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Salts in Solution. Hydrolysis and Buffers. Introduction. Strong acids added to water produce a weak conjugate base . HCl(g) + H 2 O(l) ➜ Cl - (aq) + H 3 O + (aq). strong acid. weak base. Strong bases added to water produce a weak conjugate acid . NaOH(s) ➜ Na + (aq) + OH - (aq).

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salts in solution
Salts in Solution
  • Hydrolysis and Buffers
introduction
Introduction
  • Strong acids added to water produce a weak conjugate base.
      • HCl(g) + H2O(l) ➜ Cl-(aq) + H3O+(aq)

strong

acid

weak

base

  • Strong bases added to water produce a weak conjugate acid.
      • NaOH(s) ➜ Na+(aq) + OH-(aq)

strong

base

weak

acid

introduction1
Introduction
  • Weak acids added to water produce a relatively strong conjugate base.
      • HNO2(aq) + H2O(l) ➜ NO2-(aq) + H3O+(aq)

weak

acid

strong

base

  • Weak bases added to water produce a relatively strong conjugate acid.
      • NH3(aq) + H2O(l) ➜ NH4+(aq) + OH-(aq)

weak

base

strong

acid

salt hydrolysis
Salt Hydrolysis
  • When acids and bases become involved in neutralizations, they form salts and water.
      • HNO2(aq) + NaOH(aq) ➜ NaNO2(aq) + H2O(l)

nitrous

acid

sodium

hydroxide

sodium

nitrite

water

  • HCl(aq) + NH3(aq) ➜ NH4Cl(aq)

hydrochloric

acid

ammonium

chloride

ammonia

  • HCl(aq) + NaOH(aq) ➜ NaCl(aq) + H2O(l)

hydrochloric

acid

sodium

hydroxide

sodium

chloride

water

salt hydrolysis1
Salt Hydrolysis
  • When the salts themselves are dissolved in water, they hydrolyze.
      • NaNO2(aq) ➜ Na+(aq) + NO2-(aq)

sodium

nitrite

sodium

nitrite

  • NH4Cl(aq) ➜ NH4+(aq) + Cl-(aq)

ammonium

chloride

ammonium

chloride

  • NaCl(aq) ➜ Na+(aq) + Cl-(aq)

sodium

chloride

sodium

chloride

salt hydrolysis2
Salt Hydrolysis
  • The ions that are weak conjugate acids and bases have no other effect on the solution.
salt hydrolysis3
Salt Hydrolysis
  • The ions that are relatively strong conjugate acids and bases have effects on the solution.
salt hydrolysis4
Salt Hydrolysis
  • Strong conjugate acids hydrolyze in solution, donate hydrogen ions, and lower the pH.
    • NH4+(aq) + H2O(l) ➜ NH3(aq) + H3O+(aq)
  • Strong conjugate bases hydrolyze in solution, accept hydrogen ions, and raise the pH.
    • CH3COO-(aq) + H2O(l) ➜ CH3COOH(aq) + OH-(aq)
salt hydrolysis5
Salt Hydrolysis
  • If we have a salt which results from the neutralization of a strong acid and a strong base
    • the resulting solution is neutral.
      • NaCl(aq) ➜ Na+(aq) + Cl-(aq)
          • Na+(aq) + H2O(l) ➜ no reaction
          • Cl-(aq) + H2O(l) ➜ no reaction
salt hydrolysis6
Salt Hydrolysis
  • If we have a salt which results from the neutralization of a strong acid and a weak base
    • the resulting solution is acidic.
      • NH4Cl(aq) ➜ NH4+(aq) + Cl-(aq)
          • NH4+(aq) + H2O(l) ➜ NH3(aq) + H3O+(aq)
          • Cl-(aq) + H2O(l) ➜ noreaction
salt hydrolysis7
Salt Hydrolysis
  • If we have a salt which results from the neutralization of a weak acid and a strong base
    • the resulting solution is basic.
      • NaClO(aq) ➜ Na+(aq) + ClO-(aq)
          • Na+(aq) + H2O(l) ➜ noreaction
          • ClO-(aq) + H2O(l) ➜ HClO(aq) + OH-(aq)
salt hydrolysis8
Salt Hydrolysis
  • If we have a salt which results from the neutralization of a weak acid and a weak base
    • the resulting solution may be acidic, basic, or neutral.
  • It depends on the relative strengths of the acid and base.
      • NH4ClO(aq) ➜ NH4+(aq) + ClO-(aq)
          • NH4+(aq) + H2O(l) ➜ NH3(aq) + H3O+(aq)
          • ClO-(aq) + H2O(l) ➜ HClO(aq) + OH-(aq)
buffers
Buffers
  • A buffer is a solution in which the pH remains relatively constant when small amounts of acid or base are added.
  • A buffer is prepared with a solution of
      • a weak acid and one of its salts
          • CH3COOH and NaCH3COO
      • a weak base and one of its salts
          • NH3 and NH4Cl
buffers1
Buffers
  • Buffers are better able to resist pH changes than is pure water.
  • Add 10 mL of 0.1 M HCl to 50 mL of
      • pure water
          • pH goes from 7.00 to 1.78 (∆pH = 5.22)
      • acetic acid/acetate buffer
          • pH goes from 4.74 to 4.57 (∆pH = 0.18)
buffers2
Buffers
  • The equilibrium set up between the acetic acid (CH3COOH) and its acetate salt (CH3COO-) allows the solution to absorb excess acid or base.
      • H3O+ + CH3COO- ⇄ CH3COOH + H2O
      • OH- + CH3COOH ⇄ CH3COO- + H2O
  • The concentrations of the acid and the salt act as reservoirs of neutralizing power.
buffers3
Buffers
  • A buffer cannot control pH when too much acid or base is added.
      • The reservoirs of neutralizing power are used up.
  • When this happens, we exceed the buffering capacity of the system.
  • Our bodies keep blood at pH = 7.35-7.45 using
      • carbonic acid/hydrogen carbonate
      • dihydrogen phosphate/hydrogen phosphate