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Entropy of mixing

The entropy of mixing is the increase in thermodynamic entropy that accompanies the combination of two or more distinct substances, such as ideal gases or solutes, resulting from the enhanced disorder and multiplicity of microscopic arrangements in the homogeneous mixture compared to the separated components.^[1] This quantity, often denoted as ΔS_mix, quantifies a key driving force for spontaneous mixing processes in physical chemistry and materials science, where it contributes to the negative Gibbs free energy change under conditions of constant temperature and pressure.^[2] In the context of ideal gases, the entropy of mixing arises during the irreversible diffusion of gases initially confined at the same temperature and pressure behind a removable partition in an insulated container.^[1] Upon mixing, each gas expands freely into the total volume without heat exchange or work done, leading to an entropy increase for the system as ΔS = ∑ n_i R \ln(V_f / V_i), where n_i is the number of moles of gas i, R is the gas constant, V_f is the final total volume, and V_i is the initial volume occupied by gas i.^[3] This expression reflects the greater accessibility of phase space for each molecular species, ensuring ΔS_mix > 0 and underscoring the second law of thermodynamics, as the process cannot reverse spontaneously without external intervention.^[1] For ideal solutions, such as dilute mixtures or lattice models of polymers and alloys, the entropy of mixing is purely configurational and given by ΔS_mix = -n R ∑ x_i \ln x_i, where n is the total number of moles and x_i are the mole fractions of the components (with ∑ x_i = 1).^[2] In these cases, assuming no intermolecular interactions beyond random placement (i.e., ΔH_mix = 0), the positive ΔS_mix solely drives miscibility, as derived from statistical mechanics where the number of ways to arrange species on lattice sites follows a combinatorial form akin to the binomial coefficient.^[2] This entropy term is maximal for equimolar mixtures (x_i = 0.5) and approaches zero as one component dominates, influencing phenomena like phase stability in geochemical systems such as Earth's mantle alloys.^[4] Beyond ideal cases, the entropy of mixing can include excess contributions from vibrational or rotational modes in real mixtures, potentially positive or negative depending on structural ordering, but the ideal configurational component remains a foundational benchmark for predicting solubility and phase separation.^[5] Applications span from atmospheric gas diffusion to polymer blend design, where balancing this entropy against enthalpic interactions determines material properties.^[4]

Thermodynamic Foundations

Definition and Basic Formula

The entropy of mixing, denoted as ΔS_mix, quantifies the increase in entropy that occurs when two or more distinguishable substances are combined at constant temperature and pressure, under the assumption of ideal behavior where there are no intermolecular interactions beyond random collisions.^[6] This change arises from the enhanced configurational disorder as the molecules of each component gain greater freedom to occupy the available space in the mixture compared to their separate states.^[6] For ideal systems, such as dilute solutions or non-interacting gases, the process is spontaneous and reflects the second law of thermodynamics, which dictates that entropy increases in irreversible mixing.^[7] The concept originated in the 19th century within classical thermodynamics, with foundational contributions from Rudolf Clausius and Josiah Willard Gibbs. Clausius introduced the notion of entropy in 1865 as a measure of energy transformation unavailable for work, linking it to processes like diffusion that underpin mixing.^[8] Gibbs further developed the entropy of mixing in his 1875–1878 memoir "On the Equilibrium of Heterogeneous Substances," where he analyzed the entropy change for isothermal mixing of ideal gases as a state function derived from reversible paths.^[9] Their work established entropy of mixing as a key element in understanding chemical equilibrium and phase behavior without relying on molecular theories.^[8] The basic formula for the entropy of mixing in an ideal system with multiple components is given by

\Delta S_\text{mix} = -nR \sum_i x_i \ln x_i,

where n is the total number of moles, R is the gas constant (8.314 J/mol·K), and x_i are the mole fractions of each component (\sum_i x_i = 1).^[6] This expression is derived thermodynamically from the second law by considering a reversible path for the irreversible mixing process: each pure component is imagined to expand isothermally and reversibly from its initial volume to the total mixture volume, with the entropy change for each expansion calculated as \Delta S = n_i R \ln (V_\text{total}/V_i), leading to the summed form after substituting mole fractions via the ideal gas law.^[7] The negative sign ensures the overall value is positive, as \ln x_i < 0 for x_i < 1. The units of ΔS_mix are joules per kelvin (J/K), consistent with entropy as an extensive property.^[6] By convention, ΔS_mix is always positive for the spontaneous mixing of ideal, distinguishable species, signifying the irreversible increase in configurational freedom and aligning with the second law's requirement for processes to proceed toward higher entropy states.^[7]

