A B Reaction Rates [A]& [B] determined by monitoring the change in concentration of either reactants or products as a function of time. [A] [B] -[A] t [B] t Rate = = {RxRateIntro} Spectrometer

Reaction Rates C4H9Cl(aq) + H2O(l)  C4H9OH(aq) + HCl(aq) -[A] t butyl chloride butanol -[A] t -[A] t -[A] t [B] t Rate = =

Reaction Rates C4H9Cl(aq) + H2O(l)  C4H9OH(aq) + HCl(aq) The average rate of the reaction over each interval is the change in concentration divided by time: Ave. rate = -[C4H9Cl] t Note that the average rate decreases as the reaction proceeds. This is because as the reaction goes forward, there are fewer collisions between reactant molecules.

Reaction Rates C4H9Cl(aq) + H2O(l)  C4H9OH(aq) + HCl(aq) Concentration vs. Time Graph All reactions slow down over time. Therefore, the best indicator of the rate of a reaction is the instantaneous rate near the beginning. The slope of a line tangent to the curve at any point is the instantaneous rate at that time.

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. Rate = -[C4H9Cl] t = [C4H9OH]

Reaction Rates and Stoichiometry What if the ratio is not 1:1? 2 HI(g)  H2(g) + I2(g) How do rates compare? Rate1 = − [HI] t Rate2= + [I2] ≠ Rate = − 1 2 [HI] t = [I2] Rate = − [HI] t = 2 [I2] To generalize, then, for the reaction 5 A + 2 B 4 C + 3 D

Practice Problems

Measuring rate of reaction Two common ways: Measure how fast the products are formed Colour change Gas formation Measure how fast the reactants are used up

Slows down as time passes as reactants are used up Rate of reaction graph Amount of product formed Slows down as time passes as reactants are used up Reaction begins rapidly Stops after a certain time – graph levels out the reactants have been used up. Time

How does Concentration affect Rate? NH4+(aq) + NO2−(aq) N2(g) + 2 H2O(l) The data demonstrates: Rate  [NH4+] Rate  [NO2−] Rate  [NH4+] [NO2−] Data shows the relationship between the reaction rate and the conc. of reactants. This equation is the rate law, and k is the rate constant @ particular temp. or Rate = k [NH4+] [NO2−]

Concentration and Rate Compare Experiments 1 and 2: when [NH4+] doubles, the initial rate doubles. Likewise, compare Experiments 5 and 6: when [NO2- ] doubles, the rate doubles.

Concentration and Rate adding the exponents on the reactants in the rate law. This reaction is second-order overall. This equation is called the rate law, and k is the rate constant. The value of k is determined experimentally. Constant is relative here- k is unique for each rxn and k changes with T

Determination of Rate Law using Initial Rate Reaction: S2O82-(aq) + 3I- (aq) → 2SO42- (aq) + I3- (aq)  Expt. [S2O82-] [I-] Initial Rate, # (mol/L) (mol/L) (mol/L.s) 1 0.036 0.060 1.5 x 10-5 2 0.072 0.060 2.9 x 10-5 3 0.036 0.120 2.9 x 10-5 Calculate x by keeping y constant Rate = k[S2O82-]x[I-]y Rate Line2 Rate Line3 X

Determination of Rate Law using Initial Rate (b) Calculation of rate order, y by keeping x constant: This reaction is first order w.r.t. [S2O82-] and [I-] Rate = k[S2O82-][I-]

Calculating rate constant and rate at different concentrations of reactants Rate constant, k = = 6.6 x 10-3 L.mol-1.s-1 If [S2O82-] = 0.20 M, [I-] = 0.20 M, and k = 6.6 x 10-3 L.mol-1.s-1 Rate = (6.6 x 10-3 L.mol-1.s-1)(0.20 mol/L)2 = 2.6 x 10-4 mol/(L.s)

You try Overall order of reaction: 6 Rate = k [H2]1[I2]1

O2 + 2 NO  2NO2

Reaction Mechanisms Reactions may occur all at once or through several discrete steps. Each of these processes is known as an elementary reaction or elementary process.

Reaction Mechanisms The molecularity of a process tells how many molecules are involved in the process. The rate law for an elementary step is written directly from that step. Which , one must agree, is far easier than calculating it from experiment

Multistep Mechanisms In a multistep process, one of the steps will be slower than all others. The overall reaction cannot occur faster than this slowest, rate-determining step.

Slow Initial Step NO2 (g) + CO (g) NO (g) + CO2 (g) The rate law for this reaction is found experimentally to be Rate = k [NO2]2 CO is necessary for this reaction to occur, but the rate of the reaction does not depend on its concentration. This suggests the reaction occurs in two steps.

Slow Initial Step A proposed mechanism for this reaction is Step 1: NO2 + NO2 NO3 + NO (slow) Step 2: NO3 + CO NO2 + CO2 (fast) The NO3 intermediate is consumed in the second step. As CO is not involved in the slow, rate-determining step, it does not appear in the rate law.

Example 2 Fast Initial Step What is the rate law? The rate of the overall reaction depends upon the rate of the slow step.

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