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Equilibrium

Equilibrium. L. Scheffler Lincoln High School 2010. 1. Equilibrium Systems. Many chemical reactions are reversible. Such reactions do not go to completion. A state of balance exists between the products and the reactants

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Equilibrium

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  1. Equilibrium L. Scheffler Lincoln High School 2010. 1

  2. Equilibrium Systems • Many chemical reactions are reversible. • Such reactions do not go to completion. • A state of balance exists between the products and the reactants • When the concentration of neither the reactants nor the products is changing, the system is in equilibrium. • Note: this does NOT mean the reactions have stopped. Just the rates are equal. 2

  3. Chemical Equilibrium Chemical equilibrium occurs in chemical reactions that are reversible. In a reaction such as: CH4(g) + H2O(g)  CO(g) + 3H2 (g) The reaction can proceed in both directions CO(g) + 3H2 (g)  CH4(g) + H2O(g) 3

  4. Equilibrium Conditions • At equilibrium, the rate of reaction in the forward direction and the rate in the reverse direction are equal. 4

  5. An Equilibrium System CH4(g) + H2O(g) CO(g) + 3H2 (g) • After some of the products are created, the products begin to react to form the reactants. • At equilibrium, there is NO net change in the concentrations of the reactants and products. • The concentrations do not change but they are not necessarily equal. 5

  6. H2O NaCl (s)  NaCl (aq) Dynamic Equilibrium • At equilibrium, two opposing processes are taking place at equal rates. • In other words, the rate in the forward direction = the rate in the reverse direction • Examples H2O (l)  H2O (g) CO (g) + 2 H2 (g)  CH3OH (g) 6

  7. Equilibrium Conditions H2O + CO  H2 + CO2 7

  8. Law of Mass Action • Given the reaction aA + bB cC + dD The rate in the forward direction is rate forward = kf[A]a [B]b The rate in the reverse direction is rate reverse = kr [C]c [D]d k is the rate constant. It is unique for each reaction @ a certain temperature. 8

  9. Law of Mass Action • At equilibrium these rates are equal rate forward = rate reverse kf[A]a [B]b = kr [C]c [D]d • The ratio of the rate constants is Keq = kf/kr 9

  10. Writing Equilibrium Expressions 1. N2(g) + 3 H2 (g) 2NH3 (g) 2. 2 SO2(g) + O2 (g)  2SO3 (g) 3. H2(g) + Br2 (g) 2 HBr(g) 4. 2N2O (g) 2 N2 (g) + O2 (g) 10

  11. Answers Check your work against the following: 1. N2(g) + 3 H2 (g) 2NH3 (g) 2. 2 SO2 (g) + O2 (g)  2SO3 (g) 3. H2 (g) + Br2 (g)  2 HBr(g) 4. 2N2O (g)  2 N2 (g) + O2 (g) 11

  12. Calculating Equilibrium Constants Nitrogen dioxide decomposes at high temperatures according to this equation: 2 NO2 (g)  2 NO (g) + O2(g) If the equilibrium concentrations are as follows: [NO2]= 1.20 M, [NO] = 0.160 M, and [O2] = 0.080 M; calculate the equilibrium constant. 12

  13. Calculating Equilibrium Constants The equilibrium equation for the oxidation of sulfur dioxide is as follows: 2SO2(g) + O2 (g) 2 SO3 (g) If the equilibrium concentrations are as follows: [SO2 ]= 0.44 M, [O2] = 0.22 M, and [SO3] = 0.78 M, Calculate the equilibrium constant 13

  14. Practice Problem 1 The equilibrium equation for the carbon monoxide with steam to produce hydrogen gas is as follows: CO2(g) + H2(g)  CO (g) + H2O (g) If the equilibrium concentrations are as follows: [CO] = 1.00 M, [H2O] = 0.025 M, [CO2] = 0.075 M and [H2] = 0.060 M, calculate the equilibrium constant. 14

  15. The Meaning of the Equilibrium Constant • Kc>>1: The reaction is product-favored; equilibrium concentrations of products are greater than equilibrium concentrations of reactants. • Kc<<1: Thereaction is reactant-favored; equilibrium concentrations of reactants are greater than equilibrium concentrations of products. 15

  16. The Meaning of the Equilibrium Constant So, what do those Keq values mean? 2 NO2 (g)  2 NO (g) + O2(g) Kc =0.00142 Kc<<1: Thereaction is reactant-favored 2SO2(g) + O2 (g)  2 SO3 (g)Kc = 14.3 Kc>>1: The reaction is product-favored CO2(g) + H2(g)  CO (g) + H2O (g) Kc= 5.6 Kc~1: The reaction is slightly product-favored 16

