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Unformatted text preview: Chapter 14: Chemical Kinetics How to make reactions as fast as we want them to be. Topics Covered
q Defining Reaction Rate q Calculating Rates q Using Initial Rates q Factors that Affect Rate q Collision Theory
qOrientation qActivation Energy Reaction Rates: Why Care?
q Reactions happen at a variety of time scales, from millions of years to millionths of seconds q The same reaction can happen at very different rates given different conditions Reactions: Fast and Slow
q 2C8H18 (g)+ 25O2 (g) 16CO2 (g)+ 18H2O (g) Fast Slow (or stopped) Reactions: Fast and Slow
q H2CO3 (aq) Slow H2O (l) + CO2 (g) Too Fast! Reactions: Fast and Slow
q 235U + 1n Slow....
92 Kr + 141Ba + 3 1n Faster.... Fastest What Affects Rate?
q To know how to change reaction rates, we need to understand what goes on in the reaction (and vice versa) q What is a chemical reaction, generally?
qAtomic Rearrangement qChange in Energy Factors Affecting Rate
q Reactions usually require collisions between reactant molecules q Increasing collisions increases the reaction rate q Collisions must also have:
qSufficient energy qSuitable orientation Factors Affecting Rate
q Rate is affected by:
qPhysical state of reactants
qCan affect number of molecules available to collide qConcentration of reactants
qHigher concentrations often increase collisions qTemperature of reaction
qHigher temperature means more collisions with more energy qPresence of catalyst
qAffect mechanism of reactions or orientation of molecules In your car's engine, which factors are regulated to control the speed of combustion of gasoline?
1. Reaction Temperature 2. Physical State of Reactants 3. Reactant Concentration 4. All of the Above
3 4 1 What Is Rate?
q In the last example, interested in change in gasoline vs. time
q Generally, Rate= positive) Rate = moles t (or moles, so Rate is -moles C8H18 t q In solution, can use concentration instead of moles q In gas phase, can use partial pressures 1 Defining Rates
q Consider the reaction: A B -[ A] [ B] qRate= t = t qWhen ratios are not 1:1, use stoichiometric coefficients to relate rates qN2 (g) + 3H2 (g) 2NH3 (g) qRate =
-[N 2 ] t =
1 [H 2 ] - 3 t = 1 [NH 3 ] 2 t 1 Rates
N2 + 3H2
3.5 3 2.5 Conc. 2 1.5 1 0.5 0 0 5 10 Time (hr) 15 20 2NH3 N2 H2 NH3 0.285 0.095 0.190 1 Calculating Rates
q 2NO2 2NO + O2 q Start with 0.1 M NO2 and measure concentrations at certain time intervals 1 [NO ] 2 - q Rate Expression: Rate= 2 t q Calculate between time = 300s and 400s 1 Calculating Rates
0 50 100 200 300 * 400 * 500 [NO2] (M)
0.1000 0.0269 0.0156 0.0084 0.0058 0.0044 0.0036 q Rate= q Rate= 1 [NO 2 ] - 2 t - 1 [NO 2 ]2 - [NO 2 ]1 2 t2 - t1 1 0.0044M - 0.0058M 2 400 s - 300 s
1 -0.0014 M M = 7.00 x10-6 2 100 s s q Rate= - q Rate=
- 1 Calculating Rates
q This rate is an average rate; can also determine instantaneous rates from graph q Rate changes with time (usually decreases) 1 Initial Rates
q Can compare reactions based on initial rates C q Use as t approximation q Often, concentrations known better at beginning anyway
1.2 1 0.8
[N2] 0.6 0.4 0.2 0 0 5 10 Time (hr) 15 20 1 From the given data, calculate the initial rate of the reaction, and the rate between t=10h and t=15h
N2 + 3H2 2NH3
Time (hr) pNH3 (atm) 0 1 3 5 10 15 20
1 0.000 0.190 0.518 0.787 1.264 1.554
2 1. 2. 3. 4. 0.19, 0.290 0.095, 0.145 0.285, 0.058 0.095, .029 0% 1.729 0%
4 1 Urban Legend + = ? q Fact: methane does combust in oxygen: CH4 (g) + 2O2 (g) CO2 (g) + 2H2O (g) 1 Busting the Myth And nothing happens 2Light a spark
Methane goes In Measure [CH4], reads 15% 2 q Try #2 (or 6)
Portapotty goes boom! 2Continuous Flame
Methane goes In Measure [CH4], reads 1% 2 Let's take a closer look
High [CH4] Low [CH4] 2 Concentration Effects
q Rate depends on concentration q As reactant molecules are consumed, fewer reactants collide, so rate decreases
1.2 1 0.8 [N2] . 0.6 0.4 0.2 0 0 -0.2 Time (hr) 5 10 15 20 2 Rate Laws
q Concentration dependence can be described by a rate law: Rate = k[A]x[B]y... q x and y are not stoichiometric coefficients q Determine k, x, y from experiments:
qMeasure initial rate at different [A], [B] to determine x and y qk is called rate constant qx, y are reaction order for each species qx+y+...is overall reaction order 2 Determining Order
q Change concentration of one species, and measure rate; compare ratio of rates to ratio of concentrations NH3 q Example: N2 + 3H2
pN2 (atm) 1 3 3 pH2 (atm) 2 2 4 pNH3 (atm) 0 0 0 Rate (atm/hr) 0.095 0.280 1.12 2 Determining Order
q N2 + 3H2
1 3 3 NH3, Rate=k*(pN2)x(pH2)y
Initial Conditions pH2 (atm)
2 2 4 pNH3 (atm)
0 0 0 Rate (atm/hr)
0.095 0.280 1.12 q Compare Experiment 1 and 2: pN /pN =3, 2 Compare Experiment 2 and 3; what is y? Rate=k*(pN2)x(pH2)y
1 3 3 pH2 (atm)
2 2 4 pNH3 (atm)
0 0 0 Rate (atm/hr)
0.095 0.280 1.12 1. 2. 3. 4. 5. 0 1 2 3 1/2 0 %
1 0 %
2 0 %
3 0 %
4 0 %
5 2 Determining Rate Laws
q Rate=k*(pN2)x(pH2)y q x=1, y=2; Rate=k(pN2)(pH2)2 q How do we find k?
