Your blood sits at pH 7.4 and stays there, even though you produce acid all day long. Shift it by a few tenths and you are seriously ill. Something is holding it steady, and that something is a buffer — the same flat shoulder you saw on the weak acid titration curve, put to work.
📘 What you need to know
A buffer resists changes in pH when small amounts of acid or alkali are added.
An acidic buffer is a weak acid plus its salt, for example CH3COOH with CH3COONa.
A basic buffer is a weak base plus its salt, for example NH3 with NH4Cl.
It works because there are large reserves of both halves of a conjugate pair in the same solution.
Add H+ and the A− reserve mops it up, shifting the equilibrium left.
Add OH− and it removes H+, so the HA reserve dissociates further, shifting the equilibrium right.
Buffer capacity is how much acid or alkali it can absorb before the pH shifts sharply.
Diluting a buffer barely changes its pH, because the ratio stays the same — but it reduces the capacity.
A strong acid can never make a buffer, because it has no undissociated reserve.
Making one
There are two standard recipes, and both end up with the same thing: plenty of a weak acid and plenty of its conjugate base sitting in the same beaker.
Mix them directly. Dissolve ethanoic acid and sodium ethanoate together. The acid is weak so most stays as CH3COOH; the salt is ionic so it fully dissociates and floods the solution with CH3COO−.
Half-neutralise a weak acid. Add strong alkali until about half the acid has reacted. What is left is half acid and half salt — which is exactly the half-equivalence point from the titration curve.
The two components
CH3COOH(aq) ↔ H+(aq) + CH3COO−(aq) (weak, so mostly left)
CH3COONa(aq) → Na+(aq) + CH3COO−(aq) (a salt, so fully split)
Notice that neither reserve is ever exhausted by a small addition. That is the whole trick — the ratio between the two shifts slightly, and the pH follows only slightly.
What happens when you attack it
Adding acid
Extra H+ arrives. By Le Chatelier the equilibrium shifts left to remove it, and the huge reserve of CH3COO− is there to react with it and form CH3COOH.
Because that reserve is large, using a bit of it up hardly dents the concentration. And because the acid reserve is also large, adding a bit to it hardly changes that either. The ratio barely moves, so [H+] barely moves, so the pH barely moves.
Adding alkali
OH− arrives and reacts with the H+ in solution to make water. That removes a product, so the equilibrium shifts right and more CH3COOH dissociates to replace the lost H+.
Both panels describe the same equilibrium. All that changes is which side is being disturbed, and therefore which way the system moves to oppose it.
Write “there is a large reserve of CH3COO−, so its concentration does not change much”. Answers that only say “the equilibrium shifts left” miss the point — a shift alone would not hold the pH steady if the reserve were tiny.
Basic buffers
Ammonia with ammonium chloride works the same way, one level up the pH scale. Ammonia is weak so most of it stays as NH3, and the salt fully dissociates to give plenty of NH4+.
A basic buffer at work
add acid: NH3(aq) + H+(aq) → NH4+(aq)
add alkali: NH4+(aq) + OH−(aq) → NH3(aq) + H2O(l)
In your bloodstream the main buffer is carbonic acid with hydrogencarbonate, H2CO3 and HCO3−. It is a weak acid with its conjugate base — exactly the pattern above — and it holds blood at pH 7.35 to 7.45.
Capacity, dilution and temperature
Capacity is limited. Add enough acid to use up all the A− and there is nothing left to absorb the next drop, so the pH collapses. A buffer only handles small additions.
Dilution hardly moves the pH, because both concentrations are cut by the same factor and the ratio is unchanged. But it does lower the capacity, since there are fewer particles of each reserve.
Temperature changes Ka, and Ka sets the pH, so a buffer must be kept at a steady temperature to stay reliable.
Worked examples
WORKED EXAMPLE
Which of these makes a buffer? (a) HCl + NaCl (b) CH3COOH + CH3COONa (c) NH3 + NH4Cl (d) NaOH + NaCl
Step 1: test each for a weak acid or weak base plus its salt(a) HCl is strong, so there is no undissociated reserve. No.(b) weak acid with its salt. Yes.(c) weak base with its salt. Yes.(d) NaOH is strong and NaCl is neutral. No.Step 2: say why (a) fails, since it looks plausible
HCl is fully dissociated, so nothing is left to release more H+. Cl− is far too weak a base to absorb any either.
Only (b) and (c) are buffersthe test is always “weak partner plus its conjugate”, never “acid plus salt”
WORKED EXAMPLE
Explain, with an equation, how a CH3COOH / CH3COONa buffer resists a change in pH when a little hydrochloric acid is added.
Step 1: say what the buffer contains
A large reserve of CH3COOH from the weak acid, and a large reserve of CH3COO− from the salt.
Step 2: write what the added H+ doesCH3COO−(aq) + H+(aq) → CH3COOH(aq)
The equilibrium shifts left to remove the added H+.
Step 3: explain why the pH holds
The ethanoate reserve is large, so removing a little of it changes its concentration only slightly, and the same is true for the acid it forms.
The ratio of acid to salt barely shifts, so [H+] and the pH barely shiftthree marks here: the reserves, the equation, and the ratio staying nearly constant
WORKED EXAMPLE
A student dilutes a buffer with an equal volume of water. Predict the effect on its pH and on its buffer capacity.
Step 1: see what dilution does to each concentrationBoth [acid] and [salt] are halved.Step 2: look at what the pH actually depends on
It depends on the ratio of the two, and halving both leaves that ratio unchanged.
Step 3: now think about capacity
There are only half as many moles of each reserve in every dm3, so less acid or alkali can be absorbed before they run out.
The pH stays almost the same, but the buffer capacity falls“pH depends on the ratio, capacity depends on the amounts” is the sentence to remember
💡 Exam tip
Always name both components: the weak acid and its salt. One alone is not a buffer.
Use the phrase “large reserve” in explanations. It is usually a mark on its own.
Give the equation for the ion that reacts, not just a statement about the equilibrium shifting.
Keep pH and capacity separate: pH comes from the ratio, capacity from the amounts.
A buffer only copes with small additions. Say so — it is often part of the definition mark.
Half-neutralising a weak acid gives a buffer with pH = pKa, which links straight back to the titration curve.
⚠ Common mix-up
Thinking a buffer keeps the pH exactly constant. It only keeps the change small.
Trying to build one from a strong acid. There is no undissociated reserve, so it cannot work.
Saying dilution changes the pH a lot. The ratio is untouched, so the pH barely moves.
Confusing capacity with pH. Two buffers can share a pH and have completely different capacities.
Forgetting the salt fully dissociates. That is where the large reserve of A− comes from.
Writing the buffer equation with a reversible arrow. When you add H+ to A−, that reaction goes essentially to completion.
Up next: Buffer Calculations. You can now explain why a buffer holds its pH. The last page of the topic puts a number on it, and it turns out to be one line of algebra you have already half derived.
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