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Chemical Equilibrium

Chemical Equilibrium. Reversible Reactions. Most chemical reactions are able to proceed in both directions. They are reversible . Example: Fe 3 O 4 (s) + 4 H 2 (g) ↔ 3 Fe (s) + 4 H 2 O (g) https://www.youtube.com/watch?v=g5wNg_dKsYY. Reversible Reactions cont.

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Chemical Equilibrium

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  1. Chemical Equilibrium

  2. Reversible Reactions • Most chemical reactions are able to proceed in both directions. They are reversible. • Example: Fe3O4 (s) + 4 H2 (g)↔ 3 Fe(s) + 4 H2O(g) https://www.youtube.com/watch?v=g5wNg_dKsYY

  3. Reversible Reactions cont. • As products are produced they will react in the reverse reaction until the rates of the forward and reverse reactions are equal. Ratefwd = Raterev • This is called chemical equilibrium.

  4. Chemical Equilibrium • Rateforward reaction = Ratereverse reaction • Concentration of reactants and products are constant NOT necessarily equal. • [reactants] and [products] is constant.

  5. The Concept of Equilibrium • As a system approaches equilibrium, both the forward and reverse reactions are occurring. • At equilibrium, the forward and reverse reactions are proceeding at the same rate.

  6. Le Chatelier’s Principle • Whenever stress is applied to a reaction at equilibrium, the reaction will shift its point of equilibrium to offset the stress. • A “shift” means that the forward or reverse direction will occur more. • Stresses include: • Temperature, pressure, changes in reactant or product concentrations

  7. Example: The Haber Process N2 (g) + 3 H2 (g) 2 NH3 (g) + heat •  [N2] •  [H2] •  [NH3] •  [NH3] •  pressure •  pressure •  temperature •  temperature shift towards products (right) shift towards reactants (left) shift towards reactants (left) shift towards products (right) shift towards products (right) shift towards reactants (left) shift towards reactants (left) shift towards products (right)

  8. Equilibrium shifts due to stresses: • Concentration increase: shift away from increase • Concentration decrease: shift toward decrease •  pressure: shifts in direction of fewer gas molecules. •  pressure: shifts in direction of more gas molecules •  temperature favors endothermic reaction • Shift away from heat •  temperature favors exothermic reaction • Shift towards heat

  9. The Haber Process • Application of LeChatelier’s Principle N2 (g) + 3 H2 (g) 2 NH3 (g) + 92 kJ increase pressure Shift  decrease Temp Shift  remove NH3 add N2 and H2 Shift  ****Maximum yields of NH3 occurs under high pressures, low temperatures and by constantly removing NH3 and adding N2 & H2

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