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Le Chatelier’s Principle

Le Chatelier’s Principle. Chem 12 - Unit 3. Le Chatelier’s Principle. The French chemist Henri Le Chatelier (1850-1936) studied how the equilibrium position shifts as a result of changing conditions. LeChatelier’s Principle.

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Le Chatelier’s Principle

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  1. Le Chatelier’s Principle Chem 12 - Unit 3

  2. Le Chatelier’s Principle • The French chemist Henri Le Chatelier (1850-1936) studied how the equilibrium position shifts as a result of changing conditions

  3. LeChatelier’s Principle • When a system at equilibrium is placed under stress, the system will undergo a change in such a way as to relieve that stress.

  4. Le Chatelier Translated: • When you take something away from a system at equilibrium, the system shifts in such a way as to replace what you’ve taken away. • When you add something to a system at equilibrium, the system shifts in such a way as to use up what you’ve added.

  5. What items did he consider to be stress on the equilibrium? • Concentration • Temperature • Pressure

  6. Examples: • Let’s look at some examples, there will be a summary table at the end!

  7. Le Chatelier Example #1 A closed container of ice and water at equilibrium. The temperature is raised. Ice + Energy  Water The equilibrium of the system shifts to the _______ to use up the added energy. right

  8. Le Chatelier Example #2 A closed container of N2O4 and NO2 at equilibrium.NO2 is addedto the container. N2O4 (g) + Energy  2 NO2(g) The equilibrium of the system shifts to the _______ to use up the added NO2. left

  9. Le Chatelier Example #3 A closed container of water and its vapor at equilibrium.Vapor is removedfrom the system. water + Energy  vapor The equilibrium of the system shifts to the _______ to replace the vapor. right

  10. Le Chatelier Example #4 A closed container of N2O4 and NO2 at equilibrium. Thepressureisincreased. N2O4 (g) + Energy  2 NO2(g) The equilibrium of the system shifts to the _______ to lower the pressure, because there are fewer moles of gas on that side of the equation. left

  11. Pressure Changes to system: • If the volume decreases, the concentration increases, and there will be a shift to the side with the less amount of moles. • If the volume increases, the concentration decreases, and there will be a shift to the side with the more amount of moles.

  12. Another Example: • If I increase the pressure, where is the shift? (right) • If I decrease the pressure, where is the shift? (left) 2SO2 + O2<--> 2SO3 (3moles) (2moles)

  13. Effect of Concentration: • If you add more reactant, it shifts to the right increasing the formation of product, using up the reactants. • If you add product, it shifts to the left • If you remove product, it shifts to the right, increasing the formation of product. • If you remove reactant, it shifts to the left

  14. Effect of temperature: • Remember: Energy is treated as a reactant if endothermic equation and as a product if exothermic equation. • If heating a system, it shifts so extra heat is used up. • If cooling a system, then it shifts so more heat is produced.

  15. Example Endothermic Reaction: • Heating this reaction causes the system to shift to the right = more products, because you treat energy like a reactant (this reaction needs energy to go forward) 2NaCl +H2SO4 + energy < -- > 2HCl + Na2SO4 • Cooling this reaction causes the system to shift to the left = less reactants, so need to make up more

  16. Example Exothermic Reaction: • Heating this reaction causes the system to shift to the left, to use up the extra heat. 2SO2 + O2<--> 2SO3 + energy • Cooling this reaction causes the system to shift to the right, to make up for the lost heat.

  17. Common Ion Effect: • Involves adding of an ion to a solution in which the ion is already present in the solution • Increases concentration of that ion • Eq shifts away from that ion Example: Cu2+(aq) + 4Cl-1 <-> CuCl42+ (aq) Adding NaCl would increase the Cl-1 concentration and shift Eq right!

  18. Try it : • Handouts • Page 529 #33-37

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