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Preparation and Properties of Buffers: A Chemistry Lab Manual - Prof. Teresa E. Eberhardt, Lecture notes of Translation Theory

Albany State University (ASU)Translation Theory
Prof. Teresa E. Eberhardt

This lab manual provides a comprehensive guide to preparing and understanding buffer solutions. It covers the direct and indirect methods of buffer preparation, explores the concept of buffer capacity, and includes detailed instructions for conducting experiments to determine the ph and buffering capacity of various buffer solutions. The manual also includes post-lab questions to reinforce key concepts and encourage further exploration of buffer systems.

Typology: Lecture notes

2022/2023

Uploaded on 09/04/2024

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PREPARATON AND PROPERTIES OF BUFFERS
Pre-Lab
Your pre-lab is to complete the calculations for PART A. 1 and 2 and PART B. 1 for this
experiment!!
Introduction:
A buffer solution is one that is resistant to change in pH when small amounts of strong
acid or base are added. For example, when 0.01 mole of strong acid or base are added to
distilled water, the pH drops to 2 with the acid and rises to 12 with the base. If the same amount
of acid or base is added to an acetic acid – sodium acetate buffer, the pH may only change a
fraction of a unit.
Buffers are important in many areas of chemistry. When the pH must be controlled
during the course of a reaction, the solutions are often buffered. This is often the case in
biochemistry when enzymes or proteins are being studied. Our blood is buffered to a pH of 7.4.
Variations of a few tenths of a pH unit can cause illness or death. Acidosis is the condition when
pH drops too low. Alkalosis results when the pH is higher than normal.
Two species are required in a buffer solution, most often, a weak acid (HA) and its
conjugate base (A-). The weak acid component of the buffer is capable of reacting with strong
base, (OH- ) and the conjugate base component will react with strong acid (H3O+) as shown
below:
HA + OH- A- + H2O
A- + H3O+ H2O + HA
Note that the result is effectively the elimination of the strong acid or base.
The two species must not react with each other. An example of a buffer system a weak acid and
its conjugate is acetic acid and sodium acetate. Another, less common type of buffer system may
1
Exp 2
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Download Preparation and Properties of Buffers: A Chemistry Lab Manual - Prof. Teresa E. Eberhardt and more Lecture notes Translation Theory in PDF only on Docsity!

PREPARATON AND PROPERTIES OF BUFFERS

Pre-Lab Your pre-lab is to complete the calculations for PART A. 1 and 2 and PART B. 1 for this experiment!! Introduction: A buffer solution is one that is resistant to change in pH when small amounts of strong acid or base are added. For example, when 0.01 mole of strong acid or base are added to distilled water, the pH drops to 2 with the acid and rises to 12 with the base. If the same amount of acid or base is added to an acetic acid – sodium acetate buffer, the pH may only change a fraction of a unit. Buffers are important in many areas of chemistry. When the pH must be controlled during the course of a reaction, the solutions are often buffered. This is often the case in biochemistry when enzymes or proteins are being studied. Our blood is buffered to a pH of 7.4. Variations of a few tenths of a pH unit can cause illness or death. Acidosis is the condition when pH drops too low. Alkalosis results when the pH is higher than normal. Two species are required in a buffer solution, most often, a weak acid (HA) and its conjugate base (A-). The weak acid component of the buffer is capable of reacting with strong base, (OH-^ ) and the conjugate base component will react with strong acid (H 3 O+) as shown below: HA + OH-  A-^ + H 2 O A-^ + H 3 O+^  H 2 O + HA Note that the result is effectively the elimination of the strong acid or base. The two species must not react with each other. An example of a buffer system a weak acid and its conjugate is acetic acid and sodium acetate. Another, less common type of buffer system may

Exp 2

contain a weak base (B) and its conjugate acid (BH+) e.g. ammonia (B) and ammonium chloride (BH+). In general, the pH range in which a buffer solution is effective is +/- one pH unit on either side of the pKa. The Henderson–Hasselbalch provides the information needed to prepare a buffer.

pH = pKa + log

[ conjugatebase ] [ weakacid ] There is a limit to the amount of acid or base that can be added to a buffer solution before one of the components is used up. This limit is called the buffer capacity and is defined as the moles of acid or base necessary to change the pH of one liter of solution by one unit. Buffer Capacity = (number of moles of OH-^ or H 3 O+^ added) (pH change)(volume of buffer in L) There are two major ways to prepare a buffer

  1. The Direct Method The quantity of both components of the conjugate acid-base pair are determined using the Hendersen Hasselbach equation to obtain the desired ratio and then dissolved in water.
  2. The Indirect Method Both components are obtained from a prescribed amount of only one component, with the second being produced by adding a specified amount of strong acid or strong base to yield the desired ratio. In this experiment, the Henderson-Hasselbalch equation will be used to determine the amount of acetic acid and sodium acetate required to prepare a series of buffer solutions. Once the buffer solutions have been prepared, their buffer capacity will be determined. A buffer will also be prepared using the indirect method

Procedure : PART A. Direct Method of Preparing Buffer Solutions.

