Chemical equilibrium

In a dynamic equilibrium ^[1] ^[2]

Concepts in Chemical Equilibria
Acid dissociation constant
Binding constant
Chemical equilibrium
Dissociation constant
Distribution coefficient
Distribution ratio
Equilibrium constant
Equilibrium unfolding
Equilibrium stage
Liquid-liquid extraction
Phase diagram
Phase rule
Reaction quotient
Relative volatility
Solubility equilibrium
Stability constant
Thermodynamic equilibrium
Theoretical plate
Vapor-liquid equilibrium
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Introduction

In a chemical reaction, when reactants are mixed together in a reaction vessel (and heated if needed), the whole of reactants do not get converted into the products. After some time (which may be shorter than millionths of a second or longer than the age of the universe), there will come a point when a fixed amount of reactants will exist in harmony with a fixed amount of products, the amounts of neither changing anymore. This is called chemical equilibrium.

The concept of chemical equilibrium was developed after Berthollet (1803) found that some chemical reactions are reversible. For any reaction such as

$\alpha A + \beta B \rightleftharpoons \sigma S + \tau T$

to be at equilibrium the chemical equation with harpoon 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 used up and 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:

$\mbox{forward reaction rate} = k_+ {A}^\alpha{B}^\beta \,\!$

$\mbox{backward reaction rate} = k_{-} {S}^\sigma{T}^\tau \,\!$

where A, B, S and T are rate constants. Since forward and backward rates are equal:

$k_+ {A}^\alpha{B}^\beta = k_{-} {S}^\sigma{T}^\tau \,$

and the ratio of the rate constants is also a constant, now known as an equilibrium constant.

$K=\frac{k_+}{k_-}=\frac{\{S\}^\sigma \{T\}^\tau } {\{A\}^\alpha \{B\}^\beta}$

By convention the products form the numerator. Unfortunately, the 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 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.^[3] ^[4]

Although the macroscopic equilibrium concentrations are constant in time reactions do occur at the molecular level. For example, in the case of hydronium ions,

CH₃CO₂H + H₂O ⇌ CH₃CO₂⁻ + H₃O⁺

a proton may hop from one molecule of ethanoic acid on to a water molecule and then on to an ethanoate ion to form another molecule of ethanoic acid and leaving the number of ethanoic acid molecules unchanged. This is an example of dynamic equilibrium. Equilibriums, like the rest of thermodynamics, are statistical phenomena, averages of microscopic behaviour.

Le Chatelier's principle (1884) is a useful principle that gives a qualitative idea of an equilibrium system's response to changes in reaction conditions. If a dynamic equilibrium is disturbed by changing the conditions, the position of equilibrium moves to counteract 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 ethanoic 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:

$K=\frac{\{CH_3CO_2^-\}\{H_3O^+\}} {\{CH_3CO_2H\}\{H_2O \}}$

if {H₃O⁺} increases {CH₃CO₂H} must increase and {CH₃CO₂⁻} must decrease.

A quantitative version is given by the reaction quotient.

Gibbs energy change for the reaction by the equation

$\Delta G^\ominus = -RT \ln K_{eq}$

where R is the universal gas constant and T the temperature.

When the reactants are activity coefficients may be taken to be constant. In that case the concentration quotient, K_c,

$K_c=\frac{[S]^\sigma [T]^\tau } {[A]^\alpha [B]^\beta}$

where [A] is the partial pressure. These constants are easier to measure and encountered in high-school chemistry courses.

Thermodynamics

The relationship between the Gibbs energy and the equilibrium constant can be found by considering chemical potentials. The thermodynamic condition for chemical equilibrium is^[6]

At constant pressure ΔG=0 (ΔG is the change Gibbs free energy for the reaction)
At constant volume ΔA=0 (ΔA is the change in Helmholtz free energy for the reaction)

In this article only the constant pressure case is considered. The constant volume case is important in 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.^[7]

In general an equilibrium system is defined by writing an equilibrium equation for the reaction

$\alpha A + \beta B \rightleftharpoons \sigma S + \tau T$

To meet the thermodynamic condition for equilibrium the Gibbs energy must be stationary, meaning that the derivative of G with respect to reaction coordinate (ΔG) must be zero. It can be shown that ΔG is in fact equal to the difference between the chemical potentials of the products and those of 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.

$\alpha \mu_A + \beta \mu_B = \sigma \mu_S + \tau \mu_T \,$

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.

