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Solutions of a Weak Acid or Base Acid-Ionization Equilibria Acid-Base Equilibria Chapter 16.1 and 16.3 16.6 Essentials of General Chemistry, 2nd Ed., Ebbing and Gammon Acid-ionization Equilibria Consider an equilibrium involving a weak acid in H2O: HA (aq) + H2O (l) H3O+ (aq) + A- (aq) Acid-Ionization Constant, Ka (acid-dissociation constant) - equilibrium constant for the ionization of a weak acid Ka = [H3O+][A-] [H2O] Acid dissociation - acid reacts with water to produce hydronium ion and the conjugate base ion Consider acetic acid HC2H3O2 (aq) + H2O (l) H3O+ + C2H3O2- (aq) - since acetic acid is weak electrolyte, it ionizes only to small amount (about 1% or less/partially ionized) Note: this is not the same thing as a dilute sol n of a strong acid (100% dissociated) AcidIonization Constants @ 25oC Table 16.1 AcidIonization Constants @ 25oC Calculations of and using Ka Experimental determination of Ka - using pH of a solution of a weak acid - degree of ionization - percent ionization Calculations with Ka - calculate the equilibrium concentrations of HA, Aand H3O+ for solutions of different molarities - use simplifying assumption or quadratic formula Table 16.1 1 Percent Ionization (Degree of ionization) Use of Percent Ionization Degree of Ionization - the concentration of the acid that dissolves divided by the initial concentration of the acid Consider HA (aq) + H2O (l) H3O+ (aq) + A- (aq) - degree of ionization of weak acid depends of both Ka and the concentration of the acid solution i.) for given concentration - the larger the Ka, the greater the degree of ionization ii.) for a given value of Ka - the more dilute the solution, the greater the degree of ionization How do you know when to use the simplifying assumption? - it can be shown that this simplifying assumption gives an error of less than 5% if the concentration of acid, Ca, divided by Ka equals 100 of more Degree of Ionization = [HA]dissociated = [H3O+]eq [HA]initial [HA]initial Percent Ionization = degree of ionization x 100% Base-Ionization Equilibria Consider the reaction NH3 (aq) + H2O (l) Base-Ionization Constants @ 25oC NH4+ (aq) + OH- (aq) Kb = [NH4+][OH-] [NH3] Base-Ionization Constant, Kb (base-dissociation constant - equilibrium constant for the ionization of a weak acid B (aq) + H2O (l) HB+ (aq) + OH- (aq) Kb = [BH+][OH-] [B] amines - derivatives of ammonia in which one or more hydrogen atoms are replaced by another group CH3NH2 (aq) + H2O (l) CH3NH3+ (aq) + OH- (aq) Acid-Base Properties of Salt Solutions Salt - formed when an acid neutralizes a base - salt solutions can be neutral, basic or acidic (with acidity or basicity dependent upon individual ions of the salt) Eg. NaCN - gives Na+ and CN- ions - Na+ is unreactive with water (neutral) - CN- reacts with water CN (aq) + H2O (l) HCN (aq) + OH- (aq) base acid (Kb) to form OH- (basic solution) 2 Hydrolysis (of an ion) - the reaction of an ion with water to produce the conjugate acid and hydroxide ion or the conjugate base and hydronium ion Eg. NH4Cl - gives NH4+ and Cl- ions - Cl- is unreactive with water (neutral) - NH4+ reacts with water NH4+ (aq) + H2O (l) NH3 (aq) + H3O+ (aq) acid base (Ka) Predicting the Acidity/Basicity of a Salt Consider hydrolysis of CN- to produce HCN - here, HCN is a weak acid and CN- acts as a base CN- (aq) + H2O (l) HCN (aq) + OH- (aq) generally: - the anions of weak acids are basic - the anions of strong acid have hardly any basic character eg. HCl Cl- (aq) + H2O (l) no reaction to form H3O+ (acidic solution) Consider NH4+ coming from NH3 - a cation conjugate of a weak base - NH4+ behaves as an acid with NH3 acting as a base NH4+ (aq) + H2O (l) NH3 (aq) + H3O+ (aq) generally: - the cations of weak bases are acidic - the cations of strong bases (metal ions of Groups IA and IIA elements except Be) have hardly any acidic character eg. NaOH Na+ (aq) + H2O (l) no reaction Predicting the Acidity/Basicity of a Salt These rules apply to normal salts (those in which the anion has no acidic hydrogen) 1.) A salt of a strong base and a strong acid. - the salt has no hydrolyzable ions and so gives a neutral aqueous solution - eg. NaCl 2.) A salt of a strong base and a weak acid. - the anion of the salt is the conjugate of the weak acid. It hydrolyzes to give a basic solution - eg. NaCN - aqueous metal ions other than the cations of strong bases usually hydrolyze by acting as acids eg. hydrated aluminum ion, Al(H2O)63+ (aq) - e- are