CHAPTER 3 CONCEPTS OF ACID-BASE NEUTRALIZATION

1. CHAPTER 3 CONCEPTS OF ACID-BASE NEUTRALIZATION

2. ELECTROLYTIC PROPERTIES • An electrolyte is a substance that, when dissolved in water, results in a solution that can conduct electricity. • A nonelectrolyte is a substance that, when dissolved, results in a solution that does not conduct electricity. nonelectrolyte strong electrolyte weak electrolyte

3. METHOD OF DISTINGUISHING BETWEEN ELECTROLYTES AND NONELECTROLYTES • A pair of inert electrodes (Cu or Pt) is immersed in a beaker of water. • To light the bulb, electric current must flow from one electrode to the other, thus completing the circuit. • By adding NaCl (ionic compound), the bulb will glow. • NaCl breaks up into Na+ and Cl- ions when dissolves in water. • Na+ are attracted to the negative electrode. • Cl- are attracted to the positive electrode. • The movement sets up an electric current that is equivalent to the flow of electrons along a metal wire.

4. Strong Electrolyte • – 100% dissociation (breaking up of compound into • cations and anions H2O NaCl (s) Na+ (aq) + Cl- (aq) • Weak Electrolyte • – not completely dissociated CH3COOH CH3COO- (aq) + H+ (aq) A reversible reaction. The reaction can occur in both directions.

6. Hydrationis the process in which an ion is surrounded by water molecules arranged in a specific manner. • Hydration helps to stabilize ions and prevents cations from combining with anions. d- d+ d+ H2O

7. H2O C6H12O6 (s) C6H12O6 (aq) Nonelectrolyte does not conduct electricity? No cations (+) and anions (-) in solution

8. 2HCl (aq) + Mg (s) MgCl2 (aq) + H2 (g) 2HCl (aq) + CaCO3 (s) CaCl2 (aq) + CO2 (g) + H2O (l) PROPERTIES ACIDS • Have a sour taste. • - Vinegar owes its taste to acetic acid. • - Citrus fruits contain citric acid. • Cause color changes in plant dyes. • React with certain metals to produce hydrogen gas. • React with carbonates and bicarbonates to produce carbon dioxide gas • Aqueous acid solutions conduct electricity.

9. PROPERTIES OF BASES • Have a bitter taste. • Feel slippery. Many soaps contain bases. • Cause color changes in plant dyes. • Aqueous base solutions conduct electricity. • Examples:

10. ROLE OF WATER TO SHOW PROPERTIES OF ACIDS • Anhydrous pure acid (without water) does not show acidic properties. • In dry form, acids exist as neutral covalent molecules. • Dry acids do not dissociate to form hydrogen ion (H+). • When a pure acid is dissolved in water, it will show the properties of acids. • This is because acids will dissociate in water to form H+ or hydroxonium/hydronium ion, H3O+ which are free to move. • For example: i) HCl in liquid methylbenzene (organic solvent) - does not show acidic properties. ii) HCl in water – show acidic properties

11. ROLE OF WATER TO SHOW PROPERTIES OF ALKALI • Dry base does not show alkaline properties. • A base in dry form, contains hydroxide ions (OH-) that are not free to move. Thus, the alkaline properties cannot be shown. • In the presence of water, bases can dissociate in water to form hydroxide ions, OH-, which are free to move. Thus, alkaline properties are shown. • For example: i) ammonia in tetrachlomethane (organic solvent) – do not show alkaline properties ii) ammonia in water – show alkaline properties

12. DEFINITION OF ACID AND BASE Brønsted-Lowry Arrhenius Lewis

13. DEFINITION OF ACID AND BASE BY ARRHENIUS • Arrhenius acid is a substance that produces H+ (hydrogen ion) or hydronium ion(H3O+) in water • Arrhenius base is a substance that produces OH- in water