Gibbs Free Energy of Mixing for Ideal Species

The Gibbs free energy of mixing, \Delta G_{\text{mix}}, quantifies the change in Gibbs free energy upon mixing ideal species at constant temperature and pressure, serving as the key thermodynamic criterion for the spontaneity of the process. It is expressed as \Delta G_{\text{mix}} = \Delta H_{\text{mix}} - T \Delta S_{\text{mix}}, where \Delta H_{\text{mix}} is the enthalpy change of mixing, T is the absolute temperature, and \Delta S_{\text{mix}} is the entropy change of mixing. For ideal mixtures, which include ideal gases and dilute ideal solutions assuming no volume change upon mixing, the enthalpy of mixing is zero (\Delta H_{\text{mix}} = 0) due to the absence of intermolecular interactions beyond those in the pure components.^[10] This simplifies the expression to \Delta G_{\text{mix}} = -T \Delta S_{\text{mix}}. Substituting the entropy of mixing for ideal species, \Delta S_{\text{mix}} = -R \sum_i n_i \ln x_i (where R is the gas constant, n_i is the number of moles of component i, and x_i is its mole fraction), yields the explicit form \Delta G_{\text{mix}} = RT \sum_i n_i \ln x_i.^[11] Since each \ln x_i < 0 for $0 < x_i < 1, \Delta G_{\text{mix}} is inherently negative, indicating that mixing is spontaneous under constant temperature and pressure conditions.^[10] The magnitude of this negative value increases with the number of components and their even distribution in mole fractions, reflecting greater disorder. To assess phase stability, one plots the molar Gibbs free energy of mixing, g_{\text{mix}} = RT \sum_i x_i \ln x_i, versus composition; for ideal binary mixtures, this curve is strictly convex (with positive second derivative RT (1/x + 1/(1-x))), lying entirely below any common tangent line, which precludes phase separation and confirms complete miscibility.^[11] As an illustrative example, consider a binary ideal gas mixture of 1 mol each of two components (total 2 mol) at 298 K, yielding x_1 = x_2 = 0.5. Here, \Delta G_{\text{mix}} = 2RT (0.5 \ln 0.5 + 0.5 \ln 0.5) = RT \ln 0.5 = -RT \ln 2 \approx -1.72 kJ/mol (or -3.44 kJ total), confirming the favorable thermodynamics of mixing.^[10]

Behavior in Solutions

Ideal and Regular Solutions

In ideal solutions, the entropy of mixing follows the standard expression derived from statistical mechanics for indistinguishable particles, ΔS_mix = -nR (x_1 \ln x_1 + x_2 \ln x_2), where n is the total number of moles, R is the gas constant, and x_i are the mole fractions of components i. These solutions exhibit no heat of mixing, ΔH_mix = 0, implying that intermolecular interactions between unlike molecules are identical to those between like molecules, leading to random mixing without energetic preferences. As a result, ideal solutions obey Raoult's law over the entire composition range, where the partial vapor pressure of each component is P_i = x_i P_i^, with P_i^ being the vapor pressure of the pure component. A classic example is the benzene-toluene system, where the structural similarity of the molecules results in nearly ideal behavior, with deviations from Raoult's law being minimal across all compositions. Regular solutions extend the ideal model by incorporating non-ideal enthalpic contributions while retaining the ideal configurational entropy of mixing, ΔS_mix as above. Introduced by Hildebrand in his foundational work on solubility, regular solution theory posits that the enthalpy of mixing arises from differences in cohesive energies between components, expressed as ΔH_mix = B x_1 x_2, where B is the interaction parameter related to pairwise molecular attractions. This parameter B can be estimated using Hildebrand's solubility parameters, δ_i = \sqrt{ΔE_i / V_i}, where ΔE_i is the energy of vaporization and V_i the molar volume of pure component i; miscibility is favored when |δ_1 - δ_2| is small, as the free energy of mixing ΔG_mix = ΔH_mix - T ΔS_mix remains negative. Unlike ideal solutions, regular solutions display activity coefficients γ_i > 1 for positive B (endothermic mixing), but the excess entropy S^E = 0, ensuring the entropic term drives mixing despite enthalpic opposition. The excess Gibbs free energy in regular solutions simplifies to G^E = B x_1 x_2, since the excess enthalpy H^E = B x_1 x_2 and the excess entropy is zero, contrasting with ideal solutions where all excess functions vanish. This non-ideality manifests in phase diagrams as symmetric miscibility gaps for systems with sufficiently large positive B, where the binodal curve exhibits mirror symmetry about x = 0.5 due to the quadratic form of ΔH_mix, leading to phase separation into conjugate phases of equal but opposite deviations from equimolar composition below a critical temperature. Such gaps are observed in polymer blends and certain metal alloys modeled as regular solutions, highlighting how enthalpic interactions modulate solubility without altering the entropic contribution.