drawn away from O atoms by Al atom so O atoms draw e- from O-H bonds, making them highly polar - Al(H2O)63+ hydrolyzes in aqueous sol n by donating proton (from acidic H2O molecules on ion) to water Predicting the Acidity/Basicity of a Salt 3.) A salt of a weak base and a strong acid. - the cation of the salt is the conjugate of the weak base. It hydrolyzes to give an acidic solution - eg. NH4Cl 4.) A salt of a weak base and a weak acid. - both ions hydrolyze. You must compare the Ka of the cation with the Kb of the anion - if the Ka of the cation is larger the solution is acidic - if the Kb of the anion is larger, the solution is basic 3 Predicting the Acidity/Basicity of a Salt - to predict the acidity or basicity of a salt sol n, examine acidity or basicity of ions composing salt Consider KC2H3O2 (potassium acetate) - potassium is Group IA element does not hydrolyze - acetate ion is conjugate of acetic acid (weak acid); therefore, acetate ion is basic K+ (aq) + H2O (l) no reaction C2H3O2 (aq) + H2O (l) HC2H3O2 (aq) + OH- (aq) The pH of a Salt Solution Consider pH of 0.10 M NaCN - already seen that sol n should be basic due to the hydrolysis of CN- ion - require the Kb for CN- (related to Ka for HCN) Relation between Ka and Kb - relationship between Ka and Kb for conjugate acidbase pairs HCN (aq) + H2O (l) H3O+ (aq) + CN- (aq) Ka CN- (aq) + H2O (l) HCN (aq) + OH- (aq) Kb ___________________________________________________ 2 H2O (l) H3O+ (aq) + OH- (aq) Kw - by the rule of multiple equilibria (chpt. 14) KaKb = Kw Consider three possible cases: a.) Ka > Kb - if Ka for cation is greater than Kb for anion, the solution will contain excess of H3O+ ions (pH < 7) b.) Ka < Kb - if Ka for cation is less than Kb for anion, the solution will contain excess of OH- ions (pH > 7) c.) Ka Kb - if Ka for cation and Kb for anion are comparable, the solution contains approximately equal concentrations of H3O+ and OH- ions (pH 7) Buffer Solutions - a solution characterized by the ability to resist changes in pH when limited amounts of acid or base are added to it - contain either a weak acid and its conjugate base or a base and it conjugate acid - so, buffer solution contains both an acid species and a base species in equilibrium Consider buffer system containing approximately equal amounts of weak acid, HA, and conjugate base, A- The Common-Ion Effect - the shift in an ionic equilibrium caused by the addition of a solute that provides an ion that takes part in the equilibrium Note: the common-ion effect is just another example of Le Chatelier s principle (applies to weak acids and bases Consider a solution of HC2H3O2 HC2H3O2 (aq) + H2O (l) C2H3O2- (aq) + H3O+ (aq) - now add HCl or C2H3O2HC2H3O2 (aq) + H2O (l) C2H3O2- (aq) + H3O+ (aq) added or added *** degree of ionization of acetic acid is decreased by the addition of a strong acid OH- added to buffer (strong base) HA (aq) + OH- (aq) A- (aq) + H2O (l) - the pH increases slightly H3O+ added to buffer (strong acid) A- (aq) + H3O+ (aq) HA (aq) + H2O (l) - the pH decreases slightly 4 Two important characteristics of a buffer: 1.) pH 2.) buffer capacity - the amount of acid or base the buffer can react with before giving a significant pH change - depends on the amount of acid and conjugate base in the solution Figure 16.10 - now, use the preceding equation to derive an equation for the pH of a buffer - take negative logarithm of both sides of equation so, - log [H3O+] = - log Ka x [HA] = - log Ka - log [HA] [A-] [A-] Note: left side equals pH and can simplify right side pKa of weak acid - defined in same manner as pH and pOH pKa = -log Ka Now, pH = pKa log [HA] = pKa + [A-] [A-] [HA] Henderson-Hasselbalch Equation How do you prepare a buffer of given pH? - show that buffer must be prepared from conjugate acidbase pair in which acid-ionization constant is approximately equal to desired H3O+ concentration Consider a buffer made of weak acid, HA, and conjugate base, A-: HA (aq) + H2O (l) H3O+ (aq) + A- (aq) + to give Ka = [H3O ][A ] [HA] and rearrange to give [H3O+] = Ka x [HA] [A-] Henderson-Hasselbaslch equation - an equation relating the pH of a buffer for different concentrations of conjugate acid and base Generally, pH = pKa + log [base] [acid] - by substituting the value of pKa for the conjugate acid and the ratio [base]/[acid], you obtain the pH of the buffer 5 This document was created with Win2PDF available at http://www.daneprairie.com. 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