Understanding Buffers in Liquid Chromatography

Part 1. Why buffers in chromatography?

The pH value of the mobile phase (eluent) is adjusted to improve peaks separation and extend the lifetime of your column. If possible, this pH adjustment should be performed using a buffer solution (liquid), rather than simply adding few drops of acid or a base (alkali) solution to the solvent bottle in uncontrolled fashion. Most certainly, it is not possible to manage reproducibility of “buffered” Mobile phase preparation when there is no buffering capacity.

Buffer solutions are prepared by combining a weak acid and its salt (e.g., sodium salt) or a weak base and its salt. Commonly used preparation methods can be divided in only 2 methods:

• Method 1 includes adding the acid (or base) into an aqueous solution of the salt while measuring the pH value with calibrated pH meter
• Method 2 involves combining the acid aqueous solution with the same concentration as the salt’s aqueous solution and mixing the two while measuring the pH value with a pH meter prior to combining them with organic solvent.

One of the most common mistakes is to assume that you can premix Mobile Phase (buffer and organic solvent) and measurement the pH of combined buffer/organic entity with a standard pH meter. The pH measurement of organic solutions requires specialized pH sensors and not that many scientists are willing to interchange standard pH sensors with organic solvents compatible pH sensors. Usually there is a solution to every problem, should you require more information about organic compatible pH sensors please check the Mettler recommendation here (pH Measurement of Organic Solvents).

When using a buffer solution as an HPLC mobile phase, however, a slight error in the pH value may adversely affect the separation reproducibility. Every time the pH meter shall be used it must be carefully inspected and calibrated.

Some buffer preparation methods will not require a pH meter, refer to table 1 below for generic buffer recipes. The methods involve accurate weighing of theoretically calculated amounts of the salt and acid (or base and acid) which guarantee the final pH value of prepared solution. Table 2 below contains 13x detailed preparations of buffers ranging from pH of 2.1 to 9.6.

 

Part 2. Buffer Solutions naming convention

It is hard to find a guideline to help us with buffer naming. Typical buffered solution name will contain salt name, pH value and buffer concentration at minimum, i.e., the notation “100 mM phosphate buffer solution, pH = 2.1” is not descriptive enough to help us in the buffer solution preparation. However, “100 mM sodium phosphate pH = 2.1” is very descriptive and annotates a buffer solution for which phosphoric acid is the acid, sodium ions act as counterions, the total concentration of the phosphoric acid radical is 100 mM, and the pH value of the buffer solution is 2.1.

We really need to be careful when annotating the buffer solutions, the best example is phosphate buffer. It would be huge mistake to specify only phosphate buffer solution without stating the associated pH value. The phosphoric acid is so called triprotic acid which has 3x ionizable hydrogen atoms in the same molecule thus exhibiting 3x different pKa values of pK₁ = 2.12, pK₂ = 7.21 and pK₃ = 12.44, which means you can use phosphoric acid for preparation of technically 3x different buffers.

 

Part 3. Buffering effect is Best near acid’s (or salt’s) pKa value

Let’s just consider acetic-acid (sodium) buffer solution created by mixing 1:1 portion of acetic acid and sodium acetate. The pH value of this buffer solution will be approx. 4.7, which is close to the pKa value of acetic acid (pKa 4.76), thus maximizing the buffering effect. This means that small changes in pH of the “system” (incl. mobile phase and sample) will have no effect on the pH value of buffer within the mobile phase and this will have no impact on your chromatography.

 

Part 4. Buffering capacity is better with higher concentration

The same situation exists when it comes to buffering capacity considering the buffer concentration. Look at the acetic acid (sodium) buffer solution in matter of two different concentrations, i.e., 100 mM and 10 mM. Higher concentration buffer will exhibit larger capacity to accommodate any addition of acid or base and still maintain the buffering capability and consistent pH value. However, all the advantage of higher concentration comes with a risk. At higher concentrations, there is a greater likelihood of salt crystals forming when the organic content of mobile phase meets the aqueous buffer solution.

 

Part 5. Buffered Salt solubility

The solubility of different salts always varies, e.g., potassium salt or sodium salt. Also, with higher slat concentration there is a greater likelihood of salt crystals forming when mixed with an organic solvent, as we discussed before.

Check below the example of the table with solubilities of various chloride salts at different temperatures (source) base on this example you can clearly see that even potassium and sodium salt of the same chloride very in their solubilities. Therefore, the maximum buffering concentration and the buffering capacity.

 

Table 1. Chloride salts solubility table 

Compound Name	Formula	Solubility in g/100 H2O at given temperature
		0oC	20oC	40oC	60oC	80oC
Ammonium chloride	NH4Cl	29.7	37.56	46	55.3	65.6
Barium chloride dihydrate	BaCl2 · 2H2O	30.7	35.7	40.8	46.4	52.5
Calcium chloride dihydrate	CaCl2 · 2H2O	–	–	128.1	136.8	147
Copper (I) chloride (at 25oC)	CuCl	–	1.52 *	–	–	–
Copper (II) chloride dihydrate	CuCl2 · 2H2O	70.65	77	83.8	91.2	99.2
Iron (III) chloride hexahydrate	FeCl3 · 6H2O	74.5	91.94	–	–	–
Iron (II) chloride tetrahydrate	FeCl2 · 4H2O	–	62.35	68.6	78.3	–
Lead chloride	PbCl2	0.67	0.99	1.45	1.98	2.62
Lithium chloride monohydrate	LiCl · H2O	–	82.82	90.4	100	113
Magnesium chloride hexahydrate	MgCl2 · 6H2O	52.8	54.57	57.5	60.7	65.87
Potassium chloride	KCl	28.15	34.24	40.3	45.6	51
Sodium chloride	NaCl	35.6	35.8	36.42	37.05	38.05

In addition to the above, when performing high-sensitivity analysis using UV short wavelengths, you should avoid organic acids, such as carboxylic acid, buffer solutions if possible. It is always necessary to consider these and other analysis conditions to use an appropriate buffer solution with the correct concentration and buffering capacity.