Temperature-Dependent Miscibility and Lower Critical Solution Temperature

In partially miscible liquid systems, the lower critical solution temperature (LCST) represents the critical temperature below which the components are fully miscible and above which they phase separate into two liquid phases.^[12] This behavior contrasts with the upper critical solution temperature (UCST), where miscibility decreases upon cooling, and arises primarily from temperature-dependent interactions that render mixing entropically unfavorable at higher temperatures.^[13] In LCST systems, specific interactions such as hydrogen bonding promote miscibility at lower temperatures by forming ordered structures between unlike molecules, but thermal disruption at elevated temperatures leads to an overall decrease in mixing entropy.^[14] The thermodynamic foundation for LCST lies in the Gibbs free energy of mixing, \Delta G_\text{mix}, where the temperature dependence is given by \left( \frac{\partial \Delta G_\text{mix}/T}{\partial (1/T)} \right)_P = \Delta H_\text{mix} and \frac{d \Delta G_\text{mix}}{dT} = -\Delta S_\text{mix}.^[12] For LCST behavior, \Delta H_\text{mix} is typically negative (exothermic mixing favored by attractive interactions), while \Delta S_\text{mix} is negative (due to reduced configurational freedom from ordered associations), causing the -T \Delta S_\text{mix} term to dominate and make \Delta G_\text{mix} positive above the LCST.^[14] This enthalpy-entropy crossover distinguishes LCST from UCST systems, where positive \Delta H_\text{mix} drives demixing upon cooling; in LCST cases, the unfavorable entropy effect intensifies with rising temperature.^[13] A classic example is the nicotine-water system, which exhibits an LCST around 60°C, below which the components mix completely due to hydrogen bonding between nicotine's amine groups and water, but separates above this temperature as thermal energy disrupts these bonds, amplifying the entropic penalty.^[15] Polymer blends, such as poly(ethylene oxide) with poly(methyl methacrylate), also display LCST behavior, where phase separation occurs upon heating due to weakened intermolecular attractions and increased chain entanglements that reduce mixing entropy.^[13] In the Flory-Huggins theory for polymer solutions, the temperature-dependent miscibility is captured by the interaction parameter \chi, which governs the non-ideal contribution to \Delta G_\text{mix}; phase separation ensues when \chi > 2/N for symmetric blends with equal degree of polymerization N (or \chi N > 2).^[16] For LCST systems, \chi increases with temperature, often modeled as \chi = A/T + B where A < 0 (reflecting negative enthalpic interactions) and B > 0 (positive entropic contribution), such that rising T reduces the magnitude of the negative A/T term, pushing \chi above the critical value and inducing demixing.^[17] This form highlights the entropic origins of LCST in polymer contexts, where specific interactions like hydrogen bonding contribute to the temperature sensitivity.^[13]

