1. Rates of Reaction Learning Goal: I will understand units of measure of reaction rates, and know the factors that increase reaction rate.

2. Chemical Reaction Rates Chemical kinetics is the study of the rate at which chemical reactions occur. The term reaction rate, or rate of reaction refers to: the speed that a chemical reaction occurs at, or the change in amount of reactants consumed or products formed over a specific time interval

3. Determining Reaction Rates The reaction rate is often given in terms of the change in concentration of a reactant or product per unit of time. The change in concentration of reactant A was monitored over time.

4. Average and Instantaneous Reaction Rates Average rate of reaction: change in [reactant] or [product] over a given time period (slope between two points) Instantaneous rate of reaction: the rate of a reaction at a particular point in time (slope of the tangent line) Average rate of reaction and instantaneous rate of reaction can be determined from a graph of concentration vs. time.

5. A B 13.1 rate = -  [A] tt rate = [B][B] tt time

6. Determining Reaction Rates Example # 1 a)Using the data below, determine the instantaneous reaction rate that occurs between 40.0 and 50.0 seconds. b) Determine the reaction rate between time 0 and t=60.0s. The change in concentration of reactant A was monitored over time.

7. Example #2 For the reaction 2A + B  3C, it was found that the rate of disappearance of B was 0.300molL -1 s -1. What is the rate of disappearance of A and appearance of C?

8. Learning Check At a certain temperature, the rate of decomposition of N 2 O 5 is 2.5 x 10 -6 molL -1 s -1. How fast is each product being formed? N 2 O 5  NO 2 + O 2

9. Learning Check #2 If NaOH and HCl react to form 25 g of water in 3 seconds, express the reaction rate in g/s, then mol/s

10. Learning Check #3 If two moles of nitrogen monoxide reacts with oxygen gas to form nitrogen dioxide, a)Write a balanced chemical equation for this reaction b)If 0.58 mol of oxygen reacts in 30s, calculate the: i) grams of oxygen consumed per second ii) moles of nitrogen monoxide consumed per second iii) moles of nitrogen dioxide formed per second

11. Factors Affecting Reaction Rate 1. Nature of reactants reactions of ions tend to be faster than those of molecules 2. Concentration a greater number of effective collisions are more likely with a higher concentration of reactant particles 3. Temperature with an increase in temperature, there are more particles with sufficient energy needed for a reaction

12. 4. Pressure for gaseous reactants, the number of collisions in a certain time interval increases with increased pressure 5. Surface area a greater exposed surface area of solid reactant means a greater chance of effective collisions 6. Presence of a catalyst a catalyst is a substance that increases a reaction rate without being consumed by the reaction

13. Apply Your Knowledge 1.What a campfire burn faster with kindling or large logs? Why…which factor does this relate to? 2.Why do we use a refrigerator to keep our food fresh longer? Which factor does this relate to? 3.How can a magnifying glass be used to start a fire? Which factor does this relate to? 4.If you spill concentrated HCl on your skin, will it have a stronger/weaker effect? Why? What factor does this relate to?

14. Homework Pg 361 Q 1, 2, 5 Pg 365 Q1

15. Collision Theory and Factors Affecting Rates of Reaction According to collision theory, a chemical reaction occurs when the reacting particles collide with one another. Only a fraction of collisions between particles result in a chemical reaction because certain criteria must be met.

16. Effective Collision Criteria 1: The Correct Orientation of Reactants For a chemical reaction to occur, reactant molecules must collide with the correct orientation relative to each other (collision geometry). Five of many possible ways that NO(g) can collide with NO 3 (g) are shown. Only one has the correct collision geometry for reaction to occur.

17. Effective Collision Criteria 2: Sufficient Activation Energy The shaded part of the Maxwell- Boltzmann distribution curve represents the fraction of particles that have enough collision energy for a reaction (ie the energy is ≥ E a ). For a chemical reaction, reactant molecules must also collide with sufficient energy. Activation energy, E a, is the minimum amount of kinetic energy the reactants must possess in order to have an effective collision (energy required for products to form) Collision energy depends on the kinetic energy of the colliding particles.

18. In other words, there is a minimum amount of energy required for reaction: the activation energy, E a. Just as a ball cannot get over a hill if it does not roll up the hill with enough energy, a reaction cannot occur unless the molecules possess sufficient energy to get over the activation energy barrier.

