2 I. Introduction A. Definition of Chemical Kinetics Rusting of Iron “The study of the speed or rate of reactions and the nanoscale pathways or processes by which reactants are transformed into products.B. Examples of Reactions and RatesRusting of IronCombustion ReactionC. Significance of Studying Kinetics
3 D. Factors Affecting Reaction Rate 1. Concentration of Reactants2. Temperature3. Presence of a Catalyst4. Surface Area of a Solid Reactant or Catalyst5. Properties of Reactants and Products
4 II. Understanding Reaction Rates A. Kinetic Molecular TheoryMatter composed of particles in constant motion.Increase in temperature increases particle’s kinetic energy.B. Collision TheoryFor a reaction to occur , reactant molecules must collide with the proper orientation and with an energy greater than some minimum value.Activation Energy (Ea) – minimum energy required for reaction to occur.
5 Importance of Orientation One hydrogen atom can approach another from any direction …Effective collision; the I atom can bond to the C atom to form CH3I… and reaction will still occur; the spherical symmetry of the atoms means that orientation does not matter.Ineffective collision; orientation is important in this reaction.
6 Distribution of Kinetic Energies At higher temperature (red), more molecules have the necessary activation energy.
7 C. Transition State Theory 1. Note Reaction Profile:CO(g) + NO2(g) CO2(g) + NO(g)
8 2. Note Ea for either forward or reverse reaction. “At a given temperature, the higher the energy barrier, the slower the reaction.”3. Transition State or “Activated Complex”“Transition structure (between reactants andproducts) which is always found on top of theenergy hill (energy of activation).4. Is reaction, exothermic or endothermic as written from left to right?
10 III. Rates of Reactions A. Definition 1. Reaction rate: expresses how much product is appearing or how much reactant in disappearing per unit time.2. Units for reaction rates:(Examples)Ms-1 or M/s or mol L-1s-1
11 For Reaction: A → PRate of Disappearance of “A” = - ΔA / Δt Rate of Formation of “P” = + ΔP / Δt
12 B. Example Average Rate Determination For Rxn of Cisplatin and Water:H2O + Pt(NH3)2Cl2 Pt(NH3)2Cl(H2O) + Cl-CisplatinTime (min) [Cisplatin] (mol/L)What is average rate of disappearance of cisplatin(in mol L-1 min-1) for the first 200 min?(in mol L-1 min-1) for the next 200 min?
13 C. Instantaneous Rate Determination Significance of Measuring Instantaneous Rates!!
14 D. Reaction Rates and Stoichiometry Given the reaction: 2 N2O5(g) 4 NO2(g) O2 (g)If the rate of NO2 formation is mol L-1 s-1:1. What is the rate of disappearance of N2O5?2. What is the rate of formation of O2?
15 IV. Concentration and Rxn Rate A. Rate Law Equation:1. An equation that relates the rate of areaction to the concentrations of reactants(and catalyst) raised to various powers.2. Must be experimentally determined!!
16 3. For reaction:A + B CRate is proportional to reactant concentrationsRate = k[A]m [B]n k = rate constant(exponents “m” and “n” must be experimentally determined).4. For reaction:2 NO2(g) + F2(g) 2 NO2F(g)(experimentally determined Rate Law is:)Rate = k [NO2]1[F2]1 = k [NO2][F2]Exponents not necessarily same as rxn coefficients!!
17 4. For hypothetical reaction: 2 A(g) + B2(g) 2 AB(g)(experimentally determined Rate Law is:)Rate = k [A]2Not all reactants necessarily show up in the rate law equation!!
18 For the general equation: aA + bB pP The rate equation is: B. Reaction OrderFor the general equation: aA + bB pPThe rate equation is:Rate = k [A]m [B]nm and n are experimentally determined and are usually integers (0, 1, 2, 3, …). They may be fractions.
19 This reaction is said to m th order with respect to A and n th order with respect to B.The overall reaction order is the sum of the individual orders, orOverall Reaction Order = m + n
20 For the following reaction: 2 NO(g) + Cl2(g) 2 NOCl(g) Example:For the following reaction:2 NO(g) + Cl2(g) 2 NOCl(g)The observed rate law is:Rate = k [NO]2 [Cl2]What is the reaction order with respect to NO? What is the reaction order with respect to Cl2? What is the overall reaction order?How would the rate of the reaction be affected by doubling the concentration of both NO and Cl2?
21 C. Determination of Rate Law Exponents Done experimentally by measuring initial rates for several different known concentrations of reactants.Consider the reaction:2 NO(g) H2(g) N2(g) H2O(g)Given the information on the next slide:1. Determine the rate law.2. What is the order of the reaction?3. What is the value of the rate constant? “units”
22 Initial Concentration (M) Rate Experiment [NO] [H2] mol / L.s10.1001.23 x 10-320.2002.45 x 10-334.93 x 10-340.3001.11 x 10-2
23 1. Determine the rate law. a. Determine general form of rate law. rate = k [NO]m[H2]nb. Determine exponents.For each reactant, compare two experiments or trials where its concentration is changing and all other reactant concentrations are held constant.
24 Methods For Determining Rate Law Exponents Method 1 – solve analyticallySubstitute data into rate law and compareDivide equation with larger rate by eq. withsmaller rate. Cancel terms and solve.or simplifyingMethod 2 - solve by inspectionHow does changing conc. affect rate?
25 2. Determine order of the reaction. 3. Determine rate constant (include units).Use rate law and either set of data.4. What is the rate of the reaction when[NO] = [H2] = M ?
