How to Prepare a Buffer Solution

Knowing the theory of buffers is one thing. Standing at a lab bench with a bottle of sodium ethanoate, a pH meter and a target of pH 5.00 is another. This guide covers how buffers are actually made in school, university and research labs, from the calculation to the final adjustment.

If you need a refresher on the equation used throughout, see the Henderson–Hasselbalch equation.

Before you start: choose the buffer system

1. Decide the target pH and the volume you need.
2. Choose a weak acid with pKa within about ±1 of the target. Closer is better, because capacity is highest when pH = pKa. See buffer capacity.
3. Decide the total buffer concentration, based on how much acid or base the buffer will have to absorb. 0.05–0.2 mol/dm³ is typical for school and general lab work.
4. Check that the buffer won’t interfere with whatever you’re using it for.

For this guide, the running example is 250 cm³ of an ethanoate buffer at pH 5.00 (ethanoic acid pKa = 4.76).

Method 1: weak acid + its salt

Mix a solution of the weak acid with a solution (or solid) of its sodium or potassium salt.

Calculation

Target ratio: [A⁻] ÷ [HA] = 10^(5.00 − 4.76) = 10⁰·²⁴ = 1.74

Suppose we want a total buffer concentration of 0.100 mol/dm³ in 250 cm³, so total moles = 0.0250 mol.

• Moles HA + moles A⁻ = 0.0250
• Moles A⁻ = 1.74 × moles HA
• So 2.74 × moles HA = 0.0250 → moles HA = 9.12 × 10⁻³ mol
• Moles A⁻ = 0.0250 − 0.00912 = 1.588 × 10⁻² mol

Using stock solutions of 0.100 mol/dm³ ethanoic acid and 0.100 mol/dm³ sodium ethanoate:

• Volume of acid = 9.12 × 10⁻³ ÷ 0.100 = 91.2 cm³
• Volume of salt solution = 1.588 × 10⁻² ÷ 0.100 = 158.8 cm³
• Total = 250.0 cm³ ✔

Using solid sodium ethanoate (anhydrous, M = 82.03 g/mol) with ethanoic acid:

• Mass of sodium ethanoate = 1.588 × 10⁻² × 82.03 = 1.30 g
• Ethanoic acid from a 1.00 mol/dm³ stock: 9.12 cm³
• Make up to 250 cm³ with distilled water

If you’re using the trihydrate, CH₃COONa·3H₂O (M = 136.08 g/mol), the mass is 2.16 g. Always check which form is on the label.

Lab steps

1. Measure the acid solution with a burette or measuring cylinder (burette for accuracy).
2. Measure the salt solution, or weigh the solid on a balance and dissolve it in a little distilled water in a beaker.
3. Combine in a beaker, leaving some room.
4. Check the pH with a calibrated pH meter and adjust (see below).
5. Transfer to a 250 cm³ volumetric flask, rinse the beaker into the flask, and make up to the mark with distilled water. Stopper and invert to mix.
6. Label with the contents, pH, concentration, date and your initials.

Method 2: partial neutralisation

Start with the weak acid and add a strong base (usually sodium hydroxide) until part of the acid has been converted to its conjugate base. This is often the most convenient method, because you only need one weak acid and a standard alkali.

Calculation

For the pH 5.00 ethanoate buffer, total 0.0250 mol ethanoic acid in 250 cm³:

• We need 1.588 × 10⁻² mol of A⁻, all made from HA by adding OH⁻.
• So add 1.588 × 10⁻² mol NaOH.
• Remaining HA = 0.0250 − 0.01588 = 9.12 × 10⁻³ mol ✔

Using 1.00 mol/dm³ NaOH: volume = 15.9 cm³.

So: take 25.0 cm³ of 1.00 mol/dm³ ethanoic acid, add 15.9 cm³ of 1.00 mol/dm³ NaOH, and dilute to 250 cm³.

