1. Unit 10: Kinetics, Thermodynamics, & Equilibrium
Essential Questions:
How do chemical reactions happen in the first place?
How fast do they go?
Unit Exam is Friday May 8th

2. Kinetics = study of the rate or speed at which reactions occur
A REACTION is the breaking and reforming of bonds to make entirely new compounds as products

3. Just like when we bake a cake we must follow directions
Reaction Mechanism = step by step process needed to make a product; how you get from “a” to “b” (like a recipe)
REACTANTS  PRODUCTS
Just like when we bake a cake we must follow directions
CAN’T OMIT any STEPS!
CAN’T CHANGE THE ORDER of the steps!
CAN’T OMIT any REACTANTS (ingredients)

4. Styrofoam decomposing Weathering of rocks Bleach removing color
Determine whether each of the following chemical reactions is an example of a slow or fast reaction. Explain why knowing this relative rate of rxn is significant.
Rusting
Alka seltzer in water
Styrofoam decomposing
Weathering of rocks
Bleach removing color

5. What determines the rate of a reaction?
1. Number of Steps = more steps can mean a slower reaction
2. Rate Determining Step = the SLOWEST step of the reaction; most important factor influencing reaction rate

6. Collision Theory: In order for a reaction to occur, reactant particles must collide and have the following when doing so:
1. Proper amount of energy
2. Proper alignment/direction/orientation
*Only when particles collide with these two conditions are met will there be an EFFECTIVE COLLISION, resulting in a reaction.

7. THE 6 factors Affecting Rate of Reaction:
How Rate is Affected
Ionic substances react faster
Covalent substances react slower
Factor 1.
Nature of Reactants
Why does it increase the rate
Ionic = small (Less bonds to break; less steps)
NaCl(aq) + AgNO3(aq) → NaNO(aq) + AgCl(s)
Na+(aq) + Cl-(aq) + Ag+(aq) + NO3-(aq) →
Na+(aq) + NO3-(aq) AgCl(s) (1 step)
Covalent = larger (More bonds to break; more steps)
CH4(g) + O2(g) → CO2(g) + 2H2O(l)
(break 4 C-H bonds, 1 O-O bond,
form 2 C-O bonds, and 4 O-H bonds)

8. INCREASE concentration, increase rxn rate
The MORE PARTICLES in a given space, the LESS SPACE b/w particles  MORE COLLISIONS

9. INCREASE pressure, INCREASES rxn rate (affects GASES ONLY!)
Increasing pressure DECREASES VOLUME which DECREASES SPACE b/w particles  MORE COLLISIONS 4.
Temperature
INCREASE temperature, increase rxn rate
Greater SPEED  MORE total COLLISIONS
Greater AVERAGE Kinetic Energy (KE)  collisions take place with MORE energy

10. 5. Surface Area
INCREASE the surface area (by making PIECES SMALLER) INCREASES the rxn rate
Increasing surface area EXPOSES MORE REACTANT PARTICLES to possible collisions

11. Genie in a Bottle: https://www.youtube.com/watch?v=5q5bzHckSIM
6. Catalyst
SPEEDS UP THE RXN WITHOUT CHANGING THE NATURE OF THE REACTANTS/PRODUCTS
Provides a SHORTCUT or ALTERNATIVE PATHWAY for the mechanism
Lowers the ACTIVATION ENERGY for the reaction
Genie in a Bottle:
MnO2
H2O  O2 + H2O

12. Elephant Toothpaste Demo

13. Unit 10: Kinetics, Thermodynamics, & Equilibrium
Essential Question:
Why do you have to strike a match just right to get it to light?
Unit Exam is Friday May 8th

14. Potential Energy Diagrams
Recall, we have talked about chemical bonds having stored energy (AKA potential energy). For that reason, chemists use diagrams called Potential Energy Diagrams to illustrate the potential (or stored) energy changes that occur during specific chemical reactions.
Recall: A reaction is the breaking and reforming of bonds
BREAK BONDS  FORM BONDS
A B  C D

15. Heat of Reaction (ΔH) = the amount of HEAT ENERGY LOST or GAINED throughout a REACTION (ΔH = entHalpy)
PE OF THE PRODUCTS - PE OF THE REACTANTS
ΔH = HP – HR
Also recall, there are two (2) types of reactions:
1. Reactions that release energy  EXOTHERMIC
ΔH = negative value (-)
energy is released (on right)
A B  C D ENERGY
Example: Sodium in water – lots of heat (and fire!) produced as product; heat felt on a test tube during a reaction

