Chemical KineticsCollision Theory:

How reactions takes placeReaction Rates:

How fast reactions occurReaction Mechanisms

Resource: www.mwiseman.com

Why are kinetics important?

In order to control processes. speed up useful reactions that occur too slowly slow down reactions that are harmful

Example: Catalysts are used in our cars to

rapidly convert toxic substances into safer substances

Refrigerators are used to slow the process of spoiling in food

Collision Theory

How do reactions occur at the molecular level? Molecules collide with each other Form activated complex

http://www.chem.iastate.edu/group/Greenbowe/sections/projectfolder/animations/NO+O3singlerxn.html

collisions http://www.mhhe.com/physsci/chemistry/

essentialchemistry/flash/collis11.swf

correct and incorrect collisions

The area under the curve is a measure of the total number of particles present.

Svante Arrhenius Did some fancy math to figure out that

number of collisions alone don’t account for reaction rates

He found that reactants also require:Activation energy (Ea - energy to break bonds) Right orientation http://www.mhhe.com/physsci/chemistry/essentialch

emistry/flash/activa2.swftransition state

Not all collisions leads to a reaction For effective collisions proper orientation ofthe molecules must be possible

What affects reaction rate?

Temperature http://www.sciencepages.co.uk/keystage4/GCSEChemistry/rate5

concentration and temperature

Increased number of collisions More molecules have enough activation energy Remember Maxwell-Boltzmann distribution

Increased temperature, distribution flattens out More molecules

have Ea

What affects reaction rate?

Higher concentration Number of collisions increased

http://www.chem.iastate.edu/group/Greenbowe/sections/projectfolder/animations/O2+NO2%20kinetics8.html

concentration

Increased surface area Number of collisions increased

What affects reaction rate?

Catalysts Def’n: substance that speeds up a rxn w/o being used

up itself Number of collisions with Ea increase

Ea lowers Catalysts hold molecules in right orientation

• Homogeneous catalyst (same phase of matter) Demo: Catalysis by Co2+

• Heterogeneous catalyst (different phase)

http://www.chem.iastate.edu/group/Greenbowe/sections/projectfolder/animations/Catalyst2NOO2N28.html

catalyst

What is this?

How do we measure rxn rates?

Rates must be measured by experiment Indicators that a reaction is happening

Color change Gas formation Precipitate formation Heat and light

Many ways to measure the rate Volume / time Concentration / time Mass / time Pressure / time

How do we measure rxn rate?

A B How fast product appears

How fast reactant disappears

t

A

t

B

Forward vs Reverse Rxn

Some rxns are reversibleAfter a sufficient amount of product is

made, the products begin to collide and form the reactants

We will deal only w/ rxns for which reverse rxn is insignificant

2 N2O5(aq) 4 NO2(aq) + O2 (g)Why is reverse rxn not important here?

Rate Law

Math equation that tells how reaction rate depends on concentration of reactants and products

Rates = k[A]n

K = rate constant / proportionality constant n = order of reaction

Tells how reaction depends on concentration• Does rate double when concentration doubles?

• Does rate quadruple when concentration doubles?

2 kinds of rate laws

Both determined by experimentDifferential Rate Law

How rate depends on [ ]

Integrated Rate Law How rate depends on time

Differential Rate Law

2 methods Graphical analysis Method of initial rates

Graphical Analysis

1. Graph [ ] vs. time

2. Take slope at various pts

3. Evaluate rate for various concentrations

[N2O5]

(M)

Rate

(M/s)

1.0 2

0.5 1.0

0.25 0.5

Graphical Analysis

When concentration is halved… Rate is halved Order = 1 Rate = k[N2O5]1

[NO2]

(M)

Rate

(M/s)

1.0 2

2.0 8

4.0 32

Graphical Analysis

When concentration is doubled… Rate is quadrupled Order = 2 Rate = k[N2O5]2

Method of Initial Rates

Initial rate calculated right after rxn begins for various initial concentrations

