-
1/28/2015
1
Kinetics: Part II
Collision Theory,
Reaction Mechanisms,
& the Arrhenius Equation
How to reactions actually happen?
Collision Theory &
Reaction Mechanisms
-
1/28/2015
2
Collision Theory of Kinetics
• For most reactions, for a reaction to take place, the reacting molecules must collide with each other.
(On average, there are about 109 collisions per second!)
• Once molecules collide they may react together or they may not, depending on two factors:
– Whether the collision has enough kinetic energy to "break the bonds holding reactant molecules together"
– Whether the reacting molecules collide in the proper orientation for new bonds to form
Effective Collisions: Kinetic Energy Factor
For a collision to lead
to overcoming the energy barrier, the reacting molecules
must have sufficient kinetic energy so that
when they collide the
activated complex can form.
-
1/28/2015
3
Effective Collisions: Orientation Effect
Effective Collisions
• Collisions in which these two conditions are met (and therefore lead to reaction) are called effective collisions.
• The higher the frequency of effective collisions, the faster the reaction rate.
• When two molecules have an effective collision, a temporary, high energy (unstable) chemical species is formed— the activated complex.
-
1/28/2015
4
What factors affect Reaction Rate?
Try to think about these answers in light of collision theory.
Factors Affecting Reaction Rate:1. Reactant Concentration
• Generally, the larger the concentration of reactant molecules, the faster the reaction. – This increases the frequency of reactant molecule
contact.
– Concentration of gases depends on the partial pressure of the gas.
Higher Pressure = Higher Concentration
• Concentrations of solutions depend on the solute-to-solution ratio (molarity).
-
1/28/2015
5
Factors Affecting Reaction Rate: 2. Nature of the Reactants
• Nature of the reactants means what kind of reactant molecules and what physical condition they are in. – Small molecules tend to react faster than large molecules.
– Gases tend to react faster than liquids, which react faster than solids.
– Powdered solids are more reactive than “blocks.”
• More surface area for contact with other reactants
– Certain types of chemicals are more reactive than others.
• For example, potassium metal is more reactive than sodium
– Ions react faster than molecules.
• No bonds need to be broken.
Factors Affecting Reaction Rate:3. Catalysts
• Catalysts are substances that affect the speed of a reaction without being consumed.
• Most catalysts are used to speed up a reaction; these are called positive catalysts.
– Catalysts used to slow a reaction are called negative
catalysts.
• Homogeneous = present in same phase
• Heterogeneous = present in different phase
• How catalysts work will be examined later.
-
1/28/2015
6
Factors Affecting Reaction Rate:4. Temperature
• Increasing temperature increases the reaction rate.
This should make sense to us:
Temperature is kinetic energy – higher KE means more forceful
collisions, and a higher percentage of effective collisions.
– Chemist’s rule—for each 10 °C rise in temperature, the speed of the reaction doubles.
• There is a mathematical relationship between the absolute temperature and the speed of a reaction discovered by Svante Arrhenius, which will be examined… right now!
• Changing the temperature changes the rate constant of the rate law.
• Svante Arrhenius investigated this relationship and showed the following:
R is the gas constant in energy units, 8.314 J/(mol · K).
where T is the temperature in kelvins.
A is called the frequency factor, the rate the reactant energy approaches the activation energy.
Ea is the activation energy, the extra energy needed to start the molecules reacting.
The Effect of Temperature on Rate
-
1/28/2015
7
Section 12.6
A Model for Chemical Kinetics
Activation Energy, Ea
� Energy that must be overcome to produce a chemical
reaction.
Copyright © Cengage Learning. All rights reserved 13
-
1/28/2015
8
Activation Energy and the Activated Complex
• There is an energy barrier to almost all reactions.
• The activation energy is the amount of energy needed to convert reactants into the activated complex.
– The activated complex is also know as, transition state
• The activated complex is a chemical species with partially broken and partially formed bonds.
