Titration Notes

Titration

• Analytical technique to determine an unknown concentration of a solution by reacting it with a known concentration.
• In an acid–base titration, a solution of unknown concentration (titrant) is slowly added to a solution of known concentration from a burette until the reaction is complete.
  • When the reaction is complete, we have reached the endpoint of the titration.
  • An indicator may be added to determine the endpoint -- a chemical that changes color when the pH changes.
• When the moles of $H3O+$ = moles of $OH−$, the titration has reached its equivalence point.

Titration Curve

• Plot of pH versus the amount of added titrant.
• The inflection point of the curve is the equivalence point of the titration.
• Prior to the equivalence point, the known solution in the flask is in excess, so the pH is closest to its pH.
• The pH of the equivalence point depends on the pH of the salt solution.
  • Equivalence point of neutral salt: pH = 7
  • Equivalence point of acidic salt: pH < 7
  • Equivalence point of basic salt: pH > 7
• Beyond the equivalence point, the unknown solution in the burette is in excess, so the pH approaches its pH.

Strong Acid-Strong Base Titration

• Neutralization Reaction: $HCl + NaOH \rightarrow H2O + NaCl$
• Net ionic equation: $H3O^+ (aq) + OH^- (aq) \rightarrow 2 H2O(l)$

Calculating pH

• Starting pH: $pH= -log[H3O^+]$
• Before Equivalence:
  • Calculate moles of $H3O^+$ and $OH^-$
  • Subtract moles of $OH^-$ from moles of $H3O^+$ to find remaining $H3O^+$
  • Calculate new molarity of $H3O^+$
  • Calculate pH: $pH = -log[H3O^+]$
• Equivalence Point:
  • Moles of $H3O^+$ = moles of $OH^-$
  • pH = 7
• After Equivalence:
  • Calculate moles of $OH^-$ in excess
  • Calculate new molarity of $OH^-$
  • Calculate pOH: $pOH = -log[OH^-]$
  • Calculate pH: $pH = 14 - pOH$

Weak Acid-Strong Base Titration

• Neutralization Reaction: $CH3CO2H + NaOH \rightarrow H2O + NaCH3CO2$
• Net ionic equation: $CH3CO2H(aq) + OH^- (aq) \rightarrow H2O(l) + CH3CO2^- (aq)$

Calculating pH

• Starting pH: Use an ICE table to find equilibrium concentrations and calculate $[H3O^+]$ and pH.
• Before Equivalence:
  • Use stoichiometry to determine the remaining moles of weak acid and the moles of conjugate base formed.
  • Use the Henderson-Hasselbalch equation to calculate pH: $pH = pKa + log \frac{[A^-]}{[HA]}$
• Half-Equivalence Point: pH = pKa
• Equivalence Point:
  • All of the weak acid has been converted to its conjugate base. Use an ICE table to find equilibrium concentrations and calculate
  • $[OH^-]$ then calculate pOH and pH.
• After Equivalence:
  • The pH is determined by the excess $OH^-$ from the strong base. Calculate $[OH^-]$ and then calculate pOH and pH as with a strong acid/strong base titration after equivalence.

Weak Base/Strong Acid Titration

• Starting pH: pH of a weak base (ICE Table)
• Before Equivalence:
  • Moles table
  • pH of a buffer – ICE Table or HH equation
  • Must Test x with HH equation
• Half- Equivalence:
  • Moles table
  • pH of a buffer – ICE Table or HH equation
  • pH=pKa because concentration of B and HB+ are equal
• Equivalence Point:
  • Moles table
  • pH of a salt – ICE Table ONLY
  • pH will be acidic – only conj. acid left
• After Equivalence:
  • Moles table
  • pH of a strong acid

Acid-Base Indicators

• Indicators are weak acids where the weak acid is a different color than its conjugate base.
• The pH at which the indicator changes color is called the end point.
• Choose an indicator with an end point that coincides with the pH at the equivalence point.