14. Examples of bases: • NaOH (s) Na+ (aq) + OH- (aq) • N2H4 (aq) + H2O N2H5+ (aq) + OH- (aq) • metal oxides + H2O bases Na2O (s) + H2O (l) 2NaOH (aq) • Examples of acid: • CO2 (g) + H2O (l) H2CO3 (aq) H2CO3 (aq) + H2O(l) H3O+ (aq) + HCO3- (aq) • nonmetal oxides + H2O acid * Limited only to aqueous solutions

15. DEFINITION OF ACID AND BASE BY BRØNSTED-LOWRY • A Brønsted acid is a proton donor • A Brønsted base is a proton acceptor • Example: HCl (aq) +H2O (l) → H3O+ (aq) + Cl- (aq) acid base acid base • HCl is a acid because it donates proton to H2O • H2O is a base because it accepts proton from HCl • A Brønsted acid must contain at least one ionizable proton!

16. Brønsted acids and bases • Conjugate acid-base pair: i) Conjugate base of a Brønsted acid - the species that remains when one proton has been removed from the acid ii) Conjugate acid - addition of a proton to a Brønsted base

17. Examples: • HCl (aq) +H2O (l) H3O+ (aq) + Cl- (aq) • acid1base2acid2base1 • Cl- is a conjugate base of HCl and HCl is a conjugate acid of Cl- • H2O is a base conjugate of H3O+ and H3O+ is a acid conjugate of • H2O • NH3 (aq) + H2O (l) NH4+ (aq) + OH- (aq) • base1acid2 acid1base2 • subscripts 1 and 2 = two conjugate acid-base pair

18. When a strong acid react with a strong base in Brønsted acid-base reaction, it will give a weak conjugate acid and conjugate base. • Examples: • HCl (aq) + H2O (l) H3O+ (aq) + Cl- (aq) • strong acid strong base weak conjugate weak conjugate • acid base • NH3 (aq) + H2O (l) NH4+ (aq) + OH- (aq) • Weak base weak acid strong conjugate strong conjugate • acid base • H2O can function as acid or base which called amphoteric • Amphoteric or amphiprotic substance is one that can react as either an acid or base

19. HI (aq) H+ (aq) + I- (aq) CH3COO- (aq) + H+ (aq) CH3COOH (aq) H2PO4- (aq) H+ (aq) + HPO42- (aq) H2PO4- (aq) + H+ (aq) H3PO4 (aq) Identify each of the following species as a Brønsted acid, base, or both. (a) HI, (b) CH3COO-, (c) H2PO4- Brønsted acid Brønsted base Brønsted acid Brønsted base

21. + OH- H H H O H + H+ H N H N H • • • • • • • • • • • • H H DEFINITION OF ACID AND BASE BY LEWIS • A Lewis acid is a substance that can accept a pair of electrons • A Lewis base is a substance that can donate a pair of electrons H+ acid base + acid base

22. H F F F B F B N H F F H H N H H • • Examples of Lewis Acids and Bases reactions: a) + acid base b) Ag+ (aq) + 2NH3 (aq) Ag(NH3)2+ (aq) acidbase c) Cd+ (aq) + 4I- (aq) CdI2-4 (aq) acidbase d) Ni (s) + 4CO (g) Ni(CO)4 (g) acidbase

25. TYPES OF ACIDS-BASES • Acids i) Strong acids: • Acids that completely ionized in solution. • Example: HCl (aq) → H+ (aq) + Cl- (aq) ii) Weak acids • Acids that incompletely ionized in solution • Example: CH3COOH (aq) CH3COO- (aq) + H+ (aq)

27. HCl H+ + Cl- HNO3H+ + NO3- CH3COOH H+ + CH3COO- H2SO4H+ + HSO4- HSO4-H+ + SO42- • Monoprotic acid: - each unit of the acid yields one hydrogen ion upon ionization Strong electrolyte, strong acid Strong electrolyte, strong acid Weak electrolyte, weak acid • Diprotic acid: • - each unit of the acid gives up two H+ ions, in two separate steps Strong electrolyte, strong acid Weak electrolyte, weak acid

