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Add a single drop of concentrated hydrochloric acid to a litre of pure water and the pH plunges from 7 to about 3. Add the same drop to a litre of blood and the pH barely moves. Blood isn’t special because it lacks acid; it’s special because it’s buffered.
What a buffer is
A buffer is a solution that resists changes in pH when small amounts of acid or base are added. It contains two partners in similar amounts:
- a weak acid and its conjugate base, or
- a weak base and its conjugate acid.
Classic examples:
- Acetic acid (CH₃COOH) + sodium acetate (CH₃COO⁻)
- Ammonia (NH₃) + ammonium chloride (NH₄⁺)
- Carbonic acid (H₂CO₃) + hydrogen carbonate (HCO₃⁻) — the main buffer in blood
- Dihydrogen phosphate (H₂PO₄⁻) + hydrogen phosphate (HPO₄²⁻) — important inside cells
How it works
Think of the two partners as a pair of sponges, one for each kind of threat.
Add acid (H⁺): the conjugate base soaks it up. CH₃COO⁻ + H⁺ → CH₃COOH
The strong acid is replaced by a weak acid, which barely ionizes, so [H⁺] hardly rises.
Add base (OH⁻): the weak acid neutralizes it. CH₃COOH + OH⁻ → CH₃COO⁻ + H₂O
The strong base is replaced by a weak base, so [OH⁻] hardly rises either.
In both cases the added acid or base is converted into more of one buffer partner. The ratio of the two partners changes only slightly, and as we’ll see, pH depends on that ratio.
Why a strong acid can’t make a buffer
HCl with NaCl isn’t a buffer. Chloride ions are such a weak base (see strong vs. weak acids) that they won’t take H⁺ back. You need a weak acid, whose conjugate base actually grabs protons.
The Henderson-Hasselbalch equation
For a buffer made from a weak acid (HA) and its conjugate base (A⁻):
pH = pKa + log([A⁻] ÷ [HA])
Three useful consequences:
- When [A⁻] = [HA], the log term is zero and pH = pKa. So you choose a weak acid whose pKa is close to the pH you want.
- pH depends on the ratio of the two partners, not their absolute amounts. Diluting a buffer with water barely changes its pH.
- A buffer works best within about ±1 pH unit of its pKa, where neither partner is running low.
Worked example 1: pH of a buffer
A buffer contains 0.20 M acetic acid and 0.30 M sodium acetate. Acetic acid’s pKa is 4.74.
pH = 4.74 + log(0.30 ÷ 0.20) = 4.74 + 0.18 = 4.92
Worked example 2: adding acid
0.010 mol of HCl is added to 1.0 L of that buffer.
The acetate absorbs the H⁺:
- Acetate: 0.30 − 0.010 = 0.29 mol
- Acetic acid: 0.20 + 0.010 = 0.21 mol
pH = 4.74 + log(0.29 ÷ 0.21) = 4.74 + 0.14 = 4.88
The pH fell by only 0.04. Adding the same 0.010 mol of HCl to 1.0 L of pure water would take it from pH 7.00 to pH 2.00 — a change of five units.
Buffer capacity
A buffer can only absorb so much. Once one partner is used up, the protection is gone and the pH swings sharply. The amount of acid or base a buffer can take is its buffer capacity, and it’s larger when:
- the concentrations of both partners are higher, and
- the two partners are present in roughly equal amounts.
A buffer made of 1.0 M acid and 1.0 M conjugate base has the same pH as one made of 0.01 M of each, but can absorb a hundred times more.
The buffer in your blood
Blood is held between about pH 7.35 and 7.45. The main buffer is the carbonic acid–hydrogen carbonate system, linked to the carbon dioxide you breathe out:
CO₂ + H₂O ⇌ H₂CO₃ ⇌ H⁺ + HCO₃⁻
This buffer is unusual because the body can adjust both ends: your lungs control CO₂ by changing your breathing rate, and your kidneys control HCO₃⁻. Hyperventilating blows off too much CO₂ and pushes blood pH up; conditions that stop you breathing effectively let CO₂ build up and push it down.
Other everyday buffers
- Seawater is buffered by the carbonate system around pH 8.1. As the ocean absorbs more atmospheric CO₂, that buffer is slowly being used up — which is the chemistry behind ocean acidification.
- Shampoos, cosmetics and medicines are buffered to stay gentle on skin and stable on the shelf.
- Laboratories use buffers like phosphate-buffered saline to keep enzymes and cells working.
Quick answers
Does a buffer keep pH exactly constant? No. It keeps the change small, up to its capacity.
What’s the best pH for an acetate buffer? Around its pKa, 4.74 — usable from roughly pH 3.7 to 5.7.
Can I make a buffer by partly neutralizing a weak acid? Yes. Adding half as many moles of NaOH as there are moles of weak acid converts half of it to the conjugate base, giving pH = pKa.
Related
The pH calculator finds the pH of weak acids and bases on their own; for the background, see the pH scale explained and Le Chatelier’s principle.
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