Lab 8 11600 lab manual_CH 6_Acid-Base Equilibria_2019-2020

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103 CHAPTER 6 ACID–BASE EQUILIBRIA Introduction: Acid–Base Titrations Reactions between acids and bases that are dissolved in water occur almost instantaneously; they occur as fast as the two solutions can be mixed. These reactions also tend to go to completion, reacting until all of the limiting reagent is consumed. When exact stoichiometric amounts of acid and base have been mixed, the reaction is said to have reached the equivalence point . Essentially all of the acid has reacted with the base, and vice versa. The technique of slowly adding an acid to a base—or vice versa—until the reaction has reached the equivalence point is known as a titration . In this lab, you will perform two types of acid–base titrations: colorimetric titrations and pH (or potentiometric) titrations. Colorimetric Titrations One way to determine the equivalence point of a titration is the use of an indicator to show when the equivalence point of the reaction has been reached. Indicators take many forms, but often are substances whose solutions change color due to changes in pH. The most familiar indicators are litmus and phenolphthalein. Phenolphthalein is colorless in acid and pink in base. In theory, the indicator should turn color at the equivalence point. In practice, each indicator has an endpoint — the pH at which it turns color—that might be slightly different from the pH for the equivalence point of the reaction. Titrations of a strong acid with a strong base that use phenolphthalein as the indicator, for example, should be stopped just before the solution turns a permanent pink color. In this fashion, the endpoint of the indicator (pH 8.3) is brought as closely as possible to the equivalence point of the reaction (pH 7 for strong acid– strong base titration). In this lab, your first colorimetric titration will be a “ scout ” titration , in which you will add the titrant continuously at a moderate rate to find the approximate volume at the equivalence point. This “scout volume” provides an estimate which allows you to conduct subsequent titrations with better accuracy. The second colorimetric titration will be done more carefully and slowly to find a more accurate measurement of the volume at the equivalence point. pH (Potentiometric) Titrations In a potentiometric (pH) titration, the equivalence point is determined using a pH titration curve, which is a plot of pH (y-axis) vs. volume of titrant added (x-axis). The equivalence point is estimated from the inflection point

STUDENT NOTES 104 Chemistry 11600 Laboratory Manual of the titration curve. In this lab, you will use the Logger Pro program to record and plot the data. Indicators are not used in potentiometric titrations since they have acid–base characteristics. Figure 6.1 shows a pH titration curve for the titration of acetic acid with NaOH. The sharp rise in pH indicates the equivalence point, the point at which all of the acetic acid has reacted with the NaOH that has been added. Figure 6.1. pH Titration curve for the titration of acetic acid with NaOH. Finding the Equivalence Point A mathematical method to determine the equivalence point is by using the 1st or 2nd derivative of the pH titration curve data. 0

STUDENT NOTES Chapter 6 • Acid–Base Equilibria 105 The Logger Pro program used in this experiment allows you to calculate and view the 1 st and 2 nd derivative data. Using the 1 st derivative data, the equivalence point is the plot maximum. Find the x-value (mL of titrant) where the 1 st derivative curve reaches its maximum. Using the 2 nd derivative data, the equivalence point is the point where the curve crosses zero on the x-axis between the curve’s maximum and minimum y-value. Find the value (mL of titrant) where the 2 nd derivative curve crosses zero between the maximum and minimum y-value. Determining the pK a of a weak acid from a titration curve A pH titration curve can also be used to determine the equilibrium constant (K a ) for a weak acid (see Figure 6.2). At the halfway point between the beginning of the titration and the equivalence point, half of the acid has been converted to the base. Therefore the concentration of remaining acid is equal to the concentration of conjugate base present. At this point (the ½ equivalence point), pH = pK a , as can be seen when you consider the K a expression for the weak acid. K a = at the ½ equivalence point, [HA] = [A – ] so K a = [H 3 O + ] and pK a = pH Figure 6.2. Determining the K a from a potentiometric titration curve. More information about acid–base titrations, equivalence points, and the use of indicators can be found in your textbook. [H 3 O + ][A – ] [HA] pK a pK a

