TITRATIONwebsites.rcc.edu/lugo/files/2019/10/1B-Titration-Curves-Fall-2017.pdf · 2019-10-01  · INTRODUCTION: Titration is a common laboratory technique (one type of volumetric

TITRATION

Revised 06/22/15

OBJECTIVES:

• Standardize a solution of sodium hydroxide

• Generate and compare the pH curves of a strong acid/strong base titration and a weak acid/strong base titration

• Determine the identity of an unknown monoprotic weak acid based on its molar mass and pKa value

INTRODUCTION:

Titration is a common laboratory technique (one type of volumetric analysis) used in the chemistry laboratory to determine the concentration of an unknown (the analyte) by quantitatively adding a known volume of a reagent known as the titrant. Two common titrations often performed in chemistry labs are acid-base and oxidation-reduction titrations. In this experiment, several acid-base titrations will be performed using the set up shown below.

A buret is used to quantitatively deliver a variable volume of the titrant (NaOH) into a beaker containing the analyte (acid) until they have fully reacted according to stoichiometric ratio. This point in the titration (when all the acid has been neutralized by the added base) is known as the equivalence point. Since information about the acid is dependent on the volume and concentration of the NaOH used, it is critical to have the exact molarity of the titrant.

In Part 1 of this experiment, a solution of sodium hydroxide will be standardized using

potassium hydrogen phthalate (KHC8H4O4). KHC8H4O4 is an acidic salt that dissociates completely into potassium ion (K+) and hydrogen phthalate ion (HC8H4O4

–) in water. The hydrogen phthalate ion acts as a weak monoprotic acid and reacts with sodium hydroxide according to the following net ionic equation:

HC8H4O4

– + OH– → H2O + C8H4O42– (1)

On Off

Stirrer

Analyte

Titrant

Figure 1 Titration Set Up

TITRATION

Standardization is a process that is performed because the molarity of the NaOH is often given as an approximate value. The approximate molarity is due to several factors, one of which is that prepared sodium hydroxide solution contains impurities. Second, over time, the molarity of NaOH decreases due to reaction with the carbon dioxide in the air. Therefore, it is important to determine the exact molarity of NaOH through a standardization procedure immediately before the solution is used in titrations. The calculation involved in this process is shown below:

Since all the reactants and products in the reaction are clear and colorless solutions, it is

impossible to know when the reaction is complete (i.e. have reached the equivalence point). One method to aid visualization of the equivalence point is through the use of an acid-base indicator such as phenolphthalein. Acid-base indicators are dyes derived from plants that display different colors at different pH values. For instance, phenolphthalein is colorless in solutions with pH < 8.3 but turns a reddish-pink color in solutions with pH > 9.5. The pH range in which an indicator changes color is known as its end point. Since different indicators have different end points (as shown in Table 1), the selection of a specific indicator for use in titrations to visualize the equivalence point must be done carefully. In other words, if the pH at the equivalence point is 8.5, phenolphthalein would be a good choice but bromophenol blue would be a poor choice.

Table 1: Some Common Acid-Base Indicators Indicators Colors pH range

(end point) Crystal violet Yellow-Blue 0.3-1.7 Thymol Blue Red-Yellow 1.5-2.5 Bromophenol blue Yellow-Blue 3.2-4.0 Methyl red Red-Yellow 4.6-6.0 Bromothymol Blue Yellow-Blue 6.2-7.0 Phenolphthalein Colorless-Pink 8.3-9.5 Thymolphthalein Colorless-Blue 9.5-10.2

A second way to visualize the equivalence point is the construction of a pH curve. In

Part 2 of the experiment, a pH meter will be used to monitor the progress of two titrations: HCl (strong acid)/NaOH and HC2H3O2 (weak acid)/NaOH. Initially the pH of an acidic solution is low since there are plenty of hydronium ions (H3O+) in solution. However, as a base is added, neutralization of the acid occurs, lowering the amount of H3O+ in solution and thereby raising the pH of the solution. The equivalence point is the inflection point (the point where the slope changes) in the graph (see Figure 2). Notice that the pH rises sharply with a small addition of a base at the equivalence point as shown in the figure below. As expected, reaction between a strong base and a strong acid result in a neutral salt and thus the pH at the equivalence point is 7.00.

