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Copyright 1999, PRENTICE HALL Chapter 14 1
Chemical KineticsChemical Kinetics
Chapter 14Chapter 14
David P. WhiteDavid P. White
University of North Carolina, WilmingtonUniversity of North Carolina, Wilmington
Copyright 1999, PRENTICE HALL Chapter 14 2
Chemical KineticsChemical Kinetics• 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, and
– surface area.
• Goal: to understand chemical reactions at the molecular level.
Copyright 1999, PRENTICE HALL Chapter 14 3
Chemical KineticsChemical KineticsReaction RatesReaction Rates• Speed of a reaction is measured by the change in
concentration with time.• For a reaction A B
• Suppose A reacts to form B. Let us begin with 1.00 mol A.
t
B of moles
time in changeB of moles ofnumber the in change
rate Average
Copyright 1999, PRENTICE HALL Chapter 14 4
Chemical KineticsChemical KineticsReaction RatesReaction Rates
Copyright 1999, PRENTICE HALL Chapter 14 5
Chemical KineticsChemical KineticsReaction RatesReaction Rates– At t = 0 (time zero) there is 1.00 mol A (100 red spheres)
and no B present.
– At t = 20 min, there is 0.54 mol A and 0.46 mol B.
– At t = 40 min, there is 0.30 mol A and 0.70 mol B.
– Calculating,
mol/min 026.0min 0 - min 10mol 0 - mol 26.0
min 0 - min 100at B of moles10at B of moles
B of molesrate Average
ttt
Copyright 1999, PRENTICE HALL Chapter 14 6
Chemical KineticsChemical KineticsReaction RatesReaction Rates• For 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), or
– the speed at which the reactants disappear (i.e. the change in moles of A per unit time).
• A plot of number of moles versus time shows that as the reactants (red A) disappear, the products (blue B) appear.
Copyright 1999, PRENTICE HALL Chapter 14 7
Chemical KineticsChemical KineticsReaction RatesReaction Rates
Copyright 1999, PRENTICE HALL Chapter 14 8
Chemical KineticsChemical KineticsRates in Terms of ConcentrationsRates in Terms of Concentrations• For the reaction A B there are two ways of• Most useful units for rates are to look at molarity.
Since volume is constant, molarity and moles are directly proportional.
• Consider:
C4H9Cl(aq) + H2O(l) C4H9OH(aq) + HCl(aq)
Copyright 1999, PRENTICE HALL Chapter 14 9
Chemical KineticsChemical KineticsRates in Terms of ConcentrationsRates in Terms of Concentrations
C4H9Cl(aq) + H2O(l) C4H9OH(aq) + HCl(aq)
Copyright 1999, PRENTICE HALL Chapter 14 10
Chemical KineticsChemical KineticsRates in Terms of ConcentrationsRates in Terms of Concentrations
C4H9Cl(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.
Copyright 1999, PRENTICE HALL Chapter 14 11
Chemical KineticsChemical KineticsRates in Terms of ConcentrationsRates in Terms of Concentrations
Copyright 1999, PRENTICE HALL Chapter 14 12
Chemical KineticsChemical KineticsReaction Rates and StoichiometryReaction Rates and Stoichiometry• For the reaction
C4H9Cl(aq) + H2O(l) C4H9OH(aq) + HCl(aq)
we know
• In general for
aA + bB cC + dD
t
OHHC
t
ClHCRate 9494
tD1
tC1
tB1
tA1
Rate
dcba
Copyright 1999, PRENTICE HALL Chapter 14 13
The Dependence of Rate on ConcentrationThe Dependence of Rate on Concentration• In general rates increase as concentrations increase.
NH4+(aq) + NO2
-(aq) N2(g) + 2H2O(l)
Copyright 1999, PRENTICE HALL Chapter 14 14
The Dependence of Rate on ConcentrationThe Dependence of Rate on Concentration• For the reaction
NH4+(aq) + NO2
-(aq) N2(g) + 2H2O(l)
we note – as [NH4
+] doubles with [NO2-] constant the rate doubles,
– as [NO2-] doubles with [NH4
+] constant, the rate doubles,
– We conclude rate [NH4+][NO2
-].
• Rate law:
Rate = k[NH4+][ NO2
-].
• The constant k is the rate constant.
