 Chemical equilibrium

In a chemical reaction, chemical equilibrium is the state in which the concentrations of the reactants and products have not yet changed with time. It occurs only in reversible reactions, and not in irreversible reactions. Usually, this state results when the forward reaction proceeds at the same rate as the reverse reaction. The reaction rates of the forward and reverse reactions are generally not zero but, being equal, there are no net changes in the concentrations of the reactant and product. This process is called dynamic equilibrium.^{[1]}^{[2]}
Contents
Introduction
The concept of chemical equilibrium was developed after Berthollet (1803) found that some chemical reactions are reversible. For any reaction mixture to exist at equilibrium, the rates of the forward and backward (reverse) reactions are equal. In the following chemical equation with arrows pointing both ways to indicate equilibrium, A and B are reactant chemical species, S and T are product species, and α, β, σ, and τ are the stoichiometric coefficients of the respective reactants and products:
The equilibrium position of a reaction is said to lie "far to the right" if, at equilibrium, nearly all the reactants are consumed. Conversely the equilibrium position is said to be "far to the left" if hardly any product is formed from the reactants.
Guldberg and Waage (1865), building on Berthollet’s ideas, proposed the law of mass action:
where A, B, S and T are active masses and k_{+} and k_{−} are rate constants. Since at equilibrium forward and backward rates are equal:
and the ratio of the rate constants is also a constant, now known as an equilibrium constant.
By convention the products form the numerator. However, the law of mass action is valid only for concerted onestep reactions that proceed through a single transition state and is not valid in general because rate equations do not, in general, follow the stoichiometry of the reaction as Guldberg and Waage had proposed (see, for example, nucleophilic aliphatic substitution by S_{N}1 or reaction of hydrogen and bromine to form hydrogen bromide). Equality of forward and backward reaction rates, however, is a necessary condition for chemical equilibrium, though it is not sufficient to explain why equilibrium occurs.
Despite the failure of this derivation, the equilibrium constant for a reaction is indeed a constant, independent of the activities of the various species involved, though it does depend on temperature as observed by the van 't Hoff equation. Adding a catalyst will affect both the forward reaction and the reverse reaction in the same way and will not have an effect on the equilibrium constant. The catalyst will speed up both reactions thereby increasing the speed at which equilibrium is reached.^{[1]}^{[3]}
Although the macroscopic equilibrium concentrations are constant in time reactions do occur at the molecular level. For example, in the case of acetic acid dissolved in water and forming acetate and hydronium ions,
 CH_{3}CO_{2}H + H_{2}O ⇌ CH_{3}CO_{2}^{−} + H_{3}O^{+}
a proton may hop from one molecule of acetic acid on to a water molecule and then on to an acetate anion to form another molecule of acetic acid and leaving the number of acetic acid molecules unchanged. This is an example of dynamic equilibrium. Equilibria, like the rest of thermodynamics, are statistical phenomena, averages of microscopic behavior.
Le Chatelier's principle (1884) gives an idea of the behavior of an equilibrium system when changes to its reaction conditions occur. If a dynamic equilibrium is disturbed by changing the conditions, the position of equilibrium moves to partially reverse the change. For example, adding more S from the outside will cause an excess of products, and the system will try to counteract this by increasing the reverse reaction and pushing the equilibrium point backward (though the equilibrium constant will stay the same).
If mineral acid is added to the acetic acid mixture, increasing the concentration of hydronium ion, the amount of dissociation must decrease as the reaction is driven to the left in accordance with this principle. This can also be deduced from the equilibrium constant expression for the reaction:
If {H_{3}O^{+}} increases {CH_{3}CO_{2}H} must increase and {CH_{3}CO_{2}^{−}} must decrease. The H_{2}O is left out as it is a pure liquid and its concentration is undefined.
A quantitative version is given by the reaction quotient.
J. W. Gibbs suggested in 1873 that equilibrium is attained when the Gibbs energy of the system is at its minimum value (assuming the reaction is carried out under constant pressure). What this means is that the derivative of the Gibbs energy with respect to reaction coordinate (a measure of the extent of reaction that has occurred, ranging from zero for all reactants to a maximum for all products) vanishes, signalling a stationary point. This derivative is usually called, for certain technical reasons, the Gibbs energy change.^{[4]} This criterion is both necessary and sufficient. If a mixture is not at equilibrium, the liberation of the excess Gibbs energy (or Helmholtz energy at constant volume reactions) is the “driving force” for the composition of the mixture to change until equilibrium is reached. The equilibrium constant can be related to the standard Gibbs energy change for the reaction by the equation
where R is the universal gas constant and T the temperature.