19. Potential Energy Curves From left to right on a potential energy curve for a reaction: potential energy increases as reactants become closer when collision energy is ≥ maximum potential energy, reactants will transform to a transition state products then form (or reactants re-form if ineffective) ExothermicEndothermic

20. Activation Energy and Enthalpy The E a for a reaction cannot be predicted from ∆H. ∆H is determined only by the difference in potential energy between reactants and products. E a is determined by analyzing rates of reaction at differing temperatures. Reactions with low E a occur quickly. Reactions with high E a occur slowly. Potential energy diagram for the combustion of octane.

21. Analyzing Reactions Using Potential Energy Diagrams The BrCH 3 molecule and OH - collide with the correct orientation and sufficient energy and an activated complex forms. When chemical bonds reform, potential energy decreases and kinetic energy increases as the particles move apart.

22. Activation Energy for Reversible Reactions Potential energy diagrams can represent both forward and reverse reactions. follow left to right for the forward reaction follow right to left for the reverse reaction

23. Draw a potential energy curve for a reversible exothermic reaction. Label the reactants, products, transition state (activated complex), activation energy (for both the forward and reverse reaction (E a(fwd) & E a(rev) ) and the change in enthalpy. L EARNING C HECK

24. Reaction Mechanisms Reaction Mechanisms are individual steps that led to the overall chemical change Ex. HBr + O 2  H 2 O + Br 2 … is not likely to occur in 1 single step Instead: HBr + O 2  H 2 OFast HOOBr + HBr  HOBrSlow HOBr HBr  H 2 O + Br 2 Slow

25. Transition-State Theory This explains the reaction in terms of the collision of two high energy species – activated complexes ▫ An activated complex (transition state) is an unstable grouping of atoms that can break up to form products ▫ A simple analogy would be the collision of three billiard balls on a billiard table  Suppose two balls are coated with a slightly stick adhesive  We’ll take a third ball covered with an extremely sticky adhesive and collide it with our joined pair 25

26. Reaction Mechanisms Steps in the overall reaction that detail how reactants change into products ▫ Reaction Mechanism – set of elementary reactions that leads to overall chemical equation ▫ Reaction Intermediate – species produced during a chemical reaction that do not appear in chemical equation ▫ Elementary Reactions – single molecular event resulting in a reaction ▫ Molecularity – number of molecules on the reactant side of elementary reaction ▫ Rate Determining Step (RDS) – slowest step in the reaction mechanism This is the reaction used to construct the rate law; it is not necessarily the overall reaction 26

27. In a multistep process, one of the steps will be slower than all others. The overall reaction cannot occur faster than this slowest, rate-determining step.

28. Rate Determining Step Which would be the rate determining step for the following potential energy curve?

29. A Catalyst Influences the Reaction Rate A catalyst lowers the E a of a reaction. this increases the percentage of reactants that have enough kinetic energy to overcome the activation energy barrier a catalyzed reaction has the same reactants, products, and enthalpy change as the uncatalyzed reaction A catalyst decreases both E a(fwd) and E a(rev).

30. Catalysts in Industry A metal catalyst is used for industrial-scale production of ammonia from nitrogen and hydrogen. Hydrogen and nitrogen molecules break apart when in contact with the catalyst. These highly reactive species then recombine to form ammonia. Many industries use biological catalysts, called enzymes, which are most often proteins. For example: the use of enzymes decreases the amount of bleach (an environmental hazard) needed to whiten fibres used in paper production.

31. Homework Pg 372 Q1, 4, 5ab, 6

32. Rate Law Learning Goals: I will be able to express an average or instantaneous reaction rate for a given reaction, as well as understand the rate law and its components and be able to determine the reaction order, and solve for the rate (k)

33. “How fast does the reaction occur?” Kinetics is the study of the rate of chemical reactions ▫rate is a time dependent process  Rate units are concentration over time Consider the reaction A  B ▫Reactant concentrations [A] decrease while product concentrations [B] increase

34. The rate of reaction for the formation of the products is the same as for the reactants, but opposite, that is the concentrations are increasing The rate of change of Ethane (C 2 H 4 ) and Ozone (O 3 ) is the same, but exactly opposite for Acetaldehyde (C 2 H 4 O) and Oxygen (O 2 ) Product concentration increases at the same rate that the reactant concentrations decrease The curves have the same shape, but are inverted C 2 H 4 + O 3 → C 2 H 4 O + O 2 Rate = -∆[ C 2 H 4 ] = - ∆ [O 3 ] = + ∆ [ C 2 H 4 O] = +∆[O 2 ] ∆t ∆t ∆t ∆t

