Chemical Kinetics Collision Theory: How reactions takes place Reaction Rates: How fast reactions occur Reaction 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/essentialchemistry/flash/activa2.swf transition state

Not all collisions leads to a reaction For effective collisions proper orientation of the 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

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

Forward vs Reverse Rxn Some rxns are reversible After 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 experiment Differential 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 Graph [ ] vs. time Take slope at various pts Evaluate rate for various concentrations

Graphical Analysis When concentration is halved… [N2O5] (M) Rate (M/s) Rate is halved Order = 1 Rate = k[N2O5]1 [N2O5] (M) Rate (M/s) 1.0 2 0.5 0.25

Graphical Analysis When concentration is doubled… [NO2] (M) Rate (M/s) Rate is quadrupled Order = 2 Rate = k[N2O5]2 [NO2] (M) Rate (M/s) 1.0 2 2.0 8 4.0 32

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 2 0.2 4 6

[NH4] [NO2-] Rate 0.1 2 0.2 4 6 [NH4] [NO2-] Rate 0.1 2 0.2 4 8 When [NO2] doubles, rate doubles, First order with respect to (wrt) NO2 n = 1 When [NO2] doubles, rate doubles, First order with respect to (wrt) NO2 m = 1 Rate = k[NH4+] [NO2-]

Calculate k, using any of the trials, you should get the same value Try this one: [NH4+] [NO2-] Rate (M/s) 0.1 2 0.2 8 Rate = k [NO2-]2 Calculate k, using any of the trials, you should get the same value

Integrated Rate Law Tells how rate changes with time Laws are different depending on order Overall 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: 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: 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: 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 rxn At t1/2 [A]=[A]0/2

Half-Life First order Second order Zero Order

Reaction Mechanism Reactions occur by a series of steps = 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 molecule Bimolecular – 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? Rate law is determined using experiment (method of initial rates, etc.) Chemist uses intuition to come up w/ various mechanisms Narrows down choices using rules for mechanisms No mechanism is ever absolutely proven