Grams of KHP Moles of KHP Moles of NaOH Molarity of NaOH

Volume of NaOH at equivalence point

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Unlike a strong acid-strong base titration, the reaction between a weak acid and strong

base produces a weakly basic salt and hence the pH at the equivalence point is above 7.00 as shown in Figure 3. Another observation that is important to note on the titration curve is the buffering region. Recall that a buffer is composed of a weak acid and its conjugate base. In the course of the titration, some of the initial weak acid is converted into the conjugate base through the reaction between the weak acid and sodium hydroxide.

Figure 2 Strong Acid-Strong Base Titration

Volume NaOH Added (mL)

pH

0.0 1.00 5.0 1.18 10.0 1.37 15.0 1.60 20.0 1.95 22.0 2.20 24.0 2.69 26.0 11.29 28.0 11.75 30.0 11.96 35.0 12.22 40.0 12.36 45.0 12.46 50.0 12.52

!

!

Volume NaOH Added (mL)

pH

0.0 2.87 5.0 4.14 10.0 4.57 15.0 4.92 20.0 5.35 22.0 5.61 24.0 6.12 26.0 10.29 28.0 11.75 30.0 11.96 35.0 12.22 40.0 12.36 45.0 12.46 50.0 12.52

!

!

Figure 3 Weak Acid-Strong Base Titration

equivalence point

buffering region

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At the half-way to the equivalence point (when the volume of base added is ½ the amount needed to reach the equivalence point), half of the initial moles of acid has reacted with the added NaOH and has been converted into its conjugate base. At this point, the concentration of the weak acid and its conjugate base are the same (the ratio is 1) and according to the Henderson-Hasselbalch equation (equation 2), the log term drops out and the pH of the solution is equal to the pKa of the acid.

Since pKa is a physical property of a weak acid, it can be used to determine the identity of a weak acid. In Part 3 of the experiment, a measured mass of an unknown, monoprotic weak acid will be titrated against a standardized solution of sodium hydroxide. By constructing a pH curve from the titration data, the pKa (see discussion above) and the molar mass of the acid (equation 3) can be obtained. The moles of the acid can be calculated from the volume and molarity of the sodium hydroxide at the equivalence point since the mole of the acid is equal to the mole of the added base (1:1 stoichiometric ratio).

In summary, we will be performing a total of four different titrations in this experiment. In Part 1 of the experiment, a solution of sodium hydroxide will be standardized using KHP and the titration will be monitored using an acid-base indicator. In Part 2, equimolar and equal volumes of hydrochloric acid and acetic acid solutions will be titrated with the standardized sodium hydroxide solution, and the titration will be monitored with a pH meter. In Part 3, an unknown weak acid will be titrated with standardized sodium hydroxide in order to determine its identity by comparing the experimental molar mass and pKa to a list of known acids.

PRE-LAB QUESTIONS:

1. What is the difference between end point and equivalence point? How is the equivalence point determined in Part 1? Part 2? Part 3?

2. Compare the two titration curves (Figure 2 and 3 in the Introduction) to answer the

following questions. a. Write the balanced chemical equations for the reaction between HCl/NaOH and

HC2H3O2/NaOH.

b. Which figure shows a lower initial pH on the pH curve? Why? c. Why is the pH at the equivalence point different on the two graphs? Explain.

d. What volume of NaOH is needed to reach the equivalence in both graphs? Explain why they are the same. Assume that both the volume and molarity of the strong acid and the weak acid, and the molarity of NaOH used in the titration are the same.

€

pH = pKa + log [conjugate base][weak acid]

(2)

€

Molar mass of the acid = mass of the acidmoles of the acid

(3)

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3. For the reaction between KHC8H4O4 and NaOH in Part 1, a. Approximately how many grams of KHC8H4O4 should be weighed out if

approximately 10 mL of 0.20 M NaOH solution is used to reach the equivalence point?

b. Calculate the pH at the equivalence point for the reaction between KHC8H4O4 and NaOH from part (a). Assume that the total volume of the solution at equivalence point is 40 mL. Look up the pKa of KHC8H4O4 from a reference table.

c. Based on your answer from above, which acid-base indicator (see Table 1 in the Introduction) should be used for the titration? Explain your reasoning.

