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Unit 6 – Kinetics

Unit 6 – Kinetics. AP CHemistry. Collision Theory. Key Ideas of Collision Theory: For a chemical reaction to occur, the reacting particles must collide. Not all collisions are successful (successful = collision results in a reaction). What determines if a collision is successful ?.

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Unit 6 – Kinetics

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  1. Unit 6 – Kinetics AP CHemistry

  2. Collision Theory Key Ideas of Collision Theory: • For a chemical reaction to occur, the reacting particles must collide. • Not all collisions are successful (successful = collision results in a reaction)

  3. What determines if a collision is successful? • Particles need to have proper orientation (direction in which they collide)

  4. What determines if a collision is successful? • Particles need to have enough energy, or speed • ActivationEnergy (EA) – Particles need to have a minimum amount of energy when they collide to have a successful collision • EA needed to break existing bonds and form new bonds

  5. Potential energy diagrams

  6. Potential energy diagrams • At all times, the matter has an amount of potential energy associated with it.

  7. Potential energy diagrams • Diagram shows: ΔH, EA

  8. Nature of the reactants –Some reactant molecules react in a hurry, others react very slowly. • Physical state- • gasoline (l) vs. gasoline (g) ; • K2SO4(s) + Ba(NO3)2(s) → no rxn; while both of these in the aqueous state will react. Heterogeneous reactants Homogeneous reactants

  9. Factors that Affect Reaction rates • Concentration of reactants –more molecules, more collisions. • Temperature – heat >em up & speed >em up; the faster they move, the more likely they are to collide. • An increase in temperature produces more successful collisions that are able to overcome the needed activation energy, therefore, a general increase in reaction rate with increasing temperature.

  10. Factors that Affect Reaction rates Catalysts – accelerate chemical reactions but are not themselves transformed. • Biological catalysts are proteins called enzymes. • A catalyst is a substance that changes the rate of reaction by altering the reaction pathway. Most catalysts work by lowering the activation energy needed for the reaction to proceed, therefore, more collisions are successful and the reaction rate is increased. • Remember! The catalyst is not part of the chemical reaction and is not used up during the reaction. • (May be homogeneous or heterogeneous catalysts.)

  11. Note: A catalyst lowers the activation energy barrier. Therefore, the forward and reverse reactions are both accelerated to the same degree.

  12. Factors that Affect Reaction rates Surface area of reactants – exposed surfaces affect speed. • Except for substances in the gaseous state or solution, reactions occur at the boundary, or interface, between two phases. • The greater surface area exposed, the greater chance of collisions between particles, hence, the reaction should proceed at a much faster rate. Ex. coal dust is very explosive as opposed to a piece of charcoal. Solutions are ultimate exposure! Adding an inert gas has NO EFFECT on the rate [or equilibrium] of the reaction.

  13. Chemical reaction rates • The speed of a reaction is expressed in terms of its “rate” = some measurable quantity is changing with time • The rate of a chemical reaction is measured by the decrease in concentration of a reactant or an increase in concentration of a product in a unit of time: • When writing rate expressions, they can be written in terms of reactant disappearance or product appearance

  14. aA + bB cC + dD • Generalized rate of reaction: *The change in rate of a reactant is negative (because it’s disappearing) *The change in rate of a product is positive (because it’s appearing)

  15. Average Rate of the Reaction – calculate using any time interval (Δt) • You can choose your time interval range, but the longer the interval, the slower the rate. • This is because the reaction slows down as it proceeds.

  16. What is the rate of the reaction with respect to A: • 0 – 10 seconds? • 1 – 20 seconds? • 0 – 30 seconds?

  17. Consider this balanced chemical equation: H2O2 (aq) + 3I- (aq) + 2H+ (aq)  I3- (aq) + 2H2O (l) In the first 10.0 seconds of the reaction, the concentration of I- dropped from 1.000 M to 0.868 M. a) Calculate the average rate of this reaction in the time interval.