  1. Your instructor has prepared a 2.0 M stock solution of acetic acid. Prepare 100 mL of 0. M acetic acid by dilution of this stock solution.
  2. Also prepare 100 mL of stock solution of 0.25 M sodium acetate by dissolving the solid salt. Record all pertinent data.
  3. Using the 0.250 M solutions prepared, prepare in clean 50 mL beakers the following buffers. mL of 0.250 M HC 2 H 3 O 2 (HA) mL of 0.250 M NaC 2 H 3 O 2 (A-) Buffer 1 16.0^ 4. Buffer 2 12.0 8. Buffer 3 10.0 10. Buffer 4 6.0 14. Buffer 5 2.0^ 18.
  4. Determine the pH of each buffer. Before using the pH meter you should always check that it is properly calibrated. Use a standard pH 7.0 solution to check this. If the pH meter does not give the correct pH reading to within 0.1 pH units it will have to be calibrated. The procedure for calibrating pH meters varies from instrument to instrument. Follow your instructor’s directions for this step. Once the pH meter is calibrated, Measure the pH and record the results for each solution.
  5. Tabulate the measured pH and the calculated pH for each solution. PART B: Indirect method of preparing the Buffer. (Using one component of the buffer and adding acid or base to produce the other component of the buffer).
    1. Determine the quantity of sodium acetate in grams required to prepare 100 mL of the assigned molarity of your buffer solution MW (sodium acetate = )

Assigned Buffer solutions to prepare Group 1 0.15 M pH 4. Group 2 0.15 M pH 5. Group 3 0.25 M pH 4. Group 4 0.25 M pH 5. Group 5 0.35 M pH 4. Group 6 0.35 M pH 4. Group 7 0.45 M pH 4. Group 8 0.45 M pH 4.

  1. Weigh out the quantity of sodium acetate determined in step 1 and transfer to a beaker and dissolve in ~90 mL of water and measure the pH with a pH meter.
  2. If pH is not the desired pH, use a 3.0 M HCl or 3.0 M NaOH to adjust to required pH. Please note: Be careful not to overshoot the required pH! ADD dropwise! As you get closer to the required pH you may also use a less concentrated (1M) HCL or NaOH solutions.
  3. Pour solution into the appropriate volumetric flask and bring volume up to the 100 mL mark. PART C. pH changes by adding Strong acid or Base to buffered and unbuffered solutions.
  4. Transfer 15 mL of your buffer prepared in PART B. to a small beaker.
  5. Add 0.5 mL of 1.0 M HCl and mix well. Record the original pH of the buffer and the pH after addition of the HCl.
  6. Transfer 15 mL of water to a beaker. Record the original pH of the water
  7. Add 0.5 mL of 1.0 M HCl to the water; mix well and record the pH
  8. Repeat step 1 – 3 using 0.5 mL of 1.0 M NaOH instead of the HCl.
  9. Compare the pH of the buffered solutions and water. PART D. Concentration and Buffering Capacity
  10. You will need 0.5 M NaOH and 0.5 M HCl for this section. Prepare 50 mL of 0.5 M solutions of each of these from the STOCK solutions provided in the lab.

EXPERIMENT 2: PREPARING BUFFERS AND BUFFER CAPACITY

LAB REPORT

Write up the lab report using the guidelines that have been provided. Report should have a ‘calculations section that shows at a minimum one example of each type of calculation done in the experiment. Data/Results Report Part A: Use the Henderson Hasselbalch equation to calculate the theoretical pH of each of the buffers prepared. Report the calculated and measured pH of each buffer solution Part B. Report the concentration of your buffer and the final pH, the acid or base that was necessary to obtain the desired pH and the predominant component of the in the solution Part C. Report the pH of the buffered and unbuffered solutions Part D. Tabulate the initial pH of each solution, the final pH the change in pH and the buffering capacity of each solution. Part E. Graph the titration curves for your buffer. Label the following regions; pKa, buffering range, region where [HA] > [A-]; [HA]<[A-]; [HA]=[A-] NOTE: As you formulate the discussion section of this lab report you should consider the following, providing specifics for your buffer in the process. What is a buffer? Why was your buffer not necessarily at the required pH when first prepared by dissolving buffer salt in water? Why did adding acid or base get your buffer to the required pH? At what pH is a buffer most effective? Within what pH range values should it be used? What does the different regions of the titration curve of the buffer show? What determine the capacity of a buffer? How is the buffering capacity affected by dilution? Were your results as expected? If not what reasons may have contributed to error or to your experiment not generating the results expected. How might this be corrected if he experiment was to be repeated?

Post Lab Questions:

  1. How does the concentration of the buffer affect the buffer capacity?
  2. What differences would be observed if HCl were used in place of NaOH to measure buffering capacity?
  3. Write equations to show which component of the acetate buffer system reacted with a) the HCl and b) NaOH.
  4. Calculate the pH of a solution prepared by mixing 25.0 mL of 0.60 M HC 2 H 3 O 2 and 15 mL of 0.60 M NaC 2 H 3 O 2.
  5. Calculate the pH of a solution prepared by mixing 25.0 mL of 0.60 M HC 2 H 3 O 2 and 15.0 mL of 0.60 M NaOH.