\frac{\partial H}{\partial T}\right)_p} \left(\frac{\partial H}{\partial p}\right)_T \end{align}$ $C_p = \left(\frac{\partial H}{\partial T}\right)_p$ $\Rightarrow \left(\frac{\partial T}{\partial p}\right)_H = -\frac{1}{C_p} \left(\frac{\partial H}{\partial p}\right)_T$

Derivation of heat capacity (constant volume)

Since $dU = \delta Q_{rev} - \delta W_{rev} ,$ (where δWrev is the work done by the system), $\delta S = \frac{\delta Q_{rev}}{T}, \delta W_{rev}= p \delta V$ $d U = T \delta S- p\delta V$ $\left(\frac{\partial U}{\partial T}\right)_V = T\left(\frac{\partial S}{\partial T}\right)_V - p\left(\frac{\partial V}{\partial T}\right)_V ; C_V = \left(\frac{\partial U}{\partial T}\right)_V$ $\Rightarrow C_V = T\left(\frac{\partial S}{\partial T}\right)_V$

Thermal transfer

Physical situation	Nomenclature	Equations
Net intensity emission/absorption	Texternal = external temperature (outside of system) Tsystem = internal temperature (inside system) ε = emmisivity	$I = \sigma \epsilon \left ( T_\mathrm{external}^4 - T_\mathrm{system}^4 \right ) \,\!$
Internal energy of a substance	CV = isovolumetric heat capacity of substance ΔT = temperature change of substance	$\Delta U = N C_V \Delta T\,\!$
Meyer's equation	Cp = isobaric heat capacity CV = isovolumetric heat capacity n = number of moles	$C_p - C_V = nR \,\!$
Effective thermal conductivities	λi = thermal conductivity of substance i λnet = equivalent thermal conductivity	Series $\lambda_\mathrm{net} = \sum_j \lambda_j \,\!$ Parallel $\frac{1}{\lambda}_\mathrm{net} = \sum_j \left ( \frac{1}{\lambda}_j \right ) \,\!$

Thermal efficiencies

Physical situation	Nomenclature	Equations
Thermodynamic engines	η = efficiency W = work done by engine QH = heat energy in higher temperature reservoir QL = heat energy in lower temperature reservoir TH = temperature of higher temp. reservoir TL = temperature of lower temp. reservoir	Thermodynamic engine: $\eta = \left |\frac{W}{Q_H} \right|\,\!$ Carnot engine efficiency: $\eta_c = 1 - \left | \frac{Q_L}{Q_H} \right | = 1-\frac{T_L}{T_H}\,\!$
Refrigeration	K = coefficient of refrigeration performance	Refrigeration performance $K = \left | \frac{Q_L}{W} \right | \,\!$ Carnot refrigeration performance $K_C = \frac{|Q_L|}{|Q_H|-|Q_L|} = \frac{T_L}{T_H-T_L}\,\!$

See also

• Antoine equation
• Bejan number
• Bowen ratio
• Bridgman's equations
• Clausius–Clapeyron relation
• Departure functions
• Duhem–Margules equation
• Ehrenfest equations

• Gibbs–Helmholtz equation
• Gibbs' phase rule
• Kopp's law
• Kopp–Neumann law
• Noro–Frenkel law of corresponding states
• Onsager reciprocal relations
• Stefan number
• Triple product rule

References

• Atkins, Peter and de Paula, Julio Physical Chemistry, 7th edition, W.H. Freeman and Company, 2002 [ISBN 0-7167-3539-3].
  • Chapters 1 - 10, Part 1: Equilibrium.
• Bridgman, P.W., Phys. Rev., 3, 273 (1914).
• Landsberg, Peter T. Thermodynamics and Statistical Mechanics. New York: Dover Publications, Inc., 1990. (reprinted from Oxford University Press, 1978).
• Lewis, G.N., and Randall, M., "Thermodynamics", 2nd Edition, McGraw-Hill Book Company, New York, 1961.
• Reichl, L.E., "A Modern Course in Statistical Physics", 2nd edition, New York: John Wiley & Sons, 1998.
• Schroeder, Daniel V. Thermal Physics. San Francisco: Addison Wesley Longman, 2000 [ISBN 0-201-38027-7].
• Silbey, Robert J., et al. Physical Chemistry. 4th ed. New Jersey: Wiley, 2004.
• Callen, Herbert B. (1985). "Thermodynamics and an Introduction to Themostatistics", 2nd Ed., New York: John Wiley & Sons.