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UNIT 3

UNIT 3. Chemical Kinetics. Things to Review for Unit 3. Metric multipliers Molarity Natural logarithms ln (logarithms to the base e ). Reaction Rates. Some reactions are fast. Reaction Rates. Some reactions are slow. Reaction Rates. Some processes (or reactions) are really slow.

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UNIT 3

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  1. UNIT 3 Chemical Kinetics

  2. Things to Review for Unit 3 • Metric multipliers • Molarity • Natural logarithms ln (logarithms to the base e)

  3. Reaction Rates Some reactions are fast.

  4. Reaction Rates Some reactions are slow.

  5. Reaction Rates Some processes (or reactions) are really slow.

  6. Reaction Rates • What happens in a chemical reaction? • New bonds are formed and some existing bonds are broken. • This can only happen when the reactants get close enough together for intermolecular attractive forces to come into play. NO Cl + NOCl Cl2

  7. Factors Affecting Reaction Rates • The reaction rate is a measure of how much product is formed in a given time, so a reaction rate is affected by: • Concentration of reactants. • More reactants in a given volume mean more collisions per second and therefore more chances of forming product.

  8. Factors Affecting Reaction Rates • Temperature of the reaction system. • The higher the temperature of a system, the faster the reactants move and the more collisions per second . Also, the collisions are higher energy. Both increase the number of products formed in a given time.

  9. Factors Affecting Reaction Rates • The physical state of the reactants. • Gases can easily mix, collide and react. When a solid is involved in a reaction, only the surface of the solid will be involved in collisions. Increasing the area of the solid surface will increase the number of collisions. grain burns slowly grain dust can explode

  10. Factors Affecting Reaction Rates • The presence of a catalyst. Catalysts are agents that increase the rate of a reaction without themselves being used up. Metals such as Pd are catalysts. Enzymes are catalysts.

  11. Reaction Rate The average rate of a reaction is the change in concentration of a product or a reactant with time. For the reaction A  B, average rate = Δ[B] = - Δ[A] Δt Δt where [B] means the molarity of B. Δt = tfinal - tinitial

  12. Reaction Rate For the reaction A  B, The rate of formation of B is Δ[B] . Δt The rate of consumption of Ais -Δ[A] . Δt The rate of disappearance of A is- Δ[A] . Δt The negative sign makes the rate positive.

  13. Determining the Average Reaction Rate from Concentration versus Time Data For the reaction C4H9Cl(aq) + H2O(l)  C4H9OH(aq) + HCl(aq) average rate = Δ[C4H9OH] = Δ[HCl] = - Δ[C4H9Cl] ΔtΔt Δt If we had experimental data showing how the concentration of a reactant or product changed with time, we could calculate the average rate of the reaction.

  14. Determining the Average Reaction Rate from Concentration versus Time Data C4H9Cl(aq) + H2O(l)  C4H9OH(aq) + HCl(aq) For the interval 0-50 s, average rate = 0.0095 – 0 50.0 – 0 = 1.9 x 10-4mol/L-s -OR- average rate = - (0.0905 – 0.1000) 50.0 – 0 = 1.9 x 10-4mol/L-s same as M/s Is the rate different for the interval 50-100 s?

  15. Instantaneous Reaction Rate C4H9Cl(aq) + H2O(l)  C4H9OH(aq) + HCl(aq) instantaneous rate = d[C4H9OH] = d[HCl] = - d[C4H9Cl] dt dt dt = slope of the graph at time t rate at t=330s is the negative of the slope of this line rate at t=330s is the slope of this line

  16. Reaction Rate C4H9Cl(aq) + H2O(l)  C4H9OH(aq) + HCl(aq) • Instantaneous rates can be determined graphically. • The initial rate is often the rate we are most interested in knowing.

  17. Reaction Rates and Stoichiometry The reaction rate is related to the rate of disappearance of reactant or the rate of appearance of product. For the reaction aA  bB, the average reaction rate is average rate = 1Δ[B] = - 1Δ[A] b Δt a Δt This can be extended to multiple reactants and multiple products. this is the rate of formation of B

  18. Reaction Rates and Stoichiometry 2N2O5(g)  4NO2(g) + O2(g) For the interval 0-15 s and using the NO2(g) data, average rate of appearance of NO2: ΔNO2 = 0.500 - 0 Δt 15.0 - 0 = 3.33 x 10-2 mol/L-s average rate of reaction is: 1ΔNO2= 8.33 x 10-3 mol/L-s 4Δt Can we talk about the molarity of a gas? Yes, if we define the molarity of a gas as moles of gas per liter of the container. For an ideal gas: PV = nRT M = n = P V RT

