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 $CH<em>3COONa$ is added to a solution of $CH</em>3COOH$.

Example Calculation: Common Ion Effect

• Calculate the pH of a 0.30 M $CH<em>3COOH$ solution. Also, calculate the pH of a solution containing 0.30 M $CH</em>3COOH$ and 0.20 M $CH<em>3COONa$. The $K</em>a$ 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 $NH<em>3$ and 0.2 M $NH</em>4Cl$? The $K<em>b$ for $NH</em>3 = 1.8 \times 10^{-5}$. To solve for pH, first find pOH with the Henderson-Hasselbalch equation, then convert
  $pOH= pK<em>b + log\frac{[NH</em>4^+]}{[NH_3]}$

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:
  • $CH<em>3COONa/CH</em>3COOH$
  • $KH<em>2PO</em>4/K<em>2HPO</em>4$
  • $HCN/KCN$

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 $HC<em>2H</em>3O<em>2$ and 0.1 mol of $KC</em>2H<em>3O</em>2$. The value of $K<em>a$ for $HC</em>2H<em>3O</em>2$ 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 $HC<em>2H</em>3O<em>2$ and add it to the moles of $KC</em>2H<em>3O</em>2$, 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. $HClO<em>2$ $pK</em>a = 1.95$
  • 2. $HNO<em>2$ $pK</em>a = 3.34$
  • 3. $HCHO<em>2$ $pK</em>a = 3.74$
  • 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.