International Baccalaureate Chemistry

1. International Baccalaureate Chemistry Topic 7 – Chemical Equilibrium

2. Equilibrium • In general, equilibrium is the state in which the rate of the forward process/reaction equals the rate of the reverse process/reaction.

3. Physical Equilibria • Liquid - Vapor Equilibria • The following equation represents Bromine liquid with Bromine gas in equilibrium. Br(l)→ Br(g) • At equilibrium, the rate of evaporation equals the rate of condensation. • NOTE: This does not mean that the number of molecules in each state is the same!

4. Physical Equilibria • Solute – Solution Equilibria • The following equation represents solid sodium chloride in equilibrium with its ions in a saturated solution: NaCl(s) ↔ Na+(aq) + Cl-(aq) • At equilibrium, the rate of dissolving equals the rate of crystalization. • NOTE: This does not mean that the number of molecules in each state is the same!

5. What Determines the Equilibrium Point of a Substance? • Nature of the reactants • Type of Liquid (in liquid – vapor equilibria) • Type of Solid (in solute – solution equilibria) • Temperature • In liquid – vapor equilibriamore particles will be found in the gas phase with an increase in temperature. • In a solute-solution equilibria, more particles will be found in the dissolved phase with an increase in temperature. In a physical equilibrium, the system must be closed!

6. Reversible Reactions • Many chemical reactions are reversible and never go to completion. • Consider the following reaction: (Haber Process) N2 + 3H2↔ 2NH3 • If a certain quantity of each reactant is placed in a closed container, initially, N2 and H2 are the only species present in the container. • N2 and H2 will begin to react at their maximum rate because their concentrations are at a maximum. Forward Reaction → maximum rate Reverse Reaction → no rate

7. N2 + H2 NH3 • As time passes, the forward reaction will start to decrease because [N2] and [H2] are decreasing (used up). • As this happens, an opposing change begins to occur. • NH3 begins to decompose into N2 and H2 through the reverse reaction. • The rate of the reverse reaction will steadily increase (as the [NH3] is increasing). • Eventually, the forward reaction rate becomes equal to the reverse reaction rate. • The system is said to be in a state of Dynamic Equilibrium.

8. Chemical Equilibrium

9. Characteristics of a System in a State of Equilibrium • Equilibrium can be approached from both directions. • The rate of the forward reaction equals the rate of the reverse reaction and the concentrations of all reactants and products remains constant (NOT EQUAL!) • The system must be closed. • Macroscopic properties remain constant (although microscopically there is lots going on!)

10. The Equilibrium Constant Expression • For all chemicals, Kc = [Products] [Reactants] Equilibrium Constant Equilibrium Expression For the reaction aA + bB ↔ cC + dD, The Equilibrium Constant expression is: Kc = [Products] = [C]c [D]d [Reactants] [A]a [B]b Where: a,b,c,d are the coefficients from the chemical equation

11. Characteristics • The coefficients from the equation become a power in the equilibrium expression. • Only gasses and aqueous solutions are involved in the equilibrium expression since solids and liquids cannot be expressed as concentrations. • NOTE: • Solids cannot be compressed so their density cannot be changed. • Therefore, the molar concentrations cannot be changed. • Liquids cannot be compressed either so that they have a constant density and molar concentration. • If there is another liquid present that can dilute the first liquid, then the liquid is not “pure” and can have its concentration changed by dilution.

12. Practice: Write Kc Expressions for the Following Reactions • N2(g) + 3H2(g)↔ 2NH3(g) • Cu(s) + 2AgNO3(aq) ↔ 2Ag(s) + Cu(NO3)2(aq) • P4(s) + 5O2(g) ↔ P4O10(s) • Br2(l) + H2(g) ↔ 2HBr(g) • CH3COCH3(l) + Cl2(g) ↔ CH3COCH2Cl(l) + HCl(g) 1. Kc = [NH3]2 2. Kc = [Cu(NO3)2] 3. Kc = 1 [N2] [H2]3 [AgNO3]2 [O2]5 4. Kc = [HBr]2 5. Kc = [HCl] [H2] [Cl2]

13. Kc Values • Whenever Kc is large (Kc > > 1): • Products are favored and the reaction goes almost to completion. • Kc = [Products] Kc = [BIG #] > > 1 [Reactants] [small #] • Whenever Kc is small (Kc < < 1): • Reactants are favored and the reaction hardly proceeds. • Kc = [Products]Kc = [small #] < < 1 [Reactants] [BIG #] • Note: The value of Kcis affected only by temperature.

14. Changes in Equilibrium • Changes will occur in a closed system if it has been disturbed or stressed. • To explain these changes, we use Le Chatelier’s Principle which states: • “Whenever a stress is applied to a system at equilibrium, it will shift to relieve that stress applied.”

