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# Gibbs-Duhem equation

The Gibbs-Duhem equation in thermodynamics describes the relationship between changes in chemical potential for components in a thermodynamical system [1] :

$\sum_{i=1}^I N_i\mathrm{d}\mu_i = - S\mathrm{d}T + V\mathrm{d}p \,$

where $N_i\,$ is the number of moles of component $i\,$, $\mathrm{d}\mu_i\,$ the incremental increase in chemical potential for this component, $S\,$ the entropy, $T\,$ the absolute temperature, $V\,$ volume and $p\,$ the pressure. It shows that in thermodynamics intensive properties are not independent but related, making it a mathematical statement of the state postulate. When pressure and temperature are variable, only $I-1\,$ of $I\,$ components have independent values for chemical potential and Gibbs' phase rule follows. The law is named after Josiah Gibbs and Pierre Duhem.

## Derivation

Deriving the Gibbs-Duhem equation from basic thermodynamic state equations is straightforward[2]. The total differential of the Gibbs free energy $G\,$ in terms of its natural variables is

$\mathrm{d}G =\left. \frac{\partial G}{\partial p}\right | _{T,N}\mathrm{d}p +\left. \frac{\partial G}{\partial T}\right | _{p,N}\mathrm{d}T +\sum_{i=1}^I \left. \frac{\partial G}{\partial N_i}\right | _{p,N_{j \neq i}}\mathrm{d}N_i \,$.

With the substitution of two of the Maxwell relations and the definition of chemical potential, this is transformed into:[3]

$\mathrm{d}G =V \mathrm{d}p-S \mathrm{d}T +\sum_{i=1}^I \mu_i \mathrm{d}N_i \,$

As shown in the Gibbs free energy article, the chemical potential is just another name for the partial molar (or just partial, depending on the units of N) Gibbs free energy, thus

$G = \sum_{i=1}^I \mu_i N_i \,$.

The total differential of this expression is[3]

$\mathrm{d}G = \sum_{i=1}^I \mu_i \mathrm{d}N_i + \sum_{i=1}^I N_i \mathrm{d}\mu_i \,$

Subtracting the two expressions for the total differential of the Gibbs free energy gives the Gibbs-Duhem relation:[3]

$\sum_{i=1}^I N_i\mathrm{d}\mu_i = - S\mathrm{d}T + V\mathrm{d}p \,$

## Applications

By normalizing the above equation by the extent of a system, such as the total number of moles, the Gibbs-Duhem equation provides a relationship between the intensive variables of the system. For a simple system with $I\,$ different components, there will be $I+1\,$ independent parameters or "degrees of freedom". For example, a gas cylinder filled with nitrogen is at room temperature (298 K) and at 2500 psi, we can determine the gas density, entropy or any other intensive thermodynamic variable. If instead the cylinder contains a nitrogen/oxygen mixture, we require an additional piece of information, usually the ratio of oxygen-to-nitrogen.

If multiple phases of matter are present, the chemical potential across a phase boundary are equal.[4] Combining expressions for the Gibbs-Duhem equation in each phase and assuming systematic equilibrium (i.e. that the temperature and pressure is constant throughout the system), we recover the Gibbs' phase rule.

One particularly useful expression arises when considering binary solutions[5]. At constant P (isobaric) and T (isothermal) it becomes:

$0= N_1\mathrm{d}\mu_1 + N_2\mathrm{d}\mu_2 \,$

or, normalizing by total number of moles in the system $N_1 + N_2 \,$, substituting in the definition of activity coefficient $\gamma\$ and using the identity $x_1 + x_2 = 1\,$:

$x_1 \left. \frac{\mathrm{d}ln \gamma_1}{\mathrm{d}x_1} \right |_{p,T} =x_2 \left. \frac{\mathrm{d}ln \gamma_2}{\mathrm{d}x_2} \right |_{p,T} \,$

This equation is instrumental in the calculation of thermodynamically consistent and thus more accurate expressions for the vapor pressure of a fluid mixture from limited experimental data.

## References

1. ^ A to Z of Thermodynamics Pierre Perrot ISBN 0198565569
2. ^ Fundamentals of Engineering Thermodynamics, 3rd Edition Michael J. Moran and Howard N. Shapiro, p. 538 ISBN 0-471-07681-3
3. ^ a b c Salzman, William R. (2001-08-21). Open Systems (English). Chemical Thermodynamics. University of Arizona. Archived from the original on 2007-07-07. Retrieved on 2007-10-11.
4. ^ Fundamentals of Engineering Thermodynamics, 3rd Edition Michael J. Moran and Howard N. Shapiro, p. 710 ISBN 0-471-07681-3
5. ^ The Properties of Gases and Liquids, 5th Edition Poling, Prausnitz and O'Connell, p. 8.13, ISBN 0-07-011682-2