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  1. Chemical Kinetics http://www.chem1.com/acad/webtext/dynamics/dynamics-3.htmlhttp://ibchem.com/IB/ibnotes/brief/kin-sl.htm http://www.docbrown.info/page03/3_31rates.htm#TheeffectofaCatalysthttp://www.practicalchemistry.org/experiments/the-rate-of-reaction-of-magnesium-with-hydrochloric-acid,100,EX.html 1.

  2. What do we mean by kinetics? Kinetics is the study of • therateat which chemical reactions occur. • The reaction mechanismor pathway through which a reaction proceeds. 2.

  3. Reaction Rate • The change in the concentration of a reactant or a product with time (M/s). Reactant → Products A → B • Since reactants go away with time: Rate = - Δ[A]/Δt = Δ[B]/Δt mol dm-3 s-1

  4. How an Airbag Works: http://www.youtube.com/watch?v=dZfLOnXoVOQ • When sodium azide, NaN3, is ignited by a spark, it releases nitrogen gas which can instantly inflate an airbag.(one 25th of a second)  SiO2 10NaN3 (s) + 2KNO3 (s)=> K2O (s) + 5 Na2 O(s) + 16N2 (g)

  5. A => B Rate = - Δ[A]/Δt = Δ[B]/Δt until completion

  6. Reaction Rates C4H9Cl(aq) + H2O(l) C4H9OH(aq) + HCl(aq) • The reaction slows down with time because the concentration of the reactant decreases. 6.

  7. Collision Theory • http://www.saskschools.ca/curr_content/chem30_05/2_kinetics/kinetics2_1.htm • For a reaction to occur: Particles must collide with each other and these collisions must be effective: • Sufficient energy(Activation energy) and proper orientation.

  8. Activation Energy? Minimum energy required. Maxwell Boltzmann? Read CC

  9. Maxwell Boltzman • As you increase the temperature the rate of reaction increases. As a rough approximation, for many reactions happening at around room temperature, the rate of reaction doubles for every 10°C rise in temperature.

  10. Factors That Affect Reaction Rateshttp://www.docbrown.info/page03/3_31rates.htmhttp://www.blackgold.ab.ca/ict/Divison3/reactionrates/reactionrates.htm • The Nature of the Reactants • Chemical compounds vary considerably in their chemical reactivities. • Concentration of Reactants • As the concentration of reactants increases, so does the likelihood that reactant molecules will collide. • Temperature • At higher temperatures, reactant molecules have more kinetic energy, move faster, and collide more often and with greater energy. • Catalysts • Change the rate of a reaction by changing the mechanism. • Surface Area • http://www.physchem.co.za/OB12-che/rates.htm .

  11. http://www.gcsescience.com/rc4-sodium-thiosulfate-hydrochloric.htmhttp://www.gcsescience.com/rc4-sodium-thiosulfate-hydrochloric.htm • investigate the effect that the temperature of water has on the reaction rate of antacid tablets. • investigate the change in reaction rate when varying sizes of zinc particles are reacted with sulfuric acid. • explore the relationship between concentration and reaction rate. • observing how quickly reactants are used up or how quickly products are forming. This can be calculated in three ways: of precipitation - When the products of the reaction is a precipitate, which will cloud the solution. Observing a marker through the solution and timing how long it takes for this marker to disappear will determine the rate of reaction. Hydrochloric Acid and Sodium Thiosulphate HCl(aq)    +   Na2S2O3(aq)              NaCl(aq)       +     SO2(g)       +       S(s)  +  H2O(l)

  12. The graph lines W, X, original, Y and Z on the left diagram are typical of when a gaseous product is being collected • X ,a catalyst was added, forming the same amount of product, but faster. • Z could represent taking half the amount of reactants or half a concentration. The reaction is slower and only half as much gas is formed • W might represent taking double the quantity of reactants, forming twice as much gas e.g. same volume of reactant solution but doubling the concentration, so producing twice as much gas, initially at double the speed (gradient twice as steep). • W > X > original > Y > Z See SG graphs

  13. Measuring the rate of a reaction The change in concentration can be measured by using any property that changes during conversion of reactants to products: CaCO3(s) + HCl(aq) => CaCl2(aq) + H2O(l) + CO2(g) 1. mass and volume changes for gaseous reactions 2. change in pH for acids and bases 3. change in conductivity for metals 4. colorimetry for color changes