  18. Consider this balanced chemical equation: H2O2 (aq) + 3I- (aq) + 2H+ (aq)  I3- (aq) + 2H2O (l) In the first 10.0 seconds of the reaction, the concentration of I- dropped from 1.000 M to 0.868 M. b) Determine the rate of change in the concentration of H+ (that is, Δ[H+]/Δt). )

  19. Consider this balanced chemical equation: H2O2 (aq) + 3I- (aq) + 2H+ (aq)  I3- (aq) + 2H2O (l) • For the reaction shown above (example 1), predict the rate of change in concentration of • a) H2O2(Δ[H2O2]/Δt) • b) I3- (Δ[I3-]/Δt)

  20. Group Practice: Expressing Reaction Rates In groups, • express the rate of reaction in terms of the change in concentration for each of the reactants and products, • When the first reactant is decreasing at a rate of 0.100 M/s, how fast are the other reactants decreasing? How fast are the products increasing?

  21. Group Practice: Expressing Reaction Rates • 2 NO2(g) + F2(g)  2 NO2F(g) • CH3Cl(g) + 2Cl2(g)  CCl4(g) + 3HCl(g) • 2H2O2(aq)  2H2O(l) + O2(g) • NO2(g)  NO(g) + ½ O2(g) • Cl2(g) + 3F2(g)  2ClF3(g)

  22. Warm-up #1 • For the reaction A + 2B  C under a given set of conditions, the initial rate is 0.100 M/s. What is Δ[B]/Δt under the same conditions? • -0.0500 M/s • -1.000 M/s • -0.200 M/s

  23. Warm-up #2 • Dinitrogen monoxide decomposes into nitrogen and oxygen when heated. The initial rate of the reaction is 0.022 M/s. What is the initial rate of change of the concentration of N2O (that is, Δ[N2O]/Δt)? 2 N2O (g)  2 N2 (g) + O2 (g) • -0.022 M/s • -0.011 M/s • -0.044 M/s • +0.022 M/s

  24. Warm-up #3 A burning splint will burn more vigorously in pure oxygen than in air because • oxygen is a reactant in combustion and concentration of oxygen is higher in pure oxygen than it is in air. • oxygen is a catalyst for combustion. • oxygen is a product of combustion. • nitrogen is a product of combustion and the system reaches equilibrium at a lower temperature. • nitrogen is a reactant in combustion and its low concentration in pure oxygen catalyzes the combustion.

  25. Warm-up #4 As the temperature of a reaction is increased, the rate of the reaction increases because the • reactant molecules collide less frequently • reactant molecules collide more frequently and with greater energy per collision • activation energy is lowered • reactant molecules collide less frequently and with greater energy per collision • reactant molecules collide more frequently with less energy per collision

  26. Warm-up #5 5. Which one of the following is not a valid expression for the rate of the reaction below? 4NH3 + 7O2 4NO2 + 6H2O • All of the above are valid expressions of the reaction rate.

  27. The Rate law • The rate of a reaction often depends on the concentration of one or more of the reactants. A  Products • We can express the relationship between the rate of the reaction and the concentration of the react: k = rate constant (“proportionality constant”) n = reaction order

  28. Reaction order • Reaction order (n) – determines how the rate depends on the concentration

  29. Reactant concentration as a function of time for different reaction orders

  30. Determining the order of a reaction • The order of a reaction can be determined ONLY by experiment! • Method of initial rates – initial rates measured for different initial reactant concentrations • Rate measured for a short period of time at beginning of the experiment • Use to determine the effect of concentration for each reactant on the rate.

  31. A  Products • In an experiment, the initial rate is measured at several different initial concentrations with the following results: Doubling [A] resulted in a doubled rate. Quadrupling [A] resulted in a quadrupled rate(x4). ***Rate is directly proportional to initial concentration. Therefore, “first order in A”

  32. Solving for k • We can determine the value of the rate constant, k, by solving the rate law for k and substituting the concentration and the initial rate from any one of the three measurements:

  33. Zero order (n=0) • Zero order reaction, the initial rate is independent of the reactant concentration – the rate is the same at all measured initial concentrations!

  34. Second order (n=2) • Second order reaction, the initial rate quadruples for a doubling of the reactant concentration – the relationship is quadradic!