4. How will you dispose of all excess and neutralized solution?

CHEMICALS AND MATERIALS Solid KHC8H4O4 Indicators (in solution): 0.20 M NaOH crystal violet bromothymol blue 0.100 M HCl thymol blue phenolphthalein 0.100 M HC2H3O2 bromphenol blue thymolphthalein Unknown acid methyl red pH 4, 7, and 10 buffers

EQUIPMENT TO BE CHECKED OUT 25 mL buret + buret reading card Magnetic stir bar + Magnetic stirrers Pipet pump

SAFETY Goggles must be worn in the laboratory.

WASTE DISPOSAL Dispose all excess and used solutions in the chemical waste container.

EXPERIMENTAL PROCEDURE:

• Obtain approximately 150 mL of 0.10M NaOH solution in a clean and dry 250 mL Nalgene plastic bottle. The bottle should be kept closed when not in use to minimize exposure to CO2. This NaOH solution will be kept in your drawer and used for all 3 parts, so be sure to use it sparingly.

• Check out the following items from your professor: a buret, a buret reading card, and a magnetic stir bar. (buret clamps, ring stands, and magnetic stir plate can be found in the lab room).

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Cleaning, Condition, and Setting up the buret

• Clean and set up the buret as follows: ! Check the buret to be sure that it is not leaking and the stopcock turns easily.

Rinse the buret with tap water and then do a final rise with DI H2O. If water does not drain properly or beading occurs inside the buret, clean the buret with soapy water and a buret brush.

! Condition the buret with approximately 5 mL of 0.20M NaOH solution. Be sure to drain the NaOH through the buret tip. The used NaOH solution should be placed in a beaker labeled “Waste NaOH solution.” Repeat the conditioning two more times. Place the “Waste NaOH solution” beaker underneath the buret to catch any drips when the buret is not in use.

! Clamp the conditioned buret to the ring stand and fill the buret with a fresh 0.10M NaOH solution (use a clean, dry funnel as needed) so that the liquid level is above the 0.00mL mark. Drain some NaOH solution through the tip of the buret so it is filled with the fresh NaOH solution and the level of the NaOH in the buret is slightly below the 0.00 mL mark. Check to be sure that there are no air bubbles in the tip.

! Have your instructor check your set up. Record the initial volume of the NaOH solution (use the buret reading card to aid in the reading) in your notebook.

Part 1: Standardization of NaOH Solution (Optional)

1. Check with your professor to see whether you will be performing this part of the experiment. If you are doing this part, check and see whether you are performing multiple trials or doing a single trial and pooling your data with those of your classmates working at your lab bench.

2. Weigh out the calculated mass of KHC8H4O4 (see answer to Pre-Lab Question #3) on a weighing boat or paper and quantitatively transfer all the solids into a clean 125 mL Erlenmeyer flask. Record the exact mass of the KHC8H4O4 used.

3. Measure 30.0 mL of DI H2O in a graduated cylinder and add to the KHC8H4O4 in the flask. Add three drops of the chosen indicator (see answer to Pre-Lab Question #3c). Add the magnetic stir bar to the flask and place the flask on the magnetic stir plate to dissolve the KHC8H4O4.

4. Titrate the dissolved KHC8H4O4 until the end point is reached using the procedure below.

a. Place the Erlenmeyer flask underneath the buret so that the tip of the buret is slightly inside the flask.

b. Open the stopcock of the buret and drain the NaOH solution into the flask. Based on your calculations, it should take approximately 10 mL of the NaOH to reach the equivalence point, so add approximately 8 mL of NaOH solution quickly.

TITRATION

c. Add NaOH solution drop wise as you near the end point—you can see this visually as the color of the solution stays longer and longer in the flask. As you near the end point, consider adding “partial drops”—open the stopcock and let a “partial drop” hang on the tip of the buret. Squirt DI H2O from a wash bottle so that the drop falls into the flask). Be sure to stop the addition of NaOH and swirl completely before adding the next “partial drop.” The end point is reached when a faint color persists for more than 30 seconds. Record the color of the solution in the flask and the final volume of the NaOH.

5. Based on the exact mass of KHC8H4O4 and the volume of NaOH used in the titration, calculate the exact concentration of NaOH (be careful with significant figures). Repeat the titration for additional trials or report the concentration to your professor for pooled data. (Note: An average value of the concentration of NaOH will be computed using student data for each lab bench and written on the board. Copy this average concentration down in your notebook. This concentration will be used for Data Analysis for Part 2 and 3, as well as for Post-Lab Question #1).

