Le Chatelier’s Principle
1. For each of the following chemical systems at equilibrium, use Le Chatelier’s principle to predict the effect of the change imposed on the chemical system. Indicate the direction in which the equilibrium is expected to shift. Assume that the systems are closed and that they are initially at equilibrium.
a) H2O(l) +energy ↔ H2O(g)
The container is heated.
b) H2O(l) ↔H+(aq) + OH-(aq)
A few crystals of NaOH are added to the container.
c) CaCO3(s) + energy↔ CaO(s) + CO2(g)
CO2(g) is removed from the container.
d) CH3COOH(aq) ↔ H+(aq) + CH3COO-(aq)
A few drops of CH3COOH(aq) are added to the system.
2. The following equation represents part of the industrial production of nitric acid. Predict the direction of the equilibrium shift for each of the following changes.
4NH3(g) + 5O2(g) ↔ 4NO(g) + 6H2O(g) + Energy
a) O2(g) is added to the system
b) The temperature of the system is increased
c) NO(g) is removed from the system
d) The pressure of the system is increased by decreasing the volume
3. The following chemical equilibrium sustem is part of the Haber process for the production of ammonia.
N2(g) +3H2(g) ↔ 2NH3(g) + energy
Suppose you are a chemical process engineer. Use Le Chatelier’s principle to predict four specific changes that you might impose on the equilibrium system to increase the yield of ammonia.
4. Nitrogen monoxide, a major pollutant, is formed in automobile engines from the endothermic reaction of nitrogen and oxygen gases.
a) Write the equilibrium reaction including the term “energy” in the equation.
b) How will the equilibrium shift if oxygen is added?
c) How will the equilibrium shift if the volume is decreased?
5. Use the following equation to predict what color change would be visible when stresses are applied.
Fe3+ (aq) + SCN-(aq) ↔ FeSCN2+(aq)
Yellow Colorless Red
a) A crystal of KSCN(s) is added
b) A crystal of FeCl3(aq)
c) Pressure is increased
Equilibrium Position
1. For each of the following write the chemical equation with the appropriate equilibrium arrows.
a) pH measurements indicate the acetic acid in vinegar is approximately 1% ionized into hydrogen ions and acetate ions.
b) Analysis of the reaction of sodium sulfate and calcium chloride solutions shows that the products are favored
c) Aluminum sulfate solution reacts quantitatively with sodium hydroxide solution.
d) The Haber process is used to manufacture ammonia fertilizer from hydrogen and nitrogen gases. Under less then desirable conditions, only an 11% yield of ammonia is obtained at equilibrium.
e) Because of the cost of silver, many science departments recover silver metal from waste solutions containing silver compounds or silver ions. A quantitative reaction of waste silver ion solutions with copper metal results in the production of silver metal and copper (II) ions.
f) One step in the industrial process used to manufacture sulfuric acid is the production of sulfur trioxide from sulfur dioxide and oxygen gases. Under certain conditions the reaction produces a 65% yield of products.
Equilibrium Constant- Keq
1. In an experiment, 0.500 mol/L of hydrogen bromide gas is decomposed into hydrogen and bromine gases.
a) Write the equilibrium equation and equilibrium law for this reaction.
b) The equilibrium concentrations in this system are [HBr(g)] =0.240 mol/L and [H2(g)]=0.130 mol/L and [Br2(g)] = 0.130 mol/L. Calculate Keq.
2.