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Chapter 18

Chapter 18. Chemical Equilibrium. Sect. 18-1: The Nature of Chemical Equilibrium. Reversible reaction – a chemical reaction in which the products can react to re-form the reactants

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Chapter 18

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  1. Chapter 18 Chemical Equilibrium

  2. Sect. 18-1: The Nature of Chemical Equilibrium • Reversible reaction – a chemical reaction in which the products can react to re-form the reactants • Chemical equilibrium – when a reversible reaction’s rate of its forward reaction equals the rate of its reverse reaction and the concentrations of reactants and products remain unchanged

  3. Some reactions will reach a state of equilibrium very quickly, while others require the forward reaction to be almost complete before the rate of the reverse reaction becomes high enough to reach equilibrium • Some cases can be represented by showing equal length arrows • The 2nd case can be represented by showing the arrows as different lengths, a longer arrow for the reaction that is favored

  4. Equilibrium constant (K) – the ratio of the mathematical product of the concentrations of substances formed at equilibrium to the mathematical product of the concentrations of reacting substances. Each concentration is raised to the power equal to the coefficient of that substance in the balanced equation.

  5. For a general equation: nA + mB  xC + yD • The equilibrium constant would be: K = [C]x [D]y [A]n [B]m

  6. Ex: An equilibrium mixture of N2, O2, and NO gases at 1500 K is determined to consist of 6.4x10-3 mol/L of N2, 1.4x10-3 mol/L of O2, and 1.1x10-5 mol/L of NO. What is the equilibrium constant for the system at this temperature? • N2 + O2 NO

  7. Sect 18-2: Shifting Equilibrium • Pressure only affects reactions that involve a different number of moles of gas on the reactants & products side • An increase in pressure would cause the reaction to shift in such a way to lower the pressure • Ex. N2 + 3H2 2NH3; 4 moles of gas on the left & 2 on the right, so if pressure is increased, it will shift right to lower the pressure

  8. Changes in Concentration • An increase in concentration of one reactant means a shift to the right, but when it reaches equilibrium there will be a lower concentration of the other reactant • Changes in concentration do not affect the equilibrium constant

  9. Concentration (con’t) • The concentration of solids & liquids are not changed by adding/removing the substance because concentration is density-dependent • A pure substance in a liquid or solid form can be removed from the equilibrium constant equation because it is assumed to remain constant

  10. Changes in Temperature • In reversible reactions, one direction is endothermic & the other exothermic • Addition of heat favors the endothermic reaction • Catalysts increase reaction rates equally, so they do not affect K

  11. Reactions that go to Completion • Reactions where ions form a gas • Reactions where ions form a precipitate • Reactions where ions form a product that is only slightly ionized (ex. acid/base reaction where water is a product)

  12. Common-ion Effect • Common-ion effect – when the addition of an ion common to two solutes brings about precipitation or reduced ionization • Ex. – adding HCl gas to saturated solution of NaCl increases [Cl-], causing them to combine with Na+, thus lowering [Na+] and raising [Cl-]; product of [Na+] and [Cl-] stays the same • Also when an ion of a weak electrolyte is added to a solution, it will reduce the ionization of the electrolyte

  13. Sect. 18-3: Equilibria of Acids, Bases, and Salts • Acid-ionization constant (Ka) – ratio of ion concentration to acid concentration at equilibrium • Even when additional anions are added and a new equilibrium is reached, the Ka value will be the same.

  14. Buffers • Buffered solution – a solution that resists changes in pH • Usually made up of a weak acid or weak base and a common ion • Ex. Acetic acid and sodium acetate

  15. Ionization Constant of Water • Remember that self-ionization of water is an equilibrium equation • Kw = [H3O+] [OH-]

  16. Hydrolysis of Salts • Hydrolysis – a reaction between water and ions of a dissolved salt • If anions react with water, it results in a more basic solution • HA (aq) + H2O  H3O++ A- • Ka = [H3O+] [A-] / [HA] • Hydrolysis reaction: A- + H2O  HA + OH- • The lower the Ka value of HA, the stronger the A- attraction for protons & the greater the formation of OH-

  17. If cations react with water, it results in a more acidic solution • B + H2O  BH+ + OH- • Kb = [BH+] [OH-] / [B] • Hydrolysis reaction: BH+ + H2O  H3O++ B • The lower the Kb value of B, the stronger the donation of protons the BH+ will have and the great the production of H3O+

  18. Hydrolysis in Acid-Base Reactions • Salts of strong acid/strong base produce neutral solutions • Salts of weak acid/strong base produce basic solutions • Salts of strong acid/weak base produce acidic solutions • Salts of weak acid/weak base can produce acidic, basic, or neutral solutions depending on what the salts are

  19. Sect. 18-4: Solubility Equilibrium • Solubility product constant (Ksp) – product of the molar concentrations of a substance’s ions in a saturated solution, each raised to the power of its coefficient • Used with sparingly soluble salts • Ex. AgCl (s)  Ag+ (aq) + Cl- (aq) • Ksp = [Ag+][Cl-], remember the solid AgCl is not included because its concentration is assumed to be constant

  20. Numerical values of Ksp can be calculated using solubility data from pg. 901 • Ex. CaF2 has a solubility of 1.7x10-3g/100g water and this must be converted to mol/L, which is 2.2x10-4 mol/L • When it ionizes there are 1 Ca and 2 F ions, so their concentrations are 2.2x10-4 mol/L and 2(2.2x10-4) mol/L • These values are then put into the Ksp equation

  21. Calculating Solubilities • If the Ksp is known, you can write the equation, put in the known value for Ksp, and then solve for the concentration of the ions

  22. Precipitation Calculations • If a calculated value for Ksp for a solution is greater than the known value for Ksp for a sparingly soluble salt, then precipitation will occur

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