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Equilibrium

Equilibrium. AP Chem Mr. Nelson. The Concept of Equilibrium. The term reversible reaction is used to describe reactions that can go in either the forward or the reverse direction Equilibrium is used to specify that a reversible reaction has reached an equal rate for both directions.

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Equilibrium

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  1. Equilibrium AP Chem Mr. Nelson

  2. The Concept of Equilibrium • The term reversible reaction is used to describe reactions that can go in either the forward or the reverse direction • Equilibrium is used to specify that a reversible reaction has reached an equal rate for both directions

  3. The Concept of Equilibrium • Chemical equilibrium occurs when opposing reactions are proceeding at equal rates • Concentrations of the reactants/products remains constant

  4. The Concept of Equilibrium • Its important to note that forward and reverse reactions still occur, although outwardly no apparent change occurs

  5. The Concept of Equilibrium • Note that the equilibrium concentrations are independent of initial concentrations • A mathematical relationship exists between concentration of the reactants and products once equilibrium has been reached

  6. The Concept of Equilibrium • Fish Tank Example

  7. Connection between Equilibrium & Kinetics • For generalized reaction: A⇋B • Forward reaction: AB • Forward Rate Law: Ratef= kf[A] • Reverse reaction: BA • Reverse Rate Law: Rater = kr[B]

  8. Connection between Equilibrium & Kinetics • By definition, at equilibrium forward and reverse rates must be equal, so: Ratef = Rater kf[A] = kr[B] • Also, concentrations remain constant,so:

  9. The Equilibrium Constant • The equilibrium constant (Keq) can be determined at constant temperature for any reaction at equilibrium: aA+bBcC+ dD [C]c [D]d[A]a [B]b Keq= Keqis unitless

  10. The Equilibrium Constant • Note that onlyequilibriumconcentrations can be placed into an equilibrium constant expression! • Also, no solids or liquids are included in an equilibrium constant expression • The reason for this is their concentrations remain fairly constant in equilibrium • Note that they are still required for the reaction to proceed however

  11. The Equilibrium Constant

  12. Concentration vs. Pressure Equilibria • Equilibrium expressions can be written in terms of concentration (Kc) measured in molarity or pressure (Kp) measured in atm • Example: or

  13. Concentration vs. Pressure Equilibria • In general, Kc is not equal to Kp • The following equation converts between the two: n = moles of gas products – moles of gas reactants • Note: When no change in moles, Kp=Kc

  14. Homogeneous vs. Heterogeneous Equilibria • The term homogeneous equilibrium is used to describe reactions where all the species are in the same phase • Example: • The term heterogeneous equilibrium is used to describe reactions where species are in different phases • Example: H2CO3 (aq) + H2O (l) HCO3-(aq) + H3O+(aq)

  15. Multiple Equilibria • When the products of one reaction are the reactants of another • Example: A + B  C + D C + D  E + F K1 K2 A + B  E + F Kc

  16. Multiple Equilibria • Using their equilibrium expressions, we can derive a mathematical relationship:

  17. Manipulating Equilibria Expressions • If an equilibrium is written in the reverse, the following relationship exists: PCl3 + Cl2 PCl5 PCl5  PCl3 + Cl2

  18. Manipulating Equilibria Expressions • If an equilibrium is multiplied by a factor, the following relationship exists: PCl3 + Cl2 PCl5 2PCl3 + 2Cl2 2PCl5

  19. The Magnititude of Equilibrium Constants • What does Keq tell us? • If Keq is very large, products will be favored • This implies the numerator is large in the Keq expression compared to denominator • If Keq is very small, reactants will be favored • This implies the denominator is large in the Keq expression compared to numerator • If Keq is close to 1, then roughly equal amounts of reactants and products are present at equilibrium

  20. The Reaction Quotient (Q) • Recall: In electrochemistry, we utilized “Q” which is similar in all respects to “Keq” except in definition • Q is utilized when given initial concentrations, not concentrations at equilibrium

  21. Predicting the Direction of Reversible Reactions • If given an initial set of conditions for a reversible reaction, utilizing Q and Keq, a prediction of the direction of the reaction (towards reactants/towards products) can be made

  22. Comparing Kc and Qc

  23. Comparison of Kc and Qc Qc Kc Qc Kc Kc Qc

  24. Calculating Equilibrium Concentrations • Tabulate all known initial and equilibrium concentrations • When both are known, calculate the change • Using coefficients, calculate changes in concentration for all participants in a rxn • Calculate equilibrium concentrations

  25. The ICE Method • This is the process used for determining equilibrium concentrations Initial molarity Change in molarity Equilibrium molarity

  26. Example #1: Calculating Equilibrium Concentrations A mixture of 0.500 mol H2 and 0.500 mol I2 was placed in a 1.00 L stainless-steel flask at 430 °C. The equilibrium constant Kc for the reaction: H2 (g) + I2 (g)  2 HI (g) is 54.3 at this temp. Calculate the concentrations of H2, I2, and HI at equilibrium.

