Chemical Equilibrium

The Equilibrium Condition For all of your chemistry career, especially when performing stoichiometry calculations, we assumed that all reactions proceed to completion. In other words, until one of the reactants ran out.

Many reactions, however, do not run to completion. Consider the dimerization of nitrogen dioxide: NO2(g) + NO2(g) → N2O4(g) The reactant, NO2, is a dark brown gas, and the product, N2O4, is a colorless gas. Notice when NO2 is placed in a sealed container, the initial brown color decreases in intensity as it is converted to colorless N2O4. However, the contents of the container do not become colorless. Instead the intensity of the brown color becomes constant, which mean the concentration of NO2 is no longer changing.

This reaction is illustrated below. The observation that the brown color does not completely fade is an indication the reaction has not reached completion (all NO2 is gone). The system has come to equilibrium. Chemical equilibrium: the state where the concentrations of all reactants and products remain constant with time.

Equilibrium is not static but a highly dynamic condition. Consider the flow of cars across a bridge connecting two cities with the traffic flow in both directions even. There is motion since the cars are moving, but the number of cars in each city is not changing because equal numbers of cars are entering and leaving. The result is no net change in the car population.

The same concept applies to a reaction at equilibrium. It may appear that everything has stopped since no changes occur in the concentrations of reactants or products. However, on the molecular level, there is frantic activity because both the forward and reverse reactions are occurring at an equal rate.

Consider the reaction between H2O and CO: H2O(g) + CO(g) ⇌ H2(g) + CO2(g) ⇌ (double arrows) indicate reaction can occur in either direction In (a) H2O and CO are mixed in equal numbers and begin to react (b) to form CO2 and H2. After time has passed, equilibrium is reached (c) and the numbers of reactant and product molecules then remain constant over time (d).

H2O(g) + CO(g) ⇌ H2(g) + CO2(g) Once equilibrium occurs, the concentration of the reactants and products remain unchanged. Unless the system is disturbed, no further changes in concentrations will occur. Remember that the reaction is still occurring but the forward and reverse rates are equal.

The Equilibrium constant In 1864 Guldberg and Waage proposed the law of mass action as a general description of the equilibrium condition. For a reaction of the type jA + kB ⇌ lC + mD where A, B, C, and D represent chemical species and j, k, l, and m are their coefficients, the law of mass action is represented by the following equilibrium expression: The square brackets indicate the concentration of the chemical species at equilibrium and K is a constant called the equilibrium constant.

Heterogeneous Equilibria Homogeneous equilibria: where all reactants and products are in the gaseous phase. Heterogeneous equilibria: involves more than one phase. Consider the reaction below: CaCO3(s) ⇌ CaO(s) + CO2(g) Experimental results show that the position of a heterogeneous equilibrium does not depend on the amounts of pure solids or liquids present.

Fundamental reason: concentrations of pure solids or liquids present cannot change. For any pure liquid or solid, the ratio of the amount of substance to volume of substance is a constant. Therefore the equilibrium expression for CaCO3(s) ⇌ CaO(s) + CO2(g) is written as K = [CO2]

The Meaning of K The value of K tells us how far a reaction proceeds to reach equilibrium. A value of K much larger than 1 means that at equilibrium, the reaction system will consist of mainly products – the equilibrium lies to the right. For example, consider a general reaction of the type A(g) → B(g) where K = [B]/[A]

If K for this reaction is 10,000 (104), then at equilibrium, [B]/[A] = 10,000 or = [B]/[A] = 10,000/1 That is, at equilibrium [B] is 10,000 times greater than [A]. Reaction strongly favors the product B, so essentially the reaction goes to completion. A small value of K means that the system at equilibrium consists largely of reactants – equilibrium lies far to the left.

The Meaning of K K > 1 → the equilibrium position is far to the right K < 1 → the equilibrium position is far to the left