6. Remove the magnetic stir bar from the flask. Dispose the neutralized KHC8H4O4 solution into the waste container provided in the fume hood. Rinse the flask and magnetic stir bar with tap water and do a final rinse with DI H2O.

7. Refill the buret with more NaOH as needed so that the initial volume in the buret is a little bit below the 0.00 mL mark.

Part 2: Titration of a Strong Acid and Weak Acid with NaOH

8. Obtain approximately 35 mL of 0.100 M HCl in a clean Erlenmeyer flask. Condition a 10mL volumetric pipet with the HCl solution. Pipet 20.00 mL of 0.100 M HCl solution into a clean, dry 150 mL beaker.

9. Obtain pH 4 and 7 buffers using the small vials near the buffer containers. Use the pH 4 and 7 buffers to calibrate the pH meter (see the calibration handout for details).

10. Set up the titration apparatus as shown in the figure.

• Add a magnetic stir bar into the beaker and place it on top of the stir plate. Begin gentle stirring.

• Clamp the pH electrode to the ring stand with the bulb of the electrode completely submerged in the solution in the beaker. Be sure that the magnetic stir bar does not hit the electrode.

• Lower the buret so that the tip of the buret is close to the solution but not submerged in the solution.

TITRATION

11. Once the pH of the HCl solution is stabilized, record both the pH and the initial volume of the NaOH solution in your lab notebook as shown below. Add more spaces until titration is complete.

pH Buret reading (mL)

2.34 0.10 Continue recording pH Continue recording volume of

NaOH

12. Add NaOH solution slowly into the beaker until the pH unit rises by approximately 0.20 units. Stop the addition of NaOH and record the buret reading and the pH in your lab notebook as shown above. The pH measurement by the electrode lags behind the addition of NaOH so be sure to wait for the pH reading to stabilize.

13. Continue the above process (add NaOH, stop when pH increases by ~0.20 unit, record the buret reading and the pH) until pH reaches an upper plateau region and changes very little (~ pH 12.00). The volume of NaOH required to increase the pH will decrease as you approach the equivalence point. Note: It is strongly recommended that the NaOH solution be added drop wise or even by partial drops near the equivalence point.

14. Remove the stir bar from the flask. Dispose the neutralized solution into the waste container. Wash, rinse, and dry the beaker.

15. Refill the buret with fresh NaOH solution and repeat the above titration procedure using 20.00 mL of 0.100 M HC2H3O2 in a clean, dry beaker.

16. Dispose of any remaining unused HCl and HC2H3O2 into the waste container.

Part 3: Titration of an Unknown Acid

17. Obtain an unknown acid from your professor. Record the weak acid unknown number in your lab notebook, along with initial observations of the acid. Measure approximately 0.2-0.3 grams (record the exact mass) of the unknown acid in a clean and dry150 mL beaker.

18. Using a graduated cylinder, measure 30.0 mL of DI H2O and add to the acid. Place the magnetic stir bar into the beaker and stir gently on the stir plate to dissolve all the acid. If the acid does not dissolve within 5-10 minutes, you may need to heat the solution to help the dissolution process.

19. Titrate the unknown acid using the procedure from Part 2 until pH reaches an upper plateau region and changes very little (~ pH 12.00). Remove the magnetic stir bar from the flask. Dispose the neutralized unknown acid solution into the waste container provided in the fume hood. Rinse the flask and magnetic stir bar with tap water and do a final rinse with DI H2O before performing a second trial of the titration of the unknown if time allows or as directed by your professor.

TITRATION

20. Dispose of the solid unknown by emptying out the unknown solid into the waste container in the fume hood. The empty plastic baggie can be thrown into the trashcan.

21. Empty the contents of the buret into the “Waste NaOH solution” beaker and dispose the solution into the waste container. Wash all the glassware with tap water and then do a final rinse with distilled water. Return the checked out items back to your professor.

DATA ANALYSIS

Part 1. Standardization of NaOH (Optional)

1. If you are using pooled class data, how does your calculated molarity of NaOH compare with the average? Compute the percent error using the average as the “Theoretical” value. Is the percent error within the expected range based on the limitation of the glassware used? Explain.

2. If you performed multiple trials for the titration, compute the average and the standard deviation. Is the standard deviation within the expected range based on the limitation of the glassware/instrumentation used? Explain.

