1. TOPIC 11: KINETICS AND EQUILIBRIUM
Part 1 – Kinetics
Part 2 – General Equilibrium

2. AIM#1 : How does Kinetics affect the rate of chemical reactions?
Kinetics: the branch of chemistry that deals with the rates (how fast the reaction is) of chemical reactions
Collision Theory:
In order for a reaction to occur, reactants must collide with each other
An effective collision is when reactants come together with the correct amount of energy and in the correct position to form a product
Example: Fender Bender vs. Head On Collision

3. As the amount of effective collisions increases, the faster products are formed
How can we increase the reactions rate (increase the # of effective collisions??

4. FACTORS THAT AFFECT THE RATE OF CHEMICAL REACTIONS
Nature of Reactants
Concentration
Surface Area
Pressure
Temperature
Adding a Catalyst

5. 1. NATURE OF REACTANTS:
Reactions involve the breaking of old bonds and the formation of new bonds. In general:
Covalent bonds are slower to react than ionic bonds
Breaking more bonds requires more energy than making bonds during collisions

6. 2. CONCENTRATION:
Increase concentration increase rate of reaction
(More particles available for collisions)
(especially if volume is decreased)

7. 3. SURFACE AREA:
the more surface area that is exposed the more chances there are for collisions (effective collisions) and will increase rate of reaction
Demo: Lycopodium lump vs. lycopodium poder
Sugar lump vs. granulated sugar
Alka Seltzer tablet vs. ground tablet

8. No effect on solids and liquids, only gases
4. PRESSURE:
No effect on solids and liquids, only gases
Increasing pressure, decreases the volume, increasing the rate of effective collisions – increases the rate of reaction
Demo Example: Poston/Blocking Syringe

9. 5. TEMPERATURE:
Increasing temperature increases kinetic energy of molecules, leading to an increase in the amount of effective collisions – increasing the rate of reaction
Demo examples:
Dissolving copper sulfate crystals in cold water and hot water
Glow stick in hot water and cold water – hot water will get really bright fast and then burn out increasing the temperature increases the rate of the reaction, in cold water it will stay dim but will glow longer then the one in hot water since we slowed down the rate of the reaction

10. 6. CATALYST
Addition of a catalyst increases the rate of the reaction by providing a different and easier pathway for the reaction

11. AIM #2: HOW CAN WE CLASSIFY ENERGY IN CHEMICAL REACTIONS?
Enthalpy (H) – flow of energy (heat exchange) at constant pressure when two systems are in contact
Measure only the change in enthalpy ΔH (the difference between the potential energies of the products and reactants)

12. Δ H = + reaction is endothermic and heat energy is added into the system
Δ H = - reaction is exothermic and heat energy is lost from the system

13. Look at the reactions (where is E term. ) CH4 + 2O2 CO2 + 2H2O + 890
Look at the reactions (where is E term?) CH4 + 2O CO2 + 2H2O kJ N2 + O kJ NO2
Look at the H = heat of the reaction (Table I)
Potential Energy Diagrams
H = PE products – PE reactants

15. Practice using Table I
A negative value means that the reaction is exothermic
If we were to rewrite the equation with the heat of the reaction shown it would look like this
It is added to the right side of the equation because it is exothermic reaction in which heat is released as a product
Exothermic can be also called spontaneous

16. A positive value means that the reaction is endothermic
If we were to rewrite the equation with the heat of the reaction shown it would look like this
It is added to the left side of the equation because it is endothermic reaction in which heat is absorbed as a reactant

17. ENTROPY ΔS: disorder or randomness of the matter and energy of a system (more disordered/dispersal is favored)
Nature favors CHAOS (high entropy low energy)

18. Thermodynamically favored processes or reactions are those that involve a decrease in internal energy of the components and increase in entropy
These are spontaneous or thermodynamically favored

19. AIM: What is entropy?
Ex)Predict which has the largest increase in entropy:
CO2(s)  CO2(g)
H2(g) + Cl2(g)  2HCl(g)
KNO3(s)  KNO3(l)
C(diamond)  C(graphite)

20. 2 Conditions for a reaction to be considered spontaneous 1
2 Conditions for a reaction to be considered spontaneous 1. Tendency toward lower energy (PE) exothermic ΔH = (-) Table I * products have lower energy and more stable