28. H3PO4H+ + H2PO4- H2PO4-H+ + HPO42- HPO42-H+ + PO43- • Triprotic acids: • - yield three H+ ions Weak electrolyte, weak acid Weak electrolyte, weak acid Weak electrolyte, weak acid

29. Bases i) Strong bases: • Bases that completely ionized in solution. • Example: NaOH (s) → Na+ (aq) + OH- (aq) ii) Weak bases • bases that incompletely ionized in solution • Example: NH3 (aq) + H2O (l) NH4+ (aq) + OH- (aq)

30. HCl (aq) + H2O (l) H3O+ (aq) + Cl- (aq) HNO3 (aq) + H2O (l) H3O+ (aq) + NO3- (aq) HClO4 (aq) + H2O (l) H3O+ (aq) + ClO4- (aq) H2SO4 (aq) + H2O (l) H3O+ (aq) + HSO4- (aq) Acids and bases as electrolytes • Strong acids such as HCl and HNO3 are strong electrolytes, while weak acid such as acetic acid (CH3COOH) is a weak electrolyte.

31. HF (aq) + H2O (l) H3O+ (aq) + F- (aq) HNO2 (aq) + H2O (l) H3O+ (aq) + NO2- (aq) HSO4- (aq) + H2O (l) H3O+ (aq) + SO42- (aq) H2O (l) + H2O (l) H3O+ (aq) + OH- (aq) HCl (aq) + H2O (l) H3O+ (aq) + Cl- (aq) HNO3 (aq) + H2O (l) H3O+ (aq) + NO3- (aq) HClO4 (aq) + H2O (l) H3O+ (aq) + ClO4- (aq) H2SO4 (aq) + H2O (l) H3O+ (aq) + HSO4- (aq) Acids and bases as electrolytes Strong Acids are strong electrolytes Weak Acids are weak electrolytes

32. F- (aq) + H2O (l) OH- (aq) + HF (aq) NO2- (aq) + H2O (l) OH- (aq) + HNO2 (aq) H2O NaOH (s) Na+ (aq) + OH- (aq) H2O KOH (s) K+ (aq) + OH- (aq) H2O Ba(OH)2 (s) Ba2+ (aq) + 2OH- (aq) Strong Bases are strong electrolytes Weak Bases are weak electrolytes

33. Conjugate acid-base pairs: • The conjugate base of a strong acid has no measurable strength. • H3O+ is the strongest acid that can exist in aqueous solution. • The OH- ion is the strongest base that can exist in aqueous solution.

35. H2O (l) H+(aq) + OH-(aq) ACID-BASE PROPERTIES OF WATER • Can act either as a acid or as a base. • Water functions as a base with acids such as HCl and CH3COOH and function as acid in reaction with bases. • Water is a very weak electrolyte and undergo ionization to a small extent: autoionization of water

36. [H+][OH-] Kc = [H2O] H2O (l) H+ (aq) + OH- (aq) The Ion Product of Water [H2O] = constant Kc = equilibrium constant Kc[H2O] = Kw = [H+][OH-] The ion-product constant (Kw) is the product of the molar concentrations of H+ and OH- ions at a particular temperature. Solution Is [H+] = [OH-] neutral At 250C Kw = [H+][OH-] = 1.0 x 10-14 [H+] > [OH-] acidic [H+] < [OH-] basic

37. = [OH-] = 1 x 10-14 Kw 1.3 [H+] What is the concentration of OH- ions in a HCl solution whose hydrogen ion concentration is 1.3 M? Kw = [H+][OH-] = 1.0 x 10-14 [H+] = 1.3 M = 7.7 x 10-15M

