1. Kinetics: study of rate or speed at which Reactions occur.
Unit 8 :Chapter 16 and 17
Kinetics: study of rate or speed at which Reactions occur.
a reaction is the BREAKING and REFORMING of bonds to make enitirely new compounds as products.
Reaction Mechanism= Step by Step needed to make a product (like a recipe)
Just like you bake a cake we must follow directions
CAN’T OMIT ANY STEPS CAN’T CHANGE ORDER OF STEPS CAN’T OMIT any REACTANTS (ingredients)

2. What Determines the Rate of Reaction?
# of steps: more steps slower the reaction
Rate Determining Step: the slowest step of the reaction; most importat factor influencing reaction rate.
Collision Theory: Reactions only take place if there are effective collisions. They need:
1. Proper amount of energy
2. Proper orientation/alignment
*****NEED EFFECTIVE COLLISIONS for reactions to take place.
Activation Energy(Eact) minimum amount of energy needed to start a reaction.

3. Activated Complex: (AC) immediate substance form waiting to see if effective collisions take place. Unstable Ex Reactants AC

4. Factor
How Rate Affected
Why does it increase the rate?
Nature of
Reactants
IONIC:substances react Faster
Covalent: substances react slower
IONIC: Smaller less bonds to break
Covalent: larger more bonds to break (more steps)

5. 2. Concentration
Increase Concentration=
Increase rate
3. Pressure
INCREASE press =
Increase Rate
GASES ONLY!!!
4. Temperature
Increase Temp = Increase rate

6. 5. Surface Area
Increase SA (make pieces smaller)=
Increase rate
6. Catalyst

7. PE of the Products-PE of the Reactants
Heat of Reaction =the amount of heat ENERGY LOST or GAINED
throughout a Reaction enthalpy
PE of the Products-PE of the Reactants
Also recall, there are 2 types of reactions:
1. Reactions that release energy exothermic
is negative (-) A + B  C + ENERGY
Ex Sodium in water  heat (fire) as product
2. Reactions that absorb/ gain energy  endothermic
iis positive (+) A + B + ENERGY  C
Ex baking (need oven to supply heat)

8. How do I interpret Table I: Look at the equation on your paper:
1. If it’s written the same keep the sign you see on Table I
2. If it’s written in reverse you must reverse the sign you see on Table I
whatever you do to a chemical reaction you also must do to the
Ex 1 If you REVERSE a reaction (flip the products and reactants) then you must
Reverse the sign for
Ex2 If you double the equation (or coefficients) then you must double the

9. Ex#1 H2O(l)  H+(aq) + OH-(aq) Is it written the same?
Do you reverse the sign? H=?

15. III. Potential Energy Diagram1
1=PE Reactants = EACT REV RXN
2=EACT FWD RXN =
3= Activated Complex (AC) 6= PEProducts

16. Stays the same Changes with catalyst added
PE Reactants EACT FWD RXN
PEproducts EACT REV RXN AC

17. Equilibrium: when the rate of the FWD RXN=rate REV RXN in a closed system

18. Types of Equilibrium
Phase Equilibrium: occurs during a phase change Ex 2.

19. 3. Haber: Process used by Germans during WW1 to pull nitrogen out of the air to make ammonia

20. Lechatelier’s Principle:
When a stress is applied to a system, the reaction will try to shift in a direction that will relieve that stress
Right=shifts means that the fwd rxn is predominate until eq. is reestablished. Arrow=  more product is made
Left= shift means that the reverse reaction is predominate until eq. is reestablished. Arrow= more reactant is made

21. [ ]: Concentration [ ]
Increase [ ]=shifts to other direction to make less
Decrease [ ]= stays on that side to make more
Ex. N2(g) + 3H2(g) > 2NH3(g)
Increase [N2] shifts? Decrease [H2] shifts?
Increase [NH3] shifts?

22. Increase pressure shifts? Decrease pressure shifts?
Effect Pressure Gases ONLY!!!!!!!!!!
Increase pressure=shifts to side with lesser # of moles
Decrease pressure = shifts to side with greater # of moles
Ex: N2(g) + 3H2(g) > 2NH3(g)
Increase pressure shifts? Decrease pressure shifts?

23. H2(g) + I2(g) <-> 2HI (g)
Increase pressure shifts? Decrease pressure shifts?

24. 4Al(s) + 3O2(g) <-> 2Al2O3 (s)
Increase pressure shifts? Decrease pressure shifts?

25. Effect of temperature=same as [ ]
Increase temperature= shift to other side
Decrease temperature = stay on that side
Ex H2(g) + O2(g) <-> 2H2O(g) kJ
Increase Temperature shifts? Decrease temperature shifts?
2H2O(g) kJ <-> 2H2(g) + O2(g)
Increase temp shifts? Decrease temp shifts?

27. Balance the equation and apply the stresses.
N2(g) + H2(g_  NH3(g) kJ
N2(g) increased____ 7. Pressure Increased___
NH3(g) increased ___ 8. Pressure Decreased___
H2(g_ increased___ Temp Increased___
N2(g) decreased____ 10. Temp Decreased___
NH3(g) decreased ___
H2(g_ decreased___

28. 2HgO(s) + 10kcal 2Hg(l) + O2(g)
Increase Pressure=
Increase temp=
[HgO] increased=
[Hg] Increased=
Temp Decreased=
Presure decreased=
[O2]decreased=
Catalyst=
[HgO] decreased=

29. 2CO(g) + O2(g)  2CO2(g) kJ
Press increased= catalyst=
Temp increased= [CO] decreased-
[CO] increased=
[CO2] increased=
Temp decreased=
Pressure decreased=
[O2] decreased=