🧊Thermodynamics II Unit 10 – Phase Equilibrium and Stability

Key Concepts and Definitions

• Phase equilibrium occurs when two or more phases coexist in a stable state with no net transfer of mass or energy between them
• Stability refers to a system's ability to maintain its state when subjected to small perturbations
• Gibbs free energy ($G$) is a thermodynamic potential that determines the stability and spontaneity of a process at constant temperature and pressure
• Chemical potential ($\mu$) is the partial molar Gibbs free energy and represents the change in $G$ when a component is added to a system
• Fugacity ($f$) is a measure of a component's tendency to escape from a phase and is related to its chemical potential
• Activity ($a$) is the effective concentration of a component in a mixture and is defined as the ratio of fugacity to a standard reference state
• Raoult's law states that the vapor pressure of an ideal solution is directly proportional to the mole fraction of each component
• Henry's law describes the solubility of a gas in a liquid and states that the partial pressure of the gas is directly proportional to its mole fraction in the liquid phase

Thermodynamic Foundations

• The first law of thermodynamics states that energy cannot be created or destroyed, only converted from one form to another
• The second law of thermodynamics introduces the concept of entropy ($S$) and states that the total entropy of an isolated system always increases over time
• Fundamental equation of thermodynamics relates the change in internal energy ($U$) to changes in entropy, volume ($V$), and composition ($n_i$) through the equation $dU = TdS - PdV + \sum_i \mu_i dn_i$
• Maxwell relations are derived from the fundamental equation and provide useful relationships between thermodynamic properties
• Gibbs-Duhem equation relates changes in chemical potentials of components in a mixture and is given by $\sum_i n_i d\mu_i = 0$ at constant temperature and pressure
• Partial molar properties describe the contribution of each component to the total property of a mixture and are defined as $\bar{X}_i = (\partial X / \partial n_i)_{T,P,n_j}$
• Ideal gas law relates pressure ($P$), volume ($V$), amount of substance ($n$), and temperature ($T$) through the equation $PV = nRT$, where $R$ is the ideal gas constant
• Equations of state (van der Waals, Redlich-Kwong, Peng-Robinson) are used to describe the behavior of real gases and account for intermolecular interactions and volume of molecules

Phase Diagrams and Their Interpretation

• Phase diagrams graphically represent the equilibrium states of a system as a function of thermodynamic variables (pressure, temperature, composition)
• Pure component phase diagrams show the regions of stability for solid, liquid, and gas phases and the transitions between them (melting, boiling, sublimation)
• Binary phase diagrams depict the equilibrium between two components and can include features such as eutectic points, azeotropes, and immiscibility regions
• Ternary phase diagrams represent systems with three components and are often used in liquid-liquid extraction and crystallization processes
• Critical point is the highest temperature and pressure at which vapor and liquid phases can coexist in equilibrium
• Triple point is the unique condition where all three phases (solid, liquid, gas) coexist in equilibrium
• Tie lines connect the compositions of coexisting phases in a two-phase region and are used to determine the relative amounts of each phase
• Lever rule is a graphical method for calculating the relative amounts of phases in a two-phase region based on the position along a tie line

Gibbs Phase Rule and Its Applications

• Gibbs phase rule relates the number of components ($C$), phases ($P$), and degrees of freedom ($F$) in a system through the equation $F = C - P + 2$
• Degrees of freedom represent the number of independent variables that can be changed without altering the number of phases in equilibrium
• For a single-component system, the maximum number of phases that can coexist in equilibrium is three (at the triple point)
• In binary systems, the maximum number of phases that can coexist in equilibrium is four (at the invariant point)
• Gibbs phase rule is used to determine the variance of a system and the number of intensive variables needed to fully specify its state
• Phase rule is applied in the construction and interpretation of phase diagrams, as well as in the design of separation processes (distillation, extraction, crystallization)
• Gibbs-Konovalov theorem states that the partial derivative of temperature with respect to pressure along a phase boundary is equal to the difference in molar entropies divided by the difference in molar volumes of the two phases
• Clapeyron equation relates the slope of a phase boundary in a $P$-$T$ diagram to the latent heat and volume change associated with the phase transition