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STUDENT NOTES 106 Chemistry 11600 Laboratory Manual Part I. Analysis of Acids Using a Base Although stomach acid contains hydrochloric acid, a strong acid, most acids found in biological systems are weak acids. Among the simplest of these acids is acetic acid, CH 3 COOH, the component that gives white vinegar its sharp odor and sour taste. Other types of vinegar are also solutions of acetic acid, but they contain other compounds which give them various colors and more complex flavors. All commercial vinegars are required by law to contain at least 4% by mass of acetic acid. If we examine the ingredients on the label of a bottle of vinegar, we generally find that the vinegar contains about 5% by mass of acetic acid in water (often listed as “5% Acidity”). Percent by mass (mass percent) is calculated as follows: mass percent = mass of solute (acetic acid) mass of solute (acetic acid) + mass of solvent (water) x 100 After you determine the concentration of acetic acid in the vinegar solution in M (moles of acetic acid per L of solution), you can convert that value to mass percent. For this calculation, you will assume that the density of vinegar is 1.00 g/mL. Experimental Design In the first titration sequence, you will titrate a solution of hydrochloric acid of known concentration with a sodium hydroxide solution of unknown concentration and calculate the concentration of sodium hydroxide. Then, in the second titration sequence, you will use this standardized sodium hydroxide solution to titrate and determine the concentration of acetic acid in the vinegar. The vinegar is too concentrated to titrate directly, so you must titrate a carefully diluted solution. Each group is to complete a minimum of three titrations of each acid: two colorimetric and one pH titration. Groups should split into two subgroups to collect data as outlined in the table below. The subgroups will compare the results obtained from the two titration methods. Ste p Subgroup A (1 st buret) Subgroup B (2 nd buret) 1 Do scout colorimetric titration of HCl with NaOH (p. 108) Set up LabQuest 2 & calibrate pH electrode (pp. 108-111) 2 Subgroup A provides estimated equivalence point volume to Subgroup B 3 Do careful colorimetric titration of HCl with NaOH (p. 108) Do pH titration of HCl with NaOH; one person record data (volume & pH) in lab notebook (pp. 112-113)

STUDENT NOTES Chapter 6 • Acid–Base Equilibria 107 4 Make and print graphs; Prepare diluted vinegar solution for second titration 5 Work on calculations while waiting for estimated equivalence point for diluted vinegar titration (p. 115) Do scout colorimetric titration of diluted vinegar with NaOH (p. 114) 6 Subgroup B provides estimated equivalence point volume to Subgroup A 7 Do pH titration of diluted vinegar with NaOH; one person record data (volume & pH) in lab notebook (p. 114) Do careful colorimetric titration of diluted vinegar with NaOH (p. 114) 8 Make and print graphs; Complete calculations and report (pp. 115-118) SAFETY Wear your goggles at all times in the laboratory. PROCEDURE Lab work is done in groups. Each person must record a complete set of data, including (1) volumes of acid and base for each titration, (2) HCl concentration from reagent bottle, (3) brand of vinegar, and (4) dilution of vinegar, in his/her lab notebook and turn in the copy on the perforated pages from the lab notebook at the end of lab. Each group is allotted 300 mL of NaOH solution of unknown concentration, 100 mL of standardized HCl, and 20 mL of the assigned commercial vinegar. Buffers are reused. NEVER pipet directly from the reagent bottle! Pour a sample of the solution to be measured into a clean beaker and pipet from the beaker. Ź Review proper use of volumetric measuring techniques (burets, pipets, etc.) in Appendix C. All volumes should be reported to 2 decimal places. (Graduated cylinders should not be used for measuring in this experiment.) DATA COLLECTION Titrate Strong Acid with Strong Base Goal: Determine the exact concentration of NaOH in a NaOH solution by titrating it with standardized HCl. Record the letter or number of the NaOH solution of unknown concentration and the exact concentration of the HCl in your lab notebook.

STUDENT NOTES 108 Chemistry 11600 Laboratory Manual Colorimetric Titration Using a clean, rinsed pipet measure exactly 25.00 mL of the HCl solution into a clean Erlenmeyer flask. Add one or two drops of phenolphthalein to the acid. Either set up a magnetic stir plate or plan to swirl the flask throughout the titration. Support a 50.00-mL buret on a ring stand with a buret clamp. Fill the buret with NaOH before setting it over the Erlenmeyer flask containing HCl. For each new titration the buret must be filled with titrant so that the volume reading is near 0.00 mL. Ź Review the procedure for filling burets and removing air bubbles in Appendix C. Position the buret so that the stream from the buret can be directed into the flask on a magnetic stir plate, if using one. “Scout” titration: With stirring, add NaOH from the buret about 1 mL at a time to the acid sample until the color change from colorless to light pink persists for 15–30 seconds. Record all necessary measurements. Careful titration: Repeat with a second sample of HCl, adding the NaOH about 1 mL at a time until you are within 1-2 mL of the equivalence point as determined in the “scout” titration. Then decrease the addition of base to 2 drops at a time as the endpoint is reached. pH Titration Instrumentation can be a great advance in data collection. However, it seems that any gain in accuracy and precision is offset by additional time needed to care for, calibrate, and operate the equipment along with an increase in cost. At the current time, the cost to replace a pH electrode of the type you will use is $100. Care and Handling of pH Electrodes Keep the electrode moist (wet) at all times. When you are not using the electrode, it should be in the storage bottle or tube containing storage solution. Be sure the electrode is stored properly when you have finished using it. Be very careful of the tip of the electrode. The electrode has a thin membrane at the bottom that is easily broken. Never dry an electrode by rubbing it with a paper towel. Just “dab” water droplets with a lint-free paper such as a Kimwipe or AccuWipe. When using a magnetic stirring bar do not allow it to hit the electrode.

