Chapter Seven

1. Fundamentals of General, Organic and Biological Chemistry 5th Edition Chapter Seven Chemical Reactions: Energy, Rates, and Equilibrium James E. Mayhugh Oklahoma City University 2007 Prentice Hall, Inc.

2. There are many questions about chemical reactions 2 Au (s) + 3 H2O (l)  Au2O3 (s) + 3 H2 (g) This reaction is balanced but does not occur. Your gold jewelry is safe in the shower. Why do reactions occur? If it is balanced does that mean it will occur? Chapter Seven

3. Chapter 7 Goals 1. What Energy changes take place during a reaction? Explain the 2 factors that influence energy in a chemical reaction 2. What is “Free Energy,” and what is the criterion for spontaneity? Define enthalpy, entropy, and free energy change and explain how these values affect a chemical reaction 3. What determines the rate of a chemical reaction? Explain activation energy, and the other factor that determines the rate. Chapter Seven

4. Chapter 7 Goals 4. What is a chemical equilibrium? What occurs in a reaction at equilibrium and write the equilibrium constant expression. 5. What is Le Châtelier’s Principle? State Le Châtelier’s Principle, use it to predict the effect of temperature, pressure, and concentration on a reaction. Chapter Seven

5. Outline • 7.1 Energy and Chemical Bonds • 7.2 Heat Changes during Chemical Reactions • 7.3 Exothermic and Endothermic Reactions • 7.4 Why Do Chemical Reactions Occur? Free Energy • 7.5 How Do Chemical Reactions Occur? Reaction Rates • 7.6 Effects of Temperature, Concentration, and Catalysts on Reaction Rates • 7.7 Reversible Reactions and Chemical Equilibrium • 7.8 Equilibrium Equations and Equilibrium Constants • 7.9 Le Châtelier’s Principle: The Effect of Changing Conditions on Equilibria Chapter Seven

6. Chapter 7 2 factors determine if a reaction goes to products • Total Energy • Is the reaction fast or slow Can we influence the 2 factors above? Chapter Seven

7. 7.1 Energy and Chemical Bonds • There are two fundamental kinds of energy. • Potential energy is stored energy. The water in a reservoir behind a dam, an automobile poised to coast downhill, and a coiled spring have potential energy waiting to be released. • Kinetic energy is the energy of motion. When the water falls over the dam and turns a turbine, when the car rolls downhill, or when the spring uncoils and makes the hands on a clock move, the potential energy in each is converted to kinetic energy. Chapter Seven

8. 1. What Energy changes take place during a reaction? Explain the 2 factors that influence energy in a chemical reaction The first factor is Heat. The term in chemistry is called enthalpy, and it’s symbol is “H”. Mathematically, it is always treated as ΔH (Delta H, mean change in heat). We always compare the amount of heat at the beginning of a reaction, we call it bond disassociation, and compare it to the amount of heat at the end, we call it bond forming, of a chemical reaction; (the second factor is “S”, slide 19). Chapter Seven

9. 7.2 Heat Changes During Chemical Reactions • The first step in any reaction is Bond dissociation energy: The amount of energy that must be supplied to break a bond and separate the atoms in an isolated gaseous molecule. The second step will be Bond Formation energy. • The triple bond in N2 has a bond dissociation energy 226 kcal/mole, while the single bond in Cl2 has a bond dissociation energy 58 kcal/mole. Chapter Seven

10. Measure of Heat in a reaction We always want to know if the total heat is: Endothermic: A process or reaction that absorbs heat and has a positive DH. Exothermic: A process or reaction that releases heat and has a negative DH. Examples of each follow Chapter Seven

11. 7.3 Exothermic and Endothermic Reactions The seconds step is bond formation. When the total strength of the bonds formed in the products is greater than the total strength of the bonds broken in the reactants, energy is released and a reaction is exothermic. Chapter Seven

12. When the total energy of the bonds formed in the products is less than the total energy of the bonds broken in the reactants, energy is absorbed and the reaction is endothermic. Chapter Seven

13. A thermite reaction is: Fe2O3 + Al  Fe + Al2O3 ΔH = -851.3 KJ/mol Is this endothermic or exothermic? Reactions that give off heat are _________________? exothermic

14. When potassium is added to water contained in a beaker, the reaction shown below occurs, and the beaker feels hot to the touch. K(s) + 2 H2O(l)  2 KOH(aq) + H2(g)This reaction is • endothermic and H = – • endothermic and H = + • exothermic and H = – • exothermic and H = +

15. When potassium is added to water contained in a beaker, the reaction shown below occurs, and the beaker feels hot to the touch. K(s) + 2 H2O(l)  2 KOH(aq) + H2(g)This reaction is • endothermic and H = – • endothermic and H = + • exothermic and H = – • exothermic and H = +

16. This desert has 550 kcal of potential energy, heat releasing calories. Food is “potential energy.” Does our body use food in an endothermic or exothermic manner?

