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Chemical Kinetics

Chemical Kinetics. An Introduction. CO(g) + NO 2 (g)  CO 2 (g) + NO. H 2 O 2 (aq)  H 2 O(l) + O 2 (g). S 2 O 8 2- + 2 I -  I 2 + 2 SO 4 2-. Questions about reactions. ?. What’s happening?. The chemical equation.

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Chemical Kinetics

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  1. Chemical Kinetics An Introduction CO(g) + NO2(g)  CO2(g) + NO H2O2(aq)  H2O(l) + O2(g) S2O82- + 2 I- I2 + 2 SO42-

  2. Questions about reactions ? • What’s happening? The chemical equation 2 IO3- + 5 HSO3-→ I2 + 5 SO42- + 3 H+ + H2O 2. How fast is it happening? Kinetics H2O2(aq) → H2O(l) + O2(g) -- slow H2O2(aq) + catalyst → H2O(l) + O2(g) -- FAST 3. To what extent does it occur? Equilibrium Chapters 15-18 HC2H3O2 + H2O C2H3O2- + H3O+

  3. Chemical Kinetics The study of reaction rates and the sequence of steps by which a reaction occurs A Definition

  4. Rates of Reaction • Very fast • Explosions, neutralizations • Almost instantaneous • Medium • Cooking, rusting • minutes to years • Slow • Formation of diamonds, decay of 235U • Up to millions of years

  5. Expressing Reaction Rates Speed of a car: mph = For a chemical reaction we want to track concentration of products or reactants over time: Concentration of reactants decreases Concentration of products increases

  6. Expressing Reaction Rate Cont’d For the reaction A → B Rate of reaction = Or… Note the sign: A is disappearing

  7. Δ[B] Δ[A] 1 1 = - = - b a Δt Δt Δ[D] Δ[C] 1 1 = = d c Δt Δt General Rate of Reaction a A + b B → c C + d D Rate of reaction = rate of disappearance of reactants = rate of appearance of products

  8. An Example C2H4 (g) + O3 (g) C2H4O(g) + O2 (g) [C2H4] t [O3] t Rate = - = - Time(s) [O3](mol/L) [O3] t 0.0 3.20x10-5 10.0 2.42x10-5 20.0 1.95x10-5 30.0 1.63x10-5 40.0 1.40x10-5 50.0 1.23x10-5 60.0 1.10x10-5 Rate = - (1.10x10-5mol/L) - (3.20x10-5mol/L) Rate = - = -3.50 x 10-7mol/L·s 60.0 s - 0.0 s The reaction between ethylene and ozone:

  9. This rate is the average rate for a time period Does not show that rate is changing with time Does not show rate at a given instant

  10. Plot of [O3] vs. Time

  11. But Wait! There’s More… Rate is dependent on concentration! Can see this experimentally

  12. Plot of [C2H4] and [O2] vs. Time

  13. Instantaneous rate • Use smaller and smaller increments of time • The slope of a tangent line to the curve at any point is the instantaneous rate • Note that reaction rate usually refers to the instantaneous rate

  14. Plot of [O3] vs. Time

  15. Reaction Rate Law For a chemical reaction: aA + bB + . . . →cC + dD + . . . The rate law for the forward reaction has the form: Rate = k [A]m[B]n . . . • k = the reaction rate constant • exponents m & n are the reaction orders • defines how rate is affected by concentration For example, if the rate doubles when the concentration of A doubles, the rate depends on [A]1, so m = 1; if the rate quadruples when the concentration of B doubles, the rate depends on [B]2, so n = 2. More on this next week

  16. So, what affects reaction rate?  Concentration Molecules must collide in order to react. Reaction rate is proportional to the concentration of reactants. Rate = k (collision frequency) = k (concentration)

  17. Factors Affecting Reaction Rate  Physical state Molecules must mix in order to collide. The physical state (solid, liquid, gas) will affect frequency of collisions, as well as the physical size of droplets (liquid) or particles in the case of solids.

  18. Factors Affecting Reaction Rate  Temperature Molecules must collide with enough energy to react. Raising the temperature increases the reaction rate by increasing the number of collisions per time unit, and especially, the energy of the collisions.

  19. Factors Affecting Reaction Rate  Nature of the reactants Some species are more reactive than others. You have seen this with the periodicity of reactivity. For example the reactivity of the group 1 metals.

  20. Factors Affecting Reaction Rate  Presence of a catalyst Catalysts can provide alternate, lower energy, reaction pathways. Catalysts generally reroute the pathway of a chemical reaction so that this “alternate” path, although perhaps more circuitous, has a lower activation energy for reaction than the un-catalyzed reaction.

  21. Lab This Week • Explore several of the factors which affect reaction rate • Perform in any order • Perform either B or C • CHECK REAGENT CONCENTRATIONS! • Caution with strong acids • Use waste beakers • Part A.1 use 2M H3PO4 • Part D—Note H2O2 decomposes slowly on its own.

  22. Evidences of a Chemical Reaction • Color change • Precipitate formation • Temperature change • Gas evolution • Formation of a weak electrolyte

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