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STUDENT NOTES Chapter 6 • Acid–Base Equilibria 109 Calibration of pH Electrode You must calibrate the equipment for a pH titration. As you go through the calibration procedure, the display will not show a reasonable pH value that corresponds to the pH of the buffer solutions until both buffers have been used for the calibration process. You only need to recalibrate the equipment if you need a different pH range for a titration or if you turn off the equipment. Since we are titrating a sample of acid in both experiments in Part I, you will not need to recalibrate between titrations. The pH electrode is kept in a tube of storage solution next to the computer; connect it to the LabQuest 2 interface via the port labeled CH1. Then connect the LabQuest 2 interface to the computer via the USB port. Turn on the LabQuest 2 by pressing the power button. Open the Logger Pro program, which you will use to record pH and volume of base added as you perform pH titrations. The LabQuest 2 interface should display the LabQuest 2 logo as an indication that the equipment is connected correctly. LabQuest 2 Troubleshooting Tips Issue: Battery will not charge or hold power. Solution: Batteries are not supplied for the LabQuest 2. You must use the power supply and connect the LabQuest 2 to an electrical power outlet. Note: The LabQuest 2 does not receive power through the USB connection. Issue: Screen is not responsive to taps or acts as if you tapped the wrong location. Solution: Press and hold the Home button (not the icon) for 5 seconds or until the calibration screen appears. Tap the calibration target + in each corner and center of the screen. Logger Pro Troubleshooting Tips Issue: Logger Pro is not recognizing the Lab Quest 2 or sensor or is just generally not responding. Solution: Close all programs, disconnect the Lab Quest 2 and restart computer. Force the log off of all previous users. Then, log back in, connect the LabQuest 2 and start Logger Pro again. If this does not work, try connecting to a different USB port on the computer.

STUDENT NOTES 110 Chemistry 11600 Laboratory Manual Issue: Logger Pro does not respond or unexpectedly quits. Solution: Do not open multiple windows or instances of Logger Pro . You will need jars of pH 4 and 7 buffer for a two point calibration of the electrode. Use the buffer solutions in the jars provided. Do not transfer the buffer solutions to another container. Do not discard the buffer solutions. 1. On the toolbar, click Experiment > Calibrate > LabQuest 2: 1 CH1: pH . This will open a new window called Sensor Settings . Click Calibrate Now . Note: Once you have opened the Sensor Settings window, the pH reading in the lower left is inactive. You will only be able to see the voltage value associated with the pH electrode’s response to the hydronium ions in the buffer solution. See below. 2. Remove the pH electrode from the storage solution. Rinse the electrode with de-ionized water (use a small beaker to collect the rinses), blot it dry using a KimWipe and place the electrode in the pH 4 buffer solution. Swirl the jar gently for 15 seconds. In the box for Reading 1: enter a value of “ 4 ” and click Keep to store the calibration point. Voltage value

STUDENT NOTES Chapter 6 • Acid–Base Equilibria 111 3. Remove the electrode from the pH 4 buffer, rinse it with de-ionized water and blot dry. Place the electrode in the pH 7 buffer solution and swirl gently for 15 seconds. Enter “ 7 ” as the value for Reading 2 and click Keep to store the 2 nd calibration point. Click Done to exit the Sensor Settings window. 4. Remove the electrode from the buffer, rinse with de-ionized water and place the electrode in the tube of storage solution or de-ionized water. Your pH electrode is now calibrated. Buret (in buret clamp) NaOH solution To LabQuest 2 and computer pH electrode (in clamp) Acid solution Magnetic bar Magnetic stirrer Figure 6.3. pH titration setup. Ź Check with the subgroup doing the colorimetric titration to find out the volume of NaOH needed to reach the equivalence point. Pipet 25.00 mL of acid into a clean 100 mL beaker containing a magnetic stir bar. Add enough distilled water (about 15 mL) so that the tip of the pH electrode is immersed and far enough above the magnetic stir bar that the stir bar does not hit the electrode as the bar spins. Set the beaker on a magnetic stir plate next to the computer. Support a 50.00-mL buret on a ring stand. Fill the buret with NaOH before setting it over the beaker containing HCl. For