  27. H2 (g) + I2 (g)  2 HI (g)

  28. Substituting, we get: • Taking the square root of both sides: x = 0.393 M

  29. So, the concentrations at equilibrium would be the following: • [H2] = (0.500 – 0.393) M = 0.107 M • [I2] = (0.500 – 0.393) M = 0.107 M • [HI] = 2 x 0.393 M = 0.786 M

  30. Example #2: Calculating Equilibrium Concentrations • When 3.0 mol of I2 and 4.0 mol of Br2 are placed in a 2.0 L reactor at 150oC, the following reaction occurs until equilibrium is reached: I2(g) + Br2(g) 2IBr(g) Chemical analysis then shows that the reactor contains 3.2 mol of IBr. What is the value of the equilibrium constant Kc for the reaction?

  31. I2 (g) + Br2 (g)  2 IBr (g)

  32. When initial and final concentrations are known, calculate the change: • 0 + 2x = 1.6 • x = 0.8 • Then determine all equilibrium concentrations: • [I2] = 1.5 – 0.8 = 0.7 M • [Br2] = 2 – 0.8 = 1.2 M • [IBr] = 1.6 M

  33. Plug equilibrium concentrations into equation and solve:

  34. Calculating Equilibrium Concentrations when Keq is small • When solving equilibrium concentrations and Keq is small, assume the effect of the change (“x”) is negligible (for rcts only)

  35. Example: Calculating Equilibrium Concentrations when Keq is small • At 100 °C the equilibrium constant (Kc) for the reaction: COCl2 (g)  2 CO (g) + Cl2 (g) is 2.19 x 10-10. If the initial concentration of [COCl2] = 1.5 M, calculate the concentrations of all species at equilibrium.

  36. COCl2 (g)  2 CO (g) + Cl2 (g)

  37. Le Chatelier’s Principle • If a stress is applied to a system in equilibrium, the equilibrium shifts to relieve that stress • Forms of “stress”: • Concentration • Pressure • Temperature

  38. Change in Reactant or Product Concentration • Adding a substance shifts a system at equilibrium away from the added substance • System must consume added item • Removing a substance from a system at equilibrium shifts the reaction toward the removed substance • System must replace removed item

  39. Changes in Pressure • Increase in pressure forces an equilibrium to shift in the direction that reduces the number of moles of gas in the system • Example: • An increase in pressure causes a shift to the left N2O4 (g)  2 NO2 (g)

  40. Changes in Pressure • Recall: Changes in volume can also affect changes in pressure • Increase volume = decrease pressure • Decrease volume = increase pressure • Also the addition of an inert gas (one not in the reaction) can cause an increase in pressure as well

  41. Changes in Temperature • Enthalpy of the reaction plays a major roll in how heat effects the equilibrium • Recall: Endothermic (+∆H) means heat is a reactant, while exothermic (- ∆H) means it is a product • Treat changes in heat in the same manner that you treat changes in concentration

  42. Effects of Catalysts on Equilibrium • Recall: Catalysts lower energy barrier between reactants and products • Ea for both forward and reverse are lowered to the same extent • Catalysts increase the rate at which equilibrium is reached • However, it does NOT change the composition of the equilibrium mixture.

  43. Solubility Equilibria • Saturated solutions of salts are another type of chemical equilibria • Slightly soluble salts establish a dynamic equilibrium with the hydrated cations and anions in solution

  44. The Solubility Product, Ksp • The equilibrium constant, the Ksp, is no more than the product of the ions in solution (Recall: solids do not appear) • For a saturated solution of AgCl, the equation would be: • The solubility product expression would be:

  45. Sequence of Steps for Solubility Equilibria

  46. Solubility • Solubility can be expressed two ways: • Molar solubilityis the moles of solute per 1L of saturated solution (mol/L) • Solubility is the number of grams of solute per 1 L of saturated solution (g/L) • Ksp expressions require molar solubility

  47. Determining Ksp by Experimental Measurements • Use ICE table, but with “s” for change (instead of “x”) • “s” is not only the change, but the molar solubility

  48. Determining Ksp: Example #1 The solubility of CaSO4 is found to be 0.67 g/L. Calculate the value of Ksp for calcium sulfate.

  49. Determining Ksp: Example #1

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