1. Unit 9 Kinetics and Equilibrium
Ms. Randall

2. Lesson 2: Reaction Rates Objective: To compare and contrast factors that affect the rate of a chemical reaction

3. 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.
Demo water to milk to juice

4. Reaction Mechanism
STEP BY STEP PROCESS needed to make a product; how you get from “a” to “b” (like a recipe).

5. 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)

6. Collision Theory
In order for a reaction to occur, reactant PARTICLES MUST COLLIDE
Proper amount of ENERGY
Proper ALIGNMENT/ DIRECTION/ ORIENTATION

7. Only when particles collide and these two conditions are met will there be an EFFECTIVE COLLISON, resulting in a reaction
Model

8. What Determines The Rate of a Reaction?
NUMBER OF STEPS= more steps cam mean a slower reaction
RATE DETERMINING STEP= the SLOWEST STEP of the reaction; most important factor influencing reaction rate.

10. Factors Affecting the Rate of Reaction
SIX FACTORS that affect the rate of reaction by changing the number of effective collisions that take place between particles.
The MORE EFFECTIVE COLLISIONS, THE FASTER THE REACTION!

11. Type of substance:
Ionic substances react faster: bonds require less energy to break
AgNO3 (aq)+NaCl(aq)AgCl(s)+NaNO3 (aq)
In solution ionic solids dissociate
into ions:
Ag+ NO Na+ Cl-

12. Covalent react more slowly: bonds require more energy to break
H2 (g)+I2 (g)2 HI (g)
Bonds must be broken then be reformed. (takes more time)

13. IONIC substances react FASTER
COVALENT substances react SLOWER

14. Concentration
Increases rate due to the fact that more particles are in a given volume, which creates more collisions.
Low Concentration
High Concentration

15. Pressure Increases the rate of reactions involving gases only
As pressure  Volume  so:
spaces between molecules 
 frequency of effective collisions
Low Pressure
High Pressure

16. Temperature
INCREASE temperature, INCREASE Average Kinetic Energy. More particles colliding!
High Temp
Low Temp

17. Surface Area
Increases rate due to increased reactant interaction or collisions (powder vs. lump)

18. Catalyst
SPEEDS UP THE REACTION WITHOUT CHANGING THE NATURE OF THE REACTANTS/PRODUCTS

19. Provides a SHORTCUT or ALTERNATIVE PATHWAY for the mechanism
**NOTE: catalysts DO NOT change the nature of the reaction (the same reaction is still occurring), the reaction is just completed in a shorter amount of
time!
Activation energy: amount of energy required to “start” a reaction

20. “Real World Example” Catalytic Converter in vehicles
The NO molecules Bind to the Catalyst Surface which Weakens their Bonds.
The job of the catalytic converter is to convert harmful pollutants into less harmful emissions before they ever leave the car's­ exhaust system.

21. Quick Review – Factors that affect reactions
Ionic solutions have faster reactions than molecule compounds. (bonding)
Temp.  Rate
 conc. rate
 surface area  rate
 Pressure  rate,  P  rate
Catalysts speed up reactions.

22. Check your understanding and practice

23. Lesson 3: Energy in reactions Objective: To apply changes during a reaction in energy to a PE diagram

24. A reaction is the breaking and reforming of bonds

25. 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
Enthalpy (ΔH = Hproducts – Hreactants)

26. Two types of reactions Reactions that release energy EXOTHERMIC
Δ H = negative value (-)
energy released (on right)
A + B C + D + ENERGY

27. 2. Reactions that absorb/gain energy ENDOTHERMIC Δ H = positive value (+) energy absorbed (on left) A + B + ENERGY C + D

28. Table I
Includes heats of reaction for combustion, synthesis (formation) and solution reactions.
You must remember equation stoichiometry (balanced equations).
Endothermic: heat is a reactant
Exothermic: heat is a product

29. Forward Reaction
Reading LEFT TO RIGHT in a reaction; reaction moves toward the right

30. Reverse Reaction
Reading RIGHT TO LEFT in a reaction; reaction moves toward the left

31. Table I
Tells us if particular reactions are exothermic or endothermic based on sign of the Δ H value
+ 66.4
ENDOthermic
-91.8
EXOthermic
ENDOthermic
+ 91.8

32. Potential Energy Diagrams
Chemists use diagrams called Potential Energy Diagrams to illustrate the potential (or stored) energy changes that occur during specific chemical reactions.