  17. Reaction Quotient • The equilibrium constant (Keqor Kc) is a constant ratio only when the system is in equilibrium. • Similar Equation: Q =[prod]/[reactants] • If the system it not at equilibrium, the ratio is known as a Reaction Quotient. (while reactions are “in progress”). • If the reaction quotient is equal to the equilibrium constant then the system is at equilibrium. 17

  18. Le Chatelier’s Principle 18

  19. Le Chatelier’s Principle • Le Chatelier's Principle states: When a system in chemical equilibrium is disturbed by a change of temperature, pressure, or a concentration, the system shifts in equilibrium composition in a way that tends to counteract this change of variable. • A change imposed on an equilibrium system is called a stress • The equilibrium always responds in such a way so as to counteract the stress 19

  20. Le Chatelier’s Principle • A stress is any sudden change in conditions that drives the system out of equilibrium. • When a stress is placed on an equilibrium system, the system will shift in such a way so as to lessen or mitigate the stress and restore equilibrium. • A stress usually involves a change in the temperature, pressure, or in the concentration of one or more of the substances that are in equilibrium. • Le Chatelier's principle predicts the direction of the change in the equilibrium. 20

  21. Applications ofLe Chatelier’s Principle N2 (g) +3 H2 (g)  2NH3(g) DH = - 92 kJ • Haber’s process for the production of ammonia is an example of an industrial equilibrium system. We will use this equilibrium as a model to explain the how Le Chatelier’s principle operates with the following stresses: • Change in the concentration of one of the components • Changes in pressure • Changes in temperature • Use of a catalyst 21

  22. Applications of Le Chatelier’s Principle N2 (g) +3 H2 (g)  2NH3(g) DH = - 92 kJ The equilibrium constant for this reaction is Any change in the concentration (or partial pressure) of any component will cause the equilibrium to shift in such a way so as to return the equilibrium constant to its original value. 22

  23. Le Chatelier’s Principle –The Concentration Effect N2 (g) +3 H2 (g)  2NH3(g) DH = - 92 kJ An increase in the concentration of N2 results in a decrease H2 and an increase in NH3 in such a way to keep the equilibrium constant the same 23

  24. Le Chatelier’s Principle – The Concentration Effect N2 (g) +3 H2 (g)  2NH3(g) DH = - 92 kJ Likewise an increase in the concentration ofH2 results in a decrease in N2 and an increase in NH3 in such a way to keep the equilibrium constant the same 24

  25. Le Chatelier’s Principle – The Concentration Effect N2 (g) +3 H2 (g)  2NH3(g) DH = - 92 kJ An increase in the concentration ofNH3 results in a increase in N2 and an increase in H2 in such a way to keep the equilibrium constant the same. 25

  26. Le Chatelier’s Principle – The Temperature Effect N2 (g) +3 H2 (g)  2NH3(g) DH = - 92 kJ • The reaction is exothermic in the forward direction. • An increase in the temperature would trigger a response in the heat consuming (endothermic) direction. • An increase in temperature therefore causes the reaction to shift in the reverse direction. Some NH3 decomposes to N2 and H2. 26

  27. Le Chatelier’s Principle –The Pressure Effect N2 (g) +3 H2 (g)  2NH3(g) DH = - 92 kJ • All molecules in the equilibrium are gases • When the reaction proceeds in the forward direction the number of moles are reduced from 4 to 2. • Pressure is proportional to the number of moles of gas • An increase in pressure therefore causes the reaction to move in the forward direction. Some N2 and H2 combine to form more NH3 27

  28. Le Chatelier’s Principle – The Effect of Catalysts N2 (g) +3 H2 (g)  2NH3(g) DH = - 92 kJ • Catalysts lower the activation energy 28

  29. Le Chatelier’s Principle – The Effect of Catalysts N2 (g) +3 H2 (g)  2NH3(g) DH = - 92 kJ • Catalysts lower the activation energy (prior knowledge) • A catalyst affects the forward and the reverse direction equally • There is no change in the equilibrium position from a catalyst • A catalyst decreases the time required for the system to achieve equilibrium 29

  30. Le Chatelier’s Principle – Summary 30

  31. Different types of equilibrium constants • Keq General designation for an equilibrium constant • Kc Equilibrium constant based on concentration • Kp Equilibrium constant based on pressure (gases) • Ksp Solubility product • Ka Acid equilibrium constant • Kb Base equilibrium constant • Kw Ion product of water 31

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