21. 2. Tendency toward randomness ΔS = (+) Entropy  measure of randomness or disorder Physical changed and entropy :
Solid
Liquid
Gas
Intermediate Entropy
Low Entropy
High Entropy
Table I (stability) greater exothermic reaction the more stable the products are ( lower PE)

22. Spontaneous reactions: DOESN’T REQUIRE EFFORT (proceed on their own without intervention) Ex) ice melting, nuclear reactions- natural transmutation
WILD FIRES- dry and hot in southern california

23. Chemical changes and Entropy:
Free elements (ex. O2, Na, Fe) high entropy
Compounds (ex. H2O, NH3) lower entropy
***NATURE FAVORS REACTIONS THAT HAVE LOW ENERGY AND HIGH ENTROPY***
ΔH  (negative, exothermic)
ΔS  (positive, high entropy)

24. AIM # 3: HOW CAN WE INTERPRET POTENTIAL ENERGY DIAGRAMS?
Show how the potential energy of reactant particles changes to chemical potential energy stored in bonds
PE diagrams keep track of PE changes during a chemical reaction in stages

25. For Example: A + B AB
The forward reaction is read from left to right
Compare the potential energy of the reactants to the potential energy of the products in the forward reaction (PE diagram #1)
PE REACTANTS: 25 Joules PE PRODUCTS: 75Joules
*Energy must have been absorbed during the reactions
PE diagrams with this pattern represent endothermic reactions

26. AIM: HOW CAN WE INTERPRET POTENTIAL ENERGY DIAGRAMS?
Label 1ST line: H ( -)
2nd line: exothermic reaction
FORWARD REACTION!!

27. AIM: HOW CAN WE INTERPRET POTENTIAL ENERGY DIAGRAMS?
Lets Label the Diagram
PE of Reactants
PE of Products
H (heat of reaction) = H = PE products – PE reactants
If H is positive [PE products  PE reactants] – endothermic
If H is negative [PE products  PE reactants] – exothermic
Table I – shows different chemical reactions and H for each one

28. AIM: HOW CAN WE INTERPRET POTENTIAL ENERGY DIAGRAMS?
Label the Diagram
PE of Activated Complex – intermediate molecule that forms when reactants have an effective collision. It is unstable and temporary
Activation Energy of forward reaction – amount of energy needed to start the reaction in the forward direction
Activation Energy (with catalyst) – amount of energy needed to start the reaction if a catalyst is added
* Catalysts – speed the reaction rate by lowering the activation energy needed to start a reactions (gives an alternate pathway)

29. For Example: A + B AB
Compare the potential energy of the reactants to the potential energy of the products in the reverse reaction (PE diagram #2)
PE REACTANTS: 75 Joules PE PRODUCTS: 25Joules
*Energy must have been released during the reactions
PE diagrams with this pattern represent exothermic reactions

30. 1st line H (+)
2nd line Endothermic

32. Which diagram is endothermic? Exothermic?

33. EQUILIBRIUM AIM # 4: What is equilibrium?
Equilibrium: when the forward and reverse reactions occur at the same rate – it is a state of balance
Physical Equilibrium: the changes that take place in chemical reactions during physical processes such as changes of state or dissolving.
Phase Equilibrium: Equilibrium between phases
Solid and liquid phase – during melting the rate of dissolving is equal to the rate of crystallization in a closed container (system) (H20 (s) H2O (l))

34. EQUILIBRIUM
2. Liquid and Gas phase – during this phase the rate of evaporation is equal to the rate of condensation in a closed container (H20 (l) H2O (g))

35. Solution Equilibrium:
1. Solid/Liquid Solution – saturated solutions are examples of solid/liquid solution equilibrium in a closed system (C12H6O11 (s) / C12H22O11 (aq)
2. Gas/Liquid Solution – in a closed system or container, there is equilibrium between the gaseous and dissolved state of the gas
Chemical Equilibrium
as time progress in a reaction, the concentration of the reactants can decrease causing the overall reaction to slow. As this occurs the concentration of the products can increase causing the reaction to reverse (indicated by a double arrow)

36. Le Chateliers Principle
AIM # 5: What are external factors that affect a reaction at equilibrium?
Le Chateliers Principle
If a stress is applied to a system at equilibrium the equilibrium will shift to release the effects of the stress
Stressors:
Temperature
Concentration
Pressure