Equilibrium Conditions and Stability Criteria

• Thermodynamic equilibrium is achieved when a system minimizes its Gibbs free energy at constant temperature and pressure
• Chemical equilibrium occurs when the chemical potentials of each component are equal in all phases, i.e., $\mu_i^\alpha = \mu_i^\beta = ... = \mu_i^\pi$
• Mechanical equilibrium requires that the pressures in all phases are equal, i.e., $P^\alpha = P^\beta = ... = P^\pi$
• Thermal equilibrium is reached when the temperatures of all phases are equal, i.e., $T^\alpha = T^\beta = ... = T^\pi$
• Stability of a phase is determined by the curvature of the Gibbs free energy surface with respect to composition
  • A phase is stable if the Gibbs free energy surface is convex (positive curvature)
  • A phase is metastable if the Gibbs free energy surface has a local minimum but is not the global minimum
  • A phase is unstable if the Gibbs free energy surface is concave (negative curvature)
• Spinodal curve separates the metastable and unstable regions in a phase diagram and is defined by the points where the second derivative of Gibbs free energy with respect to composition is zero
• Binodal curve (or coexistence curve) represents the thermodynamic equilibrium between two phases and is determined by the equality of chemical potentials and pressures

Calculation Methods and Problem-Solving Techniques

• Equation of state methods use models (van der Waals, Redlich-Kwong, Peng-Robinson) to calculate thermodynamic properties and phase equilibria of real fluids
• Activity coefficient models (Margules, van Laar, Wilson, NRTL, UNIQUAC) are used to describe the non-ideality of liquid mixtures and predict liquid-liquid equilibria
• Gibbs free energy minimization is a powerful technique for determining the equilibrium state of a system by minimizing the total Gibbs free energy subject to mass balance constraints
• Flash calculations determine the equilibrium composition and amounts of vapor and liquid phases for a given feed composition, temperature, and pressure
• Dew point and bubble point calculations find the temperature or pressure at which the first drop of liquid or vapor appears in a system
• Azeotropic point calculations identify compositions where the liquid and vapor phases have the same composition and cannot be separated by conventional distillation
• Residue curve maps are graphical tools used to visualize the composition trajectories of batch distillation processes and aid in the design of separation schemes
• Reactive phase equilibria involve the simultaneous consideration of chemical reactions and phase equilibria, requiring the minimization of Gibbs free energy subject to both mass balance and reaction equilibrium constraints

Real-World Applications and Case Studies

• Distillation is a common separation process that relies on the differences in volatility between components and is used in the production of petroleum products, alcoholic beverages, and specialty chemicals
• Liquid-liquid extraction exploits the differences in solubility of components in two immiscible liquid phases and is used in the recovery of antibiotics, rare earth elements, and nuclear fuel reprocessing
• Crystallization is a purification technique that separates a solid phase from a liquid solution based on differences in solubility and is used in the production of pharmaceuticals, semiconductors, and specialty chemicals
• Gas absorption is a process where a soluble gas is removed from a gas stream by contacting it with a liquid solvent and is used in acid gas removal, carbon capture, and air pollution control
• Adsorption is a surface phenomenon where molecules adhere to a solid adsorbent and is used in gas purification, water treatment, and energy storage applications
• Steam distillation is a special case of distillation where steam is used to volatilize and separate temperature-sensitive compounds, such as essential oils and fragrances
• Supercritical fluid extraction uses fluids above their critical point to selectively extract components from a mixture and is used in the decaffeination of coffee, extraction of hops for beer brewing, and cleaning of semiconductor wafers
• Reservoir engineering applies phase equilibria concepts to the production of oil and gas from underground reservoirs, considering the effects of pressure, temperature, and composition on the flow and recovery of hydrocarbons

Common Pitfalls and Misconceptions

• Assuming that all mixtures behave ideally can lead to significant errors in phase equilibrium calculations, especially for systems with strong intermolecular interactions or large differences in component properties
• Neglecting the effect of pressure on liquid phase properties can result in inaccurate predictions of liquid-liquid equilibria and miscibility gaps
• Using inappropriate or inconsistent reference states for chemical potentials and fugacities can cause discrepancies in equilibrium calculations and lead to incorrect conclusions
• Extrapolating thermodynamic models and parameters outside their range of applicability can produce unreliable results and should be avoided or done with caution
• Overlooking the presence of azeotropes in distillation processes can lead to inefficient or infeasible separation schemes and require alternative strategies, such as pressure swing distillation or extractive distillation
• Ignoring the potential for multiple stable phases in a system can result in the failure to predict important phenomena, such as liquid-liquid immiscibility or solid-solid phase transitions
• Confusing the concepts of stability and metastability can lead to misinterpretation of phase diagrams and incorrect assessment of the feasibility of separation processes
• Neglecting the impact of impurities or trace components on phase equilibria can lead to unexpected behavior and deviations from predicted performance in real-world applications