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STUDENT NOTES 112 Chemistry 11600 Laboratory Manual each new titration the buret must be filled with titrant so that the volume reading is near 0.00 mL. Ź Review the procedure for filling burets and removing air bubbles in Appendix C. Adjust the 50.00-mL buret filled with NaOH so that the stream from the buret can be directed into the beaker on the magnetic stir plate. Adjust the stirring rate to a moderate rotation speed. Collecting pH Titration data with the Logger Pro program The default setting for the Logger Pro program is to collect data at time intervals, but for a titration, you will want to collect the pH data and compare it to the volume of titrant instead. Change this setting by clicking on Experiment > Data Collection . From the Mode drop- down menu, select Events With Entry . Enter the following in the corresponding fields: Column Name enter “volume” Short Name enter “vol” Units enter “mL” Click Done . The pH value will be displayed in the lower left corner of the screen. To collect titration data: 1. Obtain the equivalence point of the colorimetric scout titration performed by your subgroup. This is the approximate equivalence point for your potentiometric titration. 2. Click . When the pH stabilizes, click . You will be prompted to enter the cumulative volume of titrant that has been dispensed by the buret, which at this point is 0.00. 3. Begin the titration by adding the solution of unknown base to the acid about 1 mL at a time until you are within 1–2 mL of the predicted equivalence point (see step #1 above). Wait for the pH to stabilize, and then click Keep and enter the total cumulative volume dispensed by the buret. If you erroneously click Keep before the pH has stabilized, click on the corresponding volume in the data table and delete the entry. If you make a mistake while entering the volume, you can edit the entry by double clicking on the value in the data table. If you stop the data collection prematurely, click Collect again and choose Append Run to continue recording data. 4. Continue adding titrant 1 mL at a time until you are within 1-2 mL of the predicted equivalence point , and then decrease the addition of base to 2 drops at a time . Wait only 10–15 seconds for the pH

STUDENT NOTES Chapter 6 • Acid–Base Equilibria 113 to stabilize before selecting Keep and entering the volume fater each addition. Prolonged stirring introduces CO 2 from the air and will change the pH of the solution. 5. When you are 1 mL past the equivalence point, return to adding NaOH about 1 mL at a time until you are approximately 10 mL beyond the equivalence point. 6. When you have finished the titration, press . Rinse the electrode and immerse it in the tube of storage solution. Pour the titrated solution down the drain. BE CAREFUL: Do not drop the magnetic stirring bar down the drain. Rinse the magnetic stirring bar with distilled water for reuse. Saving Your Data 1. If you are performing your data analysis in the lab using Logger Pro , choose Save As from the File dropdown menu. Give your data an appropriate file name and save to the desired destination. 2. If you will analyze your data away from the lab, choose Export As from the File drop down menu and select CSV . Give your data an appropriate file name and save to the desired destination. This will allow you to open your data with Excel or some other spreadsheet program. Your only access to the Logger Pro software outside of lab is in an ITaP computer lab, so make sure that any data you need, including any plots of the first or second derivative, are exported rather than saved. See Plotting the Derivative later . 3. Choose Clear all Data from the Data drop down menu. Titrate the Weak Acid in Vinegar with Strong Base Goals Determine the concentration of acetic acid in • a diluted solution of a commercial vinegar using the NaOH solution standardized in the previous procedure, • and the original, undiluted, commercial vinegar using the dilution factor Determine an experimental value for K a of acetic acid. In this experiment, you use the same techniques that you used for titrating a strong acid with a strong base, but this time you’ll be titrating a weak acid.

STUDENT NOTES 114 Chemistry 11600 Laboratory Manual Dropper pipet with water Add water to the 100.00 mL mark Volumetric flask Figure 6.4. Final Step to Prepare Dilute Solution in Volumetric Flask Because commercial vinegar is too concentrated to titrate directly you will need to prepare a diluted solution. Ź Review of proper volumetric measuring techniques can be found in Appendix C. Prepare the diluted solution as follows: Use a beaker to obtain about 20 mL of the vinegar brand assigned to your group. Use a 10.00-mL volumetric pipet to transfer 10.00 mL of vinegar from the beaker to a 100.00-mL volumetric flask. Add some deionized water to the flask and mix well, then carefully fill the flask to the mark with deionized water and mix well again. (To mix, cover the flask with a small square of Parafilm and invert several times.) What is the dilution factor for this dilution? Perform two colorimetric titrations and one pH titration with the standardized NaOH solution to determine the concentration of acetic acid in the diluted vinegar solution. Use 25.00 mL samples of the diluted vinegar solution for the titrations. Follow the same general procedure you used for the titration of a strong acid with a strong base. WASTE DISPOSAL AND CLEANUP The titration solutions and excess acid and base solutions can be poured down the drain and rinsed with plenty of water. Do NOT discard the pH 4 or pH 7 buffer solutions. Rinse the stir bar and return it to your instructor. BE CAREFUL: Do not pour the magnetic stir bar into the sink with the solution.

