ASRJC 2025 H2 Chem Acid-Base Equilibria Notes

2025 JC2 H2 Chemistry of Aqueous Solutions: Acid–base Equilibria (Student) 2025/ASRJC/Chemistry 1 ANDERSON SERANGOON JUNIOR COLLEGE JC2 H2 CHEMISTRY ACID–BASE EQUILIBRIA CONTENT AND SECTION Remarks 1 Theories of Acids & Bases Lecture 1 (SDL) 13 Jan 2025 1.1 : Arrhenius Theory of Acids & Bases 1.2 : Brønsted–Lowry Theory of Acids & Bases 1.3 : Lewis Theory of Acids & Bases 2. The pH Scale 3. The Ionic Product of Water (Kw) 4. Strengths of Acids & Bases 4.1 : Calculations involving a Strong Acid or Strong Alkali Lecture 2 4.2 : Weak Acids & Weak Bases (A) Strength of Acids and Bases & Dissociation Constants (B) Calculations involving a Weak Acid or Weak Alkali 4.3 : Complementary Strengths of a Conjugate Acid–Base Pair 5 Salt Solutions Lecture 3 5.1 : Salt Hydrolysis 5.2 : Calculation of pH of salt solution 6 Buffer Solutions 6.1 : Types of Buffer Solutions 6.2 : How a Buffer System Controls pH 6.3 : Calculating pH of Buffer Solutions 6.4 : Buffer Capacity and Buffer Range 6.5 : Importance & Uses of Buffer Solutions Lecture 1 (SDL) 13 Jan 2025 7 Acid–Base Titration Lecture 4 7.1 : Indicators 7.2 : Titration Curves (A) Strong acid – strong base (SA–SB) titration (B) Weak acid – strong base (WA–SB) titration (C) Strong acid – weak base (SA–WB) titration Lecture 5 (D) Weak acid – weak base (WA–WB) titration Examples of titration curves (E) Titration involving polybasic acids Summary – how to use the last 2 pages of the notes

2025 JC2 H2 Chemistry of Aqueous Solutions: Acid–base Equilibria (Student) 2025/ASRJC/Chemistry 2 Content Map Content (i) Acid dissociation constants, Ka and the use of pKa (ii) Base dissociation constants, Kb and the use of pKb (iii) The ionic product of water, Kw (iv) pH: choice of pH indicators (v) Buffer solutions Learning Outcomes Students should be able to: (a) explain qualitatively the differences in behaviour between strong and weak acids and bases in terms of the extent of dissociation (b) explain the terms pH; Ka; pKa; Kb; pKb; Kw and apply them in calculations, including the relationship Kw = KaKb (c) calculate [H+(aq)] and pH values for strong acids, weak monobasic (monoprotic) acids, strong bases, and weak monoacidic bases [Calculations involving weak acids/bases will not require solving of quadratic equations] (d) describe the changes in pH during acid–base titrations and explain these changes in terms of the strengths of the acids and bases (e) explain the choice of suitable indicators for acid–base titrations, given appropriate data (f) (i) explain how buffer solutions control pH (ii) describe and explain their uses, including the role of H2CO3/HCO3– in controlling pH in blood (g) calculate the pH of buffer solutions, given appropriate data References 1. Chemistry for Advanced Level, Cann and Hughes, Murray 2. Chemistry, The Molecular Nature of Matter and Change (Fourth Edition), Silberberg, McGraw Hill 3. Chemistry & Chemical Reactivity (Sixth Edition), Kotz, Treichel and Weaver, Thomson 4. Chemistry The Central Science (Ninth Edition), Brown, LeMay, Bursten, Prentice Hall 5. Understanding Advanced Physical Inorganic Chemistry, The Learner’s Approach, Jeanne Tan, Kim Seng Chan, WS Education