16. 2. Reactions that absorb/gain energy ENDOTHERMIC
ΔH = positive value (+)
energy absorbed (on left)
A + B + ENERGY  C + D
Example: chemical cold packs (feels cold because it’s absorbing heat energy)

17. Table I (of the Reference Tables) tells us if particular reactions are exothermic or endothermic based on the sign of the ΔH value.
1. N2(g) + 2O2(g) 2NO2(g)
+66.4 kJ
Endothermic
2. N2(g) + 3H2(g)  2NH3(g)
-91.8 kJ
Exothermic
3. 2NH3(g) N2(g) + 3H2(g)
+91.8 kJ

18. Forward Reaction = reading LEFT TO RIGHT in a reaction; reaction moves toward the right
A + B  C + D
Reverse Reaction = reading RIGHT TO LEFT in a reaction; reaction moves toward the left
A + B  C + D

19. Activation Energy = amount of ENERGY needed to get a reaction STARTED or form the ACTIVATED COMPLEX of a reaction (you must get over the “bump” in order for a reaction to be successful)
How exactly does a catalyst shorten the reaction time needed for a reaction to complete?
The ACTIVATED COMPLEX is lowered which also lowers the ACTIVATION ENERGY needed for the reaction to get started (gets rxn started easier and makes the rxn complete itself faster)

20. ENDOTHERMIC Potential Energy Diagrams  Positive ΔH
Product side (to the Right) is always HIGHER than the reactant side (to the Left) meaning that energy is absorbed.
Label the following:
A = PE of Reactants
B = PE of Products
C = PE of Activated Complex
D = Activation E of Forward rxn
E = Activation E of Reverse rxn
F = Heat of Reaction ΔH
Activated Complex = Highest energy point of the reaction; this is where full rearrangement of the reactants occurs.

21. EXOTHERMIC Potential Energy Diagrams  Negative ΔH
*Product side (to the R) is always LOWER than the reactant side (to the L) meaning that energy is released.
Label the following:
A = PE of Reactants
B = PE of Products
C = PE of Activated Complex
D = Activation E of Forward rxn
E = Activation E of Reverse rxn
F = Heat of Reaction (ΔH = Hp – Hr)
Question: If a catalyst were added to the above diagram, which letter quantities would change within the diagram?
Question: How does the addition of a catalyst change the heat of reaction (ΔH)?
Question: What are the benefits to adding a catalyst?

22. Unit 10: Kinetics, Thermodynamics, & Equilibrium
Essential Question:
What is equilibrium?
How can we use it to our advantage?
Unit Exam is Friday May 8th

23. EquilibRium A
T E S
Some physical and chemical reactions are capable of reaching equilibrium.
Equilibrium occurs WHEN THE RATE OF THE FORWARD REACTION EQUALS THE RATE OF THE REVERSE REACTION in a closed system.
When equilibrium is reached, IT DOES NOT MEAN that the reactants and products are of equal QUANTITIES. So…

24. EquilibRium A
T E S
Equilibrium is represented by DOUBLE ARROWS instead of a single arrow. This allows us to illustrate that the reactions are proceeding in both directions (forward and reverse).
Equilibrium is DYNAMIC which means that it is constantly CHANGING or FLUCTUATING
Equilibrium means that reactant and product CONCENTRATIONS are CONSTANT.
*Equilibrium does NOT mean that reactant and product concentrations are equal.
*MEMORIZATION TRICK: “Con Con Requal” - Concentrations are constant, Rates are Equal

25. Equilibrium Video
Chemical Equilibrium Lesson:
If Molecules were people:

26. Equilibrium
Define equilibrium in terms of reactant and product concentrations:
Define equilibrium in terms of forward and reverse reaction rates:
Reactant and product concentrations are constant.
Forward and reverse reaction rates are equal.

27. TYPES OF EQUILIBRIUM (all occur in Closed Systems)
TYPES OF EQUILIBRIUM (all occur in Closed Systems) *It’s all about the equal rates!
1. Physical Equilibrium : Equilibrium that involves physical changes
a) Phase Equilibrium – occurs during a phase change
(s) (l) Rate of Melting = Rate of Freezing
(sealed 0oC)
(l) (g) Rate of Evaporation = Rate of Condensation
(sealed container at 100oC)

28. b) Solution Equilibrium – occurs at a solution’s SATURATION POINT
RATE of DISSOLVING = RATE of CRYSTALLIZATION
Example NaCl(s) NaCl(aq)

29. 2. Chemical Equilibrium Rate of FORWARD RXN = RATE of the REVERSE RXN
Rate of BREAKING BONDS = Rate of FORMING BONDS

30. Equilibrium Video
Dynamic Equilibrium:

31. LE CHATELIER’s PRINCIPLE
Video:
LE CHATELIER’s PRINCIPLE
Le Chatelier’s principle explains HOW A SYSTEM at equilibrium WILL RESPOND TO stress.
STRESS = Any change in TEMPERATURE, CONCENTRATION, or PRESSURE put upon an system at equilibrium
When a STRESS is added to a system at equilibrium, the system will SHIFT in order to relieve that stress and reach a new equilibrium.