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STUDENT NOTES Chapter 6 • Acid–Base Equilibria 115 Disconnect the pH electrode and return it to the tube of storage solution. Return the LabQuest 2 instrument box to your TA, after you have checked that all of the parts are present. Rinse shared glassware with deionized water and return to the appropriate location(s). Return shared equipment to the appropriate location. Keep your splash goggles on until you have completed the data analysis, turned in your report, and are leaving the lab. Lock your lab drawer before leaving lab. DATA ANALYSIS Strong Acid with Strong Base Titrations Write the balanced chemical equation to represent the acid–base reaction that occurred in the titration of a strong acid with strong base. Numerical Analysis Component (using data from the colorimetric titration) Use data from the colorimetric titrations to calculate the molarity of NaOH in the NaOH solution. (You can review the fundamentals of solution stoichiometry calculations in your textbook.) Graphical Analysis Component (using data from the pH titration) Calculate and Graph the Derivative of your pH Titration Curve To calculate the 1st or 2nd derivative data: 1. Click on Data > New Calculated Column . Provide an appropriate title for the new column in the Name and Short Name boxes such as 2 nd derivative and d2. 2. Click inside the Expression box; then click Functions . Select the appropriate function ( calculus > derivative or calculus > secondDerivative ). 3. NOTE: The variables are separated by a comma with the y variable first then x variable. Choose Variables (Columns) > pH. Type a comma (,) then click Variables (Columns) > Volume . Click Done .

STUDENT NOTES 116 Chemistry 11600 Laboratory Manual To view the graph of the derivative and find the equivalence point: 1. Select or highlight the derivative column on the data table, then click Insert > Graph . 2. To find the equivalence point using one of the derivative methods, it is necessary to focus in on a small range of values around the equivalence point on the x-axis. The best range to view the curve is within ±0.50 mL of the equivalence point. Click and drag on the graph to select portions of the curve (see figure below). Your selection should include the maximum of the first derivative graph or the region between the maximum and minimum if you are using the second derivative graph. Click Zoom In, t o expand the axes and find the mL of titrant added to reach the equivalence point. 3. Record the volume of NaOH added to within two places beyond the decimal. 4. Print a graph showing how you determined the equivalence point from the pH titration data. a. Change orientation to landscape by selecting File > Page Setup > Landscape > OK. b. To print, click on the window containing the graph and select File > Print Graph > OK. Or to print all of the windows (i.e. data table and graphs) on the screen on a single page: File > Print > OK. c. Graphs are printed with the same relative dimensions shown on the screen. (They are not expanded to fill the page.) d. Identify and label (by hand) the equivalence point and volume of titrant at the equivalence point on your printed graph. Numerical Analysis Component (using data from the pH titration) Use information obtained from the graph and other measurements from the pH titration to calculate the concentration of the NaOH in the strong base solution.

STUDENT NOTES Chapter 6 • Acid–Base Equilibria 117 Weak Acid with Strong Base Titrations Write the balanced chemical equation to represent the acid–base reaction that occurred in the titration of the weak acid with strong base. Numerical Analysis Component (using data from the colorimetric titration) Use data from the colorimetric titration and the molarity of NaOH determined in the previous titration to calculate the molarity of acetic acid in the dilute vinegar solution. Then calculate the molarity of acetic acid in the original commercial vinegar. (Use the dilution factor you determined earlier.) Graphical Analysis Component (using data from the pH titration) Print a graph showing how you determined the equivalence point from the pH titration data. Identify and label the equivalence point and volume of titrant at the equivalence point. Numerical Analysis Component (using data from the pH titration) Using the data obtained from the graph, calculate the concentration of the acetic acid in the dilute vinegar solution. Calculate the concentration of the vinegar before it was diluted (i.e., the original commercial vinegar). (Use the dilution factor you determined earlier.) Calculate an experimental value of the % by mass (mass percent) “acidity” in the vinegar brand you analyzed. Refer to the information provided in the Introduction. Assume the density of vinegar is 1.00 g/mL. Determine the experimental K a value for the weak acid, acetic acid. (Refer to the information given in the Introduction.) Identify and label information on the graph to illustrate how you obtained the values for [H 3 O + ] and ½ equivalence point. Write your results (vinegar brand, original acetic acid concentration, and % acetic acid by mass) on the board to share with the class. Record at least one set of results for a different brand of vinegar than the one your group analyzed. RESULTS Summarize your experimental results for NaOH concentration, acetic acid concentration in vinegar and K a of acetic acid. Look up the accepted value for the K a of acetic acid in your textbook or ebook.