33. 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!)

34. The ACTIVATED COMPLEX is lowered which also lowers the ACTIVATION ENERGY needed for the reaction to get started (gets reaction started easier and makes the reaction complete itself faster)

35. HOW EXACTLY DOES A CATALYST SHORTEN THE REACTION TIME NEEDED FOR A REACTION TO COMPLETE?

36. ENDOTHERMIC Potential Energy Diagrams POSITIVE ΔH
Product side (to the R) always HIGHER than the reactant side (to the L) meaning that ENERGY is ABSORBED

37. Label the following: A = Potential Energy of the Reactants
B = Potential Energy of the Products
C = Potential Energy of the Activated Complex
D = Activation energy of the Forward rxn
E = Activation Energy of the Reverse rxn
F = Heat of the Reaction (ΔH = Hp – Hr)

38. D E F C B A

39. EXOTHERMIC Potential Energy Diagrams NEGATIVE ΔH
Product side (to the R) always LOWER than the reactant side (to the L) meaning that ENERGY is RELEASED

40. Label the following: A = Potential Energy of the Reactants
B = Potential Energy of the Products
C = Potential Energy of the Activated Complex
D = Activation energy of the Forward rxn
E = Activation Energy of the Reverse rxn
F = Heat of the Reaction (ΔH = Hp – Hr)

41. D C E F A B

42. Check your understanding and practice
Check your understanding and practice

43. Lesson 4:Equilibrium Objective: To compare and contrast the different types of system equilibriums

44. Reversible Reactions
Most chemical reactions are able to proceed in both directions under the appropriate conditions.
Example:
Fe3O4 (s) H2 (g) ↔ 3 Fe(s) + 4 H2O(g)

45. Reversible Reactions cont.
In a closed system, as products are produced they will react in the reverse reaction until the rates of the forward and reverse reactions are equal.
Ratefwd = Raterev
This is called chemical equilibrium.

46. EQUILIBRIUM
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…

47. Equilibrium is represented by DOUBLE ARROWS instead of a single arrow.
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.

48. TYPES OF EQUILIBRIUM Phase equilibrium Solution Equilibrium
All occur in CLOSED SYSTEMS
*ITS ALL ABOUT THE EQUAL RATES!
Phase equilibrium
Solution Equilibrium
Chemical Equilibrium

49. Equilibrium
Equilibrium is dynamic condition where rates of opposing processes are equal.
Types of Equilibrium:
Phase equilibrium
Solution Equilibrium
Chemical Equilibrium

50. Phase Equilibrium
Rate of one phase change is equal to the rate of the opposing phase change.
Occurs when two phases exist at the same temperature.

51. Solution Equilibrium Rate of dissolving = rate of crystallization
Occurs in saturated solutions

52. Chemical Equilibrium Rateforward reaction = Ratereverse reaction
Concentration of reactants and products are constant NOT necessarily equal.
[reactants] and [products] is constant.

53. The Concept of Equilibrium
As a system approaches equilibrium, both the forward and reverse reactions are occurring.
At equilibrium, the forward and reverse reactions are proceeding at the same rate.

54. Check your understanding and practice

55. Lesson 5: Le Chatelier’s Principle Objective: To apply Le Chatelier’s principle to a system at equilibrium and predict the changes on the system

56. LE CHATELIER’s PRINCIPLE
Le Chatelier’s principle explains HOW A SYSTEM AT EQUILIBRIUM WILL RESPOND TO STRESS.

57. Le Chatelier’s Principle
Whenever stress is applied to a reaction at equilibrium, the reaction will shift its point of equilibrium to offset the stress.
Stresses include:
Temperature, pressure, changes in reactant or product concentrations

58. SHIFT
An increase in the RATE of EITHER the forward OR the reverse reaction.

60. Demo of LeChatelier’s Principle
NO2  N2O4
Cooling will cause a shift that produces heat (in this case the forward reaction). Thus, the forward reaction is Exothermic.

