Buffer Solutions - How to prepare buffer solutions

Buffer Solutions - How to prepare buffer solutions - Hydrochloric Acid Buffer

Buffer solutions play important roles in controlling the solubility of ions in solution and in maintaining the pH in biochemical and physiological processes. Many life processes are sensitive to pH and require regulation withing a small range of H₃O⁺ and OH^- concentrations. Human blood has a pH near 7.4 that is maintained by a combination of carbonate, phosphate, and protein buffer systems.

How can we prepare buffer solutions in the laboratory?

A buffer solution is any solution that maintains approximately constant pH upon small addditions of acid or base. Adding as little as 0.1 ml of concentrated HCl to a liter of H₂O shifts the pH from 7.0 to 3.0. The same addition of HCl to a liter solution that is 0.1 M in both a weak acid and its conjugate weak base (a buffer solution) , however, results in only a negligible change in pH.

A broader definition of buffered solutions is that these solutions resist changes in the activity of an ion on the addition of substances that are expected to change the activity of that ion. Buffered solutions are systems in which the ion is in equilibrium with substances capable of removing or releasing the ion. Their buffering action is a consequence of the relationship between pH and the relative concentrations of the conjugate weak acid / weak base pair.

An example of a buffer is a mixture of acetic acid and sodium acetate. The equilibrium position of the buffer is governed by the reaction:

As mentioned above control of pH is vital in synthetic and analytical chemistry and in living organisms. Procedures that work well at a pH of 5 may fail when the pH becomes 4. Fortunately, it is possible to prepare buffer solutions (using buffers) that maintain the pH close to any desired value by the proper choice of a weak acid and its conjugate base concentration.

The following procedure can be used to prepare any buffer (different buffer) solutions:

Determine the optimal pH (the required pH)
Select a weak acid with a pka near the desired pH
Calculate the ratio of salt to acid required to produce the desired pH (Henderson-Hasselbach equation): pH = pk_a - log [H_A]_O/[A^-]_O
Determine the desired buffer capacity of the solution
Calculate the total buffer concentration required to produce this buffer capacity ß (Van Slyke equation): ß = 2.3* C* (k_a * [H₃O⁺]) / (k_a + [H₃O⁺] )²
Determine the pH and the buffer capacity of the final buffer solution using a reliable pH meter.

The preparation of standard buffer solutions for various ranges between pH 1.2 and 10.0 is described below. The volumes shown in the table are for 200 ml of buffer solution except where otherwise stated:

Hydrochloric Acid Buffer, Standard Buffer Solution

Prepare the following solutions:

Hydrochloric Acid , 1 M:  

Prepare a 1 M HCl solution: Dilute 85 ml of HCl (d=36.46) with water to 1000 ml.
Standarize the above solution as follows: Weight 5.0 g of tromethamine, dried according to the label instructions or, if this information is not available, dried at 105 C for 3 h. Dissolve in 50 ml of water and add 2 drops of bromocresol green. Titrate with 1M HCl to a pale yellow endpont:

M = mg tromethamine / (121.14 * ml HCl) (M, Molarity of solution)

Hydrochloric Acid , 0.2 M:  

Prepare a 0.2 M HCl solution: Dilute appropriatelly the 1 M HCl solution

 

Potassium Chloride, 0.2 M:  

Dissolve 14.91 g of potassium chloride KCl in water
Dilute with water to 1000 ml

 

Place 50 ml of the 0.2 M KCl solution in a 200 ml volumetric flask
Add the volume of the 0.2 M HCl shown in the table below
Add water to volume 200 ml

Hydrochloric Acid Buffer - Standard Buffer Solution
	pH	0.2 M HCl (ml)	0.2 M KCl (ml)
	1.2	85.0	50
	1.3	67.2	50
	1.4	53.2	50
	1.5	41.4	50
	1.6	32.4	50
	1.7	26.0	50

1.8	20.4	50
1.9	16.2	50
2.0	13.0	50
2.1	10.2	50
2.2	7.8	50

 

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Relevant Posts - Relevant Videos

Preparing Standard Buffer Solutions -Phthalate Acid and Neutralized Buffers

Standard Buffer Solutions - Phosphate Buffer(Monobasic)

Standard Buffer Solutions - Preparing Na Acetate Buffers

 

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References

1. CRC Handbook of Chemistry and Physics, 52nd edition, The Chemical Rubber Co., (1971)
2. U.S. Pharmacopeia, 68, USP 36
3. David W. Oxtoby, H.P. Gillis, Alan Campion, “Principles of Modern Chemistry”, Sixth Edition, Thomson Brooks/Cole, 2008
4. Steven S. Zumdahl, “Chemical Principles”  6th Edition, Houghton Mifflin Company, 2009

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Key Terms

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