1. Kinetics , Thermodynamics and Equilibrium
Regents Chemistry

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
Measured in:
#moles of reactant used/unit time Or
# moles of product formed/unit time
Frequency of collisions: more collisions = faster rate
Effective collisions: must have 1) proper orientation and 2) enough energy

4. Factors Affecting Rate
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)

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

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

7. Factors Affecting Rate
4. Surface Area Increase
Increases rate due to increased reactant interaction or collisions (powder vs. lump)

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

9. Factors Affecting Rate
6. Catalyst: substance that increases rate of reaction, provides a shorter or alternate pathway by lowering the activation energy of the reaction.
Catalysts remain unchanged during the reaction and can be reused.
Activation energy: amount of energy required to “start” a reaction

10. 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.

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

12. 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

13. 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)
Activation Energy (Ea)* reverse reaction
PE of Activated Complex

14. Catalysts

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

16. ΔH = Hproducts – Hreactants
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

17. 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

18. Table I- Practice Which reaction gives off the most energy?
Which reaction gives off the least energy?
Which reaction requires the most energy to occur?

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

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

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

22. 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.

23. 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.

24. Chemical Equilibrium
Regents Chemistry

25. 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)

26. 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.

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

28. 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.
Example: Ratemelting = Ratefreezing
H2O (s)  H2O (l)

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

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

31. 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.

32. 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

33. Example: The Haber Process
N2 (g) + 3 H2 (g)  2 NH3 (g) + heat
 [N2]
 [H2]
 [NH3]
 [NH3]
 pressure
 pressure
 temperature
 temperature

34. 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)

35. 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

36. 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.

37. 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)

38. 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