Phase rule

Last updated May 11, 2024

In thermodynamics, the phase rule is a general principle governing "pVT" systems, whose thermodynamic states are completely described by the variables pressure ( $p$ ), volume ( $V$ ) and temperature ( $T$ ), in thermodynamic equilibrium. If $F$ is the number of degrees of freedom, $C$ is the number of components and $P$ is the number of phases, then

The number of degrees of freedom is the number of independent intensive variables, i.e. the largest number of thermodynamic parameters such as temperature or pressure that can be varied simultaneously and arbitrarily without determining one another. An example of one-component system is a system involving one pure chemical, while two-component systems, such as mixtures of water and ethanol, have two chemically independent components, and so on. Typical phases are solids, liquids and gases.

Foundations

A phase is a form of matter that is homogeneous in chemical composition and physical state. Typical phases are solid, liquid and gas. Two immiscible liquids (or liquid mixtures with different compositions) separated by a distinct boundary are counted as two different phases, as are two immiscible solids.
The number of components (C) is the number of chemically independent constituents of the system, i.e. the minimum number of independent species necessary to define the composition of all phases of the system. [2]
The number of degrees of freedom (F) in this context is the number of intensive variables which are independent of each other.

The basis for the rule^[2]^{: 122–126} is that equilibrium between phases places a constraint on the intensive variables. More rigorously, since the phases are in thermodynamic equilibrium with each other, the chemical potentials of the phases must be equal. The number of equality relationships determines the number of degrees of freedom. For example, if the chemical potentials of a liquid and of its vapour depend on temperature (T) and pressure (p), the equality of chemical potentials will mean that each of those variables will be dependent on the other. Mathematically, the equation $μ liq (T, p) = μ vap (T, p)$ , where $μ$ , the chemical potential, defines temperature as a function of pressure or vice versa. (Caution: do not confuse $p$ as pressure with $P$ , number of phases.)

To be more specific, the composition of each phase is determined by C − 1 intensive variables (such as mole fractions) in each phase. The total number of variables is (C − 1)P + 2, where the extra two are temperature T and pressure p. The number of constraints is C(P − 1), since the chemical potential of each component must be equal in all phases. Subtract the number of constraints from the number of variables to obtain the number of degrees of freedom as F = (C − 1)P + 2 − C(P − 1) = C − P + 2.

The rule is valid provided the equilibrium between phases is not influenced by gravitational, electrical or magnetic forces, or by surface area, and only by temperature, pressure, and concentration.

Consequences and examples

Pure substances (one component)

For pure substances C = 1 so that F = 3 − P. In a single phase (P = 1) condition of a pure component system, two variables (F = 2), such as temperature and pressure, can be chosen independently to be any pair of values consistent with the phase. However, if the temperature and pressure combination ranges to a point where the pure component undergoes a separation into two phases (P = 2), F decreases from 2 to 1. When the system enters the two-phase region, it becomes no longer possible to independently control temperature and pressure.

In the phase diagram to the right, the boundary curve between the liquid and gas regions maps the constraint between temperature and pressure when the single-component system has separated into liquid and gas phases at equilibrium. The only way to increase the pressure on the two phase line is by increasing the temperature. If the temperature is decreased by cooling, some of the gas condenses, decreasing the pressure. Throughout both processes, the temperature and pressure stay in the relationship shown by this boundary curve unless one phase is entirely consumed by evaporation or condensation, or unless the critical point is reached. As long as there are two phases, there is only one degree of freedom, which corresponds to the position along the phase boundary curve.

The critical point is the black dot at the end of the liquid–gas boundary. As this point is approached, the liquid and gas phases become progressively more similar until, at the critical point, there is no longer a separation into two phases. Above the critical point and away from the phase boundary curve, F = 2 and the temperature and pressure can be controlled independently. Hence there is only one phase, and it has the physical properties of a dense gas, but is also referred to as a supercritical fluid.

Of the other two-boundary curves, one is the solid–liquid boundary or melting point curve which indicates the conditions for equilibrium between these two phases, and the other at lower temperature and pressure is the solid–gas boundary.

