Chapter 17:

1. Chapter 17: Equilibrium

2. Students will learn about… • How does a reaction occur? • What affects the reaction rate? • Activation energy • What is equilibrium? • Equilibrium constant • Heterogeneous equilibria • Le Chatelier’s principle • Solving equilibrium constant problems • Solubility equilibrium

3. Collision Model • Chemical reaction is breaking bonds and forming new bonds • Breaking bonds = endothermic • Forming bonds = exothermic • Molecules must collide in order for a reaction to occur. • The rate of collisions depends on • concentrations of reactants • kinetic energy of molecules (= temperature)

4. Factors Affecting Rate • External factors • Concentrations of reactants • Temperature • Pressure if gas substances • Surface area – smaller particles, higher the rate • Catalyst • Internal factor • Activation Energy: the energy barrier the reactants must overcome in order to become products • Orientation of molecules

5. Activation Energy (Ea) (Ex) 2BrNO  2NO + Br2 (1) The just right orientation: Br---Br (2) The formation of Br-Br bond (3) The breaking of Br-N bond *The peak of the curve is called the activated state or transition state.

6. Ea lowered by catalyst • Catalyst: speeds up a reaction by lowering Ea and is never altered during a chemrxn (Ex) enzymes

7. Equilibrium • The exact balancing of two processes, one of which is the opposite of the other. (Ex) vapor pressure: the pressure when rate of evaporation = rate of condensation • Dynamic vs. static equilibrium

8. Chemical Equilibrium • Rate of forward reaction = rate of reverse reaction • Shown with double arrows, • Signals the end of reaction (Ex) N2O4 (colorless gas) 2NO2 (brown gas)

10. Law of Chemical Equilibrium • Doesn’t apply to pure substances because their concentrations do not change • Applies only to solutions and gases • Equilibrium constant, K aA + bBcC + dD (Be sure to use concentrations @ • K is unitlessequilibrium) • The value of K is independent of concentrations as long as the temp and pressure remain the same • The expression depends on the balanced equation

11. (Ex) N2(g) + 3H2(g) 2NH3(g) At 500 °C and 1 atm, K = 0.062

12. Heterogeneous Equilibria *Equilibrium between different states of matter (Ex) Write the K expression. (a) 2KClO3(s) 2KCl(s) + 3O2(g) (b)2H2O(l) 2H2(g) + O2(g)

13. Le Châtelier’s Principle • When a system is a equilibrium is disturbed, the system adjusts itself in a way to reduce the change. The disturbances are: • Adding or removing reactants or products • Increasing or decreasing temperature • Increasing or decreasing volume • Increasing or decreasing pressure • Adding a catalyst

14. Adding or Removing A Substance • Adding a substance shifts the reaction to the direction that uses up the substance • Removing a substance shifts the reaction to the direction that produces the substance (Ex) 2NO2(g) N2O4(g) If more NO2 is added, the reaction shifts to right If N2O4 is added, the reaction shifts to left If N2O4 is removed, the reaction shifts to right

15. Temperature Changes • If the temperature is raised, the reaction shifts to the direction that removes heat • If the temperature is lowered, the reaction shifts to the direction that produces heat (Ex) N2(g) + 3H2(g) 2NH3(g) + heat Lowering the temperature shifts the reaction to right Raising the temperature shifts the reaction to left (Ex) N2O4(g)+ heat 2NO2(g) At the lowered temp, the reaction shifts to left At the higher temp, the reaction shifts to right

16. Increasing or Decreasing Volume • Has no affect on solid or liquid substances • Increasing (decreasing) volume shifts the reaction to the direction that has higher(lower) amount of gas (Ex) 2NOCl(g) 2NO(g) + Cl2(g) If volume is increased, the reaction shifts to right (3 moles of gas) If volume is decreased, the reaction shifts to left (2 moles of gas) (Ex) 4Na(s) + O2(g) 2Na2O(s) Increasing the volume shifts the reaction to left Decreasing the volume shifts the reaction to right

17. Increasing or Decreasing Pressure • Has no affect on liquids or solids • Increasing pressure has the same affect as decreasing the volume • Decreasing pressure has the same affect as increasing the volume (Ex) 2NOCl(g) 2NO(g) + Cl2(g) Increasing the pressure shifts the reaction to left (lower volume) Decreasing the pressure shifts the reaction to right (higher volume)

18. Adding a Catalyst • No affect at equilibrium • Establishes the equilibrium quicker because it lowers the activation energy

19. Example One method for the production of hydrogen gas can be described by the following endothermicreaction: CH4(g) + H2O(g) CO(g) + 3H2(g) Which of the following changes would decrease the amount of hydrogen gas (H2) produced? I. H2O(g) is added to the reaction vessel. II. The volume of the container is doubled. III. CH4(g) is removed from the reaction vessel. IV. The temperature is increased in the reaction vessel.

20. Application of K • Large K value = products dominant rxn • At equilibrium, mostly products • Small K value = reactants dominant rxn • At equilibrium, mostly reactants

21. Solving for a concentration using K At a given temperature, K = 50 for the reaction: H2(g) + I2(g) 2HI(g) Calculate the equilibrium concentration of H2 given: [I2] = 1.5 × 10–2M and [HI] = 5.0 × 10–1M

22. Solubility Equilibria • Equilibrium between an insoluble reactant and its ions MX (s) aM+(aq) + bX-(aq) • Solubility product (Ksp) • Equilibrium constant; has only one value for a given solid at a given temperature • Used to calculate the solubility of a reactant

23. (Ex) Bi2S3(s) 2Bi3+(aq) + 3S2–(aq)

24. (Ex) If a saturated solution of PbCl2 is prepared by dissolving some of the salt in distilled water and the concentration of Pb2+ is determined to be 1.6 × 10–2M, what is the value of Ksp?

25. (Ex) Calculate the solubility of silver chloride in water. Ksp = 1.6 × 10–10