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1 CHEM 152 WINTER 2010 POTENTIOMETRIC TITRATIONS Fill-in, Prelab attached (p 12) Name Turn in your graphs and pages 9-12 ONLY LEARNING OBJECTIVES: After completing this experiment, you should feel comfortable: Calibrating a ph electrode to prepare for a potentiometric titration. Titrating a strong acid with a strong base, weak acid with a strong base, and a polyprotic acid with a strong base. Differentiating acid strength by the shape of a titration curve. Using titration data to determine the concentration of an unknown solution of an acid. Identifying equivalence and half-equivalence points on a titration curve. Relating molarity to normality for a polyprotic acid. Determining acid strength/pk a values off of a titration curve. TO EARN YOUR FINAL STAMP: The following items must be completed in lab. You may complete the entire assignment in the lab; this reflects the minimum required to earn your final stamp. Collect and process all data for the titration of HCl with NaOH Collect and process all data for the titration of H 3 with NaOH Determine the concentration of HCl and H 3 from your titration data, and enter the data in the lab computer for the computation of class averages. Determine the pk a1 and pk a2 for H 3, based on your titration curve. Introduction In the experiment today you will be performing a potentiometric titration to determine the unknown concentrations of two acids. In order to perform this titration you will utilize a ph meter, an instrument which allows you to directly measure the strength and concentration of an acid. Discussion ph meters (see Figure 1) operate by measuring the difference in voltage generated between and indicator electrode and a reference electrode. The reference electrode or probe is not sensitive to H + concentration changes and is simply used as a reference point. This electrode produces a constant voltage (E = +0.2 V) which is provided by the following reaction: Hg 2 Cl 2 + 2e - 2Hg Cl - The sensing probe, which is H + ion sensitive, consists of a glass bulb filled with dilute HCl into which is inserted a silver or platinum wire. (See Figure 1) Exp #8 Potentiometric Titrations Rev W10 NF Winter 2010 Page 1 of 12

2 The glass bulb is permeable to H + ions but not to other ions. If the solution into which the electrode is placed is more concentrated than that in the bulb, H + ions will move from the solution into the bulb. The HCl solution inside the bulb will then have an excess of H + ions and will be positive with respect to the solution being measured This potential difference across the glass membrane can be measured then compared to the reference voltage and a ph determined for the measured solution. Quite often the reference and indicator electrodes are combined into one probe called a combination probe or electrode. It operates the same way as the two probes. Method for Determining Concentration In this experiment you will be titrating both a strong monoprotic acid and a weak polyprotic acid with a standardized strong base. You will be adding the standardized base to the acid and measuring the resulting ph. When the volume of base added is plotted on a graph as a function of ph, a titration curve is produced. From this curve it is possible to determine the milliliters of base needed to neutralize the amount of acid present in the solution of unknown concentration. The shape of the titration curve is dependent on the type of acid being titrated (i.e., strong or weak acid). The two situations are described below. Strong Monoprotic Acids Figure 1 When a strong monoprotic acid, such as HCl, is dissolved in water it totally ionizes into H + and Cl - ions. As a strong base, such as sodium hydroxide, is added to the acid solution, the available hydroxide ions combine with some of the available H + ions to form water. If we continue to add base and record the resulting ph, we produce, upon graphing a titration curve similar to that shown in Figure 4. HCl + H 2 O H +1 (aq) + Cl -1 (aq) H + + OH - H 2 O Exp #8 Potentiometric Titrations Rev W10 NF Winter 2010 Page 2 of 12

3 Figure 2 (accessed February 27, 2008) Notice that until we reach the neutralization or end point, the ph of the solution stays nearly constant. We know that ph is a measure of H + ion concentration (ph = log[h + ]). Since our strong acid totally ionizes, we have a large reserve of H + (actually H 3 O + ) ions available in solution to instantly react with any added base. Because we have such a large reserve of H +, the addition of a few milliliters of base will cause only a small change in the total H + ion concentration. We also know that ph + poh = 14. Since any hydroxide ion that is added is immediately neutralized, the ph remains relatively unchanged. Once the neutralization point is reached, however, the available reserve of H + ions is depleted. The concentration of H + drops dramatically, while the concentration of hydroxide ion increases. This causes the ph value to rise. The ph value continues to rise until the ph of the base being added is reached. Once reached, addition of more base has little effect and the ph again becomes relatively constant. It is possible to determine the endpoint for the neutralization reaction, and therefore the base needed to neutralize the acid, simply by extending a line from the vertical portion of the titration curve to the x-axis (see Figure 2). Since we are given the normality of the standard base and the volume of acid used, the normality of the acid can be calculated using the relationship: Weak Polyprotic Acids N acid V acid = N base V base Weak polyprotic acids such as phosphoric acid (H 3 ) do not totally ionize in solution. In fact, most of the phosphoric acid remains as unionized H 3 molecules in solution. The small amount that does ionize exists as a series of equilibrium equations, each of which produce some H + ions. Reaction #1 H 3 H + + H 2 Reaction #2 Reaction #3 H 2 H 2 H + + H 2 H Ka 1 = [H 2PO 4 ][H + ] = [H 3 ] Ka 2 = [HPO 2 4 ][H + ] = [H 2 PO 4 ] Ka 3 = [PO 3 4 ][H + ] = [HPO 2 4 ] Exp #8 Potentiometric Titrations Rev W10 NF Winter 2010 Page 3 of 12