 

Table 2. Preparation details – common buffers

100 mM Phosphoric-Acid (Sodium) Buffer Solution	pH = 2.1
Sodium dihydrogen phosphate dihydrate (M.W. = 156.01)	50 mmol (7.8 g)
Phosphoric acid (85%, 14.7 mol/L)	50 mmol (3.4 mL)
Add the above to water to create a solution of Total volume of 1 L
10 mM Phosphoric-Acid (Sodium) Buffer Solution	pH = 2.6
Sodium dihydrogen phosphate dihydrate (M.W. = 156.01)	5  mmol (0.78 g)
Phosphoric acid (85%, 14.7 mol/L)	5 mmol (0.34 mL)
Add the above to water to create a solution of Total volume of 1 L
Alternatively, dilute 100 mM phosphoric acid (sodium) buffer solution (pH = 2.1) by a factor of 10
50 mM Phosphoric-Acid (Sodium) Buffer Solution	pH = 2.8
Sodium dihydrogen phosphate dihydrate (M.W. = 156.01)	40 mmol (6.24 g)
Phosphoric acid (85%, 14.7 mol/L)	10 mmol (0.68 mL)
Add the above to water to create a solution of Total volume of 1 L
10 mM Tartaric Acid (Sodium) Buffer Solution	pH = 2.9
Tartaric acid (M.W. = 150.09)	7.5 mmol (1.13 g)
Sodium tartrate dihydrate (M.W. = 230.08)	2.5 mmol (0.58 g)
Add the above to water to create a solution of Total volume of 1 L
20 mM Citric Acid (Sodium) Buffer Solution	pH = 3.1
Citric acid monohydrate (M.W. = 210.14)	16.7 mmol (3.51 g)
Trisodium citrate dihydrate (M.W. = 294.10)	3.3 mmol (0.97 g)
Add the above to water to create a solution of Total volume of 1 L
10 mM Tartaric Acid (Sodium) Buffer Solution	pH = 4.2
Tartaric acid (M.W. = 150.09)	2.5 mmol (0.375 g)
Sodium tartrate dihydrate (M.W. = 230.08)	7.5 mmol (1.726 g)
Add the above to water to create a solution of Total volume of 1 L
20 mM Citric Acid (Sodium) Buffer Solution	pH = 4.6
Citric acid monohydrate (M.W. = 210.14)	10 mmol (2.1 g)
Trisodium citrate dihydrate (M.W. = 294.10)	10 mmol (2.94 g)
Add the above to water to create a solution of Total volume of 1 L
100 mM Acetic Acid (Sodium) Buffer Solution	pH = 4.7
Acetic acid (glacial acetic acid, 99.5%, 17.4 mol/L)	50 mmol (2.87 mL)
Sodium acetate trihydrate (M.W. = 136.08)	50 mmol (6.80 g)
Add the above to water to create a solution of Total volume of 1 L
100 mM Phosphoric-Acid (Sodium) Buffer Solution	pH = 6.8
Sodium dihydrogen phosphate dihydrate (M.W. = 156.01)	50 mmol (7.8 g)
Disodium hydrogen phosphate dodecahydrate (M.W. = 358.14)	50 mmol (17.9 g)
Add the above to water to create a solution of Total volume of 1 L
10 mM Phosphoric Acid (Sodium) Buffer Solution	pH = 6.9
Sodium dihydrogen phosphate dihydrate (M.W. = 156.01)	5 mmol (0.78 g)
Disodium hydrogen phosphate dodecahydrate (M.W. = 358.14)	5 mmol (1.79 g)
Add the above to water to create a solution of Total volume of 1 L
100 mM Boric Acid (Potassium) Buffer Solution	pH = 9.1
Boric acid (M.W. = 61.83)	100 mmol (6.18 g)
Potassium hydroxide (M.W. = 56.11)	50 mmol (2.81 g)
Add the above to water to create a solution of Total volume of 1 L
100 mM Boric Acid (Sodium) Buffer Solution	pH = 9.1
Boric acid (M.W. = 61.83)	100 mmol (6.18 g)
Sodium hydroxide (M.W. = 40.00)	50 mmol (2.00 g)
Add the above to water to create a solution of Total volume of 1 L
20 mM (Acetic Acid) Ethanolamine Buffer Solution	pH = 9.6
Monoethanolamine (M.W. = 61.87, d = 1.017)	20 mmol (1.22 mL)
Acetic acid (glacial acetic acid, 17.4 mol/L)	10 mmol (0.575 mL)
Add the above to water to create a solution of Total volume of 1 L

 

Part 6. Summary and conclusion

I hope that the above helps you to understand a little bit more about buffers, their capacity and few little things which can have a tremendous effect on your analysis. Additionally, the above can help you troubleshooting some nasty chromatography and understand that also instrumentation (HPLC) can be impacted by the incorrect choice of the buffer concentration.

 

Source: Tips for practical HPLC analysis – Separation Know-how. LC World Talk Special Issue Volume 2. Shimadzu.

Liquid Chromatography – Master the Basics

This article is part of our “Liquid Chromatography – Master the Basics” series, your go-to resource for comprehensive and insightful updates on the world of liquid chromatography. Each month in 2024 we will dive into a Liquid Chromatography topic, offering content that is both accessible to beginners and beneficial for experienced scientists.

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