Buffer Solutions: MDCAT Chemistry notes
Buffer Solutions MDCAT notes: acidic and basic buffers, how buffers resist pH change, Henderson equation, blood pH, pH calculations and Bronsted acids.
What is a buffer?
A buffer solution resists a change in pH when a small amount of acid or base is added, or when it is diluted. It keeps the pH of a solution nearly constant.
| Type | Made from | Examples | pH |
|---|---|---|---|
| Acidic buffer | Weak acid + its salt with a strong base | $\mathrm{CH_3COOH/CH_3COONa}$, $\mathrm{HNO_2/NaNO_2}$, $\mathrm{C_2H_5COOH/C_2H_5COONa}$, $\mathrm{H_2CO_3/NaHCO_3}$ | Below 7 |
| Basic buffer | Weak base + its salt with a strong acid | $\mathrm{NH_4OH/NH_4Cl}$ | Above 7 |
A strong acid with its salt (HCl/NaCl, $\mathrm{HNO_3/NH_4NO_3}$) or a strong base with its salt (NaOH/NaCl) is not a buffer, because there is no weak partner to hold a reserve.
How a buffer works
Acetate buffer
$\mathrm{CH_3COOH \rightleftharpoons CH_3COO^- + H^+}$, with a large reserve of $\mathrm{CH_3COO^-}$ from the salt.
- Added $\mathrm{H^+}$ combines with $\mathrm{CH_3COO^-}$ to give undissociated $\mathrm{CH_3COOH}$. The pH scarcely changes.
- Added $\mathrm{OH^-}$ reacts with $\mathrm{CH_3COOH}$ to give water and acetate. The pH scarcely changes.
Ammonium buffer
$\mathrm{NH_4OH \rightleftharpoons NH_4^+ + OH^-}$. Added $\mathrm{H^+}$ is neutralized by $\mathrm{OH^-}$ (and hence by $\mathrm{NH_4OH}$), so more $\mathrm{NH_4OH}$ ionizes to replace it. Added $\mathrm{OH^-}$ combines with $\mathrm{NH_4^+}$ to form $\mathrm{NH_4OH}$, shifting this equilibrium in reverse.
Henderson equation
For an acidic buffer: $$\mathrm{pH} = \mathrm{p}K_a + \log\frac{[\text{salt}]}{[\text{acid}]}$$
For a basic buffer: $$\mathrm{pOH} = \mathrm{p}K_b + \log\frac{[\text{salt}]}{[\text{base}]}$$
- The pH of an acidic buffer depends on the concentration of acid, the concentration of salt and the $\mathrm{p}K_a$ of the acid.
- Raising the salt-to-acid ratio raises the pH. Adding extra $\mathrm{CH_3COONa}$ increases the pH.
- When $[\text{salt}] = [\text{acid}]$, pH = $\mathrm{p}K_a$; this is where buffer capacity is greatest.
Worked example
An acetic acid buffer with $\mathrm{p}K_a = 4.74$ has $[\text{salt}] = 0.20\ \mathrm{M}$ and $[\text{acid}] = 0.020\ \mathrm{M}$. Then pH $= 4.74 + \log 10 = 5.74$.
pH basics
- $\mathrm{pH} = -\log[\mathrm{H^+}]$, $\mathrm{pOH} = -\log[\mathrm{OH^-}]$, $\mathrm{pH + pOH} = 14$ at 25 °C.
- Example: $0.001\ \mathrm{M}$ NaOH gives pOH 3, pH 11.
- The larger the $K_a$, the stronger the acid.
- In the Bronsted-Lowry concept, water is amphoteric: it can donate or accept a proton.
Uses of buffers
- Human blood is buffered at pH 7.35 to 7.45, mainly by $\mathrm{H_2CO_3/HCO_3^-}$ (and phosphate and proteins).
- Calibration of pH meters, preserving biological specimens, fermentation, enzyme and industrial processes.
- A buffer does not measure the concentration of a substance.
Common MDCAT traps
- HCl + NaCl is not a buffer: HCl is a strong acid.
- The buffer with the highest pH (same acid) is the one with the largest salt-to-acid ratio.
- Adding a little acid to a buffer has almost no effect on pH; it does not "decrease the pH" noticeably.
- Henderson equation uses $\mathrm{p}K_a$ of the acid, not its pH.
Quick revision
- Acidic buffer: weak acid + salt of strong base.
- Basic buffer: $\mathrm{NH_4OH + NH_4Cl}$.
- Blood pH is 7.35 to 7.45.
- pH = $\mathrm{p}K_a$ when salt and acid concentrations are equal.
- Buffer action is based on the common ion effect.