1. Kinetics and Thermodynamics
Cartoon courtesy of
NearingZero.net

2. Enthalpy and Entropy
Reactions tend to proceed in the direction that lowers the energy of the system (H, enthalpy).
Reactions tend to proceed in the direction that increases the disorder of the system (S, entropy).
Entropy
Enthalpy

4. Endothermic Reactions

5. Exothermic Reactions

6. A Bomb Calorimeter

7. A Cheaper Calorimeter

8. Specific Heat
The amount of heat required to raise the temperature of one gram of substance by one degree Celsius.

9. Calculations involving Specific HeatOR
cp = Specific Heat
q = Heat lost or gained
T = Temperature change
m = Mass

11. J
J/gx0C
J/g
J/g

13. Table of Specific Heats

14. Molar Heat of Fusion
The energy that must be absorbed in order to convert one gram of solid to liquid at its melting point.
The energy that must be removed in order to convert one gram of liquid to solid at its freezing point.

15. 33400.0J 2g 20.0 g/J 668,000.0 J q = mHf m = mass Hf = heat of fusion
The heat of fusion is the heat needed to melt 1 gram of a solid at constant T. The quantity of heat that is absorbed during fusion or melting is released during solidification or freezing.
q = mHf  
m = mass
Hf = heat of fusion
Example
The heat of fusion of water is 334 J/g. Therefore, to melt 1 g of ice at  0 oC 334 J are required .
Problems  
1. What is the heat needed to melt grams of ice?
2. How many grams of ice can be melted with 668 J?
3. What is the heat of fusion of substance X if J are required to melt 10.0 g of X?
4. What is the heat released when 2.0 kg of water freezes? J 2g
20.0 g/J
668,000.0 J

16. Heat of Vaporization
The energy that must be absorbed in order to convert one gram of liquid to gas at its boiling point.
The energy that must be removed in order to convert one gram of gas to liquid at its point of condensation.

17. The heat of vaporization is the heat needed to vaporize 1 gram of a liquid at constant T. The quantity of heat that is absorbed during vaporization is released during condensation.
Heat: q = mHv  
q = heat
m = mass
Hv = heat of vaporization
Example
The heat of vaporization of water is 2260 J/g. Therefore, to vaporize 1 g of water at 100 oC 2260 J are required .
Problems 
1. What is the heat needed to vaporize grams of water?
2. How many grams of water can be vaporized with 6780 J?
3. What is the heat of vaporization of substance Y if J are required to vaporize g of X?
4. What is the heat released when mg of water condenses?

18. Chemical Kinetics Key Idea: Molecules must collide to react.
The area of chemistry that concerns reaction rates.
Key Idea: Molecules must collide to react.
However, only a small fraction of collisions produces a reaction. Why?

19. Collision Model
Collisions must have enough energy to produce the reaction (must equal or exceed the activation energy).
Orientation of reactants must allow formation of new bonds.

20. Reaction Mechanism The series of steps by which a chemical
reaction occurs.
A chemical equation does not tell us
how reactants become products
It is a summary of the overall process.
Example:
has many steps in the reaction mechanism

21. Rate-Determining Step
In a multi-step reaction, the slowest step is the rate-determining step. It therefore determines the rate of reaction.

22. Catalysis
Catalyst: A substance that speeds up a reaction without being consumed
Enzyme: A large molecule (usually a protein) that catalyzes biological reactions.
Homogeneous catalyst: Present in the same phase as the reacting molecules.
Heterogeneous catalyst: Present in a different phase than the reacting molecules.

23. Factors Affecting Rate
Temperature
Increasing temperature always makes the molecules or atoms move faster thus increasing the opportunity for collisions to occur. Surface Area
Increasing surface area increases the rate of a reaction
ConcentrationIncreasing the concentration of the reactants will increase the frequency of collisions between the two reactants.
Increasing concentration USUALLY increases the rate of a reaction
Presence of Catalysts

24. Catalysts Increase the Number of Effective Collisions

25. Endothermic Reaction with a Catalyst

26. Exothermic Reaction with a Catalyst

27. Chemical Equilibrium Reversible Reactions:
A chemical reaction in which the products
can react to re-form the reactants
Chemical Equilibrium:
When the rate of the forward reaction
equals the rate of the reverse reaction
and the concentration of products and
reactants remains unchanged
2HgO(s)  2Hg(l) + O2(g)
Arrows going both directions (  ) indicates equilibrium in a chemical equation

30. LeChatelier’s Principle
When a system at equilibrium is placed
under stress, the system will undergo a
change in such a way as to relieve that
stress.

31. Le Chatelier Translated:
When you take something away from a system at equilibrium, the system shifts in such a way as to replace what you’ve taken away.
When you add something to a system at equilibrium, the system shifts in such a way as to use up what you’ve added.

32. LeChatelier Example #1
A closed container of ice and water at equilibrium. The temperature is raised.
Ice + Energy  Water
The equilibrium of the system shifts to the _______ to use up the added energy.
right

33. LeChatelier Example #2
A closed container of N2O4 and NO2 at equilibrium. NO2 is added to the container.
N2O4 (g) + Energy  2 NO2 (g)
The equilibrium of the system shifts to the _______ to use up the added NO2.
left

34. LeChatelier Example #3
A closed container of water and its vapor at equilibrium. Vapor is removed from the system.
water + Energy  vapor
The equilibrium of the system shifts to the _______ to replace the vapor.
right

35. LeChatelier Example #4
A closed container of N2O4 and NO2 at equilibrium. The pressure is increased.
N2O4 (g) + Energy  2 NO2 (g)
The equilibrium of the system shifts to the _______ to lower the pressure, because there are fewer moles of gas on that side of the equation.
left

36. Practice EOC Problems

37. Practice EOC Problems

38. Practice EOC Problems

39. Practice EOC Problems

40. Practice EOC Problems

41. Practice EOC Problems

42. Practice EOC Problems