Chemical Reactions: Energy, Rates, and Equilibrium

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. 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

3. 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

4. 7.2 Heat Changes During Chemical Reactions • 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 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

5. Enthalpy of Reaction DH DH = Energy of bonds formed – Energy of bonds broken Chapter Seven

6. 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. • Law of conservation of energy: Energy can be neither created nor destroyed in any physical or chemical change. • Heat of reaction: Represented by DH, is the difference between the energy absorbed in breaking bonds and that released in forming bonds. DH is also known as enthalpy change. Chapter Seven

7. 7.3 Exothermic and Endothermic Reactions 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

8. 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

9. Energy from foodkcal = Calorie Fat 9 kcal/g Protein 4 kcal/g Carbohydrate 4 kcal/g Ethanol 7.1 kcal/g Hamburger patty 245 kcal Pizza slice 290 kcal Cup of ice cream 270 kcal Chapter Seven

10. Nutrition label

11. 7.4 Why Do Chemical Reactions Occur? Free Energy • Spontaneous process: A process that, once started, proceeds without any external influence. • 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. Chapter Seven

12. Entropy – a measure of disorder Chapter Seven

13. 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

14. DG = DH - TDS Chapter Seven

15. 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

16. 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

17. 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. 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. Chapter Seven

18. 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

19. Problem 7.6 Melting of solid ice to liquid water has: DH = 1.44 kcal/mol DS = +5.26 cal/mol.K What is DG at -10oC, 0oC and 10oC? DG = DH - TDS Chapter Seven

20. At -10oC DG = DH – TDS DH = 1.44 kcal/mol T = -10 +273 = 263K DS = 5.26 cal/mol.K x 1 kcal/1000 cal = 0.00526 kcal/mol.K DG = 1.44kcal/mol – [(263 K)(0.00526 kcal/mol.K)] = + 0.06 kcal/mol Not spontaneous Chapter Seven

21. At 0oC DG = DH – TDS DG = 1.44kcal/mol – [(273 K)(0.00526 kcal/mol.K)] = -0.09 kcal/mol spontaneous Chapter Seven

22. 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

23. 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

24. 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 catalyst lowers the activation energy. Chapter Seven

25. 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

26. 7.7 Reversible Reactions and Chemical Equilibrium Imagine the situation if you mix acetic acid and ethyl alcohol. The two begin to form ethyl acetate and water. But as soon as ethyl acetate and water form, they begin to go back to acetic acid and ethyl alcohol. Such a reaction, which easily goes in either direction, is said to be reversible and is indicated by a double arrow in equations. Chapter Seven

27. Both reactions occur until the concentrations of reactants and products reach constant values. The reaction vessel contains both reactants and products and is said to be in a state of chemical equilibrium. A state in which the rates of forward and reverse reactions are the same. Chapter Seven

28. 7.8 Equilibrium Equations and Equilibrium Constants • Consider the following general equilibrium reaction: aA + bB + …  mM + nN + … • Where A, B, … are the reactants; M, N, …. Are the products; a, b, ….m, n, …. are coefficients in the balanced equation. At equilibrium, the composition of the reaction mixture obeys an equilibrium equation. Chapter Seven

29. The equilibrium constant K is the number obtained by multiplying the equilibrium concentrations of the products and dividing by the equilibrium concentrations of the reactants, with the concentration of each substance raised to a power equal to its coefficient in the balanced equation. • The value of K varies with temperature. Chapter Seven

30. K larger than 1000: Reaction goes essentially to completion. • K between 1 and 1000: More products than reactants are present at equilibrium. • K between 1 and 0.001: More reactants than products are present at equilibrium. • K smaller than 0.001: Essentially no reaction occurs. Chapter Seven

31. Equilibrium Equation Keq Chapter Seven

32. Write the equilibrium equation Chapter Seven

33. Problem 7.57 Chapter Seven

34. Problem 7.59 Having solved for Keq, Determine the [O2] at equilibrium when [OF2] = 0.18 mol/L [F2] = 0.0200 mol/L Chapter Seven

35. Solve for [O2] Chapter Seven

36. Problem Solved! Chapter Seven

37. 7.9 Le Châtelier's Principle: The Effect of Changing Conditions on Equilibria • Le Châtelier's Principle: When a stress is applied to a system at equilibrium, the equilibrium shifts to relieve the stress. • The stress can be any change in concentration, pressure, volume, or temperature that disturbs original equilibrium. Chapter Seven

38. What happens if the concentration of CO is increased? • To relieve the “stress” of added CO, according to Le Châtelier’s principle, the extra CO must be used up. In other words, the rate of the forward reaction must increase to consume CO. • Think of the CO added on the left as “pushing” the equilibrium to the right: Chapter Seven

39. The forward and reverse reaction rates adjust until they are again equal and equilibrium is reestablished. • At this new equilibrium state, the value of [H2] will be lower, because more has reacted with the added CO, and the value of [CH3OH] will be higher. • The changes offset each other, however, so the value of the equilibrium constant K remains constant. Chapter Seven

40. Le Châtelier’s principle predicts that an increase in temperature will cause an equilibrium to shift in favor of the endothermic reaction so the additional heat is absorbed. • You can think of heat as a reactant or product whose increase or decrease stresses an equilibrium just as a change in reactant or product concentration does. Chapter Seven

41. Pressure influences an equilibrium only if one or more of the substances involved is a gas. As predicted by Le Châtelier’s principle, increasing the pressure shifts the equilibrium in the direction that decreases the number of molecules in the gas phase and thus decreases the pressure. • For the ammonia synthesis, increasing the pressure favors the forward reaction because 4 moles of gas is converted to 2 moles of gas. Chapter Seven

42. The effects of changing reaction conditions on equilibria are summarized below. Chapter Seven

43. Chapter Summary • The strength of a covalent bond is measured by its bond dissociation energy. • If heat is released, H is negative and the reaction is said to be exothermic. If heat is absorbed, H is positive and the reaction is said to be endothermic. • Spontaneous reactionsare those that, once started, continue without external influence; nonspontaneous reactions require a continuous external influence. • Spontaneity depends on two factors, the amount of heat absorbed or released in a reaction and the entropy change. Chapter Seven

44. Chapter Summary Contd. • Spontaneous reactions are favored by a release of heat, H <0, and an increase in entropy, S >0. • The free-energy change, G = H - T S, takes both factors into account. • G<0 indicates spontaneity, G>0 indicates nonspontaneity. • Chemical reactions occur when reactant particles collide with proper orientation and energy. The exact amount of collision energy necessary is the activation energy. • Reaction rates can be increased by raising the temperature, by raising the concentrations of reactants, or by adding a catalyst. Chapter Seven

45. Chapter Summary Contd. • At equilibrium, the forward and reverse reactions occur at the same rate, and the concentrations of reactants and products are constant. Every reversible reaction has an equilibrium constant, K. The forward reaction is favored if K>1; the reverse reaction is favored if K<1. • Le Châtelier’s principlestates that when a stress is applied to a system in equilibrium, the equilibrium shifts so that the stress is relieved. • Applying this principle allows prediction of the effects of changes in temperature, pressure, and concentration. Chapter Seven