1. Kinetics
By: Anderson Clark

2. Kinetics
Kinetics is defined as the study of the speeds or rates of chemical reactions
Important to industry
Expiration dates
Pharmaceuticals
Large scale production

3. Factors that affect rate of reaction
Chemical properties of reactants
EN, EA, IE, size, red. or ox.
Ability to collide SA
Concentration of measured species
Temperature
Particles collide more and move with greater E, making them more likely to collide with enough E to meet Ea
Catalyst/Inhibitors
NOT consumed in rxn.
Lower Ea

4. Rates With respect to one component
Rate-change in quantity per unit time
M/s, m/s
Always positive (+)
Once one rate is found, it can be used to find others

5. Formulas Rate of X = Δ[X]/Δt
Rate of Rxn. = Δ[X]/Δt x 1/coefficient of X

6. More stuff Rates change as reactants are used up
Instantaneous rate – rate at which a reactant is being used up at any given moment
Steep slope = high rate
Y2-y1/x2-x1 a time 0
Use coefficients

7. Rate Laws Rate = k [A]a [B]b
Exponents are experimentally determined, NOT FROM BCE
k is experimental constant, varies with T
units are L/mol sec or L mol-1 sec-1
Units for Rate are mol/ L sec or mol L-1 sec-1
Exponents are “order of rxn”
Usually 0, 1st, 2nd
Can be fractions or negative – controls rate inversely

8. Overall Order of Rxn
Add exponents of all reactants

9. Table logic Obtaining rate laws from data Look for patterns
Compare rates with only 1 variable

10. Integrated Rate Laws Shows how [conc.] changes over TIME
Graphs will all have time as x-axis
Formula depends on order of reaction

11. First Order DIfferential Integrated Half-life
Rate = k [A]a = -Δ[A] / Δt
Integrated
ln[A]o-ln[A]f = kt
Half-life
t1/2 = ln 2/ t = .693 / k
Not affected by [conc.] of reactants

13. Second Order Differential Integrated Half-Life Rate = k [B]2
1/ [B]f – 1/ [B]o = kt
Half-Life
t1/2 = 1/k [B]o
Depends on [conc.] of reactants

15. Zero Order
Means that reaction rates are independent of the [conc.] of any reactant
Usually involves a small amount of catalyst
Ex. Blood alcohol through liver
Formula
Rate = k
Units are mol/ L sec
Depends heavily on catalyst
SA, amount, quality

16. More Zero Order Differential Integrated Half-Life Rate = k
[C]f = -kt + [C]o
Half-Life
t1/2 = (1/k) [C]o/2

19. !Cat Break!

21. Collision Theory
States that the rate of rxn. Is proportional to the number of EFFECTIVE COLLISIONS per second among reactants
An effective collision is one that yields product molecules
Factors that affect collision rate
Temperature
Concentration
Molecular orientation
Molecular kinetic energy (must meet Ea)
KE changes to PE by forming bonds

22. Transition State Theory
Total KE decreases as it changes to PE, then back to KE
PE diagram = reaction coordination diagram
Path that the reaction follows
Smaller Ea, faster reaction
k = A e-Ea/RT (Arrhenius Equation!!!!)
Also written as Ln k = -Ea/RT + ln A
T = temp in Kelvin, A = proportionality constant, R J/mol K
Rule of thumb: at room temp, rxn. Rates double for every 10 degree C increase

24. Reaction Mechanisms BCE shows total net change, but not all steps
Reaction mechanism- series of individual steps that add up to the overall observed reaction.
Elementary Process (EP) – a reaction involving collisions between molecules
Rate laws for EP can be written based on the coefficients in the BCE, will add up to overall rate law for BCE
Exponents in EP are = to the coefficients of the reactants in chemical equation for that EP

25. Rate Laws and Rate Determining Steps
Make intermediate chemicals
Add the EP BCE’s for overall BCE
One step is usually slower – called the rate determining step
Rate Law for RDS is directly related to overall rate law
Overall order should match the RDS order (from coefficients)
Rate laws cannot be proven correct/true
We can only see beginning and end, not middle steps

26. Etc Etc
Usually only see the collision of two bodies, not three or more
3 way collision is unlikely
When there are 3+ sequential steps, the first can be assumed to be establishing equilibrium, and therefore unstable intermediates

27. Catalysts Homogeneous (MnO2)
Heterogeneous (solid, promotes reaction only on surface)

28. Almost the End R is always 8.314 J/ mol K, never atm
Rates are dependent on temperature

29. Practice

30. Practice

31. Practice

32. Practice

33. Links http://www.youtube.com/watch?v=Tv6_IsdnaGg