35. The Rate expression must be consistent with stoichiometry When the stoichiometric molar ratios are not 1:1, the reactants still disappear and the products still appear, but at different rates H 2 (g) + I 2 (g) → 2HI (g) For every molecule of H 2 that disappears, one molecule of I 2 disappears and 2 molecules of HI appear The rate of H 2 decrease is the same as the rate of I 2 decrease, but both are only half the rate of HI increase Rate = -∆[H 2 ] = - ∆[I 2 ] = 1 ∆[HI] ∆t ∆t2 ∆t

36. Summary equation for any reaction: aA + bB → cC + dD Rate = -1 ∆[A] = -1 ∆[B] = +1 ∆[C] = +1 ∆[D] a ∆t b ∆t c ∆t d ∆t

37. Rate Law The dependence of reaction rate on concentrations is expressed mathematically by the rate law The rate law expresses the rate as a function of reactant concentrations, product concentrations, and temperature Only the Reactants Appear in the Rate Law For a general reaction at a fixed temperature: aA + bB + …  cC + dD + … the rate law has the form: Rate = k[A] m [B] n …

38. aA + bB + …  cC + dD + … Note: The Stoichiometric Coefficients – a, b, c – are not used in the rate law equation; they are not related to the reaction order terms – m, n, p, etc

39. ▫The rate law “k” changes with temperature; thus it determines how temperature affects the rate of the reaction ▫The exponents (m, n, p, etc.) are called reaction orders, which must be determined experimentally ▫Reaction orders define how the rate is affected by the reactant concentration  If the rate doubles when [A] doubles, the rate depends on [A] raised to the first power, i.e., m =1 (a 1 st order reaction)  If the rate quadruples when [B] doubles, the rate depends on [B] raised to the second power, i.e., n = 2 (a 2 nd order reaction)  If the rate does not change even though [A] doubles, the rate does not depend on the concentration of A and p = 0 (zero order reaction)

40. Reaction Order Tells How Changing Reactant Concentration Affects Rate O th Order; rate =k[A] 0 ; double [A], rate same 1 st Order; rate = k[A] 1 ; double [A], double rate 2 nd Order; rate =k[A] 2 ; double [A]; quadruple rate 3 rd Order; rate =k[A] 3 ; double [A]; rate 8X

41. Example Given: NO(g) + O 3 (g)  NO 2 (g) + O 2 (g) Given: Rate = k[NO] 1 [O 3 ] 1 What is the reaction order with respect to a)NO b)O 3 c)Overall?

42. Overall reaction order: Sum of Exponents in Rate Equation Order of RxnPossible Expression of Rate Law 1k[A] 2k[A] 2 2k[A][B] 3 k[A] 2 [B] 3k[A][B][C]

43. Example Experiments are performed for the reaction: A  B + C and the rate law has the been determined to be of the form Rate = k [A] x Determine the value of the exponent “x” for each to determine the rate law of each of the following: a) [A] is tripled and you observe no rate change b) [A] is doubled and the rate doubles c) [A] is double and the rate quadruples

44. F 2 (g) + 2ClO 2 (g) 2FClO 2 (g) rate = k [F 2 ] x [ClO 2 ] y Double [F 2 ] with [ClO 2 ] constant Rate doubles x = 1 Quadruple [ClO 2 ] with [F 2 ] constant Rate quadruples y = 1 rate = k [F 2 ][ClO 2 ] 13.2

45. Initial Reactant Concentration (mol/L) ExperimentO2O2 NO Initial Rate Mol/L  s 11.10x10 -2 1.30x10 -2 3.21x10 -3 22.20x10 -2 1.30x10 -2 6.40x10 -3 33.30x10 -2 1.30x10 -2 9.60x10 -3 41.10x10 -2 2.60x10 -2 12.8x10 -3 51.10x10 -2 3.90x10 -2 28.8x10 -3 As [O 2 ] doubles, the reaction rate _________, therefore m=___ As [NO] doubles, the reaction rate _________, therefore n= ___

46. Example A certain reaction that follows the equation A + B  C +D gave the following data: Calculate the reaction order a)reactant A b) reactant B c)the overall reaction order d)the rate law e)the rate law constant Initial Reactant Concentration (mol/L) Experimen t [A][B] Initial Rate Mol/L  s 10.40.31.0x10 -4 20.80.34.0x10 -4 30.80.61.6x10 -3

47. Homework pg 382 Q1, 2, 4abc