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Chemical Kinetics

Chemical Kinetics. AP Chem. Chemical Kinetics. Area of chemistry that deals with rates or speeds at which a reaction occurs The rate of these reactions are affected by several factors. Describing Rates. Such as: Concentration Color Bubbles Temp pH. Whatever is appropriate: Hours

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Chemical Kinetics

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  1. Chemical Kinetics AP Chem

  2. Chemical Kinetics • Area of chemistry that deals with rates or speeds at which a reaction occurs • The rate of these reactions are affected by several factors

  3. Describing Rates Such as: Concentration Color Bubbles Temp pH Whatever is appropriate: Hours Minutes Seconds

  4. Describing Rates • Candle Wax Example

  5. Collision Theory • Collisions cause reactions! • Breaking of bonds directly linked to rate • Must overcome repulsion of electron clouds • Also correct orientation sometimes needed • Example: Chalk dropping

  6. Factors Affecting Rates of Reaction • Concentration of reactants • With gases, pressure used instead • Temperature at which the reaction occurs • The presence of a catalyst • Surface area of solid or liquid reactants

  7. Factors Affecting Rates • Demo: • Chalk + 1 M HCl / Chalk + 1 M Acetic Acid • Prediction? • Factor? • Chalk + 1 M HCl / Chalk + 6 M HCl • Prediction? • Factor? • Chalk + 1 M HCl (Room Temp) / Chalk + 1 M HCl (Heated) • Prediction? • Factor?

  8. Average Reaction Rates • Speed of reaction or reaction rate is the time over which a change occurs • Consider the reaction A  B • Reaction rate is a measure of how quickly A is consumed or B is produced

  9. Average Reaction Rates • Average rate of reaction can be written: • This is a measure of the average rate of appearance of B

  10. Average Reaction Rates • Average rate can also be written in terms of A: • This is the rate of disappearance of A (equal to B only negative) • Average Rates can only be positive

  11. Average Reaction Rates • Start with one mole of A at time zero, measure amounts of A and B at given time intervals

  12. Average Reaction Rate • Data for Reaction A  B

  13.  Moles A  Moles B

  14. Rates and Stoichiometry • When mole ratios of equations are not 1:1 • For the reaction: aA + bB cC + dD

  15. Rates and Stoichiometry • Example • If the rate of decomposition of N2O5 in a reaction vessel is 4.2 x 10-7 M/s, what is the rate of appearance of NO2 and O2 2 N2O5(g)  4 NO2(g) + O2(g)

  16. Instantaneous Rate of Reaction • Consider the reaction between butyl chloride and water: C4H9Cl(aq) + H2O(l) C4H9OH(aq) +HCl(aq)

  17. Instantaneous Rate of Reaction • Using the curve created from the data, we can determine the instantaneous rate for any given point on the curve • Recall: slope is rise over run!

  18. Instantaneous Rate vs. Average Rate • Analogy: Distance between Fall River and Norton is 29.7 mi along a certain route. It takes Mr. N 30 minutes to get to school. His average rate is 59.4 mph. • But at t = 15, Mr. N’s instantaneous rate is 95 mph, and at t = 1 Mr. N’s instantaneous rate is 25 mph.

  19. Concentration and Rates • Increasing concentration of reactants gives increasing rate • Decreasing rates of reactions over time is typical • Due to decreasing concentration of reactants

  20. Rate Laws • Rates of a reaction can be related to concentrations with a rate constant (k) • For reaction: • Rate laws are defined by reactant (not product) concentrations aA + bB cC + dD

  21. Reaction Order • For the rate law expression: • The overall order of reaction is the sum of the powers (x + y) • However, rate with respect to [A] is only x

  22. Reaction Order • In most rate laws reaction orders are 0, 1, or 2 • Can be fractional or negative at times • Most commonly 1 or 2 • Reaction orders are determined experimentally, and do not necessarily relate to coefficients of a balanced equation

  23. Reaction Order • Example: What is the overall order of reaction for the reaction below? CHCl3(g) + Cl2(g) CCl4(g) + HCl(g) Rate= k[CHCl3][Cl2]1/2 A.) ½ B.) 2 C.) 3/2 D.) 2/2

  24. Meaning of Reaction Order • Zero orderfor a reactant means concentration changes have no effect on reaction rate • Example: Drinking • 1st ordermeans concentration changes give proportional changes in reaction rate • Double the concentration, double the rate

  25. Meaning of Reaction Order • 2nd orderrate law, increasing in concentration results in a rate increase equal to the concentration increased to the second power • Example: • Double conc. = 22 = 4 (rate increase) • Triple conc. = 32 = 9 (rate increase)

