




Study with the several resources on Docsity
Earn points by helping other students or get them with a premium plan
Prepare for your exams
Study with the several resources on Docsity
Earn points to download
Earn points by helping other students or get them with a premium plan
Community
Ask the community for help and clear up your study doubts
Discover the best universities in your country according to Docsity users
Free resources
Download our free guides on studying techniques, anxiety management strategies, and thesis advice from Docsity tutors
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
1 / 8
This page cannot be seen from the preview
Don't miss anything!
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
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.
[ 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
Procedure : PART A. Direct Method of Preparing Buffer Solutions.
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.
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: