Chemical Equilibrium Lecture Notes PDF

Chemical equilibrium Equilibrium in electrolytes  Electrolytes Electrolyte: a substance that gives an electrically conducting solution when dissolved in water Other substances are called nonelectrolytes. A strong electrolyte exists in solution almost entirely as ions. HC...

Chemical equilibrium Equilibrium in electrolytes  Electrolytes Electrolyte: a substance that gives an electrically conducting solution when dissolved in water Other substances are called nonelectrolytes. A strong electrolyte exists in solution almost entirely as ions. HCl + H2O  H3O+ + Cl A weak electrolyte dissolves in water to give an equilibrium between a molecular substance and a small amount of ions. NH3 + H2O NH4+ + OH The most common electrolytes Acids and bases Salts Water  Acid-base concepts 1. Bronsted-Lowry concept 2. Lewis concept  Bronsted-Lowry concept of acids and bases An acid is a species that donates a proton in a proton-transfer reaction. A base is a species that accepts a proton in a proton-transfer reaction. Conjugate acid-base pairs: HCl and Cl B2 A2 H2O and H3O+ HCl + H2O H3O+ + Cl Conjugate acid: in a conjugate acid-base A1 B1 pair, the species that can donate a proton. Conjugate base: in a conjugate acid-base pair, the species that can accept a proton. Not restricted to aqueous solutions. Amphoteric compound (ampholyte): a species that can act either as an acid or a base.  Acid-base reactions according to the Bronsted-Lowry concept Every chemical reaction can theoretically be in acid1 + base2 base1 + acid2 equilibrium. Every reaction is reversible. HNO3 + H2O NO3 + H3O+ H2O is amphoteric. H2O + NH3 OH + NH4+ CH3COOH + H2O CH3COO + H3O+ HClO4 + HNO3  ClO4 + H2NO3 + Acidity - basicity depends on the partner. H2O + CH3COO OH + CH3COOH The acidic or basic nature NH4+ + H2O NH3 + H3O+ of salts can be explained.  Lewis concept of acids and bases I. An acid is a species that can form a covalent bond by accepting an electron pair from another species. A base is a species that can form a covalent bond by donating an electron pair to another species. BF3 + NH3 F3B:NH3 F electron F H pair N F B N H B F F H H H F H Lewis Lewis acid base Coordinate (dative) covalent bond forms Lewis concept of acids and bases II. Coordinate (dative) covalent bond Lewis concept – complex ions The formation of complex ions is a Lewis acid-base reaction. Cu2+ + 4 :NH3 [Cu(NH3)4]2+  Molecular structure and acid strength I. Two factors are important in determining relative acid strengths. 1. The polarity of the bond to which the H atom is attached. +  HX The more polarized the bond is, the more easily the proton is removed. Inductive effect. O Other examples: HOS=O < HOS=O HNO2 < HNO3 HClO < HClO2 < HClO3 < HClO4 OH OH HOI < HOBr < HOCl Polarity of the bond increases Acid strength increases  Molecular structure and acid strength II. 2. The strength of the bond: how tightly the proton is held. +  HX This depends on the size of the X atom. The larger the X atom, the weaker is the bond and the greater the acid strength. HF < HCl < HBr < HI Size of the X atom increases Strength of the bond decreases Acid strength increases End of MTO1 Answer the questions  Equilibrium in electrolytes Self-ionization of water Acid and base ionization equilibrium Equilibrium in salt solutions - solubility  Equilibrium in electrolytes Self-ionization of water Acid and base ionization equilibrium Equilibrium in salt solutions - solubility  Self-ionization of water Pure water has a very small conductivity. This results from self-ionization (or autoionization), a reaction in which two like molecules react to give ions. H2O(l) + H2O(l) H3O+(aq) + OH(aq) hydronium hydroxide ion ion [H3O+][OH] Kc = [H2O] = 56 M ~ constant [H2 O]2 Kc [H2O]2 = [H3O+][OH] constant = Kw  Ion-product constant for water Kw = [H3O+][OH] Ion-product constant for water Kw = 1.0  1014 (at 25ºC) In a neutral solution, [H3O+] = [OH] [H3O+] = 1.0  107 M In an acidic solution, [H3O+] > [OH] [H3O+] > 1.0  107 M In a basic solution, [H3O+] < [OH] [H3O+] < 1.0  107 M  The pH of a solution pH = log[H3O+] In a neutral solution, [H3O+] = 1.0  107 M pH = 7 In an acidic solution, [H3O+] > 1.0  107 M pH < 7 In a basic solution, [H3O+] < 1.0  107 M pH > 7 [H3O+][OH] = 1014 log(ab) = loga+logb log [H3O+]  log[OH] = 14 pH + pOH = 14 where pOH =  log[OH] The pH scale  Equilibrium in electrolytes Self-ionization of water Acid and base ionization equilibrium Equilibrium in salt solutions - solubility  Acid and base ionization equilibrium Weak acids and bases dissociate or ionize to a small extent. Equilibrium Weak acids HA + H2O H3O+ + A [H3O+][A] Kc = [H2O] ~ constant [HA][H2O] [H3O+][A] Kc[H2O] = [HA] constant = Ka  Acid ionization (or dissociation) constant [H3O+][A] Ka = acid ionization (or dissociation) constant [HA] pKa = logKa The bigger the Ka, the stronger the acid; the smaller the pKa, the stronger the acid. An example: HF (hydrofluoric acid): Ka = 6.8  10-4 pKa = 3.17 HClO (hypochlorous acid): Ka = 2.9  10-8 pKa = 7.54 HF is stronger acid than HClO  Polyprotic weak acids Polyprotic weak acids contain multiple acidic protons that can sequentially dissociate from the compound with unique acid dissociation constants for each proton. Examples: H2CO3 (carbonic acid) is a diprotic, and H3PO4 (phosphoric acid) is a triprotic weak acid. Dissociation (ionization) of phosphoric acid: H3PO4 H+ + H2PO4 Ka1 = 6.9  103 H2PO4 H+ + HPO42 Ka2 = 6.2  108 Ka1 > Ka2 > Ka3 HPO42 H+ + PO43 Ka3 = 4.8  1013  Base ionization (or dissociation) constant Weak bases B + H2 O HB+ + OH [HB+][OH] Kc = [H2O] ~ constant [B][H2O] [HB+][OH] Kc[H2O] = [B] constant = Kb [HB+][OH] Kb = base ionization (or dissociation) constant [B] pKb = logKb  Acid and base ionization constants of conjugate acid- base pairs Ka  Kb = Kw Ka  Kb = 1014 (at 25ºC) pKa + pKb = 14 (at 25ºC) An example: Ka for HF: 6.8  10-4 pKa for HF: 3.17 Kb for F- : 1.5  10-11 pKb for F- : 10.83 Tips for calculating pH  Weak acids HA + H2O H3O+ + A 1 : 1  [A] = [H3O+] [H3O+][A] [H3O+]2 [HA]e = c-[H3O+] Ka = = [HA] c-[H3O+] (c: initial concentration of the acid) [H3 O +] 2 [H3O+]2 Ka = If c/Ka > 103: Ka = c-[H3 O+] c-[H3O+] negligible [H3O+]2 Ka = [H3O+] =  Ka  c c  Weak bases B + H2O BH+ + OH 1 : 1  [BH+] = [OH] [BH+][OH] [OH]2 [B]e = c-[OH] Kb = = [B] c-[OH] (c: initial concentration of the base) [OH]2 [OH]2 Kb = If c/Kb > 103: Kb = c-[OH] c-[OH] negligible [OH]2 Kb = [OH] =  Kb  c c  Equilibrium in electrolytes Self-ionization of water Acid and base ionization equilibrium Equilibrium in salt solutions - solubility  Solubility of electrolytes Strong electrolyte NaCl Na+ + Cl Weak electrolyte AgCl Ag+ + Cl (Slightly soluble or [Ag+][Cl] Kc = [AgCl] ~ constant nearly insoluble salt.) [AgCl] Kc[AgCl] = [Ag+][Cl] constant = Ksp solubility product constant Ksp(AgCl) = [Ag+][Cl] PbI2 Pb2+ + 2 I Ksp(PbI2) = [Pb2+][I]2 Password Quiz