qOnce reaction orders are known, pick an experiment and solve for k
pN2 (atm) 1 3 3 pH2 (atm) 2 2 4 pNH3 (atm) 0 0 0 Rate (atm/hr) 0.095 0.280 1.12 2 According to the given data, which is correct rate law for this reaction? [F ] (M) [ClO ] (M) Rate (M/s)
2 2 F2 + 2ClO2 2FClO2 0.100 0.100 0.200 0.0100 0.0400
25% 0.0100 25% 0.0012 0.0048
25% 25% 0.0024 1. Rate= 1.2*[F2]2[ClO2] 2. Rate= 2.4*[F2][ClO2]2 3. Rate= 2.4*[F2]2[ClO2] 4. Rate= 1.2*[F ][ClO ]2
1 2 3 4 2 14.4 Concentration vs. Time
q Rate Laws in the form Rate = k[A]x[B]y tell us about the rate of a reaction at a specific instant, and about the dependence of initial concentration on initial rate. q This form of the rate law does not tell us information about time, such as "how long it will take for a specific concentration to react completely, or half way, ..." q Integrated rate laws will give us this information. 3 14.4 The Change of Concentration with Time How do we determine rate laws? The dependence of concentration on reaction rate for each reactant is either zero order, 1st order, or 2nd order (usually). q Zero Order: Rate = - [A]/ t = k[A]0 = k q 1st Order: Rate = - [A]/ t = k[A]1 q 2nd order: Rate = - [A]/ t = k[A]2 q Converting these equations to linear equations as a function of time will allow us to graph the three possibilities (Method 3) 3 14.4 The Change of Concentration with Time Zero Order in A: (n=0) q Rate = - [A]/ t = k[A]0 = k q - [A]/ t = k Rearranging and taking the integral of both sides gives the integrated rate law: [A] = [A] - kt q 3 How do we determine rate laws?
q If zero order ( [H2]/ t = k), a graph of [H2] vs. time should be linear. [H2] = [H2]o + kt
Why +kt in this case? 3 How do we Determine Rate Laws?
1st Order in A (n=1) q Rate = - [A]/ t = k[A] Rearranging and taking the integral of both sides gives [A] integrated rate law: ln [A] = ln the - kt q
o 3 How do we Determine Rate Laws?
1st Order in A (n=1) q ln [A] = ln [A]o - kt q 2N2O5 4NO2 + O2(g) q If n = 1, a graph of ln [N2O5] vs. time should be linear. 3 How do we determine rate laws?
2nd Order in A (n=2) q Rate = - [A]/ t = k[A]2 Rearranging and taking the integral of both sides gives the+ kt 1/[A] = 1/[A] integrated rate law: q
o 3 How do we determine rate laws?
2nd Order in A (n=2) q 1/[A] = 1/[A]o + kt q 2UO2+ + 4H+ U4+ + UO22+ + 2H2O q If n = 2, a graph of 1/[UO2+] vs. time should be linear. 3 How do we determine rate laws?
q Summary: n = 0 - [A]/ t = k n = 1 - [A]/ t = k[A] n = 2 - [A]/ t = k[A]2 [A] = [A]o - kt ln [A] = ln [A]o - kt 1/[A] = 1/[A]o + kt Another Method for Determining Rate Law: Graph [A] vs. t, ln [A] vs. t, and 1/[A] vs. t and see which graph is linear. 3 What is the order? 3 What is the order? 4 ...
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This note was uploaded on 04/06/2011 for the course CHM 116 taught by Professor Unknown during the Spring '08 term at ASU.
- Spring '08