37. TEMPERATURE
If you increase temperature, ALL reactions will speed up. The endothermic reaction will speed up the most (away from where the heat is) Increase in temperature favors the endothermic reaction Decrease in temperature will favor the exothermic reaction * *The side you shift towards will increase and the side you shift away from will decrease**
NOTE: Only temperature can disturb equilibrium and cause a change. The system will shift to establish a new equilibrium
Demonstration: Equilibrium involving Cobalt (II) Chloride ions-
Temperature:
MATERIALS:
1. Prepare a 0.2M cobalt chloride solution but dissolving 2.6 g of CoCl2- 6H2O in 100mL of distilled water
2. Hot water bath
3. Ice water bath
Cobalt equilibrium system is represented by the following equation
CoCl2 * 6H2O <-> CoCl2 + 6H2O ΔH = + 50kJ
(PINK) (BLUE)
FORWARD ENDO
REVERSE EXO
PROCEDURE :
Add some pink from the system to a test tube and place it into the hot water bath, or carefully warm the solution using a Bunsen burner. The solution will turn blue
Place the blue test tube in the ice water bath; it will return back to its pink color
EXPLANATION:
- THE FORWARD REACTION IS AN ENDOTHERMIC REACTION. According LeChateliers Principle adding heat favors the endothermic direction. Thus warming the system will favor the blue cholor complex side of the reactions. When placed in the ice bath, the exothermic direction will now be favored, shifting to the pink hydrated side

38. 4Nh3 + 5O2  4NO + 6H2O + heat

39. CONCENTRATION 4Nh3 + 5O2  4NO + 6H2O + heat
If you increase the concentration of a substance the reaction will shift away from the substance
4Nh3 + 5O2  4NO + 6H2O + heat

40. Le Chateliers Principle
Increase conc. of reactants
Favors FORWARD reaction
Speeds up the forward direction
SHIFTS RIGHT
Decrease conc. of reactants
Favors REVERSE reaction
Speeds up the REVERSE direction
SHIFTS LEFT
Increase conc. of products
Favors REVERSE reactions
Speeds up the REVERSE direction
Decrease conc. of the products
Speeds up the FORWARD direction
Demonstration: Equilibrium involving Cobalt (II) Chloride ions-
Concentration
MATERIALS:
Prepare a 0.2M cobalt chloride solution but dissolving 2.6 g of CoCl2- 6H2O in 100mL of distilled water
12M HCl
Cobalt equilibrium system is represented by the following equation
CoCl2 * 6H2O <-> CoCl2 + 6H2O ΔH = + 50kJ
(PINK) (BLUE)
FORWARD ENDO
REVERSE EXO
PROCEDURE :
Begin with 50mL of the pink hydrated form of the cobalt ion in a larger 250mL Erlenmeyer flask
Carefully add HCl to the solution until the blue form of the equilibrium system appears
Add distilled water to the solution in the beaker or flask. The equilibrium will shift back to the pink form
EXPLANATION:
When HCl is added to the pink hydrated form of the cobalt ion, the concentration of Cl- increases (H+ remains in solution as a spectator ion) The increase in Cl- causes the equilibrium to shift to the product side, and the solution turns blue
Adding water will shift the equilibrium to the left, forming more pink hydrated cobalt ion.
Q? What does this pink color tell us?
A: have reactant in the test tube
**add HCl to test tube see color change **
Q? Why did adding HCl turn the solution blue?
A: More product was formed
Conclusion: changing the concentration of the reactant or product will affecgt the direction of equilibrium
Notes: adding HCl – increase in Cl ions

41. CONCENTRATION ADD AWAY (increase) TAKE TOWARDS (decrease)
Pen/pencil will go up on which ever side you shift towards
Trick only some students will like/use this method --- can copy it down in their notes if they choose to

42. PRESSURE
Need to know how many gas molecules are on the reactant side and on the product side Equal # of gas molecules – pressure will have no effect!! Increase pressure – shift from more gas molecules  towards less gas molecules Decrease pressure – shift from less  toward more

43. PRESSURE
Ex) 4NH (g) + 5O2(g) 4NO(g) + 6H2O(g) = 9 gas molecules 4+6 = 10 gas molecules Increase pressure: shift to the left (more to less) 9 10 Decrease pressure: shift to the right (less to more) 9 10

44. What are the conditions at equilibrium?
At equilibrium the rate of the forward reaction is equal to the rate of the reverse reaction
Concentrations are constant not equal
Adding a catalyst would speed up the forward and reverse reactions to the same extent