Chemical Equilibrium Chapter 15

15.1 – The Concept of Equilibrium Chemical Equilibrium = when the forward and reverse reactions proceed at equal rates (and the concentrations of reactants and products remain constant). Ex. For the reaction H 2 + I 2 2HI

15.2 – The Equilibrium Constant The Equilibrium Constant Expression = a ratio of the concentrations (M) of substances present in a system at equilibrium raised to powers equal to their mole coefficients. This ratio is a constant (K c ). It doesn’t depend on how equilibrium is reached. Ex. For the reaction aA + bB cC + dD

Example: Consider the Haber Process, which is the industrial preparation of ammonia: N 2 (g) + 3 H 2 (g) ⇌ 2 NH 3 (g) The equilibrium constant expression would be written as follows: K c depends only on the particular reaction and the temperature K p = the equilibrium constant when pressures are used K eq = the Thermodynamic Equilibrium Constant. Uses a combination of M and pressures K a and K b are for weak acids and bases equilibria (chp 16) K sp is for solubility equilibrium (chp 17) **no units are used for equil. constants See sample exercise 15.1

15.3 – Understanding and Working with Equilibrium Constants Significance of the Magnitude of an Equilibrium Constant: A very large K c or K p means that the forward reaction goes to completion or very nearly so (the equilibrium lies to the right and the equilibrium mixture contains mostly products). A very small K c or K p means that the forward reaction does not occur to any significant extent (the equilibrium lies to the left and the equilibrium mixture contains mostly reactants). See sample exercise 15.3

Relationship of K c (or K p ) to the balanced equation:  for reverse reactions, invert K c  if the coefficients are multiplied by a common factor, raise K c to the corresponding power  if the coeff. are divided by a common factor, take the corresponding root of K c.  if equations are combined, multiply their K c values to get K c for the net reaction.  See sample exercise 15.4

15.4 – Heterogeneous Equilibria Homogeneous equilibria occur when all reactants and products are in the same phase. Heterogeneous equilibria occur when something in the equilibrium is in a different phase. Whenever a pure solid or a pure liquid is involved in a heterogeneous equilibrium, its concentration is not included in the equilibrium constant expression because pure solid and liquid concentrations do not change. See sample exercises 15.5 and 15.6

15.5 – Calculating Equilibrium Constants A) If you are given the equilibrium concentrations of all the reactants and products, simply plug all the values into the equilibrium-constant expression see sample exercise 15.7 B) If you are given some initial concentrations and some equilibrium concentrations, then use the ICE approach (initial/change/equilibrium) See sample exercise 15.8

15.6 – Applications of Equilibrium Constants 1) Predicting the direction a reaction will go to reach equilibrium: use the reaction quotient (Q). Q uses the same expression as K c, but non-equilibrium concentrations are plugged into the expression. If Q c < K c, the reaction proceeds in the forward direction to reach equilibrium. If Q c > K c, the reaction proceeds in the reverse direction to reach equilibrium. If Q c = K c, the reaction is at equilibrium. See sample exercise 15.9

2) Using equilibrium constants to calculate equilibrium concentrations or partial pressures: --use the ICE approach with x representing one of the changes in concentration. Relate everything else to x, and solve for x. --Sometimes you will have to do a Q c to determine whether to add or subtract x. For both types, partial pressures can be used instead of concentrations. See sample exercises and 15.11

15.7 – Le Chatelier’s Principle When a change is imposed on a system at equilibrium, the system responds to minimize the impact of the change. Reactant and product concentrations change to accommodate the new situation (we say the equilibrium “shifts” to the right or left). The equilibrium constant remains the same.  If a reactant is added, the equilib. shifts 

 If a reactant is removed, the equilib. shifts   If a product is added, the equilib. shifts   If a product is removed, the equilib. shifts →  If a pure solid or liquid is added, there is no change

 If the pressure is increased on a gaseous system (by decreasing the volume of the container), the equilibrium shifts to the side with fewer moles of gas.  However, if the total pressure is increased by adding an inert gas to the mixture, no shift in equilibrium occurs, because the partial pressures of the reactants and products remain the same.

 Changes in temperature: depends on whether the reaction is endothermic or exothermic. Endothermic: Heats acts like a reactant; adding heat drives a reaction toward products. Exothermic: Heat acts like a product; adding heat drives a reaction toward reactants. Temperature changes also affect the value of K.  A catalyst does not affect the equilibrium position. It just increases the rate at which equilibrium is reached.  See sample exercises 15.12, 15.13