Lab steps

1. Pipette the weak acid into a beaker and add about half the final volume of distilled water.
2. Place a calibrated pH electrode in the solution with gentle stirring.
3. Add the calculated volume of NaOH from a burette, then fine-tune by adding it dropwise while watching the meter, until the target pH is reached.
4. Transfer to a volumetric flask and make up to the mark.

This “titrate to the target pH” approach automatically corrects for small errors in the calculation, in reagent purity and in pKa.

Method 3: two salts of a polyprotic acid

For phosphate buffers near pH 7, mix two salts rather than an acid and a salt. Sodium dihydrogenphosphate (NaH₂PO₄) is the acid and disodium hydrogenphosphate (Na₂HPO₄) is the conjugate base. pKa₂ = 7.20.

Example: 500 cm³ of 0.100 mol/dm³ phosphate buffer at pH 7.40.

• Ratio [HPO₄²⁻] ÷ [H₂PO₄⁻] = 10^(7.40 − 7.20) = 1.585
• Total moles = 0.0500 mol
• Moles H₂PO₄⁻ = 0.0500 ÷ 2.585 = 0.01934 mol
• Moles HPO₄²⁻ = 0.0500 − 0.01934 = 0.03066 mol

Masses (anhydrous salts): NaH₂PO₄ (M = 119.98) → 2.32 g; Na₂HPO₄ (M = 141.96) → 4.35 g.

Phosphate salts come in several hydrated forms. The monohydrate, dihydrate, heptahydrate and dodecahydrate all have different molar masses, so always use the one printed on the bottle.

Calibrating the pH meter

A pH meter is only as good as its calibration.

1. Rinse the electrode with distilled water and blot gently (don’t rub the glass bulb).
2. Calibrate with at least two standard buffers that bracket your target, for example pH 4.00 and 7.00 for an ethanoate buffer.
3. Make sure standards and your buffer are at the same temperature, ideally 25 °C, or use the meter’s temperature compensation.
4. Rinse between every solution.
5. Store the electrode in its storage solution, never in distilled water, which can damage the glass membrane over time.

Adjusting the final pH

Calculated recipes rarely hit the target exactly, because real pKa values depend on temperature and ionic strength, and reagents are never perfectly pure. Fine-tune like this:

• pH too high? Add a few drops of the weak acid or a dilute strong acid (for example, 1 mol/dm³ HCl or ethanoic acid).
• pH too low? Add a few drops of dilute NaOH.

Add slowly, stir, and wait for the reading to stabilise before adding more. Adjust before making up to the final volume.

Troubleshooting

Problem	Likely cause	Fix
pH far from calculated value	wrong hydrate form used; wrong pKa; meter not calibrated	check labels and pKa; recalibrate
Reading drifts	electrode not equilibrated; temperature changing	wait longer; control temperature
Cloudy buffer	precipitation (e.g. phosphate with calcium or magnesium in tap water)	use distilled or deionised water
Mould or bacteria growth after days	phosphate and acetate buffers support microbes	store in the fridge; make fresh; filter-sterilise for biology
pH shifts on dilution for use	ionic strength effects	prepare at working concentration, or re-check pH after dilution

Safety

• Wear eye protection throughout; sodium hydroxide and concentrated acids are corrosive.
• Glacial ethanoic acid is corrosive and has strong fumes; handle it in a fume cupboard if using it undiluted.
• Wash spills with plenty of water.

Common mistakes

• Choosing a weak acid with the wrong pKa, so the buffer has little capacity at the target pH.
• Ignoring hydrate water in molar masses.
• Making up to volume before adjusting pH.
• Using tap water, which introduces calcium, magnesium and dissolved CO₂.
• Trusting the calculation over the meter. Always verify.

Key takeaways

• Choose a weak acid with pKa close to the target pH, then calculate the ratio with Henderson–Hasselbalch.
• Three methods: acid + salt, partial neutralisation with a strong base, or two salts of a polyprotic acid.
• Calibrate the pH meter with two standards and fine-tune the pH before making up to volume.
• Watch out for hydrated salts, temperature and water quality.
• For how the buffer will respond once made, see buffer calculations.