1. Chemical kinetics: In what way do chemical reactions occur
Important in medicine, cooking, industry
Reaction mechanism (reaction pathway):
Steps by which atoms rearrange themselves during a chemical reaction
Can have several intermediate steps
Reaction rate:
How fast a reaction takes place
Measured in moles of reactant disappearing or moles of product formed per unit time

2. Collision theory:
In order for a reaction to occur, particles must collide with sufficient energy and the proper orientation
Effective collision: a collision that has enough energy and the proper orientation, thus `resulting in a reaction taking place
Activation energy (Ea): the minimum energy a collision must have to result in a reaction
Rate of reaction depends on the number of effective collisions

4. Factors affecting rate of reaction:
Nature of the reactants:
The more tightly the atoms are bound, the slower the rate
In solution, ionics react faster than covalents
Gases react faster than liquids, which are faster than solids
Simple molecules react faster than more complex ones

5. B) Concentration of reactants:
The greater the concentration, the faster the rate of reaction
The greater the concentration, the closer the particles are to each other, the greater the likelihood of collisions
What about pressure?
The higher the pressure, the closer the particles are to each other, the greater the concentration
The higher the pressure, the faster the reaction

6. D) Temperature of a system:
C) Surface area:
The more surface area, the faster the rate
Increases area of contact between reactants
Ex: Steel bar versus steel wool
Chewing of food
D) Temperature of a system:
The higher the temperature, the faster the rate
Collisions are more likely to have sufficient energy to be effective

7. Catalyst speeds up reaction but is not consumed by the reaction
E) Presence of a catalyst:
Catalyst speeds up reaction but is not consumed by the reaction
Lowers activation energy needed
Provides alternate, lower energy pathway
F) Presence of an inhibitor:
Slows down reaction
Ties up reactants
Increases activation energy
Ex: Preservatives, medicine

8. Energy in reactions (Thermodynamics)
Heat: a form of energy; the ability to do work
Potential energy (PE): stored energy
Activation energy (Ea):
Amount of energy for there to be an effective collision
Amount of energy necessary to form activated complex
Activated complex:
A high energy, intermediate product
Can continue to form final product

9. Heat of reaction (H, or enthalpy) H = heatproducts – heatreactants
The amount of heat lost or gained during a chemical reaction
(+) in endothermic reactions
(-) in exothermic reactions
H tells how endo or exothermic a reaction is

10. Potential energy (PE) diagrams:
Reaction coordinate: How far a long a reaction has progressed
Exothermic reaction:
Gives off energy to surroundings
So, heat of products is less than heat of reactants
H is negative
Heat is a product of the reaction
Heat is written on right side of reaction arrow
Ex: 2H2 + O2  2H2O + heat

11. Endothermic reaction:
Absorbs energy from surroundings
So, products have more energy than reactants
H is positive
Heat is a reactant
Heat is written on the left side of reaction arrow
Ex: 2H2O + heat  2H2 + O2

12. Effect of a catalyst on a chemical reaction:
Lowers activation energy needed
Provides alternate, lower energy pathway
Speeds up the reaction in both directions
Reaction just reaches equilibrium faster
ΔH is unchanged

13. Chemical equilibrium:
Dynamic equilibrium:
The rate of one process is equal to the rate of the reverse process with no net change
Types of equilibrium so far…
Phase equilibrium:
Melting = freezing
or Evaporation = condensation
B) Solution equilibrium:
Dissolving = crystallizing

14. C) Chemical equilibrium: A condition in which…
Rate of forward reaction = rate of reverse reaction
*Must have a closed system
Ex: N2 + 3H2 ⇌ 2NH3 + heat
*The double-ended arrow indicates this is a reversible reaction
At equilibrium, the rate at which N2 and H2 combine to form NH3 is equal to the rate at which NH3 breaks down into N2 and H2
There is no net change in the concentration or amount of N2, H2, or NH3

16. Lechâtelier’s Principle:
If a stress is placed on a system at equilibrium, that equilibrium will shift to absorb the stress until a new equilibrium is reached
Possible stress factors:
1) Increasing or decreasing concentration (Gases and aqueous) ONLY:
Equilibrium will shift away from the side where concentration increases and toward the side where concentration decreases
Ex: N2(g) + 3H2(g) ⇌ 2NH3(g) + heat
Increase [H2], equilibrium shifts to the right
Decrease [NH3], equilibrium shifts to the right

17. 2) Increasing or decreasing pressure or volume (Gases ONLY):
Increasing pressure will shift equilibrium to the side with the fewest moles of gas
Decreasing pressure will shift equilibrium to the side with the more moles of gas
Ex: N2(g) + 3H2(g) ⇌ 2NH3(g) + heat
1 mole moles moles
4 moles vs. 2 moles
So, increasing pressure shifts equilibrium to the right

18. 3) Increasing or decreasing temperature
Treat “heat” as a reactant
Increasing temperature shifts equilibrium away from side with “heat” on it
Ex: N2(g) + 3H2(g) ⇌ 2NH3(g) + heat
Increasing temperature shifts equilibrium to the left

19. OVERALL SUMMARY ON HOW TO PREDICT SHIFTS IN EQUILIBRIUM:
4) Presence of a catalyst:
No change
Rate of reaction increases equally in both directions
Equilibrium is simply reached faster
OVERALL SUMMARY ON HOW TO PREDICT SHIFTS IN EQUILIBRIUM:
Equilibrium shifts away from side one adds to and toward the side one takes away from
Changes in amounts of solids and liquids has no effect

21. Ex: H2(g) + Cl2(g) ⇌ 2HCI(g) + energy
Ex: 2 ZnS(s) + 3 O2(g) ⇌ 2 ZnO(s) SO2(g) ∆H = +65Kcal
Ex: C(s) ­+ H2O(g) ⇌ CO(g) + H2(g)
Ex: 2 SO2(g) + O2(g) ⇌ 2 SO3(g)
Ex: NH4CI(s) + heat ⇌ NH3(g) + HCI(g)

22. Spontaneous versus non-spontaneous reactions:
Two basic natural tendencies:
1) Energy changes (H):
Nature favors lower energy (-H)
2) Entropy changes (S):
Entropy: randomness, disorder, chaos
Nature favors higher entropy (+S)
If a reaction has both +S and -H, it is always spontaneous
If a reaction has both -S and +H, it is never spontaneous

23. Ways that entropy in a system can increase:
Phase change
solid liquid gas
increasing entropy
Ex: 2H2(g) + O2(g)  2H2O(l) indicates lower entropy
B) Decomposition:
Increase in the number of moles of particles
Ex: 2H2O2  2H2O + 3O2 indicates higher entropy

24. C) Dissolving:
Particles have more entropy in aqueous phase
Ex:
D) Number of moles of gas
Ex: 2NH3(g)N2(g) + 3H2(g)
2 moles  4 moles, so an increase in entropy
E) Temperature:
Higher temperature means more randomness of particles, so an increase in entropy