Reaction Mechanism
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Transcript of Reaction Mechanism
Reaction Mechanism
The reaction mechanism is the series of elementary steps by which a chemical reaction occurs.
The sum of the elementary steps must give the overall balanced equation for the reaction
The mechanism must agree with the experimentally determined rate law
Rate-Determining Step
In a multi-step reaction, the slowest step is the rate-determining step. It therefore determines the rate of the reaction.The experimental rate law must agree with the rate-determining step
Identifying the Rate-Determining Step
For the reaction:2H2(g) + 2NO(g) N2(g) +
2H2O(g)The experimental rate law is:
R = k[NO]2[H2]Which step in the reaction mechanism is the rate-determining (slowest) step?
Step #1 H2(g) + 2NO(g) N2O(g) + H2O(g)Step #2 N2O(g) + H2(g) N2(g) + H2O(g)
Step #1 agrees with the experimental rate law
Identifying IntermediatesFor the reaction:
2H2(g) + 2NO(g) N2(g) + 2H2O(g)
Which species in the reaction mechanism are intermediates (do not show up in the final, balanced equation?)
Step #1 H2(g) + 2NO(g) N2O(g) + H2O(g)Step #2 N2O(g) + H2(g) N2(g) + H2O(g)
2H2(g) + 2NO(g) N2(g) + 2H2O(g)
N2O(g) is an intermediate
Identifying CatalystsFor the reaction:
O3(g) + O(g) 2O2
Intermediates are products in step 1 and reactants in step 2Catalysts are reactants in step 1 and products in step 2
Step #1 O3(g) + NO(g) NO2(g) + O2(g)Step #2 NO2(g) + O(g) NO(g) + O2(g)
O3(g) + O(g) 2O2
NO is the catalyst
Collision ModelKey Idea: Molecules must collide to react.However, only a small fraction of collisions produces a reaction. Why?
Collision ModelCollisions must have sufficient energy to produce the reaction (must equal or exceed the activation energy).Colliding particles must be correctly oriented to one another in order to produce a reaction.
Factors Affecting RateIncreasing temperature always increases the rate of a reaction.
Particles collide more frequently Particles collide more energeticallyIncreasing surface area increases
the rate of a reactionIncreasing Concentration USUALLY increases the rate of a reactionPresence of Catalysts, which lower the activation energy by providing alternate pathways
Arrhenius Equation k = Ae-E
a/RT , k is rate constant, A is
frequency factor, e-Ea/RT is the
fraction of collisions with sufficient energy to produce a reaction, Ea is activation energy, R is 8.31 J/mol K
ln(k) = -Ea/R(1/T) + ln(A) linear equation
ln(k2/k1) = Ea/R(1/T1 – 1/T2)
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.
Lowering of Activation Energy by a Catalyst
Catalysts Increase the Number of Effective Collisions
Heterogeneous Catalysis
Step #1: Adsorption and
activation of the reactants.
Carbon monoxide and nitrogen
monoxide adsorbed on a
platinum surface
Heterogeneous Catalysis
Step #2: Migration of the
adsorbed reactants on the
surface.
Carbon monoxide and nitrogen
monoxide arranged prior to
reacting
Heterogeneous Catalysis
Step #3: Reaction of the
adsorbed substances.
Carbon dioxide and nitrogen form
from previous molecules
Heterogeneous Catalysis
Step #4: Escape, or
desorption, of the products.
Carbon dioxide and nitrogen gases escape
(desorb) from the platinum surface