In the Laboratory
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Preparation of Buffers An Experiment for Quantitative Analysis Laboratory Paul T. Buckley Department of Chemistry, Hartwick College, Oneonta, NY 13820;
[email protected] The topics of acid–base buffers, buffer capacity, and buffer preparation are often granted several hours of class lecture in a typical quantitative analysis course, and in general chemistry courses as well. Buffer action plays a critical role in many aspects of biochemistry, industrial chemistry, analytical chemistry, and environmental systems. However, very few published laboratory experiments pertaining to buffers exist. Modern quantitative analysis texts seldom contain buffer experiments, and relatively few have been published in this Journal (1–3). Furthermore, none of these experiments gives students hands-on experience in the preparation of buffer solutions. We have noticed in our department that even students who do quite well on homework and exam questions dealing with buffers and buffer preparation are often at a loss when they actually need to prepare a real buffer to a certain pH in a research setting. The following experiment developed for our quantitative analysis laboratory gives students hands-on experience in buffer preparation and the application of the Henderson–Hasselbalch equation. This experiment only requires 1–2 hours, but we like to perform the experiment of Russo and Hanania (1) on the same day. The Russo and Hanania experiment examines the effect of varying the ratio of base concentration to acid concentration on buffer capacity. The two experiments complement each other well, and nicely fill a 3–4-hour lab. Experimental Procedure Before the laboratory, the instructor writes random numbers ranging from 3.0 to 12.0, to one decimal place, on slips of paper, which are then folded over and placed in a large beaker. Each student draws one of these slips from the beaker and students are then instructed to prepare 100 mL of a buffer in a volumetric flask, 0.05 M in the conjugate acid, to the pH they have selected. They are then directed to an assortment of conjugate acid–base pairs that represent a wide range of pK a values. It is up to the student to select the acid–base pair to work with, to calculate the amounts of reagents to mix, and to prepare the buffer to the instructor’s specifications. If both the weak acid and conjugate base are available in solid form the student should calculate, with the aid of the Henderson–Hasselbalch equation, the mass of each reagent to add. If only one of the pair is available as a solid, the student should calculate the mass of the solid needed and the volume of 0.50 M HCl or NaOH to add. Once the buffer has been prepared, the students measure the pH of their buffers with a calibrated pH meter. They are often surprised to find that while some solutions give reasonably good agreement between the theoretical and the actual pH, others may miss the target by as much as half a pH unit. The students are then asked to give reasons for these differences, such as neglect of activity 1384
coefficients in their calculations, or temperature effects in the laboratory. They then recalculate the theoretical pH using activity coefficients and see that the agreement with the actual pH has improved. They can calculate the ionic strength of their buffer from the concentrations of all reagents added, and determine activity coefficients from an activity coefficients table such as the one found in Harris (4 ). The final step in buffer preparation is to bring the solution to the desired pH by the addition of strong acid or strong base. The students complete this step by delivering, from a buret, the necessary amount of 1 M HCl or NaOH. The new volume of their solution is the original 100 mL in the volumetric flask plus the volume of acid or base added from the buret. The new theoretical pH is calculated by first calculating the new ratio of weak base to conjugate acid and is compared to the actual pH with and without the use of the proper activity coefficients. Hazards The handling of some strongly acidic or basic substances in solid form warrants the use of protective equipment such as goggles and light gloves. Most other substances or solutions present a negligible hazard during handling. Conclusion This experiment gives students a practical lesson in the preparation of buffers. It also gives them an opportunity to apply their theoretical understanding of calculations involving the Henderson–Hasselbalch equation, activity coefficients, and the effect of adding acid or base to a buffer. We find that students tend to feel more comfortable with buffer theory after performing this experiment and are better prepared for topics in courses such as biochemistry, where a working knowledge of buffer theory is essential. Also, students gain the experience required so that they feel competent in research situations requiring the preparation and use of buffers. W
Supplemental Material
A worksheet for students, notes for the instructor, and list of reagents are available in this issue of JCE Online. Literature Cited 1. Russo, S. O.; Hanania, G. I. H. J. Chem. Educ. 1987, 64, 817–819. 2. Russo, S. O.; Wiger, G. R.; de la Camp, U. J. Chem. Educ. 1978, 55, 401–402. 3. Clark, R. W.; White, G. D.; Bonicamp, J. M.; Watts, E. D. J. Chem. Educ. 1995, 72, 746–750. 4. Harris, D. C. Quantitative Chemical Analysis, 5th ed.; Freeman: New York, 1999; p 180.
Journal of Chemical Education • Vol. 78 No. 10 October 2001 • JChemEd.chem.wisc.edu
