Presentation on theme: "Topic # 15 Chemical Kinetics"— Presentation transcript:
1Topic # 15 Chemical Kinetics Conditions of Reaction RatesRate of a reactionGeneral Rate LawOrders of ReactionsHalf-LifeReaction Mechanisms(Zumdahl Chapter 12)
2A reaction will only take place if three conditions are met: Collision TheoryA reaction will only take place if threeconditions are met:1. the reactants come into contact (collide)2. the collision happens with enough energy,(activation energy)3. the reactants hit at the best possibleorientation to facilitate a reaction.
3Factors that control the rate of a reaction A. Concentrationputting more reactants in the same spacewill increase the collision frequency, producinga faster reaction rate. Same effect occursby increasing the pressure in a gaseousenvironment.B. TemperatureIncreasing temperature will increase therate of a reaction. GENERALLY, a 10oC risewill result in approx. doubling the rate.
4As the system is heated, the position of the average kinetic energy must shift to the right in order to reflect that change.As the system cools, the average kinetic energy must respond by shifting position to the left.At higher temperatures, the particles are more evenly distributed over a range of kinetic energy values.
5Reaction rates are temperature dependent. Rate and temperatureReaction rates are temperature dependent.Here are rate constantsfor N2O5 decompositionat various temperatures.T, oC k x 104, s-1k x 104 (s-1)Temperature (oC)
6C. Solid particle sizeincrease in surface area of the solid willincrease the reaction rateD. CatalysisCatalysts are substances that increase thethe rate of a reaction while remaining unchangedchemically. They work by providing andifferent reaction route that needs a loweractivation energy.
8Rates of reactions Rate of a chemical reaction. The change in the quantity of a reactant or product that takes place in a period of time.rate ==concentrationlater - concentrationearliertimelater - timeearlierD [ ]D tMolarity
9Rates of reactions To study rates of reaction, you must: Identify the reactants and products.Carry out the reaction.Measure the concentrations of one of the reactants or products at known intervals.There MUST be DATA!There needs to be a way to measure at least one of the species involved.
10The example reaction Decomposition of N2O5 Dinitrogen pentoxide is known to decompose completely by the following reaction.2N2O5 (g) N2O4 (g) + O2 (g)This reaction can be conducted in an inert solvent like carbon tetrachloride.When N2O5 decomposes, N2O4 remains in solution and O2 escapes and can be measured.
11An example reactionWe can easily measure the oxygen as dinitrogen pentoxide decomposes.Temperature must be maintained to within oC.The reaction flask must be shaken to keep oxygen from forming a supersaturated solution.It is found that the reaction initially occurs very rapidly but gradually slows down.
12An example reactionGasburetConstant temperature bath
13An example reaction Time (s) STP O2 in mL 0 0 300 1.15 600 2.18 Here are the results for an experiment.
14An example reaction Volume, mL O2 The rate of O2 production slows down with time.Time, s
15Calculations with Average rates We can calculate the average rate of oxygen formation during any time interval as:Average rate ofO2 formationD V O2D t=Time, s Rate O2**The rates shown here have units O2 mL(at STP)/s.Notice how the rate decreases with time.
16Looking at the N2O5 reaction again….. Since we know the stoichiometry for our reaction, we can calculate the concentration of N2O5 during the reaction.2N2O5 (g) N2O4 (g) + O2 (g)For each mole of O2 produced, two moles of N2O5 will have decomposed.The rate of reaction will be:rate of reaction = = -D[O2]Dt12D[N2O5]Dt
17This method can be used for any equation……. Using the equation:2 A + B C + DThe rate of reaction can be expressed in terms of D:d[D]dtd[C]dt--1 d[A]2 dtd[B]dt==rate ==Note the use of coefficients and choosing a componentto base the rates on!