Copyright 1999, PRENTICE HALL Chapter 14 15
The Dependence of Rate on ConcentrationThe Dependence of Rate on Concentration• For a general reaction with rate law
Rate = k[reactant 1]m[reactant 2]n,
we 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.
Copyright 1999, PRENTICE HALL Chapter 14 16
The Dependence of Rate on ConcentrationThe Dependence of Rate on ConcentrationUsing Initial Rates to Determine Rate LawsUsing Initial Rates to Determine Rate Laws• A 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 reaction is second order if doubling the
concentration results in a 22 increase in rate. • 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.
Copyright 1999, PRENTICE HALL Chapter 14 17
The Change of Concentration with TimeThe Change of Concentration with TimeFirst-Order ReactionsFirst-Order Reactions• Goal: 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.• We can show that
• A plot of ln[A]t versus t is a straight line with slope -k and intercept ln[A]0.
• In the above we use the natural logarithm, ln, which is log to the base e.
0AlnAln ktt
Copyright 1999, PRENTICE HALL Chapter 14 18
The Change of Concentration with TimeThe Change of Concentration with TimeFirst-Order ReactionsFirst-Order Reactions
0AlnAln ktt
Copyright 1999, PRENTICE HALL Chapter 14 19
The Change of Concentration with TimeThe Change of Concentration with TimeHalf-LifeHalf-Life• Half-life is the time taken for the concentration of a
reactant to drop to half its original value.
• That is, half life, t1/2 is the time taken for [A]0 to reach ½[A]0.
• Mathematically,
kkt
693.0ln21
21
Copyright 1999, PRENTICE HALL Chapter 14 20
The Change of Concentration with TimeThe Change of Concentration with TimeHalf-LifeHalf-Life
Copyright 1999, PRENTICE HALL Chapter 14 21
The Change of Concentration with TimeThe Change of Concentration with TimeSecond-Order ReactionsSecond-Order Reactions• For a second order reaction with just one reactant
• A plot of 1/[A]t versus t is a straight line with slope k and intercept 1/[A]0
• For a second order reaction, a plot of ln[A]t vs. t is not linear.
0A1
A1 kt
t
Copyright 1999, PRENTICE HALL Chapter 14 22
The Change of Concentration with TimeThe Change of Concentration with TimeSecond-Order ReactionsSecond-Order Reactions
Copyright 1999, PRENTICE HALL Chapter 14 23
The Change of Concentration with TimeThe Change of Concentration with TimeSecond-Order ReactionsSecond-Order Reactions• We can show that the half life
• A reaction can have rate constant expression of the form
rate = k[A][B],
i.e., is second order overall, but has first order dependence on A and B.
0A1
21
kt
Copyright 1999, PRENTICE HALL Chapter 14 24
Temperature and RateTemperature and Rate• Most reactions speed up as temperature increases.
(E.g. food spoils when not refrigerated.)• When two light sticks are placed in water: one at
room temperature and one in ice, the one at room temperature is brighter than the one in ice.
• The chemical reaction responsible for chemiluminescence is dependent on temperature: the higher the temperature, the faster the reaction and the brighter the light.
• As temperature increases, the rate increases.
Copyright 1999, PRENTICE HALL Chapter 14 25
Temperature and RateTemperature and Rate
Copyright 1999, PRENTICE HALL Chapter 14 26
Temperature and RateTemperature and Rate• As temperature increases, the rate increases.• Since the rate law has no temperature term in it, the
rate constant must depend on temperature.
• Consider the first order reaction CH3NC CH3CN. – As temperature increases from 190 C to 250 C the rate
constant increases from 2.52 10-5 s-1 to 3.16 10-3 s-1.
• The temperature effect is quite dramatic. Why?
Copyright 1999, PRENTICE HALL Chapter 14 27
Temperature and RateTemperature and RateThe Collision ModelThe Collision Model• Observations: rates of reactions are affected by
concentration and temperature.• Goal: develop a model that explains why rates of
reactions increase as concentration and temperature increases.
• The collision model: in order for molecules to react they must collide.
• The greater the number of collisions the faster the rate.
