1. Unit 10: Kinetics, and Equilibrium
Clock Reaction…Make observations!

2. I. Kinetics Kinetics: Collision Theory:
- Study of how chemical reactions occur and reaction rates
- It is collisions between reactants that allow a reaction to occur
- In order for a reaction to occur, there has to be the correct amount of energy and orientation of the collisions
How to Speed Up Chemical Reactions (and get a date) Video Clip

3. Ii. Six factors affecting rate of a reaction
1) Nature of the Reactants:
2) Concentration:
3) Surface Area:
- Ionic reactants, especially aqueous, react very quickly (easier to break apart ions)
Covalent reactants take a long time
(hard to break apart the atoms that are sharing e-)
- If the concentration of the reactants increases, then the reaction will occur faster
- If the surface area of the reactants increases, then the reaction will occur faster

4. CHECK FOR UNDERSTANDING
1) Why would putting batteries in the refrigerator make them last longer?
Decrease in temperature, decrease in the amount of collisions.
2) At STP, which 4.0 g zinc sample will react fastest with dilute hydrochloric acid?
Lump C) bar
Powdered D) sheet metal
B – powdered – more surface area
3) Which substancee would react fastest with hydrochloric acid?
C6H12O6 C) MgCl2
H2O D) SF2
C – ionic compound

5. Ii. Six factors affecting rate of a reaction
4) Pressure (ONLY EFFECTS FOR GASEOUS REACTANTS)
5) Presence of Catalyst:
6) Temperature:
If the pressure of the reactants INCREASES, then the reaction will occur FASTER
A substance that increases the rate of a reaction by creating a different pathway.
It is NOT used in the reaction equation, so it is NOT a reactant or product. It appears above the reaction
- If the temperature of the reactants INCREASES, then the reaction will occur FASTER

6. IIi. Rates and equilibrium
Example: NaCl (aq) Na+ (aq) + Cl – (aq)
Not all reactions can exist at equilibrium ( _________________ ).  
Some reactions ___________________________ which means all the reactants get _____________ and ___________________ go backwards
Rates can refer to how fast the forward reaction occurs or how fast the reverse reaction occurs.
Forward = backward
go to completion
used up
does not

7. IIi. Rates and equilibrium
A reaction will ____________________ if it is not a closed system. Another reason it may not have a reverse rxn is if it forms: ______________________________
Example: Pb(NO3)2 (aq) + 2 KI (aq)  PbI2 (s) + 2 K(NO3) (aq)
go to completion
A precipitate (insoluble solid)

8. IV. Energy changes in chemical reactions
absorbed
Endothermic Reactions: heat is _______________ from the surroundings and the temperature of the surroundings ___________
General Equation:
Using Table _______:
Example: N2(g) + O2 (g)  2 NO (g) ∆ H (kJ) = ________
Can be rewritten as:
decreases
A + B + Heat/energy  C + D I kJ
N2(g) + O2(g) kJ  2 NO (g)
Excess stored energy in the products makes them
UNSTABLE and more reactive

9. IV. Energy changes in chemical reactions
released
Exothermic Reactions: heat is _______________ INTO the surroundings and the temperature of the surroundings ___________
General Equation:
Using Table _______:
Example: 2 C (s) + 3H2 (g)  C2H6 (g) ∆ H (kJ) = ________
Can be rewritten as:
increases
A + B  C + D + Heat/energy I kJ
2 C (s) + 3H2 (g)  C2H6 (g) kJ
The products have less stored energy making them more STABLE

10. V. Potential Energy Diagrams (VIDEO) (ANIMATION)
Endothermic Reactions: Energy (heat) is _________ by the reactants, so the net amount of potential energy ________________
absorbed
increases
A + B +50 kJ C
Reactants PE (60 kJ)
A) Potential Energy (PE) of the Reactants: amount of energy the reactants have

11. A + B +50 kJ C
Activation Energy (70 kJ)
Reactants PE (60 kJ)
B) Activation energy: amount of energy it takes to BREAK EXISTING BONDS and get the reaction STARTED

12. V. Potential Energy Diagrams
A + B +50 kJ C
Activation Energy(70 kJ)
PE of AC (130 kJ)
Reactants PE (60 kJ)
C) PE of Activation Complex: amount of energy activated complex has

13. V. Potential Energy Diagrams
A + B +50 kJ C
Activation Energy (70 kJ)
PE of AC (130 kJ)
PE of Products (110 kJ)
Reactants PE (60 kJ)
D) PE of Products: amount of energy the products have

14. V. Potential Energy Diagrams
A + B +50 kJ C
Activation Energy (70 kJ)
+ ∆ H for ENDO (+50 kJ)
PE of AC (130 kJ)
Products PE (110 kJ)
Reactants PE (60kJ)
E) Heat of Reaction (∆H) : the NET CHANGE IN ENERGY for the reaction (Difference between Products and Reactants)