61. 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

62. When the concentration of a reactant or product is INCREASED:
The reaction will SHIFT AWAY from the increase (use up the excess)
LEFT
INCREASE
RIGHT
INCREASE
LEFT
DECREASE
[NO2], [H2O]

63. 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)
LEFT
[C6H12O6]
DECREASE
RIGHT
[C6H12O6]
INCREASE
[CO2], [H2O]
[CO2], [H2O]

65. 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

66. When temperature (or HEAT) is increased:
The reaction will SHIFT AWAY from the rxn side containing “HEAT”
(in the ENDOTHERMIC direction)

67. When temperature (or HEAT) is decreased:
The reaction will SHIFT TOWARD the rxn side containing “HEAT” (in the EXOTHERMIC direction)

68. Example #1:
[NO2], [H2O]
[NH3], [O2]

69. Example #2:
[CH4], [O2]
[CO2], [H2O]

70. 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.

71. INCREASE PRESSURE: rxn shifts to side with LEAST # GAS MOLECULES (or least # of moles of gas)
DECREASE PRESSURE: rexn shifts to side with GREATEST # GAS MOLECULES (or greatest # moles of gas)

72. NOTE: If the rexn 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

73. Example 1:
CO2 (aq)
CO2 (g)
CO2 (g)
CO2 (aq)

74. The Haber Process Application of LeChatelier’s Principle
N2 (g) + 3 H2 (g)  2 NH3 (g) + 92 kJ
increase pressure
Shift 
decrease Temp
remove NH3 add N2 and H2
****Maximum yields of NH3 occurs under high pressures, low temperatures and by constantly removing NH3 and adding N2 & H2

75. Equilibrium shifts due to stresses:
Concentration increase shift away from increase
Concentration decrease shift toward decrease
 pressure shifts in direction of fewer gas molecules.
 pressure shifts in direction of more gas molecules
 temperature favors endothermic reaction
Shift away from heat
 temperature favors exothermic reaction
Shift towards heat

76. Effect of Catalyst:
Addition of catalysts changes the rate of both the forward and reverse reactions.
There is no change in concentrations but equilibrium is reached more rapidly.

77. Reactions that go to completion:
Equilibrium is not reached if one of the products is withdrawn as quickly as it is produced and no new reactants are added.
Reaction continues until reactants are used up.
Products are removed if:
Gases in liquid solution
Insoluble products (precipitate)

78. Check your understanding and practice!

79. Lesson 6: Entropy/ Enthalpy Objective: To define the conditions necessary for a spontaneous reaction to occur

80. Entropy (ΔS) Randomness, disorder in a sample of matter
Nature tends to proceed to a state of GREATER entropy, or disorder.
Gases have high entropy
Solids have low entropy

81. ENTROPY (ΔS):
Degree of RANDOMNESS or CHAOS or DISORDER or “MESSINESS” in a system

83. Increasing ΔS Phase change from s  l  g Mixing gases
Dissolving a substance

84. What physical state of matter has the highest entropy?
GAS aq g s l

85. Entropy INCREASES when a compound is broken down
*Entropy INCREASES when a compound is broken down. *Entropy DECREASES when a compound is created and bonds are formed.

86. For the following determine if there is an increase, decrease, or no change in entropy:

87. NONE
NOTE: If all states are the same in reaction, count up the # of molecules on each side

88. ENTHALPY (ΔH)
the ENERGY in a system; nature tends to proceed to a state of LOWER enthalpy, or energy
Exothermic reactions RELEASE energy & move to LOWER energy state

89. ** Most common types of reactions because less energy has to be put in to get the reaction started (LOWER activation energy)

90. Endothermic reactions ABSORB energy & move to a HIGHER energy state (HIGHER ACTIVATION ENERGY NEEDED)

91. ** Not as common because more energy must be put in to get the reaction started (HIGHER activation energy)

92. Spontaneous Reactions
Nature favors low energy (more stable) and high entropy
Reactions are spontaneous when heat (ΔH) decreases and entropy (ΔS) increases
ΔH = (-)
ΔS= (+)

93. Stability of Products and H
Help determine if a reaction is spontaneous
Products tend toward Lower energy (-ΔH)
Products tend toward more randomness (+ΔS)
Products of exothermic reactions are usually more stable. Result in lower amounts of heat.
The more negative the H, the more stable the product is.
Gas products result in increased Entropy.

94. Analogy: Your Bedroom
You like to have low enthalpy (low energy) when it comes to household chores.
As a result, your room tends to have high entropy (very messy, disorderly).
This is what nature prefers: low enthalpy and high entropy.

95. Check your understanding and practice