Le Chatelier’s Principle

 When a chemical system at equilibrium is disturbed by a change in a property of the system, the system always appears to react in the direction that opposes the change (until a new equilibrium is reached)

In my own words:

Concentration Change  If you increase the concentration on one side, the shift will be in the opposite direction

 That is, if you add reactant, the equilibrium will shift towards the products  If you add products, the equilibrium will shift towards the reactants

 This is because when you add concentration, more molecules are available to react, creating an increased forward reaction

Example  The production of freon-12 (a CFC refrigerant) involves the following equilibrium reaction: CCl4(l) + 2HF(g) --> CCl2F2(g) + 2HCl(g)  To improve the yield of freon-12 (CCL2F2), more hydrogen fluoride is added to the initial equilibrium system, shifting it to the right.

Temperature Change  Whether energy is added or removed, the equilibrium shifts to minimize the change in energy  Depends on if the reaction is exothermic or endothermic

Example  In the salt-sulfuric acid process, used to produce HCl, the system is heated in order to increase the percent yield of hydrogen chloride gas: 2NaCl(s) + H2SO4(l) + energy --> 2HCl(g) + Na2SO4  Adding energy shifts the equilibrium to the right to absorb some of the energy

Example 2  In the production of sulfuric acid, the key reaction step is the equilibrium represented by the following: 2SO2(g) + O2(g) --> 2SO3(g) + energy  Products are increased by removing energy, causing the system to replace the energy lost and shift to the right

Pressure and Volume Change  According to Boyle’s Law, the concentration of a gas is directly proportional to its pressure

 If the volume is decreased, the concentration increases and the number of molecules will decrease (when possible)

Example  In the equilibrium reaction of sulfur dioxide and oxygen, three moles of gaseous reactants produce 2 moles of gaseous products: 2SO2(g) + O2(g) --> 2SO3(g)  If the volume is decreased, the overall pressure increases and this causes the reaction to shift right, which decreases the number of gas molecules

Example 2  A system with equal numbers of gas molecules on each side (i.e. H2 + I2 --> 2HI) is not affected by change in volume

NOTE on Gases  Adding or removing gas not involved in the equilibrium will not influence the equilibrium

Catalyst Reactions  Catalysts decrease the time required to reach an equilibrium position, but does not affect the final position of equilibrium.

 Catalysts affect both forward and reverse reactions at the same rate  Does not influence the equilibrium

Graphing Changes  Changes result in a quick spike  Le Chatelier’s allows for the gradual return to equilibrium by shifting in the opposite direction

Example  The Haber-Bosch process produces ammonia from nitrogen and hydrogen gas. It is an important process for adding nitrates to fertilizers and was used in the manufacture of explosives during the Second World War. Graph the effects of the following changes: N2(g) + 3H2(g) --> 2NH3(g) + heat

 Increase N2  Response = decrease N2 (forward reaction)

N2(g) + 3H2(g) --> 2NH3(g) + heat  Cool reaction  Response = Increase temp (forward reaction)

N2(g) + 3H2(g) --> 2NH3(g) + heat  Decrease NH3  Response = increase NH3 (forward reaction)

N2(g) + 3H2(g) --> 2NH3(g) + heat  Add Catalyst  Response = increase rxn rate – no change

N2(g) + 3H2(g) --> 2NH3(g) + heat  Decrease pressure  Response = increase pressure to more moles (reverse reaction)

N2(g) + 3H2(g) --> 2NH3(g) + heat  Decrease the volume  Response = increase pressure, shift towards less moles (forward reaction)