Kinetics and Equilibrium

1. Kinetics and Equilibrium

2. Kinetics • Kinetics - concerned with the rates of chemical reactions. • Basic concept; in order for a reaction to occur, reactant particles must collide • Collision Theory– collisions between particles can produce a reaction if both the spatial orientation and energy of the colliding particles are conducive to a reaction.

3. Collision Theory • An effective collision of oxygen and hydrogen molecules produces water molecules.

4. Collision Theory • An ineffective collision of oxygen and hydrogen molecules produces no reaction; the reactants bounce apart unchanged.

5. Collision Theory Particles must collide in a specific arrangement for the collision to be effective, allowing products to form.

6. Collision Theory • Properly positioned particles during a collision allows the activated complex to form • Activated complex – temporary, intermediate product. • Allows for the rearrangement of atoms to form new products • OR the reformation of reactants

7. Collision Theory • The minimum energy that colliding particles must have in order to react is called the activation energy. • Think of it as a barrier that reactants must cross before products can form. • When two reactant particles collide, they may form an activated complex. • An activated complexis an unstable arrangement of atoms that forms for a moment at the peak of the activation-energy barrier.

8. Collision Theory • When two reactant particles collide, they may form an activated complex. • The activated complex forms only if the colliding particles have enough energy and if the atoms are oriented properly. • The lifetime of an activated complex is typically about 10-13 seconds. • Its brief existence ends with the reformation of the reactants or with the formation of products. • Thus, the activated complex is sometimes called the transition state.

9. Rates of Reaction Potential Energy Diagrams • The activation-energy barrier must be crossed before reactants are converted to products. • Note: Energy is measured in kJ (kilojoules)

10. Potential Energy Diagram • Recall: chemical bonds contain energy  specifically that energy is potential energy. • These diagrams illustrate the potential energy change that occurs during a chemical reaction. • Y-axis: change in potential energy • X-axis: reaction coordinate – the progress of the reaction • In order for a reaction to occur: • reactants must have sufficient potential energy to collide effectively

11. The figure below illustrates the progress of a typical reaction. Over time, the amount of reactant decreases and the amount of product increases.

12. Collision / Reaction Process 1. • Energy is applied to reactants. Reactant particles absorb energy; kinetic energy. • Remember that reactants are more unstable then their products. • Effective collisions allow kinetic energy to be converted to potential 2. 3. 1.

13. Potential Energy (P.E.) Diagram E F • P.E. of reactants • Activation Energy (A.E.) • Heat of Rx./ Difference in P.E. b/w reactants and products • P.E. of products • A.E. of forwardRx. /Difference b/w the reactants and the activated complex • A.E. of reverseRx./ Difference b/w the products and the activated complex

14. P.E. of activated complex P.E. of Products P.E. of Reactants

15. Rates of Reaction Potential Energy Diagrams • Remember: An endothermic reaction absorbs heat, and an exothermic reaction releases heat. Endothermic or Exothermic

16. Exothermic • Products have less potential energy then reactants • Energy is released, and the temperature of the surrounding increases • Heat of the Rx is always negative

17. Exothermic • Reactant has more P.E. then products. • Net loss of energy. • More energy was released to form the product than was gained to form the activated complex. • ΔH = negative

18. Endothermic • Products have more potential energy then the reactants • Energy is absorbed, temperature of the surrounding decreases • Heat of reaction is always positive

19. Endothermic • Product has more P.E. then reactants. • Net gain of energy. • More energy was absorbed to form the activated complex than was released to form the product. • ΔH = positive

20. Topic 9 - Kinetics & Equilibrium Note Packet • Complete Practice Problems #9-14 on page 161. • Complete Practice Problems #23-25 on page 163. • Complete Practice Problems #26-32 on page 165.

21. Factors Affecting Rates of Rx • Nature of the reactants • Concentration • Surface area/ Particle size • Pressure – gasses only • Presence of a catalyst • Temperature

22. Factors Affecting Rates of Rx • Nature of the Reactants • Reactions involve the breaking of existing bonds and the formation of new ones. • Ionic bonds – faster to react • low bond dissociation energy/ easy to break • Covalent bonds – slower to react • high bond dissociation energy/ hard to break • Greater number of bonds must be broken before the reaction can occur. • Breaking more bonds requires that the particles must have more energy when they collide

23. Factors Affecting Rates of Rx • Concentration • If one or more of the reactants is increased, the reaction will proceed faster.

24. Factors Affecting Rates of Rx • Surface Area/ Particle Size • More surface area exposed, the greater chance for reactant particles to collide. Only atoms at the surface of the metal are available for reaction. Dividing the metal into smaller pieces increases the surface area and the number of collisions.

25. Factors Affecting Rates of Rx • Pressure • An increase in pressure on a gas, increases the concentration of gaseous particles, increasing the rate of reaction for gases • Pressure has little or no effect on solids and liquids

26. Factors Affecting Rates of Rx • Presence of a Catalyst • a catalyst is a substance that increases the rate of a reaction without being used up during the reaction. • Catalysts permit reactions to proceed along a lower energy path. Pt 2H2(g) + O2(g) 2H2O(l)

27. Advantage of a catalyst Heat of the reaction – difference between the P.E. of the reactants and the P.E. of the products • ΔH does not change with the presence of a catalyst. • Activation Energy of the forward reaction or the reverse reaction are both lowered.

28. Catalysts • At normal body temperature (37C), reactions in the body would be too slow without catalysts. • The catalysts that increase the rates of biological reactions are called enzymes. • When you eat a meal containing protein, enzymes in your digestive tract help break down the protein molecules in a few hours.