15. V. Potential Energy Diagrams
A + B +50 kJ C
Reverse Activation Energy (20 kJ)
Activation Energy (70kJ)
+ ∆ H for ENDO (+50 kJ)
PE of AC (130 kJ)
Products PE (110kJ)
Reactants PE (60kJ)
F) Activation Energy of Reverse Reaction: amount of energy it takes to get THE REVERSE REACTION GOING

16. V. Potential Energy Diagrams
A + B +50 kJ C
Reverse Activation Energy (20 kJ)
Activation Energy (70kJ)
+ ∆ H for ENDO (+50 kJ)
- ∆ H for EXO (-50 kJ)
PE of AC (130 kJ)
Products PE (110kJ)
Reactants PE (60kJ)
G) Heat of Reaction (∆H) of Reverse Reaction : the NET CHANGE IN ENERGY for the REVERSE reaction

17. V. Potential Energy Diagrams (VIDEO) (ANIMATION)
RELEASED
Exothermic Reactions: Energy (heat) is _________ by the reactants, so the net amount of potential energy ________________
DECREASES
X + Y Z + 40 kJ
Reverse Activation Energy (120 kJ)
Activation Energy (80kJ)
- ∆ H for EXO (- 40 kJ)
+ ∆ H for ENDO (+40 kJ)
PE of AC (150kJ)
Products PE (30 kJ)
Reactants PE (70kJ)

18. V. Potential Energy Diagrams
Effects of a Catalyst with Potential Energy Diagrams:
Catalysts often _______________a chemical reaction by providing a ___________ with a ____________________. However, the heat of reaction (∆ H) remains_________________
SPEED UP
NEW PATHWAY
LOWER ACTIVATION ENERGY
THE SAME
AE without
AE with
WITH CATALYST

19. Part C: Equilibrium I. types of equilibrium
- Occurs when rate of forward reaction equals the rate of reverse reaction
- Does NOT mean there is the same amount of reactants and products. However, the CONCENTRATION of them REMAINS CONSTANT
“CON CON and REqual”
(CONcentration CONstant and REaction Equal)

20. Part C: Equilibrium I. types of equilibrium
Phase Equilibrium:
Example: H2O (s)  H2O (l) or H2O (l)  H2O (g)
- Occurs when rate of the forward phase equals the rate of the reverse phase

21. Part C: Equilibrium I. types of equilibrium
B. Solution Equilibrium:
Example: NaCl (s)  Na + (aq) + Cl – (aq) (animation)
When sodium chloride is first placed in the water, the salt dissolves. As the concentration of dissolved ions increases, some of those dissolved ions will rejoin each other and form a precipitate, which is almost immediately re-dissolved. Eventually, all of the water molecules will be engaged in holding ions apart and no more salt can go into solution until some ions come out of solution as precipitate. The rate of dissolving equals the rate of precipitating. This is called a saturated solution. If you place a salt cube in a beaker of saturated sodium chloride solution, the cube will change shape over time as ions precipitate onto the cube from solution as salt from the cube dissolves. This demonstrates very well the dynamicity of systems at equilibrium.
- Formed when a saturated solution has the rate of dissolving equal to the rate of precipitating

22. Part C: Equilibrium I. types of equilibrium
C. Chemical Equilibrium:
 Example: N2 (g) + 3 H2 (g)  2 NH3 (g) + heat
- Reached when rate of the forward reaction is equal to the rate of the reverse reaction and the concentration of each substance remains constant
- Forward Reaction: N2 + 3H2  2NH3 + heat
- Reverse Reaction: 2NH3 + heat  N2 + 3H2

23. Part C: Equilibrium II. Le Chatelier’s Principle
Stress – any change in _________________ , _____________________, or ________________ on an equilibrium system is called a stress.
LeChatelier’s Principle – if stress is applied to a equilibrium system, the system will _______________________
A. Stress #1: Concentration Changes:
concentration
temperature
pressure
Shift to relieve the stress
- Reaction will always shift away from something that is added and towards something that is taken away (Concentration [ ] = M)

24. Part C: Equilibrium II. Le Chatelier’s Principle
Example:
4 NH3 (g) + 5 O2 (g)  4 NO (g) + 6 H2O (g) + Heat
Increase [NH3] results in: Shifts to the __________________,
____ [O2], _____ [NO], _____ [H2O] _____ heat
Less [NH3] results in: Shifts to the __________,
____ [O2], _____ [NO], _____ [H2O,]_____ heat +
--- + + +
RIGHT ↓ ↑ ↑ ↑
--- +
---
---
---
LEFT ↑ ↓ ↓ ↓