  35. Determine Rate Order • When it’s not as obvious to see how the rate changes, you can substitute any two initial concentrations and the corresponding initial rates into a ratio of the rate laws:

  36. Another way to Determine Rate Order • Because k is in both rate laws, and it is the same value for k (it’s the same reaction just different concentrations), simplify the ratio of rate laws:

  37. Rate Constants • Rate Constant (k) – depends on the reaction • Experimentally determined • A constant value that lets us mathematically compare RATE to INITIAL CONCENTRATION • Units are different for each reaction order: • Zero order, k = M/s (M· s-1) • First order, k = 1/s (s-1) • Second order, k = M-1· s-1 • If you forget which is which, just look at the units required in the rate law to so that Rate = M/s Second order: First order: Zero order:

  38. Rate Constants • Units are different for each reaction order: • Zero order, k = M/s (M· s-1) • First order, k = 1/s (s-1) • Second order, k = M-1· s-1 • If you forget which is which, just look at the units required in the rate law to so that Rate = M/s Second order: Zero order: First order:

  39. This reaction is first order in N2O5: N2O5(g)  NO3(g) + NO2(g) • The rate constant for the reaction at a certain temperature is 0.053/s. • Calculate the rate for the reaction when [N2O5] = 0.055 M. • What would the rate of the reaction be at the concentration indicated in part a if the reaction were second order? Zero order?

  40. Consider the data showing the initial rate of a reaction (A  Products) at different concentrations of A. • What is the order of the reaction? • Write a rate law for the reaction including the value of the rate constant, k.

  41. Consider the data showing the initial rate of a reaction (A  Products) at different concentrations of A. • What is the order of the reaction? • Write a rate law for the reaction including the value of the rate constant, k. 2nd order

  42. Reaction order for multiple reactants aA + bB cC + dD • For the generic reaction: • Each reactant has its own reaction order. • The overall order is the sum of the exponents: overall order = m + n

  43. Determine the order experimentally! Use method of initial rates (like before). • When looking at reaction order for A, choose two experiments where [A] changes but all other concentrations stay constant. Do likewise for B (choose where [B] changes, holding [A] constant). When looking for reaction order of NO2…

  44. Determine the order experimentally! Use method of initial rates (like before). • When looking at reaction order for A, choose two experiments where [A] changes but all other concentrations stay constant. Do likewise for B (choose where [B] changes, holding [A] constant). When looking for reaction order of CO…

  45. Consider the reaction between nitrogen dioxide and carbon monoxide: NO2(g) + CO(g) --> NO(g) + CO2(g) • (a) the rate law for the reaction • (b) the rate constant (k) for the reaction.

  46. Rate = k[A][B]2 A reaction in which A, B, and C reacto to form products is first order in A, second order in B, and zero order in C. • Write a rate law for the reaction. • What is the overall order of the reaction? By what factor does the reaction rate change... • if [A] is doubled? • if [B] is doubled? • if [C] is doubled? • By what factor does the reaction rate change if the concentration of all three reactants are doubled? 3rd order 2 (doubled) 4 (quadrupled) 1 (no change) 8 (doubled, then quadrupled)

  47. Warm-up #1 For the following reaction: NO2(g) + CO(g) → NO(g) + CO2(g), the rate law is: Rate = k[NO2]2. If a small amount of gaseous carbon monoxide (CO) is added to a reaction mixture that was 0.10 molar in NO2 and 0.20 molar in CO, which of the following statements is true? • Both k and the reaction rate remain the same. • Both k and the reaction rate increase. • Both k and the reaction rate decrease. • Only k increases, the reaction rate remains the same. • Only the reaction rate increases; k remains the same.

  48. Warm-up #2 Changes in which of the factors affect both rate and rate constant? I- temperature II- concentration • I only • II only • Both I and II • Neither I or II

  49. Warm-up #3 A reaction was found to be second order in carbon monoxide concentration. The rate of the reaction __________ if the [CO] is doubled, with everything else kept the same. • doubles • remains unchanged • triples • increases by a factor of 4 • is reduced by a factor of 2

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