Factors Affecting Buffer pH
Note- In this article, we use Henderson-Hasselbalch equation to reason the intuition behind buffer pH.
- In IB Chemistry (First Examination 2025), you don't have to remember this equation.
- However, according to the syllabus, it is essential to understand that the pH of a buffer solution depends on both:
- the $pK_a$ or $pK_b$ of its acid or base
- the ratio of the concentration of acid or base to the concentration of the conjugate base or
acid.
The Henderson-Hasselbalch Equation: A Tool for Predicting Buffer pH
- As discussed earlier, buffers consist of a weak acid and its conjugate base (or a weak base and its conjugate acid), which work together to resist pH changes.
- The precise pH of a buffer solution can be calculated using the Henderson-Hasselbalch equation: $$\text{pH} = \text{pKa} + \log{\frac{[\text{A}^-]}{[\text{HA}]}}$$ where:
- $[A^-]$: Concentration of the conjugate base.
- $[HA]$: Concentration of the weak acid.
- $\text{pKa}$: The negative logarithm of the acid dissociation constant $K_a$ of the weak acid.
This equation highlights two critical factors influencing buffer pH:
- The $ \text{pKa} $ of the weak acid.
- The ratio of the concentrations of the conjugate base $[A^-]$ to the weak acid $[HA]$.
The Role of $\text{pKa}$
- The $ \text{pKa} $ of a weak acid is a measure of its tendency to donate protons $ \text{H}^+ $.
- A lower $ \text{pKa} $ corresponds to a stronger weak acid, while a higher $ \text{pKa} $ indicates a weaker acid.
Buffers are most effective when their pH is close to the $ \text{pKa} $ of the weak acid, which occurs when $[A^-] = [HA]$.Example
Selecting a Buffer for pH 4.8
- To prepare a buffer with a pH of 4.8, choose a weak acid with a $ \text{pKa} $ value near 4.8, such as ethanoic acid $ \text{pKa} = 4.76 $.
- By adjusting the ratio of $[A^-]$ (ethanoate ion) to $[HA]$ (ethanoic acid), you can fine-tune the pH to 4.8.