When the reactants are dissolved in a medium of high ionic strength the quotient of activity coefficients may be taken to be constant. In that case the concentration quotient, K_{c},
where [A] is the concentration of A, etc., is independent of the analytical concentration of the reactants. For this reason, equilibrium constants for solutions are usually determined in media of high ionic strength. K_{c} varies with ionic strength, temperature and pressure (or volume). Likewise K_{p} for gases depends on partial pressure. These constants are easier to measure and encountered in highschool chemistry courses.
Thermodynamics
The relation between the Gibbs energy and the equilibrium constant can be found by considering chemical potentials.^{[5]} At constant temperature and pressure the function G Gibbs free energy for the reaction, depends only with the extent of reaction: ξ and can only decrease according to the second law of thermodynamics. It means that the derivative of G with ξ must be negative if the reaction happens; at the equilibrium the derivative being equal to zero.
 : equilibrium
At constant volume, one must consider the Helmholtz free energy for the reaction: A.
In this article only the constant pressure case is considered. The constant volume case is important in geochemistry and atmospheric chemistry where pressure variations are significant. Note that, if reactants and products were in standard state (completely pure), then there would be no reversibility and no equilibrium. The mixing of the products and reactants contributes a large entropy (known as entropy of mixing) to states containing equal mixture of products and reactants. The combination of the standard Gibbs energy change and the Gibbs energy of mixing determines the equilibrium state.^{[6]}^{[7]}
In general an equilibrium system is defined by writing an equilibrium equation for the reaction
In order to meet the thermodynamic condition for equilibrium, the Gibbs energy must be stationary, meaning that the derivative of G with respect to the extent of reaction: ξ, must be zero. It can be shown that in this case, the sum of chemical potentials of the products is equal to the sum of those corresponding to the reactants. Therefore, the sum of the Gibbs energies of the reactants must be the equal to the sum of the Gibbs energies of the products.
where μ is in this case a partial molar Gibbs energy, a chemical potential. The chemical potential of a reagent A is a function of the activity, {A} of that reagent.
 , ( is the standard chemical potential ).
Substituting expressions like this into the Gibbs energy equation:
 in the case of a closed system.
Now
 ( corresponds to the Stoichiometric coefficient and is the differential of the extent of reaction ).
At constant pressure and temperature we obtain:
 which corresponds to the Gibbs free energy change for the reaction .
This results in:
 .
By substituting the chemical potentials:
 ,
the relationship becomes:

 : which is the standard Gibbs energy change for the reaction. It is a constant at a given temperature, which can be calculated, using thermodynamical tables.

 ( is the reaction quotient when the system is not at equilibrium ).
Therefore
At equilibrium

 ; the reaction quotient becomes equal to the equilibrium constant.
leading to:
and
Obtaining the value of the standard Gibbs energy change, allows the calculation of the equilibrium constant
Addition of reactants or products
For a reactional system at equilibrium: ; .
 If are modified activities of constituents, the value of the reaction quotient changes and becomes different from the equilibrium constant:
and
then
 If activity of a reagent increases
, the reaction quotient decreases.
 then
and : The reaction will shift to the right (i.e. in the forward direction, and thus more products will form).
 If activity of a product increases
 then
and : The reaction will shift to the left (i.e. in the reverse direction, and thus less products will form).
Note that activities and equilibrium constants are dimensionless numbers.
Treatment of activity
The expression for the equilibrium constant can be rewritten as the product of a concentration quotient, K_{c} and an activity coefficient quotient, Γ.
[A] is the concentration of reagent A, etc. It is possible in principle to obtain values of the activity coefficients, γ. For solutions, equations such as the DebyeHückel equation or extensions such as Davies equation^{[8]} Specific ion interaction theory or Pitzer equations^{[9]} may be used.^{Software (below)}. However this is not always possible. It is common practice to assume that Γ is a constant, and to use the concentration quotient in place of the thermodynamic equilibrium constant. It is also general practice to use the term equilibrium constant instead of the more accurate concentration quotient. This practice will be followed here.
