Buffers-Day 2 (More About Buffers)

Name: Objectives Buffers-Day 2 (More About Buffers) 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 can be prepared from a weak acid-conjugate base pair. Buffer solutions are able to maintain a relatively constant ph since they contain both a weak acid and a weak base component. Reacting with the weak base neutralizes strong acids. Reacting with the weak acid neutralizes strong bases. Buffers work by the action of the common ion effect. The equilibrium reaction, for any buffer, may be written as the acidionization (K a ) equilibrium HA(aq) + H 2 O(l) H 3 O + (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 2 O(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 3 O + (aq) HA(aq) + H 2 O(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 [base] $ # & (4) " [acid] % gives a good approximation for determining the ph of a buffer solution. In this equation [acid] is the concentration of the weak acid component of the buffer, [base] 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. There are many recipes for preparing buffers, in the Handbook of Chemistry and Physics, for example. 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 = pk a + log(0.1) to ph = pk a + log(10). There are three easy methods to prepare simple buffer solutions: 1. 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 [base]/[acid]: 2. 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). 3. 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). Foothill College-Larson/Daley 1 Revised/Printed 1/14/08

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 3 PO 4, and (4) test buffering capacity and effective buffering range by adding HCl and NaOH to a buffer. Reagents Available 1-M HCl Solid sodium acetate, NaCH 3 COO (May be a hydrate, check label!) 1-M NaOH 0.4-M NH 4 Cl 0.1-M HCl 0.4-M Na 2 HPO 4 0.1-M NaH 2 PO 4 /0.1-M Na 2 HPO 4 buffer (prepared by stockroom) Solution concentrations given above are approximate. The stock solutions will be standardized and the actual concentrations written on the containers. You must record the actual concentrations in your lab notebook. Procedure NOTE: 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. Prepare a ph sensor for data collection. See the LabPro Quick Start Guide for details. 2. Calibrate the ph sensor with ph 4 and ph 10 buffers. 3. When not in use, place the ph sensor back into its storage solution. 4. Rinse the ph electrode thoroughly with distilled water and gently pat it dry before placing it in a solution to measure the ph. 5. DO NOT use the ph electrode to stir solutions. 6. As you observe the ph readings allow at least 15 seconds, or until the reading stabilizes before recording your measurements. 7. Read and record the ph values from your calculator screen to two decimal places! (That means two digits after the decimal point.) Part 1: Buffer Capacity and Range-An Investigation Using a H 2 PO 4 and HPO 4 2 Buffer Solution Undiluted H 2 PO 4 and HPO 2 4 buffer 1. In your lab notebook, record the concentration (from the container) of the H 2 PO 4 /HPO 2 4 stock buffer solution, the HCl stock solution and the NaOH stock solution. 2. Measure two different 25-mL portions of the stock buffer solution into two different 50-mL beakers. 3. Using one of these 25 ml portions: b. Add drop wise 10 drops of the HCl solution, recording the ph after each drop. 4. Using the second 25 ml portion: b. Add drop wise 10 drops of the NaOH solution, recording the ph after each drop. Diluted H 2 PO 4 and HPO 2 4 buffer 1. Measure 10-mL of the H 2 PO 4 /HPO 2 4 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 H 2 PO 4 and HPO 2 4 buffer into two different 50-mL beakers. 5. Using one of these 25 ml portions: b. Add drop wise 10 drops of the HCl solution, recording the ph after each drop. 6. Using the second 25 ml portion: b. Add drop wise 10 drops of the NaOH solution, recording the ph after each drop. Foothill College-Larson/Daley 2 January 14, 2008

