Choosing a Buffer and pH for Peptide Solutions

Why pH matters for a peptide

Every peptide carries ionisable groups: the free N- and C-termini and side chains such as Asp, Glu, His, Lys, Arg, Cys and Tyr. The pH decides which of them are charged, and so the net charge of the molecule. Solubility is usually lowest near the isoelectric point, where the net charge is close to zero, which is why the solubility guide recommends working away from it. The peptide properties calculator estimates net charge across pH for a sequence.

pH also changes chemistry. In a model hexapeptide containing Asn-Gly, the rate of deamidation depended strongly on pH, temperature and buffer composition, and the pathway itself changed with pH: from pH 5 to 12 it ran through a cyclic imide and gave both Asp and isoAsp products, while under acidic conditions direct hydrolysis gave only the Asp form.^[1] Other pathways are summarised in the storage stability guide.

What a buffer is and where it works

A buffer is a weak acid together with its conjugate base. Added acid is taken up by the base and added base by the acid, so the pH moves much less than it would in water. The pKa of the weak acid is the pH at which the two forms are present in equal amounts, and buffering is strongest there. The useful range is generally taken as the pKa plus or minus one unit.^[2] Approximate pKa values at 25 °C for common choices are:

• Acetate about 4.8, for roughly pH 3.8 to 5.8.
• Citrate three pKa values near 3.1, 4.8 and 6.4, covering a broad acidic range.
• Histidine about 6.0 for the imidazole side chain.
• Phosphate about 7.2 for the second dissociation, the one used near neutral pH.
• HEPES about 7.5.
• Tris about 8.1, falling as the temperature rises.

In 1966 Good and colleagues described a set of zwitterionic buffers chosen for biological work, with pKa values between about 6 and 8, high solubility in water, low absorbance in the ultraviolet and limited binding of metal ions; HEPES and MES are among them.^[3] A later paper from the same group added further buffers to the series.^[4]

A worked example: Henderson-Hasselbalch

The Henderson-Hasselbalch equation links pH to the ratio of the two forms: pH = pKa + log([base]/[acid]).

Suppose you want 50 mM acetate buffer at pH 5.0, using a pKa of 4.76. The ratio is 10^(5.0 - 4.76) = 10^0.24 ≈ 1.74. The base therefore makes up 1.74 / 2.74 ≈ 63.5% of the total: about 31.7 mM sodium acetate and 18.3 mM acetic acid. The molarity calculator converts these into grams.

The equation is an approximation. It uses concentrations in place of activities, so ionic strength and temperature move the real pH; measure the finished solution and adjust it.

Buffer capacity: concentration matters

The pH at which a buffer works is set by its pKa; how much acid or base it can absorb is set by its concentration. That second property is the buffer capacity.

An illustrative calculation shows the difference. Take sodium phosphate at pH 7.21, so that the two forms are equal, and add strong acid to a final 2 mM. In a 10 mM buffer, 5 mM base and 5 mM acid become 3 mM and 7 mM, and the pH falls to about 7.21 + log(3/7) ≈ 6.84, a drop of 0.37. In a 50 mM buffer, 25 and 25 mM become 23 and 27 mM, and the pH falls only to about 7.14. A dilute buffer can be overwhelmed by the acidity of whatever is dissolved in it, including residual acid in a peptide salt (see the counter-ion guide), which is one reason to measure the final pH after the peptide has been added.

Temperature and freezing move pH

pKa values change with temperature. Amine buffers such as Tris are the most sensitive: the pKa of Tris falls by roughly 0.03 for each degree Celsius of warming, so a Tris solution set to pH 8.0 at room temperature reads noticeably higher in a cold room. A pH should therefore be set at, or reported with, the temperature at which the solution will be used.

Freezing adds larger effects. As ice forms, solutes are concentrated in the remaining liquid, pKa values shift and some buffer salts crystallise. In sodium phosphate buffers cooled to -10 °C, precipitation of the disodium salt caused abrupt drops in pH; 50 and 100 mM solutions starting at pH 7.4 reached about pH 4.2, and an 8 mM solution about pH 5.2, while solutions starting at pH 5.7 shifted less.^[5] A survey of common formulation buffers from +25 to -30 °C found the largest change in sodium phosphate, an increase of less than one unit for histidine, acetate, citrate and succinate, and an increase of about 1.2 units for Tris. In that study, aggregation of a frozen antibody solution depended on the presence of a cryoprotectant rather than on the pH change.^[6] The freeze-thaw guide covers aliquoting to limit repeated cycles.

The buffer is not always inert

A buffer species can take part in the chemistry it is meant to control. In the hexapeptide study above, buffer catalysis of deamidation was seen between pH 7 and 11.^[1] In a series of small Asn-containing peptides, deamidation rates were not affected by ionic strength but showed general base catalysis by small bases such as ammonia; the residue following Asn had a large effect, with Asn-Gly the fastest, and at neutral and alkaline pH the products formed in about a 3:1 ratio of isoAsp to Asp.^[7]

A review of buffers in protein solutions describes effects well beyond pH control: some buffers increase conformational stability through binding or by being excluded from the protein surface, some change colloidal stability or behaviour at interfaces, some destabilise, and the buffers themselves can degrade or interact with other components.^[8] Basic chemistry adds others: phosphate precipitates with calcium, and the primary amine of Tris reacts with amine-reactive labelling reagents. For a metal complex such as GHK-Cu, any buffer that binds the metal ion competes for it.

Buffers for mass spectrometry and HPLC

Non-volatile salts such as phosphate deposit in an electrospray (ESI) source and suppress ionisation, so samples for LC-MS are usually prepared in volatile additives such as formic acid, acetic acid, ammonium acetate or ammonium bicarbonate. These are not interchangeable buffers. Ammonium acetate dissolved in water gives about pH 7, but ammonium and acetate are not a conjugate acid-base pair, so the solution buffers only near pH 4.75 and 9.25 and its neutral pH is easily shifted; the author of that analysis recommends calling it an ammonium acetate solution rather than a pH 7 buffer.^[2] The mass spectrometry guide explains what the spectrum then shows.

Measuring pH and keeping records

Calibrate the meter with at least two standards that bracket the target pH, at the temperature of use. Adjust the pH after all components, including any cosolvent, have been added, then bring the solution to volume. Then record:

• Buffer. Species, concentration and counter-ion, for example 20 mM sodium phosphate rather than "phosphate buffer".
• pH and temperature. The measured value and the temperature at which it was measured.
• Additives. Salts, chelators, cosolvents and their concentrations.
• History. Preparation date, storage temperature and number of freeze-thaw cycles.

This guide describes general laboratory methods and published literature only. It is not a statement about the stability or behaviour of any particular material. Terms used here are defined in the glossary, and questions can be sent to support.