Chemical Kinetics Chapter 14. Reaction Mechanisms The overall progress of a chemical reaction can be...
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Chemical Kinetics
Chapter 14
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Reaction Mechanisms
The overall progress of a chemical reaction can be represented at the molecular level by a series of simple elementary steps or elementary reactions
The sequence of elementary steps that leads to product formation is the reaction mechanism.
2NO (g) + O2 (g) 2NO2 (g)
N2O2 is detected during the reaction!
Elementary step: NO + NO N2O2
Elementary step: N2O2 + O2 2NO2
Overall reaction: 2NO + O2 2NO2
+
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Elementary step: NO + NO N2O2
Elementary step: N2O2 + O2 2NO2
Overall reaction: 2NO + O2 2NO2
+
Intermediates - species that appear in a reaction mechanism but not in the overall balanced equation
An intermediate is always formed in an early elementary step and consumed in a later elementary step.
Molecularity of a reaction - the number of molecules reacting in an elementary step.
• Unimolecular reaction – elementary step with 1 molecule
• Bimolecular reaction – elementary step with 2 molecules
• Termolecular reaction – elementary step with 3 molecules
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Unimolecular reaction A products rate = k [A]
Bimolecular reaction A + B products rate = k [A][B]
Bimolecular reaction A + A products rate = k [A]2
Rate Laws and Elementary Steps
Writing plausible reaction mechanisms:
• The sum of the elementary steps must give the overall balanced equation for the reaction.
• The rate-determining step should predict the same rate law that is determined experimentally.
Rate-determining step - the slowest step in the sequence of steps leading to product formation.
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The experimental rate law for the reaction between NO2 and CO to produce NO and CO2 is rate = k[NO2]2. The reaction is believed to occur via two steps:
Step 1: NO2 + NO2 NO + NO3
Step 2: NO3 + CO NO2 + CO2
What is the equation for the overall reaction?
NO2+ CO NO + CO2
What is the intermediate?
NO3
What can you say about the relative rates of steps 1 and 2?
rate = k[NO2]2 is the rate law for step 1 so step 1 must be slower than step 2
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Catalyst - a substance that increases the rate of a chemical reaction without itself being consumed
k = A • exp( -Ea/RT ) Ea k
uncatalyzed catalyzed
ratecatalyzed > rateuncatalyzed
Ea < Ea'
Fig 14.17
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Heterogeneous catalysis - the reactants and the catalysts are in different phases.
Homogeneous catalysis - reactants and the catalysts are dispersed in a single phase, usually liquid.
• Haber synthesis of ammonia
• Ostwald process for the production of nitric acid
• Catalytic converters
• Acid catalysis
• Base catalysis
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N2 (g) + 3H2 (g) 2NH3 (g)Fe/Al2O3/K2O
catalyst
Haber ProcessFig 14.18
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Ostwald Process for Making Nitirc Acid
Hot Pt wire over NH3 solutionPt-Rh catalysts used
in Ostwald process
4NH3 (g) + 5O2 (g) 4NO (g) + 6H2O (g)Pt catalyst
2NO (g) + O2 (g) 2NO2 (g)
2NO2 (g) + H2O (l) HNO2 (aq) + HNO3 (aq)
Fig 14.19
Fig on pg 438
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Catalytic Converters
CO + Unburned Hydrocarbons + O2 CO2 + H2Ocatalytic
converter
2NO + 2NO2 2N2 + 3O2
catalyticconverter
Fig 14.20
Fig 14.21
Platinum, palladium, rhodium
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Enzyme CatalysisFig 14.22
Fig 14.23
glucose
hexokinasehexokinase
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Basic reaction steps in enzyme catalysis
E + S ⇌ ES (fast)
ES → P + E (slow)k
uncatalyzedenzyme
catalyzed
Fig 14.24
rate = Δ[P]Δt
rate = k [ES]
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Fig 14.25 Plot of rate of product formation vs
substrate concentration
First order
Zero order
rate k [S]