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Acid-Base Equilibrium 1

Acid-Base Equilibrium 1. Acid-Base Theory and Equilibrium Constants. AP Chemistry. Models of Acids and Bases. Arrhenius Concept : Acids produce H + (or H 3 O + , hydronium ion) in solution; bases produce OH − ion. HCl + H 2 O Cl − + H 3 O + acid base.

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Acid-Base Equilibrium 1

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  1. Acid-Base Equilibrium 1 Acid-Base Theory and Equilibrium Constants AP Chemistry

  2. Models of Acids and Bases • Arrhenius Concept: Acids produce H+ (or H3O+, hydronium ion) in solution; bases produce OH− ion. • HCl + H2O Cl− + H3O+ • acid base • Brønsted-Lowry: Acids are proton (H+ ) donors bases are proton acceptors. 2

  3. Conjugate Acid-Base Pairs • HA(aq) + H2O(l) H3O+(aq) + A−(aq) conj conj conj conjacid 1 base 2 acid 2 base 1 • conjugate base: everything that remains of the acid molecule after a proton is lost. • conjugate acid: formed when the proton is transferred to the base. 3

  4. Acid Dissociation Constant, Ka [H+][A -] [H3O+][A-] Ka = = [HA] [HA] • HA(aq) + H2O(l) H3O+(aq) + A−(aq) 4

  5. Acid Strength • Its equilibrium position lies far to the right. (HNO3) • Large Ka value • Yields a weak conjugate base. (NO3−) • HNO3 + H2O H3O+ + NO3− Strong Acid: 5

  6. Acid Strength • Its equilibrium lies far to the left. (CH3COOH) • Small Ka value • Yields a much stronger (it is relatively strong) conjugate base than water. (CH3COO−) • CH3COOH + H2O H3O+ + CH3COO− Weak Acid: 6

  7. Base Dissociation Constant, Kb [OH-][HB+] Kb = [B] • B(aq) + H2O(l) OH-(aq) + HB+(aq) 7

  8. Bases • “Strong” and “weak” are used in the same sense for bases as for acids. • strong = complete dissociation (hydroxide ion supplied to solution) • NaOH(s) → Na+(aq) + OH−(aq) 8

  9. Bases (con’t) • weak = very little dissociation (or reaction with water) • H3CNH2(aq) + H2O(l) H3CNH3+(aq) + OH−(aq) 9

  10. Water is amphoteric. • Water can behave either as an acid or a base. • H2O + H2O H3O+ + OH-conj conjacid 1 base 2 acid 2 base 1 Kw = [H+] [OH−] Kw = 1 × 10−14 at 25°C 10

  11. Acid-Base Equilibrium 2 The pH Scale AP Chemistry

  12. The pH Scale • pH ≈−log[H+] and pOH = −log[OH-] • pH in water ranges from 0 to 14. • Kw = 1.00 × 10−14 = [H+] [OH−] • pKw = 14.00 = pH + pOH • As pH rises, pOH falls (sum = 14.00). 12

  13. pH =−log[H+] HNO3 + H2O H3O+ + NO3− pH = −log(0.0020) pH =2.70 Example Problem What is the pH, pOH, and [OH-] of a 0.0020M solution of nitric acid? pH + pOH = 14.002.70 + pOH = 14.00pOH = 11.30 [H+] [OH−] = 1.00 × 10−14(0.0020) [OH−] = 1.00 × 10−14[OH−] = 5.0 × 10−12 M

  14. Acid-Base Equilibrium 3 Solving Problems AP Chemistry

  15. Solving Weak Acid Equilibrium Problems • List major species in solution. • Choose species that can produce H+ and write reactions. • Based on K values, decide on dominant equilibrium. • Write equilibrium expression for dominant equilibrium. • List initial concentrations in dominant equilibrium. 15

  16. (continued) • Define change at equilibrium (as “x”). • Write equilibrium concentrations in terms of x. • Substitute equilibrium concentrations into equilibrium expression. • Solve for x the “easy way.” • Calculate [H+] and pH. 16

  17. Amount dissociated (M) x 100% % dissociation = Initial concentration (M) % Dissociation or % Ionization 17

  18. Acid-Base Equilibrium 4 Ka and Kb Relationships &Polyprotic Acids AP Chemistry

  19. [H3O+][A-] Ka x Kb [OH-][HA] = Kw = [H3O+] [OH-] [HA] [A] Ka and Kb Relationship • HA(aq) + H2O(l) H3O+(aq) + A−(aq) Ka A-(aq) + H2O(l) OH-(aq) + HA(aq) Kb • H2O(l) + H2O(l) H3O+(aq) + OH-(aq)Kw

  20. H2CO3 H+ + HCO3- Ka1 Polyprotic Acids • . . . can furnish more than one proton (H+) to the solution. HCO3- H+ + CO32- Ka2 H2CO3 2H+ + CO32- Ka Ka1 x Ka2 = Ka(overall) 20

  21. Ka and Kb Relationship for Polyprotic Acids For triprotic acids Ka1 x Kb3 = Kw Ka2 x Kb2 = Kw Ka3 x Kb1 = Kw For diprotic acids Ka1 x Kb2 = Kw Ka2 x Kb1 = Kw

  22. Acid-Base Equilibrium 5 Miscellaneous Acid-Base Concepts AP Chemistry

  23. Acidic and Basic Salts 23

  24. Molecular Structure and Acidic Properties • Two factors affecting acidity in binary compounds • Bond Polarity (higher is more acidic) • Bond Strength (lower is more acidic) 24

  25. Oxides of Metal and Nonmetals • 1. Acidic Oxides (Acid Anhydrides): • Nonmetal Oxides (Examples: SO2, CO2, CrO3) • O−X bond is strong and covalent. • 2. Basic Oxides (Basic Anhydrides): • Metal Oxides (Examples: K2O, CaO) • O−X bond is ionic. 25

  26. Examples of Anhydrides • 1. Acidic Oxides (Acid Anhydrides): • H2O + CO2 H2CO3 • N2O5 + H2O 2HNO3 • 2. Basic Oxides (Basic Anhydrides): • K2O + H2O 2KOH • CaO + H2O Ca(OH)2 26

  27. Coordinate Covalent Bond Lewis Acids and Bases • Lewis Acid: electron pair acceptor • Lewis Base: electron pair donor H H F F B +N F B N H : : F H F Lewis LewisAcid Base F H H 27

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