For reactions in the gas phase partial pressure is used in place of concentration and fugacity coefficient in place of activity coefficient. In the real world, for example, when making ammonia in industry, fugacity coefficients must be taken into account. Fugacity, f, is the product of partial pressure and fugacity coefficient. The chemical potential of a species in the gas phase is given by
so the general expression defining an equilibrium constant is valid for both solution and gas phases.
Concentration quotients
In aqueous solution, equilibrium constants are usually determined in the presence of an "inert" electrolyte such as sodium nitrate NaNO_{3} or Potassium perchlorate KClO_{4}. The ionic strength, I, of a solution containing a dissolved salt, X^{+}Y^{}, is given by
where c stands for concentration, z stands for ionic charge and the sum is taken over all the species in equilibrium. When the concentration of dissolved salt is much higher than the analytical concentrations of the reagents, the ionic strength is effectively constant. Since activity coefficients depend on ionic strength the activity coefficients of the species are effectively independent of concentration. Thus, the assumption that Γ is constant is justified. The concentration quotient is a simple multiple of the equilibrium constant.^{[10]}
However, K_{c} will vary with ionic strength. If it is measured at a series of different ionic strengths the value can be extrapolated to zero ionic strength.^{[9]} The concentration quotient obtained in this manner is known, paradoxically, as a thermodynamic equilibrium constant.
To use a published value of an equilibrium constant in conditions of ionic strength different from the conditions used in its determination, the value should be adjusted^{Software (below)}.
Metastable mixtures
A mixture may be appear to have no tendency to change, though it is not at equilibrium. For example, a mixture of SO_{2} and O_{2} is metastable as there is a kinetic barrier to formation of the product, SO_{3}.
 2SO_{2} + O_{2} 2SO_{3}
The barrier can be overcome when a catalyst is also present in the mixture as in the contact process, but the catalyst does not affect the equilibrium concentrations.
Likewise, the formation of bicarbonate from carbon dioxide and water is very slow under normal conditions
 CO_{2} + 2H_{2}O HCO_{3}^{} +H_{3}O^{+}
but almost instantaneous in the presence of the catalytic enzyme carbonic anhydrase.
Pure compounds
When pure substances (liquids or solids) are involved in equilibria they do not appear in the equilibrium equation^{[11]}
Applying the general formula for an equilibrium constant to the specific case of acetic acid one obtains
It may be assumed that the concentration of water is constant. This assumption will be valid for all but very concentrated solutions. The equilibrium constant expression is therefore usually written as
where now
a constant factor is incorporated into the equilibrium constant.
A particular case is the selfionization of water itself
The selfionization constant of water is defined as
It is perfectly legitimate to write [H^{+}] for the hydronium ion concentration, since the state of solvation of the proton is constant (in dilute solutions) and so does not affect the equilibrium concentrations. K_{w} varies with variation in ionic strength and/or temperature.
The concentrations of H^{+} and OH^{} are not independent quantities. Most commonly [OH^{}] is replaced by K_{w}[H^{+}]^{−1} in equilibrium constant expressions which would otherwise hydroxide.
Solids also do not appear in the equilibrium equation. An example is the Boudouard reaction^{[11]}:
for which the equation (without solid carbon) is written as:
Multiple equilibria
Consider the case of a dibasic acid H_{2}A. When dissolved in water, the mixture will contain H_{2}A, HA^{} and A^{2}. This equilibrium can be split into two steps in each of which one proton is liberated.
K_{1} and K_{2} are examples of stepwise equilibrium constants. The overall equilibrium constant,β_{D}, is product of the stepwise constants.
Note that these constants are dissociation constants because the products on the right hand side of the equilibrium expression are dissociation products. In many systems, it is preferable to use association constants.
β_{1} and β_{2} are examples of association constants. Clearly β_{1} = 1/K_{2} and β_{2} = 1/β_{D}; lg β_{1} = pK_{2} and lg β_{2} = pK_{2} + pK_{1}^{[12]} For multiple equilibrium systems, also see: theory of Response reactions.
