What it is
A buffer is a solution containing a weak acid, HA, and its conjugate base, A⁻, in comparable amounts. Add a little acid and A⁻ takes up the hydrogen ions; add a little base and HA takes up the hydroxide ions. The addition is consumed by one member of the pair instead of changing the pH.
Examples of buffers
- Acetic acid and acetate. 0.10 mol acetic acid and 0.10 mol acetate in 1 dm³ give pH = pKa = 4.756 (pKa from the acid constants table), so about 4.76.
- The same mixture after 0.0010 mol hydrochloric acid. The acetate falls to 0.099 mol and the acetic acid rises to 0.101 mol, and the Henderson–Hasselbalch equation gives pH = 4.756 + log(0.099/0.101) ≈ 4.747. The pH moves by 0.009.
- The same amount of hydrochloric acid in 1 dm³ of pure water gives a concentration of 0.0010 mol/dm³ and pH 3, a change of about four pH units from 7.
- With added base. 0.0010 mol sodium hydroxide in the same buffer gives pH = 4.756 + log(0.101/0.099) ≈ 4.765.
Not buffers
- Pure acetic acid. It contains almost no acetate, so nothing is there to take up added acid.
- Hydrochloric acid, HCl. A strong acid has no conjugate base that can take hydrogen ions up again.
- Table salt in water. Neither Na⁺ nor Cl⁻ takes up or gives off hydrogen ions to any real extent.
The ratio of A⁻ to HA sets the pH. At 1:1, pH = pKa. At 10:1, pH = pKa + 1 (5.756 for acetic acid), and at 1:10, pH = pKa − 1 (3.756). That is why a buffer is said to work roughly within pKa ± 1.
Why it matters
Buffers are needed wherever pH has to stay put: in body fluids, in baking and brewing, in swimming pools, in laboratory solutions and in medicines. Buffer action is a regular topic in upper-secondary chemistry.
- The Henderson–Hasselbalch equation gives the pH from the pKa and the ratio of base to acid.
- The buffer capacity measures how much acid or base fits before the pH moves.
- The pKa decides which pair suits which pH. Choose a pair whose pKa is close to the pH you want to hold.
- What acids and bases are: acids and bases.
Common misconceptions
- “A buffer holds the pH perfectly constant.” It damps the change, it does not prevent it: in the example above the pH moves by 0.009 instead of 4.
- “A buffer can take any amount of acid.” It lasts until one member of the pair is used up. Add 0.15 mol HCl to 1 dm³ of the buffer with 0.10 mol of each, and the acetate runs out (0.10 mol), 0.05 mol HCl is left over, and the pH falls to about 1.3.
- “Diluting a buffer changes its pH.” The ratio of base to acid is the same after dilution, so the pH hardly changes, but the capacity drops.
- “Any acid can make a buffer.” The acid must be weak and stay next to its base. A strong acid has given up everything and does not form a buffer.
Common questions
Why does a buffer keep the pH constant?
Added acid meets A⁻, added base meets HA. Both reactions use up the addition and change the ratio A⁻/HA only
slightly, and the pH depends on that ratio, not on the amounts separately.
How do you make a buffer?
Mix a weak acid with its salt (for example acetic acid and sodium acetate), or a weak acid with just enough
strong base to turn part of it into its conjugate base. The ratio sets the pH.
What is the difference between buffer action and buffer capacity?
Buffer action is the pH staying nearly constant. Buffer capacity is a number for how much acid or base can
be added before the pH changes by a given amount.