Statistical Mechanics Derivation

Microscopic Origin for Ideal Gases

In statistical mechanics, the entropy of an ideal gas mixture originates from the fundamental relation proposed by Ludwig Boltzmann, where the entropy S of a system is given by S = k \ln W, with k being Boltzmann's constant and W the multiplicity, or number of accessible microstates consistent with the macroscopic constraints of the system.^[18] This perspective links the macroscopic thermodynamic entropy to the underlying microscopic configurations of particles, emphasizing that entropy measures the disorder or the number of ways energy can be distributed among particles. For ideal gases, which are non-interacting point particles, the microstates are determined by their positions and momenta in phase space, and mixing increases W due to the greater positional freedom available to each species in the combined volume.^[19] The configurational entropy component specifically accounts for the increased number of ways to arrange distinguishable particles of different species within a shared volume V, without altering their individual kinetic energies or momenta distributions. In the unmixed state, particles of species i are confined to a subvolume V_i, limiting their positional microstates; upon mixing at constant temperature and total volume, each particle type gains access to the full V, multiplicatively expanding the total configurational multiplicity. This expansion arises because the phase space for positions is proportional to V^N for N particles, but for a mixture, the total W becomes the product over species of their individual W_i evaluated at V, reflecting the independence of distinguishable species.^[20] The assumption of distinguishability between different species—such as different atomic or molecular types—ensures that interspecies permutations do not overcount microstates, unlike for identical particles within a species, where the $1/N! correction in the partition function prevents spurious entropy increases.^[19] This microscopic framework is formalized in the Sackur-Tetrode equation for the entropy of a pure monatomic ideal gas in the microcanonical ensemble, given by

S = nR \left[ \ln\left( \frac{V}{n} \left( \frac{4\pi m E}{3n h^2} \right)^{3/2} \right) + \frac{5}{2} \right],

where n is the number of moles, R is the gas constant, V is the volume, m is the particle mass, E is the total internal energy, and h is Planck's constant.^[21] For a mixture of ideal gases, the total entropy is the sum of such terms for each component, each evaluated using the total volume V and their respective n_i, m_i, and partial energies E_i, but the logarithmic dependence on V/n_i introduces an additional configurational contribution equivalent to \Delta S_\text{mix} = -R \sum_i n_i \ln x_i, where x_i = n_i / n is the mole fraction of species i. This mixing term, purely configurational, quantifies the entropy increase from the enhanced positional arrangements and vanishes for identical gases due to their indistinguishability.^[22]

Proof Using Phase Space and Combinatorics

The phase space formulation in classical statistical mechanics provides a rigorous basis for deriving the entropy of mixing for ideal gases. Consider two distinct ideal gases, A and B, initially confined to separate volumes V_A and V_B at the same temperature T and pressure P, with N_A indistinguishable particles of species A and N_B indistinguishable particles of species B. The total phase space volume for the positions of the particles in the unmixed state is proportional to V_A^{N_A} V_B^{N_B}, while the full multiplicity (including momenta and accounting for indistinguishability) is \Omega_\text{unmixed} = \frac{1}{N_A! N_B!} \left( \frac{V_A^{N_A}}{h^{3N_A}} \right) \left( \frac{(2\pi m_A k T)^{3N_A/2}}{(2\pi)^{3N_A/2}} \right) \left( \frac{V_B^{N_B}}{h^{3N_B}} \right) \left( \frac{(2\pi m_B k T)^{3N_B/2}}{(2\pi)^{3N_B/2}} \right), where h is Planck's constant and k is Boltzmann's constant; the momentum integrals remain unchanged upon mixing at constant T. Upon removing the partition, the gases mix irreversibly in the total volume V = V_A + V_B, yielding a mixed multiplicity \Omega_\text{mix} = \frac{1}{N_A! N_B!} \left( \frac{V^{N_A + N_B}}{h^{3(N_A + N_B)}} \right) \left( \frac{(2\pi m_A k T)^{3N_A/2} (2\pi m_B k T)^{3N_B/2}}{(2\pi)^{3(N_A + N_B)/2}} \right). The entropy change is then \Delta S = k \ln (\Omega_\text{mix} / \Omega_\text{unmixed}), which simplifies to the configurational contribution since momentum factors cancel. The combinatorial aspect arises from the factorial terms enforcing indistinguishability within each species, while the different species are treated as distinguishable from each other. The ratio of multiplicities reduces to \frac{\Omega_\text{mix}}{\Omega_\text{unmixed}} = \frac{V^{N_A + N_B}}{V_A^{N_A} V_B^{N_B}}, as the N! and momentum terms are identical in numerator and denominator. For large N_A and N_B, Stirling's approximation \ln N! \approx N \ln N - N is applied, though it cancels out in the ratio here; the entropy of mixing becomes \Delta S = k [(N_A + N_B) \ln V - N_A \ln V_A - N_B \ln V_B]. Assuming constant P and T, the initial partial volumes satisfy V_A / N_A = V_B / N_B = V / N where N = N_A + N_B, so V_A = x_A V and V_B = x_B V with mole fractions x_A = N_A / N and x_B = N_B / N. Substituting yields \Delta S = -k (N_A \ln x_A + N_B \ln x_B) = -N k \sum_i x_i \ln x_i. This derivation holds under key assumptions: the gases are classical ideal gases with no interparticle interactions beyond hard-sphere exclusions (neglected in the point-particle limit), particles within the same species are indistinguishable, different species are distinguishable, and the process occurs at constant temperature with either constant total volume or constant pressure (where volumes are additive). The phase space approach highlights why entropy is extensive yet increases upon mixing due to the enhanced positional freedom, resolving apparent violations of additivity for distinguishable subsystems.

