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Concepts of Acid-Base and Electrolytic Properties

Explore the properties of electrolytes, nonelectrolytes, acids, and bases. Learn about the methods of distinguishing between electrolytes and nonelectrolytes. Understand the definitions of acids and bases by Arrhenius and Brønsted-Lowry. Discover the role of water in showing properties of acids and alkalis.

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Concepts of Acid-Base and Electrolytic Properties

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  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.

  5. 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

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

  7. 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.

  8. 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:

  9. 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

  10. 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

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

  12. 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

  13. 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

  14. 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!

  15. 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

  16. 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

  17. 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

  18. 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

  19. + 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

  20. 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

  21. 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)

  22. 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

  23. 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

  24. 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)

  25. 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.

  26. 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

  27. 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

  28. 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.

  29. 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

  30. [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

  31. = [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

  32. 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+]

  33. 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

  34. 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

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

  36. 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

  37. 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

  38. 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

  39. = 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.

  40. 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.

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