Even for a pure substance, it is possible that three phases, such as solid, liquid and vapour, can exist together in equilibrium (P = 3). If there is only one component, there are no degrees of freedom (F = 0) when there are three phases. Therefore, in a single-component system, this three-phase mixture can only exist at a single temperature and pressure, which is known as a triple point. Here there are two equations μ_sol(T, p) = μ_liq(T, p) = μ_vap(T, p), which are sufficient to determine the two variables T and p. In the diagram for CO₂ the triple point is the point at which the solid, liquid and gas phases come together, at 5.2 bar and 217 K. It is also possible for other sets of phases to form a triple point, for example in the water system there is a triple point where ice I, ice III and liquid can coexist.

If four phases of a pure substance were in equilibrium (P = 4), the phase rule would give F = −1, which is meaningless, since there cannot be −1 independent variables. This explains the fact that four phases of a pure substance (such as ice I, ice III, liquid water and water vapour) are not found in equilibrium at any temperature and pressure. In terms of chemical potentials there are now three equations, which cannot in general be satisfied by any values of the two variables T and p, although in principle they might be solved in a special case where one equation is mathematically dependent on the other two. In practice, however, the coexistence of more phases than allowed by the phase rule normally means that the phases are not all in true equilibrium.

Two-component systems

For binary mixtures of two chemically independent components, C = 2 so that F = 4 − P. In addition to temperature and pressure, the other degree of freedom is the composition of each phase, often expressed as mole fraction or mass fraction of one component.

As an example, consider the system of two completely miscible liquids such as toluene and benzene, in equilibrium with their vapours. This system may be described by a boiling-point diagram which shows the composition (mole fraction) of the two phases in equilibrium as functions of temperature (at a fixed pressure).

Four thermodynamic variables which may describe the system include temperature (T), pressure (p), mole fraction of component 1 (toluene) in the liquid phase (x_1L), and mole fraction of component 1 in the vapour phase (x_1V). However, since two phases are present (P = 2) in equilibrium, only two of these variables can be independent (F = 2). This is because the four variables are constrained by two relations: the equality of the chemical potentials of liquid toluene and toluene vapour, and the corresponding equality for benzene.

For given T and p, there will be two phases at equilibrium when the overall composition of the system (system point) lies in between the two curves. A horizontal line (isotherm or tie line) can be drawn through any such system point, and intersects the curve for each phase at its equilibrium composition. The quantity of each phase is given by the lever rule (expressed in the variable corresponding to the x-axis, here mole fraction).

For the analysis of fractional distillation, the two independent variables are instead considered to be liquid-phase composition (x_1L) and pressure. In that case the phase rule implies that the equilibrium temperature (boiling point) and vapour-phase composition are determined.

Liquid–vapour phase diagrams for other systems may have azeotropes (maxima or minima) in the composition curves, but the application of the phase rule is unchanged. The only difference is that the compositions of the two phases are equal exactly at the azeotropic composition.

Aqueous solution of 4 kinds of salts

Consider an aqueous solution containing sodium chloride (NaCl), potassium chloride (KCl), sodium bromide (NaBr), and potassium bromide (KBr), in equilibrium with their respective solid phases. Each salt, in solid form, is a different phase, because each possesses a distinct crystal structure and composition. The aqueous solution itself is another phase, because it forms a homogeneous liquid phase separate from the solid salts, with its own distinct composition and physical properties. Thus we have P = 5 phases.

There are 6 elements present (H, O, Na, K, Cl, Br), but we have 2 constraints:

The stoichiometry of water: n(H) = 2n(O).
Charge balance in the solution: n(Na) + n(K) = n(Cl) + n(Br).

giving C = 6 - 2 = 4 components. The Gibbs phase rule states that F = 1. So, for example, if we plot the P-T phase diagram of the system, there is only one line at which all phases coexist. Any deviation from the line would either cause one of the salts to completely dissolve or one of the ions to completely precipitate from the solution.

Phase rule at constant pressure

For applications in materials science dealing with phase changes between different solid structures, pressure is often imagined to be constant (for example at one atmosphere), and is ignored as a degree of freedom, so the formula becomes:^[4]

F=C-P+1

This is sometimes incorrectly called the "condensed phase rule", but it is not applicable to condensed systems which are subject to high pressures (for example, in geology), since the effects of these pressures are important.^[5]

Phase rule in colloidal mixtures

In colloidal mixtures quintuple^[6]^[7] and sixtuple points^[8]^[9] have been described in violation of Gibbs phase rule but it is argued that in these systems the rule can be generalized to $F=M+C-P+1$ where $M$ accounts for additional parameters of interaction among the components like the diameter of one type of particle in relation to the diameter of the other particles in the solution.