Part 2: Preparation of Buffer Solutions In this part of the experiment you will prepare three different buffer solutions: one with ph = 9.50, a second with ph = 5.00 and a third with ph= 2.00. For each of these buffers you will be required to: choose which weak acid/conjugate base pair will be best suited for the assigned ph; calculate the amounts of reagents needed; 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: 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 1 M 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 Na 2 HPO 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, 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. Measure and record the ph of the buffer. 5. Pour the buffer solution into your waste collection beaker ph = 5.00 Buffer 1. From the choices given, determine which one the three buffer solutions should be used to make a buffer of ph = 5.00. 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. a. Record the actual concentrations of the available stock solutions. c. Using the actual stock solution concentrations, calculate the amount of each reagent needed to prepare this buffer. 3. Prepare the buffer. 4. Measure and record the ph of the buffer. 5. Pour the buffer solution into your waste collection beaker. ph = 2.00 Buffer 1. From the choices given, determine which one the three buffer solutions should be used to make a buffer of ph = 2.00. Record your choice in your notebook. How did you decide? 2. In your notebook: a. Write the chemical equation for acid-ionization equilibrium reaction for this buffer solution. a. Write the net-ionic equation for the chemical reaction that will be used to prepare this buffer. b. Record the actual concentrations of the available stock solutions. d. Using the actual stock solution concentrations, c alculate the amount of each reagent needed to prepare this buffer. 3. Prepare the buffer. 4. Measure and record the ph of the buffer. 5. Pour the buffer solution into your waste collection beaker. 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 and then return it to the stockroom. Foothill College-Larson/Daley 3 January 14, 2008

Report Sheet: Buffers Day 2 Name: Partner: For calculations, remember to always show your work, with units in the set-up, and report answers to the correct significant figures. Part 1: Buffer Capacity and Buffer Range-Investigated Using a H 2 PO 4 and HPO 4 2 Buffer Solution 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: 1. How do the measured phs compare to the pk a? a) Is this what you expected? Explain. 2. How does the ph of the diluted buffer compare to the ph of the undiluted buffer? a) Is this what you would expect? Explain. 3. Make a graph of ph versus drops of added HCl and added NaOH for each buffer, the undiluted and the diluted. Plan the scale of your graph so that all four curves can be shown on the same graph. Include a legend on the graph to identify each curve. Attach your graph to this report. a) Which buffer had the greater capacity? Is this what you expect? Explain. b) Based on your graph, at approximately how many ph units from pk a of the weak acid component does the buffer begin to fail? c) Is this what you would expect? Explain. Foothill College-Larson/Daley 4 January 14, 2008

Part 2: Preparation of Buffer Solutions ph = 9.50 Buffer 1. What weak acid/conjugate base pair did this buffer contain? 2. Write the chemical equation for the acid-ionization equilibrium reaction for this buffer solution. 3. Write the net-ionic equation for the chemical reaction that was used to prepare this buffer. 4. List all of the reagents used to make this buffer along with the amounts of each reagent and, where appropriate, the actual concentrations of the solutions used. 5. Report your measured ph of this buffer 6. Does the measured ph agree with the calculated ph, within ±0.10 ph units? 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? 2. Write the chemical equation for the acid-ionization equilibrium reaction for this buffer solution. 3. Write the net-ionic equation for the chemical reaction that was used to prepare this buffer. 4. List all of the reagents used to make this buffer along with the amounts of each reagent and, where appropriate, the actual concentrations of the solutions used. 5. Report your measured ph of this buffer 6. Does the measured ph agree with the calculated ph, within ±0.10 ph units? 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? Foothill College-Larson/Daley 5 January 14, 2008

ph = 2.00 Buffer 1. What weak acid/conjugate base pair did this buffer contain? 2. Write the chemical equation for the acid-ionization equilibrium reaction for this buffer solution. 3. Write the net-ionic equation for the chemical reaction that was used to prepare this buffer. 4. List all of the reagents used to make this buffer along with the amounts of each reagent and, where appropriate, the actual concentrations of the solutions used. 5. Report your measured ph of this buffer 6. Does the measured ph agree with the calculated ph, within ±0.10 ph units? 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? Follow-up Questions 1. Write down at least three methods you could use to prepare a H 2 PO 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. Foothill College-Larson/Daley 6 January 14, 2008

2. A buffer is made by dissolving a combination of Na 3 PO 4 and Na 2 HPO 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? b) If the concentration of PO 4 3 is 0.400 M, what mass of Na 2 HPO 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. Foothill College-Larson/Daley 7 January 14, 2008