1. Chemical Equilibrium and Reaction Rates
Chapter 14 & 16

2. Chemical Equilibrium
In an equilibrium reaction, the rate of the reaction of one direction will equal the rate of the reaction in the opposite direction.

3. Chemical Equilibrium
Reversible reactions: Products take part in a separate reaction to reform the reactants.
Some reactions can reverse direction; however, some need added energy to proceed in the reverse direction.
A double arrow is used to express a reversible reaction. ↔↔

4. Chemical Equilibrium ↔ or
The rate of a reaction is determined by the following factors: concentration, temperature, pressure.
As the rate of the forward reaction is equal to the rate of the reverse reaction, the reaction has reached a state of equilibrium.
↔ or

5. Chemical Equilibrium
Chemical Equilibrium: is the state in which the concentration of the reactants and products remain constant with time because the rate at which they are formed in each reaction equal the rate they are consumed in the opposite reaction.
Chemical equilibrium is called dynamic because it is always changing.

6. Chemical Equilibrium

7. Equilibrium

8. Concentration becomes Constant NOT Equal
Rates become Equal

9. 1. What is in the reaction vessel at time = 0?
H2 + N2
2. Write the forward reaction:
3H2 + N2 → 2NH3
3. What kind of reaction is the forward reaction?
synthesis
4. Write the reverse reaction:
2NH3 → 3H2 + N2
5. What kind of reaction is the reverse reaction?
decomposition
6. Over time the concentration of which substance(s) decreases?
7. Over time the concentration of which substance(s) increases?
NH3
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?
10. Is the forward or reverse reaction favored?
reverse

10. Equilibrium System and Stress
LeChatelier’s Principle states that if a change in conditions is imposed on a system at equilibrium, the equilibrium position will shift on a system that tends to reduce that change in conditions.
Pure solids and liquids are not affected by changes in equilibrium.

11. Changes in Concentration
Adding a substance to a system drive the system to consume the substance.
Removing a substance from a system at equilibrium drives the system to the production of the substance.
N2O4(g) NO2

12. Changes in Pressure
If the pressure is increased, the reaction will shift in the direction that produces the least molecules.
2NO(g) N2O2
2H2O O2 + 2H2
H2 + Cl2 2HCl
Shift to the right
Shift to the left
No shift

13. Changes in Temperature Exothermic Reactions
H2(g) + I2(g) HI(g) + heat
Lowering the temperature will produce a higher yield of HI
Increasing the temperature will produce a higher yield of H2 & I2

14. Changes in Temperature Endothermic Reactions
Heat + NH4Cl(s) NH3(g) + HCl
Lowering the temperature will produce a higher yield of NH4Cl
Increasing the temperature will produce a higher yield of NH3 & HCl

15. Homogeneous and Heterogeneous Equilibria
Homogeneous equilibria: reactions where reactants and products are in the same states.
Heterogeneous equilibria: reactions where the reactants and products are in different states.

16. Collision Theory
Collision Theory – Molecules must collide with each other in order to react.
Particle must collide with another particle or the container wall.
Types of collisions
Effective Collisions: Lead to the formation of products
Ineffective Collisions: Do not lead to the formation of products

17. Factors which determine effectiveness of collisions:
1) Energy of particles
2) Orientation of particles
In a chemical reaction, the kinetic energy of the reactants is converted to potential energy for the products.

18. Ineffective Collision – Insufficient Energy - No Products
Effective Collision – Sufficient Energy – Forms Products

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

20. Activation Energy (EA)
Particles must possess a minimum amount of energy in order to react.
Analogy: Rolling a ball up a hill
Activation Energy: Energy needed to start a reaction; difference between the energy at the peak and the energy of the reactants

21. Activation Energy

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

23. Endothermic
Exothermic

24. Activated Complex
Transition State: A short-lived complex which lives as neither a reactant or product. This state occurs when old bonds are broken and new bonds are formed.
This complex is called the activated complex.
Activation energy is the energy needed to form the activated complex.

25. Activated Complex (cont.)
Most reactions occur in a series of steps called a reaction mechanism.
A B
B C
C D
Often one step is slower than the others. It is called the rate determining step.
A substance that increases the rate of a chemical reaction by providing a mechanism with a lower energy of activation is called a catalyst.

26. Factors Affecting Reaction Rates
Reaction rate is determined by measuring a change in concentration of reactants and products in a certain amount of time. Reaction rates are determined experimentally.
Rate-influencing factors are those factors that affect rate of reactions by altering the frequency, orientation, or energy levels at which the particles collide.

27. Factors affecting the rate of reaction
A) Nature of Reactants
1) Structure: The more complex, the slower it reacts.

28. A) Nature of Reactants
2) State: Gases react the fastest and solids react the slowest.

29. B) Temperature: As temperature increases, the rate of the reaction increases.
Rule of Thumb: Rate doubles every 10ºC increase

30. C) Concentration: As concentration increases, the rate of the reaction increases.

31. D) Surface Area: As surface area increases, the rate of the reaction increases.

32. E) Catalyst increases the rate of the reaction without being used in the reaction.
Catalysts speed up the reaction rate by reducing the activation energy needed for the reaction to occur.
Uncatalyzed reaction (slow)
Catalyzed reaction (fast)

33. Inhibitors: slow the rate of the reaction

34. Energy diagram involving a catalyst: