1. Kinetics , Thermodynamics and Equilibrium

2. Kinetics and Thermodynamics
Kinetics: deals with rates of reactions (how quickly a reaction occurs)
Thermodynamics: involves changes in energy that occur in reactions

3. Kinetics: Collision Theory
Collision theory states that a reaction is most likely to occur if reactant particles collide with the proper energy and orientation.
An ineffective collision reaction does not occur
An effective collision reaction occurs

4. Which event must always occur for a chemical reaction to take place?
(1) formation of a precipitate
(2) formation of a gas
(3) effective collisions between reacting particles
(4) addition of a catalyst to the reaction system

5. Factors Affecting Rate

6. 1. 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-
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)

7. 2. Temperature increase
Average kinetic energy increases and the number of collisions increases. Reactants have more energy when colliding. This increases rate.

8. 3. Concentration increase
Increases rate due to the fact that more particles are in a given volume, which creates more collisions.

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

10. 5. Pressure Increases
Increases the rate of reactions involving gases only
As pressure 
Volume  so:
spaces between
molecules 
which 
the frequency of
effective collisions

11. 6. Catalyst
Catalyst: substance that increases rate of reaction, provides a shorter or pathway for the reaction to occur.
Catalysts remain unchanged during the reaction and can be reused.

12. 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 (gas)
Catalysts speed up reactions.

15. What are the two things that
must happen in order to have an
effective collision?

19. Potential Energy Diagrams
Graphs heat during the course of a reaction.

20. Exothermic: PE of products is less because energy was lost.
PE of reactants (ER)
Activation Energy (Ea)
PE of Activated Complex
PE of products (EP)
Heat of reaction (ΔH) = Ep - ER
Activation Energy (Ea)* reverse reaction

21. Endothermic: PE of products is more because energy was gained.
PE of products (EP)
PE of reactants (ER)
Activation Energy (Ea)
Heat of reaction (ΔH)
PE of Activated Complex
Activation Energy (Ea)* reverse reaction

22. Catalysts

25. Thermodynamics
Heat content (Enthalpy): amount of heat absorbed or released in a chemical reaction
Enthalpy (ΔH = Hproducts – Hreactants)

26. ΔH = PEproducts – PEreactants
ΔH is positive when the reaction is endothermic. Heat of products are greater than reactants
ΔH is negative when the reaction is exothermic. Heat of reactants were greater than the products

27. Exothermic Releases heat Endothermic Absorbs energy PE decreases
H is negative
Energy is on the right of equation
2H2 + O2 2H2O + energy
Endothermic
Absorbs energy
PE increases
H is positive
Energy is on the left of the equation
2H2O + energy H2 + O2
J Deutsch 2003

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 (left)
Exothermic: heat is a product (right)

30. Table I- Check for understanding
Which reaction gives off the most energy?
Which reaction gives off the least energy?
Which reaction requires the most energy absorbed to occur?

31. 3

32. 1

33. 2

34. This is called the Heat of Summation,ΔH
AP topic
Hess’s Law
Hess’s Law states that the heat of a whole reaction is equivalent to the sum of it’s steps.
This is called the Heat of Summation,ΔH

35. Plan a Strategy
Evaluate the given equations.
Rearrange and manipulate the equations so that they will produce the overall equation. (change coefficients, reverse the reaction).
Cross out common reactants/ products in all steps.
Change enthalpy (ΔH) terms accordingly and add the (ΔH ) terms.
ΔH1 = ΔH2 + ΔH3

36. For example, suppose you are given the following data:
Use the data above to obtain the enthalpy change for the following reaction:

37. Work Area

38. This reaction can also be carried out in two steps:
Example #1
Our reaction of interest is:
N2(g) + 2O2(g) 2NO2(g) ΔH1 = 68 kJ
This reaction can also be carried out in two steps:
N2 (g) + O2 (g) 2NO(g) ΔH2 = 180 kJ
2NO (g) + O2 (g) 2NO2(g) ΔH3 = -112 kJ

39. 2NO (g) + O2 (g) 2NO2(g) ΔH3 = -112 kJ
If we take the previous two reactions and add them, we get the original reaction of interest:
N2 (g) + O2 (g) 2NO(g) ΔH2 = 180 kJ
2NO (g) + O2 (g) 2NO2(g) ΔH3 = -112 kJ
N2 (g) + 2O2 (g) 2NO2(g) ΔH1 = 68 kJ

