SCH4U Unit #1: Energy Changes & Rates of Reactions (Cont’d)

1. SCH4UUnit #1: Energy Changes &Rates of Reactions (Cont’d) Ms. Cornacchione Wed Feb 26th 2014

2. RateC= D[A] D[C] Dt Dt RECAP Stoich. Rate Relationships RateA = −

3. RECAP Stoich. Rate Relationships C3H8 + 5 O2 3 CO2 + 4 H20 • How does the rate of O2 compare to that of C3H8? • How does the rate of CO2 compare to that of O2? • If the rate of consumption of C3H8 is 2.3 x 10-3mol/Ls, calculate the rate of production of CO2?

4. Reaction Rates - Practice Makes Perfect • “Intro To Rates, Collision Theory and Potential Energy Diagrams” Worksheet • Pg 350 Practice #1 • Pg 352 Practice #1 • Pg 357 Practice #2 • Pg 360 Practice #2 • Page 361 #3, 4

5. Unit #2: Rates of ReactionTOPICS • Reaction Rates • Factors Affecting Rates • Rate Law & Reaction Order • Collision Theory • Reaction Mechanisms & Rate Determining Step

6. Rate Law Equation & Reaction Order • The Rate Law is a mathematical expression used to measure the rate of a reaction when you are provided the initial concentrations of the reactants kis the “rate constant” xis the “order with respect to A” yis the “order with respect to B” x + yis the “total order of the reaction” aA + bB products Rate = k[A]x[B]y

7. Rate Law Equation Example Rate = k[A]1[B]2 Example: When the initial concentration of reactant A is doubled, the rate of the reaction doubles (x2) When the initial concentration of reactant B is doubled, the rate of the reaction is quadruples (x4)

8. Experimental Data to Order of Reaction Match the graph with the correct statement: 1. 2. 3. 4.

9. Rate Law Example (HI) Rate = k[A]x[B]y At elevated temperatures, HI reacts according to the chemical equation: 2HI (g)  H2(g) + I2(g) At 443°C, the initial rate of the reaction increases with concentration of HI, as shown in the table below: • Determine the rate law • Calculate the rate constant • What is the order of each reactant

10. Rate Law Example (NO) Rate = k[A]x[B]y The rate of the reaction of the following reaction was measured for different initial concentrations of reactants and shown below: 2NO (g) + O2(g)  2NO2(g) • Determine the rate law • Calculate the rate constant • What is the order of each reactant

11. Units for the Rate Constant, k Read Page 381 Rate = k[A]x[B]y

12. Rate Law Equation & Reaction Order

13. Rate Law Equation & Reaction OrderPractice Makes Perfect • Study Sample Problem 1 & 2 (Pg 378 & Pg 379) • Try Practice Problems (Page 380) • “Rate Law” Worksheet

14. Unit #2: Rates of ReactionTOPICS • Reaction Rates • Factors Affecting Rates • Rate Law & Reaction Order • Collision Theory • Reaction Mechanisms & Rate Determining Step

15. Collision Theory

16. Collision Theory In order for reactions to occur: • Particles must collide • Particles must collide in the correct orientation (collision geometry) • Particles must collide with sufficient energy to break the bonds of the reactants (this energy is called the ACTIVATION ENERGY) ANALOGY: Air bags only go off if 1) you hit something, 2) you hit at a high enough speed, 3) you hit the appropriate sensor

17. 2. Orientation (Collision Geometry) + If the bromine atoms in the nitrosyl bromide molecules do not make direct contact, the reaction cannot form products and there will be no reaction

18. 2. Orientation (Collision Geometry)

19. 2. Orientation (Collision Geometry) RXN

20. 2. Orientation (Collision Geometry) RXN

21. 2. Orientation (Collision Geometry) • The larger the molecule or ion, or the more complex the molecule, the slower the rate of reaction • Why? These complex molecules have more bonds and are less likely to collide in the correct orientation

22. 3. Activation Energy (Ea) • For collisions to be successful the reactants must collide with enough kinetic energy to ultimately break the bonds in the reactants • Ea represents the minimum amount of kinetic energy that colliding molecules must posses to react • Ea varies from reaction to reaction

23. Potential Energy Diagrams & Ea Shows energy of the products, reactants and intermediates Ea Ea

24. Activated Complex (Transition State) An unstable arrangement of atoms containing partially formed and partially broken bonds that represents the maximum potential energy point Bonds being broken Bonds being formed

25. Activation Energy Affects Rate • The magnitude of the activation energy (energy to get over the hill) is inversely related to rate of the reaction • It is measured by taking the difference in energy between the reactants and the maximum energy (transition state) Which has the greater energy barrier, Ea? Ea

26. Maxwell-Boltzman Distribution • Chemical systems have particles in constant random motion at various speeds • Recall that Earepresents the minimum amount of kinetic energy that colliding molecules must posses to react, so only the particles with MORE energy than Ea have enough energy for a successful collision (reaction) reaction no reaction Ea

27. Maxwell-BoltzmanDistribution & Temperature • The average kinetic energy of the particles is proportional to the temperature of the sample • At a higher temperature, T2, there is a larger fraction of reactants that have enough kinetic energy needed to initiate a reaction If T increases by 10°C, then rate doubles If T decreases by 10°C, then r is halved

28. Maxwell-BoltzmanDistribution & Catalyst Uncatalyzed What do you think the potential energy diagram will look like when comparing catalyzed and uncatalyzed? Label Ea1 and Ea2 Catalyzed Does a catalyst change ΔH? NO!

29. Maxwell-BoltzmanDistribution & Concentration What do you think the MBD diagram will look like for two different concentrations? Label the Ea Number of Particles Kinetic Energy

32. Collision Theory & Factors Affecting Rates - Practice Makes Perfect • Do #1-5 on Pg 365 • Do #1-5 on Pg 372 • Complete “Factors • Affecting Rate Of Change” Worksheet (also the “Intro To Rates, Collision Theory …” Worksheetwhich was already assigned)

33. Unit #2: Rates of ReactionTOPICS • Reaction Rates • Factors Affecting Rates • Rate Law & Reaction Order • Collision Theory & Rate Determining Step • Reaction Mechanisms

34. Reaction Mechanism • Most reactions occur in a series of steps (called elementary steps) • These steps make up a reaction mechanism • Each step has a different reaction rate • The slowest is called the rate determining step

35. The Rate Determining Step Reaction Intermediate The Rate Determining Step is: __________________ Potential Energy Time (or Reaction Progress)

36. Molecularity & Rate of Elementary Steps Molecularity -the number of reactants (ions, molecules or atoms) that are involved in the reaction: UNIMOLECULAR (1 reactant) BIMOLECULAR (2 reactants) TERMOLECULAR: (3 reactants) The rate law equation for elementary steps is very easy! The order with respect to each reactant is the stoichiometric coefficient 2A

37. Reaction Mechanism is possible IF AND ONLY IF: • Elementary steps must be unimolecular, bimolecular, or termolecular • Elementary steps must add up to the overall reaction • The rate law of the slowest elementary step must be equal to the rate law of the overall reaction

38. Reaction Mechanism Example

40. Reaction Mechanism - Practice Makes Perfect • Complete “Reaction Mechanism” Worksheet • Study Sample Problems 1 & 2 on Pg 385 • Do Practice #1-3 on Pg 386