Part 2. Titration of a Strong Acid and Weak Acid with a Strong Base

3. Use Excel to calculate the volume of NaOH added and organize your data in a table similar to the one below (note: only a portion of the table is shown; extend as needed). Plot an XY (Scatter) graph for the titration data (HCl with NaOH) with data points connected by a “smooth line”. Include minor gridlines on your graph so that you can precisely interpolate the equivalence and half equivalence points. Print the data table and graph on separate pages.

pH Buret reading (mL) Volume of NaOH added (mL)

2.34 0.10 0.00

2.55 0.35 0.25

Continue Continue Continue

4. Repeat step 3 for HC2H3O2 vs NaOH titration data. 5. Mark and label the following points on the graphs (use x and y gridlines on your graph).

a. Point A: Before any NaOH has been added b. Point B: When ½ way to the equivalence point has been reached

c. Point C: When equivalence point has been reached d. Point D: When 1.00 mL of additional NaOH has been added beyond the

equivalence point.

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6. Create a table similar to the one shown below using data from both titration graphs.

HCl vs NaOH HC2H3O2 vs NaOH

Volume of NaOH added (mL)

Experimental pH

Volume of NaOH added (mL)

Experimental pH

Point A

Point B

Point C

Point D

7. How do the pH values at point A compare in the two graphs? Explain. 8. Since the same volume and molarity of HCl and HC2H3O2 were used, the volume of

NaOH required to reach the equivalence point should be the same. Compare the experimental volumes of NaOH at Point C in both graphs. Explain any discrepancies.

9. How do the pH values at point C compare in the two graphs? Explain. 10. Which point on the graph corresponds to the pKa of the acid (only for the weak acid)?

Explain. 11. Compare the experimental pKa of HC2H3O2 with the theoretical value. Calculate the

percent error.

Part 3: Identification of an Unknown Weak Acid.

12. Use Excel to generate a table and pH curve for the unknown acid titration as in Part 2. If necessary, create an additional “zoomed-in” display of only the portion of the graph that revolves around equivalence point and half-way to the equivalence point in order to correctly identify the NaOH volume or pH value. Label the equivalence point and the half-way to the equivalence point on the curve.

13. Calculate the molar mass of the unknown acid using the titration data. Reminder: be sure to use the average concentration for NaOH from Part 1.

14. Determine the pKa of the unknown acid from the titration graph.

TITRATION

15. Determine the identity of the unknown acid based on its molar mass and pKa. (See table below). Calculate the percent error for the molar mass and the pKa.

Possible Acids Formula MM (g/mol) pKa benzoic acid C6H5COOH 122.13 4.19 potassium hydrogen tartrate KHC4H4O6 188.18 4.34 potassium hydrogen phthalate KHC8H4O4 204.23 5.41 citric acid H3C6H5O7 192.14 3.16 lactic acid HC3H5O3 90.09 3.85 sodium hydrogen phosphate Na2HPO4 141.96 12.38 sodium dihydrogen phosphate NaH2PO4 119.98 7.21 sodium hydrogen carbonate NaHCO3 84.01 10.25 ammonium chloride NH4Cl 53.50 9.25

POST-LAB QUESTIONS:

1. In Part 2 of the experiment, 20.00 mL of 0.100M HCl was titrated with a standardized solution of NaOH.

a. Using the exact molarity of NaOH found through the standardization (the value obtained from Data Analysis for Part 1), calculate the volume of NaOH needed to reach the equivalence point.

b. Calculate the theoretical pH values at points A, B, C, and D using the molarity and volume of NaOH from part “a”.

c. Organize your results in a table such as the one shown below. Compare the theoretical and experimental pH values and calculate % error.

Theoretical pH Experimental pH % Error

Point A

Point B

Point C

Point D

2. Repeat question #1 for the titration between 20.00 mL of 0.100M HC2H3O2 and NaOH.

3. How would the pH graph of a polyprotic acid such as H2SO4 differ from a monoprotic acid such as HCl? Include a rough sketch in your explanation.

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4. How would the pH graph differ if the NaOH was the analyte and HCl was the titrant? Include a rough sketch in your explanation.

5. How would the molar mass of the unknown be affected (increase, decrease, no change) if the following occurred during the lab? Explain.

a. Some unknown solids were spilled while transferring into the beaker. b. The buret was rinsed with DI H2O but not conditioned with the 0.20 M NaOH

solution.

REFERENCES:

Written by Jenny Leung, April 2011 (Reviewed and edited by Janet Truttmann and Terri Beam June 2011)