Chapter 14 Chemical Kinetics. Review Section of Chapter 14 Test Net Ionic Equations

Chapter 14

Chemical Kinetics

Review Section of Chapter 14 Test

• Net Ionic Equations

Reaction Rate• The rate of a chemical reaction is measured as

the decrease in the concentration of a reactant or the increase in the concentration of a product in a unit of time.



∆[ ]

∆timeRate =



∆[ ]

∆timeRate =

What units would What units would we use for rate?we use for rate?



2H2H22OO22(aq) → 2H(aq) → 2H22O(l) + OO(l) + O22(g)(g)

∆[ ]

∆timeRate =




∆[ ]

∆timeRate =

How could the rate be expressed for How could the rate be expressed for this reaction in terms of Hthis reaction in terms of H22OO22??




What is the rate of the reaction from 0s to 2.16 x 10What is the rate of the reaction from 0s to 2.16 x 1044s?s?


What is the average rate of appearance of OWhat is the average rate of appearance of O22 from 0s to 2.16 x 10 from 0s to 2.16 x 1044s?s?

General Rate of Reactiona A + b B → c C + d D

Rate of reaction = rate of disappearance of reactants

We can use the coefficients in the equation to compare the reaction rates for all the substances in the reaction.

Rate of reaction = rate of appearance (formation) of productsor

15-1 The Rate of a Chemical Reaction

• Rate is change of concentration with time.

2 Fe3+(aq) + Sn2+(aq) → 2 Fe2+(aq) + Sn4+(aq)

t = 38.5 s [Fe2+] = 0.0010 M

Rate of formation of Fe2+= = = 2.6 x 10-5 M s-1Δ[Fe2+]

Δt

0.0010 M

38.5 s

∆t = 38.5 s ∆[Fe2+] = (0.0010 – 0) M

Rates of Chemical Reaction2 Fe3+(aq) + Sn2+(aq) → 2 Fe2+(aq) + Sn4+(aq)

Rate of formation of Fe2+ = 2.6 x 10-5 mol L-1 s-1

What is the rate of formation of Sn4+?

What is the rate of disappearance of Fe3+?

What does the slope of the line represent?

What is the concentration at 100s for the reaction: 2H2O2(aq) → 2H2O(l) + O2(g)?

Given:

-Δ[H2O2] = (1.7 x 10-3 M s-1) (∆t)

Rate = 1.7 x 10-3 M s-1

Δt=

- Δ[H2O2]

∆[H2O2] = -(1.7 x 10-3 M s-1)(100 s) = -0.17M

= 2.15 M

= 2.32 M - 0.17 M [H2O2]100 s

[H2O2]i = 2.32 M

What does it mean when the rate of a reaction reaches zero?

• For a normal reaction it means that one or more of the reactants are used up and the reaction has stopped.

• For a reversible reaction it means that the reaction has reached equilibrium.

Factors Affecting Reaction Rates

1.1. The nature of the The nature of the reacting substances.reacting substances.

Factors Affecting Reaction Rates2.2. The state of subdivision of the The state of subdivision of the

reacting substances.reacting substances.

Factors Affecting Reaction Rates3. The temperature of the 3. The temperature of the

reacting substances.reacting substances.

http://en.wikipedia.org/wiki/File:Large_bonfire.jpg


4. The concentration of the reacting substances.4. The concentration of the reacting substances.

Air (21% oxygen)Air (21% oxygen) 100% oxygen 100% oxygen


5.5. The presence of a catalyst.The presence of a catalyst.

CatalystsCatalysts speed up reactions speed up reactions but are left unchanged by the but are left unchanged by the reaction.reaction.

The Rate Law

a A + b B …. → g G + h H ….

Rate = k [A]m[B]n ….

Rate constant = k (k is constant for a particular reaction at a specific temperature)

Order of A = m Order of B = n

Overall order of reaction = m + n + ….

Temperature and Rate• Generally, as temperature

increases, so does the reaction rate.

• This is because k is temperature dependent.

