Name__________________________________________ Period__ Date__________
Lab 13: Determination of the Equivalent Mass and pKa of an Unknown Acid Objectives: • To find the equivalent mass of an unknown weak acid • To plot a titration curve of standardized sodium hydroxide solution and the acid • To determine the equilibrium constant for the dissociation of the acid. Introduction: In this experiment you will determine the equivalent mass of an unknown acid, that is, the mass of the acid that supplies one mole of hydrogen ions. The acid, a solid crystalline substance, will be weighed out and titrated with a standardized solution of sodium hydroxide. From the moles of base used and the mass of the acid, you will be able to determine the equivalent mass of the acid. Next, you will plot the titration curve of the acid, with pH on the vertical axis and the volume of NaOH on the horizontal axis. From this graph you will be able to determine the value of the equilibrium constant for the dissociation of the acid. Acids are substances that contain ionizable hydrogen atoms in the molecule. Strong acids ionize totally, weak acids partially. The value of Ka, the equilibrium constant for the dissociation of the acid, is an indication of the strength of the acid. We can also use pKa, the –log(Ka), as an indication of acid strength. An acid may contain one or more ionizable hydrogen atoms in the molecule. The equivalent mass of an acid is the mass that provides one mole of hydrogen ions. It can be calculated from the molecular mass divided by the number of ionizable hydrogen atoms in a molecule. For example: hydrochloric acid, HCl, contains one ionizable hydrogen atom; the molecular mass is 36.45 g/ mol; the equivalent mass is also 36. 45 g/mol. Sulfuric acid, H2SO4, contains 2 ionizable hydrogen atoms; the molecular mass is 98.07 g/mol; the equivalent mass is 49.04 g/mol. Thus 36.45 g of HCl or 49.04 g of H2SO4 would provide you with one mole of H+ ions. The equivalent mass may be determined by titrating an acid with a standardized solution of NaOH. Since one mole of NaOH will react with one mole of hydrogen ion, at the equivalence point the following relationship is true: Vb x Mb = moles of OH- = moles H+ Equivalent Mass = ma/ moles of H+ where Vb is the volume of base (NaOH) in liters, Mb is the molarity of the base in moles per liter, and ma is the grams of acid used. A graph of pH versus mL of NaOH added can be created by following the titration with a pH meter. There should be a significant change in pH in the vicinity of the equivalence point. Note that since you are titrating a weak acid with a strong base, the equivalence point will NOT be at pH of 7, but should be slightly basic. The value of the equilibrium constant for the dissociation of the acid can be obtained from the graph. If we represent the dissociation of the acid as: HA H+ + Athe equilibrium expression is:
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When the acid is half neutralized, [HA] = [A-] so these terms cancel in the above equation:
This point can be found from the graph, by finding the half-equivalence point. The equivalence point must be identified. Then, you must find the pH for the volume that is half-way to that equivalence point. From the pH at that volume, which is you pKa, the Ka can be calculated.