  26. Units of Rate Constant (k) • The units for the rate constant depend on the order of the rate law Z = overall order of reaction

  27. Units of Rate Constant (k) • What is the unit for the rate constant for the reaction below? CHCl3(g) + Cl2(g) CCl4(g) + HCl(g) Rate= k[CHCl3][Cl2]1/2 A.) M½/s B.) M/s C.) M2/s D.) M-1/2/s

  28. Determining Order of Reactants from Experimental Data • A particular reaction was found to depend on the concentration of the hydrogen ion. The initial rates varied as a function of [H+] as follows: [H+] (M) 0.0500 0.100 0.200 Initial rate(M/s) 6.4x10-7 3.2x10-7 1.6x10-7 • What is the order of the reaction in [H+] • A.) 1 • B.) 2 • C.) -1 • D.) -1/2

  29. Determining Rate Law by Experimental Data What is the rate law expression for the reaction?

  30. Concentration and Time • Rate law tells how rate changes with changing concentrations at a particular temperature • We can derive equations that can give us the concentrations of reactants or products at any time during a reaction (instantaneous) • These are known as integrated rate laws

  31. Integrated Zero Order Reactions Rate = k • Using calculus, the integrated rate law is: • [A]t is concentration of reactant at time t • [A]0 is initial concentration of reactant

  32. Integrated Zero Order Reaction • This has the same form as the general equation for a straight line • Graphically, the slope is equal to -k

  33. Half-life (t1/2) for Zero Order • Separate equations can be derived relating to time required for reactants to decrease to half of initial concentration (aka half-life or t1/2) • When t = t1/2, [A]t is half of [A]0 ([A]t =[A]0/2)

  34. Integrated First Order Reaction Rate = k [A] • Using calculus, the integrated rate law becomes: • This equation is of the general equation for a straight line (like before)

  35. Graph of Int. First Order • Note that only the second graph is used so that the slope can be determined • Also, y-intercept is ln [A]0

  36. Int. First Order Reaction • Example 2N2O5(g) 4NO2(g) + O2(g) • The decomposition of dinitrogenpentoxide is a 1st order reaction with a rate constant of 5.1 x 10-4 s-1 at 45ºC. If initial conc. is 0.25M, what is the concentration after 3.2 min.?

  37. Int. First Order Reaction • Example #2 2N2O5(g) 4NO2(g) + O2(g) The decomposition of dinitrogenpentoxide is a 1st order reaction with a rate constant of 5.1 x 10-4 s-1 at 45ºC. How long will it take for the concentration of N2O5 to decrease from 0.25M to 0.15M?

  38. Half life for First Order Reactions • For first order reactions: • Note it is independent of concentration! • This is used to describe radioactive decay and elimination of medications from the body

  39. Int. Second Order Reaction Rate = k [A]2 • Using calculus, the integrated rate law becomes: • Just like the previous two, this equation is of the general equation for a straight line

  40. Half Life of Second Order Reaction • Unlike first order, second order does depend on initial concentrations:

  41. Summary

  42. Relationship between Temperature & Rate • Most reactions increase in rate with increasing temperature • This is due to an increase in the rate constant with increasing temperature

  43. Activation Energy, Ea • Minimum amount of energy required to initiate a chemical reaction • Varies from reaction to reaction • This is the kinetic energy required by colliding molecules in order to begin a reaction • Remember, even with sufficient KE, orientation is still important

  44. Activation Energy, Ea • Activation energy must be enough to overcome initial resistance for a reaction to take place

  45. Activation Energy Diagram • Diagram can be used to determine if reaction is exothermic (- ∆H) or endothermic (+∆H) • Activated complex (or transition state) is the arrangement of atoms at the peak of the Ea barrier • Unstable and only appears briefly

  46. Observations made by Arrhenius • The relationship between rate and temperature was non-linear • Reaction rate obeyed an equation based on 3 factors: • Fraction of molecules that possess Ea • # of collisions per second • Fraction of collisions with proper orientation

  47. Arrhenius Equation • k = the rate constant • R = gas constant (8.314 J/mol*K) • T = Absolute temperature (K) • Ea = the activation energy • A = frequency factor • A is mostly constant with variations in temperature

  48. Arrhenius Equation • Taking the natural log of both sides gives a formula in straight line form: • Graph of ln k versus 1/T will be a straight line with a slope of –Ea/R and a y-intercept of ln A

  49. Arrhenius Equation • In order to compare different rates at different temperatures the equation can be rearranged:

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