39. pH-A MEASURE OF ACIDITY • pH – the negative logarithm of the hydrogen in • concentration (in mol/L) pH = -log [H+] Solution Is At 250C neutral [H+] = [OH-] [H+] = 1 x 10-7 pH = 7 [H+] > 1 x 10-7 pH < 7 acidic [H+] > [OH-] basic [H+] < [OH-] [H+] < 1 x 10-7 pH > 7 pH [H+]

40. Other important relationships pOH = -log [OH-] [H+][OH-] = Kw = 1.0 x 10-14 -log [H+] – log [OH-] = 14.00 pH + pOH = 14.00 pH Meter

41. 1) The pH of rainwater collected in a certain region of the northeastern United States on a particular day was 4.82. What is the H+ ion concentration of the rainwater? pH = -log [H+] = 10-4.82 = 1.5 x 10-5M [H+] = 10-pH 2) The OH- ion concentration of a blood sample is 2.5 x 10-7 M. What is the pH of the blood? pH + pOH = 14.00 pOH = -log [OH-] = -log (2.5 x 10-7) = 6.60 pH = 14.00 – pOH = 14.00 – 6.60 = 7.40

44. CALCULATION OF pH FOR SOLUTION CONTAINING A STRONG ACID AND A SOLUTION OF A STRONG BASE

45. HNO3 (aq) + H2O (l) H3O+ (aq) + NO3- (aq) Ba(OH)2 (s) Ba2+ (aq) + 2OH- (aq) 1) What is the pH of a 2 x 10-3 M HNO3 solution? HNO3 is a strong acid – 100% dissociation. 0.0 M 0.0 M Start 0.002 M 0.0 M 0.002 M 0.002 M End pH = -log [H+] = -log [H3O+] = -log(0.002) = 2.7 2) What is the pH of a 1.8 x 10-2 M Ba(OH)2 solution? Ba(OH)2 is a strong base – 100% dissociation. 0.0 M 0.0 M Start 0.018 M 0.0 M 0.018 M 0.036 M End pH = 14.00 – pOH = 14.00 + log(0.036) = 12.6

46. HA (aq) + H2O (l) H3O+ (aq) + A- (aq) HA (aq) H+ (aq) + A- (aq) [H+][A-] Ka = [HA] Weak Acids (HA) and Acid Ionization Constants Ka is the acid ionization constant weak acid strength Ka

48. HF (aq) H+ (aq) + F- (aq) = 7.1 x 10-4 = 7.1 x 10-4 = 7.1 x 10-4 [H+][F-] x2 x2 Ka Ka = Ka = 0.50 - x [HF] 0.50 1) What is the pH of a 0.5M HFsolution (at 250C)? HF (aq) H+ (aq) + F- (aq) Initial (M) 0.50 0.00 0.00 Change (M) -x +x +x Equilibrium (M) 0.50 - x x x Ka << 1 0.50 – x 0.50 x2 = 3.55 x 10-4 x = 0.019 M pH = -log [H+] = 1.72 [H+] = [F-] = 0.019 M [HF] = 0.50 – x = 0.48 M

49. = 7.1 x 10-4 0.019 M 0.006 M x2 x 100% = 3.8% x 100% = 12% 0.50 M 0.05 M Ka 0.05 When can I use the approximation? 0.50 – x 0.50 Ka << 1 When x is less than 5% of the value from which it is subtracted. Less than 5% Approximation ok. x = 0.019 1) What is the pH of a 0.05M HFsolution (at 250C)? x = 0.006 M More than 5% Approximation not ok. Must solve for x exactly using quadratic equation or method of successive approximations.

50. Solving weak acid ionization problems: • Identify the major species that can affect the pH. • In most cases, you can ignore the autoionization of water. • Ignore [OH-] because it is determined by [H+]. • Use ICE to express the equilibrium concentrations in terms of single unknown x. • Write Kain terms of equilibrium concentrations. Solve for x by the approximation method. If approximation is not valid, solve for x exactly. • Calculate concentrations of all species and/or pH of the solution.