2025 JC2 H2 Chemistry of Aqueous Solutions: Acid–base Equilibria (Student) 2025/ASRJC/Chemistry 3 1 Theories of Acids and Bases 1.1 Arrhenius Theory of Acids & Bases • An Arrhenius acid is a substance that releases H + ions and an Arrhenius base is a substance that releases OH– when they are dissolved in water. • An acid–base reaction occurs when H + reacts with OH− to produce H2O. This reaction is also known as neutralisation. • Example: NaOH(aq) + HCl(aq) → NaCl(aq) + H2O(l) 1.2 Brønsted–Lowry Theory of Acids & Bases • A Brønsted acid is a proton (H+) donor. • A Brønsted base is a proton (H+) acceptor. • An acid–base reaction involves transfer of proton (H+) from the acid to the base. • Example: HCl(aq) + NH 3(aq) → NH4+(aq) + Cl–(aq) 1.2.1 Conjugate Acid-Base Pairs HA(aq) + : B(aq) → A–(aq) + BH+(aq) acid base conjugate base conjugate acid • When a Brønsted−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 the conjugate acid of B. • Consider the following equilibrium: • An acid and a base such as CH3COOH and CH3COO−, which differ from each other by only one proton, H+, are referred to as a conjugate acid-base pair. CH3COOH(aq) + H 2O(l) CH3COO−(aq) + H 3O+(aq) Acid1 Base2 Conjugate base1 Accepted H+ Conjugate acid2 Donated H+ ⇌

2025 JC2 H2 Chemistry of Aqueous Solutions: Acid–base Equilibria (Student) 2025/ASRJC/Chemistry 4 Exercise 1. Identify the acid, base, conjugate acid and base in the following reactions: (a) HNO3(aq) + H 2O(l) → NO3 – (aq) + H3O+(aq) acid base conj ugate base conj ugate acid (b) HCOOH(aq) + OH –(aq) → HCOO–(aq) + H2O(l) acid base conjugate _____ conjugate _______ (c) NH4+(aq) + S2−(aq) → HS−(aq) + NH3(aq) acid base conjugate _____ conjugate _______ 2. Circle the conjugate acid-base pairs from the following sets of substances. (a) NH4+ NH3 NH2+ (b) CH3O+ CH3OH CH3O− 1.3 Lewis Theory of Acids & Bases • A Lewis acid is an electron–pair acceptor. • A Lewis base is an electron–pair donor. • An acid–base reaction involves the donation of an electron pair by a Lewis base to a Lewis acid to form a covalent bond. • Example: BF3 acts as a Lewis acid and F− acts as a Lewis base to form BF4−. B F F F + F- B F F F F- Summary ACID BASE Arrhenius releases H+ ions in aqueous solution releases OH– in aqueous solution Brønsted– Lowry proton (H+) donor proton (H+) acceptor Lewis electron–pair acceptor electron–pair donor This topic deals with acid –base equilibria in aqueous solution s. Hence, the Brønsted– Lowry definitions of acid/base are primarily used in the understanding of the pH of solutions.

2025 JC2 H2 Chemistry of Aqueous Solutions: Acid–base Equilibria (Student) 2025/ASRJC/Chemistry 5 2 THE pH SCALE The pH scale was developed to measure the acidity of solutions based on the concentration of hydrogen ions, H+ (or hydronium ions, H3O+) in the solution. Experimentally, the pH value is found to lie between 0 and 14 for most aqueous solutions at 298 K (25 oC). The pH of a solution is defined as: pH = – log10[H+] or – log10[H3O+] Conversely, pOH is a measure of the concentrati on of hydroxide ions, OH– in an aqueous system. pOH is defined as: pOH = – log10[OH–] 3 The Ionic Product of Water (Kw) Pure water conducts electricity slightly as it is ionised (or dissociated) to a small extent, to form proton and hydroxide ion. H2O(l) H+(aq) + OH −(aq) Since H+ has a very high charge density, it attracts another water molecule, forming H3O+ and the dissociation equation can also be represented by equation below. 2H2O(l) H3O+(aq) + OH–(aq) The above reaction is also known as the auto–ionisation of water, and the Kc expression is: Kc = 2 2 3 ]OH[ ]OH][OH[ −+ Since [H2O(l)] is a constant, we can incorporate it into the equilibrium constant, Kc, and define a new constant, Kw: Kc x [H2O]2 = [H3O+][OH–] Kw = [H3O+][OH–]

2025 JC2 H2 Chemistry of Aqueous Solutions: Acid–base Equilibria (Student) 2025/ASRJC/Chemistry 6 The constant Kw is known as the ioni