Chemical Kinetics

1. Chemical Kinetics Reaction Rates and Equilibrium

2. Objectives • Know and understand collision theory. • Be able to explain the role of an activated complex in a chemical reaction. • Be able to draw a potential energy diagram that includes activation energy.

3. Collision Theory • collision theory: particles must collide with enough KE to break bonds and form new bonds. • Many reactions require high temperatures! • For example, N2 and O2 only react to make NO2 in hot auto engines. Not enough KE Enough KE

4. Activated Complex activated complex: a high PE, short-lived molecule that forms as reactants change to products activated complex (H2Cl2) H2 + Cl2 2 HCl

5. Activation Energy • activation energy (Ea): the additional energy needed to produce an activated complex • example: a spark is needed to get gasoline to burn DH

6. Activation Energy DH

7. Objectives • Understand the concept of a rate. • Describe the factors that affect the rate of a chemical reaction. • Understand how catalysts and inhibitors affect reaction rates. • Explain the factors that affect explosions.

8. Reaction Rates time ↓ rate ↑

9. Factors Affecting Reaction Rates temperature (T↑ R↑) concentration (C↑ R↑) particle size (size↓ R↑) catalyst: substance that reduces Ea (but not consumed) MnO2 2H2O2(l) → 2H2O(l) + O2(g) inhibitor: increases Ea and decreases reaction rate (example: preservatives)

10. The Catalytic Converter A catalytic converter reduces pollution. IN: HC (unburned fuel), CO, and NOx OUT: CO2, H2O, and N2

11. Explosives Explosions are (1) very exothermic reactions with (2) gaseousproducts that occur at a (3) high rate. 10 KNO3(s) + 3S(s) + 8C(s) black powder -DH 2K2CO3(s) + 3K2SO4 (s) + 6CO2(g)+ 5N2(g) high explosives contain nitrogen and oxygen TNT nitroglycerin

12. Zinc Pyrotechnics! • ingredients: NH4NO3(s), Zn(s), NH4Cl(s), and H2O(l) • water added to the zinc produces heat to initiate • the reaction water dissolves ammonium chloride (makes Cl-)… NH4NO3(s) → N2O(g) + H2O(g) + 23 kJ heat melts ammonium nitrate so 2nd reaction can occur… Zn(s) + NH4NO3(l ) → N2(g) + ZnO(s) + 2H2O(g) + 466 kJ Cl-

13. Reversible Reactions

14. Objectives • Understand the concept of a reversible chemical reaction. • Understand the concept of chemical equilibrium and how systems respond to stresses to change the equilibrium position. • Be able to apply LéChâtelier’s principle to determine the effect that changes in concentration, temperature, or pressure have on the equilibrium position.

15. Reversible Reactions • Many reactions can go forward or backward depending on conditions. • forward: 2NO2(g) → N2O4(g) (as T ↓) • reverse: 2NO2(g) ← N2O4(g) (as T↑) 2NO2(g) ↔ N2O4(g) red (warm) clear (cool)

16. Dynamic Equilibrium dynamic equilibrium: forward rate = reverse rate, but relative amounts will differ!

17. Equilibrium Position • equilibrium position: tells which side is favored • reactants ↔ products “products favored” < 50% > 50% • reactants ↔ products “reactants favored” > 50% < 50%

18. LéChâtelier’s Principle LéChâtelier’s Principle: if a “stress” is applied to a system in dynamic equilibrium, the system will respond to reduce the stress

19. LéChâtelier and Concentration • As a substance is introduced to an chemical equilibrium, the reaction is favored away from that substance. • If a substance is removed, the reaction adjusts to replace it. D Add B B C A D A A + B → C + D, products favored A C B B B A D B A Remove A C A + B ← C + D, reactants favored A + B ↔ C + D

20. LéChâtelier and Temperature Just treat energy like a substance! 2NO2(g) ↔ N2O4(g) + energy red clear higher T (add energy) ← sky turns redder lower T (remove energy) → sky appears clearer

21. LéChâtelier and Pressure • Under high pressure (less volume), the side with fewer gas particles (or moles) is favored—reducing the stress. 2 NO2(g) ↔ 1 N2O4(g) high pressure low pressure 2NO2(g) ← N2O4(g) 2NO2(g) →N2O4(g) reactant favored product favored

22. Applying LéChâtelier • N2(g) + 3H2(g) ↔ 2NH3(g) + energy • Add hydrogen gas? products favored • Remove nitrogen? reactants favored • Increase temperature? reactants favored • Increase pressure? products favored

23. Common Ion Effect • The addition of a “common ion” can affect an equilibrium… • PbCrO4(s) ↔ Pb2+ (aq) + CrO42-(aq) • adding PbCl2(aq), which contains Pb2+ , will favor the production of PbCr4(s) • adding K2CrO4 also produces more reactant

24. Objectives • Understand the concept of an equilibrium constant (Keq) and what it reveals about the equilibrium position of a system in chemical equilibrium. • Be able to calculate an equilibrium constant.