Copyright 1999, PRENTICE HALL Chapter 14 28
Temperature and RateTemperature and RateThe Collision ModelThe Collision Model
Copyright 1999, PRENTICE HALL Chapter 14 29
Temperature and RateTemperature and RateThe Collision ModelThe Collision Model• The more molecules present, the greater the
probability of collision and the faster the rate.• The higher the temperature, the more energy
available to the molecules and the faster the rate.• Complication: not all collisions lead to products. In
fact, only a small fraction of collisions lead to product.• In order for reaction to occur the reactant molecules
must collide in the correct orientation and with enough energy to form products.
Copyright 1999, PRENTICE HALL Chapter 14 30
Temperature and RateTemperature and RateActivation EnergyActivation Energy• Arrhenius: molecules must posses a minimum amount
of energy to react. Why?– In order to form products, bonds must be broken in the
reactants.
– Bond breakage requires energy.
• Activation energy, Ea, is the minimum energy required to initiate a chemical reaction.
Copyright 1999, PRENTICE HALL Chapter 14 31
Temperature and RateTemperature and RateActivation EnergyActivation Energy
Copyright 1999, PRENTICE HALL Chapter 14 32
Temperature and RateTemperature and RateActivation EnergyActivation Energy• Consider the rearrangement of acetonitrile:
– In H3C-NC, the C-NC bond bends until the C-N bond breaks and the NC portion is perpendicular to the H3C portion. This structure is called the activated complex or transition state.
– The energy required for the above twist and break is the activation energy, Ea.
– Once the C-N bond is broken, the NC portion can continue to rotate forming a C-CN bond.
H3C N CC
NH3C H3C C N
Copyright 1999, PRENTICE HALL Chapter 14 33
Temperature and RateTemperature and RateActivation EnergyActivation Energy
Copyright 1999, PRENTICE HALL Chapter 14 34
Temperature and RateTemperature and RateActivation EnergyActivation Energy• The change in energy for the reaction is the difference
in energy between CH3NC and CH3CN.
• The activation energy is the difference in energy between reactants, CH3NC and transition state.
• The rate depends on Ea.
• Notice that if a forward reaction is exothermic (CH3NC CH3CN), then the reverse reaction is endothermic (CH3CN CH3NC).
Copyright 1999, PRENTICE HALL Chapter 14 35
Temperature and RateTemperature and RateActivation EnergyActivation Energy• Consider the reaction between Cl and NOCl:– If the Cl collides with the Cl of NOCl then the products are
Cl2 and NO.
– If the Cl collided with the O of NOCl then no products are formed.
• We need to quantify this effect.
Copyright 1999, PRENTICE HALL Chapter 14 36
Temperature and RateTemperature and RateActivation EnergyActivation Energy
Copyright 1999, PRENTICE HALL Chapter 14 37
Temperature and RateTemperature and RateThe Arrhenius EquationThe Arrhenius Equation• Arrhenius discovered most reaction-rate data obeyed the Arrhenius
equation:
– k is the rate constant, Ea is the activation energy, R is the gas constant (8.314 J/K-mol) and T is the temperature in K.
– A is called the frequency factor.
– A is a measure of the probability of a favorable collision.
– Both A and Ea are specific to a given reaction.
• If we have a lot of data, we can determine Ea and A graphically by rearranging the Arrhenius equation:
• If we do not have a lot of data, then we can use
RTaE
Aek
Copyright 1999, PRENTICE HALL Chapter 14 38
Temperature and RateTemperature and RateThe Arrhenius EquationThe Arrhenius Equation• If we have a lot of data, we can determine Ea and A
graphically by rearranging the Arrhenius equation:
• If we do not have a lot of data, then we can use
ART
Ek a lnln
122
1 11ln
TTR
E
kk a
Copyright 1999, PRENTICE HALL Chapter 14 39
Reaction MechanismsReaction Mechanisms• The balanced chemical equation provides information
about the beginning and end of reaction.• The reaction mechanism gives the path of the
reaction.• Mechanisms provide a very detailed picture of which
bonds are broken and formed during the course of a reaction.
Copyright 1999, PRENTICE HALL Chapter 14 40
Reaction MechanismsReaction MechanismsElementary StepsElementary Steps• Elementary step: any process that occurs in a single
step.• Molecularity: the number of molecules present in an
elementary step.– Unimolecular: one molecule in the elementary step,
– Bimolecular: two molecules in the elementary step, and
– Termolecular: three molecules in the elementary step.
• It is not common to see termolecular processes (statistically improbable).
