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Acid-Base Equilibria

Acid-Base Equilibria. BLB 10 th Chapter 16. Examples of acids & bases. 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

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Acid-Base Equilibria

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  1. Acid-Base Equilibria BLB 10th Chapter 16

  2. Examples of acids & bases

  3. 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-

  4. 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.

  5. 16.2 Brønsted-Lowry Acids & Bases • Brønsted-Lowry definitions acid – proton donor • Neutral (HNO3), anionic (HCO3-), cationic (NH4+) • Must have a removable (acidic) proton base – proton acceptor • Neutral (NH3), anionic (CO32-) • Must have a lone pair of electrons

  6. Acid-Base Reactions HCl(aq) + H2O(l) → H3O+(aq) + Cl-(aq) NH3(aq) + H2O(l) ⇌ NH4+(aq) + OH-(aq) HCl(aq) + NH3(aq) → NH4+(aq) + Cl-(aq)

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

  8. 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: acid + base → salt + water • Conjugate acid/base pairs – differ by a single proton HA(aq) + H2O(l) → H3O+(aq) + A-(aq) acid + base conj. acid + conj. base

  9. Relative Acid/Base Strength • 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. • 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.

  10. p. 672

  11. Acid/base reactions proceed from the stronger acid-base pair to the weaker acid-base pair. • Common strong acids (p. 679): 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. 680): M(OH)n, where M = group I (n=1) & II (n=2) metals, except Be

  12. Acid/Base Reactions

  13. 16.3 The Autoionization of Water H2O(l) + H2O(l) ⇌ H3O+(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 • For an aqueous solution: [H3O+] > [OH-] acidic [H3O+] = [OH-] neutral [H3O+] < [OH-] basic

  14. Working with Kw

  15. 16.4 The pH Scale • pH represents a solution’s acidity (@ 25°C. 0 ← 7 → 14 acid neutral base • See Table 16.1, p. 678 for summary. • See Figure 16.5, p. 679 for examples. • pH = −log[H3O+] = −log[H+] [H3O+] = 10-pH pOH = −log[OH-] pH + pOH = 14 [OH-] = 10-pOH

  16. p. 676

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

  18. pH calculations

  19. More about pH • pH does not necessarily indicate strength. • Measuring pH • pH meters • Acid-base indicators

  20. p. 679

  21. 16.5 Strong Acids and Bases • Strong acids & bases completely ionize. [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)

  22. pH problems End Test #1 material

  23. 16.6 Weak Acids & 16.7 Weak Bases • Weak acids & bases do not completely ionize. • 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-]

  24. 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↓

  25. From p. 682 + more in Appendix D, p. 1115-1116

  26. Weak Bases & Base-dissociation Constant • Weak bases establish an equilibrium in aqueous solution. B(aq) + H2O(l) ⇌ BH+(aq) + OH-(aq)

  27. From p. 691 + more in Appendix D, p. 1115-1116

  28. % Dissociation (or ionization) • % dissociation decreases as concentration increases (p. 686)

  29. 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.

  30. The pH of a 0.10 M solution of propanoic acid (CH3CH2CO2H) is 2.94. Calculate the Ka for propanoic acid.

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

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

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

  34. 16.8 Relationship between Ka and Kb • For a conjugate acid/base pair: Ka x Kb = Kw (derivation p. 693) pKa + pKb = pKw = 14.00

  35. 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)

  36. 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

  37. Cation + Anion ►Acidic, basic, or neutral?

  38. 16.10 Acid-Base Behavior and Chemical Structure • Binary Acids (HX) • As bond strength increases, acid strength decreases. • Group: size of X ↑ acid strength ↓ • Period: electronegativity of X ↑ acid strength↑

  39. Oxyacids – acidic H attached to an oxygen atom • Same # of OH groups and O atoms: central atom electronegativity ↑ acid strength ↑ HClO > HBrO > HIO • Same central atom, Y: # O atoms ↑ acid strength ↑ HClO4 > HClO3 > HClO2 > HClO • Carboxylic acids – contain −COOH or CO2H • # electronegative atoms ↑ acid strength ↑

  40. Oxyacids

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