Chemical Kinetics

1. Chemical Kinetics

2. Chemical Kinetics Do Now: TRUE / FALSE: By definition, spontaneous reactions are fast. What is meant by the term “rate of reaction”? List three factors that may affect the rate of reaction, and indicate how/why.

3. Section Objectives • Explain the rate of reaction is the decrease in concentration of reactants or the increase in concentration of products with time • Describe how reaction rates depend on such factors as concentration, temperature, and pressure • Explain the role a catalyst plays in increasing the reaction rate

4. Reaction Rates • Chemical reactions don’t all happen at the same rate… • Some reactions, like the explosion of dynamite, are completely finished in less than a second. • Others, like the rusting of iron, may not completely finish for many years.

5. Chemical Kinetics is a study of the rate at which a successful chemical reactions occur. • Several factors can influence the rate or a chemical reaction. • Each factor must have a discernible effect on the microscopic collisions that lead to a successful chemical reaction

6. Factors that Affect Reaction Rate • Concentration • Temperature • Solid particle size • Catalysts

7. Concentration • Increasing the reactants concentration by putting more reactants into the same space will increase rate • If the particles are more concentrated, they will collide more and react faster. • **similar effect is observed when increasing the pressure (*reducing volume*) in a gaseous reaction

8. Temperature • A rise in temperature will result in an increased rate of reaction • As a general rule, a 10 oC rise in temperature will double the rate of the reaction. • Higher temperature means particles are moving faster • Therefore hit (collide with) each other more often and with more energy; • they will react faster.

9. Maxwell Boltzmann graph • Consider the Maxwell-Boltzmann distribution plot of energies • The area underneath the curve represents the total number of particles. • The vertical line separating the shaded region indicates the activation energy (Eact)

11. Eact: Only particles with energies greater than activation energy will have successful collisions • At higher temperatures, the average energy of the molecules is greater • As a result, greater number of molecules have energies > Eact (notice areas under shaded region of the curves above) • A reaction will always be much faster at higher temperatures • More particles possess the required, minimum energy on collsion • Therefore, greater number of successful collisions occur. • Additionally, increase in temperature results in increase in frequency of collisions.

12. Surface Area • When a solid reacts, only the particles on the surface of the solid are available for reaction. • The greater the surface area, the greater the number of particles that can collide • And result in successful collisions. • When a solid is broken up into smaller pieces, its surface area gets larger and more particles are available for successful collisions • Example: Small crystals of NaCl dissolve in water faster than rock salt. • As a result, reaction rate increases.

13. Catalyst: Some substances aren’t reactants in a reaction, but they can speed up the reaction by being present. • Discussed in topic 4D

14. Reaction Rates • There are two ways to measure the rate of a reaction… • Either measure the rate at which a REACTANT DISSAPPEARS (decrease in concentration of reactants over time)or… • …measure the rate at which a PRODUCT APPEARS (increase in concentration of products over time) • For example: Beer’s Law can be used to determine change in concentration of colored solutions

15. Reaction Rates • Often, chemists are trying to control the rate of a chemical reaction • Sometimes, we wish to speed a reaction up, like when we design a quick-drying paint formula. • Other times, we want to slow reactions down, like when we want the paint to resist fading for many years.

16. Collision Theory • Forms the basis of the study of the speed or rate of chemical reactions (chemical kinetics) • Indicates that a reaction will only occur (be successful) if THREE conditions are met.

17. Conditions for successful collisions • Reactants must come in contact (be able to collide) • Collision must occur with a certain minimum energy (activation energy, Eact) • Collisions must have correct molecular orientation • Reactants must collide in a certain physical, three-dimensional orientation such that a specific part of one reactant species comes in contact with a specific part of another reactant species

19. Collision Theory • If reactants do not collide, or collide with energies lower than activation energy, or collide without correct molecular orientation, then NO reaction can occur. • reaction will be unsuccessful, • Reactants will remain unchanged

20. Factors in Successful Collisions • There are multiple factors that affect the rate of a chemical reaction. • Collision Frequency: • By increasing concentration and/or temperature we can increase the frequency at which molecules collide and increase the rate of reaction. • WHY? • Concentration: __________ • Temperature: ___________

21. Collision Energy: • For a reaction to occur, the molecules must collide with enough energy to form the new bonds. Increase temperature to increase the energy. • Greater temperature = more # molecules with greater energy = greater probability of molecules with E > Eact

22. Collision Geometry: • For a reaction to occur, the molecules must be oriented in the proper geometry for the reaction to occur. • A Catalyst can help with the geometry. It is NOT consumed by the reaction

23. Solid Particle Size: • When a solid reacts, only particles on the surface of the solid are available to react. • Solid broken up into smaller pieces results in greater surface area • Therefore more particles are available for collision, and reaction rate increases.

24. Energy Barriers in Reactions • For a chemical reaction to occur, the reactants must collide with sufficient energy to react. • This energy is required to achieve the transition state required to form the products (a). • WITHOUT sufficient energy, the reaction does not occur (b).

25. Reaction Profiles • A reaction profile shows the energy of reactants and products during a reaction. • The highest point on a reaction profile is the transition stateor the activated complex.

26. The energy required for reactants to achieve the transition state is the activation energy, Eact. • The energy difference between reactants and products is the heat of reaction, ΔH.