Connections to Broader Concepts

Relationship to Information Theory

The entropy of mixing in thermodynamics exhibits a direct mathematical analogy to the Shannon entropy defined in information theory, both quantifying uncertainty in a system's composition or state. For an ideal mixture, the entropy change upon mixing is \Delta S_\text{mix} = -N R \sum_i x_i \ln x_i, where N is the total number of particles (or moles, with R = N_A k), x_i are the mole fractions of components, and the sum is over species i. This form mirrors the Shannon entropy H = -\sum_i p_i \log p_i, where p_i represent probabilities; treating mole fractions x_i as probabilities p_i and using natural logarithms yields the thermodynamic expression, while base-2 logarithms give bits in information theory. To equate units, dividing \Delta S_\text{mix} by k \ln 2 (with k Boltzmann's constant) converts the thermodynamic entropy to an information measure in bits, highlighting how mixing increases compositional uncertainty akin to informational entropy.^[23] This analogy underscores the conceptual overlap, where the entropy of mixing represents the loss of information about the specific locations or identities of particles before mixing. In an unmixed state, one possesses greater knowledge of which particles occupy which regions; upon mixing, this specificity is lost, elevating uncertainty to a maximum at uniform composition. This informational interpretation aligns with the maximum entropy principle, as articulated by Edwin T. Jaynes, who reformulated thermodynamics as statistical inference under incomplete knowledge. Jaynes argued that the ideal mixing assumption—uniform random distribution of particles—arises from maximizing entropy subject to known constraints (e.g., total composition), reflecting maximal ignorance of microstates beyond macroscopic observables. Thus, \Delta S_\text{mix} embodies the additional uncertainty introduced by ignorance of particle arrangements in the mixture.^[9]^[24] Historically, Ludwig Boltzmann's H-theorem provides a foundational link, defining H = \int f(\mathbf{v}) \ln f(\mathbf{v}) \, d\mathbf{v} (where f(\mathbf{v}) is the velocity distribution function) as a measure whose decrease corresponds to entropy increase toward equilibrium, including diffusive mixing processes. The negative of Boltzmann's H-function parallels the Shannon entropy form, interpreting irreversible mixing as an effective loss of information about initial correlations in particle velocities and positions. This connection influenced Claude Shannon's 1948 formulation of information entropy, which borrowed the logarithmic structure to quantify uncertainty in message sources, extending it beyond physics to communication. In mixing contexts, the H-theorem illustrates how apparent information loss drives the system to higher entropy states, reinforcing the uncertainty parallel without implying true destruction of information.^[23]

