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Chapter: Modern Analytical Chemistry: Equilibrium Chemistry

Le Chatelier’s Principle

The equilibrium position for any reaction is defined by a fixed equilibrium constant, not by a fixed combination of concentrations for the reactants and products.

Le Chatelier’s Principle

The equilibrium position for any reaction is defined by a fixed equilibrium con- stant, not by a fixed combination of concentrations for the reactants and products. This is easily appreciated by examining the equilibrium constant expression for the dissociation of acetic acid.

As a single equation with three variables, equation 6.26 does not have a unique so- lution for the concentrations of CH3COOH, CH3COO, and H3O+. At constant temperature, different solutions of acetic acid may have different values for [H3O+], [CH3COO] and [CH3COOH], but will always have the same value of Ka.

If a solution of acetic acid at equilibrium is disturbed by adding sodium acetate, the [CH3COO] increases, suggesting an apparent increase in the value of Ka. Since Ka must remain constant, however, the concentration of all three species in equa- tion 6.26 must change in a fashion that restores Ka to its original value. In this case, equilibrium is reestablished by the partial reaction of CH3COO and H3O+ to pro- duce additional CH3COOH.

The observation that a system at equilibrium responds to a stress by reequili- brating in a manner that diminishes the stress, is formalized as Le Châtelier’s prin- ciple. One of the most common stresses that we can apply to a reaction at equilib- rium is to change the concentration of a reactant or product. We already have seen, in the case of sodium acetate and acetic acid, that adding a product to a reaction mixture at equilibrium converts a portion of the products to reactants. In this in- stance, we disturb the equilibrium by adding a product, and the stress is diminished by partially reacting the excess product. Adding acetic acid has the opposite effect, partially converting the excess acetic acid to acetate.

In our first example, the stress to the equilibrium was applied directly. It is also possible to apply a concentration stress indirectly. Consider, for example, the fol- lowing solubility equilibrium involving AgCl

The effect on the solubility of AgCl of adding AgNO3 is obvious,* but what is the ef- fect of adding a ligand that forms a stable, soluble complex with Ag+? Ammonia, for example, reacts with Ag+ as follows

Adding ammonia decreases the concentration of Ag+ as the Ag(NH3)2+ complex forms. In turn, decreasing the concentration of Ag+ increases the solubility of AgCl as reaction 6.27 reestablishes its equilibrium position. Adding together reactions 6.27 and 6.28 clarifies the effect of ammonia on the solubility of AgCl, by showing that ammonia is a reactant.

Increasing or decreasing the partial pressure of a gas is the same as increasing or decreasing its concentration. The effect on a reaction’s equilibrium position can be analyzed as described in the preceding example for aqueous solutes. Since the concentration of a gas depends on its partial pressure, and not on the total pressure of the system, adding or removing an inert gas has no effect on the equilibrium po- sition of a gas-phase reaction.

Most reactions involve reactants and products that are dispersed in a solvent. If the amount of solvent is changed, either by diluting or concentrating the solu- tion, the concentrations of all reactants and products either decrease or increase. The effect of these changes in concentration is not as intuitively obvious as when the concentration of a single reactant or product is changed. As an example, let’s consider how dilution affects the equilibrium position for the formation of the aqueous silver-amine complex (reaction 6.28). The equilibrium constant for this reaction is

where the subscript “eq” is included for clarification. If a portion of this solution is diluted with an equal volume of water, each of the concentration terms in equa- tion 6.30 is cut in half. Thus, the reaction quotient becomes

Since Q is greater than β2, equilibrium must be reestablished by shifting the reac- tion to the left, decreasing the concentration of Ag(NH3)2+. Furthermore, this new equilibrium position lies toward the side of the equilibrium reaction with the greatest number of solutes (one Ag+ ion and two molecules of NH3 versus the sin- gle metal–ligand complex). If the solution of Ag(NH3)2+ is concentrated, by evapo- rating some of the solvent, equilibrium is reestablished in the opposite direction. This is a general conclusion that can be applied to any reaction, whether gas-phase, liquid-phase, or solid-phase. Increasing volume always favors the direction pro- ducing the greatest number of particles, and decreasing volume always favors the direction producing the fewest particles. If the number of particles is the same on both sides of the equilibrium, then the equilibrium position is unaffected by a change in volume.


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Modern Analytical Chemistry: Equilibrium Chemistry : Le Chatelier’s Principle |

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