25. Part C: Equilibrium II. Le Chatelier’s Principle
B. Stress #2: Temperature Changes:
Adding Heat (warming): When heat is added, the _______________ reaction is always ______
Reaction shifts ____________________________________
Example:
4 NH3 (g) + 5 O2 (g)  4 NO (g) + 6 H2O (g) + Heat
Increase Heat results in: Shifts to the __________,
_____ [NH3], _____ [O2], ___ [NO], _____ [H2O]
ENDOTHERMIC
FAVORED
AWAY FROM WHAT IS ADDED
--- + + +
---
LEFT ↑ ↑ ↓ ↓

26. Part C: Equilibrium II. Le Chatelier’s Principle
B. Stress #2: Temperature Changes:
Removing Heat (cooling): When heat is removed, the ____________ reaction is always ______
Reaction shifts ____________________________________
Example:
4 NH3 (g) + 5 O2 (g)  4 NO (g) + 6 H2O (g) + Heat
Decrease Heat results in: Shifts to the __________,
_____ [NH3] _____ [O2], _____ [NO], ___ [H2O]
EXOTHERMIC
FAVORED
TOWARD WHERE HEAT IS REMOVED + +
---
---
---
RIGHT ↓ ↓ ↑ ↑

27. Part C: Equilibrium II. Le Chatelier’s Principle
C. Stress #3: Pressure Changes:
Increasing Pressure (decreasing volume):
Reaction shifts ____________________________________
Example: N2 (g) + 3 H2 (g)  2 NH3 (g)
Increase Pressure results in: Shifts to the __________,
_____ [N2], _____ [H2], _____ [NH3]
TO THE SIDE WITH LESS GAS MOLES
4 total
2 total
RIGHT ↓ ↓ ↑

28. Part C: Equilibrium II. Le Chatelier’s Principle
C. Stress #3: Pressure Changes:
DECREASING Pressure (INCREASING volume):
Reaction shifts ____________________________________
Example: N2 (g) + 3 H2 (g)  2 NH3 (g)
DECREASE Pressure results in: Shifts to the __________,
_____ [N2], __ [H2,], _____ [NH3]
Important Note: If the same number of gaseous reactant and products molecules, pressure changes have _________________ on system
TO THE SIDE WITH MORE GAS MOLES
4 total
2 total
LEFT ↓ ↑ ↑
NO EFFECT

29. Part C: Equilibrium II. Le Chatelier’s Principle
D. Effect of a Catalyst:
Changes rate of _______________ the forward and reverse reaction
BOTH
EQUALLY
Does NOT change any equilibrium concentrations

30. Part C: Equilibrium II. Le Chatelier’s Principle
1. H2(g) + I2 (g) + 53 kJ  HI (g)
What happens if:
A) Increase Temperature: Shifts to the _______, ______ HI, _____ H2, _____ I2
B) Increase [H2 (g)]: Shifts to the _______, ______ HI, _____ I2
C) Increase [HI (g)]: Shifts to the _______, _____ H2, _____ I2
D) Decrease Pressure: +
right
---
----
right +
---
left + +
No change. Equal number of moles of gas on both sides

31. Part C: Equilibrium II. Le Chatelier’s Principle
PCl5 (g) + heat  PCl3 (g) + Cl2(g)
 What happens if:
A) Adding Cl2: Shifts to the _______, ______ PCl5, _____ PCl3
B) Increasing Pressure: Shifts to the ____, ____ PCl5, _____ PCl3, _____ Cl2
C) Lowering Temperature: Shifts to the ____, ___ PCl5, _____ PCl3, ____ Cl2
D) Removing PCl3: Shifts to the _______, ______ PCl5, _____ Cl2 +
---
left
left +
---
----
left +
----
----
left
--- +

32. Part C: Equilibrium VI. Entropy
state of greater randomness or disorder
Nature tends to want to change to _________________________________
Entropy:
( _______________ disorder = ______________ entropy)
Systems in nature tend to undergo changes toward _____________________
and __________________________________
Measure of disorder or randomness of a system
MORE
MORE
Lower energy (enthalpy)
Higher entropy (disorder)

33. Part C: Equilibrium VI. Entropy
Physical Examples of Changes in Entropy:
Entropy INCREASES and is favored
Entropy DECREASES and is not favored

34. Part C: Equilibrium VI. Entropy
Chemical Examples of Changes in Entropy: 1
EXOTHERMIC (less energy) favored and entropy INCREASES 1
ENDOTHERMIC (more energy) not favored and entropy DECREASES

35. Check for understanding
Complete Chart A on pg 22 in Work Packet #2
Check for understanding
Increases
Favored
decreases
unfavored
Increases
Favored
decreases
unfavored
decreases
unfavored
decreases
unfavored