Resolution of Gibbs' Paradox

The Gibbs paradox refers to an apparent inconsistency in classical thermodynamics concerning the entropy of mixing ideal gases. The formula for the entropy change upon mixing two gases at constant temperature and pressure predicts a positive ΔS_mix = -R ∑ n_i \ln x_i, where R is the gas constant, n_i are the moles, and x_i the mole fractions, even when the gases are identical, such as two separate volumes of oxygen gas. However, mixing identical gases is a reversible process with no net change in the system's state, implying ΔS_mix should be zero, which contradicts the formula and appears to violate the second law of thermodynamics by suggesting an irreversible entropy increase where none occurs. This paradox was first articulated by J. Willard Gibbs in his 1876 paper on the equilibrium of heterogeneous substances.^[8] The resolution emerges from statistical mechanics by accounting for the indistinguishability of particles within the same species. In the combinatorial approach to multiplicity, for distinguishable particles (different species A and B), the number of microstates in the mixed state exceeds that of the unmixed state by the factor (N_A + N_B)! / (N_A! N_B!), where N_A and N_B are the particle numbers; using Stirling's approximation, this yields ΔS_mix ≈ -Nk (x_A \ln x_A + x_B \ln x_B), with N = N_A + N_B, k Boltzmann's constant, and x_i = N_i / N, reflecting the increased disorder from intermingling distinct types. For identical particles, however, no such combinatorial factor arises because permutations among identical particles do not create new microstates; thus, the multiplicity of the mixed configuration equals that of the unmixed (W_mix = W_unmixed), resulting in ΔS_mix = 0, consistent with reversibility. This indistinguishability principle, incorporated via the 1/N! correction in the phase space volume, ensures the entropy is extensive and resolves the paradox without altering thermodynamic predictions for distinct species.^[25] From a modern perspective, the classical resolution relies on postulating indistinguishability to correct overcounting in phase space, as developed in early statistical mechanics. Quantum mechanics provides a more fundamental explanation, where identical particles are inherently indistinguishable due to the symmetrization postulate for bosons or antisymmetrization for fermions, enforced by the form of the wavefunction; this eliminates the paradox naturally, as the Hilbert space for identical particles lacks the extra states assumed in classical distinguishable treatments. Experimental verification aligns with quantum statistics, confirming no entropy increase for identical quantum gases. In the limiting case of closely similar but non-identical species, such as isotopes of the same element (e.g., ^{16}O and ^{18}O), the entropy of mixing is near zero and approaches the identical case continuously as the physical properties (e.g., masses, leading to de Broglie wavelengths λ = h / \sqrt{2π m k T} comparable to interparticle distances) become indistinguishable; this gradual transition, rather than a sharp discontinuity, further validates the resolution by ensuring thermodynamic continuity across the boundary of identity.^[8]

Variations in Mixing Conditions

Mixing at Constant Total Volume

In the scenario of mixing at constant total volume and temperature, the process involves combining components such that the overall volume V remains fixed, while the temperature T is maintained constant. Under these conditions, the partial pressures of the components adjust according to p_i = x_i P_{\text{total}}, where x_i is the mole fraction of component i and P_{\text{total}} is the total pressure, which may differ from the initial pressures if the components started at unequal densities.^[26]^[27] For ideal mixtures, the change in internal energy upon mixing is zero, \Delta U_{\text{mix}} = 0, because the internal energy of ideal components depends solely on temperature, with no intermolecular interactions or volume-dependent contributions. The total entropy change upon mixing \Delta S_{\text{mix}} under these conditions is \Delta S_{\text{mix}} = R \sum_i n_i \ln (V / V_i), where R is the gas constant, n_i the moles of component i, V the total volume, and V_i the initial volume of component i. This expression simplifies to the standard configurational form -R \sum_i n_i \ln x_i when the initial partial pressures are equal (i.e., V_i \propto n_i).^[26]^[28] To derive \Delta S_{\text{mix}}, consider a reversible path for the irreversible mixing process, using the thermodynamic relation dS = \delta Q_{\text{rev}} / T. For each ideal component, treat the mixing as a sequence of isothermal expansions from initial volume V_i to the total volume V at constant T, yielding \Delta S_i = n_i [R](/page/R) \ln (V / V_i) per component. Summing over components gives \Delta S_{\text{mix}} = [R](/page/R) \sum_i n_i \ln (V / V_i), which simplifies to the standard form when initial conditions align with equal partial pressures (i.e., V_i \propto n_i). Since the process is isothermal and \Delta U_{\text{mix}} = 0, no heat is exchanged in the actual irreversible mixing (Q = 0), but the entropy increase reflects the irreversible diffusion.^[26]^[28] Mechanically, this mixing can be realized in a rigid container of fixed total volume V, initially divided by impermeable partitions separating the pure components in sub-volumes V_i (with \sum V_i = V). Removing the partitions allows diffusion without any change in external volume or pressure-volume work, so the entropy increase stems purely from the enhanced configurational possibilities of the molecules, rather than any expansion work.^[27]^[26] In contrast to mixing at constant pressure, where volume may adjust and PdV work contributes to the energetics, constant total volume eliminates such work terms, isolating the entropy change as an internal, irreversible configurational effect with no net heat or work exchange for the system.^[28]