32. SHIFT = an increase in the RATE of EITHER the forward OR the reverse rxn
SHIFT TO RIGHT (TOWARD PRODUCTS):
Rate of FORWARD reaction INCREASES (→)
Reactants Products
*Favors products
SHIFT TO LEFT (TOWARD REACTANTS):
Rate of REVERSE reaction INCREASES (←)
*Favors reactants

33. DIFFERENT TYPES OF STRESSES:
1) Concentration as initial stress: Equilibrium changes (or shifts) when a reactant or product is added (introduced) or decreased (taken away) in a reaction that is at equilibrium
When the concentration of a reactant or product is INCREASED: the reaction will SHIFT AWAY from the increase (use up the excess)

34. Example 1: 4NH3(g) + 5O2(g) 4NO(g) + 6H2O(g) + HEAT
If we add H2O(g), the system would shift to the _______ and the [NH3] would _________________.
If we add O2(g), the system would shift to the _______ and the [NO] would _________________.
If we add H2O(g), the system would shift to the _______ and the [NO] would _________________.
If we added NO(g), which concentration(s) would decrease? _________________
Left
increase
Right
increase
Left
decrease
H2O & Heat

35. Le Chatelier’s Principle
Part 1 (Conc. & Pressure) Video:

36. *TRICK: “AA” – what YOU ADD, the SYSTEM shifts AWAY
When the concentration of a reactant or product is DECREASED: the reaction will SHIFT TOWARD the side that has experienced the decrease in concentration (replaces what was taken)
Example 2: 2804 kJ + 6CO2(g) + 6H2O(l) C6H12O6(s) + 6O2(g)
1. If we remove water, the system will shift to the _______ and the [C6H12O6] will _________________.
2. If we remove O2(g), the system will shift to the _______ and the [C6H12O6] will _________________.
3. If we remove glucose, which concentrations will decrease?
4. If we remove CO2, which concentration(s) would increase?
Left
decrease
Right
increase
CO2 and H2O
CO2 and H2O
*TRICK: “AA” – what YOU ADD, the SYSTEM shifts AWAY
“TT” – what YOU TAKE, the SYSTEM shifts TOWARD

37. Le Chatelier’s Principle
Part 2 (TEMP) Video:

38. 2) Temperature as initial stress: (involves increasing or decreasing the “HEAT” component of a reaction)
Note: HEAT/ENERGY/J/KJ will either be a reactant or a product
A + B C + D + HEAT
A + B + energy C + D
When temperature (or HEAT) is increased: the reaction will SHIFT AWAY from the rxn side containing “HEAT” (in the ENDOTHERMIC direction)
When temperature (or HEAT) is decreased: the reaction will SHIFT TOWARD the rxn side containing “HEAT” (in the EXOTHERMIC direction)

39. 1. If we added heat, which concentration(s) will decrease?
Example #1: 4NH3(g) + 5O2(g) NO(g) + 6H2O(g) + HEAT
1. If we added heat, which concentration(s) will decrease?
2. If we added heat, which concentration(s) will increase?
Example #2: CO2(g) + H2O(l) kJ CH4(g) + 2O2(g)
3. If we remove heat, which concentration(s) will decrease?
4. If we remove heat, which concentration(s) will increase?
NO and H2O
NH3 and O2
CH4 and O2
CO2 and H2O

40. 3) Pressure as initial stress: Recall, pressure affects GASES ONLY
3) Pressure as initial stress: Recall, pressure affects GASES ONLY! So every other state (s, l, aq) in the reaction is UNAFFECTED for this type of stress
INCREASE PRESSURE: rxn shifts to side with LEAST # GAS MOLECULES (or least # moles of gas)
DECREASE PRESSURE: rxn shifts to side with GREATEST # GAS MOLECULES (or greatest # moles of gas)
NOTE: If the rxn contains NO GAS MOLECULES or if the rxn has the SAME # GAS MOLECULES on each side, there is NO EFFECT and NO SHIFT results from an increase or decrease in pressure