– Always very high in energy because of its partial bonds
Section 12.6
A Model for Chemical Kinetics
Change in Potential Energy
Copyright © Cengage Learning. All rights reserved 16
-
1/28/2015
9
Section 12.6
A Model for Chemical Kinetics
Transition States and Activation Energy
Copyright © Cengage Learning. All rights reserved 17
To play movie you must be in Slide Show ModePC Users: Please wait for content to load, then click to play
Mac Users: CLICK HERE
© 2014 Pearson Education, Inc.
An example of an Activated Complex: Isomerization of Methyl Isonitrile
Methyl isonitrile rearranges to acetonitrile.
For the reaction to occur, the
H3C─N bond must break, and a new H3C─C bond must form.
-
1/28/2015
10
© 2014 Pearson Education, Inc.
As the reaction
begins, the C─N
bond weakens
enough for the
C≡N group to
start to rotate.
An example of an Activated Complex: Isomerization of Methyl Isonitrile – Energy Profile
The activated complex is a chemical species with partial bonds.
Section 12.6
A Model for Chemical Kinetics
Arrhenius Equation
A = frequency factor
Ea = activation energy
R = gas constant (8.3145 J/K·mol)
T = temperature (in K)
Copyright © Cengage Learning. All rights reserved 20
/ =
−a
E RTk Ae
-
1/28/2015
11
© 2014 Pearson Education, Inc.
The Arrhenius Equation: The Exponential Factor
• The exponential factor in the Arrhenius equation is a number between 0 and 1. It represents the fraction of reactant molecules with sufficient energy so they can make it over the energy barrier.
The Ea term:
– The higher the energy barrier (larger activation energy), the fewer molecules that have sufficient energy to overcome it.
The T term:
– Increasing the temperature increases the average kinetic energy of the molecules, which leads to a greater # of effective collisions.
– In an effective collision, the kinetic energy of the particles is transformed into potential energy of the activated complex.
– Therefore, increasing the temperature will increase the reaction rate.
© 2014 Pearson Education, Inc.
-
1/28/2015
12
© 2014 Pearson Education, Inc.
How to reactions actually happen?
Collision Theory &
Reaction Mechanisms
© 2014 Pearson Education, Inc.
Reaction Mechanisms
• We generally describe chemical reactions with an equation listing all the reactant molecules and product molecules.
• But most reactions occur in a series of smaller elementary reactions involving one, two, or (at most) three molecules.
• Describing the series of reactions that occurs to produce the overall observed reaction is called a reaction mechanism.
• Knowing the rate law of the reaction helps us understand the sequence of reactions in the mechanism.
-
1/28/2015
13
© 2014 Pearson Education, Inc.
An Example of a Reaction Mechanism
• Overall reaction:
H2(g) + 2 ICl(g) → 2 HCl(g) + I2(g)
• Mechanism:
1. H2(g) + ICl(g) → HCl(g) + HI(g)2. HI(g) + ICl(g) → HCl(g) + I2(g)
• The reactions in this mechanism are elementary steps, meaning that they cannot be broken down into simpler steps and that the molecules actually interact directly in this manner without any other steps.
© 2014 Pearson Education, Inc.
H2(g) + 2 ICl(g) → 2 HCl(g) + I2(g)
1) H2(g) + ICl(g) → HCl(g) + HI(g)
2) HI(g) + ICl(g) → HCl(g) + I2(g)
Elements of a Mechanism Intermediates
• Notice that the HI is a product in step 1, but then a
reactant in step 2.
• Because HI is made but then consumed, HI does not
show up in the overall reaction.
• Materials that are products in an early mechanism step,
but then reactants in a later step, are called
intermediates.
-
1/28/2015
14
© 2014 Pearson Education, Inc.
Molecularity
• The number of reactant particles in an elementary step is called its molecularity.
• A unimolecular step involves one particle.
• A bimolecular step involves two particles.– However, they may be the same kind of particle.