4 We can also perform a potentiometric titration on H 3 (see Figure 3). Note in this graph that the ph value slowly increases and that the ph transitions through the end points of the acid are not abrupt but instead are gradual. The third equivalence point is not observed. Note the volumes (and relationships) at the equivalence points. Figure 3 ½ equivalence points: the volume between (½ way between each eq. pt) If the volume to get to the first eq. pt = ml then the ½ eq. pt = 5.00 ml. At this point, ph = pka 1. If the volume to get to the second eq. pt = ml then the ½ eq pt = ( ml)/2 = ml. At this point, ph = pka 2. On this graph, ph = pka 1 = 1.9 and ph = pka 2 = 7.3 Exp #8 Potentiometric Titrations Rev W10 NF Winter 2010 Page 4 of 12

5 When we titrate H 3 with a strong base (i.e., NaOH), the following reactions would occur. Or, more accurately, H 3 + NaOH NaH 2 + H 2 O NaH 2 + NaOH Na 2 H + H 2 O Na 2 H + NaOH Na 3 + H 2 O H 3 H + + H 2 NaOH (aq) NaH 2 + H 2 O Where there is no large readily available reserve of H + ions as were present with the totally ionized strong acid, there are three sources where H + ions can be obtained as needed. We have shown that the weak polyprotic acids exist as a series of equilibrium equations. If we shift the equilibrium out of balance, the systems will adjust to reestablish the equilibrium. When the available H + is used to neutralize added base, some unionized acid from each of the three equilibrium equations must ionize to reestablish the equilibrium of the hydrogen ion to near the original value. Since the [H + ] is changing very little, the ph also changes very little. Mixtures of weak acids and their salts stabilize hydrogen ion concentration upon addition of small amount of base by shifting the various equilibrium equations. This is also known as buffering the solution. The size of the ionization (equilibrium) constant is a measure of the degree of ionization. The larger the value, the greater the ionization. For this reason, K 1 is going to supply more H + ions than is K 2 or K 3. Therefore, as the sodium hydroxide is added, reaction #1 will provide more equilibrium concentrations of H + than will reaction #3. Reaction #1 will then be used up first (i.e., its endpoint will be reached first). We do not, however, get a sharp vertical rise at the endpoint as we did with the strong, monoprotic acid. This is because reactions #2 and #3 are still helping to moderate the change in [H + ] and thus, the ph. After the H + from reaction #1 is used, the ph slowly rises to the H + ion value that can be supplied by the equilibrium reactions #2 and #3. Finally, after all H + from reaction #2 is used, we see a third rise in ph. However, because of the very small size of K 3 it is often difficult to obtain an accurate measurement for the titration of reaction #3. You will be asked to perform a potentiometric titration of both HCl and H 3 using the strong base, NaOH, in this experiment. Exp #8 Potentiometric Titrations Rev W10 NF Winter 2010 Page 5 of 12