  14. Measuring the volume of gas produced with time

  15. Mass Loss Method

  16. Graphs • The graph below shows the volume of carbon dioxide gas produced against time when excess calcium carbonate is added to x cm3 of 2.0 mol dm−3 hydrochloric acid. (i)      Write a balanced equation for the reaction. (ii)      State and explain the change in the rate of reaction with time.  Outline how you would determine the rate of the reaction at a particular time. (iii)     Sketch the above graph on an answer sheet. On the same graph, draw the curves you would expect if: I.       the same volume (x cm3) of 1.0 mol dm−3  HCl is used. II.      double the volume (2x cm3) of 1.0 mol dm−3  HCl is used.           Label the curves and explain your answer in each case.

  17. Graph mass x time

  18. SG page 36

  19. Example 1: Reaction Rates C4H9Cl(aq) + H2O(l) C4H9OH(aq) + HCl(aq) [C4H9Cl] M In this reaction, the concentration of butyl chloride, C4H9Cl, was measured at various times, t. [C4H9Cl] t Rate = 20.

  20. Reaction Rates Calculation 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: Average Rate, M/s 21.

  21. Reaction Rate Determination 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 the reacting molecules. 22.

  22. 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. 23.

  23. Reaction Rates C4H9Cl(aq) + H2O(l) C4H9OH(aq) + HCl(aq) • The reaction slows down with time because the concentration of the reactants decreases. 24.

  24. -[C4H9Cl] t Rate = = [C4H9OH] t Reaction Rates and Stoichiometry C4H9Cl(aq) + H2O(l) C4H9OH(aq) + HCl(aq) • In this reaction, the ratio of C4H9Cl to C4H9OH is 1:1. • Thus, the rate of disappearance of C4H9Cl is the same as the rate of appearance of C4H9OH. 25.

  25. Reaction Rates & Stoichiometry Suppose that the mole ratio is not 1:1? Example H2(g) + I2(g)  2 HI(g) 2 moles of HI are produced for each mole of H2 used. The rate at which H2 disappears is only half of the rate at which HI is generated 26.

  26. Concentration and Rate Each reaction has its own equation that gives its rate as a function of reactant concentrations. This is called its Rate Law The general form of the rate law is Rate = k[A]x[B]y Where k is the rate constant, [A]and [B]are the concentrations of the reactants. X and y are exponents known as rate orders that must be determined experimentally To determine the rate law we measure the rate at different starting concentrations. 27.

  27. Concentration and Rate NH4+ (aq) + NO2- (aq)  N2 (g) + 2H2O (l) Compare Experiments 1 and 2:when [NH4+] doubles, the initial rate doubles. 28.

  28. Concentration and Rate NH4+ (aq) + NO2- (aq)  N2 (g) + 2H2O (l) Likewise, compare Experiments 5 and 6: when [NO2-] doubles, the initial rate doubles. 29.

  29. Concentration and Rate NH4+ (aq) + NO2- (aq)  N2 (g) + 2H2O (l) This equation is called the rate law, andk is the rate constant. 30.

  30. The Rate Law • A rate law shows the relationship between the reaction rate and the concentrations of reactants. • For gas-phase reactants use PA instead of [A]. • k is a constant that has a specific value for each reaction. • The value of k is determined experimentally. Rate = K [NH4+][NO2- ] “Constant” is relative here- the rate constant k is unique for each reaction and the value of k changes with temperature 31.

  31. The Rate Law • Exponents tell the order of the reaction with respect to each reactant. • This reaction is First-order in [NH4+] First-order in [NO2−] • The overall reaction order can be found by adding the exponents on the reactants in the rate law. • This reaction is second-order overall. Rate = K [NH4+]1[NO2- ]1 32.

  32. [NO] mol dm-3 [O2] mol dm-3 Rate mol dm-3 s-1 0.40 0.50 1.6 x 10-3 0.40 0.25 8.0 x 10-4 0.20 0.25 2.0 x 10-4 Determining the Rate constant and Order The following data was collected for the reaction of substances A and B to produce products C and D. Deduce the order of this reaction with respect to A and to B. Write an expression for the rate law in this reaction and calculate the value of the rate constant. 33.

