week 10.3 chemical kinetics

Prepared by:

Mrs Faraziehan Senusi

PA-A11-7C

Collision Model

Catalysis

Chapter 5

Chemical Kinetics

Reaction Rates

Reference: Chemistry: the Molecular Nature of Matter and Change,

6th ed, 2011, Martin S. Silberberg, McGraw-Hill

Rate Laws

Reaction mechanism

Reaction Mechanisms

• The sequence of events that describes the actual

process by which reactants become products is called

the reaction mechanism.

• It is the step-by-step pathway by which a reaction

occurs.

• Reactions may occur all at once or through several

discrete steps.

• Each of these processes is known as an elementary

reaction or elementary step.

• Some reactions take place in a single step, but most

reactions occur in a series of elementary steps.

Elementary Reaction/Step

The molecularity of a process tells how many

molecules are involved in the process.

• A reaction mechanism is defined as a proposed set

of elementary steps, which account for the overall

features of the reaction.

• Each of the reactions that comprises the

mechanism is called an elementary step.

• We believe it is elementary because it takes place

in a single reactive encounter between the

reactants involved.

• These elementary steps are the basic building

blocks of a complex reaction and cannot be broken

down any further.

Example of Reaction Mechanism

Rate Determining Step in

Reaction Mechanism

• In a reaction mechanism, one of the elementary steps will be

slower than all others.

• The overall reaction cannot occur faster than this slowest,

rate-determining step.

• Therefore, this elementary, rate-determining step

establishes the rate of the overall reaction.

• The speed at which the slow step occurs limits the rate at

which the overall reaction occurs.

Slow Initial Step

• The rate law for this reaction is found experimentally

to be

Rate = k [NO2]2

• CO is necessary for this reaction to occur, but the rate

of the reaction does not depend on its concentration.

• This suggests the reaction occurs in two steps.

NO2 (g) + CO (g) NO (g) + CO2 (g)

• A proposed mechanism for this reaction is

Step 1: NO2 + NO2 NO3 + NO (slow)

Step 2: NO3 + CO NO2 + CO2 (fast)

• In this proposed mechanism two molecules of NO2 collide

to produce one molecule each of NO3 and NO.

• The reaction intermediate NO3, then collides with one

molecule of CO and reacts very rapidly to produce one

molecule each of NO2 and CO2.

• The NO3 intermediate is consumed in the second step.

• As CO is not involved in the slow, rate-determining step, it

does not appear in the rate law.

Fast Initial Step

• The rate law for this reaction is found to be

Rate = k [NO]2 [Br2]

• Because termolecular processes are rare, this rate

law suggests a two-step mechanism.

2 NO (g) + Br2 (g) 2 NOBr (g)

• A proposed mechanism is

Step 2: NOBr2 + NO 2 NOBr (slow)

• The first step involves the collision of one NO

molecule (reactant) and one Br2 molecule

(reactant) to produce the intermediate species

NOBr2.

• The NOBr2 can react rapidly, however, to re-form

NO and Br2. We say that this is an equilibrium step

includes the forward and reverse reactions.

• Eventually another NO molecule (reactant) can

collide with a short-lived NOBr2 molecule and

react to produce two NOBr molecules (product).

Step 1: NO + Br2 NOBr2 (fast)

• The rate of the overall reaction depends upon the

rate of the slow step.

• To analyze the rate law that would be consistent

with this proposed mechanism, we again start with

the slow (rate-determining) step, step 2.

• Denoting the rate constant for this step as k2, we

could express the rate of this step as

Rate = k2 [NOBr2] [NO]

But how can we find [NOBr2]? NOBr2 is a reaction intermediate, so its

concentration at the beginning of the second step

may not be easy to measure directly.

• NOBr2 can react two ways:

– With NO to form NOBr

– By decomposition to reform NO and Br2

• The reactants and products of the first step are in

equilibrium with each other.

• Therefore,

Ratef = Rater

• Because Ratef = Rater ,

k1f [NO][Br2] = k1r [NOBr2]

Step 2: NOBr2 + NO 2 NOBr (slow)

Step 1: NO + Br2 ↔ NOBr2 (fast)

• Solving for [NOBr2] gives us

k1f [NO][Br2] = k1r [NOBr2]

• Substituting this expression for [NOBr2] in the rate law for the rate-determining step gives

k1f

k1r [NO] [Br2] = [NOBr2]

k2 k1f

k1r Rate = [NO] [Br2] [NO]

Rate = k [NO]2 [Br2]

rate law for the rate-determining step:

Rate = k2 [NOBr2] [NO]

Catalysts • Catalysts are substances that can be added to reacting

systems to increase the rate of reaction.

• They allow reactions to occur via alternative pathways that

increase reaction rates by lowering activation energies.

• Catalysts change the mechanism by which the process

occurs.

• A catalyst does take part in the reaction, but all of it is

re-formed in later steps.

• Thus, a catalyst does not appear in the balanced

equation for the reaction.

• How does activation energy affects rate of reaction??

k = A e−Ea/RT

Arrhenius Equation

When a catalyst is present,

the energy barrier is

lowered. Thus, more

molecules possess the

minimum kinetic energy

necessary for reaction.

CATALYSIS

• A catalyst changes the rate of a chemical

reaction.

• Two categories of catalysts:

(1) homogeneous catalysts

(2) heterogeneous catalysts

Homogeneous catalysts

• A homogeneous catalyst exists in the same phase as the

reactants.

• Catalyst can operate by increasing the number of effective

collisions.

• That is, from the Arrhenius equation: catalyst increase k by

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 must be lower than the activation energy for the

uncatalyzed reaction.

Heterogeneous catalysts

• A heterogeneous catalyst is present in a different phase

from the reactants.

• Such catalysts are usually solids, and they lower activation

energies by providing surfaces on which reactions can

occur.

• The first step in the catalytic process is usually adsorption,

in which one or more of the reactants become attached to

the solid surface.

• Some reactant molecules may be held in particular

orientations, or some bonds may be weakened; in other

molecules, some bonds may be broken to form atoms or

smaller molecular fragments. This causes activation of the

reactants.

• As a result, reaction occurs more readily than

would otherwise be possible.

• In a final step, desorption, the product molecules

leave the surface, freeing reaction sites to be used

again.

• Most contact catalysts are more effective as small

particles, because they have relatively large

surface areas.

A schematic representation of the catalysis of the reaction on a metallic surface

(Pt, NiO)

2CO (g) + O2(g) 2CO2(g)

Enzymes

• Most enzymes are protein

molecules with large

molecular masses (10,000

to 106 amu).

• The substrate fits into the

active site of the enzyme

much like a key fits into a

lock.

• Enzymes are proteins

that act as catalysts for

specific biochemical

reactions in living

systems.

• The reactants in enzyme-

catalyzed reactions are

called substrates.

• 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.

• The products then move away from the enzyme.

• Only substrates that fit into the enzyme lock can

be involved in the reaction.

A space-filling model of the enzyme lysozyme. This

enzyme catalyzes the hydrolysis of polysaccharides

(complex carbohydrates) found in bacterial cell walls.