25. Equilibrium Constants (Keq) [A] is conc. of A, etc. aA + bB ↔ cC + dD “Equilibrium Position”… If Keq > 1, products favored If Keq < 1, reactants favored # moles

26. Equilibrium Constants • For the chemical equilibrium H2(g) + I2(g) ↔ 2HI(g) at 25oC, the following concentrations are experimentally determined to be [HI] = 0.00998 M [H2] = 0.000867 M [I2] = 0.00264 M What is the equilibrium constant at 25oC?

27. Equilibrium Constants The following chemical reaction has a Keq = 5.63 x 1018 (at 25oC): H2(g) + Br2(g) ↔ 2HBr(g). If the concentration of each reactant is 2.11 x 10-10 M at 25oC, what is the concentration of HBr?

28. Solubility Constants • Insoluble compounds actually dissolve very slightly. • Ksp indicates the solubility of a saturated compound (usually at 25oC) • lower Ksp = lower solubility Ca(OH)2 ↔ Ca2+(aq) + 2OH-(aq) See page 562 Ca(OH)2  = 6.5 x 10-6 ZnS  = 3.0 x 10-23 Ksp = [Ca2+][OH-]2 = 6.5 x 10-6 ZnS ↔ Zn2+(aq) + S2-(aq) Ksp = [Zn2+][Cl2-] = 3.0 x 10-23

29. Objectives • Understand the concept of spontaneous and non-spontaneous processes. • Understand the concept of entropy and how it affects the spontaneity of a process. • Be able to predict the spontaneity of a process by considering the influence of enthalpy and entropy.

30. Spontaneous Processes • spontaneous: occurs naturally, favors products C3H8 + 5O2 ↔ 3CO2 + 4H2O + energy → spontaneous ← non-spontaneous, why? • exothermic: tend to be spontaneous endothermic: tend to be non-spontaneous • But endothermic processes do occur spontaneously: H2O(s) + energy → H2O(l) Why? Processes that increase ENTROPY tend to be spontaneous.

31. Entropy Does your room look like this? entropy: a measure of disorder low entropy high entropy It takes a lot of effort and energy to reduce entropy! Systems tend to become more disordered (increased entropy) due to probability.

32. Entropy tend to be non-spontaneous tend to be spontaneous LOWER ENTROPY HIGHER ENTROPY solid ↔ liquid or (aq) ↔ gas fewer moles ↔ more moles

33. Predicting Reaction Outcomes CH3CH2OH(l) + 2O2(g) → 2CO2(g) + 3H2O(g) + energy exothermic = spontaneous entropy? (liquid → gas, more moles) = spontaneous spontaneous reaction! 2NaCl(s) + energy → 2Na(s) + Cl2(g) endothermic = non-spontaneous entropy? (solid → gas, more moles) = spontaneous spontaneous or non-spontaneous?

34. Objectives • Understand the concept of free energy. • Know the factors that affect the amount of Gibbs free energy that is either absorbed or released in a chemical reaction. • Be able to calculate the amount of Gibbs free energy involved in a chemical reaction (and determine the spontaneity of the reaction from this value).

35. Gibbs Free Energy free energy: energy that can do “work” (move things) • It is released when a (s) or (l)→ (g); expanding gas can do work. Exothermic reactions can do work too. • Processes that release Gibbs free energy (G0) are spontaneous. Enthalpy, entropy, and temperature must all be considered. C2H5OH(l) + 2O2(g) → 2CO2(g) + 3H2O(g) + energy - DG0 = spontaneous reaction + DG0 = non-spontaneous DG0 = DH0 – TDS0 enthalpy entropy Gibbs free energy temperature

36. Calculating Changes in Gibbs Free Energy DG0 = SDGf0 (products) - SDGf0 (reactants) Is the following reaction spontaneous at 25oC? 2NaCl(s) →2Na(s) + Cl2(g) Is the following reaction spontaneous at 25oC? 2AgNO3(aq) + CaCl2(aq) → 2AgCl(s) + Ca(NO3)2(aq) Hint: write the net ionic equation first!