Effect of temperature
The effect of changing temperature on an equilibrium constant is given by the van 't Hoff equation
Thus, for exothermic reactions, (ΔH is negative) K decreases with an increase in temperature, but, for endothermic reactions, (ΔH is positive) K increases with an increase temperature. An alternative formulation is
At first sight this appears to offer a means of obtaining the standard molar enthalpy of the reaction by studying the variation of K with temperature. In practice, however, the method is unreliable because error propagation almost always gives very large errors on the values calculated in this way.
Types of equilibrium
 In the gas phase. Rocket engines^{[13]}
 The industrial synthesis such as ammonia in the HaberBosch process (depicted right) takes place through a succession of equilibrium steps including adsorption processes.
 atmospheric chemistry
 Seawater and other natural waters: Chemical oceanography
 Distribution between two phases
 LogDDistribution coefficient: Important for pharmaceuticals where lipophilicity is a significant property of a drug
 Liquidliquid extraction, Ion exchange, Chromatography
 Solubility product
 Uptake and release of oxygen by haemoglobin in blood
 Acid/base equilibria: Acid dissociation constant, hydrolysis, buffer solutions, indicators, acidbase homeostasis
 Metalligand complexation: sequestering agents, chelation therapy, MRI contrast reagents, Schlenk equilibrium
 Adduct formation: Hostguest chemistry, supramolecular chemistry, molecular recognition, dinitrogen tetroxide
 In certain oscillating reactions, the approach to equilibrium is not asymptotically but in the form of a damped oscillation .^{[11]}
 The related Nernst equation in electrochemistry gives the difference in electrode potential as a function of redox concentrations.
 When molecules on each side of the equilibrium are able to further react irreversibly in secondary reactions, the final product ratio is determined according to the CurtinHammett principle.
In these applications, terms such as stability constant, formation constant, binding constant, affinity constant, association/dissociation constant are used. In biochemistry, it is common to give units for binding constants, which serve to define the concentration units used when the constant’s value was determined.
Composition of a mixture
When the only equilibrium is that of the formation of a 1:1 adduct as the composition of a mixture, there are any number of ways that the composition of a mixture can be calculated. For example, see ICE table for a traditional method of calculating the pH of a solution of a weak acid.
There are three approaches to the general calculation of the composition of a mixture at equilibrium.
 The most basic approach is to manipulate the various equilibrium constants until the desired concentrations are expressed in terms of measured equilibrium constants (equivalent to measuring chemical potentials) and initial conditions.
 Minimize the Gibbs energy of the system.^{[14]}
 Satisfy the equation of mass balance. The equations of mass balance are simply statements that demonstrate that the total concentration of each reactant must be constant by the law of conservation of mass.
Massbalance equations
In general, the calculations are rather complicated. For instance, in the case of a dibasic acid, H_{2}A dissolved in water the two reactants can be specified as the conjugate base, A^{2}, and the proton, H^{+}. The following equations of massbalance could apply equally well to a base such as 1,2diaminoethane, in which case the base itself is designated as the reactant A:
With T_{A} the total concentration of species A. Note that it is customary to omit the ionic charges when writing and using these equations.
When the equilibrium constants are known and the total concentrations are specified there are two equations in two unknown "free concentrations" [A] and [H]. This follows from the fact that [HA]= β_{1}[A][H], [H_{2}A]= β_{2}[A][H]^{2} and [OH] = K_{w}[H]^{−1}
so the concentrations of the "complexes" are calculated from the free concentrations and the equilibrium constants. General expressions applicable to all systems with two reagents, A and B would be
It is easy to see how this can be extended to three or more reagents.
Polybasic acids
The composition of solutions containing reactants A and H is easy to calculate as a function of p[H]. When [H] is known, the free concentration [A] is calculated from the massbalance equation in A. Here is an example of the results that can be obtained.
This diagram, for the hydrolysis of the aluminium Lewis acid Al^{3+}_{aq}^{[15]} shows the species concentrations for a 5×10^{−6}M solution of an aluminium salt as a function of pH. Each concentration is shown as a percentage of the total aluminium.
