Prepared by:Mrs Faraziehan Senusi

PA-A11-7CCollision 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.