Presentation on theme: "Chapter 13 Chemical Kinetics"— Presentation transcript:
1Chapter 13 Chemical Kinetics CHEMISTRYChapter 13 Chemical Kinetics
2Factors that Affect Reaction Rates Kinetics is the study of how fast chemical reactions occur.There are 4 important factors which affect rates of reactions:reactant concentration,temperature,action of catalysts, andsurface area.Goal: to understand chemical reactions at the molecular level.
3C4H9Cl(aq) + H2O(l) C4H9OH(aq) + HCl(aq) Reaction RatesChange of Rate with TimeConsider:C4H9Cl(aq) + H2O(l) C4H9OH(aq) + HCl(aq)
5C4H9Cl(aq) + H2O(l) C4H9OH(aq) + HCl(aq) Reaction RatesChange of Rate with TimeC4H9Cl(aq) + H2O(l) C4H9OH(aq) + HCl(aq)We can calculate the average rate in terms of the disappearance of C4H9Cl.The units for average rate are mol/L·s or M/s.The average rate decreases with time.We plot [C4H9Cl] versus time.The rate at any instant in time (instantaneous rate) is the slope of the tangent to the curve.Instantaneous rate is different from average rate.We usually call the instantaneous rate the rate.
6Reaction RatesFor the reaction A B there are two ways of measuring rate:the speed at which the products appear (i.e. change in moles of B per unit time), orthe speed at which the reactants disappear (i.e. the change in moles of A per unit time).
7Reaction Rates Reaction Rate and Stoichiometry In general for: aA + bB cC + dD
8Rate is (-) if reagent is consumed. Rate is (+) if compound is produced.Rate will ultimately be (+) because change in concentration will be negative. Two (-)’s become (+).
13Sample ProblemA. How is the rate at which ozone (O3) disappears related to the rate at which O2 appears in the reaction: 2 O3 (g) 3 O2 (g)?B. If the rate at which O2 appears, D[O2]/Dt, is 6.0 x 10-5 M/s at a particular instant, at what rate is O3 disappearing at this same time, D[O3]/Dt?
14AnswersA. “Related to” means compare, so write the rate expression comparing the compounds.B x 10-5 M/s
15Sample ProblemThe decomposition of N2O5 proceeds according to the following equation: N2O5 (g) 4 NO2 (g) O2 (g)If the rate of the decomposition of N2O5 at a particular instant in a reaction vessel is 4.2 x 10-7 M/s, what is the rate of appearance of: (a) NO2, (b) O2 ?
16Differential Rate Law Example: Rate = k[A]n - is a rate law that expresses how rate is dependent on concentrationExample:Rate = k[A]n
17Differential First Order Rate Law First Order ReactionRate dependent on concentrationIf concentration of starting reagent was doubled, rate of production of compounds would also double
18Concentration and Rate Using Initial Rates to Determines Rate LawsA reaction is zero order in a reactant if the change in concentration of that reactant produces no effect.A reaction is first order if doubling the concentration causes the rate to double.A reacting is nth order if doubling the concentration causes an 2n increase in rate.Note that the rate constant does not depend on concentration.
19Differential Rate Law Rate = k[A]n Rate = k[A]n[B]m For single reactants: A CRate = k[A]nFor 2 or more reactants: A + B CRate = k[A]n[B]mRate = k[A]n[B]m[C]p
20Concentration and Rate Exponents in the Rate LawFor a general reaction with rate lawwe say the reaction is mth order in reactant 1 and nth order in reactant 2.The overall order of reaction is m + n + ….A reaction can be zeroth order if m, n, … are zero.Note the values of the exponents (orders) have to be determined experimentally. They are not simply related to stoichiometry.
21Concentration and Rate In general rates increase as concentrations increase.NH4+(aq) + NO2-(aq) N2(g) + 2H2O(l)
22Concentration and Rate Rate law:The constant k is the rate constant.
