Download Understanding Reaction Kinetics: Rate Laws & Mechanisms for Multiple Reactant Reactions and more Study notes Chemistry in PDF only on Docsity! Integrated Rate Law for Reactions with More than one Reactant Example: aA + bB + cC products Rate = k[A]n[B]m[C]p Overall reaction order = n + m + p rxn order w.r.t. A = n w.r.t. B = m w.r.t. C = p How do we solve for all of these unknowns? Suppose we have large excess of [B]0 and [C]0 i.e. [B]0 >> [A]0 and [C]0 >> [A]0 This simplifies the rate law, which becomes pseudo-first order: Rate = k'[A]n where k' = k[B]0[C]0 We can then solve for k' and n in the usual way Determination of k' at several different [B]0 and [C]0 allows determination of m, p, and k • • • Why do we study reaction kinetics? To fully understand a chemical reaction, we must know more than just the identities of the reactants and the products. We must know: • if the reaction will occur (is it thermodynamically favorable?) • how fast it will occur (is it kinetically feasible?) For a complete understanding, we should also know how it occurs. Chapter #15 – Chemical Kinetics 15.1) Reaction Rates 15.2) Rate Laws: Introduction 15.3) Determining the Form of the Rate Law 15.4) Integrated Rate Law 15.5) Rate Laws: Summary 15.6) Reaction Mechanisms 15.7) The Steady-State Approximation 15.8) A Model for Chemical Kinetics 15.9) Catalysis Molecularity of an elementary step: The number of species (i.e. reactants) that must collide to produce the reaction indicated by that step NO2(g) + CO(g) NO(g) + CO2(g) • The overall reaction must proceed in multiple elementary steps. What have we learned about how this reaction occurs? k • The slowest step must involve collision of two NO2 molecules: NO2 + NO2 (something else) Kinetics expts show: Rate = k[NO2]2 NO2(g) + CO(g) NO(g) + CO2(g) The slowest elementary step determines the rate of the overall reaction. This step is therefore called the Rate Determining Step. k • The slowest step must involve collision of two NO2 molecules: NO2 + NO2 (something else) We would like to use these results to develop a detailed reaction mechanism. Necessary criteria for a valid proposed mechanism: 1. The sum of the elementary steps must equal the overall balanced equation for the reaction. 2. The mechanism must agree with the experimentally determined rate law. Sometimes multiple mechanisms exist that meet these two criteria. These criteria are therefore necessary but not sufficient to prove a mechanism. rate determining step k1 intermediate intermediate net reaction: NO2(g) + CO(g) NO(g) + CO2(g) Proposed Reaction Mechanism The proposed mechanism: NO2 + NO2 NO3(g) + NO(g) NO3(g) + CO(g) NO2(g) + CO2(g) Overall: NO2(g) + CO(g) NO(g) + CO2(g) k k' k k is SMALL (slow step) k' is LARGE (fast step) The predicted Rate = k[NO2]2 of the overall reaction is now consistent with experiment Note: The species NO3(g) is not a reactant or a product. It is "an intermediate." Example (cont): We now use this information to determine the rate law for the overall reaction: Rate of reaction = -d[H2]/dt = k2[H2][N2O2] (from 2nd, rate–limiting step of the multi-step mechanism) 2NO(g) + H2(g) N2O(g) + H2O(g) unknown concentration rxn intermediate k1[NO]2 = k-1[N2O2] + k2[N2O2][H2] Solving the Steady-State Equation for [N2O2] k1[NO]2 = [N2O2](k-1 + k2[H2]) [N2O2] = k-1 + k2[H2] k1[NO]2 rxn rate = k2[H2][N2O2] Substituting [N2O2] obtained above into the rxn rate expression: gives: k1[NO]2k2[H2] k-1 + k2[H2] Rate = To test this mechanism, we test the rate law by changing the concentrations of H2 and NO For large [H2], k2[H2] >> k-1 and Rate ≈ k1[NO]2 For large [NO], k2[H2] << k-1 and Rate ≈ k2[H2][NO]2 Steady-State Example (cont.): 2NO(g) + H2(g) N2O(g) + H2O(g) k1[NO]2k2[H2] k-1 + k2[H2] Rate = The limiting rate laws obtained from the steady-state approximation provide predictions that we can test