29. Catalysts An inhibitor is a substance that interferes with the action of a catalyst. • Some inhibitors work by reaction with, or “poisoning,” the catalyst itself. • Thus, the inhibitor reduces the amount of catalyst available for a reaction. • Reactions slow or even stop when a catalyst is poisoned.

30. Factors Affecting Rates of Rx • Temperature • Usually, raising the temperature speeds up a reaction. • Lowering the temperature usually slows down a reaction. • At higher temperatures, particles move faster. • The frequency of collisions increases. Reacting particles will have more energy, allowing the percentage of particles that have enough kinetic energy to slip over the activation-energy barrier. • Thus, an increase in temperature causes products to form faster.

31. Surprise Question Which of the following factors could be increased in order to decrease a reaction rate? A. Catalyst concentration B. Concentration C. Temperature D. Particle size

32. Topic 9 - Kinetics & Equilibrium Note Packet • Complete Practice Problems #1-8 on page 158.

33. Progress of Chemical Rx • N2O4(g) 2NO2(g) A B • The rate of the rx depends in part on the concentration of the reactants. • The rate at which A, (N2O4)forms B, (2NO2) can be expressed as a change in concentration A, () with time (t) • The initial concentration of A at t1 verses the final concentration of A at t2.

34. ΔA [A] Δt ΔA rate = = k× [A] Δt Rate Laws • The rate of disappearance of A is proportional to the concentration of A. • The proportionality can be expressed as the concentration of A, [A], multiplied by a constant, k. • Rate law, an expression for the rate of a reaction in terms of the concentration of the reactants. • Specific rate constant(k) for a reaction is a proportionality constant relating the concentrations of reactants to the rate of the reaction.

35. ΔA rate = = k× [A] Δt Rate Laws • The value of the specific rate constant, k, in a rate law is large if the products form quickly; the value is small if the products form slowly.

36. First Order Reaction • N2O4(g) 2NO2(g) A B • Conversion of A  B in one step reaction • Reaction rate is directly proportional to the concentration of only one reactant. • As the reaction progresses, the rate of the rx decreases

37. rate = k[A]a[B]b Higher Order Reactions • For the reaction of A with B, the rate of reaction is dependent on the concentrations of both A and B. • When each exponent in the rate law equals 1 (that is, x = y = 1) the reaction is said to be first order in A and first order in B. • The overall order of a reaction is the sum of the exponents for the individual reactants. i.e. second order reaction overall.

38. Equilibrium Reaction going “forward” • Current understanding: Potential energy diagrams depict rx in 1 direction; L  R in a “forward” direction. • Reactions can occur in both directions; Dynamic Equilibrium R P Reaction going “backward” R P Reaction in dynamic equilibrium R P

39. 2SO2(g) + O2(g) 2SO3(g) Equilibrium • Equilibrium– when both the forward and reverse reactions occur at the same rate • Rxs are shown proceeding in both directions with a double arrow • 2 opposite processes are occurring at the same time = equal rates. • Must occur in a closed system

40. Equilibrium • Quantities of reactants and products ARE NOT equal when at equilibrium. • Rates of reaction of reactants and products AREequal when at equilibrium. • Constant flux between reactants and products. If one side is increased or decrease, the other side will compensate for the change.

41. Physical Equilibrium • Phase Equilibrium • Ex. condensation, evaporation, and triple point • Solution Equilibrium • Ex. Solids in liq. or gases in liq. • If additional solute (solid or gas) is added to a saturated soln. in a closed system, the solute will still dissolve, but it dissolves at the same rate that it also recrystallizes back out of the solution.

42. If a system is not closed, equilibrium cannot be reached. Chemical Equilibrium Forward Reaction • Rate slows as reactants are consumed Reverse Reaction • Rate increases as products are formed Dynamic Equilibrium • Both forward and reverse reactions occur at the same rate.

43. Le Chatelier’s Principle • Used to explain how a system at equilibrium responds to relive any stress on the system. • If you had more to the left side (reactants), the lever will tip to the left. In order to restore balance, the system responds by making more of what is on the right (products) and vice versa. • Stressors include changes in: • Temperature • Concentration • Pressure Reactants Products

44. When achieving equilibrium … • Particles move from high concentration  low concentration (diffusion) • Heat travels from hot  cold (melting) • Pressure changes flow from high  low

45. Le Chatelier’s Principle Concentration Changes

46. Le Chatelier’s Principle Temperature Changes

47. Le Chatelier’s Principle Pressure Changes • Pressure changes only have effects on gases • As pressure , concentration of gases . Rx will shift to the right, causing more dissolved CO2 • As pressure , concentration of gases . Rx. Shifts to the left, forming more gaseous CO2 • Remember: Particles travel from high pressure to low pressure. • THINK SODA POP. Pressure in the bottle is greater that that outside the bottle. What happens when you open it?

49. Equilibrium Constants Think of it this way: • a mol of reactant A and b mol of reactant B react to produce c mol of product C and d mole of product D at equilibrium. Coefficients: amounts in moles of reactants & products

50. [C]cx [D]d • Keq= • [A]ax [B]b Equilibrium Constant • The equilibrium constant(Keq) is the ratio of product concentrations (mol/L) to reactant concentrations(mol/L) at equilibrium. • The value of Keq depends on the temperature of the reaction. • The size of the equilibrium constant indicates whether reactants or products are more common at equilibrium. • , products favored over equilibrium • , reactants favored over equilibrium