Use the data provided to write the rate law and indicate the order of the reaction with respect to HgCl2 and C2O4

2- and also the overall order of the reaction.

First determine the order of HgCl2

Next determine the order of C2O42-

Now write the rate law and determine the order of the reaction.

Calculate the rate constant “k” and its units.

Initial rate of disappearance HgCl2

mol L-1 min-1

What is the average rate of disappearance of C2O42- in trial 1?

Initial rate of disappearance HgCl2

mol L-1 min-1

Use the data provided to write the rate law and indicate the order of the reaction with respect to NO2 and CO (support your answers). Also give the overall order of the reaction.

(Example is in notebook)

Calculate the rate constant “k” and its units.

What is the average rate of disappearance of CO in trial 2?

• After finding the trials to compare:• A reaction is zero order in a reactant if the change in

concentration of that reactant produces no effect on the rate.• A reaction is first order if doubling the concentration causes

the rate to double.• A reaction is nth order if doubling the concentration causes

an 2n increase in rate.• Note that the rate constant does not depend on

concentration.

Concentration and Rate SummaryConcentration and Rate Summary

Collision Model

•Key Idea: Molecules must collide to react.

•However, only a small fraction of collisions produces a reaction. Why?

Two Factors

- Collisions must have enough energy to produce the reaction (must equal or exceed the activation energy).

- Orientation of reactants must allow formation of new bonds.

2HI → H2 + 2I

Concentration and Collision Theory

• Why does an increase in concentration cause an increase in reaction rate?

Concentration and Collision Theory

• Why does an increase in concentration cause an increase in reaction rate?

• Increasing the concentration increases the number of collisions and therefore there are more collisions leading to product.

Temperature and Collision Theory

• Why does a temperature increase cause the reaction rate to increase?

Temperature and Collision Theory

• Why does a temperature increase cause the reaction rate to increase?

• At higher temperatures there are more collisions and a greater percentage of the collisions have the energy necessary to create a successful collision.

Activation Energy

• The activation energy is the minimum amount of energy necessary for a reaction to occur.

Temperature and Activation Energy (Ea)Figure 14.12 (Page 432)

Activation Energy

• The activation energy can also be thought of as the energy necessary to form an activated complex during a collision between reactants.

Transition State Theory• The activated complex is a hypothetical species

lying between reactants and products at a point on the reaction profile called the transition state.

The activated complex is a transition state between reactants and products where old bonds have begun to break and new bonds have started to form. It cannot be isolated.

Determining the Activation Energy“The Arrhenius Equation”

- Collisions must have enough energy to produce the reaction (must equal or exceed the activation energy).

- Orientation of reactants must allow formation of new bonds.

Arrhenius Equation

• k = rate constant• A = frequency factor

• Ea = activation energy

• T = temperature• R = ideal gas constant

k Ae E RT a /

frequency factor: a value in the Arrhenius equation indicating how many collisions have the correct orientation to lead to products.

A

Arrhenius EquationDetermination of Activation Energy

• Graphical determination of activation energy (Ea).

– plot the ln k on the y-axis.– Plot 1/T (use Kelvin temperature) on the x-axis.

Arrhenius EquationDetermination of Activation Energy

• A plot of ln k versus 1/T (using Kelvin) will have:– slope of –Ea/R – y-intercept of ln A.

Slope = - Ea

R

ln A

Example 14.12 Page 433


x-axis y-axis


x1,y1 = 1.25 x 10-3, -2.593

x2,y2 = 1.78 x 10-3, -14.447


For two reactions at the same temperature, the reaction with the higher activation energy has the lower rate constant (k) and the slower rate.

2O3 3O2

- A chemical equation like the one above does not tell us how reactants become products - it is simply a summary of the overall reaction.

The reaction: 2O3 3O2

- Is proposed to occur through the two step process given below:

O3 O2 + O

O3 + O 2O2

This two step process is an example of a reaction mechanism

Reaction Mechanisms

• A reaction mechanism is a step-by-step description of a chemical reaction.

• Each step is called an elementary reaction.

Often Used Terms

•Intermediate: formed in one step and used up in a subsequent step and so is never seen as a product.