• A termolecular step involves three particles.– However, these are exceedingly rare in elementary steps.
© 2014 Pearson Education, Inc.
Rate Laws for Elementary Steps
• Each step in the mechanism is like its own little reaction with its own activation energy and own rate law.
• The rate law for an overall reaction must be determined experimentally.
• But the rate law of an elementary reaction can be deduced from its molecularity!
H2(g) + 2 ICl(g) → 2 HCl(g) + I2(g)1) H2(g) + ICl(g) → HCl(g) + HI(g) Rate = k1[H2][ICl] 2) HI(g) + ICl(g) → HCl(g) + I2(g) Rate = k2[HI][ICl]
-
1/28/2015
15
© 2014 Pearson Education, Inc.
Rate Laws of Elementary Steps
© 2014 Pearson Education, Inc.
Rate Determining Step
• In most mechanisms, one step occurs slower than the other steps.
• The result is that product production cannot occur any faster than the slowest step; the step determines the rate of the overall reaction.
• We call the slowest step in the mechanism the rate determining step.– The slowest step has the largest activation energy.
• The rate law of the rate determining step determines the rate law of the overall reaction.
-
1/28/2015
16
© 2014 Pearson Education, Inc.
Another Reaction Mechanism
The first step in this mechanism is the rate determining step.
The first step is slower than the
second step because its
activation energy is larger.
The rate law of the first step is the same as the rate law of the overall reaction.
NO2(g) + CO(g) → NO(g) + CO2(g) Rateobs = k[NO2]21. NO2(g) + NO2(g) → NO3(g) + NO(g) Rate = k1[NO2]2 Slow
2. NO3(g) + CO(g) → NO2(g) + CO2(g) Rate = k2[NO3][CO] Fast
© 2014 Pearson Education, Inc.
Validating a Mechanism
• To validate (not prove) a mechanism, two conditions must be met:
1. The elementary steps must sum to the overall reaction.
2. The rate law predicted by the mechanism must be consistent with the experimentally observed rate law.
-
1/28/2015
17
© 2014 Pearson Education, Inc.
Mechanisms with a Fast Initial Step
• When a mechanism contains a fast initial step, the rate limiting step may contain intermediates.
• When a previous step is rapid and reaches equilibrium, the forward and reverse reaction rates are equal, so the concentrations of reactants and products of the step are related and the product is an intermediate.
• Substituting into the rate law of the RDS will produce a rate law in terms of just reactants.
© 2014 Pearson Education, Inc.
An Example
1. 2 NO(g) ⇔ N2O2(g) Fast
2. H2(g) + N2O2(g) → H2O(g) + N2O(g) Slow Rate = k2[H2][N2O2]
3. H2(g) + N2O(g) → H2O(g) + N2(g) Fast
k1
k−1
2 H2(g) + 2 NO(g) → 2 H2O(g) + N2(g) Rateobs = k [H2][NO]2
-
1/28/2015
18
Section 12.5
Reaction Mechanisms
Reaction Mechanism
� Most chemical reactions occur by a series of elementary
steps.
� An intermediate is formed in one step and used up in a
subsequent step and thus is never seen as a product in
the overall balanced reaction.
Copyright © Cengage Learning. All rights reserved 35
Section 12.5
Reaction Mechanisms
A Molecular Representation of the Elementary Steps in the Reaction
of NO2 and CO
NO2(g) + CO(g) → NO(g) + CO2(g)
Copyright © Cengage Learning. All rights reserved 36
-
1/28/2015
19
Section 12.5
Reaction Mechanisms
Elementary Steps (Molecularity)
� Unimolecular – reaction involving one molecule; first
order.
� Bimolecular – reaction involving the collision of two
species; second order.
� Termolecular – reaction involving the collision of three
species; third order. Very rare.
Copyright © Cengage Learning. All rights reserved 37
Section 12.5
Reaction Mechanisms
Rate-Determining Step
� A reaction is only as fast as its slowest step.
