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Chemical Equilibrium

Learn the concept of chemical equilibrium, how it differs from regular chemical reactions, equilibrium expressions, common examples, calculating equilibrium constant, visualizing equilibrium, using Kc to find concentrations, RICE tables, and Le Châtelier's Principle.

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Chemical Equilibrium

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  1. Chemical Equilibrium The state where the concentrations of all reactants and products remain constant with time

  2. How is “equilibrium” different than chemical reactions we have studied so far? • For stoichiometry calculations, we assume reaction goes to completion. • In calling a compound “soluble,” we assume that ions remain dissociated in water: NaCl + H2O  Na+(aq) + Cl-(aq) • In a closed vessel, a chemical reaction achieves a state of equilibrium, where concentrations of both reactants and products remain constant over time.

  3. Equilibrium = Dynamic • Reactants convert continually to products AND • Products continually revert to reactants • Due to molecular collisions • Forward and reverse RATES are equal

  4. The Equilibrium Expression • Given the following reaction at equilibrium: • 2A(g) + 3B(g)  2C(g) + 4D(g) • This is the equilibrium expression: • Keq = [C]2[D]4 [A]2[B]3 • Keq is the equilibrium constant. • Include concentrations for products or reactants ONLY for gases and aqueous solutions. • Liquids and solids are not included because their concentrations do not change.

  5. Most Familiar Equilibrium Example: Water • Water molecules split other water molecules into hydrogen ions and hydroxide ions to a small extent. • This is called auto-ionization, or self-ionization of water. • This reaction is reversible and does not go to completion in the forward direction. It reverses itself after only a few ions are formed. • The equilibrium expression for this equation: Kw = [H3O+][OH-]

  6. Most Familiar Equilibrium Example: Water H2O   H3O+(aq) + OH- (aq) • Kw = 1.0 x 10-14 [OH-][H3O+] = 1.0 x 10-14 * • Since the K value is so small, this indicates that not very many ions form before the equation reverses itself and begins to form the molecules of water again. • In pure water and in aqueous solutions, there are both H+ and OH- • ACIDS: [H3O+] > [OH- ] • BASES: [OH-] > [H3O+] . *Includes concentrations for products or reactants ONLY for gases and aqueous solutions

  7. Equilibrium Positions • At a given temp, there is only one value of K but an infinite number of possible equilibrium positions. • Eq. position defined by the concentrations that satisfy the equilibrium expression. • Equilibrium position • Depends on initial concentrations (aqueous or gaseous mixtures) • Never depends on amount of pure solid or liquid in the system

  8. Problem: Calculating K At 127º C, [NH3] = 3.1 x 10-2 M [N2] = 8.5 x 10-1 M [H2] = 3.1 x 10-3 M Find K for the Haber process. N2(g) + 3H2(g)   2NH3 (g) K = [NH3]2 = (3.1 x 10-2 M)2 [N2][H2]3 (8.5 x 10-1 M)(3.1 x 10-3 M)3 K = 3.8 x 104 (no units)

  9. Problem: Calculating K for Reverse Reaction At 127º C, [NH3] = 3.1 x 10-2 M [N2] = 8.5 x 10-1 M [H2] = 3.1 x 10-3 M Find K for 2NH3 (g)  N2(g) + 3H2(g) K’ = [N2][H2]3=(8.5 x 10-1 M)(3.1 x 10-3 M)3 [NH3]2 (3.1 x 10-2 M)2 K’ = 2.6 x 10-5 (no units)

  10. Visualizing Equilibrium • Try this activity to see effects of changing concentrations on the equilibrium of a system: http://cheminfo.chem.ou.edu/~mra/CCLI2004/ERGBN.htm

  11. Important Points about K • For a balanced equation, Kc: • is constant at a given temperature • changes if the temperature changes • does not depend on initial concentrations • Kc’ is the equilibrium constant for the equation written in reverse, and is the reciprocal of Kc

  12. Reaction Quotient, Q • Is a measure of reaction progess • Uses same form as Kc • Also called mass action expression • Concentrations not necessarily at equilibrium

  13. Application: Using Kc to find concentrations PCl3(g)+ Cl2 (g) PCl5 (g) At a given temp, this reaction occurs in a 1.0-L container with 0.25 mol PCl5 and 0.16 mol PCl3. At equilibrium, what is the concentration of Cl2? Kc = [PCl5] = 0.25 =1.9 [Cl2][PCl3] [Cl2]0.16 [Cl2]=0.25 = 0.82 M (0.16)(1.9)

  14. RICE Table • A more common application gives initial concentrations and Kc and you must find the equilibrium concentrations. • Make a table showing the reaction, initial concentrations, change and equilibrium concentrations to solve this problem. • You will often need to use the quadratic formula to solve these, so refresh your memory on your calculator.

  15. If Kc = 49.0 at a given temp, and 0.400 mol of each reactant are placed in a 2.0 L container, what concentrations of all species are present at equilibrium for A + B   C + D? [A]o = [B]o = 0.400 mol/2L = 0.200 M [C]0 = [D]0 = 0 M

  16. KC = 49.0 = (x)(x) (0.200 –x) (0.200 –x)

  17. A rare occasion without quadratic formula Kc = 49.0 = (x)(x) (0.200 –x) (0.200 –x) 7.00 = x/(0.200-x) x = 1.40 – 7.00 x 8.00x = 1.40 X = 0.175 [A] = [B] = 0.200 – 0.175 = 0.025 M [C] = [D] = 0.175 M

  18. LeChâtelier’s Principle When a system at equilibrium is disturbed by application of stress, it attains a new equilibrium position that minimizes the stress. • "If stress is applied to a system at equilibrium,the system will tend to readjust so that the stress is reduced."

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