8.2 Chemical potential and Gibbs energy

Chemical potential and Gibbs energy are key concepts in thermodynamics. They help us understand how substances behave in mixtures and during phase changes. These ideas are crucial for predicting equilibrium conditions and spontaneous processes in various systems.

In this section, we'll explore how chemical potential relates to partial molar quantities and Gibbs energy. We'll also dive into the Gibbs-Duhem equation, equilibrium constants, and the Clausius-Clapeyron equation. These tools are essential for analyzing real-world thermodynamic problems.

Chemical Potential and Partial Molar Quantities

Definition and Significance of Chemical Potential

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• Chemical potential ($\mu$) represents the change in Gibbs free energy of a system when a component is added or removed while keeping temperature, pressure, and other component amounts constant
• Serves as a driving force for mass transfer and phase transitions in multicomponent systems
• Helps determine the direction of spontaneous processes and equilibrium conditions
• Plays a crucial role in understanding phase equilibria, chemical reactions, and transport phenomena

Partial Molar Quantities and Their Relationship to Chemical Potential

• Partial molar quantities describe the contribution of each component to the total thermodynamic property of a mixture
• Partial molar volume ($\bar{V}_i$) represents the change in total volume of a mixture when one mole of component $i$ is added at constant temperature, pressure, and amounts of other components
• Partial molar enthalpy ($\bar{H}_i$) and partial molar entropy ($\bar{S}_i$) are defined similarly for enthalpy and entropy, respectively
• Chemical potential is related to partial molar quantities through the relationship $\mu_i = \bar{G}_i = \bar{H}_i - T\bar{S}_i + P\bar{V}_i$, where $\bar{G}_i$ is the partial molar Gibbs free energy

Gibbs-Duhem Equation and Its Applications

• Gibbs-Duhem equation relates changes in intensive properties (temperature, pressure, and chemical potentials) of a system at equilibrium
• For a multicomponent system, the equation is given by $SdT - VdP + \sum_i n_i d\mu_i = 0$, where $S$ is entropy, $V$ is volume, $n_i$ is the number of moles of component $i$, and $\mu_i$ is the chemical potential of component $i$
• Provides a constraint on the changes in chemical potentials of components in a mixture
• Useful in deriving relationships between thermodynamic properties and determining the number of independent variables in a system
• Helps in the calculation of activity coefficients and fugacity coefficients in non-ideal mixtures (excess properties)

Gibbs Free Energy and Equilibrium

Gibbs Free Energy and Its Relation to Equilibrium

• Gibbs free energy ($G$) is a thermodynamic potential that measures the maximum reversible work that can be performed by a system at constant temperature and pressure
• Defined as $G = H - TS$, where $H$ is enthalpy, $T$ is temperature, and $S$ is entropy
• Systems tend to minimize their Gibbs free energy to achieve equilibrium
• At equilibrium, the Gibbs free energy of a system is at its minimum, and the change in Gibbs free energy ($dG$) is zero
• For a closed system at constant temperature and pressure, the condition for equilibrium is $dG = 0$

Equilibrium Constant and Its Relationship to Gibbs Free Energy

• Equilibrium constant ($K$) is a measure of the position of equilibrium in a chemical reaction
• Relates the concentrations or partial pressures of reactants and products at equilibrium
• For a general reaction $aA + bB \rightleftharpoons cC + dD$, the equilibrium constant is given by $K = \frac{[C]^c[D]^d}{[A]^a[B]^b}$, where $[X]$ represents the concentration or partial pressure of species $X$
• Gibbs free energy change of a reaction ($\Delta G$) is related to the equilibrium constant through the equation $\Delta G = -RT \ln K$, where $R$ is the universal gas constant and $T$ is the absolute temperature
• This relationship allows the determination of the direction of a reaction and the calculation of equilibrium compositions

Clausius-Clapeyron Equation and Phase Equilibrium

• Clausius-Clapeyron equation describes the relationship between the vapor pressure of a substance and temperature at phase equilibrium
• For a phase transition (e.g., liquid-vapor), the equation is given by $\frac{dP}{dT} = \frac{\Delta H_{vap}}{T\Delta V}$, where $P$ is the vapor pressure, $T$ is the absolute temperature, $\Delta H_{vap}$ is the enthalpy of vaporization, and $\Delta V$ is the change in volume during the phase transition
• Helps in understanding the effect of temperature on vapor pressure and boiling point
• Useful in the design of distillation columns, heat exchangers, and other equipment involving phase changes
• Can be integrated to obtain the vapor pressure as a function of temperature, assuming $\Delta H_{vap}$ is constant (e.g., $\ln \frac{P_2}{P_1} = -\frac{\Delta H_{vap}}{R}(\frac{1}{T_2} - \frac{1}{T_1})$)

Fugacity and Activity

Fugacity and Its Role in Non-Ideal Systems

• Fugacity ($f$) is a thermodynamic property that represents the effective partial pressure of a component in a non-ideal mixture
• Introduced to account for the deviations from ideal behavior in real systems
• For an ideal gas, fugacity is equal to the partial pressure ($f_i = P_i$)
• Fugacity coefficient ($\phi_i$) is defined as the ratio of fugacity to partial pressure, $\phi_i = \frac{f_i}{P_i}$
• In non-ideal systems, fugacity is related to the chemical potential through the equation $\mu_i - \mu_i^0 = RT \ln \frac{f_i}{f_i^0}$, where $\mu_i^0$ and $f_i^0$ are the chemical potential and fugacity at a reference state, respectively

Activity and Activity Coefficient

• Activity ($a_i$) is a measure of the effective concentration of a component in a non-ideal mixture
• Defined as the ratio of fugacity to a reference fugacity, $a_i = \frac{f_i}{f_i^0}$, where $f_i^0$ is the fugacity at a reference state (usually pure component at the same temperature and pressure)
• Activity coefficient ($\gamma_i$) is the ratio of activity to mole fraction, $\gamma_i = \frac{a_i}{x_i}$, where $x_i$ is the mole fraction of component $i$
• For an ideal solution, activity coefficients are equal to unity ($\gamma_i = 1$)
• Activity coefficients account for the non-ideality of mixtures and are used in the calculation of phase equilibria, chemical reaction equilibria, and other thermodynamic properties
• Various models and equations, such as Margules, van Laar, and Wilson equations, are used to estimate activity coefficients based on experimental data or molecular interactions