kinetics part ii - lebanon high schoollhs.sau88.net/ourpages/auto/2018/1/16/57190087/unit 6_kinetics...

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Kinetics: Part II

Collision Theory,

Reaction Mechanisms,

& the Arrhenius Equation

How to reactions actually happen?

Collision Theory &

Reaction Mechanisms

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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.

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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.

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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).

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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.

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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

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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

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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


Change in Potential Energy


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Section 12.6


Transition States and Activation Energy


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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.

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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


Arrhenius Equation

A = frequency factor

Ea = activation energy

R = gas constant (8.3145 J/K·mol)

T = temperature (in K)


/ =

−a

E RTk Ae

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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.


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How to reactions actually happen?

Collision Theory &

Reaction Mechanisms


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.

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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.


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.

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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.


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]

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Rate Laws of Elementary Steps


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.

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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


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.

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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.


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

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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.


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)


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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.


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.


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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.


Section 12.5

Reaction Mechanisms

Decomposition of N2O5




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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)


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


CONCEPT CHECK!

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Section 12.6


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.


Section 12.6


The Gas Phase Reaction of NO and Cl2




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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.


Section 12.7

Catalysis

Energy Plots for a Catalyzed and an Uncatalyzed Pathway

for a Given Reaction


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Section 12.7

Catalysis

Effect of a Catalyst on the Number of Reaction-Producing

Collisions


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.


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Section 12.7

Catalysis

Heterogeneous Catalysis




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.


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Section 12.7

Catalysis

Homogeneous Catalyst

� Exists in the same phase as the reacting molecules.

� Enzymes are nature’s catalysts.


Section 12.7

Catalysis

Homogeneous Catalysis




kinetics part ii - lebanon high schoollhs.sau88.net/ourpages/auto/2018/1/16/57190087/unit 6_kinetics...

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