Chapter 16 Flashcards

16.1-16.3 Common Ion Effect, Buffer Solutions

Common Ion Effect

• The presence of a common ion suppresses the ionization of a weak acid or weak base.
• This phenomenon can be understood through Le Chatelier’s Principle.
• For example, consider what happens to the H^+ concentration when CH3COONa is added to a solution of CH3COOH.

Example Calculation: Common Ion Effect

• Calculate the pH of a 0.30 M CH3COOH solution. Also, calculate the pH of a solution containing 0.30 M CH3COOH and 0.20 M CH3COONa. The Ka of CH_3COOH is 1.8 \times 10^{-5}.

Henderson-Hasselbalch Equation

• The Henderson-Hasselbalch equation simplifies solutions involving the common ion effect.

• It assumes that the [H^+] is small compared to the initial concentration of the acid/conjugate base.

  pH = pK_a + log \frac{[A^-]}{[HA]}

Example Calculation : Henderson-Hasselbalch Equation

• What is the pH of a solution containing 0.5 M NH3 and 0.2 M NH4Cl? The Kb for NH3 = 1.8 \times 10^{-5}. To solve for pH, first find pOH with the Henderson-Hasselbalch equation, then convert
  pOH= pKb + log\frac{[NH4^+]}{[NH_3]}

Buffer Solutions

• Buffers are chemical systems that resist pH changes by neutralizing added acid or base.
• A buffer contains significant amounts of both a weak acid and its conjugate base (or a weak base and its conjugate acid).
  • The weak acid neutralizes added base.
  • The conjugate base neutralizes added acid.
• Examples:
  • CH3COONa/CH3COOH
  • KH2PO4/K2HPO4
  • HCN/KCN

Action of a Buffer

• Buffers work through equilibrium. When acid is added, it shifts towards the production of HA. When base is added, it shifts towards the production of A^-.

Calculating pH Changes in Buffers

• A 2.0L buffer solution contains 0.1 mol of HC2H3O2 and 0.1 mol of KC2H3O2. The value of Ka for HC2H3O2 is 1.8 \times 10^{-5}. Calculate the pH of the buffer:

  • First calculate the pH of the buffer before adding KOH. This can be done using the Henderson-Hasselbalch equation:
    pH = pK_a + log \frac{[A^-]}{[HA]}

• 0. 010 moles of solid KOH are added to the buffer. Calculate the new pH of the buffer:

  • Then, calculate the new pH of the buffer after adding KOH. Because you are adding a strong base to a weak acid, this will change the concentration. Treat it like a limiting reagent, so subtract the 0.01 moles from the moles of HC2H3O2 and add it to the moles of KC2H3O2, then do the Henderson-Hasselbalch equation again with the new values.

• For comparison, calculate the pH of a 0.01 M solution of just KOH in pure water:

  • First, since [OH^-] = 0.01M, pOH=-log(0.01)=2, so pH= 14-2=12

Buffer Effectiveness

• Buffer Capacity: The amount of acid or base it can effectively neutralize.
  • Most effective when the concentrations of acid and conjugate base are equal.
  • Most effective when the concentrations of acid and conjugate base are high.
• Example: A 1.0L buffer solution is 1.0 M in HF and 0.050 M in NaF. Which action will destroy the buffer?
  • A) adding 0.05 mol of HCl
  • B) adding 0.05 mol of NaF
  • C) adding 0.050 mol of NaOH
    • The correct answer is C) adding 0.050 mol of NaOH because NaOH will neutralize HF, turning it into F^-. Because F^- is already at 0.050 M, adding 0.050 mol of NaOH will overwhelm the buffering system.
  • D) None of the above

Buffer Range

• Buffer Range: The pH range over which a particular acid and conjugate base can be effective.
• The effective range for a buffering system is generally one pH unit on either side of the pK_a.

Preparing a Buffer

• Which acid would you choose to combine with its sodium salt to make a solution buffered at pH 4.25?
  • 1. HClO2 pKa = 1.95
  • 2. HNO2 pKa = 3.34
  • 3. HCHO2 pKa = 3.74
  • 4. HClO pK_a = 7.54
    • The best choice would be #3 since it is closest to 4.25
• For the best choice, calculate the ratio of the conjugate base to the acid required to attain the desired pH.