Chem 1B Dr. White 57 Experiment 7 Buffer Capacity & Buffer Preparation Objectives To learn how to choose a suitable conjugate acid- base pair for making a buffer of a given ph To gain experience in using the Henderson- Hasselbach equation To understand the limitations of buffers: buffer range and buffer capacity Introduction In the Buffers Day 1 experiment you discovered that a buffer is made up of a weak acid- conjugate base pair (for example, NH 4 +/NH 3). Buffers maintain a relatively constant ph by converting added strong acid into the weak acid component or by converting added strong base into the conjugate base component. To more thoroughly explore how buffers work, we must explore the common ion effect. The equilibrium reaction, for any buffer, may be written as the acid ionization (K a) equilibrium: HA (aq) + H 2O (l) H 3O + (aq) + A (aq) (1) The presence of the conjugate base shifts the equilibrium concentrations to the left, decreasing the hydrogen ion concentration compared to a pure weak acid. When adding base, OH, to this mixture, the weak acid neutralizes the strong base and the conjugate base concentration increases while the weak acid concentration decreases: HA (aq)+ OH (aq) H 2O (l) + A (aq) (2) However, the hydrogen ion concentration itself is changed very little. Similarly when strong acid is added to the buffer, the base neutralizes the hydrogen ions producing more weak acid while decreasing the conjugate base concentration: A (aq) + H 3O + (aq) HA (aq) + H 2O (l) (3) Here again the hydrogen ion concentration of the solution itself is altered only slightly. Since equilibrium between the weak acid and the conjugate base (Eq. 1) is always maintained, the Henderson- Hasselbach equation: ph = pk a + log [A ] [HA] (4) gives a good approximation for determining the ph of a buffer solution. In this equation [HA] is the concentration of the weak acid component of the buffer, [A - ] is the concentration of its conjugate base and pk a refers to the pk a value for the weak acid component. Keep in mind that this equation has limitations, especially for buffers that contain polyprotic weak acids. On the other hand, when preparing a buffer solution for use, it is a good place to start. However, practical experience teaches us that one should not just rely on the Henderson- Hasselbach equation when preparing a buffer in lab but make careful measurements of ph as the buffer is prepared. In the preparation of buffers, the concentration of the buffering solutes (acid- base pair) are usually in the range of 0.1 to 0.5 molar, although there are media in which the concentrations are much lower. Very dilute buffer solutions can readily be overwhelmed by the addition of strong acids or bases. This is often the effective action that changes the colors of acid- base indicators. The ratio of the acid concentration to the conjugate base concentration is usually between 1:10 and 10:1 for the most effective buffering action. Substituting these ratios into the Henderson- Hasselbach equation gives us a ph range over which a buffer solution is most effective, that is from ph = pka + log(0.1) to ph = pka + log(10). There are three easy methods to prepare simple buffer solutions. The first is to directly mix a weak acid with the conjugate base. The second and third methods rely on chemical reactions to prepare the desired ratio of [A - ]/[HA]. If one has only the weak acid, one can add a strong base such as NaOH as the limiting reactant, thereby neutralizing a portion of the weak acid and producing the conjugate base (see equation 2). If one has only the weak base, one can add a strong acid such as HCl as the limiting reactant, thereby neutralizing a portion of the weak base and producing the conjugate acid (see equation 3). In this lab you will (1) investigate these three methods for preparing buffers; (2) practice choosing suitable buffer solutions for given ph s; (3) investigate the buffers of the polyprotic weak acid H 3PO 4, and (4) test buffering capacity and effective buffering range by adding HCl and NaOH to a buffer. Reagents Available 1 M HCl, 1 M NaOH, Solid NaC 2H 3O 2 3H 2O, 0.4 M NH 4Cl, 0.1 M HCl, 0.4 M NaH 2PO 4, 0.1 M NaH 2PO 4/0.1 M Na 2HPO 4 buffer (prepared by stockroom) The concentrations given above are approximate. The stock solutions will be standardized and the actual
