The Equilibrium Constant

Submitted by ChemPRIME Staff on Thu, 12/16/2010 - 14:41

The constancy of the ratio of the equilibrium concentration of one isomerOne of two or more substances that have the same molecular formula (are composed of elements in the same proportion and have the same molar mass), but differ in some aspect of structure. to the concentration of the other at a given temperatureA physical property that indicates whether one object can transfer thermal energy to another object. is characteristic of all gaseous equilibria between isomers, i.e.,of all reactions of the general type


\text{A}\text{ }({g})\rightleftharpoons \text{B}\text{ }({g})      (1)



The constant ratio of concentrations is called the equilibrium constantThe value of the equilibrium constant expression when equilibrium concentrations are substituted; a value greater than one indicates the position of equilibrium lies toward products (product-favored), and a value less than one indicates the position of equilibrium lies toward reactants (reactant-favored). and is given the symbol Kc. For reactions of the type given by Eq. (1) the equilibrium constant is thus described by the equation

K_{c}=\frac{[\text{ B }]}{[\text{ A }]}      (2)

where, by convention, the concentration of the productA substance produced by a chemical reaction. B appears in the numerator of the ratio. If, for some reason, we wish to look at this reaction in reverse,


\text{B}\text{ }({g})\rightleftharpoons \text{A}\text{ }({g})



then the equilibrium constant is denoted as

K_{c}=\frac{[\text{ A }]}{[\text{ B }]}


i.e., it is the reciprocal of the constant given in Eq. (2).

In general the equilibrium constant Kc varies with temperature and also differs from one reaction to another. Examples illustrating this behavior are given in Table 1 where the experimentally determined equilibrium constants for various cisDescribes the relationship between two atoms or groups of atoms, each attached to one of two doubly bonded carbon atoms and located on the same side of the double bond. Also refers to groups located adjacent to each other in an octahedral or square planar coordination complex.-transDescribes the relationship between two atoms or groups of atoms, each attached to one of two doubly bonded carbon atoms and located on the opposite side of the double bond. Also refers to groups located opposite to each other in an octahedral or square planar coordination complex. isomerization equilibria are recorded at various temperatures.


TABLE 1 The Equilibrium Constant Kc for some Cis-Trans Interconversions. Image:Equilibrium Constant Kc for some Cis-Trans Interconversions.jpg


When we turn our attention to more complexA central metal and the ligands surrounding it; also called coordination complex. equilibrium reactions, we find that the relationship between the concentrations of the various species is no longer a simple ratio. A good demonstration of this fact is provided by the dissociationThe breaking apart of one species into two or more smaller species; often applied to ions in a crystal lattice, which dissociate when the ionic solid dissolves in water. Dissociation refers to separation of particles that already exist; ionization refers to the formation of ions from neutral species, as in the ionization of a weak acid in aqueous solutoin. of dinitrogen tetroxide, N2O4. This compoundA substance made up of two or more elements and having those elements present in definite proportions; a compound can be decomposed into two or more different substances. is a colorless gas, but even at room temperature it dissociates partly into a vivid red-brown gas, NO2, according to the equation


\text{N}_{2}\text{O}_{4}\text{ }({g})\rightleftharpoons \text{2NO}_{2} \text{ }({g})      (3)


If 1 mol N2O4 contained in a flask of volume 1 dm3 is heated to 407.2 K, exactly one-half of it dissociates into NO2. If the volume is now increased, the ratio of [NO2] to [N2O4] does not remain constant but increases as more dissociates. As shown in Table 13.2, if we increase the volume still further, even more dissociation occurs. By the time we have increased the volume to 10 dm3, the fraction of N2O4 molecules dissociated has increased to 0.854 (i.e., to 85.4 percent).

Obviously the situation is now not quite so straightforward as in the previous example. Nevertheless there is a simple relationship between the equilibrium concentrations of the reactantA substance consumed by a chemical reaction. and product in this case too. We find that it is the quantity


\frac{[\text{ NO}_{2}]^{2}}{[\text{ N}_{2}\text{O}_{4}]}


rather than the simple ratio of concentrations, which is now constant. Accordingly we also call this quantity an equilibrium constant and give it the symbol Kc. Thus Kc for Eq. (3) is given by the relationship


K_{c}=\frac{[\text{ NO}_{2}]^{2}}{[\text{ N}_{2}\text{O}_{4}]}


where again by convention the product appears in the numerator. It is easy to check that Kc actually is a constant quantity with the value 2.00 mol dm–3 from the data given in Table 2. Thus if we take the result from line d, we find that when 1 mol N2O4 is placed in a 10 dm3 flask at 407


TABLE 2 The Dissociation of 1 mol N2O4 into NO2 at 407.2 K (134°C) and Various Volumes.