19. We’ve seen how heat is a factor in a chemical reaction (bond breaking, bond forming). Heat releasing (exothermic) reactions are often “favorable” to us. However, there is one more factor in a chemical reaction. 1. What Energy changes take place during a reaction? Explain the 2 factors that influence energy in a chemical reaction The second influence on energy is called Spontaneity. The symbol is S. Chapter Seven

20. Once at the top will water flow down

21. 7.4 Why Do Chemical Reactions Occur? Free Energy • Spontaneous process: A process that, once started, proceeds without any external influence. • Spontaneity does not care about how long it takes, only once it starts, it will go. For instance, it takes a long time for a car to rust, though it is spontaneous. Once gasoline is ignited, it will continue to burn on it’s own. • We measure spontaneity by measuring the “disorder” of a system. Once at the top will water flow down? Chapter Seven

22. 7.4 Why Do Chemical Reactions Occur? Free Energy • Entropy: The symbol S is used for entropy and it has the unit of cal/mole·K. The physical state of a substance and the number of particles have a large impact on the value of S. An example: solid CO2 has a lower entropy value then gaseous CO2…why? Chapter Seven

23. If I shake the beaker on the right, will they ever line up like the beaker on the left? The beaker on the right has a more “+” entropy value Chapter Seven

24. 7.4 Why Do Chemical Reactions Occur? Free Energy • Entropy: The symbol S is used for entropy and it has the unit of cal/mole·K. The physical state of a substance and the number of particles have a large impact on the value of S. An example: solid CO2 has a lower entropy value then gaseous CO2…why? Gasmolecules are more disordered then solids. Chapter Seven

25. Entropy In 7.4 your book talks about melting ice. Taking ice from the freezer and letting it melt on the table is an endothermic process—so in your head, you might think that this is “Not Favorable,” in the sense that burning a match is heat “favorable.” However, the ice does melt. It melts because liquid molecules are more disordered then solid molecules. At room temperature, the ice cube will not only melt but Evaporate into a gas…why? Going from order to disorder is a powerful driving factor in a chemical reaction. Chapter Seven

27. Entropy It goes from 2 moles of a soild to 22 moles of a gas. Prespective: it starts out as a cup of solid, and ends up as 2083 cups of gas Is this an entropy favorable, S = +, reaction? Dynamite, better known as TNT 2 C7H5N3O6 (s)  3 N2 (g) + 5 H2O (g) + 7 CO (g) + 7 C (g) Chapter Seven

28. When potassium is added to water contained in a beaker, the reaction shown below occurs, and the beaker feels hot to the touch. K(s) + 2 H2O(l)  2 KOH(aq) + H2(g)During this reaction • entropy decreases and S = – • entropy decreases and S = + • entropy increases and S = – • entropy increases and S = +

29. When potassium is added to water contained in a beaker, the reaction shown below occurs, and the beaker feels hot to the touch. K(s) + 2 H2O(l)  2 KOH(aq) + H2(g)During this reaction entropy decreases and S = – entropy decreases and S = + entropy increases and S = – entropy increases and S = +

30. 2. What is “Free Energy,” and what is the criterion for spontaneity? Define enthalpy, entropy, and free energy change and explain how these values affect a chemical reaction Both Heat and Entropy affect the Total Energy of a chemical reaction. Chapter Seven

31. Free energy change (DG):Free energy change is used to describe spontaneity of a process. It takes both DH and DS into account. • Exergonic: A spontaneous reaction or process that releases free energy and has a negative G. • Endergonic: A nonspontaneous reaction or process that absorbs free energy and has a positive G. Chapter Seven

