1 chapter 12 – chemical kinetics 1.second order rate law 2.zero order rate law 3.reaction...
TRANSCRIPT
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Chapter 12 – Chemical Kinetics
1. Second order Rate Law2. Zero Order Rate Law3. Reaction Mechanism4. Model for Chemical Kinetics5. Collision6. Catalysis7. Heterogeneous Catalysis8. Homogeneous Catalysis
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Second-Order Rate Law• For aA products in a second-order
reaction,
• Integrated rate law is
• Plot of 1/[A] vs t will produce a straight line: slope = k
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Half-Life of a 2nd-Order Reaction
•t1/2 = half-life of the reaction•k = rate constant•Ao = initial concentration of A
•The half-life is dependent upon the initial concentration.
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Second-Order Rate Law – 12.5
Butadiene reacts to form its dimer
2C4H6(g) C8H12(g)
• Data:
[C4H6] Time
0.01000 0
0.00625 1000
0.00476 1800
0.00370 2800
0.00313 3600
0.00270 4400
0.00241 5200
0.00208 6200
a) Reaction order?b) Value of k?c) Half-life?
1/[C4H6] ln[C4H6]
100 -4.605
160 -5.075
210 -5.348
270 -5.599
320 -5.767
370 -5.915
415 -6.028
481 -6.175
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Second-Order Rate Law – 12.5
Butadiene reacts to form its dimer
2C4H6(g) C8H12(g)
• Data:
[C4H6] Time
0.01000 0
0.00625 1000
0.00476 1800
0.00370 2800
0.00313 3600
0.00270 4400
0.00241 5200
0.00208 6200
a) Reaction order?
Rate = k[C4H6]2
1/[C4H6] ln[C4H6]
100 -4.605
160 -5.075
210 -5.348
270 -5.599
320 -5.767
370 -5.915
415 -6.028
481 -6.175
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Second-Order Rate Law – 12.5
Butadiene reacts to form its dimer
2C4H6(g) C8H12(g)
• Data:
[C4H6] Time
0.01000 0
0.00625 1000
0.00476 1800
0.00370 2800
0.00313 3600
0.00270 4400
0.00241 5200
0.00208 6200
b) Value of k?k = slope
1/[C4H6] ln[C4H6]
100 -4.605
160 -5.075
210 -5.348
270 -5.599
320 -5.767
370 -5.915
415 -6.028
481 -6.175
smol
L
s
molLk
.6200
381
)0.6200(
/)100481(
smol
Lk
0614.0
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Second-Order Rate Law – 12.5
Butadiene reacts to form its dimer
2C4H6(g) C8H12(g)
k = 6.14x10-2 L/mol*s[A]0= 1.000x10-2 mol/L
c) Half-life?
02/1 ][
1
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Lmol
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t22
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t 3
02/1 1063.1
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Zero-Order Rate Law
• Zero-order reaction the rate is constant.
• Rate does not change with respect to concentration
• Rate = k[A]o = k(1) = k• [A] = -kt + [A]o
• t1/2 = [A]o/2k
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Reaction Mechanism
The series of steps by which a chemical reaction occurs.
A chemical equation does not tell us how reactants become products - it is a summary of the overall process.
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Often Used Terms•Intermediate: formed in one step and used up in a subsequent step and so is never seen as a product.•Molecularity: the number of species that must collide to produce the reaction indicated by that step.
•Elementary Step: A reaction whose rate law can be written from its molecularity.
•uni, bi and termolecular
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Reaction Mechanism
We can define a reaction mechanism.It is a series of elementary steps that must
satisfy two requirements:1. The sum of the elementary steps must give
the overall balanced equation for the reactions.