41. Example 1: CO2(g) CO2(aq)
1 gas molecule
No gas molecules
1. If we increase the pressure, the concentrations of which species will increase?
2. If we increase the pressure, the concentrations of which species will decrease?
3. If we decrease the pressure, the concentrations of which species will increase?
4. If we decrease the pressure, the concentrations of which species will decrease?
CO2(aq)
CO2(g)
CO2(g)
CO2(aq)

42. Example 2: N2(g) + 3H2(g) 2NH3(g)
1. If we increase the pressure, in which direction will the equilibrium shift? (Count moles of gases on each side 1st)
2. If we increase the pressure, the concentration of which species will increase initially?
3. If we decrease the pressure, the concentration of which species will decrease initially?
4. If we decrease the pressure, the concentration of which species will increase initially?
RIGHT
NH3
N2 and H2
NH3

43. States of Matter

44. Gas Video
Invisible Properties of Gases -

45. Kinetic Molecular Theory (KMT): A theory used to explain the behavior of gases in terms of the MOTION of their particles.
Major Assumptions of KMT:
1. Gas particles have NO volume (take up no space)
2. The gas particles do NOT attract each other (NO intermolecular forces)
3. The gas particles move in STRAIGHT, continuous motion.
4. The gas particles are perfectly ELASTIC – lose NO speed after they COLLIDE w/one another or w/ the walls of the container.

46. KMT is based on the CONCEPT or MODEL of an IDEAL GAS
An IDEAL GAS is THEORETICAL and is used to PREDICT the behavior of REAL GASES (O2, H2, He, etc.)

47. *Trick: What is your ideal Boyfriend/Girlfriend?
REAL GASES: Deviate (are different) from ideal gas laws when molecules are close together.
*Trick: What is your ideal Boyfriend/Girlfriend?
Ideal Boyfriend/Girlfriend = Hot and No Pressure
How to make a “real” BF/GF into an “ideal” BF/GF  Increase Temp; Lower Pressure
Why?
Deviate from #1: Real gases contain particles that have volume.
Deviate from #2: Real gas particles are attracted to each other.
*These deviations occur especially under high PRESSURE and low TEMPERATURE.
QUESTION: When does a REAL GAS behave most nearly like an IDEAL GAS?
At high temp and low pressure

48. Practice Regents
1. According the kinetic molecular theory of gases, which assumption is correct?
Gas particles strongly attract each other.
Gas particles travel in curved paths
The volume of gas particles prevents random motion.
Energy may be transferred between colliding particles.
2. A real gas would behave most like an ideal gas under conditions of
low pressure and low temperature
low pressure and high temperature
high pressure and low temperature
high pressure and high temperature * *

49. Gas Video
ABC’s of Gases -

50. Combined Gas Law
Boyle’s Law:
P1V1 = P2V2
Charles’s Law:
V1 = V2
T T2

51. P1= 1 atm V1= 50 mL T1= constant P2= ? V2= 150 mL T2=constant
If a problem says the gas is at STP use Table A:
Temp: 273K
Pressure: kPa (1 atm)
Example #1: Question: A gas has a volume of 50 mL, at a pressure of 1 atmosphere. The volume is increased by 100 mL, and the temperature remains constant. What is the new pressure?
Solution: Use combined gas law.
P1V1 = P2V2
T T2
P1= 1 atm
V1= 50 mL
T1= constant
P2= ?
V2= 150 mL
T2=constant
1(50) = P2(150)
(50) = P2
(150)
0.33 atm = P2

52. Solution: Use combined gas law.
Example #2: Question: At 200K, the volume of a gas is 100 mL. The temperature is raised to 300K. What is the new volume?
Solution: Use combined gas law.
P1V1 = P2V2
T T2
100 = V2
P1(100mL) = P2V2
200K K
V2(200) = 30,000 .
V2 = 30,000
200 .
V2 = 150 mL.

53. Solution: Use combined gas law.
Example #3: Question: At a temperature of 273K, a 400mL gas sample has a pressure of kPa. If the pressure is changed to kPa, at which temperature will this gas sample have a volume of 551 mL?
Solution: Use combined gas law.
(101.3kPa)(400mL) = (50.65kPa)(551mL)
273K T2
P1V1 = P2V2
T T2
40520 = T2
7,618, = (40520)T2
7,618,938.6 = T2
(40520) .
188 K = T2