1. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Completion Reactions and Reversible Reactions If enough oxygen gas is provided for the following reaction, almost all of the sulfur will react: S 8 (s) + 8O 2 (g) → 8SO 2 (g) Reactions such as this one, in which almost all of the reactants react, are called completion reactions. In other reactions, called reversible reactions, the products can re-form reactants. Section 1 Reversible Reactions and Equilibrium Chapter 14

2. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Completion Reactions and Reversible Reactions, continued Reversible Reactions Reach Equilibrium One reversible reaction occurs when you mix solutions of calcium chloride and sodium sulfate. CaCl 2 (aq) + Na 2 SO 4 (aq) → CaSO 4 (s) + 2NaCl(aq) The net ionic equation best describes what happens. Section 1 Reversible Reactions and Equilibrium Chapter 14

3. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Completion Reactions and Reversible Reactions, continued Reversible Reactions Reach Equilibrium, continued Solid calcium sulfate, the product, can break down to make calcium ions and sulfate ions in a reaction that is the reverse of the previous one. Section 1 Reversible Reactions and Equilibrium Chapter 14 Use arrows that point in opposite directions when writing a chemical equation for a reversible reaction.

4. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Completion Reactions and Reversible Reactions, continued Reversible Reactions Reach Equilibrium, continued The reactions occur at the same rate after the initial mixing of CaCl 2 and Na 2 SO 4. The amounts of the products and reactants do not change. Chemical equilibrium is a state of balance in which the rate of a forward reaction equals the rate of the reverse reaction and the concentrations of products and reactants remain unchanged. Section 1 Reversible Reactions and Equilibrium Chapter 14

5. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Completion Reactions and Reversible Reactions, continued Opposing Reaction Rates Are Equal at Equilibrium The reaction of hydrogen, H 2, and iodine, I 2, to form hydrogen iodide, HI, reaches chemical equilibrium. Section 1 Reversible Reactions and Equilibrium Chapter 14 Only a very small fraction of the collisions between H 2 and I 2 result in the formation of HI. H 2 (g) + I 2 (g) → 2HI(g)

6. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Completion Reactions and Reversible Reactions, continued Opposing Reaction Rates Are Equal at Equilibrium, continued After some time, the concentration of HI goes up. As a result, fewer collisions occur between H 2 and I 2 molecules, and the rate of the forward reaction drops. Similarly, in the beginning, few HI molecules exist in the system, so they rarely collide with each other. Section 1 Reversible Reactions and Equilibrium Chapter 14

7. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Completion Reactions and Reversible Reactions, continued Opposing Reaction Rates Are Equal at Equilibrium, continued As more HI molecules are made, they collide more often and form H 2 and I 2 by the reverse reaction. 2HI(g) → H 2 (g) + I 2 (g) The greater the number of HI molecules that form, the more often the reverse reaction occurs. Section 1 Reversible Reactions and Equilibrium Chapter 14

8. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Rate Comparison for H 2 (g) + I 2 (g)  2HI(g) Section 1 Reversible Reactions and Equilibrium Chapter 14

9. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Completion Reactions and Reversible Reactions, continued Opposing Reaction Rates Are Equal at Equilibrium, continued When the forward rate and the reverse rate are equal, the system is at chemical equilibrium. If you repeated this experiment at the same temperature, starting with a similar amount of pure HI instead of the H 2 and I 2, the reaction would reach chemical equilibrium again and produce the same concentrations of each substance. Section 1 Reversible Reactions and Equilibrium Chapter 14

10. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Chemical Equilibrium Visual Concepts Chapter 14

11. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Chemical Equilibria Are Dynamic If you drop a ball into a bowl, it will bounce. When the ball comes to a stop it has reached static equilibrium, a state in which nothing changes. Chemical equilibrium is different from static equilibrium because it is dynamic. In a dynamic equilibrium, there is no net change in the system. Two opposite changes occur at the same time. Section 1 Reversible Reactions and Equilibrium Chapter 14

12. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Chemical Equilibria Are Dynamic, continued In equilibrium, an atom may change from being part of the products to part of the reactants many times. But the overall concentrations of products and reactants stay the same. For chemical equilibrium to be maintained, the rates of the forward and reverse reactions must be equal. Arrows of equal length also show equilibrium. Section 1 Reversible Reactions and Equilibrium Chapter 14

13. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Chemical Equilibria Are Dynamic, continued In some cases, the equilibrium has a higher concentration of products than reactants. This type of equilibrium is also shown by using two arrows. Section 1 Reversible Reactions and Equilibrium Chapter 14 The forward reaction has a longer arrow to show that the products are favored.

14. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Le Châtelier’s Principle Stress is another word for something that causes a change in a system at equilibrium. Chemical equilibrium can be disturbed by a stress, but the system soon reaches a new equilibrium. Le Châtelier’s principle states that when a system at equilibrium is disturbed, the system adjusts in a way to reduce the change. Section 3 Equilibrium Systems and Stress Chapter 14

15. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Le Châtelier’s Principle, continued Chemical equilibria respond to three kinds of stress: changes in the concentrations of reactants or products changes in temperature changes in pressure When a stress is first applied to a system, equilibrium is disturbed and the rates of the forward and backward reactions are no longer equal. Section 3 Equilibrium Systems and Stress Chapter 14

16. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Le Châtelier’s Principle, continued The system responds to the stress by forming more products or by forming more reactants. A new chemical equilibrium is reached when enough reactants or products form. At this point, the rates of the forward and backward reactions are equal again. Section 3 Equilibrium Systems and Stress Chapter 14

17. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Visual Concepts Le Châtelier's Principle Chapter 14

18. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Le Châtelier’s Principle, continued Changes in Concentration Alter Equilibrium Composition If you increase a reactant’s concentration, the system will respond to decrease the concentration of the reactant by changing some of it into product. Therefore, the rate of the forward reaction must be greater than the rate of the reverse reaction. The equilibrium is said to shift right, and the reactant concentration drops until the reaction reaches equilibrium. Section 3 Equilibrium Systems and Stress Chapter 14

19. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Le Châtelier’s Principle, continued Changes in Concentration Alter Equilibrium Composition, continued In a reaction of two colored complex ions: Section 3 Equilibrium Systems and Stress Chapter 14 When the reaction mixture in a beaker is pale blue, we know that chemical equilibrium favors the formation of reactants. pale blue blue-purple

20. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Le Châtelier’s Principle, continued Changes in Concentration Alter Equilibrium Composition, continued If ammonia is added, the system responds by forming more product and the solution becomes blue-purple. Section 3 Equilibrium Systems and Stress Chapter 14

21. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Le Châtelier’s Principle, continued Changes in Concentration Alter Equilibrium Composition, continued The equilibrium below occurs in a bottle of soda. Section 3 Equilibrium Systems and Stress Chapter 14 After you uncap the bottle, the dissolved carbon dioxide leaves the solution and enters the air. The forward reaction rate of this system will increase to produce more CO 2.

22. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Le Châtelier’s Principle, continued Changes in Concentration Alter Equilibrium Composition, continued This increase in the rate of the forward reaction decreases the concentration of H 3 O +. As a result, the drink gets “flat.” If you could increase the concentration of CO 2 in the bottle, the reverse reaction rate would increase, and would increase. Section 3 Equilibrium Systems and Stress Chapter 14

23. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Le Châtelier’s Principle, continued Temperature Affects Equilibrium Systems The effect of temperature on the gas-phase equilibrium of nitrogen dioxide, NO 2, and dinitrogen tetroxide, N 2 O 4, can be seen because of the difference in color of NO 2 and N 2 O 4. The intense brown NO 2 gas is the pollution that is responsible for the colored haze that you sometimes see on smoggy days. Section 3 Equilibrium Systems and Stress Chapter 14

24. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Le Châtelier’s Principle, continued Temperature Affects Equilibrium Systems, continued Recall that endothermic reactions absorb energy and have positive ∆H values. Exothermic reactions release energy and have negative ∆H values. The forward reaction is an exothermic process, as the equation below shows. 2NO 2 (g) → N 2 O 4 (g)∆H = −55.3 kJ Section 3 Equilibrium Systems and Stress Chapter 14

25. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Temperature Changes Affect an Equilibrium System Section 3 Equilibrium Systems and Stress Chapter 14

26. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Le Châtelier’s Principle, continued Temperature Affects Equilibrium Systems, continued Now suppose that you heat the flask to 100°C. The mixture becomes dark brown because the reverse reaction rate increased to remove some of the energy that you added to the system. The equilibrium shifts to the left, toward the formation of NO 2. Because this reaction is endothermic, the temperature of the flask drops as energy is absorbed. Section 3 Equilibrium Systems and Stress Chapter 14

27. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Le Châtelier’s Principle, continued Temperature Affects Equilibrium Systems, continued This equilibrium shift is true for all exothermic forward reactions: Increasing the temperature of an equilibrium mixture usually leads to a shift in favor of the reactants. The opposite statement is true for endothermic forward reactions: Increasing the temperature of an equilibrium mixture usually leads to a shift in favor of the products. Section 3 Equilibrium Systems and Stress Chapter 14

28. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Le Châtelier’s Principle, continued Temperature Affects Equilibrium Systems, continued The following is an endothermic reaction involving two colored cobalt complex ions: Section 3 Equilibrium Systems and Stress Chapter 14 The forward reaction is endothermic, so the forward reaction is favored at high temperatures. The reverse reaction is favored at low temperatures.

29. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Le Châtelier’s Principle, continued Temperature Affects Equilibrium Systems, continued Equilibrium changes with changes in temperature they affect the value of equilibrium constants. Consider K eq for the ammonia synthesis equilibrium: Section 3 Equilibrium Systems and Stress Chapter 14 The forward reaction is exothermic (∆H = −91.8 kJ), so the equilibrium constant decreases a lot as temperature increases.

30. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Visual Concepts Factors Affecting Equilibrium Chapter 14

31. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Le Châtelier’s Principle, continued Pressure Changes May Alter Systems in Equilibrium Changes in pressure can affect gases in equilibrium. The NO 2 and N 2 O 4 equilibrium can show the effect of a pressure stress on a chemical equilibrium. 2NO 2 (g) → N 2 O 4 (g) When the gas mixture has a larger concentration of N 2 O 4 than of NO 2, it has a pale color. Section 3 Equilibrium Systems and Stress Chapter 14

32. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Le Châtelier’s Principle, continued Pressure Changes May Alter Systems in Equilibrium When the gas is suddenly compressed to about half its former volume, the pressure doubles. Section 3 Equilibrium Systems and Stress Chapter 14

33. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Le Châtelier’s Principle, continued Pressure Changes May Alter Systems in Equilibrium, continued Before the system has time to adjust to the pressure stress, the concentration of each gas doubles, which results in a darker color. Le Châtelier’s principle predicts that the system will adjust in an attempt to reduce the pressure. According to the equation, 2 mol of NO 2 produce 1 mol of N 2 O 4. Section 3 Equilibrium Systems and Stress Chapter 14

34. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Le Châtelier’s Principle, continued Pressure Changes May Alter Systems in Equilibrium, continued At constant volume and temperature, pressure is proportional to the number of moles. So, the pressure reduces when there are fewer moles of gas. Thus, the equilibrium shifts to the right, and more N 2 O 4 is produced. Section 3 Equilibrium Systems and Stress Chapter 14

35. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Le Châtelier’s Principle, continued Pressure Changes May Alter Systems in Equilibrium, continued In an equilibrium, a pressure increase favors the reaction that produces fewer gas molecules, which for the following equilibrium is the reverse reaction. Section 3 Equilibrium Systems and Stress Chapter 14 When there is no change in the number of molecules, a change in pressure will not affect equilibrium.

36. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Equilibrium Visual Concepts Chapter 14

37. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Practical Uses of Le Châtelier’s Principle The chemical industry makes use of Le Châtelier’s principle in the synthesis of ammonia by the Haber Process. High pressure is used to drive the following equilibrium to the right. Section 3 Equilibrium Systems and Stress Chapter 14 The forward reaction converts 4 mol of gas into 2 mol of another gas, so it is favored at high pressures.

38. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Practical Uses of Le Châtelier’s Principle, continued The ammonia synthesis is an exothermic reaction, so the forward reaction is favored at low temperatures. 0°CK eq = 6.5 × 10 8 250°CK eq = 52 500°CK eq = 5.8 × 10 −2 The Haber Process is operated at temperatures of 500°C even though the K eq is small at that temperature. The reaction proceeds too slowly at lower temperatures. Section 3 Equilibrium Systems and Stress Chapter 14

39. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu 1.In which of these reactions is the formation of the products favored by an increase in pressure? A.2O 3 (g)  3O 2 (g) B.C(s) + O 2 (g)  CO 2 (g) C.2NO(g) + O 2 (g)  2NO 2 (g) D. Standardized Test Preparation Understanding Concepts Chapter 14

40. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu 1.In which of these reactions is the formation of the products favored by an increase in pressure? A.2O 3 (g)  3O 2 (g) B.C(s) + O 2 (g)  CO 2 (g) C.2NO(g) + O 2 (g)  2NO 2 (g) D. Standardized Test Preparation Understanding Concepts Chapter 14

41. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu 2.What is the effect of an increase in temperature on an exothermic reaction at equilibrium? F.It has no effect on the equilibrium. G.It shifts the equilibrium in favor of the forward reaction. H.It shifts the equilibrium in favor of the reverse reaction. I.It shifts the equilibrium in favor of both the forward and reverse reactions. Standardized Test Preparation Understanding Concepts Chapter 14

42. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu 2.What is the effect of an increase in temperature on an exothermic reaction at equilibrium? F.It has no effect on the equilibrium. G.It shifts the equilibrium in favor of the forward reaction. H.It shifts the equilibrium in favor of the reverse reaction. I.It shifts the equilibrium in favor of both the forward and reverse reactions. Standardized Test Preparation Understanding Concepts Chapter 14

43. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu 5.Explain how pressure can be used to maximize the production of carbon dioxide in the reaction Standardized Test Preparation Understanding Concepts Chapter 14

44. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu 5.Explain how pressure can be used to maximize the production of carbon dioxide in the reaction Answer: Increasing the pressure will cause the reaction to favor the production of carbon dioxide because there are three gas molecules in the reactants and only two in the products. Standardized Test Preparation Understanding Concepts Chapter 14