Related Research Articles

<span class="mw-page-title-main">Boiling point</span> Temperature at which a substance changes from liquid into vapor

The boiling point of a substance is the temperature at which the vapor pressure of a liquid equals the pressure surrounding the liquid and the liquid changes into a vapor.

Distillation, also classical distillation, is the process of separating the component substances of a liquid mixture of two or more chemically discrete substances; the separation process is realized by way of the selective boiling of the mixture and the condensation of the vapors in a still.

In the physical sciences, a phase is a region of material that is chemically uniform, physically distinct, and (often) mechanically separable. In a system consisting of ice and water in a glass jar, the ice cubes are one phase, the water is a second phase, and the humid air is a third phase over the ice and water. The glass of the jar is another separate phase.

<span class="mw-page-title-main">Solution (chemistry)</span> Homogeneous mixture of a solute and a solvent

In chemistry, a solution is a special type of homogeneous mixture composed of two or more substances. In such a mixture, a solute is a substance dissolved in another substance, known as a solvent. If the attractive forces between the solvent and solute particles are greater than the attractive forces holding the solute particles together, the solvent particles pull the solute particles apart and surround them. These surrounded solute particles then move away from the solid solute and out into the solution. The mixing process of a solution happens at a scale where the effects of chemical polarity are involved, resulting in interactions that are specific to solvation. The solution usually has the state of the solvent when the solvent is the larger fraction of the mixture, as is commonly the case. One important parameter of a solution is the concentration, which is a measure of the amount of solute in a given amount of solution or solvent. The term "aqueous solution" is used when one of the solvents is water.

<span class="mw-page-title-main">Triple point</span> Thermodynamic point where three matter phases exist

In thermodynamics, the triple point of a substance is the temperature and pressure at which the three phases of that substance coexist in thermodynamic equilibrium. It is that temperature and pressure at which the sublimation, fusion, and vaporisation curves meet. For example, the triple point of mercury occurs at a temperature of −38.8 °C (−37.8 °F) and a pressure of 0.165 mPa.

In physics, a vapor or vapour is a substance in the gas phase at a temperature lower than its critical temperature, which means that the vapor can be condensed to a liquid by increasing the pressure on it without reducing the temperature of the vapor. A vapor is different from an aerosol. An aerosol is a suspension of tiny particles of liquid, solid, or both within a gas.

Raoult's law ( law) is a relation of physical chemistry, with implications in thermodynamics. Proposed by French chemist François-Marie Raoult in 1887, it states that the partial pressure of each component of an ideal mixture of liquids is equal to the vapor pressure of the pure component multiplied by its mole fraction in the mixture. In consequence, the relative lowering of vapor pressure of a dilute solution of nonvolatile solute is equal to the mole fraction of solute in the solution.

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<span class="mw-page-title-main">Azeotrope</span> Mixture of two or more liquids whose proportions do not change when the mixture is distilled

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In chemistry, solubility is the ability of a substance, the solute, to form a solution with another substance, the solvent. Insolubility is the opposite property, the inability of the solute to form such a solution.

In chemical thermodynamics, activity is a measure of the "effective concentration" of a species in a mixture, in the sense that the species' chemical potential depends on the activity of a real solution in the same way that it would depend on concentration for an ideal solution. The term "activity" in this sense was coined by the American chemist Gilbert N. Lewis in 1907.

In physical chemistry, Henry's law is a gas law that states that the amount of dissolved gas in a liquid is directly proportional to its partial pressure above the liquid. The proportionality factor is called Henry's law constant. It was formulated by the English chemist William Henry, who studied the topic in the early 19th century. In simple words, we can say that the partial pressure of a gas in vapour phase is directly proportional to the mole fraction of a gas in solution.

In chemistry, colligative properties are those properties of solutions that depend on the ratio of the number of solute particles to the number of solvent particles in a solution, and not on the nature of the chemical species present. The number ratio can be related to the various units for concentration of a solution such as molarity, molality, normality (chemistry), etc. The assumption that solution properties are independent of nature of solute particles is exact only for ideal solutions, which are solutions that exhibit thermodynamic properties analogous to those of an ideal gas, and is approximate for dilute real solutions. In other words, colligative properties are a set of solution properties that can be reasonably approximated by the assumption that the solution is ideal.