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STUDENT NOTES 118 Chemistry 11600 Laboratory Manual How does your experimental value of K a compare to the expected K a ? Is it of the same magnitude? DISCUSSION 1. Obtain the actual NaOH concentration value from your instructor. Calculate the % error in your titration of strong strong acid with strong base for each of the titration methods (colorimetric and pH). Based on your results, is one titration method more accurate than the other? What are the benefits and disadvantages of using pH titrations instead of the colorimetric method? % error = | experimental value-theoretical value | x 100 (theoretical value) 2. Compare the titration curves of weak acid–strong base reactions with the curve for strong acid–strong base reactions. How are they similar? How are they different? 3. When you performed the potentiometric titration of HCl with NaOH, you added about 15 mL of deionized water to the HCl solution you were preparing to titrate. In your data analysis, you did not take this added volume of water into consideration when calculating the resulting concentration of NaOH. Why not? 4. Do each of the three vinegar brands analyzed in lab satisfy the legal requirement to contain at least 4% by weight acetic acid? If not, which brand does not satisfy the requirement?. LAB RECORDS AND REPORTS Group Portion You and your partner or group will turn in one completed report, either a formal lab report or a completed lab report form at the end of this chapter. It is your responsibility as a group to ensure that everyone whose name is on the report has participated as fully as possible in the completion of the project. The report or report form is an organized summary of your work and does not replace the need to keep a complete set of lab or field notes in your lab notebook as the lab is being done and data collected. Individual Portion Each student must attach laboratory notebook duplicate pages containing a complete data set and observations for the experiment.

Chapter 6 • Acid–Base Equilibria 119 REPORT FORM Title: Names: Date: Lab Section Number: PURPOSE OR GOAL STATEMENT The goal(s) of this lab: Data and Data Analysis: NaOH + HCl Data Ź Concentration of HCl = Table 1. Titration Data for NaOH + HCl. Titration Method Titration Type Volume of Acid, mL Volume of Base, mL Colorimetric Scout Colorimetric Careful Potentiometric Careful Data Analysis Ź Equation for the titration reaction: Ź pH titration curve ( attached, properly labeled ) Ź Sample calculations: Molarity of NaOH using data from careful colorimetric titration using data from the potentiometric titration

120 Chemistry 11600 Laboratory Manual Table 2. Data analysis summary, NaOH concentration. Titration Method Titration Type [NaOH] Colorimetric Scout Colorimetric Careful Potentiometric Careful Data and Data Analysis: CH 3 COOH + NaOH Data Ź Molarity of NaOH from previous section Ź Brand of vinegar analyzed Ź Procedure for diluting the vinegar Ź Titration Data Table 3. Titration data for acetic acid + NaOH. Titration Method Titration Type Volume of Acid, mL Volume of Base, mL Colorimetric Scout Colorimetric Careful Potentiometric Careful Data Analysis Ź Equation for the titration reaction: Ź pH titration curve ( attached, properly labeled ) Ź Sample calculations for acetic acid concentrations using data from the careful colorimetric titration Acetic acid concentration in dilute vinegar solution

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Chapter 6 • Acid–Base Equilibria 121 Sample calculations continued Acetic acid concentration in the original vinegar using data from the potentiometric titration Acetic acid concentration in dilute solution Acetic acid concentration in original vinegar Table 4. Data analysis summary, acetic acid concentration. Titration Method Titration Type [CH 3 COOH] dilute [CH 3 COOH] original Colorimetric Scout Colorimetric Careful Potentiometric Careful

122 Chemistry 11600 Laboratory Manual Ź % by mass (mass percent) acetic acid in ________ (brand) vinegar ( show calculations ) Ź Experimental K a value of acetic acid ( show calculations ) RESULTS Our analyses produced the following values: Concentration of NaOH = Concentration of acetic acid in (brand) vinegar = M Our experimental K a value for acetic acid = Table 5. Results for different brands of vinegar. Group Vinegar brand Acetic acid concentration, M % by mass acetic acid DISCUSSION Answer the questions found in your lab manual. Add extra paper if needed.

Chapter 6 • Acid–Base Equilibria 123 STUDENT NOTES Part II. Analysis of Polyprotic Acids and Bases In the second part of this activity you will plan and then carry out titrations to analyze or characterize some substance(s) that you may not have worked with in the previous labs. You will use the same techniques as those used in Part I; that is, colorimetric and pH titrations. Lab work is done in groups. Each person must record a complete set of data for solution preparation and colorimetric titrations in his/her lab notebook and turn in the copy on the perforated pages from the lab notebook at the end of lab. SAFETY Wear your goggles at all times in the laboratory. Wear gloves. If you leave the lab, take the gloves off and recycle. Get new gloves when you return to lab. Avoid inhalation of ammonia vapors. Dispense the stock solution of ammonia in the main hood. Prepare the diluted ammonia solution within a student bench hood. PROCEDURE Use the same methods and equipment used in Part I to carry out the titrations described in the goals. Refer to Part I for detailed instructions on using the LabQuest 2 device. The NH 3 ( aq ) provided for you has about the same concentration as acetic acid in vinegar (i.e. dilution of the stock solution is necessary prior to titration). Titrations are to be done beginning with 25.00 mL samples of the desired solution in a beaker or flask with an additional 15 mL of distilled water to increase the volume. By knowing or estimating the pKa of the species being analyzed, the appropriate pH buffers to calibrate the pH probe can be selected. The goal is to calibrate the pH probe for the buffer region of the titration so that an accurate measurement of the pH at the ½ equivalence point can be obtained. When choosing the appropriate buffer, select one with a pH slightly more acidic than the ½ equivalence point and one with a pH slightly more basic than the ½ equivalence point. You will need to use 2 buffers to calibrate the pH probe for each of the titrations, i.e. the weak base-strong acid titration and the weak acid-strong base titration. Indicators signal the equivalence point of a titration by a change in physical property, such as color. Phenolphthalein changes color in the pH 8–10 range and works well to signal the equivalence point of reactions between strong acids or weak acids with strong bases. Methyl red changes color (yellow to pink/red) in the pH 4–6 range and works well to signal the equivalence point of reactions of weak bases with