Amount of N2O4 Added

mol

Volume of Flask

dm3

Fraction N2O4 Dissociated Amount N2O4 at Equilibrium

mol

Amount NO2 at Equilibrium

mol

Concentration N2O4 at Equilibrium

mol dm-3

Concentration NO2 at Equilibrium

mol dm-3

Equilibrium Constant Kc

mol dm-3

  a   1     1   0.500   0.500   1.000   0.500   1.000   2.000
  b   1     2   0.618   0.382   1.236   0.191   0.618   2.000
  c   1     5   0.766   0.234   1.532   0.0468   0.3064   2.006
  d   1   10   0.854   0.146   1.708   0.0146   0.1708   1.998


K, 0.854 mol dissociate. Since from Eq. (3) each moleThat chemical amount of a substance containing the same number of units as 12 g of carbon-12. which dissociates yields 2 mol NO2,there will be


\text{0.854 mol N}_{2}\text{O}_{4}\times \frac{\text{2 mol NO}_{2}}{\text{1 mol N}_{2}\text{O}_{4}}=\text{1.708 mol NO}_{2}


present in the reaction vessel. There will also be (1 – 0.854) mol = 0.146 mol N2O4 left undissociated in the flask. Since the total volume is 10 dm3, the equilibrium concentrations are

[\text{ NO}_{2}]=\frac{\text{1.708 mol}}{\text{10 dm}^{3}}=\text{0.1708 mol dm}^{-3}

and

[\text{ N}_{2}\text{O}_{4}]=\frac{\text{0.146 mol}}{\text{10 dm}^{3}}=\text{0.0146 mol dm}^{-3}


Accordingly

K_{c}=\frac{\text{1.708 mol dm}^{-3}\times \text{ 0.1708 mol dm}^{-3}}{\text{0.0146 mol dm}^{-3}}=\text{2.00 mol dm}^{-3}


In exactly the same way, if we use the data from line a in Table 2, we find

K_{c}=\frac{[\text{ NO}_{2}]^{2}}{[\text{ N}_{2}\text{O}_{4}]}=\frac{\text{(1.00 mol dm}^{3}\text{)}^{2}}{\text{0.5 mol dm}^{-3}}=\text{2.00 mol dm}^{-3}

You can check for yourself that lines b and c also yield the same value forKc.



EXAMPLE 1 When 2 mol N2O4 gas is heated to 407 K in a vessel of volume 5 dm3, it is found that 0.656 of the molecules dissociate into NO2. Show that these data are in agreement with the value for Kc of 2.00 mol dm–3 given in the text.


Solution Many equilibrium problems can be solved in a fairly standardized fashion in three stages.


a) Calculate the amount of each substanceA material that is either an element or that has a fixed ratio of elements in its chemical formula. transformed by the reaction as it comes to equilibrium, i.e., the amount of each reactant consumed by the reaction and the amount of each product produced by the reaction. Stoichiometric ratios derived from the equation must always be used in these calculations.

In this particular example we note that 0.656 of the original N2O4 dissociates. Since 2 mol was used, a total of 0.656 × 2 = 1.312 mol N2O4 is consumed. The amount of NO2 produced is accordingly


n_{\text{NO}_{2}}=\text{1.312 mol N}_{2}\text{O}_{4}\times \frac{\text{2 mol NO}_{2}}{\text{1 mol N}_{2}\text{O}_{4}}=\text{2.624 mol NO}_{2}


b) Use the amounts calculated in the first stage to calculate the amount of each substance present at equilibrium. Dividing by the volume, we can obtain the equilibrium concentrations.

Since 1.312 mol N2O4 dissociated out of an original 2 mol, we have (2 – 1.312) mol = 0.688 mol N2O4 left. The equilibrium concentration of N2O4 is thus


[\text{ N}_{2}\text{O}_{4}] =  \frac{0.688\text{ mol N}_{2}\text{O}_{4}}{\text{ 5.00 dm}^{3}}  = \text{0.1376 mol dm}^{-3}


Since no NO2 was originally present, the amount of NO2 present at equilibrium is the amount produced by the dissociation, namely, 2.624 mol NO2. Thus


[\text{ NO}_{2}]=\frac{\text{2.624 mol NO}_{2}}{\text{5.00 dm}^{3}}=\text{0.525 mol dm}^{-3}


It is usually worthwhile tabulating these calculations, particularly in more complex examples.

Note that a negative quantity in the column headed Amount Produced indicates that a given substance (such as N2O4 in this example) has been consumed. There is less of that substance when equilibrium is reached than was present initially.

Substance Initial Amount

mol

Amount Produced

mol

Equilibrium Amount

mol

Equilibrium Concentration

mol dm-3

  N2O4   2.00   -1.312   0.688   0.688/5
  NO2   0.00   2.624   2.624   2.624/5


c) In the third stage we insert the equilibrium concentrations in an expression for the equilibrium constant:


K_{c}=\frac{[\text{ NO}_{2}]^{2}}{[\text{ N}_{2}\text{O}_{4}]}=\frac{\text{0.525 mol dm}^{-3}\times \text{ 0.525 mol dm}^{-3}}{\text{0.1376 mol dm}^{-3}}=\text{2.00 mol dm}^{-3}