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Acid-Base Equilibria

Acid-Base Equilibria. BLB 12 th Chapter 16. Expectations. Distinguish between acids and bases Definitions & properties Know common strong and weak examples Calculate pH for strong and weak systems Write chemical reactions of acids and bases Predict relative acid-base strength.

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Acid-Base Equilibria

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  1. Acid-Base Equilibria BLB 12thChapter 16

  2. Expectations • Distinguish between acids and bases • Definitions & properties • Know common strong and weak examples • Calculate pH for strong and weak systems • Write chemical reactions of acids and bases • Predict relative acid-base strength

  3. Examples of acids & bases

  4. 16.1 Acids & Bases: A Brief Review • Arrhenius Definitions • Acid – a substance that produces hydrogen ions (H+) in water HA → H+ + A- • Base – a substance that produces hydroxide ions (OH-) in water BOH → B+ + OH-

  5. 16.2 Brønsted-Lowry Acids & Bases • H+ (proton) in water: H+ + H2O → H3O+ hydronium ion • Hydronium ion can hydrogen bond with more water molecules to form large clusters of hydrated hydronium ions. • H+ and H3O+ are used interchangeably.

  6. 16.2 Brønsted-Lowry Acids & Bases • Brønsted-Lowry definitions acid – hydrogen ion (or proton) donor • Neutral (HNO3), anionic (HCO3-), cationic (NH4+) • Must have a removable (acidic) proton base – hydrogen ion (or proton) acceptor • Neutral (NH3), anionic (HCO3- , CO32-) • Must have a lone pair of electrons

  7. Acid-Base Reactions (H+-transfer reactions) HCl(aq) + H2O(l) → Cl-(aq) + H3O+(aq) NH3(aq) + H2O(l) ⇌ NH4+(aq) + OH-(aq) HCl(aq) + NH3(aq) → Cl-(aq) + NH4+(aq)

  8. Acid-base reaction in non-aqueous media:HCl(g) + NH3(g) → NH4Cl(s)

  9. amphiprotic – capable of behaving as a Brønsted acid and Brønsted base • amphoteric – capable of behaving as a Lewis acid and Brønsted base (17.5) • Neutralization • reaction in which mol acid = mol base • acid(aq) + base(aq) → salt(aq) + water(l) HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)

  10. HCl(aq) + NH3(aq) → NH4+(aq) + OH-(aq) • Conjugate acid/base pairs – reactant and product that differ by a single proton HA(aq) + H2O(l) → H3O+(aq) + A-(aq) acid + base conj. acid + conj. base

  11. Relative Strengths of Acids and Bases • Strength is a measure of the ability of an acid (or base) to donate (or accepts) a H+. • Stronger acids donate H+ more readily. • Completely dissociate in water • Conjugate bases have negligible tendency to accept protons; neutral. • Weaker acids donate H+ less readily. • Partially dissociate and establish equilibrium • Conjugate bases have some tendency to accept protons. • The stronger an acid, the weaker its conjugate base and vice versa.

  12. HA(aq) + H2O(l) → H3O+(aq) + A-(aq) HA(aq) + H2O(l)⇌ H3O+(aq) + A-(aq)

  13. p. 657

  14. Acid/base reactions proceed from the stronger acid-base pair to the weaker acid-base pair. • Common strong acids (p. 664): HClO4, HClO3, H2SO4, HI, HBr, HCl, HNO3 • Monoprotic acid – capable of donating only one H+ • Polyprotic acid – capable of donating more than one H+ • Common strong bases (p. 665): M(OH)n, where M = Group I (n=1) & heavier Group II (n=2) metals

  15. Acid/Base Reactions

  16. 16.3 The Autoionization of Water H2O(l) + H2O(l) ⇌ H3O+(aq) + OH-(aq) H2O(l) ⇌ H+(aq) + OH-(aq) • Kw = [H3O+][OH-] = [H+][OH-] = 1.0 x 10-14 (@ 25°C) • Kw – ion-product constant (or dissociation constant) • Pure water is neutral. Thus, [H3O+] = [OH-] = 1.0 x 10-7 M @ 25°C

  17. 16.3 The Autoionization of Water • For an aqueous solution:

