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Kinetics

Kinetics. Chapter 15. Introduction. Rate of Reaction – the rate at which ________ are ___________ and __________ are ______________. Chemical Kinetics – the study of the ____________ of ______________. Ways to Detect Concentration. Titrations – (Quenching) Spectrophotometer -.

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Kinetics

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  1. Kinetics Chapter 15

  2. Introduction • Rate of Reaction – the rate at which ________ are ___________ and __________ are ______________. • Chemical Kinetics – the study of the ____________ of ______________.

  3. Ways to Detect Concentration • Titrations – (Quenching) • Spectrophotometer -

  4. Factors that affect Reaction Rates • Nature of the Reactants – • Concentration – • Temperature – • Presence of a Catalyst –

  5. Rate = k[A]x[B]y • Zero order – the concentration of the reactant _________ affect the rate of production of product. The concentration of reactant is doubled, the rate of production of product __________________. • First order – the concentration of the reactant _________ affect the rate of production of product. The concentration of reactant is doubled, the rate of production of product __________________.

  6. Second order – the concentration of the reactant _________ affect the rate of production of product. The concentration of reactant is doubled, the rate of production of product __________________. • Overall order =

  7. Example 1 • Expt [A] [B] Rate (M/sec) 1 1.0 x10-2 1.0 x 10-2 1.5 x 10-6 2 1.0 x 10-2 2.0 x 10-2 3.0 x 10-6 3 2.0 x 10-2 1.0 x 10-2 6.0 x 10-6

  8. Example 2 • Expt [A] [B] [C] Rate(M/min) 1 .20 .20 .20 2.4 x 10-6 2 .40 .30 .20 9.6 x 10-6 3 .20 .30 .20 2.4 x 10-6 4 .20 .40 .60 7.2 x 10-6

  9. Example 3 • Expt [A] [B] Rate (M/min) 1 .01 .01 6.00 x 10-3 2 .02 .03 1.44 x 10-1 3 .01 .02 1.20 x 10-2

  10. CLR Graphs • Concentration vs Time • ln of Concentration vs Time • Reciprocal of Concentration vs Time

  11. First Order Reactions • ln([A]0/[A]) = akt • A0 = Initial Concentration • A = Concentration at some point • a = Coefficient of A • k = Rate Constant • t = Time

  12. Half Life of a First Order Rxn • t1/2 = ln2/ak

  13. Compound A decomposes to form B and C in a reaction that is first order with respect to A and first order overall. At 25C, the specific rate constant for the reaction is 0.0450 sec-1. What is the half life of A? A  B + C

  14. The reaction 2N2O5 2N2O4 + O2 obeys the rate law: rate = k[N2O5], in which the specific rate constant is 0.00840 sec-1 at a certain temperature. If 2.50 moles of N2O5 were placed in a 5.00 Liter container at that temperature, how many moles of N2O5 would remain after 1.00 minutes? • How long would it take for 90% of the original N2O5 to react?

  15. Second Order Rxn • 1/[A] - 1/[A]0 = akt • t1/2 = 1/(ak[A]0) • Compounds A and B react to form C and D in a reaction that was found to be second order in A and second order overall. The rate constant is 0.622 liters per mole per minute. What is the half life of A when 4.10 x 10-2 M A is mixed with excess B.

  16. The gas phase decomposition of NOBr is second order in [NOBr] with k = 0.810 M-1sec-1 at 10C. We start with 4.00 x 10-3M NOBr in a flask at 10C. How many seconds does it take to use up 1.50 x 10-3 M of this NOBr? 2NOBr  2 NO + Br2

  17. Consider the reaction in the previous problem. If we start with 2.40 x 10-3M NOB, what concentration of NOBr will remain after 5.00 minutes of reaction?

  18. Zero Order Rxn • Rate = -1/a(D[A]/Dt) = k • t1/2 = [A]0/2ak

  19. http://genchem1.chem.okstate.edu/CCLI/Startup.html

  20. Collision Theory of Reaction Rates • In order for a reaction to occur the molecules, atoms, or ions must ______________. • In order for the collision to be successful the reactants must • A. • B. • How does the presence of a Catalyst help with this?

  21. The Transition StateActivation Energy

  22. Rate Determining Step • A reaction can never proceed faster than its ____________ step. • Most reactions occur not in one step, but instead in several smaller steps that include fast moving steps as well as slow moving steps.

  23. Example 1 • NO2 + NO2 N2O4 (Slow) • N2O4 + CO  NO + CO2 + NO2 (Fast) • What are the intermediates? • What is the rate law for this reaction?

  24. Example 2 • NO + Br2 NOBr2 (fast) • NOBr2 + NO  2NOBr (slow) • What are the intermediates? • What is the rate law for this reaction?

  25. Example 3 • I2 2I (fast) • I + H2  H2I (fast) • H2I + I  2HI (slow) • What are the intermediates? • What is the rate law for this reaction?

  26. Temperature and Reactions • As the temperature goes up __________ molecules have the energy required. • Arrheius Equation • k = Ae-Ea/RT • ln k = lnA – Ea/RT • A is a constant that is equal to the # of collisions • R = 8.314 • As Ea goes up what happens to reaction speed? • As T goes up, what happens to reaction speed?

  27. ln (k2/k1) = Ea/R x (1/T1 – 1/T2) • The specific rate constant, k, for the following first order reaction is 9.16 x 10-3 sec-1 at 0C. The activation energy of this reaction is 88.0 kJ/mol. Determine the value of k at 2.0C. • N2O5 NO2 + NO3

  28. The gas phase decomposition of ethyl iodide to give ethylene and hydrogen iodide is a first order reaction. At 600K, the value of k is 1.60 x 10-5 sec-1. When the temperature is raised to 700K, the value of k increases to 6.36 x 10-3 sec-1. What is the activation energy for this reaction?

  29. Catalyst • Homogeneous Catalyst • Heterogeneous Catalyst • Enzymes • Substrates

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