26 V. Integrated Rate LawEquations derived from rate law (by using calculus) which are convenient for solving concentration versus time problems.For:1. First Order Rxns – only one covered2. Second Order Rxns3. Zero Order Rxns
27 A. Integrated First Order Rate Law For reaction: aA Productrate = - Δ[A] / Δt = k[A]and using calculus:where [A]o is the concentration of A at time zero (t = 0) and [A]t is the concentration at time t.*** A and A0 may be replaced by quantities that are proportional to concentration !!!!
28 The sugar, sucrose, will undergo the following B. ProblemThe sugar, sucrose, will undergo the following(first order) hydrolysis reactionC12H22O H2O C6H12O6 + C6H12O6sucrose glucose fructosewith a rate constant of 6.2 x 10-5 s-1 at 35oC.A sample of 0.20 mol of sucrose was initiallydissolved in a total volume of 500 mL.1. What is the sucrose conc. after 2 hours?
29 2. What will be the glucose concentration 2. What will be the glucose concentration after the 2 hours have elapsed?3. How many minutes will it take for thesucrose concentration to drop to 0.30M?
30 C. Half-Life and First Order Reactions 1. Definition (t1/2) – the time required for the concentration of a reactant to fall to one half its initial value.2. Significance:Useful in describing radioactive decay ratesUseful in describing rates of 1st order reactions3. Equation derived from Integrated Rate Law
32 4. Problem - Radioactive Iodine-131 has t1/2 of 8 days. If you had a sample of 10,000 I-131atoms initially, how many I-131 atoms wouldremain after 32 days?
33 experimental conditions. 5. Problem – The first order hydrolysis reaction ofsucroseC12H22O H2O C6H12O6 + C6H12O6sucrose glucose fructosehas a rate constant of 6.17 x 10-4 s-1 underexperimental conditions.a. What is the half-life for the hydrolysis ofsucrose?b. How many minutes are required for 75% of the initial sucrose to react?
34 VI. Temperature and Rxn Rate A. Nanoscale Explanation as to why increasing temperature increase reaction rate.B. Mathematical RelationshipArrhenius Equation(not responsible for problem solving)
35 VII. Reaction Mechanisms A. IntroductionReaction mechanism is a series of elementary reactions or simple steps whose overall effect is given by the net chemical reaction (equation).1. It is a theory of how the reaction occurs which is based on experimental data.2. Cannot be absolutely proven.3. Steps must be elementary reactions.
36 1. Definition- the simplest step in what is B. Elementary Reaction1. Definition- the simplest step in what isoften a multi-step mechanism for anobserved chemical reaction.a. The equation for an elementary reaction shows exactly which molecules, atoms, or ions take part in the elementary reaction.b. For an elementary reaction, the rate law is directly determined from the elementary reaction.
37 2. Elementary Reactions in Mechanisms Typesa. Unimolecular Reaction – structure of a single particle (atom, molecule, or ion) rearranges to produce a different particle or particles.b. Bimolecular Reaction - two particles (atoms, ions, or molecules) collide and rearrange into products.c. Termolecular Reaction – (less likely)
38 3. ProblemsIdentify the type of elementary reaction and give the rate law for the following elementaryreactions:a. Cl + Cl Cl2b. N2O5 NO2 + NO3
39 C. Properties of Valid Mechanisms 1. Must consist of only unimolecular, bimolecular, or termolecular elementary reactions. (True for any mechanisms given toyou.)***2. Sum of the elementary reactions should be equal to the overall reaction equation.***3. Should predict the experimentally observed rate law.The overall rate of the reaction is dependent on the slowest step in the mechanism - the rate-limiting step.
40 D. Example Mechanism Problems 1. For the overall reaction:2 NO2Cl 2 NO2 + Cl2the following mechanism is proposed.NO2Cl NO2 + Cl (slow)NO2Cl + Cl NO2 + Cl2 (fast)a. Does the sum of the elementary processes equal the overall reaction?b. What is the rate law for the overall rxn?c. Identify any reaction intermediates.
41 2. For the overall reaction: (CH3)3CCl + OH- (CH3)3COH + Cl- there are two proposed mechanisms.1) Concerted mechanism2) Two Step Mechanism(CH3)3CCl (CH3)3C Cl (slow)(CH3)3C+ + OH - (CH3)3COH (fast)
42 1. Which is the correct mechanism? Why? From kinetic data, the correct rate law for the overall reaction is:rate = k[(CH3)3CCl]Questions:1. Which is the correct mechanism? Why?2. What is the order of the overall reaction?3. Identify reaction intermediates in eachproposed mechanism.
43 VIII. Catalysts A. Definition Not shown in overall reaction. Catalyst – a species that increases the rate of an overall reaction but is not consumed in the reaction.Not shown in overall reaction.Will show up in rate law for catalyzed reaction.
44 B. How Do They Increase Reaction Rate? 1. Catalysts alter / participate in rxn mechanism.2. Lower activation energy. Speeds reaction.3. See Fig (next slide)4. Catalyst changes kinetics, but notthermodynamics of reaction.Increases speed of reaction.Does not change net energy of reaction, type of product produced, or direction of reaction.
46 C. Homogeneous vs. Heterogeneous Catalysts Heterogeneous Catalyst- catalyst in different phase from reacting substance.(hydrogenation of vegetable oils – next slide)(catalytic converters in autos)Homogeneous Catalyst – catalyst in same phase as reacting substance.(enzymes)
47 Heterogeneous Catalysis Hydrogen is adsorbed onto the surface of a nickel catalyst. A C=C approaches …… and is adsorbed.Hydrogen atoms attach to the carbon atoms, and the molecule is desorbed.
48 D. Enzymes (Homogeneous Catalysts) 1. Protein that catalyzes reaction. 2. Most efficient catalysts known to man.3. Specifically binds to reactants (substrates),holding them in correct position for reaction to occur.4. Lower activation energy by stabilizingtransition state or altering mechanism.5. Examples: Lysozyme Next slide