2. Kinetics, Thermodynamics and Equilibrium Regents Chemistry

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

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

5. Factors Affecting Rate 1. Type of substance: Ionic substances react faster: bonds require less energy to break Ionic substances react faster: bonds require less energy to break AgNO 3 (aq) +NaCl (aq)  AgCl (s) +NaNO 3 (aq) In solution ionic solids dissociate into ions: Ag + NO 3 - Na + Cl - Covalent react more slowly: bonds require more energy to break Covalent react more slowly: bonds require more energy to break H 2 (g) +I 2 (g)  2 HI (g) Bonds must be broken then be reformed. (takes more time)

6. 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. Average kinetic energy increases and the number of collisions increases. Reactants have more energy when colliding. This increases rate.

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

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

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

10. 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. Catalysts remain unchanged during the reaction and can be reused. Activation energy: amount of energy required to “start” a reaction Activation energy: amount of energy required to “start” a reaction

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

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

13. Heat of reaction (ΔH) = E p - E R PE of Activated Complex PE of reactants (E R )Activation Energy (E a ) PE of products (E P ) Activation Energy (E a )* reverse reaction Exothermic: PE of products is less because energy was lost.

14. PE of reactants (E R ) PE of Activated Complex Heat of reaction (ΔH)Activation Energy (E a ) PE of products (E P ) Activation Energy (E a )* reverse reaction Endothermic: PE of products is more because energy was gained.

15. Catalysts

16. Thermodynamics Heat content (Enthalpy): amount of heat absorbed or released in a chemical reaction Heat content (Enthalpy): amount of heat absorbed or released in a chemical reaction Enthalpy (ΔH = H products – H reactants ) Enthalpy (ΔH = H products – H reactants )

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

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

19. Table I- Practice 1. Which reaction gives off the most energy? 2. Which reaction gives off the least energy? 3. Which reaction requires the most energy to occur?

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

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

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

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

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

25. Chemical Equilibrium Regents Chemistry

26. Reversible Reactions Most chemical reactions are able to proceed in both directions under the appropriate conditions. Most chemical reactions are able to proceed in both directions under the appropriate conditions. Example: Example: Fe 3 O 4 (s) + 4 H 2 (g) ↔ 3 Fe (s) + 4 H 2 O (g)

27. 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. 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. Rate fwd = Rate rev This is called chemical equilibrium. This is called chemical equilibrium.

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

29. Phase Equilibrium Rate of one phase change is equal to the rate of the opposing phase change. Rate of one phase change is equal to the rate of the opposing phase change. Occurs when two phases exist at the same temperature. Occurs when two phases exist at the same temperature. Example: Rate melting = Rate freezing Example: Rate melting = Rate freezing H 2 O (s)  H 2 O (l)

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

31. Chemical Equilibrium Rate forward reaction = Rate reverse reaction Rate forward reaction = Rate reverse reaction Concentration of reactants and products are constant NOT necessarily equal. Concentration of reactants and products are constant NOT necessarily equal. [reactants] and [products] is constant. [reactants] and [products] is constant.

32. 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.

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

34. Example: The Haber Process N 2 (g) + 3 H 2 (g)  2 NH 3 (g) + heat a)  [N 2 ] b)  [H 2 ] c)  [NH 3 ] d)  [NH 3 ] e)  pressure f)  pressure g)  temperature h)  temperature

35. Example: The Haber Process N 2 (g) + 3 H 2 (g)  2 NH 3 (g) + heat a)  [N 2 ]shift towards products (right) b)  [H 2 ]shift towards reactants (left) c)  [NH 3 ]shift towards reactants (left) d)  [NH 3 ]shift towards products (right) e)  pressureshift towards products (right) f)  pressureshift towards reactants (left) g)  temperatureshift towards reactants (left) h)  temperatureshift towards products (right)

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

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

38. 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. 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. Reaction continues until reactants are used up. Products are removed if: Products are removed if: Gases in liquid solution Gases in liquid solution Insoluble products (precipitate) Insoluble products (precipitate)

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