  18. Working with Kw

  19. 16.4 The pH Scale • pH represents a solution’s acidity (@ 25°C). 0 ← 7 → 14 acidic neutral basic • See Table 16.1, p. 661 for summary. • See Figure 16.5, p. 663 for examples. • pH = −log[H3O+] = −log[H+] [H3O+] = 10-pH pOH = −log[OH-] pH + pOH = 14 [OH-] = 10-pOH

  20. p. 663

  21. More common chemicals *CaCO3 CO3- + H2O ⇌ HCO3- + OH- **CO2 + H2O → H2CO3

  22. pH calculations

  23. More about pH • pH does not necessarily indicate strength. • Measuring pH • pH meters – measures exact pH based on electrochemistry • Acid-base indicators – estimates pH based on the appearance of color

  24. p. 664

  25. Indicator Colors.

  26. 16.5 Strong Acids and Bases • Strong acids & bases completely dissociate. [HA]0 = [H3O+] → pH [MOH]0 = [OH-] → pOH → pH 2[M(OH)2]0 = [OH-] → pOH → pH • H3O+ is the strongest acid that can exist in water. (produced by all acids in water) • OH- is the strongest base that can exist in water. (produced by all bases in water)

  27. pH problems End Test #1 material

  28. 16.6 Weak Acids & 16.7 Weak Bases • Weak acids & bases do not completely dissociate. • Weak acids establish an equilibrium in aqueous solution. HA(aq) + H2O(l) ⇌ H3O+(aq) + A-(aq) HA(aq) ⇌ H+(aq) + A-(aq) • They do not readily donate or accept H+’s. • [HA]0≠ [H3O+] [MOH]0 ≠ [OH-]

  29. Weak Acids & Acid-dissociation Constant HA(aq) + H2O(l) ⇌ H3O+(aq) + A-(aq) HA(aq) ⇌ H+(aq) + A-(aq) Ka↑ acid strength ↑ For polyprotic acids: Ka1 >> Ka2 >>Ka3 pKa = −log[Ka]pKa↑ acid strength↓

  30. From p. 667 + more in Appendix D, p. 1062

  31. p. 674

  32. Weak Bases & Base-dissociation Constant Weak bases establish an equilibrium in aqueous solution. B(aq) + H2O(l) ⇌ BH+(aq) + OH-(aq) Kb↑ base strength ↑ pKb= −log[Kb]pKb↑ base strength↓

  33. From p. 674 + more in Appendix D, p. 1063

  34. % Dissociation (or ionization) • % dissociation increases as acid/base strength increases. (p. 669) • % dissociation decreases as concentration increases.

  35. Weak acid/base Problems1) Ka (or Kb) from equilibrium pH2) pH from Ka (or Kb) • Identify as weak acid or base. • Write the chemical equilibrium. • Write the equilibrium constant expression. • Set up concentration table. (Ch. 15.5) • Solve for x. • Check with 5% rule. If greater than 5%, use quadratic equation. (type 2 only) • Complete problem.

  36. The pH of a 0.10 M solution of propionicacid (CH3CH2CO2H) is 2.94. Calculate the Ka for propionic acid.

  37. Calculate the pH of a 1.0 M HF solution.

  38. Calculate the pH of a 0.0010 M HF solution.

  39. Calculate the pH of a 0.20 M solution of triethylamine N(CH2CH3)3.

  40. 16.8 Relationship between Ka and Kb • For a conjugate acid/base pair: Ka x Kb = Kw (derivation p. 679) Thus, at 25°C, Ka x Kb = 1.0 × 10-14 And, pKa+ pKb = pKw = 14.00

  41. 16.9 Acid-Base Properties of Salt Solutions • Salt – ionic compound • Salts dissolve in water to produce ions. • Ions can also affect the pH. • Hydrolysis – reaction between an ion and water to produce H3O+ or OH- F-(aq) + H2O(l) ⇌ HF(aq) + OH-(aq) NH4+(aq) + H2O(l) ⇌ H3O+(aq) + NH3(aq)

  42. Which ions will undergo hydrolysis, i.e. react with water and affect the pH of the solution? • Anion: • Conjugate base of a weak acid ► basic • Conjugate base of a monoprotic strong acid ► neutral • Cation: • Conjugate acid of a weak base ► acidic • Group I & II metal ions ► neutral (exceptions Be2+ and Mg2+ ► acidic) • Other metal ions ► acidic • See p. 683 for summary of combined effect.

  43. Effect on cations on solution pH

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