Salts in solution
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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

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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


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