Materials: unknown weak acid, solid standardized NaOH solution, ___________ M phenolphthalein Erlenmeyer flask, 125 or 250 mL medium beaker, 250-mL wash bottle with distilled water waxed weighing paper
buret analytical balance ring stand buret clamp small beaker, 100 mL large beaker, 1000 mL funnel
Safety Alert: Sodium Hydroxide is caustic and hazardous to skin and eyes. If you spill any on yourself, wash off skin with lots of water. Larger spills should be neutralized with vinegar by your teacher. The phenolphthalein solution is flammable. Keep it away from flames. Wear Safety Goggles and Chemical Resistant Apron
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Procedure: Part A: Determination of the Equivalent Mass of an Unknown Acid 1. Obtain a sample of unknown acid and use a sensitive balance to accurately mass 0.3 to 0.4 g onto a piece of waxed weighing paper. Record the exact mass on your data table. 2. Be sure to completely wash the acid sample using distilled water into an Erlenmeyer flask. Add approximately 40 mL of distilled water and swirl until it is dissolved. 3. Rinse your clean buret with tap water and then three times with small (about 7 mL) portions of your NaOH solution. While you are rinsing it, check for leaks around the stopcock. Then, secure your buret to the ring stand using a buret clamp. To be sure the tip of the buret has been filled, add NaOH solution to the buret filling it past the 0.00 mark (you may want to use a funnel). Drain the excess solution into a waste beaker until the volume is exactly 0.00 mL. 4. Add 3 drops of phenolphthalein solution to the dissolved acid in the flask. Swirl to mix. 5. Carefully, titrate the acid mixture with NaOH until the mixture stays faintly pink for 15 seconds. Remember to constantly swirl the flask and to rinse off the sides of the flask (that may be splattered with unreacted base solution) with distilled water before you reach the end point. Record the exact volume of NaOH used, estimating to the nearest 0.01 mL. 6. Refill your buret with NaOH solution. Repeat this titration process (steps #1-2 and #4-5) one more time, using enough acid to require about 45 mL of NaOH. This second titration may be done quicker since you know an approximate equivalence point. Be sure to slow down when you get close to the equivalence point so overshoot it. Part B: Determination of the pKa of the Unknown Acid 1. Obtain a sample of the same unknown acid and use a sensitive balance to accurately mass a sample of acid that will require approximately 30 mL of NaOH onto a piece of waxed weighing paper. Record the exact mass in your data table. 2. Be sure to completely wash the acid sample into a 250-mL beaker using your wash bottle. Add approximately 100 mL of distilled water and swirl until it is dissolved. Add 3 drops of phenolphthalein solution to the dissolved acid in the flask. Swirl to mix. 3. Unscrew the cap of the storage bottle on the pH sensor. After the cap has been unscrewed, slide the pH sensor out of the lid and replace the lid on the small bottle. Rinse the electrode well with distilled water over your waste beaker. Turn on your LQ and plug the sensor into one of the channel outlets at the top of it. Find Data Collection under the Sensor menu. Set the pH sensor to take 0.5 sample/second for 600 seconds. This should be more time than you need for the titration. 4. Place the pH sensor in the acid solution so the tip is submerged in your acid solution. Stir the solution with the sensor until the reading on the LQ is relatively constant (fluctuating around the same values).
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5. Situate your beaker where the buret can drip directly into the beaker (without dripping on the pH sensor) on one side and you can keep stirring the solution with the pH sensor on the other side. If the solution drips on your pH meter, have your partner rinse the side of the meter off using your wash bottle. 6. Once you have it set up correctly and are ready to begin, press Start (arrow button) and then start the buret dripping at a constant pace that is not too slow or too fast (about 1-2 drops per second). Be sure to continue to stir as the titration progresses. You can stop the titration (by pressing the arrow button again) when the pH has passed its equivalence point (the solution will turn pink and continue to get brighter) and the pH has leveled off and stayed almost constant for 30 seconds. 7. Save your data by choosing Save under the File menu. To clean up, drain your buret back into your NaOH bottle. Pour waste from the lab into your large beaker. Rinse all glassware with water. Clean your buret with a small bit of soapy water and a buret brush. Be sure to rinse all of the soap from it (use a beaker to pour water into it). 8. Get a laptop and work on the graph in Part B (instructions below). Be sure to get your graph off the LQ before you leave lab today. Data Molarity of Standardized NaOH solution: From Lab #6 in Lab Notebook- be sure that if you miscalculated it, you re-do the calculation so you have an accurate molarity for this lab. Ask for help if you need it. Unknown #: Part A
Trial 1
Trial 2
Mass of unknown (g): Volume of NaOH (mL): Part B Mass of unknown: Calculations/Discussion: Part A: 1. From the volume and molarity of the NaOH solution, calculate the number of moles of H+: Trial 1:
Trial 2:
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2. From the moles of H+ and mass of acid use, calculate the equivalent mass for your unknown acid. Trial 1:
Trial 2:
