1. IQ #1 1)What do you think could increase the rate or speed of a chemical reaction? 2)Define activation energy. Give a real-life example. 3)Define a catalyst. How does it work?

2. Chapter 18: “Reaction Rates and Equilibrium”

3. Section 18.1 Rates of Reaction OBJECTIVES – Describe how to express the rate of a chemical reaction. – Identify four factors that influence the rate of a chemical reaction.

4. Collision Theory Reactions can occur: – Very fast – such as a firecracker – Very slow – such as the time it took for dead plants to make coal A “rate” is a measure of the speed of any change that occurs within an interval of time In chemistry, reaction rate is expressed as the amount of reactant changing per unit time

5. Collision Model Key Idea: Molecules must collide to react. However, only a small fraction of collisions produces a reaction. Why?

6. Collision Model Collisions must have enough energy to produce the reaction (must equal or exceed the activation energy – the minimum energy needed to react). Think of clay clumps thrown together gently – they don’t stick, but if thrown together, they stick together!

7. The Effect of Higher Temperatures

8. Collision Model An “activated complex” is an unstable arrangement of atoms that forms momentarily (typically about 10 -13 seconds) at the peak of the activation-energy barrier. – This is sometimes called the transition state Results in either: a) forming products b) reformation of reactants – Both outcomes are equally likely

9. - Page 543

10. Collision Model The collision theory explains why some naturally occurring reactions are very slow – Carbon and oxygen react when charcoal burns, but this has a very high activation energy – At room temperature, the collisions between carbon and oxygen are not enough to cause a reaction

11. Factors Affecting Rate 1) Temperature -Increasing temperature always increases the rate of a reaction. 2) Surface Area -Increasing surface area increases the # of collisions/rate of a reaction 3) Concentration -Increasing concentration USUALLY increases the rate of a reaction 4) Presence of Catalysts

12. Catalysts Catalyst: A substance that speeds up a reaction, without being consumed itself in the reaction Enzyme: A large molecule (usually a protein) that catalyzes biological reactions. – Human body temperature = 37 o C, much too low for digestion reactions without catalysts. Inhibitors – interfere with the action of a catalyst; reactions slow or even stop

13. Endothermic Reaction with a Catalyst

14. Catalysts Increase the Number of Effective Collisions

15. Section 18.2 Reversible Reactions and Equilibrium OBJECTIVES – Describe how the amounts of reactants and products change in a chemical system at equilibrium. – Identify three stresses that can change the equilibrium position of a chemical system. – Explain what the value of K eq indicates about the position of equilibrium.

16. Reversible Reactions Some reactions do not go to completion as we have assumed – They may be reversible – a reaction in which the conversion of reactants to products and the conversion of products to reactants occur simultaneously Forward: 2SO 2(g) + O 2(g) → 2SO 3(g) Reverse: 2SO 2(g) + O 2(g) ← 2SO 3(g)

17. Reversible Reactions The two equations can be combined into one, by using a double arrow, which tells us that it is a reversible reaction: 2SO 2(g) + O 2(g) ↔ 2SO 3(g) A chemical equilibrium occurs, and no net change occurs in the actual amounts of the components of the system.

18. IQ #2 1)List the four factors that affect reaction rates. How does each do so? 2)What is an activated complex? 3)What does it mean when we say that a reaction is “reversible”?

19. Reversible Reactions Even though the rates of the forward and reverse are equal, the concentrations of components on both sides may not be equal – An equlibrium position may be shown: A B or A B 1% 99% 99% 1%  It depends on which side is favored; almost all reactions are reversible to some extent

20. Le Chatelier’s Principle The French chemist Henri Le Chatelier (1850-1936) studied how the equilibrium position shifts as a result of changing conditions Le Chatelier’s principle: If stress is applied to a system in equilibrium, the system changes in a way that relieves the stress

21. Le Chatelier’s Principle What items did he consider to be stress on the equilibrium? 1)Concentration 2)Temperature 3)Pressure Concentration – adding more reactant produces more product, and removing the product as it forms will produce more product Each of these will now be discussed in detail

22. Le Chatelier’s Principle Temperature – increasing the temperature causes the equilibrium position to shift in the direction that absorbs heat If heat is one of the products (just like a chemical), it is part of the equilibrium so cooling an exothermic reaction will produce more product, and heating it would shift the reaction to the reactant side of the equilibrium

23. Le Chatelier’s Principle Pressure – changes in pressure will only effect gaseous equilibria Increasing the pressure will usually favor the direction that has fewer molecules N 2(g) + 3H 2(g) → 2NH 3(g) For every two molecules of ammonia made, four molecules of reactant are used up – the equilibrium shifts to the right with an increase in pressure

24. LeChatelier Example #1 A closed container of ice and water is at equilibrium. Then, the temperature is raised. Ice + Energy  Water The system temporarily shifts to the _______ to restore equilibrium. right

25. LeChatelier Example #2 A closed container of N 2 O 4 and NO 2 is at equilibrium. NO 2 is added to the container. N 2 O 4 (g) + Energy  2 NO 2 (g) The system temporarily shifts to the _______ to restore equilibrium. left