32. DG = DH - TDS Chapter Seven

33. When potassium is added to water contained in a beaker, the reaction shown below occurs, and the beaker feels hot to the touch. K(s) + 2 H2O(l)  2 KOH(aq) + H2(g)This reaction is • nonspontaneous and G = – • nonspontaneous and G = + • spontaneous and G = – • spontaneous and G = +

34. When potassium is added to water contained in a beaker, the reaction shown below occurs, and the beaker feels hot to the touch. K(s) + 2 H2O(l)  2 KOH(aq) + H2(g)This reaction is • nonspontaneous and G = – • nonspontaneous and G = + • spontaneous and G = – form heat a gas and ions • spontaneous and G = +

36. When is it Spontanious? 7.40. For the reaction: H2 + Br2 2HBr ∆H = -17.4 kcal/Kmol, ∆S = 27.2 cal/Kmol Is this reaction spontaneous at all temperatures? No At what temperature is the reaction spontaneous? When ∆G = 0, anything above 0 is spontaneous, so solve for T; ∆G = ∆H – T ∆S, 0 = -17.4 kcal/Kmol– T(.0272kcal/Kmol) = 640. K or 367ºC Chapter Seven

37. 3. What determines the rate of a chemical reaction? Explain activation energy, and the other factor that determines the rate. Chapter Seven

38. 7.5 How Do Chemical Reactions Occur? Reaction Rates • The value of DG indicates whether a reaction will occur but it does not say anything about how fast the reaction will occur or about the details of the molecular changes that takes place. • For a chemical reaction to occur, reactant particles must collide, some chemical bonds have to break, and new bonds have to form. Not all collisions lead to products, however. Chapter Seven

39. One requirement for a productive collision is that the colliding molecules must approach with the correct orientation so that the atoms about to form new bonds can connect. Chapter Seven

40. Another requirement for a reaction to occur is that the collision must take place with enough energy to break the appropriate bonds in the reactant. If the reactant particles are moving slowly the particles will simply bounce apart. Chapter Seven

41. Activation energy (Ea): The amount of energy the colliding particles must have for productive collisions to occur. The size of the activation energy determines the reaction rate, or how fast the reaction occurs. • The lower the activation energy, the greater the number of productive collisions in a given amount of time, and faster the reaction. • The higher the activation energy, the lower the number of productive collisions, and slower the reaction. Chapter Seven

42. Shown is an energy diagram for a reaction with a small activation energy and products having less energy than reactants. This reaction is • nonspontaneous and fast. • nonspontaneous and slow. • spontaneous and fast. • spontaneous and slow.

43. Shown is an energy diagram for a reaction with a small activation energy and products having less energy than reactants. This reaction is • nonspontaneous and fast. • nonspontaneous and slow. • spontaneous and fast. • spontaneous and slow.

44. Draw a reaction diagram for: -50 kcal/mol 25 kcal/mol A reaction with a ∆G of -50 kcal/Kmol and Ea of 25kcal/mol. Label axis, and if the reaction is endergonic or exergonic.

45. 7.6 Effects of Temperature, Concentration, and Catalysts on Reaction Rates Reaction rates increase with temperature. With more energy the reactants move faster. The frequency of collisions and the force with which collisions occur both increase. As a rule of thumb, a 10°C rise in temperature causes a reaction rate to double. Chapter Seven

46. A second way to speed up a reaction is to increase the concentrations of the reactants. • With reactants crowded together, collisions become more frequent and reactions more likely. Flammable materials burn more rapidly in pure oxygen than in air because the concentration of molecules is higher (air is approximately 21% oxygen). • Hospitals must therefore take extraordinary precautions to ensure that no flames are used near patients receiving oxygen. Chapter Seven

47. A third way to speed up a reaction is to add a catalyst—a substance that accelerates a chemical reaction but is itself unchanged in the process. • A catalyzed reaction has a lower activation energy. Chapter Seven

49. The thousands of biochemical reactions continually taking place in our bodies are catalyzed by large protein molecules called enzymes, which promote reaction by controlling the orientation of the reacting molecules. Since almost every reaction is catalyzed by its own specific enzyme, the study of enzyme structure, activity, and control is a central part of biochemistry. Chapter Seven

50. When a catalyst is added to a reaction to increase its rate of reaction the • activation energy is lowered and G becomes more negative. • activation energy is lowered and G remains unchanged. • activation energy is raised and G becomes more positive. • activation energy is raised and G remains unchanged.