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Chapter 14

Chapter 14. Rates of Reactions Kinetics. I. I. Introduction. A) Demonstrations. B) Chemical Kinetics is the study of the rates (speeds) of chemical reactions and the mechanisms of chemical reactions.

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Chapter 14

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  1. Chapter 14 Rates of Reactions Kinetics

  2. I I. Introduction A) Demonstrations B) Chemical Kinetics is the study of the rates (speeds) of chemical reactions and the mechanisms of chemical reactions.

  3. C) The rate of a chemical reaction is a measure of how fast reactants are consumed and/or how fast products are made. D) The mechanism of a reaction is a detailed description of the way a reaction occurs. It is a sequence of elementary steps which lead from reactants to products.

  4. Mechanisms can be proven wrong through _____________, but they can never be called _________________________ since they are, in general, educated guesses.

  5. Practical reasons for studying kinetics: Some reactions we would like to speed up:drug delivery, paint drying, destruction of air pollutants in auto exhaust breakdown of materials in landfills. Some reactions we would like to slow down: food decay, rubber decay, human aging, destruction of the ozone layer, rusting, corrosion

  6. E) Some reactions take place in a fraction of a second and other take many years. What variables affect the reaction rate? 1) The characteristics of the reactants and the products.

  7. 2) The concentration of the reactants – in some reactions the rate is unaffected by the concentration of one of the reactants as long as it is there in some amount.

  8. 3) The presence of a catalyst, a substance that … 4)The temperature at which the reaction occurs. Increasing the temperature usually increases the rate. A general rule is … 5) The surface area of a solid reactant or catalyst affects the rate.

  9. G) Reaction rate is the increase in molar concentration of product of a reaction per unit time or the decrease in molar concentration of a reactant per unit time.

  10. For example, for the reaction: 2 N2O5 4 NO2 + O2 The following table shows the concentration of N2O5 as a function of time at 45 oC.

  11. 1) The rate of reaction can be written in the form: 2) Usually the rate is a rapidly changing quantity, as the reaction proceeds the reactants are used up and there remains less and less material to undergo reaction.

  12. 3)Generally we obtain for a reactant a curve which resembles the one below.

  13. If we take the first 2 points from the table above, we can find as avg. rate of decomposition of N2O5.

  14. You should be able to see that the avg. rate is decreasing, hence the curve. Even though we obtain a negative value for the rate, since N2O5 is decreasing, rates are reported as positive values.

  15. Look at the graph again, And make the time interval smaller and smaller, we can obtain an instantaneous rate.

  16. The instantaneous rate is equal to the slope of the line at that point. Calculus??? To what is the slope of the line at that point equal?

  17. 4) How is data obtained for a concentration curve? a) Monitor a color change. b) Measure pressure if a gas is produced. c) Monitor a change in pH if an acid or base reaction.

  18. 5) A look at the change in rate over time for another reaction: 2 NO2(g) 2 NO(g) + O2(g) at 300oC

  19. H) What does the balanced equation tell us about rates? The equation we will look at is: H2(g) + I2(g) ----> 2 HI(g) H2 and I2 must disappear at the same rate since 1 molecule of H2 reacts with 1 molecule of I2.

  20. In the same amount of time, 2 molecules of HI must appear.The rate of appearance of HI must equal twice the rate of disappearance of H2 (g) and I2(g).

  21. The rate of disappearance of H2(g) = the rate of disappearance of I2(g) = ½ the rate of appearance of HI.

  22. In general, for the equation: aA + bB -----> cC + dD To obtain a rate equation of the rates of the substances in relation to each other we divide through by the coefficients.

  23. But we usually want the first reagent in terms of all the others, so we multiply by the coefficient of the first reagent.

  24. II. The Rate Law (Rate expression), Rate Constant Order of Reaction A) The following equation has been studied in the gaseous state and the data at 250 K may be summarized as follows:

  25. F2(g) + 2 ClO2(g) ---> 2 ClO2F(g) B) From this data the answers to the following questions can be obtained:

  26. 1) What is the rate law of the reaction? 2) What is the order of the reaction? 3) What is the value of the rate constant k? 4) What is the rate of formation of ClO2F when [F2]is 0.010 mol/L and[ClO2] is 0.020 mol/L?

  27. C) What is a rate law? 1) In 1864 it was discovered that the rate of a reaction is proportional to some power of the concentration of reactants at constant temperature.

  28. 2) In general, for the equation: aA + bB  cC + dD

  29. rate law = rate equation = rate expression k is the specific rate constant which is independent of concentration.

  30. k depends on the nature of the reactants; fast reactions have large k's and slow reactions have small k's ordinarily k ________________ with temperature.

  31. THE EXPONENTS x AND y MUST BE EXPERIMENTALLY DETERMINED. They are not automatically obtained from the balanced equation.

  32. Some experimentally determined rate laws for equations are as follows: a) 2 N2O5(soln) 2 N2O4(soln) + O2(g)

  33. Notice that the exponent is NOT 2, the coefficient in the balanced equation, but has been experimentally determined to be 1. The reaction order with respect to a given species equals the exponent of the concentration of that species in the rate law as determined experimentally.

  34. The order of the above reaction with respect to N2O5 is 1. The order of a reaction is the sum of the exponents of the reactants in a rate law. The order of the above reaction, since 1 is the only exponent, is 1 as well. It is a first order reaction.

  35. This means that when we double the concentration of N2O5, we double the rate of reaction OR if we halve the concentration, we halve the rate. b) For the reaction: 2 NO + O2 2 NO2 What is the situation here?

  36. D) The answers to the questions then are obtained in the following manner. The rate law for F2(g) + 2 ClO2(g) ---> 2 ClO2F(g) will look like the following.

  37. Your job is to find the values of x and y from the experimental data.

  38. You need to do two division problems to find x and y. What is the order of the reaction with respect to F2? _______ What is the order of the reaction with respect to ClO2? ______ What is the overall order of the reaction? __________

  39. To find k, the rate constant, we take the experimentally determined rate law, put the data in from one of the experiments and solve for k. The units of k are important. k = _________________

  40. To find the rate of formation of ClO2F when [F2] is 0.010 M and the [ClO2] is 0.020 M, we need to look at the relationship between the rate of formation of [ClO2], and the rate of disappearance of F2. Rate = __________________________

  41. Transparency – in class problem – collect it

  42. IV. A graphical method is often used to show the order of a reaction, or from a graph we can obtain the order of a reaction. • For the general reaction of A ---> products • If the reaction is first order, we can write

  43. When we divide both sides by [A] and multiply both sides by dt we obtain the following: Those of you who have had calculus should recognize this as:

  44. Which can be changed to the following in log to the base 10.

  45. The above equations can be rewritten in a more familiar form.


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