  33. Graphical Representation • http://www.chem.purdue.edu/gchelp/howtosolveit/Kinetics/IntegratedRateLaws.html A  B

  34. Zero Order For a zero order reaction: [A] versus t (linear for a zero order reaction) rate = k (k = - slope of line)

  35. First Order For a 1st order reaction: rate = k[A] (k = - slope of line)

  36. CH3NC CH3CN First-Order Processes Consider the process in which methyl isonitrile is converted to acetonitrile. How do we know this is a first order reaction? 37.

  37. CH3NC CH3CN First-Order Processes This data was collected for this reaction at 198.9°C. Does rate=k[CH3NC] for all time intervals? 38.

  38. First-Order Processes • When Ln P is plotted as a function of time, a straight line results. • The process is first-order. • k is the negative slope: 5.1  10-5 s-1. 39.

  39. Half-Life of a Reaction • Half-life is defined as the time required for one-half of a reactant to react. • Because [A] at t1/2 is one-half of the original [A], [A]t = 0.5 [A]0. 40.

  40. Half-Life of a First Order Reaction For a first-order process, set [A]t=0.5 [A]0 in integrated rate equation: NOTE: For a first-order process, the half-life does not depend on the initial concentration, [A]0. 41.

  41. First Order Rate Calculation Example 1: The decomposition of compound A is first order. If the initial [A]0 = 0.80 mol dm-3. and the rate constant is 0.010 s-1, what is the concentration of [A] after 90 seconds? 42.

  42. First Order Rate Calculation Example 1: The decomposition of compound A is first order. If the initial [A]0 = 0.80 mol dm-3. and the rate constant is 0.010 s-1, what is the concentration of [A] after 90 seconds? Ln[A]t – Ln[A]o = -kt Ln[A]t – Ln[0.80] = - (0.010 s-1 )(90 s) Ln[A]t = - (0.010 s-1 )(90 s) + Ln[0.80] Ln[A]t = -0.90 - 0.2231 Ln[A]t = -1.1231 [A]t = 0.325 mol dm-3 43.

  43. First Order Rate Calculations Example 2:A certain first order chemical reaction required 120 seconds for the concentration of the reactant to drop from 2.00 M to 1.00 M. Find the rate constant and the concentration of reactant [A] after 80 seconds. 44.

  44. First Order Rate Calculations Example 2:A certain first order chemical reaction required 120 seconds for the concentration of the reactant to drop from 2.00 M to 1.00 M. Find the rate constant and the concentration of reactant [A] after 80 seconds. Solution k =0.693/t1/2 =0.693/120s =0.005775 s-1 Ln[A] – Ln(2.00) = -0.005775 s-1 (80 s)= -0.462 Ln A = - 0.462 + 0.693 = 0.231 A = 1.26 mol dm-3 45.

  45. First Order Rate Calculations Example 3: Radioactive decay is also a first order process. Strontium 90 is a radioactive isotope with a half-life of 28.8 years. If some strontium 90 were accidentally released, how long would it take for its concentration to fall to 1% of its original concentration? 46.

  46. First Order Rate Calculations Example 3: Radioactive decay is also a first order process. Strontium 90 is a radioactive isotope with a half-life of 28.8 years. If some strontium 90 were accidentally released, how long would it take for its concentration to fall to 1% of its original concentration? Solution k =0.693/t1/2 =0.693/28.8 yr =0.02406 yr-1 Ln[1] – Ln(100) = - (0.02406 yr-1)t = - 4.065 t = - 4.062 . - 0.02406 yr-1 t = 168.8 years 47.

  47. [A] versus t (linear for a zero order reaction) • ln [A] versus t (linear for a 1st order reaction) • 1 / [A] versus t (linear for a 2nd order reaction)

  48. Second Order For a 2nd order reaction: rate = k[A]2 (k = slope of line)

  49. NO2(g) NO (g) + 1/2 O2(g) Determining Reaction OrderDistinguishing Between 1st and 2nd Order The decomposition of NO2 at 300°C is described by the equation: A experiment with this reaction yields this data: 50.

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