Solution and precipitation
The diagram above illustrates the point that a precipitate that is not one of the main species in the solution equilibrium may be formed. At pH just below 5.5 the main species present in a 5μM solution of Al^{3+} are aluminium hydroxides Al(OH)^{2+}, Al(OH)_{2}^{+} and Al_{13}(OH)_{32}^{7+}, but on raising the pH Al(OH)_{3} precipitates from the solution. This occurs because Al(OH)_{3} has a very large lattice energy. As the pH rises more and more Al(OH)_{3} comes out of solution. This is an example of Le Chatelier's principle in action: Increasing the concentration of the hydroxide ion causes more aluminium hydroxide to precipitate, which removes hydroxide from the solution. When the hydroxide concentration becomes sufficiently high the soluble aluminate, Al(OH)_{4}^{}, is formed.
Another common instance where precipitation occurs is when a metal cation interacts with an anionic ligand to form an electricallyneutral complex. If the complex is hydrophobic, it will precipitate out of water. This occurs with the nickel ion Ni^{2+} and dimethylglyoxime, (dmgH_{2}): in this case the lattice energy of the solid is not particularly large, but it greatly exceeds the energy of solvation of the molecule Ni(dmgH)_{2}.
Minimization of free energy
At equilibrium, G is at a minimum:
For a closed system, no particles may enter or leave, although they may combine in various ways. The total number of atoms of each element will remain constant. This means that the minimization above must be subjected to the constraints:
where a_{ij} is the number of atoms of element i in molecule j and b_{i}^{0} is the total number of atoms of element i, which is a constant, since the system is closed. If there are a total of k types of atoms in the system, then there will be k such equations.
This is a standard problem in optimisation, known as constrained minimisation. The most common method of solving it is using the method of Lagrange multipliers, also known as undetermined multipliers (though other methods may be used).
Define:
where the λ_{i} are the Lagrange multipliers, one for each element. This allows each of the N_{j} to be treated independently, and it can be shown using the tools of multivariate calculus that the equilibrium condition is given by
 and
(For proof see Lagrange multipliers)
This is a set of (m+k) equations in (m+k) unknowns (the N_{j} and the λ_{i}) and may, therefore, be solved for the equilibrium concentrations N_{j} as long as the chemical potentials are known as functions of the concentrations at the given temperature and pressure. (See Thermodynamic databases for pure substances).
This method of calculating equilibrium chemical concentrations is useful for systems with a large number of different molecules. The use of k atomic element conservation equations for the mass constraint is straightforward, and replaces the use of the stoichiometric coefficient equations.^{[13]}
See also
 Autocatalytic reactions and order creation
 BenesiHildebrand method
 Determination of equilibrium constants
 Equilibrium constant
 Henderson–Hasselbalch equation
 Isotope fractionation
 MichaelisMenten kinetics
 Redox equilibria
 Thermodynamic databases for pure substances
References
 ^ ^{a} ^{b} Atkins, Peter W and Jones, Loretta Chemical Principles: The Quest for Insight 2nd Ed. ISBN 0716799030
 ^ IUPAC, Compendium of Chemical Terminology, 2nd ed. (the "Gold Book") (1997). Online corrected version: (2006–) "chemical equilibrium".
 ^ Chemistry: Matter and Its Changes James E. Brady , Fred Senese 4th Ed. ISBN 0471215171
 ^ Atkins, Peter W. and Julio de Paula Physical Chemistry, 4th Edition, WileyVCH, Weinheim 2006, ISBN 9783527315468
 ^ Atkins, Peter W. Physical Chemistry, third edition, Oxford University Press, 1985.
 ^ Schultz, Mary Jane (1999). "Why Equilibrium? Understanding Entropy of Mixing". Journal of Chemical Education 76: 1391. doi:10.1021/ed076p1391.
 ^ Clugston, Michael J. (1990). "A mathematical verification of the second law of thermodynamics from the entropy of mixing". Journal of Chemical Education 67: 203. doi:10.1021/ed067p203.
 ^ C.W. Davies, Ion Association, Butterworths, 1962
 ^ ^{a} ^{b} I. Grenthe and H. Wanner, Guidelines for the extrapolation to zero ionic strength
 ^ F.J.C. Rossotti and H. Rossotti, The Determination of Stability Constants, McGrawHill, 1961
 ^ ^{a} ^{b} ^{c} Concise Encyclopedia Chemistry 1994 ISBN 0899254578
 ^ M.T. Beck, Chemistry of Complex Equilibria, Van Nostrand, 1970. 2nd. Edition by M.T. Beck and I Nagypál, Akadémiai Kaidó, Budapest, 1990.