26. LeChatelier Example #3 A closed container of water and its vapor is at equilibrium. Vapor is removed from the system. water + Energy  vapor The system temporarily shifts to the _______ to restore equilibrium. right

27. LeChatelier Example #4 A closed container of N 2 O 4 and NO 2 is at equilibrium. The pressure is increased. N 2 O 4 (g) + Energy  2 NO 2 (g) The system temporarily shifts to the _______ to restore equilibrium, because there are fewer moles of gas on that side of the equation. left

28. Applications of Le Chatelier’s Principle 1. Glasses/Sunglasses Reaction: Stress: Add Sunlight Stress: Remove Sunlight AgCl + energy Ag + Cl dark clear Shifts left, glasses become clear Shifts right, glasses become dark

29. Biochemical Equilibrium System (Respiration) Reactions: STRESS: Exercise vigorously (____________________) STRESS: Breathe faster (_____________________) CO 2 + H 2 O H 2 CO 3 H 2 CO 3 H + + HCO 3 - More CO 2 produced Both reactions shift right [H + ] increases (acid), so the brain sends a message to exhale more rapidly. Exhale CO 2 Both reactions shift left pH returns to normal level

30. Equilibrium Constants Chemists generally express the position of equilibrium in terms of numerical values – These values relate to the amounts of reactants and products at equilibrium – This is called the equilibrium constant, and abbreviated K eq

31. Equilibrium Constants consider this reaction: a A + b B c C + d D – The equilibrium constant (K eq ) is the ratio of product concentration to the reactant concentration at equilibrium, with each concentration raised to a power (= the balancing coefficient)

32. Equilibrium Constants consider this reaction: aA + bB cC + dD – Thus, the “equilibrium constant expression” has the general form: [C] c x [D] d [A] a x [B] b (brackets: [ ] = molarity concentration) K eq =

33. Equilibrium Constants the equilibrium constants provide valuable information, such as whether products or reactants are favored: K eq > 1, products favored at equilibrium K eq < 1, reactants favored at equilibrium

34. - Page 557

35. Section 18.4 Entropy and Free Energy OBJECTIVES – Identify two characteristics of spontaneous reactions. – Describe the role of entropy in chemical reactions. – Identify two factors that determine the spontaneity of a reaction. – Define Gibbs free-energy change.

36. Free Energy and Spontaneous Reactions Many chemical and physical processes release energy, and that can be used to bring about other changes – The energy in a chemical reaction can be harnessed to do work, such as moving the pistons in your car’s engine Free energy is energy that is available to do work – That does not mean it can be used efficiently

37. Free Energy and Spontaneous Reactions Your car’s engine is only about 30 % efficient, and this is used to propel it – The remaining 70 % is lost as friction and waste heat No process can be made 100 % efficient – Even living things, which are among the most efficient users of free energy, are seldom more than 70 % efficient

38. Free Energy and Spontaneous Reactions We can only get energy from a reaction that actually occurs, not just theoretically: CO 2(g) → C (s) + O 2(g) – this is a balanced equation, and is the reverse of combustion – Experience tells us this does not tend to occur, but instead happens in the reverse direction

39. Free Energy and Spontaneous Reactions The world of balanced chemical equations is divided into two groups: 1)Equations representing reactions that actually occur 2)Equations representing reactions that do not tend to occur, or at least not efficiently

40. Free Energy and Spontaneous Reactions The first, (those that actually occur) and more important group involves processes that are spontaneous – A spontaneous reaction occurs naturally, and favors the formation of products at the specified conditions – They produce substantial amounts of product at equilibrium, and release free energy – Example: a fireworks display – page 567

41. Free Energy and Spontaneous Reactions In contrast, a nonspontaneous reaction is a reaction that does not favor the formation of products at the specified conditions – These do not give substantial amounts of product at equilibrium Think of soda pop bubbling the CO 2 out: this is spontaneous, whereas the CO 2 going back into solution happens very little, and is non-spontaneous

42. Spontaneous Reactions Do not confuse the words spontaneous and instantaneous. Spontaneous just simply means that it will work by itself, but does not say anything about how fast the reaction will take place – it may take 20 years to react, but it will eventually react. – Some spontaneous reactions are very slow: sugar + oxygen → carbon dioxide and water, but a bowl of sugar appears to be doing nothing (it is reacting, but would take thousands of years) – At room temperature, it is very slow; apply heat and the reaction is fast; thus changing the conditions (temp. or pressure) may determine whether or not it is spontaneous

43. Entropy Entropy is a measure of disorder, and is measured in units of J/mol. K; there are no negative values of entropy The law of disorder states the natural tendency is for systems to move to the direction of maximum disorder, not vice-versa – Your room NEVER cleans itself (disorder to order?) An increase in entropy favors the spontaneous chemical reaction A decrease in entropy favors the nonspontaneous reaction

44. - Page 570

45. Enthalpy and Entropy 1)Reactions tend to proceed in the direction that lowers the energy of the system (H, enthalpy). 2)Reactions tend to proceed in the direction that increases the disorder of the system (S, entropy). and,

46. Enthalpy and Entropy These are the two “drivers” to every equation. – If they both AGREE the reaction should be spontaneous, IT WILL be spontaneous at all temperatures, and you will not be able to stop the reaction without separating the reactants – If they both AGREE that the reaction should NOT be spontaneous, it will NOT work at ANY temperature, no matter how much you heat it, add pressure, or anything else!