58 Chem 1B Dr. White concentrations written on the containers. You must record the actual concentrations (with the correct number of sig figs) in your lab notebook. Procedure As you perform the lab and collect waste solutions pour them into a large beaker. This mixture should then be discarded in the appropriate waste container. DO NOT POUR ANY OF THE SOLUTIONS DOWN THE DRAIN. Use only deionized water for preparing solutions and rinsing the ph electrodes. DO NOT in any circumstance put equipment into reagent bottles or pour any unused reagents back into bottles. Use a SMALL beaker to obtain the quantity of solution needed, and refill the beaker as required. DO NOT WASTE THE CHEMICALS. Using the ph sensor: 1. When not in use, place the ph sensor back into its storage solution. 2. Rinse the ph electrode thoroughly with distilled water and gently pat it dry before placing it in a solution to measure the ph. 3. DO NOT use the ph electrode to stir solutions. 4. As you observe the ph readings allow at least 15 seconds before recording your measurements. 5. Read and record the ph values from the computer to two decimal places! (That means two digits after the decimal point.) Part A: Buffer Capacity and Range - An Investigation Using a Diluted and Undiluted H 2PO 4 and HPO 4 2 Buffer Solution 1. Prepare the ph sensor for data collection. Go to the Experiment menu and under Calibrate choose the channel with the ph probe (ex: CH1:pH ). In the window that appears make sure the Calibration tab is chosen. Click on Calibrate Now. Rinse the ph meter with copious amounts of deionized water. Carefully blot dry. Dip the ph meter in the ph 4 standard solution. Wait until the voltage reading stabilizes. Then, in the field beneath Enter Value enter the ph value of the solution and click Keep. Repeat the process with the ph 10 standard solution. When finished with this step click on Done to close the window. 2. Click on the Data Collection button. You will see a Dialog Box. Change Mode from Time Based to Events with Entry. Type Drops for column name and Drops for short name. Change Units to Drops then click Done. 3. In your lab notebook, record the concentration (from the container) of the H 2PO 4 /HPO 4 2 stock buffer solution, the HCl stock solution and the NaOH stock solution. 4. Measure two different 25- ml portions of the stock buffer solution into two different 50- ml beakers. 5. Using one of these 25 ml portions: a. BEFORE adding any 1 M HCl, click on. In the upper- left hand corner of the Logger Pro window, you will see the ph in "real time". When it appears the ph has stabilized after stirring the acid solution for a while, click. A new window will appear asking you to enter the number of drops added up to this point. Enter 0 into this window. This first data point (0 drops and corresponding ph) should now be recorded on the data table to the left, and a red dot should show up on the graph indicating this data point. b. Add a drop of the 1 M HCl solution. Stir the solution, wait 10-15 seconds and click. Enter in the total number of drops that have been added. Repeat this step until 10 drops have been added. Then click. c. Pour the buffer solution into your waste collection beaker. 6. Using the second 25 ml portion: a. BEFORE adding any NaOH, click on. A dialog box will appear. Click store latest run. In the upper- left hand corner of the Logger Pro window, you will see the ph in "real time". When it appears the ph has stabilized after stirring the acid solution for a while, click. A new window will appear asking you to enter the number of drops added up to this point. Enter 0 into this window. This first data point (0 drops and corresponding ph) should now be recorded on the data table to the left, and a red dot should show up on the graph indicating this data point. b. Add a drop of the NaOH solution. Stir the solution, wait 10-15 seconds and click. Enter in the total number of drops that have been added. Repeat this step until 10 drops have been added. c. Pour the buffer solution into your waste