Mixing with Volume Changes and Partial Volumes

In the case of mixing ideal gases assuming constant partial volumes, each component expands isothermally from its initial volume V_i to the total volume V_\text{total} = \sum V_i, leading to an entropy change of \Delta S = \sum n_i R \ln(V_\text{total}/V_i), which simplifies to \Delta S = -R \sum n_i \ln x_i where x_i = n_i / n is the mole fraction.^[29] This expression arises from the configurational contribution in the Sackur-Tetrode equation for ideal gases, reflecting the increased number of accessible microstates upon mixing.^[30] When partial volumes change during mixing while maintaining constant total volume, such as in liquid solutions or controlled gas mixtures, pressure adjustments via pistons or other mechanical means are required to achieve equilibrium. For ideal cases, the entropy of mixing retains the form \Delta S = -[R](/page/R) \sum n_i \ln x_i, but the process incorporates work terms from compression or expansion steps, ensuring reversibility. In non-ideal solutions, deviations occur if partial molar volumes vary with composition, affecting the overall volume and thus the entropy through changes in molecular interactions, though the ideal configurational term dominates for dilute systems. Mechanically controlled mixing at constant temperature often employs semi-permeable membranes or pistons to allow selective expansion of each gas component, enabling a reversible path. For instance, a piston permeable only to one gas expands it isothermally against an external pressure, with the entropy change for that component given by \Delta S_i = n_i [R](/page/R) \ln(V_\text{total}/V_i), and the total \Delta S_\text{mix} summing over components.^[29] This setup highlights how intermediate work exchanges, such as w_i = -n_i [R](/page/R) T \ln(V_\text{total}/V_i), maintain isothermal conditions without net heat effects beyond those compensating the work.^[30] The total entropy of mixing \Delta S_\text{mix} is a state function, remaining independent of the mixing path—whether irreversible diffusion or reversible mechanical control—yielding the same value for given initial and final states at constant temperature.^[29] However, intermediate steps differ: irreversible processes generate entropy internally without work recovery, while reversible ones exchange heat reversibly with surroundings, preserving the second law.^[30]

Applications

To Gaseous Mixtures

The entropy of mixing plays a fundamental role in the thermodynamics of gaseous mixtures, particularly in naturally occurring systems like Earth's atmosphere. Dry air can be approximated as an ideal gas mixture consisting of 78.08% nitrogen, 20.95% oxygen, and 0.93% argon by mole fraction. For one mole of this mixture, the entropy change upon mixing from pure components at constant temperature and pressure is given by ΔS_mix = -R ∑ x_i \ln x_i, where R is the gas constant and x_i are the mole fractions, yielding approximately 4.70 J/mol·K. This positive entropy increase reflects the greater disorder achieved when the gases intermix, contributing to the overall entropy budget in atmospheric processes such as convection and transport, where irreversible mixing generates entropy that influences weather patterns and climate dissipation.^[31]^[32]^[33] In gaseous systems, the entropy of mixing serves as the thermodynamic driving force for spontaneous processes like diffusion and effusion, which are governed by the tendency to maximize entropy. Diffusion occurs when gases intermingle due to concentration gradients, increasing entropy as molecules disperse more uniformly, while effusion involves gases escaping through a small aperture, similarly driven by the entropy gain from expansion into a larger volume. Graham's law, which states that the rate of effusion or diffusion is inversely proportional to the square root of the molecular mass, describes the kinetics of these processes but aligns with the underlying entropy maximization, as lighter gases effuse faster, leading to partial separations that are reversible only in the limit of infinitesimal gradients; in practice, these processes are irreversible, producing entropy.^[34]^[35] For non-ideal gaseous mixtures at low densities, such as those encountered in most atmospheric or dilute industrial conditions, corrections to the ideal entropy of mixing from equations of state like van der Waals are minimal. The van der Waals model accounts for excluded volume (parameter b) and intermolecular attractions (parameter a), reducing the effective accessible volume and thus slightly lowering the entropy compared to the ideal case; however, at low densities where b/V << 1 and a effects on entropy are secondary, the ideal mixing entropy approximates the real behavior closely, with deviations typically less than 1% for air-like mixtures near standard conditions.^[36] In industrial applications, the entropy of mixing informs the efficiency of processes like cryogenic air separation units (ASUs), where compressed and cooled air is distilled to produce pure oxygen, nitrogen, and argon. The minimum theoretical work required to reverse the mixing and separate one mole of air into its components at 298 K is approximately T ΔS_mix ≈ 1.40 kJ/mol, equivalent to 1.40 MJ/kmol, serving as a baseline for exergetic analysis; state-of-the-art ASUs achieve second-law efficiencies of about 35-40%, with entropy production in distillation columns representing a key inefficiency that optimization strategies aim to minimize through reduced irreversibilities in heat and mass transfer.^[32]^[37]^[38]

