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ERT 207 ANALYTICAL CHEMISTRY ACIDS AND BASES THEORIES ACID BASE EQUILIBRIA IN WATER pH SCALE 27 Jan 2011 (MISS NOORULNAJWA DIYANA YAACOB) 1 Brønsted Acids and Bases • A Brønsted acid is a species that donates a proton. (a proton donor). HCl(aq) + H2O(l) acid H3O+(aq) + hydronium ion Cl-(aq) conjugate base A Conjugate base is what remains of the acid after the donation of a proton. • A Brønsted base is a species that accepts a proton. (a proton acceptor). NH3(aq) base + H2O(l) NH4+(aq) + OH-(aq) conjugate acid A Conjugate acid is a newly formed protonated species. 2 Brønsted Acids and Bases Any reaction, using the Brønsted theory, involves both an acid and a base. Water is an amphoteric species which can act as a Brønsted acid or as a Brønsted base. 3 Brønsted Acids and Bases 4 Brønsted Acids and Bases • Exercise: Identify acids, conjugate acids, bases and conjugate bases in the reactions below: (a) H2SO4(aq) + H2O(l) acid base (b) NH4+(aq) + H2O(l) acid base HSO4‒ (aq) + conjugate base H3O+(aq) conjugate acid + conjugate acid • Exercise: ‒ What is the conjugate base of H2PO4 ? H2PO4‒(aq) + H2O(l) H3O+(aq) NH3(aq) conjugate base HPO42‒(aq) + H3O+(aq) conjugate base 5 Lewis Acids and Bases • The Brønsted base is the species that accepts a proton. • The Brønsted acid is the species that donates a proton. • A Lewis base is a species that donates a pair of electrons. • A Lewis acid is a species that accepts a pair of electrons. H+ + ‒ O H O H Lewis base Lewis acid H H+ Lewis acid H + N + H H H Lewis base H N H H 6 Lewis Acids and Bases • Lewis acid-base reactions are more general than Brønsted acid-base reactions. • Lewis acid-base reactions result in the formation of coordinate covalent bonds. 7 ARRHENIUS THEORY-H+ AND OHˉ • An acid is any substance that ionizes (partially or completely) in water to give hydrogen ions • Means associate with solvent to give hydronium ions • HA +H2O ↔ H3O+ + Aˉ • A base ionizes in water to give hydroxyl ions. • Weak (partially ionizes) bases ionizes as follows: • B + H2O ⇋ BH+ + OHˉ For strong bases e.g.: (NaOH) M(OH)n →M n+ + nOHˉ • this theory is obviously restricted to water as a solvent only Franklin’s theory • Recognizes the ionization of a solvent to give a cation and an anion • E.g: 2H2O ⇋ H3O+ + OHˉ • 2NH3 ⇋ NH4 + NH2¯ • Acid = a solute that yields the cation of the solvent • Base = a solute that yields the anion of the solvent • Eg of strong acid is NH4Cl in liquid ammonia • Eg of strong base is NaNH2 in ammonia ACID-BASE EQUILIBRIA IN WATER • When acid or base is dissolved in water, it will dissociate • Amount of ionization depends on the strength of the acid • Strong electrolyte completely dissociated • Weak electrolyte dissociated partially • Example: a) Hydrochloric acid (strong acid) HCl H 2O H 3O Cl a) Acetic acid (weak acid) HOAc H 2O H 3O OAc • Equilibrium constant K 0 a aH O aOAc 3 aHOAc aH 2O 0 K • a is the thermodynamic activity constant • H+ will be used instead of H3O+ for simplification • Molar concentration will be replaced by [] H OAc Ka HOAc K w H OH • Ka and Kw are molar equilibrium constants • [H+]=1.0×10-7 M = [OH-] The Acid-Base Properties of Water • Water is a very weak