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 effect is explained by Le Chatelier’s Principle.
• Example: When CH3COONa is added to a solution of CH3COOH, the H^+ concentration is affected.

Common Ion Effect - Calculation Example

• Problem: Calculate the pH of a 0.30 M CH3COOH solution. Then, 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.

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

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

Henderson-Hasselbalch Equation - Example

• Problem: 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}.

Buffer Solutions

• Buffers: 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

Additions to Buffer Solutions

• Action of a Buffer
  • After addition of H^+, the equilibrium shifts to consume the added acid, maintaining a stable pH.
  • After addition of OH^-, the equilibrium shifts to consume the added base, maintaining a stable pH

Calculating pH Changes in Buffers

• Problem: A 2.0 L buffer solution contains 0.1 mol of HC2H3O2 and 0.1 mol of KC2H3O2. The Ka for HC2H3O2 is 1.8 \times 10^{-5}. Calculate the pH of the buffer.
• If 0.010 moles of solid KOH are added to the buffer, calculate the new pH of the buffer.
• For comparison, calculate the pH of a 0.01 M solution of just KOH in pure water.

Buffer Effectiveness

• Buffer Capacity: The amount of acid or base a buffer 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.
• Problem: A 1.0 L 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
  • 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

• Problem: 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
• For the best choice, calculate the ratio of the conjugate base to the acid required to attain the desired pH.