47. Enthalpy and Entropy The size and direction of enthalpy and entropy changes together determine whether a reaction is spontaneous If the two drivers disagree on whether or not it should be spontaneous, a third party (Gibb’s free energy) is called in to act as the “judge” about what temperatures it will be spontaneous, and what the temp. is. – But, it WILL work and be spontaneous at some temperature!

48. Spontaneity of Reactions Reactions proceed spontaneously in the direction that lowers their Gibb’s free energy, G.  G =  H - T  S (T is kelvin temp.) If  G is negative, the reaction is spontaneous. (system loses free energy) If  G is positive, the reaction is NOT spontaneous. (requires work be expended)

49. Spontaneity of Reactions Therefore, if the enthalpy and entropy do not agree with each other as to what should happen: – Gibbs free-energy says that they are both correct, the reaction will occur – But the Gibbs free-energy will decide the conditions of temperature that it will happen Figure 18.25, page 572

50. - Page 572

51. Entropy 6. Calculation of  S rxn **Use the Tables from Thermochemistry** 7. Example calculation: Calculate the  S o rxn of the following reaction: 2 H 2 O (l) → 2 H 2(g) + O 2(g)  S o rxn =  S o Products -  S o Reactants Predict: Favorable or unfavorable  S o rxn =  S o Products -  S o Reactants  S o rxn = [2(H 2 ) + (O 2 )] - [2(H 2 O)]  S o rxn = [2 mol(130.58 J/mol·K) + (1 mol(205.0 J/mol·K)] – [2 mol(69.91 J/mol·K)]  S o rxn = + 326.3 J/K

52. C. Gibb’s Free Energy (___): Determines the ___________ of a reaction by combining the two driving forces: ________ and _______. 1. Units of Gibb’s Free Energy: ______. 2. ΔG = ___ → Reaction is ________________. 3. Δ G = ___ → Reaction is ________________. 4. Calculations of Δ G rxn Combines BOTH driving forces. G spontaneity enthalpy entropy kJ/mol + non-spontaneous - spontaneous  G =  H - T  S T = Temperature of Kelvin Standard Temp = 25 o C = 298 K

53. Example calculation: Given that the ∆H o is +571.6 KJ, calculate the ∆G o of the reaction at standard temperature. 2 H 2 O (l) → 2 H 2(g) + O 2(g)  G =  H - T  S  G = + 571.6 kJ – (298 K)(0.3263 kJ/K)  G = + 474.4 kJ  The rxn is non-spontaneous

54. + + + + + HH SS GG Condition Always Spontaneous Exothermic (favorable) FavorableSpontaneous Exothermic + or + (favorable) (unfavorable) Endothermic Favorable Unfavorable Spontaneous or Non-spon. Non- spontaneous Spontaneous or Non-spon. Always Non- spontaneous  H > T  S T  S >  H

55. 6.(a) Calculate Δ H o (exo or endo?), Δ S o (favorable or unfavorable?), and  G o (Spontaneous or non-spontaneous?) for the following reaction: Al 2 O 3(s) + 2 Fe (s) → 2 Al (s) + Fe 2 O 3(s). (b) Calculate the minimum temperature at which this reaction is spontaneous.  H o rxn = [2(Al) + (Fe 2 O 3 )] – [(Al 2 O 3 ) + 2(Fe)]  H o rxn = + 847.6 kJ Endothermic  Unfavorable  S o rxn = + 41.30 J/K  Favorable  G o =  H - T  S  H o rxn = [2 (0) + ( - 822.16)] – [( - 1669.8) + 2 (0)]  S o rxn = [2(28.32) + (89.96)] - [(51.00) + 2(27.15)]  G o = + 847.6 kJ – (298 K)(0.0413 kJ/K)  G o = + 835.3 kJ  Non-spontaneous = 20,200 K

56. 7. For the following reaction: NH 4 NO 3(s) → N 2 O (g) + 2 H 2 O (g) ∆H° f (kJ/mol) S° (J/mol K) NH 4 NO 3(s) -365.6+151.1 N 2 O (g) +82.1+219.7 H 2 O (g) -241.8+188.7 (a)Calculate ∆H° (exo or endo?), ∆S° (favorable or unfavorable?), and ∆G° (spontaneous or non-spontaneous?) @ standard temp. (b) Calculate the minimum temperature at which this reaction is spontaneous.  H o rxn = [(82.1) + 2( - 241.8)] – [( - 365.6)]  S o rxn = [(219.7) + 2(188.7)] - [(151.1)]  G o = - 35.9 kJ – (298 K)(0.446 kJ/K) = - 35.9 kJExothermic  Favorable = + 446.0 J/K  Favorable = - 169 kJ  Spontaneous Always