In chemical thermodynamics, the fugacity of a real gas is an effective partial pressure which replaces the mechanical partial pressure in an accurate computation of chemical equilibrium. It is equal to the pressure of an ideal gas which has the same temperature and molar Gibbs free energy as the real gas.

In thermodynamics, the entropy of mixing is the increase in the total entropy when several initially separate systems of different composition, each in a thermodynamic state of internal equilibrium, are mixed without chemical reaction by the thermodynamic operation of removal of impermeable partition(s) between them, followed by a time for establishment of a new thermodynamic state of internal equilibrium in the new unpartitioned closed system.

In thermodynamics and chemical engineering, the vapor–liquid equilibrium (VLE) describes the distribution of a chemical species between the vapor phase and a liquid phase.

This glossary of chemistry terms is a list of terms and definitions relevant to chemistry, including chemical laws, diagrams and formulae, laboratory tools, glassware, and equipment. Chemistry is a physical science concerned with the composition, structure, and properties of matter, as well as the changes it undergoes during chemical reactions; it features an extensive vocabulary and a significant amount of jargon.

The Duhem–Margules equation, named for Pierre Duhem and Max Margules, is a thermodynamic statement of the relationship between the two components of a single liquid where the vapour mixture is regarded as an ideal gas:

References

↑ Ness, Hendrick C. Van; Abbott, Michael; Swihart, Mark; Smith, J. M. (March 20, 2017). Introduction to Chemical Engineering Thermodynamics. Dubuque, Iowa: McGraw Hill Education. p. 422. ISBN 9781259696527. OCLC 1001316575.
1 2 3 Atkins, Peter William; Paula, Julio De; Keeler, James (2018). Atkins' Physical Chemistry (11th ed.). Oxford University Press. ISBN 9780198769866. OCLC 1013164457.
↑ Gibbs, Josiah W. (1906). Scientific Papers of J. Willard Gibbs. Longmans, Green and Co. OCLC 1136910263.
↑ Ehlers, Ernest G. "Phase - State of Matter (section Binary Systems)". Encyclopedia Britannica. Retrieved 17 November 2022. at atmospheric pressure; because the pressure variable is fixed, the phase rule is expressed as P + F = C + 1.
↑ "Phase rule". Science Education and Research Centre (SERC) at Carleton College . Retrieved 2022-10-09.
↑ Peters, V. F. D.; Vis, M.; García, Á. González; Wensink, H. H.; Tuinier, R. (2020-09-18). "Defying the Gibbs Phase Rule: Evidence for an Entropy-Driven Quintuple Point in Colloid-Polymer Mixtures". Physical Review Letters. 125 (12): 127803. arXiv: 2009.08879 . doi:10.1103/PhysRevLett.125.127803.
↑ "'Quintuple point' material defies 150-year-old thermodynamics rule". Physics World. 2020-10-22. Retrieved 2023-02-21.
↑ Opdam, J.; Peters, V. F. D.; Wensink, H. H.; Tuinier, R. (2023-01-12). "Multiphase Coexistence in Binary Hard Colloidal Mixtures: Predictions from a Simple Algebraic Theory". The Journal of Physical Chemistry Letters. 14 (1): 199–206. doi:10.1021/acs.jpclett.2c03138. ISSN 1948-7185. PMC 9841575 . PMID 36580685.
↑ "Colloidal mixture exists in up to six phases at once". Physics World. 2023-02-14. Retrieved 2023-02-21.

v t e Chemical equilibria
Concepts	Chemical stability Chelation Dynamic equilibrium Equilibrium chemistry Equilibrium stage Free energy Gibbs Helmholtz Le Chatelier's principle Phase separation Reversible reaction Thermodynamic equilibrium
Models	Equilibrium constant determination Phase diagram Predominance diagram Phase rule Reaction quotient Thermodynamic activity
Applications	Buffer solution Equilibrium unfolding Liquid–liquid extraction
Specific equilibria	Acid dissociation Hammett acidity function Binding constant Binding selectivity Coordination complexes Macrocyclic effect Dissociation constant Hydrolysis Self-ionization of water Partition Distribution coefficient Solubility Common-ion effect Vapor–liquid Henry's law