electrolyte. It undergoes ionization to a very small extent. H2O (l) H+ (aq) + OH‒ (aq) Autoionization of + ‒ 2H2O (l) H3O (aq) + OH (aq) water • The equilibrium expression of water autoionization is: Kw = [H+] [OH‒] or Kw = [H3O+] [OH‒] 14 Equilibrium Constant in Aqueous Solutions • For any aqueous solution: Kw = [H3O+] [OH‒] = 1.0 × 10-14 at 25C • Remember that Kw is always constant at a given temperature. However, there can be infinite possibilities of the equilibrium positions for aqueous solution. – When [H3O+] = [OH‒] ; the solution is neutral. – When [H3O+] > [OH‒] ; the solution is acidic. – When [H3O+] < [OH‒] ; the solution is basic. 15 Equilibrium Constant in Aqueous Solutions 16 Equilibrium Constant in Aqueous Solutions • Exercise: Calculate [OH‒] in a solution in which the concentration of protons is 0.0012 M at 25C. Is the solution acidic, basic or neutral? At 25C, Kw is always equal to 1 × 10-14 . Kw = [H3O+] [OH‒] 1 × 10-14 = (0.0012) [OH‒] [OH‒] = 1 × 10-14 / 0.0012 = 8.3 × 10-12 M The solution is acidic. 17 The pH Scale • The pH scale is a convenient way to express the acidity and basicity of a solution (the concentration of H3O+ ions). The pH of a solution is defined as: pH = -log [H3O+] or pH = -log [H+] • Example: Two solutions with hydronium ion concentrations of 1.0 × 10-4 M and 5.0 × 10-9 M. Find their pH values. 2 sig. fig. The first solution; The second solution; pH = - log (1.0 × 10-4) = 4.00 pH = - log (5.0 × 10-9) = 8.30 2 decimal places The number of sig. fig in the log is equal to 18 the number of decimal places in the answer pH scale Example 1 Calculate the pH of a 2.0 x 10-3 M solution of HCl. pH scale Solution [H+] = 2.0 x 10-3 M pH = − log10[H + ] = − log10[2.0 x 10-3 M ] = 3 - log10 2.0 = 3 – 0.30 = 2.7 The pH Scale • Exercise: Find the pH of a neutral solution. A neutral solution has [H3O+] = [OH‒] . Since Kw = [H3O+] [OH‒] = 1.0 × 10-14 then [H3O+] = 1.0 × 10-7 pH = - log (1.0 × 10-7) = 7.00 Lower pH values ( < 7.00 ) Neutral Higher pH values ( > 7.00 ) Higher [H3O+] Lower [H3O+] Acidic solution Basic solution 21 The pOH Scale • A pOH scale is analogous to the pH scale. It is defined as: pOH = - log [OH‒] Also, [OH‒] = 10-pOH 22 • Kw = [H+] [OH-] • -log Kw = - log [H+][OH-] = -log [H+] – log [OH-] pKw = pH + pOH pOH scale • PLEASE NOTE THAT AT 25˚C [H+] x [OH-] = 10-14 pH + pOH = 14 The pOH Scale • Recall that at 25C: Kw = 1.0 × 10-14 = [H3O+] [OH‒] ‒ log ([H3O+] [OH‒]) = ‒ log (1.0 × 10-14) ‒ log [H3O+] ‒ log [OH‒] = 14.00 pH + pOH = 14.00 pH 13.00 11.00 9.00 7.00 5.00 4.00 1.00 25 • Example Calculate the pOH and the pH of a 5.0 x 10-2 M solution of NaOH • Solution Method 1 [OH-] = 5.0 x10-2 M pOH = - log(5.0 x10-2 ) = 2 – log 5.0 = 1.3 pOH scale • Solution continued Since pH + 1.3 = 14.00 pH = 12.7 pH + pOH = 14 pOH scale Solution Method 2 Since [H+] x [OH-] = 10-14 [H+] = 1.0 x 10-14 = 2.0 x 10-13 5.0 x 10-12 pH = - log (2.0 x 10-13) = 13 – log 2.0 = 13 – 0.30 = 12.7 The pOH Scale • Exercise: What is the pOH of a solution that has a hydroxide ion concentration of 4.3 x 10-2 M? pOH log[OH ] pOH log[4.3 x 102 ] pOH 1.37 • Exercise: What is the hydroxide ion concentration of a solution with pOH 8.35? [OH ] 10pOH [OH ] 108.35 [OH ] 4.5x10 9 30