Presentation on theme: "Chapter 12 Chemical Kinetics"— Presentation transcript:
1 Chapter 12 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
2 Reaction RateKinetics 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.
3 Continued….2NO2(g) 2NO(g) + O2(g)The concentration of the reactant (NO2) decreases with time and the concentration of the products (NO and O2) increase with time.A change can be positive (increase) or negative (decrease) leading to a positive or negative reaction rate. However, we will always define the rate as a positive quantity.
5 Continued….The concentration of NO2 decreases with time, [NO2] is a negative quantity. The reaction rate is a positive quantityThe concentration of reactants always decrease with time. The rate expression involving reactant will include a negative sign.
6 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 productionof NO of NO of O2)
7 Rate Laws 2NO2(g) 2NO(g) + O2(g) The reaction rate will depend on the concentrations of the reactantsRate = k[NO2]nThe 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)
8 Types of Rate LawsDifferential Rate Law: expresses how rate depends on concentration.Integrated Rate Law: expresses how concentration depends on time.
9 Continued….We consider reactions where the reverse reaction is unimportant, rate laws involve only concentrations of reactants.Differential and integrated rate laws for a given reaction are related in a well-defined way, the experimental determination of either of the rate laws is sufficient.Experimental convenience dictates which types of rate law is determined experimentally.
10 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.
11 Figure 12.3 A Plot of the Concentration of N2O5 as a Function of Time for the Reaction
12 Overall Reaction Order Sum of the order of each component in therate law.rate = k[H2SeO3][H+]2[I]3The overall reaction order is = 6.
13 First-Order Rate Law ln[A] = kt + ln[A]o For aA Products in a 1st-order reaction,Integrated first-order rate law isln[A] = kt + ln[A]o
14 Continued….ln[A] = -kt + ln[A]oThe equation shows how the concentration of A depends on time. If the initial concentration of A and the rate constant k are known, the concentration of A at any time can be calculated.The above equation is the equation of straight line of the form y = mx + b, where a plot of y versus x is a straight line with slope m and intercept b.
15 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:
17 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 constantFor a first-order reaction, the half-life does not depend on concentration.
18 Figure 12.5 A Plot of (N2O5) Versus Time for the Decomposition Reaction of N2O5
19 Second-Order Rate Law For aA products in a second-order reaction, Integrated rate law isA plot of 1/[A] versus t will produce a straight line with a slope equal to kThe above equation shows how [A] depends on time and can be used to calculate [A] at any time t, provided k and [A]o are known
20 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 becomest1/2 = half-life of the reaction, k = rate constant,Ao = initial concentration of AThe half-life is dependent upon the initial concentration.
21 Figure 12.6 (a) A Plot of In(C4H6) Versus t (b) A Plot of 1/(C4H6) Versus t
22 Zero-Order Rate Laws The rate law for a zero-order reaction is Rate = k[A]o = k(1) = kFor 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
23 In this case a plot of [A] versus t gives a straight line of slope –k. continued…[A] = -kt + [A]oIn this case a plot of [A] versus t gives a straight line of slope –k.[A] = [A]o/2, when t = t1/2[A]o/2 =-kt1/2 + [A]oSolving for t1/2 gives,t1/2 = [A]o/2k
24 Figure 12.7 A Plot of (A) Versus t for a Zero-Order Reaction
25 Rate Laws for Reactions with More Than One Reactant A + B + C ProductRate = k[A]n[B]m[C]pFor such reaction, concentration of one reactant remain small compared with the concentrations of the others. So the rate law reduce toRate = k`[A]nWhere, k` = k[B]m[C]op and [B]o>>[A]o and [C]o>>[A]oThe 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.
26 A Summary1. Simplification: Conditions are set such that only forward reaction is important.2. Two types of rate law: differential rate law and integrated rate law3. Which type? Depends on the type of data collected - differential and integrated forms can be interconverted.
27 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).
28 Reaction MechanismThe series of steps by which a chemical reaction occurs.A chemical equation does not tell us how reactants become products - it is a summary of the overall process.The purpose for studying kinetics is to learn as much as possible about the steps involved in a reaction.
29 Reaction Mechanism (continued) The reactionhas many steps in the reaction mechanism.The rate law for this reaction is known from experiment to beRate = k[NO2]2The balanced equation tells us the reactants, the products, the stoichiometry.NO2(g) + NO2(g) NO3(g) + NO(g)NO3(g) + CO(g) NO2(g) + CO2(g)
30 Often Used TermsIntermediate: 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 termolecularThe sum of the elementary steps must give the overall balanced equationThe mechanism must agree with the experimentally determined rate law.
31 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
32 Collision ModelKey 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)
34 Arrhenius EquationCollisions must have enough energy to produce the reaction (must equal or exceed the activation energy).Orientation of reactants must allow formation of new bonds necessary to produce products.
35 Figure 12.13 Several Possible Orientations for a Collision Between Two BrNO Molecules
36 Arrhenius Equation (continued) k = rate constant, A = frequency factorEa = activation energy, T = temperatureR = gas constantln(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)
37 Figure 12.14 Plot of In(k) Versus 1/T for the Reaction 2N2O5(g) ® 4NO2(g) + O2(g)
38 Arrhenius Equation (continued) Activation energy (Ea) can also be calculated from the values of k at only two temperaturesAt temperature T1, where the rate constant is k1,At temperature T2, where the rate constant is k2,Ea can be calculated from k1 and k2 at temperature T1 and T2
39 CatalysisCatalyst: A substance that speeds up a reaction without being consumed it self.Enzyme: A large molecule (usually a protein) that catalyzes biological reactions.Homogeneous catalyst: Present in the same phase as the reacting molecules.Heterogeneous catalyst: Present in a different phase than the reacting molecules.
40 Figure 12.15 Energy Plots for a Catalyzed and an Uncatalyzed Pathway for a Given Reaction
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