•Molecularity: the number of species that must collide to produce the reaction indicated by that step.

•Elementary Step: A step within a reaction mechanism whose rate law can be written from its molecularity.

Elementary Steps• Molecularity: the number of molecules present in an

elementary step.– Unimolecular: one molecule in the elementary step,

– Bimolecular: two molecules in the elementary step, and

– Termolecular: three molecules in the elementary step.

• It is not common to see termolecular processes (statistically improbable).

Reaction MechanismsReaction Mechanisms

Rate Laws for Elementary Steps• The rate law of an elementary step is determined by its

molecularity:– Unimolecular processes are first order,

– Bimolecular processes are second order, and

– Termolecular processes are third order.


Rate Laws for Elementary Steps


The Rate Determining Step

Rate-Determining Step

•In a reaction mechanism, the rate determining step is the slowest step. It therefore determines the rate of reaction.

Reaction Mechanisms

• Reaction mechanisms must be consistent with:1.Stoichiometry for the overall reaction.

2.The experimentally determined rate law.

Reaction mechanism must be consistent with the stoichiometry of the overall reaction.

• Is the mechanism below consistent with the overall reaction above?

NO2(g) + NO2(g) NO3(g) + NO(g)

NO3(g) + CO(g) NO2(g) + CO2(g)

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

Determining the stoichiometry of a reaction mechanism.

Page 439Page 439

• The reaction mechanism must also support the rate law.


Rate Laws for Multistep Mechanisms

with an initial fast step.• Consider the reaction:

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


Mechanisms with an Initial Fast Step

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

• The experimentally determined rate law is

Rate = k[NO]2[Br2]

• Consider the following mechanism


NO(g) + Br2(g) NOBr2(g)k1

k-1

NOBr2(g) + NO(g) 2NOBr(g)k2

Step 1:

Step 2:

(fast)

(slow)

• The rate law is (based on Step 2):

Rate = k2[NOBr2][NO]

• The rate law should not depend on the concentration of an intermediate (intermediates are usually unstable).

• NOBr2 is an unstable intermediate, so we express the concentration of NOBr2 in terms of NOBr and Br2 Since there is an equilibrium in step 1 we have

]NO][Br[]NOBr[ 21

12

kk


k-1


Step 1:

Step 2:

(fast)

(slow)

• By definition of equilibrium:

• Therefore, the overall rate law becomes

• Note the final rate law is consistent with the experimentally observed rate law.

]NOBr[]NO][Br[ 2121 kk

][BrNO][NO][]NO][Br[Rate 22

1

122

1

12

kk

kkk

k


k-1


Step 1:

Step 2:

(fast)

(slow)

Determining the rate law of a reaction mechanism.

Page 439Page 439

Catalysts

• A catalyst is a substance that increases the rate of a chemical reaction by reducing the activation energy, but which is left unchanged by the reaction.

What is the overall reaction?

O3 O2 + O

O3 + O 2O2

What is the overall reaction?

Identify the intermediates.

Identify the intermediates.

NO is a catalyst

A homogeneous catalyst is of the same phase as the reacting substances. It lowers the activation energy by forming intermediates which allow the reaction to proceed by a different pathway.

Heterogeneous Catalysts

• A heterogeneous catalyst is of a different phase than the reacting substances. It lowers the activation energy by providing a surface on which the reaction can occur.

Inhibitor• An inhibitor decreases the rate of a reaction. It

often does this by rendering a catalyst ineffective.

Catalyst “poisoning” occurs when a catalytic converter is exposed to exhaust containing substances that coat the working surfaces, encapsulating the catalyst so that it cannot contact and treat the exhaust. The most notable contaminant is lead, so vehicles equipped with catalytic converters can only be run on unleaded gasoline.

http://upload.wikimedia.org/wikipedia/commons/b/b2/Pot_catalytique_vue_de_la_structure.jpg

Inhibitor• An inhibitor decreases the rate of a reaction.

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Chapter 14 Chemical Kinetics. Review Section of Chapter 14 Test Net Ionic Equations