Chapter 13 Lecture 1 Chemical Kinetics

1. Chapter 13 Lecture 1 Chemical Kinetics • Kinetics = The area of chemistry that is concerned with reaction rates. • What do Kinetics tell us? • Stoichiometry tells us what and how many reactants/products there are • Thermodynamics tells us if the reaction is energetically favorable (-DH) • Kinetics describes how fast or slow a reaction occurs 2 H2 + O2 2 H2O DH = -242 kJ/mol This spontaneous reaction is so slow, it essentially doesn’t happen! • We need Kinetics to fully describe chemical reactions • Kinetics help us determine how a reaction occurs = Mechanism • Reaction Rates = Change in concentration (conc) of a reactant or product per unit time. • For generic reaction: aA + bB cC + dD This is an average rate over this time period. If you find the rate for one reactant or product, you’ve found them all!

2. Example: 2 NO2 2 NO + O2

3. [NO2] = mol/L of NO2. This is decreasing with time. • [NO] and [O2] are increasing with time • Rates are always positive for product ([A]2 > [A]1) • Rates are always negative for reactants ([A]2 < [A]1) • To let us work with only + numbers, we always put a (-) before the rate if it is describing a reactant • The rate changes over the course of the reaction • Instantaneous Rate at a point = slope of a tangent line at the point • At t = 100 s, what is the instantaneous rate?

4. Rates of Product Formation 2 NO2 2 NO + O2 • 2 NO produced for every 2 NO2 consumed • Rate same at any point in the reaction • Curve is same, only inverted • Rate [NO] at 250 s = 8.6 x 10-6 mol/L.s • 1 O2 produced for every NO2 consumed • Rate = ½ the rate of NO2 consumption at any point • Rate = ½ the rate of NO production at any point • Rate [O2] at 250 s = 4.3 x 10-6 mol/L.s • O2 curve has a different shape • We can write an equation of the rates from the balanced equation • We must be specific when we talk about a reaction rate. • Rate depends on which reactant or product • Rate depends on how long the reaction has been going

5. Rate Laws • Background • Chemical reaction are reversible. Products recombine to give reactant. 2 NO2 2 NO + O2 • When you start with only reactants, the reaction only goes one way • D[NO2] at this point depends on only the forward reaction • This makes equations simple • After some products are formed, the reaction goes both directions • D[NO2] at this point depends on both forward and reverse reactions • This makes equations complicated • We will consider reactions only at times and conditions that have only the forward reaction of any significance • Rate Law Basics • A Rate Law describes how concentrations are changing in a reaction • Rate = k[A]n • k = rate constant = different for each reaction • n = the order of the reactant A

6. Products do not appear in the rate law if we use conditions giving the forward reaction only • The order n must be determined experimentally; it can’t be written directly from the balanced equation • We must specify which species we are using [A] in each rate law • Differential Rate Law tells us how rate changes with concentration • This is what we will call the rate law • Integrated Rate Law tells us how concentrations depend on time • This can be derived from the differential rate law (and vice versa) • We will look at integrated rate laws later • What do we learn from rate laws? • They help us determine if a reaction is fast enough to be useful • They help us figure out the exact steps (mechanism) of a reaction • Slowest step determines the overall rate • Chemists speed up reactions by changing that step

7. Determining the Form of a Rate Law • We must do experiments to determine the order of each reactant in the rate law • Example: 2 N2O5 4 NO2 + O2 (g) • O2 gas escapes, so there is no reverse reaction • We can determine rate at various times from the data

8. Examine how the rate changes as concentration changes [N2O5] = 0.90 M rate = 5.4 x 10-4 mol/L.s [N2O5] = 0.45 M rate = 2.7 x 10-4 mol/L.s • How does the [A] ratio compare to the rate ratio? • If the ratios are the same, the order of the reactant is 1 • When the order is 1, we call this a First Order Reaction • The general expression for a first order reaction is:

9. The Method of Initial Rates • Initial rate = rate just after a reaction starts (t = 0). This is used because we don’t have to worry about reverse reaction. • We start the reaction several time with different initial concentrations of reactants. • If the rate changes in the same ratio as the concentration, n = 1 • If the rate changes as the square of the concentration ratio, n = 2 • If the rate changes as the cube of the concentration ratio, n = 3, etc… • If the rate doesn’t change at all, n = 0 • Example with 2 reactants: NH4+ + NO2- N2 + 2 H2O • First order in both reactants, and overall second order (n + m = 2) Rate = k[NH4+]n[NO2-]m Exp1Exp2 [NH4+] same, [NO2-] doubles rate doubles so m = 1 Exp2Exp3 [NH4+] doubles, [NO2-] same rate doubles so n = 1 Rate = k[NH4+]1[NO2-]1 = k[NH4+][NO2-]

10. Now we can actually find what k is from any of the experiments since we know that: Rate = k[NH4+][NO2-] • Example: Find rate law, k, and order of the reaction BrO3- + 5 Br- + 6 H+ 3 Br2 + 3 H2O