6 Experimental The standardized base and stir bars can be found on the reagent bench. The magnetic stir plates are under the middle hood. The unknown acids will be in burets for you to use. Measure out about 20 ml and record the volume to the nearest ±0.01 ml. Refill the burets from the bottles/transfer beakers beside the apparatus. The ph electrodes and interfaces are next to the computers CAUTION: Do not spill liquids on the computer or interfaces. Put your set-up together carefully. Place magnetic stir plates as far away from the monitors as possible. PROCEDURE General description- You will be doing two titrations using a ph electrode as a sensor and typing in the volumes you read on the buret. For each acid, HCl and H 3, you will determine the volume at each equivalence point and calculate the molarity for that acid. You must calibrate the electrode before you begin. Technique Tip: Rinse the electrodes between solutions by dipping them into a beaker of distilled water or use a wash bottle to rinse the solution on the electrode into a waste beaker. Setting up the computer: 1. Obtain and wear goggles. Set up the equipment as shown in Figure 4 with the electrode attached to the ph amplifier. Instead of the ph meter you will have your computer, interface and ph amplifier. 2. At the computer, open the Vernier/LoggerPro software using the start menu. Then File Open Experiment Chemistry with Computers 25 Titration dip acid. Make sure the vertical axis of the graph is ph scaled from 0 to 14 ph units and the horizontal axis reads 0 to 50mL. Change the axes as necessary. Calibrating the ph electrode: 3. Obtain two buffer solutions, one each of buffer ph 4 and buffer ph 7. Do not contaminate or dilute these solutions! They can be reused by other groups. Place the ph probe into the ph 4 buffer solution. Go to the menu and under the Experiment open Calibrate. Click Calibrate now. The electrode will be reading a voltage when this value has stabilized, type in the ph of your buffer (4.0) into the Enter Value box under Reading 1 and then click Keep. Rinse your electrode and place it into the ph 7 buffer solution. When the voltage has stabilized again, enter that ph value (7.0) in the box under Reading 2 and click keep. Click done to return to the graph and data table screen. Rinse the electrode and place the probe back into the ph 4 buffer to make sure it is stable. If the electrode is properly calibrated, the ph of the solution should read 4.00 ± 0.05 units. If your ph value is significantly different (±0.05 ph units) than the buffer you are checking, or if the value drifts slowly to a higher or lower number, contact your lab instructor. Preparing solutions and setting up the titration apparatus: 4. Obtain a MAXIMUM of 100 ml of NaOH in a clean and dry beaker. You will use this NaOH for both experiments. Record the precise concentration of the NaOH solution in the Data and Calculations table. 5. Obtain a 50-mL buret and rinse the buret with about 5 ml of the ~0.1 N NaOH solution. Roll the NaOH in the buret to coat and rinse all sides, and drain the NaOH through the tip into a waste beaker. Use a buret clamp to attach the buret to the ring stand as shown in Figure 4. Fill the buret a little above the 0.00-mL level of the buret. Drain a small amount of NaOH solution so it fills the buret tip and leaves the NaOH at the 0.00-mL level of the buret. You do not need to start exactly at 0.00-mL- you can record the initial volume and subtract it from the final, to get your total volume. Exp #8 Potentiometric Titrations Rev W10 NF Winter 2010 Page 6 of 12

7 Use a buret clamp, not a regular clamp as shown in the figure! Figure 4 6. Place about 100 ml of distilled water into a 250 ml beaker. From a buret already set up for you at the reagent bench, measure to the nearest ±0.01 ml, 20 ml of an acid, HCl, of unknown concentration directly into the beaker. 7. Place the beaker with your acid on a magnetic stirrer and add a stirring bar. Titration 8. Use a three prong utility clamp to suspend a ph electrode on a ring stand as shown in Figure 4. Position the ph electrode in the HCl (or H 3 ) solution and adjust its position toward the outside of the beaker so that it is not struck by the stirring bar. 9. You are now ready to begin the titration. This process goes faster if one person manipulates and reads the buret while another person operates the computer and enters buret readings. Before adding any NaOH titrant, click on the Collect button and monitor ph for 5-10 seconds. After the ph has stabilized, click on the Keep button. Enter the current buret level, You have now stored the first data pair. Add enough NaOH to raise the ph by about 0.20 units. After the NaOH has been added and the ph has stabilized, click on the Keep button. Enter the new buret reading, to the nearest 0.01 ml. Be sure that the ph has stabilized! There is a delay between the addition of NaOH and a stable ph reading. Continue adding NaOH solution in increments that raise the ph about 0.20 units and enter the buret reading after each addition. By watching your graph, it will be easy to see when your ph starts jumping significantly for a small amount of base added. Adjust the amount of NaOH added to keep the ph changes at 0.20 units. Slow down your titration and collect data. Eventually you will collect every one or two drops to get good data for the equivalence point regions. Once again, be sure that the ph has stabilized! When you have reached a flat region past the equivalence point for HCl or when the ph is greater than 11 for H 3 you may Stop the titration. 10. Change the title of your graph appropriately. Click on the graph and from the File menu choose Print window and proceed to print the graph. Click on the data table window and again choose print window and print the data table. You may want to include your name in the title. 11. Double check the electrode with a buffer solution and make sure it is still calibrated. Repeat the titration with the H 3 (using about 20 ml of acid, read and recorded to the nearest 0.01 ml). **** Dispose of the beaker and buret contents and any other solutions in the collection bottles. Exp #8 Potentiometric Titrations Rev W10 NF Winter 2010 Page 7 of 12