 ^ ^{a} ^{b} NASA Reference publication 1311, Computer Program for Calculation of Complex Chemical Equilibrium Compositions and Applications
 ^ This approach is described in detail in W. R. Smith and R. W. Missen, Chemical Reaction Equilibrium Analysis: Theory and Algorithms, , Krieger Publishing, Malabar, Fla, 1991 (a reprint, with corrections, of the same title by John Wiley & Sons, 1982). A comprehensive treatment of the theory of chemical reaction equilibria and its computation. Details at http://www.mathtrek.com/
 ^ The diagram was created with the program HySS
Further reading
 F. Van Zeggeren and S.H. Storey, The Computation of Chemical Equilibria, Cambridge University Press, 1970. Mainly concerned with gasphase equilibria.
 D. J. Leggett (editor), Computational Methods for the Determination of Formation Constants, Plenum Press, 1985.
 A.E. Martell and R.J. Motekaitis, The Determination and Use of Stability Constants, WileyVCH, 1992.
 P. Gans, Stability Constants: Determination and Uses, an interactive CD, Protonic Software (Leeds), 2004
External links
Computer programs
There are n massbalance equations in n unknown free concentrations. This constitutes a set of nonlinear equations that must be solved by a method of successive approximations. The most commonlyused method is the NewtonRaphson method, which has been the subject of numerous publications. Some general computer programs are listed here.
 GeochemEZ (freeware) a multipurpose chemical speciation program, used in plant nutrition and in soil and environmental chemistry research to perform equilibrium speciation computations, allowing the user to estimate solution ion activities and to consider simple complexes and solid phases.
 HySS Titration simulation and speciation calculations.
 EQS4WIN A powerful computer program originally developed for gasphase equilibria but subsequently extended to general applications. Uses the Gibbs energy minimization approach.
 CHEMEQL A comprehensive computer program for the calculation of thermodynamic equilibrium concentrations of species in homogeneous and heterogeneous systems. Many geochemical applications.
 WinSGW A Windows version of the SOLGASWATER computer program.
 Visual MINTEQ A Windows version of MINTEQA2 (ver 4.0). MINTEQA2 is a chemical equilibrium model for the calculation of metal speciation, solubility equilibria etc. for natural waters.
 MINEQL+ A chemical equilibrium modeling system for aqueous systems. Handles a wide range of pH, redox, solubility and sorption scenarios.
Software
 Aqua solution software (Sukhno Igor, Buzko Vladimir, Polushin Alexey) A set of six computer programs for
 Specific Interaction Theory. An editable database of published SIT parameters. Estimation of SIT parameters and adjustment of stability constants for changes in ionic strength.
 Calculation of electrolyte activity coefficients, ionic activity coefficients, osmotic coefficients
 Calculation of acidbase equilibria in electrolyte solutions and sea water
 Calculation of O_{2} solubility in water, electrolyte solutions, natural fluids, and seawater as a function of temperature, concentration, salinity, altitude, external pressure, humidity
 Prediction of temperature dependence of lg K values using various thermodynamic models
 JESS:A powerful research tool for thermodynamic and kinetic modelling of chemical speciation in complex aqueous environments, with a focus on determining thermodynamic consistency by automatic means.
 Chemical Equilibrium Calculator
 Mission to Mars – A chemistry tutorial for high school students
Chemical equilibria Concepts Acid dissociation constant · Binding constant · Binding selectivity · Buffer solution · Chemical equilibrium · Chemical stability · Chelation · Determination of equilibrium constants · Dissociation constant · Distribution coefficient · Distribution ratio · Dynamic equilibrium · Equilibrium chemistry · Equilibrium constant · Equilibrium unfolding · Equilibrium stage · Henry's law · Liquidliquid extraction · Macrocycle effect · Phase diagram · Predominance diagram · Phase rule · Reaction quotient · Selfionization of water · Solubility equilibrium · Stability constants of complexes · Thermodynamic equilibrium · Vaporliquid equilibrium
Categories: Equilibrium chemistry
 Analytical chemistry
 Physical chemistry
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