3. Calculate the average equivalent mass. Circle the identity of your unknown acid using your equivalent mass.
Potassium dihydrogen phosphate Potassium hydrogen phthalate Sodium hydrogen sulfite
KH2PO4 KHC8H4O4 NaHSO3
136.09 g/mol 204.22 g/mol 104.07 g/mol
Part B: 4. Retrieve the graph you made of your Part B titration: a. Plug your LQ into a laptop using the USB cord. Find the LoggerLite program under Applications or search for it using the spotlight. Choose “Yes” when it asks to retrieve the remote data. Choose OK when it asks if it should put remote data into the current file. It will show a data table to the left that has columns for time and pH and then a large graph on the right. Check that there is a label on each axis. The y-axis should be labeled pH and the x-axis should be labeled Time (s). Save your file under “Documents”. b. Separate the data table and graph on two “pages” by selecting the data table and choosing “Edit” and then “Cut”. Select “Page”, then “Add Page”. Choose “OK”. “Paste” the data table on this new page and resize it by choosing “Page” and then “Auto Arrange”. c. Go back to your first page by choosing “Page” then “Previous Page”. Click on your graph, then, resize your graph on the first page by choosing “Page” and then “Auto Arrange”. Create a title for your graph by choosing Graph Options from the Options menu. Make your title descriptive like “Titration Curve for Unknown Weak Acid”. d. Find and select the most vertical portion of your graph (around the equivalence point). Be sure all of the area you are trying to select is highlighted in dark grey. Add a linear best fit line to that area by choosing “Linear Fit” under the Analyze menu. Be sure that the correlation of your best fit line is at least 0.98. If it is not, delete the line and trying selecting a smaller and more consistent area. e. Once you have the area selected and a best fit line, choose Statistics under the Analyze menu. This should show you the minimum and maximum values for both x and y axis in the selected area. Move the two boxes around to where the curve is completely visible. Average the minimum and maximum
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values for the x-axis to find the time at the equivalence point. Write that here: f. Write the equation for your best-fit line below: g. In your equation, plug in the time at the equivalence point (that you determined in (e)) and calculate the pH at that time. Show your work.
h. Record the time and pH at this point in the data table below. Divide the time in half to find the time for the half-equivalence point. Using the Examine tool, find the half-equivalence point on the graph. Record the time and pH at this point in the data table as well. If your half-equivalence point it not exactly at one of your data points, choose the closest data point. Time (s)
pH
Equivalence Point Half-Way Point i.
Save your file again. Now, to save your file as a PDF, choose “Print Graph” (NOT “Print”) from the File menu. Add your names as a footer and choose “OK”. In bottom left-hand corner of the box, click of the PDF drop-down menu. Choose “Save as PDF”. In the “Save” box, name your graph something like Lab 13-Graph and save it under your Documents. Go to “Next Page” under Page menu. Choose “Print Data Table” (NOT “Print”) from the File menu. Add your names as a footer and choose “OK”. In bottom left-hand corner of the box, click of the PDF drop-down menu. Choose “Save as PDF”. In the “Save” box, name your graph something like Lab 13-Data Table and save it under your Documents. Email me both files:
[email protected] so I can print them for you.
5. Using the values above, find the Ka of the weak acid. The method is described in the introduction.
6. Normally, the equivalent mass of an acid is equal to the molar mass of the acid. For which type(s) of acids would this not be true? Why?
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7. Why is phenolphthalein a good indicator for this titration?
8. Your graph is in terms of time, and not volume. What could you do with your current data to change it to be terms of volume instead of time? Describe any measurements you would need to make and how you would use them.
9. How could you change your data to be in terms Describe in detail how each of the following mistakes would affect the equivalent mass calculated at the completion of the lab. a. A student added additional water to dissolve the acid prior to the titration
b. A student shook the weighing paper into the flask but failed to rinse all acid from the weighing paper into the titration flask.
c. A student rinsed the buret with water but forgot to rinse the buret with NaOH solution before filling it.
d. A student added base beyond the equivalence point to where the solution was dark pink instead of pale pink.
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