NH4+(aq) + NO2

-(aq) N2(g) + 2H2O(l)

Rate = k [NH4+]n[NO2

-]m

[NH4+] [NO2

-] Rate (M/s)

0.1 0.1 2

0.1 0.2 4

0.2 0.2 6

[NH4] [NO2-] Rate

0.1 0.1 2

0.1 0.2 4

0.2 0.2 8

[NH4] [NO2-] Rate

0.1 0.1 2

0.1 0.2 4

0.2 0.2 6When [NO2] doubles, rate doubles,

First order with respect to (wrt) NO2

m = 1

When [NO2] doubles, rate doubles,

First order with respect to (wrt) NO2

n = 1

Rate = k[NH4+] [NO2-]

Try this one:

Rate = k [NO2-]2

[NH4+] [NO2

-] Rate (M/s)

0.1 0.1 2

0.1 0.2 8

0.2 0.2 8

Calculate k, using any of the trials, you should get the same value

Integrated Rate Law

Tells how rate changes with timeLaws are different depending on orderOverall reaction order is sum of exponents

Rate = k zero order Rate = k[A] first order Rate = k[A]2 second order Rate= k[A][B] second order

First order integrated rate law

Rearrange and use some calculus to get:

][][

Akt

A

0]ln[]ln[ AktA This is y = mx + b form

A plot of ln[A] vs time will give a straight line

If k and [A]0 (initial concentration) known, then you know the concentration at any time

Second order integrated rate law

Rearrange and use some calculus to get:

2][][

Akt

A

0][

1

][

1

Akt

A

This is y = mx + b form A plot of 1/[A] vs time will give a straight line

If k and [A]0 (initial concentration) known, then you can now the concentration at any time

Zero order integrated rate law

Rearrange and use some calculus to get:

kt

A

][

0][][ AktA This is y = mx + b form

A plot of [A] vs time will give a straight line

If k and [A]0 (initial concentration) known, then you can now the concentration at any time

Graphs give order of rxn

Use graphs to determine order If [A] vs time = zero order If ln [A] vs time = first order If 1/ [A] vs time = second order

Half-life

Def’n: time it takes for concentration to halve

Depends on order of rxnAt t1/2 [A]=[A]0/2

Half-Life

First order

Second order

Zero Order

kt

693.02/1

02/1 ][

1

Akt

k

At

2

][ 02/1

Reaction Mechanism

Reactions occur by a series of steps =

Reaction mechanism Example:

Overall reaction: NO2 + CO NO + CO2

occurs by following steps Step 1:

Step 2:

Intermediates

Two molecules of NO2 collide

Oxygen is transferred, making NO3, the intermediate Intermediates are temporarily formed during a

reaction They are neither a reactant nor a product & Get used up in reaction

Rules for Reaction Mechanisms

Sum of elementary steps = overall balanced rxn

Mechanism must agree with experimental rate law

Elementary Step

Steps in reaction from which a rate law for step can be directly written

2 molecules of NO2 need to collide, therefore…

Rate = k [NO2]2

Molecularity

Rate law written based on molecularity Number of things that have to collide

Unimolecular – rxn depends on 1 moleculeBimolecular – rxn depends on 2 molecules

Termolecular – rxn depends on 3 molecules • Very rare!

Give molecularity and rate law:

Unimolecular (first order) rate=k[A]

Bimolecular (second order) rate=k[A][B]

Rate Determining Step

The slowest step in mechanism determines overall rate

Rate cannot be faster than slowest step Demo: Filling bottle with funnel

Overall rate law can be written from molecularity of slowest step

How are mechanisms determined?

1. Rate law is determined using experiment (method of initial rates, etc.)

2. Chemist uses intuition to come up w/ various mechanisms

3. Narrows down choices using rules for mechanisms

No mechanism is ever absolutely proven