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STUDENT NOTES 124 Chemistry 11600 Laboratory Manual strong acids. Choose the appropriate indicators for the procedures. pH titrations are done without color indicators because color indicators often have acid–base properties of their own and can change the pH values slightly. pH titration curves for diprotic acids show two sigmoidal regions if the K a values differ by several orders of magnitude. Figure 6.5. pH titration curve of diprotic acid with base. Carry out pH titrations of the unknown weak acid using small increments of base throughout the entire process in order to avoid adding too much base and possibly missing the first equivalence point should the acid be diprotic. Set up one buret for HCl and one for NaOH, so you can move burets and not have to clean and refill each buret. Use a black Sharpie to mark the contents of each buret. DATA COLLECTION Titrating a Weak Base with Strong Acid Each group has been allotted 20 mL of ammonia stock solution and 150 mL of standardized HCl to complete the analyses. Because the stock solution of ammonia is too concentrated to titrate directly you will need to prepare a dilute solution. Obtain about 20 mL of the ammonia stock solution and prepare a dilute solution in the same way you diluted vinegar in Part I (a 10-fold dilution). BE CAUTIOUS WITH THE AMMONIA; DO NOT INHALE OR SMELL IT DIRECTLY. Goals: Your group will design and carry out procedures to Titrate 25.00 mL samples of the dilute weak base, NH 3 , in two ways. Calculate and compare the concentration of the dilute weak base using data from both titration methods. Buret (in clamp) Acid solution (HCl) Base solution (NH 3 ) Figure 6.6. Set-up for Titration of a Base such as Ammonia with an Acid Erlenmeyer flask (or beaker for pH titration)

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STUDENT NOTES Chapter 6 • Acid–Base Equilibria 125 Determine the value of pK b from your titration graph, and then determine an experimental value for K b of the weak base. NH 3 ( aq ) + H 2 O( l ) ֖ NH 4 + ( aq ) + OH ‒ ( aq ) Calculate the concentration of the original ammonia stock solution. Titrating an Unknown Weak Acid with Strong Base Your group will first prepare 100 mL of a solution containing an unknown weak acid. Using an analytical balance, weigh ~0.7 g of the acid into a 100 mL beaker. Record the acid mass to the nearest 0.0001 g. (Refer to Appendix B for analytical balance operation instructions.) Add approximately 25 mL of deionized water to the beaker and swirl to dissolve the solid. Carefully, transfer the solution to a 100 mL volumetric flask. Rinse the walls and bottom of beaker twice with deionized water and transfer to the volumetric flask. Cover the flask with Parafilm and invert several times to dissolve the solid. Add water exactly to the 100 mL mark on the flask; cover with Parafilm and mix. Each group has been allotted 180 mL of NaOH solution to complete the analysis. Goals: Your group will design and carry out procedures to Titrate 25.00 mL samples of the weak acid solution in two ways. Determine if the “unknown” acid is monoprotic (HA) or diprotic (H 2 A) in nature. monoprotic: HA + H 2 O ֖ H 3 O + + A ‒ K a diprotic: H 2 A + H 2 O ֖ H 3 O + + HA ‒ K a1 HA ‒ + H 2 O ֖ H 3 O + + A 2‒ K a2 Calculate and compare the concentration of acid in the solution using data from both titration methods. Calculate the molar mass of the “unknown” acid. Determine the K a value(s) for the acid. That is: K a if the acid is monoprotic; K a1 and K a2 if the acid is diprotic. [NH 4 + ][OH ‒ ] [NH 3 ] K b =