$\mu_A = \mu_{A}^{\ominus} + RT \ln\{A\} \,$

Substituting expressions like this into the Gibbs energy equation:

$\Delta G = Vdp-SdT+\sum_{i=1}^k \mu_i dN_i + \sum_{i=1}^n X_i da_i + \cdots \,$

which at constant pressure and temperature becomes:

$\Delta G =\sum_{i=1}^k \mu_i N_i$

results in:

$\Delta G = \sigma \mu_{S} + \tau \mu_{S} - \alpha \mu_{S} - \beta \mu_{S} \,$

By substituting the chemical potentials:

$\Delta G = ( \sigma \mu_{S}^{\ominus} + \tau \mu_{S}^{\ominus} ) - ( \alpha \mu_{S}^{\ominus} - \beta \mu_{S}^{\ominus} ) + ( \sigma RT \ln\{S\} + \tau RT \ln\{T\} ) - ( \alpha RT \ln\{A\} + \beta RT \ln \{B\} )$

the relationship becomes:

$\Delta G =\sum_{i=1}^k \mu_i^\ominus v_i + RT \ln \frac{\{S\}^\sigma \{T\}^\tau} {\{A\}^\alpha \{B\}^\beta}$

At equilibrium $\Delta G = 0 \,$ and therefore

$\sum_{i=1}^k \mu_i^\ominus v_i + RT \ln \frac{\{S\}^\sigma \{T\}^\tau} {\{A\}^\alpha \{B\}^\beta} = 0$

leading to:

$\Delta G_m^{\ominus} = -RT \ln K$

ΔG_m^O is the standard molar Gibbs energy change for the reaction and K is the equilibrium constant. Note that activities and equilibrium constants are dimensionless numbers.

Treatment of activity

The expression for the equilibrium constant can be re-written as the product of a concentration quotient, K_c and an activity coefficient quotient, Γ.

$K=\frac{{[S]} ^\sigma {[T]}^\tau ... } {{[A]}^\alpha {[B]}^\beta ...} \times \frac{{\gamma_S} ^\sigma {\gamma_T}^\tau ... } {{\gamma_A}^\alpha {\gamma_B}^\beta ...} = K_c \Gamma$

[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 Debye-Hückel equation or extensions such as Davies equation^[8] 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 pedantically accurate concentration quotient. This practice will be followed here.

For reactions in the gas phase ammonia industrially, 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

$\mu = \mu^{\Theta} + RT \ln \left( \frac{f}{bar} \right) + RT \ln \gamma$

so the general expression defining an equilibrium constant is valid for both solution and gas phases.

Justification for the use of concentration quotients

In aqueous solution, equilibrium constants are usually determined in the presence of an "inert" electrolyte such as ionic strength, I, of a solution containing a dissolved salt, X⁺Y^-, is given by

$I = \frac{1}{2}\left(c_X z_X^2 + c_Y z_Y^2 + \sum_{i=1}^n c_i z_i^2\right)$

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]

$K_c = \frac{K}{\Gamma}$

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 SO₃.

2SO₂ + O₂ $\rightleftharpoons$ 2SO₃

The barrier can be overcome when a Contact process, but the catalyst does not affect the equilibrium concentrations.

Likewise, the formation of carbon dioxide and water is very slow under normal conditions

CO₂ + 2H₂O $\rightleftharpoons$ HCO₃^- +H₃O⁺

but almost instantaneous in the presence of the catalytic carbonic anhydrase.

Pure compounds in equilibria

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 ethanoic acid one obtains

$CH_3CO_2H + H_2O \rightleftharpoons CH_3CO_2^- + H_3O^+$

$K_c=\frac{[{CH_3CO_2}^-][{H_3O}^+]} {[{CH_3CO_2H}][{H_2O}]}$

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

$K=\frac{[{CH_3CO_2}^-][{H_3O}^+]} {[{CH_3CO_2H}]}$

where now

$K=K_c*[H_2O]\,$

a constant factor is incorporated into the equilibrium constant.

A particular case is the self-ionization of water itself

$H_2O + H_2O \rightleftharpoons H_3O^+ + OH^-$

The self-ionization constant of water is defined as

$K_w = [H^+][OH^-]\,$

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]:

$2CO \rightleftharpoons CO_2 + C$

for which the equation (without solid carbon) is written as:

$K_c=\frac{[CO_2]} {[CO]^2}$

Multiple equilibria

Consider the case of a dibasic acid H₂A. When dissolved in water the mixture will contain H₂A, HA^- and A^2-. This equilibrium can be split into two steps in each of which one proton is liberated.

$H_2A \rightleftharpoons HA^- + H^+ :K_1=\frac{[HA^-][H^+]} {[H_2A]}$

$HA^- \rightleftharpoons A^{2-} + H^+ :K_2=\frac{[A^{2-}][H^+]} {[HA^-]}$

K₁ and K₂ are examples of stepwise equilibrium constants. The overall equilibrium constant, $β D$ , is product of the stepwise constants.

$H_2A \rightleftharpoons A^{2-} + 2H^+ :\beta_D = \frac{[A^{2-}][H^+]^2} {[H_2A]}=K_1K_2$

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.