Chem 1B Dr. White 59 collection beaker. Diluted H 2PO 4 and HPO 4 2 buffer 1. Measure 10- ml of the H 2PO 4 /HPO 4 2 stock buffer solution into a 100- ml graduated cylinder. Add water to this solution to dilute to a total volume of 100- ml (Determine the concentration of the diluted solution and record it in your notebook!). 2. Measure two different 25- ml portions of the diluted solution into two different 50- ml beakers. 3. Repeat steps 5 and 6 above. When you click latest run., a dialog box will appear. Click store 4. Once you are done, you will have a plot with 4 curves. Save your data and print it. Be sure to label the curves. Part B: Preparation of Buffer Solutions In this part of the experiment you will prepare three different buffer solutions: 1. one with ph = 9.50 2. one with ph = 5.00 3. one with ph= 2.00. Use the following table of K a values to make your selection. Table 1: Selected K a Values at 25 C Acid K a H 3PO 4 7.6 x 10-3 H 2PO - 4 6.2 x 10-8 HPO 4 2- HC 2H 3O 2 NH 4 + 2.1 x 10-13 1.8 x 10-5 5.8 x 10-10 For each of these buffers you will be required to: a. choose which weak acid/conjugate base pair will be best suited for the assigned ph; b. calculate the amounts of reagents needed; c. prepare the buffer and then measure the buffer s ph. You are given the following buffer systems and reagents to choose from for preparation of these three buffers (NOTE: You must determine which ph each of these buffer systems is best suited for; they are not necessarily listed in the same order): 1. One buffer will contain a mixture of ammonium ion and ammonia. This buffer is to be prepared by mixing 50.0- ml of 0.4 M ammonium chloride with the correct volume of 1 M NaOH. You will determine the volume of the 1M NaOH needed. 2. A second buffer will be a phosphate buffer solution. This buffer is to be prepared by mixing 50.0- ml of 0.4 M NaH 2PO 4 with the correct volume of 1 M HCl. You will determine the volume of the 1 M HCl needed. 3. The third buffer will contain a mixture of acetic acid and acetate ion. This buffer is to be prepared by mixing 25.0- ml of 0.1 M HCl with the correct mass of solid sodium acetate. You will determine the mass of solid sodium acetate needed. To assist the class in getting started, during lab lecture the lab instructor will help the class determine which buffer solution from the above choices is best suited for the ph = 9.50 buffer. The instructor will also help the class calculate the amounts of reagents needed for this buffer. ph = 9.50 Buffer 1. From the choices given above, determine which one the three buffer solutions should be used to make a buffer of ph = 9.50. Record your choice in your notebook. How did you decide? 2. In your notebook: a. Write the chemical equation for the acid- ionization equilibrium reaction for this buffer solution. b. Write the net- ionic equation for the chemical reaction that will be used to prepare this buffer. c. Record the actual concentrations of the available stock solutions. d. Using the actual stock solution concentrations, calculate the amount of each reagent needed to prepare this buffer. 3. Prepare the buffer. 4. Prepare the ph sensor for data collection. Go to the Experiment menu and under Calibrate choose the channel with the ph probe (ex: CH1:pH ). In the window that appears make sure the Calibration tab is chosen. Click on Calibrate Now. Rinse the ph meter with copious amounts of deionized water. Carefully blot dry. Dip the ph meter in the ph 7 standard solution. Wait until the voltage reading stabilizes. Then, in the field beneath Enter Value enter the ph value of the solution and click Keep. Repeat the process with the ph 10 standard solution. When finished with this step click on Done to close the window.
60 Chem 1B Dr. White 5. Measure and record the ph of the buffer. 6. Pour the buffer solution into your waste collection beaker ph = 5.00 Buffer Follow the same steps for the ph 9.50 buffer but for a buffer with a ph = 5.00. Calibrate again, but this time use the ph 4 and 7 standard solutions. ph = 2.00 Buffer Follow the same steps for the ph 9.50 buffer but for a buffer with a ph = 2.00. There should be no need to recalibrate unless the ph measurements seem off. Clean-up 1. Discard any left over reagents and the solution in your waste collection beaker in the appropriately labeled waste container. 2. Rinse the ph electrode, place it back in its storage solution.
Chem 1B Dr. White 61 Name: Lab Days/Time: Part A: Buffer Capacity and Range Experiment 7 Buffering Capacity and Preparation of Buffers Data and Results Report the actual concentration of the undiluted stock buffer solution: Report your measured initial ph of the undiluted buffer: Report your measured initial ph of diluted buffer: Report the literature pk a for the buffer s weak acid component (see table 1 on page 2): 1. How does the initial ph of the diluted buffer compare to the ph of the undiluted buffer? a) Is this what you would expect? Explain. Print your graph from Logger Pro and attach it to the report. Include a legend on the graph to identify each curve. 2. Which buffer had the greater capacity? How can you tell? a) Is this what you expect? Explain.
62 Chem 1B Dr. White Part B: Buffer Preparation ph = 9.50 Buffer 1. What weak acid/conjugate base pair did this buffer contain? 2. Write the chemical equation for the acid dissociation equilibrium reaction (i.e. the reaction of the acid with water) for this buffer solution. 3. Write the net- ionic equation for the chemical reaction that was used to prepare this buffer. 4. Determine the reagents used to make this buffer along with the amounts of each reagent and, where appropriate, the actual concentrations of the solutions used. List all of the reagents and amounts used to prepare this buffer below: 5. Report your measured ph of this buffer 6. Does the measured ph agree with the calculated ph, within ±0.10 ph units? 7. If the measured and calculated ph values do not agree, which of the reagents used could you add more of to the buffer to make the measured ph agree more closely with the calculated ph?
ph = 5.00 Buffer 1. What weak acid/conjugate base pair did this buffer contain? Chem 1B Dr. White 63 2. Write the chemical equation for the acid dissociation equilibrium reaction for this buffer solution. 3. Write the net- ionic equation for the chemical reaction that was used to prepare this buffer. 4. Determine the reagents used to make this buffer along with the amounts of each reagent and, where appropriate, the actual concentrations of the solutions used. List all of the reagents and amounts used to prepare this buffer below: 5. Report your measured ph of this buffer 6. Does the measured ph agree with the calculated ph, within ±0.10 ph units? 7. If the measured and calculated ph values do not agree, which of the reagents used could you add more of to the buffer to make the measured ph agree more closely with the calculated ph?
64 Chem 1B Dr. White ph = 2.00 Buffer 1. What weak acid/conjugate base pair did this buffer contain? 2. Write the chemical equation for the acid dissociation equilibrium reaction for this buffer solution. 3. Write the net- ionic equation for the chemical reaction that was used to prepare this buffer. 4. Determine the reagents used to make this buffer along with the amounts of each reagent and, where appropriate, the actual concentrations of the solutions used. List all of the reagents and amounts used to prepare this buffer below: 5. Report your measured ph of this buffer 6. Does the measured ph agree with the calculated ph, within ±0.10 ph units? 7. If the measured and calculated ph values do not agree, which of the reagents used could you add more of to the buffer to make the measured ph agree more closely with the calculated ph?
Chem 1B Dr. White 65 Post Lab Questions 1. Write down at three methods you could use to prepare a H 2PO 4 /HPO 4 2 buffer. In each case give the chemical formulas of the compounds you would mix and identify a limiting reactant if there is one. 2. Identify which of the following will result in a buffer solution when equal volumes of the two solutions are mixed. (Circle all that apply.) (a) 0.10 M NaNO 2 and 0.10 M HNO 2 (b) 0.10 M HCl and 0.10 M NH 3 (c) 0.10 M KNO 3 and 0.10 M HNO 3 (d) 0.20 M HCl and 0.10 M NH 3 (e) 0.10 M HCl and 0.20 M NH 3
66 Chem 1B Dr. White 3. A buffer is made by dissolving a combination of Na 3PO 4 and Na 2HPO 4 in enough water to make 250.0 ml of solution. The ph of the resulting solution is 12.25. a. Which component of the buffer is present in greatest amount? How can you tell? (HINT: use the Henderson- Hasselbach equation to determine the ratio of base:acid) b. If the concentration of PO 4 3 is 0.400 M, what mass of Na 2HPO 4 is present? c. Which component of the buffer must be added to change the ph to 12.50? What mass of this component must be added? Assume no volume change.