1. CHAPTER 12
AP CHEMISTRY

2. CHEMICAL KINETICS Speed or rate of reactions - Reaction Rate
Change in concentration of a reactant or product per unit of time
aA + bB -----> cC + dD
Rate = conc. of A at t2 - conc. of A at t1
t2 - t1
Δ [A] A = reactant or product
Δt [ ] = mol/L
Δ = change
Page
Slope of tangent line = change in y
Change in x
Rate = -slope of tangent line

3. FORM OF RATE LAW Page 551-552 First order rate = k[A]
k = constant, rate doubles if concentration of A doubles
Rate law
Any chemical reaction must be determined experimentally; it cannot be predicted
R = k[reactant 1]m[reactant 2]n
Page 554 exercise 12.1

4. INTEGRATED RATE LAW First-order rate law Integrated first-order
2N2O5(sol’n) ---> 4NO2 + O2
Rate = - Δ[N2O5] = k[N2O5]
Δ t
Integrated first-order
In[A] = -kt + ln[A]o
Reaction is first order if a plot of ln[A] versus t is a straight line
Page
y = mx + b [A]o = concentration at the start of reaction
Pressure can be used as a concentration
half-life of a first-order
t1/2 = 0.693 k
Does not depend of the CONCENTRATION
Page 560

5. SECOND-ORDER RATE LAWS
Rate = - Δ[A] = k[A]2
Δ t
Integrated second-order
1 = kt + 1
[A] [A]o
half-life of a second-order
t1/2 = 1
k[A]o
Depends on the concentration
Page 562
Zero-order
Rate = k[A]o = k
[A] = -kt + [A]o
t1/2 = [A]o 2k

6. CONTINUE Rate laws with more than one reactant
BrO3-(aq) + 5Br-(aq) + 6H+(aq) -----> 3Br2(l) + 3H2O(l)
Rate = Δ[BrO3-] = k[BrO3-][Br-][H+]2
Δ t
[BrO3-]o = 1.0 X 10-3 M, [Br-]o = 1.0M, [H+]o = 1.0M
Page 567 table 12.6
Reaction mechanism
Series of steps
Has intermediates that are not seen
Page 569
One step will be the rate determining one

7. REACTION MECHANISM Path or sequence of steps Elementary steps
Must always add to give the chemical equation of the overall process
Rate of steps = to rate constant multiplied by the concentration of each reactant
A ---> B + C Rate = k[A]
A + A -----> B + C Rate = k[A]2
Slow steps
The slow step is rate-determined
Overall rate cannot exceed the slowest step
If the step is definitely the slowest, its rate is approximately equal to the overall reaction

8. CONTINUE All intermediates MUST be eliminated
NO(g) + Cl2(g)  NOCl2(g) Fast
NOCl2(g) + NO(g) 2NOCl(g) Slow
2NO(g) + Cl2(g) ----> 2NOCl(g)
k2[NOCl2][NO]
k1[NO][Cl2] = k-1[NOCl2]
[NOCl2] = k1[NO][Cl2]
k-1
Overall = k2[k1[NO][Cl2]][NO]
[ k ]
= k2k1[NO]2 [Cl2]1

9. ACTIVATION ENERGY
THERE MUST BE A CERTAIN AMOUNT OF KINETIC ENERGY FOR A REACTION TO OCCUR
Ea > 0
Depends on nature of reaction
Independent of temperature and concentration

10. ARRHENIUS EQUATION K = Ae-Ea/RT Ea= activation energy
R = J/K·mol
T = absolute temperature
A = frequency of collisions
ln k1 = Ea(1 _ 1 )
k2 R (T2 T1)

11. CATALYSIS Increases the rate of reaction without being consumed
Heterogenous catalysis
In different phase as mixture
Initial step is absorption, outside particles in a solid will have unfilled valence shells which can attach to other particles. These places are called active sites found in Enzymes
Without digestive enzymes it is estimated it would take 50 years to digest a meal
MnO2
Homogeneous catalysis
Present in same phase