15. Types of StressesThat Can Be Placed On a System at Equilibrium • Change in Concentration • Change in Temperature • Change in Volume or Pressure • Addition of a Catalyst Affect Equil. Position Does Not Affect Equil. Position (does not cause a shift)

16. Change in Concentration • Only affects substances that are (g) or (aq) • Or if two liquids are present (l) • Example 1: Fe3+(aq) + SCN-(aq) ↔ FeSCN2+(aq) • If we add more [Fe3+] to the reaction mixture, the reaction will shift forward to use up what is added. • The result will result in a decrease of [SCN-] and an increase in [FeSCN2+] product produced. • Conclusion: Increasing the concentration of a reactant favors the forward reaction. • Eventually, a new equilibrium point will be reached and no further change in concentration will occur. • The same rules apply if a concentration is decreased. • To offset the stress, more of what is removed must be produced.

17. Change in Concentration • Example 2: N2(g)+ H2(g) ↔ NH3(g) • If the [NH3] is decreased, the system will shift forward in an attempt to replace ammonia that was removed. • As a result, N2 and H2 decrease. • Example 3: Identify the shift (forward or reverse) that occurs for the following reaction and indicate the results of the shift and identify the remaining reactants. 2A(s) + 3B(g) ↔ 2E(g) + 4D(g) • [B] is increased • [E] is increased • Amount of A is increased • [D] is decreased

18. Changes in Temperature • At equilibrium, the temperature is constant. • If the reaction vessel is cooled, the reaction will shift to produce heat. • If the reaction vessel is heated, the reaction will shift to use up heat.

19. Changes in Temperature • Example 1: Endothermic Reaction • Co(H2O)62+(aq) + 4Cl- + energy ↔ CoCl42-(aq) + 6H2O(l) • If we disturb the system by increasing the temperature, both the forward and reverse reaction rates increase, however, in this reaction the forward will increase more than the reverse. Why? • Forward reaction – uses heat (cools) • Reverse reaction – gives off heat (warms) • Results: • [Co(H2O)62+] and [Cl-] Decrease • [CoCl42-] increases • The moles of water will also increase

20. Changes in Temperature • Example 2: Exothermic Reaction • 3H2(g) + N2(g)↔ 2NH3(g) + 91 kJ • If we disturb the system by increasing the temperature, both the forward and reverse reaction rates increase, however, in this reaction the reverse reaction will increase more than the forward. • Forward reaction – gives off heat (warms) • Reverse reaction – uses heat (cools) • Results • [H2] and [N2] increase • [NH3] Decreases

21. Changes in Temperature • Generally, a temperature increase will • favor endothermic reactions • Not favor exothermic reactions.

22. Change in Volume or Pressure • Only gasses are affected by pressure/volume changes. • Recall: • If the volume of a container is decreased (pressure increased), all concentrations increase. • This will increase both the forward and reverse reactions, but the reaction will try to offset the increase in pressure . • To decrease the pressure, the reaction will favor the reaction which produces fewer gas particles.

23. Change in Volume or Pressure • Example 1: • 3H2(g) + N2(g) ↔ 2NH3(g) + 91 kJ • If the volume of the container is decreased, pressure is increased. • 4 moles of gas (3H2 and 1N2) vs. 2 moles of gas (2NH3) • Results • Forward reaction is favored since the reaction is producing fewer gas molecules and alleviates the pressure. • [H2] and [N2] decrease • [NH3] increase • What happens to the same reaction if the volume is increased?

24. Change in Volume or Pressure • Example 2: • H2(g) + I2(g)↔ 2HI(g) • If you were to either increase or decrease the pressure of the system, it would not make a difference in the reaction rate as it cannot get rid of stress in either direction (2 moles of gas on either side) • Example 3: • NH3(g) + HCl(g) ↔ NH4Cl(g) • If the pressure is decreased, which reaction is favored and what are the results?

25. Change in Volume or Pressure • Example 4: • 2A(g) + B(g) ↔ 3D(s) + C(g) • Stress: Increase the pressure by decreasing the volume: • Reaction favored: _______________ • [A] will ________________________ • [B] will ________________________ • [C] will ________________________ • Amount of [D] will _______________

26. Addition of a Catalyst • A catalyst will increase both forward and reverse rate equally. • Concentrations of all substances remain constant. • Catalysts do not affect the position of equilibrium.

27. Summary

28. Haber Process/Contact Process • Using the resources available to you, describe and explain the application of equilibrium and kinetics concepts to the Haber process and Contact process. • Start with the IB text and supplement with other texts and/or internet to make a complete set of notes.