8 Processing the Data 1. Use your graph to determine the volume of NaOH titrant used at the equivalence point of your acid. For HCl you will have one equivalence point and therefore one volume of NaOH and for H 3 you will see two equivalence points, so you will have two volumes. To do so, examine the data table to find the largest increase in ph values during the drop-wise additions of NaOH. Make sure that it coincides with the steepest slope on your graph. The equivalence point is the midpoint of the jump in the titration curve. You may use an actual data point, if you have one near that midpoint, or you may average the two points that straddle the midpoint. To do so, find the NaOH volume just before this jump. Then find the NaOH volume from the data after this ph jump. Underline both of these data pairs on the printed data table. Record these volumes in the data table. 2. Calculate the molarity of your acids using the equivalence points from your graph or data. With HCl there will be only 1 equivalence point; with H 3 there will be two so you can calculate 2 values for M. Calculate the molarity of the phosphoric acid at the second equivalence point by taking the volume at the first equivalence point and subtracting it from the volume of the second use this subtracted volume to calculate the molarity at the second equivalence point. 3. On your printed graphs, clearly label/specify the position of the equivalence point volumes you determined in Step 2, using dotted reference lines like those in Figures 2 and 3. Indicate with an x the NaOH volume of each equivalence point on the horizontal axis of the graphs. 4. On your printed graphs, clearly label the two half-equivalence points (see Figure 3 as a reference). Reading your graphs, determine pka 1 and pka 2 and report these values in the data table. (report your values to 2 sf) 5. Record your concentration data in the spreadsheet on the indicated computer and compare your results to the class average. Either the lab instructor or your lecture instructor will provide you with the class averages when they become available. Exp #8 Potentiometric Titrations Rev W10 NF Winter 2010 Page 8 of 12

9 Report Sheet Lecturer Potentiometric Titrations Name Stamp here Partner Data and Calculations -Attach all data tables and labeled graphs (see previous page) This table must be completed in lab to earn a final stamp. HCl Titration. H 3 Titration Initial volumes of Acid Concentration of NaOH Volume of NaOH at Eq.Pt. 1 Volume of NaOH at Eq.Pt. 2 Molarity of acid. Show calc. below. Molarity of acid. At eq. pt. 2 N/A for HCl N/A for HCl E.P. 1 E.P. 2 *For Vol. 2, you will subtract the Vol. at Eq. pt 1 from that at Eq. pt. 2 Volume 2* pk a1 = pk a2 = Your Average M for H 3 = Calculations for M of HCl- show formulas and all units. For instructor use only: [HCl] = Calculations for M of H 3 - show formulas and all units. [H 3 ] = initials: Class average for M of HCl Class average for M of H 3 Exp #8 Potentiometric Titrations Rev W10 NF Winter 2010 Page 9 of 12

10 Questions: (may be completed in lab or at home) 1. Calculate the % difference between your value for the concentration of the HCl and that of the class average. 2. Do the same for the H Based on your data, what is the normality of the HCl solution? 4. Taking into account the fact that H 3 is triprotic, and that you ve only examined the first two equivalence points, use your data to estimate the volume of base that would be required to reach the third equivalence point. Using that estimated value, calculate the normality of the H 3. Exp #8 Potentiometric Titrations Rev W10 NF Winter 2010 Page 10 of 12

11 5. Using the K a values provided in the lab, calculate pk a1 and pk a2. (report your values to two sf) 6. Calculate the % difference between the calculated pk a values (consider these to be the ideal values) and the values you read off your graph. (pay attention to sf in this calculation!) Exp #8 Potentiometric Titrations Rev W10 NF Winter 2010 Page 11 of 12

12 Prelab Assignment Stamp here Name 1. Write the ionic and net ionic equation for the reaction of HNO 3 (aq) with NaOH(aq). 2. The following graph shows the equivalence points for a diprotic acid. Label the axis and indicate on the graph the equivalence points and the ½ equivalence points 3. H 2 SO 3 (aq) + 2 NaOH (aq) 2 H 2 O (l) + Na 2 SO 3 (aq) a. Is sulfurous acid a strong or weak acid? b. What is the definition of a weak acid? c. Write the ionic and net ionic equation for the reaction. d. If ml of N NaOH solution is used to titrate ml of H 2 SO 3 (aq) to the first equivalence point, what is the molarity of the acid? (hints: review your acids and bases lab! And at the first eq. pt. how many H + ions have reacted? ) e. What total volume of NaOH solution would be needed to get to the second equivalence point? ml f. Use your answer to part d to compute the normality of the H 2 SO 3. Exp #8 Potentiometric Titrations Rev W10 NF Winter 2010 Page 12 of 12

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