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STUDENT NOTES 126 Chemistry 11600 Laboratory Manual WASTE DISPOSAL AND CLEANUP All leftover solutions can be poured in the sink and rinsed down the drain with lots of water. Do NOT discard the pH 4 or pH 7 buffer solutions. Rinse the stir bar and return it to the designated place. BE CAREFUL: Do not pour the magnetic stir bar into the sink with the solution. Clean up any solids spilled on the balances or on the counter around them, dispose in sink and rinse down the drain with plenty of water. Return common equipment to your instructor or to the appropriate storage location. Disconnect the pH electrode and make sure the pH electrode is stored properly in deionized water or storage solution. Return the LabQuest 2 instrument box to your TA, after you have checked that all of the parts are present. Keep your splash goggles on until you have completed the data analysis, obtained any results that are needed from other groups, and are leaving the lab. Lock your lab drawer before leaving lab. DATA ANALYSIS Use the same data analysis techniques used in Part I and complete the data analysis for NH 3 and the unknown weak acid in order to accomplish the stated goals for each substance. RESULTS Compare your experimental K b value for ammonia with the accepted values listed in your textbook by calculating the % error in your experimental values. DISCUSSION Answer the following questions in your lab report. Why is it necessary to use a buffer solution to calibrate a pH electrode or meter? Why isn’t deionized H 2 O used? Be specific with your reasons. [H 3 O + ][HA ‒ ] [HA] K a1 = [H 3 O + ][A 2‒ ] [HA – ] K a2 =

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STUDENT NOTES Chapter 6 • Acid–Base Equilibria 127 LAB RECORDS AND REPORTS Group Portion You and your partner or group will turn in one completed report, either a formal lab report or a completed lab report form at the end of this chapter. It is your responsibility as a group to ensure that everyone whose name is on the report has participated as fully as possible in the completion of the project. The report or report form is an organized summary of your work and does not replace the need to keep a complete set of lab or field notes in your lab notebook as the lab is being done and data collected. Individual Portion Each student must attach laboratory notebook duplicate pages containing a complete data set and observations for the experiment.

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STUDENT NOTES 128 Chemistry 11600 Laboratory Manual

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Chapter 6 • Acid–Base Equilibria 129 REPORT FORM Title: Names: Date: Lab Section Number: GOAL(S): Procedure: Titration of NH 3 with HCl ( Overview of the procedure you used. ) DATA AND DATA ANALYSIS: NH 3 + HCl Data Ź Concentration of HCl: ( include units ) Ź Preparation of NH 3 solution for titration: Table 1. Titration Data for NH 3 + HCl ( use appropriate significant figures ) Titration Method Titration Type Volume of Acid, mL Volume of Base, mL Colorimetric Scout Colorimetric Careful Potentiometric Careful

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130 Chemistry 11600 Laboratory Manual DATA ANALYSIS Ź Chemical equation for the titration reaction: Ź pH titration curve ( attached; axes, title, and equiv. pt. labeled ) Ź Sample Calculations: [NH 3 ] calculation of [NH 3 ] dilute using colorimetric data ( show work and units using data from the careful trial) calculation of [NH 3 ] dilute using pH titration data ( show work—include units ) calculation of [NH 3 ] original ( show work for one of the samples ) Table 2. Data Analysis Summary, Original NH 3 Concentration Titration Method Titration Type [NH 3 ] Colorimetric Scout Colorimetric Careful Potentiometric Careful Average using data from “careful” trials = Ź Experimental value of K b for NH 3 ( show calculation )

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Chapter 6 • Acid–Base Equilibria 131 Procedure : Titration of unknown acid with NaOH (Overview of the procedure you used. Include which buffers were used for calibration and which indicator was used for the titration.) DATA AND DATA ANALYSIS: UNKNOWN ACID + NaOH Data Ź Concentration of NaOH: ( include units ) Ź Preparation of acid solution for titration: g acid dissolved in mL of Table 3. Titration Data for unknown acid + NaOH ( use appropriate significant figures ) Titration Method Titration Type Volume of Acid, mL Volume of Base, mL Colorimetric Scout Colorimetric Careful Potentiometric Careful DATA ANALYSIS Ź pH titration curve ( attached; axes, title, and equiv. pt(s) labeled ) Ź Nature of the acid (monoprotic or diprotic); explain.

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132 Chemistry 11600 Laboratory Manual Ź General equation for the titration reaction ( Use HA or H 2 A as the acid ): Ź Sample Calculations: [unknown acid] calculation of [unknown acid] using colorimetric data from the careful trial calculation of [unknown acid] using pH titration data ( show work—include units ) Table 4. Data Analysis Summary, Unknown Acid Concentration Titration Method Titration Type [unknown acid] Colorimetric Scout Colorimetric Careful Potentiometric Careful Average using data from “careful” trials = ______________ Ź Molar mass of unknown acid ( show calculation )

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Chapter 6 • Acid–Base Equilibria 133 Ź K a value(s) for unknown acid ( show calculations ) RESULTS Ź Comparison of experimental K b value for NH 3 with the accepted value listed in your textbook: % error = DISCUSSION ( Answer the questions that are in your lab manual .)

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134 Chemistry 11600 Laboratory Manual

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Lab 8 11600 lab manual_CH 6_Acid-Base Equilibria_2019-2020

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