Chemical Kinetics - General Chemistry - Lecture Slides, Slides of Chemistry

Download Chemical Kinetics - General Chemistry - Lecture Slides and more Slides Chemistry in PDF only on Docsity! Chemical Kinetics • The area of chemistry that concerns reaction rates. • Goals: To understand the steps (reaction mechanism) by which a reaction takes place. Allows us to find ways to facilitate the reaction Reaction Rate Kinetics deals with the speed (rate) at which changes occur. The quantity that changes is the amount or concentration of a reactant or a product. Change in concentration (conc.) of a reactant or product per unit time. Rate = conc of A at time conc of A at time 2 1 2 1 t t t t        A t Continued…. The concentration of NO2 decreases with time, [NO2] is a negative quantity. The reaction rate is a positive quantity The concentration of reactants always decrease with time. The rate expression involving reactant will include a negative sign. Rate = change in [NO ] Time elapsed NO ] t 2 2    [ Rate = - [NO ] t 2  Continued…. 2NO2(g)  2NO(g) + O2(g) The rate can also be defined in terms of the products. In doing so we must take into account the coefficients in the balance equation for the reaction. Stoichiometry determines the relative rates. Rate of consumption = rate of production = 2(rate of production of NO2 of NO of O2)                [NO ] t [NO] t O ] t 2 2 2 [ Rate Laws 2NO2(g)  2NO(g) + O2(g) The reaction rate will depend on the concentrations of the reactants Rate = k[NO2] n The above expression shows how the rate depends on the concentration of reactants is called a rate law. k = rate constant (proportionality constant) n = rate order (integer or fraction including zero) Rate = - [NO ] t k[NO ] 2 2    n Method of Initial Rates Initial Rate: the “instantaneous rate” just after the reaction begins. The initial rate is determined in several experiments using different initial concentrations. Figure 12.3 A Plot of the Concentration of N2O5 as a Function of Time for the Reaction Overall Reaction Order Sum of the order of each component in the rate law. rate = k[H2SeO3][H +]2[I]3 The overall reaction order is 1 + 2 + 3 = 6. Continued…. • The reaction is first order in A if a plot of ln[A] versus t is a straight line. • The integrated rate law for a first order reaction can also be expressed in terms of a ratio of [A] and [A]o as follow: ln [A] [A] kt o  Figure 12.4 A Plot of In(N2O5) Versus Time Half-Life of a First-Order Reaction The time required for a reactant to reach half its original concentration is called the half-life of a reactant and is designated by the symbol t1/2. t1/2 = half-life of the reaction, k = rate constant For a first-order reaction, the half-life does not depend on concentration. ln [A] [A] kt when t = t then [A] = [A] 2 ln [A] [A] / 2 kt ln(2) = kt t = ln(2) k t = 0.693 k o 1 / 2 o o o 1 / 2 1 / 2 1 / 2 1 / 2     Half-Life of a Second-Order Reaction When one half-life of the second order reaction has elapsed (t = t1/2), by definition, [A] = [A]o/2 then the integrated rate law becomes t1/2 = half-life of the reaction, k = rate constant, Ao = initial concentration of A The half-life is dependent upon the initial concentration. 1 [A] / 2 kt + 1 [A] 2 [A] A] kt 1 [A] kt solving for t gives the expression t = 1 k[A] o 1 / 2 o o o 1 / 2 o 1 / 2 1 / 2 1 / 2 o       1 [ Figure 12.6 (a) A Plot of In(C4H6) Versus t (b) A Plot of 1/(C4H6) Versus t Zero-Order Rate Laws The rate law for a zero-order reaction is Rate = k[A]o = k(1) = k For a zero-order reaction, the rate is constant. It does not change with concentration as it does for first-order or second-order reactions. The integrated rate law for a zero-order reaction is [A] = -kt + [A]o Rate Laws for Reactions with More Than One Reactant A + B + C  Product Rate = k[A]n[B]m[C]p For such reaction, concentration of one reactant remain small compared with the concentrations of the others. So the rate law reduce to Rate = k`[A]n Where, k` = k[B]m[C]o p and [B]o>>[A]o and [C]o>>[A]o The value of n can be obtained by determining whether a plot of [A] versus t is linear (n = 0), a plot of ln[A] versus t is linear (n = 1), or a plot of 1/[A] versus t is linear (n = 2). The value of k` is determined from the slope. A Summary 1. Simplification: Conditions are set such that only forward reaction is important. 2. Two types of rate law: differential rate law and integrated rate law 3. Which type? Depends on the type of data collected - differential and integrated forms can be interconverted. A Summary (continued) 4. Most common: method of initial rates. 5. Concentration v. time: used to determine integrated rate law, often graphically. 6. For several reactants: choose conditions under which only one reactant varies significantly (pseudo first-order conditions). Often Used Terms • Intermediate: formed in one step and used up in a subsequent step and so is never seen as a product. (neither a reactant nor a product) • Molecularity: the number of species that must collide to produce the reaction indicated by that step. • Elementary Step: A reaction whose rate law can be written from its molecularity. uni, bi and termolecular • The sum of the elementary steps must give the overall balanced equation • The mechanism must agree with the experimentally determined rate law. Rate-Determining Step Multistep reaction often have one step that is much slower than all the others. Reactants can become products only as fast as they can get through this slowest step. The overall reaction can be no faster than the slowest or rate determining step. In a multistep reaction, it is the slowest step. It therefore determines the rate of reaction. Overall rate = k1[NO2] 2 NO + NO NO + NO slow(rate determining) NO + CO NO + CO fast 2(g) 2(g) 3(g) (g) 3(g) (g) 2(g) 2(g)   Collision Model Key Idea: Molecules must collide to react. However, only a small fraction of collisions produces a reaction. Why? Arrhenius: An activation energy (threshold energy) must be overcome to produce a chemical reaction. 2BrNO(g)  2NO(g) + Br2(g) Figure 12.13 Several Possible Orientations for a Collision Between Two BrNO Molecules Arrhenius Equation (continued) k = rate constant, A = frequency factor Ea = activation energy, T = temperature R = gas constant ln(k) = -Ea/R(1/T) + ln(A) A plot of ln(k) versus 1/T gives a straight line, slope = -Ea/R and intercept = ln(A) k Ae E RT   a / Figure 12.14 Plot of In(k) Versus 1/T for the Reaction 2N2O5(g) 2(g) + O2(g)