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