1. Equilibrium Part I Notes

2. The Concept of Equilibrium Ex) elevator, football game, moving walkway

3. 1. Reversible Reactions: Chemical reaction in which the products can regenerate the original reactants. A double arrow is used to express a reversible reaction. Ex) 2NO 2(g) ↔ N 2 O 4(g) can also be used reactants yield products ↔ or 2. Some reactions can reverse on their own, some under certain conditions, (temp, pressure or a catalyst) and some will not reverse. Ex) Single Replacement Rxns: 3CuCl 2(aq) + 2Al (s) → 3Cu (s) + 2AlCl 3(aq)

4. 3. Chemical Equilibrium: State of a product’s and reactant’s concentrations remaining constant, not equal. This is because the rate of the forward and reverse reactions are equal. The symbol “[ ]” denotes concentration.

5. A.Reaction rates are affected by concentration. If the concentration of the reactant increases, the rate of the forward reaction will increase. B.As the reaction proceeds, the concentration of the reactants decrease and the products increases. This will cause the forward reaction to decrease.

6. 4. Reaching equilibrium does not mean the reaction has stopped. It only means the rate of the forward and reverse reactions are equal. This is not a static equilibrium, but a dynamic equilibrium that stays constant over time.

7. Concentration becomes Constant NOT Equal Rates become Equal

8. Reaction Rates & Equilibrium: A. Collision Theory: (Theoretical Model) 1. In order for reactions to occur between substances, the particles must collide. 2. According to the Collision Theory, a successful collision occurs when A. the collision is energetic enough, and B. the particles collide with the correct orientation.

9. 3. Effective collisions lead to the formation of products; ineffective collisions do not lead to the formation of products. 4. During a collision, kinetic energy (mass & velocity) from the reactants motion is converted to potential energy within the products.

10. Ineffective Collision – Insufficient Energy - No Products Effective Collision – Sufficient Energy – Forms Products Effective and Ineffective Collisions

11. Effective collisions have enough energy, and the correct orientation to form products. Ineffective collisions revert to the original products.

12. 5. A successful collision results in bond breaking (endothermic) and bond forming (exothermic). 6. The minimum energy needed to produce an effective collision is called the activation energy for the reaction. It’s abbreviation is E a.

14. Activation Energy

15. 7. A transitional structure results from a successful collision. The structure is present while old bonds are breaking and new bonds are forming. It is called an activated complex and is unstable and short-lived. It is neither reactant nor product. 8. Most reactions occur in a series of steps called a reaction mechanism. A→ B B→ C C→ D 9. Often, one step is slower than the others. This is called the rate-determining step. Ex) relay race Elementary steps Intermediates (found in one reaction & used in a later step)

16. 10. A substance that increases the rate of a chemical reaction by providing a mechanism with a lower energy of activation is called a catalyst. 11. A system that has just one phase is called a homogeneous system, (g→g→g) while a system that has more than one phase is called a heterogeneous (s→l→g) system.

17. 12. Energy Diagram: shows changes in energy during a reaction. Endothermic Exothermic

18. Reverse Reaction Forward Reaction Energy Diagram for a Chemical Rxn. Energy of Reactants Activated Complex Energy of Products Reaction Pathway EaEa forward reaction (80 kJ) EaEa reverse reaction (140 kJ) Potential (kJ) (Exo) (Endo) ∆H = -60 kJ (for.) +60 kJ (rev.) Added Catalyst

19. Factors Affecting Reaction Rates 1. Chemical Kinetics is concerned with the rate at which a reaction occurs. 2. Reaction rate is a measure of the rate or speed of a chemical reaction. Reaction rate is determined by measuring the change in concentration of reactants and products over a certain amount of time. These reaction rates are determined experimentally. 3. Rate-influencing factors are those factors that affect rate of reactions by altering the frequency, orientation, or energy at which particles collide.

20. 4. According to the Collision Theory, the following 5 factors affect the rate of a reaction: A. Nature of Reactants: 1. Structure: complexity of bonds broken and formed & the orientation.

21. 2.State: homogeneous (faster) heterogeneous (slower) 2 liquids would usually react quickly A gas and a solid might react slowly

22. B. Temperature: Rule of thumb – every 10°C increase doubles the reaction rate. Reaction Rate

23. C. Concentration: increase concentration, increase # collisions. Aqueous solutions change [conc], gases change pressure. Solids and pure liquids (like water) cannot change concentration.

24. D. Surface Area: Increase surface area, increases frequency of collisions. This is especially true for heterogeneous systems (s→l→g).

25. E. Catalyst: Increases the rate of reaction without being consumed in the reaction. Catalysts speed up the reaction rate by lowering the activation energy needed for the reaction to occur. Uncatalyzed reaction (slow) Catalyzed reaction (fast)

27. 5. Inhibitors: Decrease the rate of the reaction by taking the place of a reactant and stopping the reaction (opposite of a catalyst.)

28. Energy diagram involving a catalyst:

29. 1. What is in the reaction vessel at time = 0? H 2 + N 2 2. Write the forward reaction: 3H 2 + N 2 → 2NH 3 3. What kind of reaction is the forward reaction? synthesis 4. Write the reverse reaction: 2NH 3 → 3H 2 + N 2 5. What kind of reaction is the reverse reaction? decomposition 6. Over time the concentration of which substance(s) decreases? H 2 + N 2 7. Over time the concentration of which substance(s) increases? NH 3 8. Mark on the graph with a dashed line when equilibrium is reached. 9. At equilibrium which substance(s) is(are) present in the greater concentration? H 2 + N 2 10. Is the forward or reverse reaction favored? reverse