kinetics and eq notes

8/2/2019 Kinetics and Eq Notes

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Kinetics and Equilibrium

Table of Contents

I. Collision Theory

II. Five Factors Affecting Reaction Rate

III. Reaction Rate Mechanisms

A. Rate Mechanisms

1. One step or

2. multistep

3. Rate Law

4. Rate determining step

5. Rate Orders

5. Rate Constant determination

B. Homogeneous v. Heterogeneous Rxns

IV. PE Diagrams

A. Exothermic Rxns

B. Endothermic Rxns

V. Spontaneous v. Nonspontaneous

A. Def.'s

B. Entropy

C. Free Energy

D. Gibbs Free Energy Equation

1. Def.

2. temperature a changing variable

3. change of state

VI. Heats of:


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A. Fusion and Vaporization

B. Solution

C. Formation

D Reaction

1. Table I, Thermochemical Data and definitions

2. Calculating from Hf

3. Hess' Law

VII. Equilibrium

A. Physical States

B. Chemical

C. Equilibrium constants

1. Keq, Ksp, Ka, Kb (heterogenous (only [aq], [g] count)

D. LeChatelier's principle

EXTRA:

PowerPoint Slides

Essential Questions:

1.Why are most exothermic reactions like the burning of magnesium spontaneous (that is they continue on

their own once initiated) and endothermic reactions are largely not?

2. Why are some endothermic reactions spontaneous even though they are gaining unwanted energy?

3. Is there any way to change the Heat of Reaction (delta H) in a given reaction?
http://var/www/apps/conversion/current/tmp/scratch18656/Kinetics%20PP%20Notes.ppthttp://var/www/apps/conversion/current/tmp/scratch18656/Kinetics%20PP%20Notes.ppt


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Kinetics and Equilibrium

I.Collision Theory

A. Particles will react if they have an EFFECTIVE COLLISION. This depends

on:

1. PROPER ORIENTATION (right place, right time)


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2. SUFFICIENT KINETIC ENERGY

If both conditions are achieved there will be an effective collision

and the particles will react.

II. Five Factors Affecting the Rate of Reaction


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III. Reaction Rate Mechanisms

A. Rate Mechanisms

Reactions can occur in one step or in multiple steps

1. ONE STEP

2. MULTISTEP

1. Reactions occur in steps. The more steps involved the slower the reaction rate.

2. Consider the following reaction: NO2 + F2 --> 2NO2F


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3. RATE LAW

The rate law is an expression that shows how necessary reactants and

their concentrations affect the overall rate of a reaction. The rate law for

a reaction has the following general expression:

xA + yB zAB

Rate = k [A]x[B]y

PRACTICE: Write the Rate Laws for the following rxns.

1. N2(g) + 3H2(g) -> 2NH3(g)

2. 2H2(g) + O2(g) 2H2O(l)

3. Mg(s) + O2(g)

4. KClO3(s)

5. C3H8(l) + O2(g) +

4. Rate Determining Step

The rate of a chemical reaction is determined by the slowest step.

(weakest link analogy)


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If you are not told that it is a one step reaction or (for multistep rxns)

which is the slowest step then the reaction rate law can not be

determined.

A + B AB slow

AB + B B2 + A fast

2B B2

The Rate Law is determined by the slowest step therefore

Rate = k [A][B]

5. Rate Order

How does a change in concentration of a reactant effect the rate of a

reaction? The rate order of a reactant/reaction is an expression of this

relationship. It can either be determined from a ONE STEP reaction or

EXPERIMENTALLY through data!

r= k[A]a [B]b

Example : 1N2 + 3H2 2NH3

The reaction rate would be 1st order with respect to N2

And 3rd order with respect to H2.


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The rate orders are as follows:

Zero Order change in [ ] has no effect on reaction rate

First Order change in [ ] has a direct effect on reaction rate

Second Order change in [ ] has a squared effect on reaction rate

Third Order change in [ ] has a cubed effect on reaction rate

[A] [B] Rate M/s Order A Order B Order Rxn

(A+B).100 .100 .0050

.200 .100 .0100

.200 .200 .0400

.400 .400

5. Rate Constants (k):

The Rate constant (k) for any reaction is whatever is necessary to make

the reaction rate be in M/s or (Ms-1)

k x [ ] = M/s so ( ) x [ ] = M/s

k

Zero Order Rate = k

First Order Rate = k[M]


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Second Order Rate = k[M]2

Third Order Rate = k[M]3

B. Homogeneous v. Heterogeneous Rxns

Reactants that are in the same phase are considered homogeneous

Ex: N2(g) + 3H2(g) 2 NH3(g)

HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l)

Reactants that are in different phases are considered heterogeneous

Ex: Na(s) + H2O(l) H2(g) + NaCl(aq)

Fe(s) + O2(g) Fe2O3(s)

IV. PE Diagrams

A. Exothermic Reactions A + B --> AB + Energy


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B.

C.


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D. PE Diagrams of a catalyzed reaction

Adding a catalyst will lower the needed activation energy and thus allow the reaction to

happen at a faster rate, and typically at a lower temperature. NOTICE that the activation

energy for the reverse reaction is also lowered by the same amount. This will come into play

when we discuss equilibrium.


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E. Endothermic Reactions

Energy + A + B --> AB

The PE diagrams for these type of reactions will be completely opposite of those for exothermic.

The heat of reaction will be POSITIVE as the system is gaining energy.

CLICK HERE for more help with endothermic PE diagrams.
http://www.kentchemistry.com/links/Kinetics/PETutorial.htmhttp://www.kentchemistry.com/links/Kinetics/PETutorial.htm


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V. Spontaneous v. Nonspontaneous

A. Definition

1. Spontaneous - rxn's favored naturally, once initiated will

continue on it's own.

Examples: burning of paper, evaporation, precip rxns

2. Non-spontaneous - rxn's not favored in nature, will only continue

with outside interference.

Examples: boiling of water, separation of metals from their ores.

B. Entropy-

1. Randomness or disorder

2. An increase in entropy is favored in nature.

3. Examples of increased entropy

a. phase changes


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b. grinding

c. electrolysis

d. temperature

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