23ProblemNH4+ + NO2- N2 + 2H2OGive the general rate law equation for rxn.Derive rate order.Derive general rate order.Solve for the rate constant k.
24To Determine the Orders of the Reaction (n, m, p, etc….) 1. Write Rate law equation.2. Get ratio of 2 rate laws from successive experiments.Ratio = rate Expt.2 = k2[NH4+]n[NO2-]mrate Expt k1[NH4+]n[NO2-]m3. Derive reaction order.4. Derive overall reaction order.
25Experimental Data Expt. [NH4]initial [NO2-]initial Initial Rate 1 30.200 M5.40 x 10-7
26Initial Rate (mol·L-1·s-1) A + B CExperiment Number[A] (mol·L-1)[B] (mol·L-1)Initial Rate (mol·L-1·s-1)10.1004.0 x 10-520.20031.6 x 10-4Determine the differential rate lawCalculate the rate constantCalculate the rate when [A]=0.050 mol·L-1 and [B]=0.100 mol·L-1
27Use the data in table 12.5 to determine 1) The orders for all three reactants2) The overall reaction order3) The value of the rate constant
28Initial Rate (mol·L-1·s-1) 2NO(g) + 2H2(g) N2(g) + 2H2O(g)Experiment Number[NO] (mol·L-1)[H2] (mol·L-1)Initial Rate (mol·L-1·s-1)10.101.23 x 10-320.202.46 x 10-334.92 x 10-3Determine the differential rate lawCalculate the rate constantCalculate the rate when [NO]=0.050 mol·L-1 and [H2]=0.150 mol·L-1
29Consider the general reaction aA + bB cC and the Sample Problem:.Consider the general reaction aA + bB cC and thefollowing average rate data over some time period Δt:Determine a set of possible coefficients to balance thisgeneral reaction.
30Problem Reaction: A + B C obeys the rate law: Rate = k[A]2[B]. A. If [A] is doubled (keeping B constant), how will rate change?B. Will rate constant k change? Explain.C. What are the reaction orders for A & B?D. What are the units of the rate constant?
31You now know that….The rate expression correlates consumption of reactant to production of product. For a reaction: 3A 2B- 1D[A] = 1D[B]3 Dt 2 DtThe differential rate law allows you to correlate rate with concentration based on the format: Rate = k [A]n
32You also know that…1. Rate of consumption of reactant decreases over time because the concentration of reactant decreases. Lower concentration equates to lower rate.2. If a graph of concentration vs. time were constructed, the graph is not a straight line
33Experiment 23You varied concentration of KIO3 and held the concentration of NaHSO3 constant.From Expt. A E, you increased the amount of KIO3.Observed result: The time it took for the reaction to occur DECREASED.
34Higher concentration of KIO3 lead to faster rate of reaction. ConclusionHigher concentration of KIO3 lead to faster rate of reaction.
41Integrated Rate LawExpresses the dependence of concentration on time
42The Change of Concentration with Time First Order ReactionsGoal: convert rate law into a convenient equation to give concentrations as a function of time.For a first order reaction, the rate doubles as the concentration of a reactant doubles.
43Integrated Rate Laws Zero Order: [A]t = -kt + [A]o First Order: ln[A]t = -kt + ln[A]oSecond Order: = kt [A]t [A]o where [A]o is the initial concentration and [A]t is the final concentration.
44Integrated First-Order Rate Law ln[A]t = -kt + ln[A]0Eqn. shows [concn] as a function of timeGives straight-line plot since equation is of the form y = mx + b
45The Change of Concentration with Time Zero Order ReactionsA plot of [A]t versus t is a straight line with slope -k and intercept [A]0.
46The Change of Concentration with Time First Order ReactionsA plot of ln[A]t versus t is a straight line with slope -k and intercept ln[A]0.
47The Change of Concentration with Time Second Order ReactionsA plot of 1/[A]t versus t is a straight line with slope k and intercept 1/[A]0.