To Liquid Solutions

The entropy of mixing in liquid solutions extends the ideal gas formulation to condensed phases, where molecular volumes and interactions play a more prominent role, particularly in dilute regimes where the solute concentration is low and the solvent dominates the mixture volume. For ideal liquid solutions, the molar entropy of mixing follows the same statistical form as for gases: ΔS_mix = -R Σ x_i ln x_i, where x_i are mole fractions and R is the gas constant, reflecting the increased disorder from intermingling components without volume change upon mixing.^[39] In dilute solutions, this simplifies due to the solvent's near-pure state; the partial molar entropy change for the solvent is approximately ΔS_solvent ≈ -R ln x_solvent, where x_solvent ≈ 1 - x_solute, leading to ΔS_solvent ≈ R x_solute for small x_solute, indicating a modest increase in solvent disorder.^[39] Henry's law governs solute behavior in such dilute solutions, stating that the partial vapor pressure of the solute p_solute = K_H x_solute, where K_H is the Henry's law constant specific to the solute-solvent pair. This law implies an ideal-dilute reference state for the solute's chemical potential, μ_solute = μ°_solute + RT ln x_solute, from which the partial molar entropy of the solute derives as S_solute = S°_solute - R ln x_solute, capturing the configurational entropy gain as the solute disperses in the solvent lattice.^[39] The total entropy of mixing in these systems thus approximates ΔS_mix ≈ -n R (x_solvent ln x_solvent + x_solute ln x_solute), but for dilute conditions (x_solute ≪ 1), it reduces to ΔS_mix ≈ n R x_solute, emphasizing the solute's dominant contribution to disorder while the solvent's remains small. This framework underpins colligative properties, where phase equilibria (ΔG = 0) link mixing entropy to observable effects like boiling point elevation ΔT_b ≈ (R T_b² / ΔH_vap) x_solute and freezing point depression ΔT_f ≈ (R T_f² / ΔH_fus) x_solute, as the entropy-driven vapor pressure lowering shifts equilibrium temperatures.^[39] Representative examples illustrate these concepts in binary liquid systems. For a dilute saltwater solution, such as 0.01 mol fraction NaCl in water at 298 K, the ideal ΔS_mix ≈ -R [0.99 ln 0.99 + 0.01 ln 0.01] ≈ 0.47 J/mol·K, reflecting modest disorder increase that contributes to freezing point depression of about 1.9 K (accounting for dissociation with van't Hoff factor i ≈ 1.85), consistent with observed colligative behavior. In an alcohol-water mixture, like 0.1 mol fraction ethanol in water, the ideal calculation yields ΔS_mix ≈ 2.7 J/mol·K, but experimental vapor-liquid equilibria show deviations due to hydrogen bonding.^[39]^[40] For non-ideal liquid solutions, the ideal entropy serves as a baseline, with corrections via excess entropy S^E that accounts for local structural deviations from random mixing. Models like UNIQUAC (UNIversal QUAsi-Chemical) decompose the excess Gibbs energy into combinatorial (entropic) and residual (enthalpic) terms, where the combinatorial contribution to S^E arises from molecular size and shape differences: S^E_comb / R = Σ x_i ln (φ_i / x_i) + (1 - φ_i) ln [(1 - φ_i)/(1 - x_i)], with φ_i as volume fractions, enabling prediction of activity coefficients in mixtures like ethanol-water where S^E can be negative due to ordering effects.^[41] This adjustment ensures accurate thermodynamic modeling while preserving the ideal ΔS_mix as the reference for dilute limits.^[42]

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