18Use the Equation to solve for concentrations: 2 A + B + 3 C + Dd[C]dt= mole/liter-mind[A]dtFind value of23Relationship of [A] will be negative and ratio-2 [0.35]= mole/liter-minis being lost
20General Rate Equation A + B C + D For an equation: rate = k [A]x[B]y x & y are ORDER of reaction for that componentthey are NOT coefficients!!k is the rate constantCalled therate law for theequationActualRated[C]dt=k [A]x [B]yCollected data is used to answerquestions on rate orders.
21Finding rate order Method of initial rates The order for each reactant is found bycompleting an experiment and using the data:Change the initial concentration of a reactant.Hold all other initial concentrations and conditions constant.Measure the initial (or starting) rates of reactionThe change in rate is used to determine the order for that specific reactant. The process is repeated for each reactant.
22N2O5 example 2N2O5 (g) 2N2O4 (g) + O2 (g) The following data was obtained for the decomposition of N2O5.Experiment [N2O5] Initial rate, M/sx 10-5x 10-5We know that the rate expression is:rate = k [N2O5]xOur goal is to determine what x (the order) is.
23N2O5 example For exp. 2 7.29 x 10-5 M/s = k (0.200 M)x We can now divide the equation for exp. two by the one for exp. one.7.29 x 10-5 M/s k (0.200 M)x3.62 x 10-5 M/s k (0.100 M)xwhich give 2.01 = (2.00)xand x = 1 (first order reaction)=rate = k [N2O5]1
24Exp. [A] [B] [C] Initial rate, M/s 1 0.030 0.010 0.050 1.7 x 10-8 The data below for rate of reaction was obtained for the following reaction:A + B + CExp [A] [B] [C] Initial rate, M/sx 10-8x 10-8x 10-8x 10-8Note that the concentrationsdouble in some experiments.Notice the productsare not needed!
25Equation: A + B + C ……… Rate Law: rate = k[A]x [B]y [C]z Pick data so calculations for two of the reactantswill cancel out……..
26x = 2 [A] is second order Order for A Use exps one and two since [B] and [C] are the same and would cancel out.6.8 x 10-8 M/s k(0.060 M)x1.7 x 10-8 M/s k(0.030 M)x4.0 = (2.0)xx = [A] is second order=
27Use experiments one and three. Order for BUse experiments one and three.4.9 x 10-8 M/s (0.020 M)y1.7 x 10-8 M/s (0.010 M)y2.9 = (2.0)yThe order is not obvious by inspection. You must take the natural logarithm of both sides and solve for y.ln 2.9 = y (ln 2.0)y = or=32
28Use experiments one and four. Order for CUse experiments one and four.Experiment [C] Initial Ratex 10-8x 10-8Here the rate did not change when [C] was doubled. This is an example of a zero order reaction.z = 0
29A more complex example We can now write the overall rate law. rate = [A]2 [B]3/2 [C]0since [C] has no effect on the rate:rate = [A]2 [B]3/2The overall order for the reaction is:x + y + z = / = /2
31Graphing rate laws Graphical method (some form of [ ] versus time!) Using integrated rate law, one can produce straight line plots. The order for a reactant is assigned if the data produces a straight line.Rate integrated Graph SlopeOrder law rate law vs. time0 rate = k [A]t = -kt + [A] [A]t -k1 rate = k[A] ln[A]t = -kt + ln[A] ln[A]t -k2 rate=k[A] = kt k1[A]t1[A]01[A]t
32Graphing rate laws As you can see from these plots of the N2O5 data, 0 order plot2nd order plot[N2O5]1/[N2O5]Time (s)Time (s)Time (s)As you can see from theseplots of the N2O5 data,only a first order plotresults in a straight line.So N2O5 is 1st order!1st order plotln[N2O5]
33Practice Graphing:Time (s) [NO2] (mol/L)1.2 x3.0 x4.5 x9.0 x1.8 xGraph:[NO2] vs. time Zero orderln[NO2] vs. time First Order1/[NO2] vs. time Second OrderStraight line tells the order !
34Special First order reactions Reactions that are first order with respect to a reactant are of great importance.Describe how many drugs pass into the blood stream or used by the body.Often useful in geochemistryRadioactive decayHalf-life (t1/2)The time required for one-half of the quantity of reactant originally present to react.