  19. Reaction Rates and Stoichiometry 2N2O5(g)  4NO2(g) + O2(g) For the interval 0-15 s and using the N2O5 data, the average rate of disappearance of N2O5(g) is - ΔN2O5 = - (0.750 - 1.000) Δt 15.0 - 0 = 1.67 x 10-2 mol/L-s average rate of the reaction is - 1ΔN2O5= 8.33 x 10-3 mol/L-s 2Δt The average rate of reaction will be the same no matter which reactant or product you choose.

  20. How Concentration Affects Rate • Four factors affect the reaction rate: • concentration of reactants • temperature of the system • physical states of the reactants and products • presence of a catalyst We will focus on how the initial concentrations of reactants affect the initial reaction rate. Different reactions will be affected differently by initial concentrations.

  21. How Concentration Affects Rate NH4+(aq) + NO2-(aq)  N2(g) + 2H2O(l) Rate data is obtained experimentally. This data is for the reacting system at 25°C. *[x] means the concentration of x in molarity.

  22. How Concentration Affects Rate -The Method of Initial Rates NH4+(aq) + NO2-(aq)  N2(g) + 2H2O(l) When [NH4+] doubled, the initial rate doubled. (Note that the concentration of [NO2-] was held constant.) Initial rate = c' [NH4+]

  23. The Method of Initial Rates NH4+(aq) + NO2-(aq)  N2(g) + 2H2O(l) When [NO2-] doubled, the initial rate doubled. (Note that the concentration of [NH4+] was held constant.) Initial rate = c" [NO2-]

  24. How Concentration Affects Rate For the reaction: NH4+(aq) + NO2-(aq)  N2(g) + 2H2O(l) the experimental data at 25°C shows that Initial rate = c' [NH4+] = c" [NO2-] = k [NH4+][NO2-] We say that the rate law for this reaction is Rate = k [NH4+][NO2-] Since each concentration is raised to the first power, we say this rate is first order in each reactant. The overall rate is second order.

  25. Rate Laws If experimental data shows that the rate law for a reaction is Rate = k [A]m[B]n then we say this rate is mth order in Aandnth order in B, and that the overall rate is (m+n)th order. k changes with temperature!!! k is called the rate constant and will have units that allow the rate to be M/s.

  26. Rate Laws Rate laws MUST be determined experimentally! 2N2O5(g)  4NO2(g) + O2(g) Rate =k [N2O5] CHCl3(g) + Cl2(g) CCl4(g) + HCl(g) Rate =k [CHCl3][Cl2]1/2 H2(g) + I2(g) 2HI(g) Rate =k [H2][I2] Although the rate constants are all shown as k,the value of k for each rate law is different.

  27. Using the Method of Initial Rates to Determine Rate Laws - Example 2NO(g) + 2H2(g)  N2(g) + 2H2O(g) 1. Determine the rate law for the reaction. 2. Calculate the rate constant. 3. Calculate the rate when [NO] = 0.050 M and [H2] = 0.150 M.

  28. Using the Method of Initial Rates to Determine Rate Laws - Example 1. Determine the rate law for the reaction. Doubling [H2] doubles the rate (runs 1 and 2), so rate = c'[H2]. Doubling [NO] quadruples the rate (runs 2 and 3). Rate 3 = 4 = [NO(run 3)]n = 2n 2n= 4 n=2 rate = c"[NO]2 Rate 1 [NO(run 1)]n The rate law is rate = k[H2][NO]2

  29. Using the Method of Initial Rates to Determine Rate Laws - Example 2. Calculate the rate constant using any of Runs 1 - 3. Run #1: 1.23 x 10-3 M/s = k (0.100 M)(0.100 M)2 k = 1.23 x 10-3 M/s = 1.23 M-2s-1 (0.100 M)(0.100 M)2 Run #2: k = 1.23 M-2s-1 Run #3: k = 1.23 M-2s-1

  30. Using the Method of Initial Rates to Determine Rate Laws - Example 3. Calculate the rate when [NO] = 0.050 M and [H2] = 0.150 M. rate = k[H2][NO]2 = (1.23 M-2s-1 )(0.150 M)(0.050 M)2 rate = 4.6 x 10-4 M/s at a temperature of 720°C

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