$A^{2-} + H^+ \rightleftharpoons HA^- :\beta_1=\frac {[HA^-]} {[A^{2-}][H^+]}$

$A^{2-} + 2H^+ \rightleftharpoons H_2A :\beta_2=\frac {[H_2A]} {[A^{2-}][H^+]^2}$

β₁ and β₂ are examples of association constants. Clearly β₁ = 1/K₂ and β₂ = 1/β_D; lg β₁ = pK₂ and lg β₂ = pK₂ + pK₁^[12]

Effect of temperature change on an equilibrium constant

The effect of changing temperature on an equilibrium constant is given by the van 't Hoff equation

$\frac {d\ln K} {dT} = \frac{{\Delta H_m}^{\Theta}} {RT^2}$

Thus, for endothermic reactions (ΔH is positive) K increases with temperature. An alternative formulation is

$\frac {d\ln K} {d(1/T)} = -\frac{{\Delta H_m}^{\Theta}} {R}$

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 and some applications

In the gas phase. Rocket engines ^[13]
The industrial synthesis such as adsorbtion processes.

atmospheric chemistry.
Seawater and other natural waters. Chemical oceanography.
Distribution between two phases.
1. LogD-Distribution coefficient Important for pharmaceuticals where lipophilicity is a significant property of a drug.
2. Liquid-liquid extraction, Ion exchange, Chromatography.
3. Solubility product.
4. Uptake and release of oxygen by haemoglobin in blood
Acid/base equilibria. indicators, acid-base homeostasis
Metal-ligand complexation. Schlenk equilibrium
Adduct formation. 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 Curtin-Hammett 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 an equilibrium 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 the total concentration of each reactant must be constant by the law of conservation of mass.

Solving the equations of mass-balance

In general the calculations are rather complicated. For instance, in the case of a dibasic acid, H₂A dissolved in water the two reactants can be specified as the 1,2-diaminoethane, in which case the base itself is designated as the reactant A:

$T_A = [A] + [HA] +[H_2A] \,$

$T_H = [H] + [HA] + 2[H_2A] - [OH] \,$

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]= β₁[A][H], [H₂A]= β₂[A][H]² and [OH] = K_w[H]^-1

$T_A = [A] + \beta_1[A][H] + \beta_2[A][H]^2 \,$

$T_H = [H] + \beta_1[A][H] + 2\beta_2[A][H]^2 - K_w[H]^{-1} \,$

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

$T_A=[A]+\sum_i{p_i \beta_i[A]^{p_i}[B]^{q_i}}$

$T_B=[B]+\sum_i{q_i \beta_i[A]^{p_i}[B]^{q_i}}$

It is easy to see how this can be extended to three or more reagents.

Composition for polybasic acids as a function of pH

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 mass-balance equation in A. Here is an example of the results that can be obtained.

This diagram, for the hydrolysis of the aluminum Lewis acid Al³⁺_aq ^[15] shows the species concentrations for a 5×10^-6M solution of an aluminium salt as a function of pH. Each concentration is shown as a percentage of the total aluminium.

Solution equilibria with precipitation

The diagram above illustrates the point that a 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 electrically neutral complex. If the complex is solvation of the molecule Ni(dmgH)₂.

Minimization of Gibbs energy

At equilibrium, G is at a minimum:

$dG= \sum_{j=1}^m \mu_j\,dN_j = 0$

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:

$\sum_{j=1}^m a_{ij}N_j=b_i^0$

where $a i j$ is the number of atoms of element i in molecule j and b_i⁰ 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:

$\mathcal{G}= G + \sum_{i=1}^k\lambda_i\left(\sum_{j=1}^m a_{ij}N_j-b_i^0\right)=0$

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

$\frac{\partial \mathcal{G}}{\partial N_j}=0$ and $\frac{\partial \mathcal{G}}{\partial \lambda_i}=0$

(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]

References

^ Atkins & Jones, 2001
^ Gold Book definition Link
^ Chemistry: Matter and Its Changes James E. Brady , Fred Senese 4th Ed. ISBN 0471215171
^ Chemical Principles: The Quest for Insight Peter Atkins, Loretta Jones 2nd Ed. ISBN 0716757010
^ Physical Chemistry by Atkins, De Paula
^ P.W. Atkins, Physical Chemistry, Oxford University Press, date
^ a) Mary Jane Schultz. Why Equilibrium? Understanding the Role of Entropy of Mixing. Journal of Chemical Education 1999, 76, 1391. b) Clugston, Michael J. A mathematical verification of the second law of thermodynamics from the entropy of mixing. Journal of Chemical Education 1990, 67, 203.
^ C.W. Davies, Ion Association,Butterworths, 1962
^ ^a ^b I. Grenthe and H. Wanner, Guidelines for the extrapolation to zero ionic strength, http://www.nea.fr/html/dbtdb/guidelines/tdb2.pdf
^ F.J,C. Rossotti and H. Rossotti, The Determination of Stability Constants, McGraw-Hill, 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 Wiley-Interscience, 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