Copyright 1999, PRENTICE HALL Chapter 14 41
Reaction MechanismsReaction MechanismsElementary StepsElementary Steps• Elementary steps must add to give the balanced
chemical equation.• Intermediate: a species which appears in an
elementary step which is not a reactant or product.
Rate Laws of Elementary StepsRate Laws of Elementary Steps• The rate law of an elementary step is determined by
its molecularity:– Unimolecular processes are first order,
– Bimolecular processes are second order, and
– Termolecular processes are third order.
Copyright 1999, PRENTICE HALL Chapter 14 42
Reaction MechanismsReaction MechanismsRate Laws of Multistep MechanismsRate Laws of Multistep Mechanisms• Rate-determining step: is the slowest of the
elementary steps.• Therefore, the rate-determining step governs the
overall rate law for the reaction.
Mechanisms with an Initial Fast StepMechanisms with an Initial Fast Step• It is possible for an intermediate to be a reactant.• Consider
2NO(g) + Br2(g) 2NOBr(g)
Copyright 1999, PRENTICE HALL Chapter 14 43
Reaction MechanismsReaction MechanismsMechanisms with an Initial Fast StepMechanisms with an Initial Fast Step
2NO(g) + Br2(g) 2NOBr(g)
• The experimentally determined rate law is
Rate = k[NO]2[Br2]
• Consider the following mechanism
for which the rate law is (based on Step 2):
NO(g) + Br2(g) NOBr2(g)k1
k-1
NOBr2(g) + NO(g) 2NOBr(g)k2
Step 1:
Step 2:
(fast)
(slow)
Copyright 1999, PRENTICE HALL Chapter 14 44
Reaction MechanismsReaction MechanismsMechanisms with an Initial Fast StepMechanisms with an Initial Fast Step• The rate law is (based on Step 2):
Rate = k2[NOBr2][NO]
• The rate law should not depend on the concentration of an intermediate because intermediates are usually unstable.
• Assume NOBr2 is unstable, so we express the concentration of NOBr2 in terms of NOBr and Br2 assuming there is an equilibrium in step 1 we have
]Br][NO[NOBr 21
12
kk
Copyright 1999, PRENTICE HALL Chapter 14 45
Reaction MechanismsReaction MechanismsMechanisms with an Initial Fast StepMechanisms with an Initial Fast Step• By definition of equilibrium:
k1[NO][Br2] = k-1[NOBr2]
• Therefore, the overall rate law becomes
• Note the final rate law is consistent with the experimentally observed rate law.
]Br[]NO[Rate 22
1
12
kk
k
Copyright 1999, PRENTICE HALL Chapter 14 46
CatalysisCatalysis• A catalyst changes the rate of a chemical reaction.• There are two types of catalyst:– homogeneous, and
– heterogeneous.
• Chlorine atoms are catalysts for the destruction of ozone.
Homogeneous CatalysisHomogeneous Catalysis• The catalyst and reaction is in one phase.• Hydrogen peroxide decomposes very slowly:
2H2O2(aq) 2H2O(l) + O2(g).
Copyright 1999, PRENTICE HALL Chapter 14 47
CatalysisCatalysisHomogeneous CatalysisHomogeneous Catalysis
2H2O2(aq) 2H2O(l) + O2(g).
• In the presence of the bromide ion, the decomposition occurs rapidly:– 2Br-(aq) + H2O2(aq) + 2H+(aq) Br2(aq) + 2H2O(l).
– Br2(aq) is brown.
– Br2(aq) + H2O2(aq) 2Br-(aq) + 2H+(aq) + O2(g).
– Br- is a catalyst because it can be recovered at the end of the reaction.
• Generally, catalysts operate by lowering the activation energy for a reaction.
Copyright 1999, PRENTICE HALL Chapter 14 48
CatalysisCatalysisHomogeneous CatalysisHomogeneous Catalysis
Copyright 1999, PRENTICE HALL Chapter 14 49
CatalysisCatalysisHomogeneous CatalysisHomogeneous Catalysis• Catalysts can operate by increasing the number of
effective collisions.• That is, from the Arrhenius equation: catalysts
increase k be increasing A or decreasing Ea.
• A catalyst may add intermediates to the reaction.
• Example: In the presence of Br-, Br2(aq) is generated as an intermediate in the decomposition of H2O2.