� The rate-determining step (slowest step) determines the
rate law and the molecularity of the overall reaction.
Copyright © Cengage Learning. All rights reserved 38
-
1/28/2015
20
Section 12.5
Reaction Mechanisms
Reaction Mechanism Requirements
� 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.
Copyright © Cengage Learning. All rights reserved 39
Section 12.5
Reaction Mechanisms
Decomposition of N2O5
Copyright © Cengage Learning. All rights reserved 40
To play movie you must be in Slide Show ModePC Users: Please wait for content to load, then click to play
Mac Users: CLICK HERE
-
1/28/2015
21
Section 12.5
Reaction Mechanisms
Decomposition of N2O5
2N2O5(g) � 4NO2(g) + O2(g)
Step 1: N2O5 NO2 + NO3 (fast)
Step 2: NO2 + NO3 → NO + O2 + NO2 (slow)
Step 3: NO3 + NO → 2NO2 (fast)
Copyright © Cengage Learning. All rights reserved 41
2( )
Section 12.5
Reaction Mechanisms
The reaction A + 2B � C has the following proposed
mechanism:
A + B D (fast equilibrium)
D + B � C (slow)
Write the rate law for this mechanism.
rate = k[A][B]2
Copyright © Cengage Learning. All rights reserved 42
CONCEPT CHECK!
-
1/28/2015
22
Section 12.6
A Model for Chemical Kinetics
For Reactants to Form Products
� Collision must involve enough energy to produce the
reaction (must equal or exceed the activation energy).
� Relative orientation of the reactants must allow
formation of any new bonds necessary to produce
products.
Copyright © Cengage Learning. All rights reserved 43
Section 12.6
A Model for Chemical Kinetics
The Gas Phase Reaction of NO and Cl2
Copyright © Cengage Learning. All rights reserved 44
To play movie you must be in Slide Show ModePC Users: Please wait for content to load, then click to play
Mac Users: CLICK HERE
-
1/28/2015
23
Section 12.7
Catalysis
Catalyst
� A substance that speeds up a reaction without being
consumed itself.
� Provides a new pathway for the reaction with a lower
activation energy.
Copyright © Cengage Learning. All rights reserved 45
Section 12.7
Catalysis
Energy Plots for a Catalyzed and an Uncatalyzed Pathway
for a Given Reaction
Copyright © Cengage Learning. All rights reserved 46
-
1/28/2015
24
Section 12.7
Catalysis
Effect of a Catalyst on the Number of Reaction-Producing
Collisions
Copyright © Cengage Learning. All rights reserved 47
Section 12.7
Catalysis
Heterogeneous Catalyst
� Most often involves gaseous reactants being adsorbed
on the surface of a solid catalyst.
� Adsorption – collection of one substance on the surface
of another substance.
Copyright © Cengage Learning. All rights reserved 48
-
1/28/2015
25
Section 12.7
Catalysis
Heterogeneous Catalysis
Copyright © Cengage Learning. All rights reserved 49
To play movie you must be in Slide Show ModePC Users: Please wait for content to load, then click to play
Mac Users: CLICK HERE
Section 12.7
Catalysis
Heterogeneous Catalyst
1. Adsorption and activation of the reactants.
2. Migration of the adsorbed reactants on the surface.
3. Reaction of the adsorbed substances.
4. Escape, or desorption, of the products.
Copyright © Cengage Learning. All rights reserved 50
-
1/28/2015
26
Section 12.7
Catalysis
Homogeneous Catalyst
� Exists in the same phase as the reacting molecules.
� Enzymes are nature’s catalysts.
Copyright © Cengage Learning. All rights reserved 51
Section 12.7
Catalysis
Homogeneous Catalysis
Copyright © Cengage Learning. All rights reserved 52
To play movie you must be in Slide Show ModePC Users: Please wait for content to load, then click to play
Mac Users: CLICK HERE