35Half-life From our N2O5 data, we can see that it takes about 1900 seconds for the concentrationto be reduced in half.It takes another 1900 secondsto reduce the concentration inhalf again.[N2O5]Time (s)
36Half-life The half-life and the rate constant are related. t1/2 = Half-life (time) can be used to calculate the first order rate constant.For our N2O5 example, the reaction took 1900 seconds to react half way so:k = = = 3.65 x 10-4 s-10.693k0.693t1/20.6931900 s
37Fluorine -21 has a half-life of 5 seconds. What fraction of the original nuclei remain after 1 minute?The answer is solved by creating the fractionWhere n = number of ½ lives.In 1 minute there will be 12 half-lives.Answer is…….
38Reaction mechanismsA detailed molecular-level picture of how a reaction might take place.activatedcomplex= bonds in theprocess ofbreaking orbeing formed
39Could also look like:R + S D but in reality…R BB + S DOr even:A + B C + D is really:A XX + B C + DNever see the activated complexin the equation!
40Reaction mechanisms Molecularity The number of particles that come together to form the activated complex in an elementary process (each step).1 - unimolecular2 - bimolecular3 - termolecular
41Reaction mechanisms rate = k [NO] [O3] For some elementary processes, the exponents for each species in the rate law are the same as the coefficients in the equation for the step.For our earlier example, Bimolecular….the rate law is:rate = k [NO] [O3](Bimolecular mechanism)
42Reaction mechanismsIn general, the rate law gives the composition of the activated complex, because that is what comes together to form the complex.The power of a species in the rate law is the same as the number of particles of the species in the activated complex.If the exponents in the rate law are not the same as the coefficients of the equation for the reaction, the overall reaction must consist of more than one step.Lets look at N2O5 - again!
43Reaction mechanisms Earlier we found that for: 2N2O5 2N2O4 + O2 The rate law was:rate = k [N2O5]According to the equation, it should be second order (coefficient of 2 in front) but the data shows it to be first order.The reaction must involve more than one step.
44Reaction Mechanisms Consider the following reaction. 2NO2 (g) + F2 (g) NO2F (g)If the reaction took place in a single step the rate law might be:rate = k [NO2]2 [F2]However, the experimentally observed rate law is:rate = k [NO2] [F2]
45Reaction MechanismsSince the observed rate law is not the same as if the reaction took place in a single step, we know two things.More than one step must be involvedThe activated complex must be produced from two species.A possible reaction mechanism might be:Step one NO2 + F NO2F + FStep two NO2 + F NO2FOverall 2NO2 + F2 2NO2F
46Reaction Mechanisms Rate-determining step. When a reaction occurs in a series of steps, with one slow step, it is the slow step that determines the overall rate.Step one NO2 + F NO2F + FExpected to be slow. It involves breaking an F-F bond.Step two NO2 + F NO2FExpected to be fast. A fluorine atom is very reactive.
47Reaction MechanismsSince step one is slow, we can expect this step to be the determiner of the overall rate of the reaction.NO2 + F NO2F + FThis would give a rate expression of:rate = k1 [NO2] [ F2]This agrees with the experimentally observed results.
48Let’s look again:R + S D but in reality…R B slowB + S D very fastThe rate only depends on the concentrationof R so the rate law only contains R!Rate = k[R]xThe Slow process determines the rate law!
49A + B C + D is really:A X Fast and equilibriumX + B C + D SLOWHere the slow step contains X and B, but theformation of X depends on A. Since X isan intermediate it isn’t in the rate law. X issubstituted with A….Rate = k[A][B]No intermediates are written in the rate law!
50Take note: the sum of all steps needs to tally up to the reaction equation.R + S DR B slowB + S D very fast
51Bimolecular first step mechanism Another mechanism example……All gases:NO2 + NO2 NO3 + NO SLOWNO3 + CO NO2 + CO2 FastRate of Formation of NO3 =∆[NO3]∆tOverall rate = k[NO2]2Bimolecular first step mechanism