2. The mechanism must agree with the experimentally determined rate law.
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Reaction Mechanism The reaction NO2(g) + CO(g) NO(g) + CO2(g)• has many steps in the reaction mechanism.• Rate = k[NO2]2
NO2(g) + NO2(g) NO3(g) + NO(g)NO3(g) + CO(g) NO2(g) + CO2(g)NO2(g) + NO2(g) + NO3(g) + CO(g)
NO3(g) + NO(g) + NO2(g) + CO2(g)Overall reaction:NO2(g) + CO(g) NO(g) + CO2(g)
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Rate-Determining Step•In a multistep reaction, it is the slowest step. It therefore determines the rate of reaction.
– Slow (rate determing)NO2(g) + NO2(g) NO3(g) + NO(g) – FastNO3(g) + CO(g) NO2(g) + CO2(g)Rate of formation of NO3 = [NO3]/t = K1[NO2]2 Overall rate = [NO3]/t = k1[NO2]2
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A Summary (continued)
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Model for Chemical KineticsCollision Model
•Key Idea: Molecules must collide to react.
•However, only a small fraction of collisions produces a reaction. Why?
•Arrhenius: An activation energy must be overcome.
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Reaction ProgressThe arrangement of atoms found at the top of
the potential energy barrier is call the activated complex or transition state.
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Reaction ProgressE has no effect on the rate of reaction.Rate depends on the size of the activation
energy Ea.
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Reaction Progress
A certain minimum energy is required for the molecules to “get over the hill”
At a given temperature only a fraction of the collisions possess enough energy to be effective
- Lower temperature
-Effective collisions - Small
- Higher temperature
-Effective collisions – Increase exponentially
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Reaction Progress - CollisionsCollisions with Ea = (Total Collisions)e-Ea/RT
Observed rate smaller than the rate of collisions with Ea
Molecular Orientations:
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Reaction Progress - Collisions
Requirements must satisfied for reactants to collide successfully1. Collisions must involve enough energy to produce the
reaction; Collision energy must equal or exceed the activation energy
2. Relative orientation of reactants must allow for the formation of any new bonds necessary to produce products.
Rate constant: k = zpe-Ea/RT
z = collision frequency R = 8.3145 J/K molp = steric factor (fraction with effective orientation)
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Catalysis
•Catalyst: A substance that speeds up a reaction without being consumed
•Enzyme: A large molecule (usually a protein) that catalyzes biological reactions.
•Homogeneous catalyst: Present in the same phase as the reacting molecules.
•Heterogeneous catalyst: Present in a different phase than the reacting molecules.
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Catalysis
Increases the number of effective collisions by providing a reaction pathway with a lower activation energy
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Heterogeneous Catalysis
Heterogeneous catalysis most often involves gaseous reactant being absorbed on the surface of a solid surface.
Hydrogenation is an example:Changes C=C into saturated H-C-C-H
1. Adsorption to addition of a substance to the surface of another.
2. Migration of absorbed reactants on the surface
3. Reaction of absorbed substances
4. Desorption of products
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Heterogeneous catalytic ethylene hydrogenation: C2H4 + H2 → C2H6
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Homogeneous Catalysis
• Reactants and Catalysis are in the same phase. Gas-Gas or Liquid-Liquid
• N2(g) + O2(g) 2NO(g)
• Product of high-temperature combustion when N2 is present. However catalytic in production of ozone
• 2NO(g) + O2(g) 2NO2(g)
• NO2(g) NO(g) + O(g) (light)
• O2(g) + O(g) O3(g)
• 3/2O2(g) O3(g)
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Homogeneous Catalysis
• In the upper atmosphere, NO has opposite effect.
• 2NO(g) + O3(g) NO2(g) + O2(g)
• O + NO2(g) NO(g) + O2(g)
• O3(g) + O(g) 2O2(g)
• Nitric Oxide is catalytic in production of O2.
• O3 required in upper atmosphere to block uv radiation.
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Homogeneous Catalysis
• Freon - ChloroFluoroCarbons
• CCl2F2(g) CClF2(g) + Cl(g) Light
• Cl + O3(g) ClO(g) + O2(g)
• O(g) + ClO(g) Cl(g) + O2(g)
• O3(g) + O(g) 2O2(g)
• Cl(g) destroyer of ozone – catalytic in destruction of O3.