40. Example #2
Using Hess’s Law, determine the ΔH for the following reaction:
4HCl(g) + O2(g) → 2 Cl2(g) + 2H2O(g)
This reaction can take place in a series of two steps:
H2(g) + Cl2(g) → 2HCl(g) ΔH =- 185 kJ
2H2(g) + O2(g) → 2H2O ΔH = kJ

41. Work Area

42. 3. Using Hess’s Law, determine the ΔH for the following reaction:
3C(s) + 4H2 (g) → C3H8 (g)
This reaction can take place in a series of three steps:
C(s) + O2 (g) → CO2 (g) ΔH = -394 kJ
C3H8 (g) + 5O2 (g) → 3CO2 (g) + 4H2O(l) ΔH = kJ
H2 (g) + ½ O2 (g) → H2O (l) ΔH = -286 kJ

43. Entropy (ΔS) Definition: randomness, disorder in a sample of matter
Gases have high entropy
Solids have low entropy

44. How to increase entropy:
Phase change from s  l  g
Increasing the moles of gases
Dissolving a substance (going from s to aq)
Decomposition reaction

45. 2

46. 2

47. 2

48. Spontaneous Reactions
This is a reaction that will occur naturally. It does not need outside help.
Nature favors low energy (more stable) and high entropy

49. 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.

50. When DH is - and entropy is +
(exothermic) (greater disorder)
the reaction would be spontaneous

51. Are all spontaneous reactions exothermic and with a greater system disorder?

52. Gibbs free energy is a measure of chemical energy
AP topic
Gibbs Free Energy
Gibbs free energy is a measure of chemical energy
Gibbs free energy, ΔG, enables us to predict whether a reaction will be spontaneous

53. The free energy change is defined as: ΔG = ΔH – TΔS
A spontaneous reaction depends on enthalpy, entropy, and temperature

54. ∆G = ∆H - T∆S
If a reaction is
exothermic (negative ∆ H) and entropy increases (positive ∆S)
then ∆G must be NEGATIVE
So the reaction is spontaneous

55. ∆G = ∆H - T∆S So the reaction is not spontaneous
If a reaction is
endothermic (positive ∆H) and
entropy decreases (negative ∆S)
then ∆Go must be POSITIVE
So the reaction is not spontaneous

56. Gibbs Free Energy, G ∆G = ∆H - T∆S ∆Ho ∆So ∆Go Reaction
exo(–) increase(+) – Prod-favored
endo(+) decrease(-) + React-favored
exo(–) decrease(-) ? T dependent
endo(+) increase(+) ? T dependent

58. Chemical Equilibrium

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

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

62. Reactions that go to completion do not reach equilibrium:
Equilibrium is not reached in an open system . The system must be closed. Gases can escape!
Equilibrium is not reached if there is an insoluble product (precipitate)

63. Phase Equilibrium
Rate of phase change is equal when two phases exist at the same temperature.
Example: H2O (l)  H2O (g)
This happens on the plateau of the heating/cooling curve.

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

67. Chemical Equilibrium
Rateforward reaction = Ratereverse reaction Concentration of reactants and products are constant NOT equal.

68. At equilibrium, the forward and reverse reaction rates are equal.

69. At equilibrium, the concentrations of the reactants and products are constant

71. The Concept of Equilibrium
RECC
That’s what equilibrium means to me!!
Rates are Equal
Concentrations are Constant

74. 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.

75. Temperature, pressure, changes in reactant or product concentrations
Stresses include:
Temperature, pressure, changes in reactant or product concentrations

76. N2 (g) + 3 H2 (g)  2 NH3 (g) + heat
The letter T for temperature goes on the
side that has heat or kJ.
The Letter P goes on the side that has a
greater number of gas moles.

77. The rule for determining which way the reaction shifts is:
Add To shifts Away (increase)
Take Out shifts Towards (decrease)

78. N2 (g) + 3 H2 (g)  2 NH3 (g) + heat
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

79. Example: The Haber Process
N2 (g) + 3 H2 (g)  2 NH3 (g) + heat
 [N2] shift towards products (right)
 [H2] shift towards reactants (left)
 [NH3] shift towards reactants (left)
 [NH3] shift towards products (right)
 pressure shift towards products (right)
 pressure shift towards reactants (left)
 temperature shift towards reactants (left)
 temperature shift towards products (right)

80. 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.

83. c State the effect on the number of moles of NH3 (g) if a catalyst is introduced into the reaction system. Explain why this occurs.

84. 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