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Transcript of Chemical Kinetics Collision Theory: How reactions takes place Reaction Rates: How fast reactions...
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Chemical KineticsCollision Theory:
How reactions takes placeReaction Rates:
How fast reactions occurReaction Mechanisms
Resource: www.mwiseman.com
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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
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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
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The area under the curve is a measure of the total number of particles present.
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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
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Not all collisions leads to a reaction For effective collisions proper orientation ofthe molecules must be possible
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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
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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
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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
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What is this?
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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
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How do we measure rxn rate?
A B How fast product appears
How fast reactant disappears
t
A
t
B
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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?
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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?
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2 kinds of rate laws
Both determined by experimentDifferential Rate Law
How rate depends on [ ]
Integrated Rate Law How rate depends on time
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Differential Rate Law
2 methods Graphical analysis Method of initial rates
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Graphical Analysis
1. Graph [ ] vs. time
2. Take slope at various pts
3. Evaluate rate for various concentrations
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[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
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[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
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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
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[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-]
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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
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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
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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
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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
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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
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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
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Half-life
Def’n: time it takes for concentration to halve
Depends on order of rxnAt t1/2 [A]=[A]0/2
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Half-Life
First order
Second order
Zero Order
kt
693.02/1
02/1 ][
1
Akt
k
At
2
][ 02/1
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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:
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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
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Rules for Reaction Mechanisms
Sum of elementary steps = overall balanced rxn
Mechanism must agree with experimental rate law
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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
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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!
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Give molecularity and rate law:
Unimolecular (first order) rate=k[A]
Bimolecular (second order) rate=k[A][B]
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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
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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