• When a catalyst adds an intermediate, the activation energies for both steps must be lower than the activation energy for the uncatalyzed reaction.
Copyright 1999, PRENTICE HALL Chapter 14 50
CatalysisCatalysisHeterogeneous CatalysisHeterogeneous Catalysis• The catalyst is in a different phase than the reactants
and products.• Typical example: solid catalyst, gaseous reactants and
products (catalytic converters in cars).• Most industrial catalysts are heterogeneous.• First step is adsorption (the binding of reactant
molecules to the catalyst surface).• Adsorbed species (atoms or ions) are very reactive.• Molecules are adsorbed onto active sites (red spheres)
on the catalyst surface.
Copyright 1999, PRENTICE HALL Chapter 14 51
CatalysisCatalysisHeterogeneous CatalysisHeterogeneous Catalysis
Copyright 1999, PRENTICE HALL Chapter 14 52
CatalysisCatalysisHeterogeneous CatalysisHeterogeneous Catalysis• Consider the hydrogenation of ethylene:
C2H4(g) + H2(g) C2H6(g), H = -136 kJ/mol.– The reaction is slow in the absence of a catalyst.
– In the presence of a metal catalyst (Ni, Pt or Pd) the reaction occurs quickly at room temperature.
– First the ethylene and hydrogen molecules are adsorbed onto active sites on the metal surface.
– The H-H bond breaks and the H atoms migrate about the metal surface.
Copyright 1999, PRENTICE HALL Chapter 14 53
CatalysisCatalysisHeterogeneous CatalysisHeterogeneous Catalysis• Consider the hydrogenation of ethylene:
C2H4(g) + H2(g) C2H6(g), H = -136 kJ/mol.– When an H atom collides with an ethylene molecule on the
surface, the C-C bond breaks and a C-H bond forms.
– When C2H6 forms it desorbs from the surface.
– When ethylene and hydrogen are adsorbed onto a surface, less energy is required to break the bonds and the activation energy for the reaction is lowered.
Copyright 1999, PRENTICE HALL Chapter 14 54
CatalysisCatalysisEnzymesEnzymes• Enzymes are biological catalysts.• Most enzymes are protein molecules with large
molecular masses (10,000 to 106 amu).• Enzymes have very specific shapes.• Most enzymes catalyze very specific reactions.• Substrates undergo reaction at the active site of an
enzyme.• A substrate locks into an enzyme and a fast reaction
occurs.
Copyright 1999, PRENTICE HALL Chapter 14 55
CatalysisCatalysisEnzymesEnzymes
Copyright 1999, PRENTICE HALL Chapter 14 56
CatalysisCatalysisEnzymesEnzymes• The products then move away from the enzyme.• Only substrates that fit into the enzyme lock can be
involved in the reaction.• If a molecule binds tightly to an enzyme so that
another substrate cannot displace it, then the active site is blocked and the catalyst is inhibited (enzyme inhibitors).
• The number of events (turnover number) catalyzed is large for enzymes (103 - 107 per second).
Copyright 1999, PRENTICE HALL Chapter 14 57
CatalysisCatalysisNitrogen Fixation and NitrogenaseNitrogen Fixation and Nitrogenase• Nitrogen gas cannot be used in the soil for plants or
animals.
• Nitrogen compounds, NO3, NO2-, and NO3
- are used in the soil.
• The conversion between N2 and NH3 is a process with a high activation energy (the NN triple bond needs to be broken).
• An enzyme, nitrogenase, in bacteria which live in root nodules of legumes, clover and alfalfa, catalyses the reduction of nitrogen to ammonia.
Copyright 1999, PRENTICE HALL Chapter 14 58
CatalysisCatalysisNitrogen Fixation and NitrogenaseNitrogen Fixation and Nitrogenase
Copyright 1999, PRENTICE HALL Chapter 14 59
CatalysisCatalysisNitrogen Fixation and NitrogenaseNitrogen Fixation and Nitrogenase• The fixed nitrogen (NO3, NO2
-, and NO3-) is consumed
by plants and then eaten by animals.• Animal waste and dead plants are attacked by
bacteria that break down the fixed nitrogen and produce N2 gas for the atmosphere.
Copyright 1999, PRENTICE HALL Chapter 14 60
End of Chapter 14End of Chapter 14
Chemical KineticsChemical Kinetics
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