27. Endothermic Reactions • An endothermicreaction absorbs heat as the reaction proceeds. • N2(g) + O2(g) + heat 2 NO(g) • Endothermic: reactants have LOWER energy than products • ΔH for endothermic reaction is positive

28. Exothermic Reactions • An exothermic reaction releases heat as the reaction proceeds. • NO(g) + O3(g)  2 NO2(g) + O2(g) + heat • The ΔH for an exothermicreaction is negative. • The products areat a lower energythan the reactants

29. Effect of a Catalyst • A catalyst is a substance that allows a reaction to proceed faster by LOWERING the activation energy. • It does this by finding an alternative path for the reaction to occur • An inhibitor is theopposite of a catalyst

30. A catalyst does not change ΔH for a reaction. • A catalyst speeds up both the forward and reverse reactions. • Examples of biological catalysts: enzymes • A catalyst can’t make a reaction occur that wouldn’t otherwise also occur.

31. Kinetics • Studies the rate at which a chemical process occurs. • Besides information about the speed at which reactions occur, kinetics also sheds light on the reaction mechanism (exactly how the reaction occurs).

32. Reaction Rates Rates of reactions can be determined by monitoring the change in concentration of either reactants or products as a function of time.

33. Factors That Affect Reaction Rates • Physical State of the Reactants • In order to react, molecules must come in contact with each other. • The more homogeneous the mixture of reactants, the faster the molecules can react.

34. Factors That Affect Reaction Rates 2. Concentration of Reactants • As the concentration of reactants increases, so does the likelihood that reactant molecules will collide.

35. Factors That Affect Reaction Rates 3. Temperature • At higher temperatures, reactant molecules have more kinetic energy, move faster, and collide more often and with greater energy.

36. Factors That Affect Reaction Rates 4. Presence of a Catalyst • Catalysts speed up reactions by changing the mechanism of the reaction. • Catalysts are not consumed during the course of the reaction. • Enzymes are biological catalysts

37. The effect of increasing the partial pressures of the reactive components of a gaseous mixture depends on which side of the chemical equation has the most gas molecules. • Increasing the partial pressures of the reactive components of a gaseous mixture has no effect on the rate of reaction if each reactant pressure is increased by the same amount. • Increasing the partial pressures of the reactive components of a gaseous mixture increases the rate of reaction. • Increasing the partial pressures of the reactive components of a gaseous mixture decreases the rate of reaction.

38. The effect of increasing the partial pressures of the reactive components of a gaseous mixture depends on which side of the chemical equation has the most gas molecules. • Increasing the partial pressures of the reactive components of a gaseous mixture has no effect on the rate of reaction if each reactant pressure is increased by the same amount. • Increasing the partial pressures of the reactive components of a gaseous mixture increases the rate of reaction. • Increasing the partial pressures of the reactive components of a gaseous mixture decreases the rate of reaction.

39. 12.1 – Reaction Rates • The speed of a chemical reaction – reaction rate – is the change in concentration of reactants or products per unit time • Units are usually molarity per second ([M]/s) • Average rate = Δ [A]ΔT**NOTE: Rate may be positive (appearance of product) or negative (disappearance of reactant). For convenience, we will always define rate at positive

41. Calculate the average rate at which A disappears over the time interval from 20 s to 40 s.

42. SAMPLE EXERCISE 1 Calculating an Average Rate of Reaction From the data given in the caption of figure above, calculate the average rate at which A disappears over the time interval from 20 s to 40 s.

43. PRACTICE EXERCISE Calculate the average rate of appearance of B over the time interval from 0 to 40 s. Answer: 1.8  10 –2M/s

44. Reaction Rates C4H9Cl(aq) + H2O(l) C4H9OH(aq) + HCl(aq) In this reaction, the concentration of butyl chloride, C4H9Cl, was measured at various times.

45. Average rate = [C4H9Cl] t Reaction Rates C4H9Cl(aq) + H2O(l) C4H9OH(aq) + HCl(aq) The average rate of the reaction over each interval is the change in concentration divided by the change in time:

46. Reaction Rates C4H9Cl(aq) + H2O(l) C4H9OH(aq) + HCl(aq) • Note that the average rate decreases as the reaction proceeds. • This is because as the reaction goes forward, there are fewer collisions between reactant molecules. • WHY?

47. Reaction Rates C4H9Cl(aq) + H2O(l) C4H9OH(aq) + HCl(aq) • A plot of concentration vs. time for this reaction yields a curve like this. • The slope of a line tangent to the curve at any point is the instantaneous rate at that time.

48. Reaction Rates C4H9Cl(aq) + H2O(l) C4H9OH(aq) + HCl(aq) • All reactions slow down over time. • Therefore, the best indicator of the rate of a reaction is the instantaneous rate near the beginning.

49. Larger triangles should be used if x and y values are three significant figures or more. • The size of the triangle is mostly a matter of convenience; different size triangles give the same ratio (slope). • The size of the triangle determines the number of significant digits in your final slope calculation. • Smaller triangles should be used only near the ends of the curve while larger triangles are used in the center only.

50. Larger triangles should be used if x and y values are three significant figures or more. • The size of the triangle is mostly a matter of convenience; different size triangles give the same ratio (slope). • The size of the triangle determines the number of significant digits in your final slope calculation. • Smaller triangles should be used only near the ends of the curve while larger triangles are used in the center only.