2022 H1 Theories of acids and bases notes (teacher)

1 St. Andrew’s Junior College H1 Chemistry 2022 Lecture Notes Theories of Acids and Bases Assessment Objectives Candidates should be able to: (a) show understanding of, and apply the Arrhenius theory of acids and bases (b) show understanding of, and apply the Bronsted-Lowry theory of acids and bases, including the concept of conjugate acids and bases; (c) explain qualitatively the differences in behaviour between strong and weak acids and bases in terms of the extent of dissociation; (d) explain the terms pH; Ka; Kb; Kw [The relationship Kw = KaKb is not required]; (e) calculate [H+(aq)] and pH values for strong acids, and strong bases; (f) explain the choice of suitable indicators for acid-base titrations, given appropriate data, in terms of the strengths of the acids and bases; (g) (i) explain how buffer solutions control pH; (ii) describe and explain the uses of buffers, including the role of H2CO3/HCO3- in controlling pH in blood. Lecture Outline: 1. Theories of acids and bases 2. Autoionisation of water 3. Dissociation of weak acid and bases 4. Acid-Base Properties of Salt Solutions 5. Indicators for acid-base titrations 6. Monoprotic Acid-Base Titration Curves 7. Buffer solutions References: 1. Chemistry by Silberberg 2. Chemistry by Raymond Chang 3. Chemistry for Advanced level by Peter Cann 4. Chemistry in context by Hill and Holman 5. A – level Chemistry by Ramsden

2 1. THEORIES OF ACIDS AND BASES 1.1 Arrhenius theory of acids and bases An Arrhenius acid is a substance that produces H+ in aqueous solution. An Arrhenius base is a substance that produces OH– in aqueous solution. E.g.: HNO3(l) + H2O(l) NO3–(aq) + H3O+(aq) Note: H3O+ is hydronium ion, which is the same as H+ (aq) NaOH(s) + aq Na+(aq) + OH–(aq)  An acid-base reaction involves the formation of salt and water 1.2 Brønsted–Lowry theory of acids and bases A Brønsted–Lowry acid is a proton (H+) donor. A Brønsted–Lowry base is a proton (H+) acceptor. base conjugate acid E.g.: NH3(aq) + H2O(l) NH4+(aq) + OH–(aq) Acid conjugate base base conjugate acid E.g.: HCl(g) + NH3 (g) NH4Cl(s) acid conjugate base  An acid–base reaction involves the transfer of a proton from the acid to the base.  When a Br ønsted–Lowry acid (HA) loses a proton, the resulting product (A –) is called the conjugate base of HA.  When a Brønsted–Lowry base (B) accepts a proton, the resulting product (BH+) is called Candidates should be able to:  show understanding of, and apply the Arrhenius theory of acids and bases  show understanding of, and apply the Brønsted–Lowry theory of acids and bases, including the concept of conjugate acids and conjugate bases

3 the conjugate acid of B.  A conjugate acid-base pair differs by a H+. Examples of conjugate pairs: HCl / Cl– NH3 / NH4+ H2O / OH– H2O / H3O+ CH3COOH / CH3COO– CH3NH2 / CH3NH3+ Exercise 1 Identify the acid, base, conjugate acid and conjugate base in the following reactions : a) CH3COOH + H2O CH3COO– + H3O+ acid base conjugate base conjugate acid b) CH3COOH + HCl CH3COOH2+ + Cl– base acid conjugate acid conjugate base c) HNO3 + H2O NO3– + H3O+ acid base conjugate base conjugate acid d) HNO3 + H2SO4 HSO4– + H2NO3+ base acid conjugate base conjugate acid  Self–Check: Q1 ٢ In Summary ٢ Arrhenius Brønsted-Lowry Acids produces H+ in water donates H+ Bases produces OH– in water accepts H+

4 1.3 Strength of Acids and Bases Strong acid completely dissociates in water. E.g.: HCl (g) + aq H+(aq) + Cl–(aq) Most mineral acids are strong acids. E.g.: HI, HBr, H2SO4, HNO3 etc. represents complete dissociation Acid Weak acid partially dissociates in water. E.g.: CH3COOH (l) + aq CH3COO–(aq) + H+(aq) Most organic acids are weak acids. E.g.: CH3COOH, H2CO3, HCN etc. ⇌ means an equilibrium is established between the undissociated molecules and the dissociated ions. Strong base completely dissociates in water. E.g.: NaOH(s) + aq Na+(aq) + OH–(aq) Examples of strong bases are NaOH, KOH, Ba(OH)2 etc. Base Weak base partially dissociates in water. E.g.: NH3 + H2O NH4+(aq) + OH–(aq) Examples of weak bases are NH3, CH3NH2, Na2CO3 etc. Candidates should be able to:  Explain qualitatively the differences in behavior between strong and weak acids and bases in terms of the extent of dissociation. Which is a stronger acid, 0.0001 mol dm–3 HCl or 1 mol dm–3 CH3COOH? HCl is a strong acid, regardless of its concentration. CH3COOH is a weak acid, regardless of its concentration . It only partially dissociates in water Strength of acid  Concentration of acid

5 1.4 The pH scale Concentration of H+ can be used as a measure of acidity and alkalinity of a solution. However, [H+] can be very small values. Hence, we can express these values as pH. pH = – log10 [H+] or [H+ ] = 10–pH [H+] in mol dm–3 A solution with low pH has a high [H+], and is more acidic than a solution with higher pH. Similarly, pOH = – log10 [OH–] Note: The logarithm used are to the base 10 (not to the base e), so make sure when doing calculations, you press the log or lg button on your calculator (not the ln button) 1.5 Basicity of acid  HCl and CH3COOH are monoprotic (monobasic) acids as each acid can only lose one H+.  H2SO4 is a diprotic (dibasic) acid as it can lose two H+. H2SO4 + H2O H3O+ + HSO4- conjugate base HSO4- + H2O H3O+ + SO42- conjugate base  Conversely for bases, NaOH is a monoacidic base and Ca(OH)2 is a diacidic base. 2. AUTOIONISATION OF WATER Water dissociates very slightly to give ions. H2O(l) H+(aq) + OH–(aq) When [H+] = [OH–], the solution is neutral. When [H+] > [OH–], the solution is acidic. When [H+] < [OH–], the solution is alkaline. Candidates should be able to:  explain pH and apply them in calculations Candidates should be able to:  explain Kw and apply them in calculations  calculate [H+(aq)] and pH values for strong acids and strong bases

6 We name the equilibrium constant for the autoionisation of water as ionic product of water, Kw. Note: [H2O] is not reflected in Kw pKw = – log10 Kw = – log10 ([H+] [OH–] ) = – log10 [H+] – log10 [OH–] pKw = pH + pOH pH = pKw – pOH Note: Derivation of pKw is not required At 25 oC in Data Booklet pKw = 14 pKw = pH + pOH 14 = pH + pOH Note: Derivation of pKw is not required In pure water at 25oC, H2O(l) H+(aq) + OH–(aq) [H+] = [OH–] = 1.00 x 10–7 mol dm–3 pH = pOH = – lg (1.00 x 10–7) = 7.0 The autoionisation of water molecules is an endothermic process. H2O(l) H+(aq) + OH–(aq) H > 0 Kw = . When temperature increases, both kf and kb increases. Since the forward reaction is endothermic and is favoured, k f increases more than k b. Hence, Kw increases when temperature increases. Table 1: Variation of ionic product of water, Kw, with temperature Temperature / oC Kw / mol2 dm–6 20 0.68 x 10–14 30 1.47 x 10–14 40 2.92 x 10–14 Kw = 1.00 x