What chemical reaction is taking place at the equivalence point of the titration in Part D that gives rise to the pH at the equivalence point? Use the generic formula HA to represent the unknown weak acid and write a chemical equation to help explain your answer.

Potentiometric Determination of Ka for a Weak Acid

Supplies:

pH meter, pH electrode, 3-250 mL beakers, 2 watch glass, 50 mL beaker, wax pencil, magnetic stir bar, stir plate, 50 mL buret, ring stand, buret clamp, 10 mL volumetric pipet, pipet pump (or bulb), transparent one foot rulers, squirt bottle of DI water, KimWipes, standard 0.10 M NaOH solution, unknown solid weak acid, 95% ethanol, buffers of pH 4.00 and 7.00.

1. Obtain about 100 mL of the standard 0.10 M NaOH solution in a clean, dry beaker and cover with a watch glass. Label the beaker. Record the exact molarity as shown on the reagent bottle.

2. Obtain a 50 mL buret and prepare it for titration by rinsing, filling, etc. with the standard 0.10 M NaOH solution as described in Appendix 5. Remember to remove any air bubbles from the buret tip.

3. Standardize the pH meter as directed in Appendix 7.

4. Place about 1.0 g of the unknown solid into a clean dry beaker.

5. Transfer 0.4-0.5 g of the unknown solid acid to a clean beaker (no larger than 250 mL).
Record the exact mass of the unknown solid acid.

Acids, Bases, Salts and Buffers Purpose

To investigate some of the properties of acids, bases, salts, and buffer solutions using pH measurement, graphical analysis, and qualitative observations. Introduction The concepts of acid and base are two of the most central in all of Chemistry. In the 1880’s Arrhenius stated that, in aqueous solution, acids produced hydrogen ions (protons) and bases produced hydroxide ions.

The reaction of a base with an acid (neutralization reaction) was thus represented by the chemical equation: H+ (aq) + OH- (aq) → H2O (l) (1) We now know that a proton in aqueous solution actually exists as a hydronium ion, H3O+. In the 1920’s, the concepts of acid and base were made more general by Brønsted in Denmark and Lowry in England. They viewed an acid-base reaction as a proton-transfer reaction, in which the acid donates or transfers a proton to another substance, and the other substance (the base) functions as the proton acceptor.

Included in their description was the idea that an acid-base reaction is a reversible process, so that any acid-base reaction could be represented by the general equation: HA + B BH+ + A- (2) In the forward reaction, the acid HA transfers a proton to the base B giving the indicated products. In the reverse reaction, the acid BH+ transfers a proton to the base A-. Thus, in any Brønsted-Lowry acid-base reaction, there are always two pairs of acids and bases, with the acid and base in any pair on opposite sides of the equation. The acid or base on the opposite side of the equation is referred to as the conjugate acid or base of the species in question.

Thus, in equation 2, BH+ is the conjugate acid of the base B, and A- is the conjugate base of the acid HA. Correspondingly, HA is the conjugate acid of A- and B is the conjugate base of BH+. Weak Acids and Bases: As for any other equilibrium, the magnitude of the equilibrium constant for the above reaction is a measure of how far to the right the reaction proceeds. That is, whether the products or reactants are favored. When a weak acid is placed in water, the equilibrium HA (aq) + H2O (l) H3O+ (aq) + A- (aq) (3) is established, and when a weak base is placed in water, the equilibrium B (aq) + H2O (l) BH+ (aq) + OH- (aq) (4) is established.

The equilibrium constants for the above reactions are termed the acid dissociation constant, Ka, and base dissociation constant. Kb, respectively. Their magnitudes can be used to set up a scale of relative strengths of acids and bases, and can also be used like any other equilibrium constant to calculate equilibrium concentrations of the species in the reaction. It is also true that the product of Ka and Kb for a conjugate acid-base pair is equal to the ion product constant for water, Kw. This means that as the acid becomes stronger (Ka increases), its conjugate base becomes weaker (Kb decreases), so that the strongest acids have the weakest conjugate bases.

As one part of this experiment, you will determine the magnitude of Ka for an unknown weak acid. Autoionization of Water: In equation 3, water is acting as a base, while in equation 4, water is acting as an acid. In pure water or any aqueous solution, the equilibrium 2 H2O (l) H3O+ (aq) + OH- (aq) (5) is established. Note that this reaction is essentially the reverse of equation 1. The equilibrium constant for the reaction in equation 5 is called the autoionization constant (AKA autoprotolysis constant or ion-product constant) for water, symbolized Kw.

The numerical value of Kw is Kw = [H3O+][OH-] = 1.0 x 10-14 (6) Buffer Solutions: In equation 3, the anion formed from the weak acid HA acts as a base, and in equation 4, the cation formed from the weak base B acts as an acid. Thus, we would expect that salts of weak acids (which contain the anion which is the conjugate base of the weak acid) and of weak bases (which contain the cation which is the conjugate acid of the weak base) would be capable of influencing the amount of H3O+ and/or OH- present in solution.

Also, a solution containing a weak acid and its salt (or a weak base and its salt), since it contains both a weak acid and a weak bases, should have interesting properties. In fact, such a solution (referred to as a buffer solution), maintains the concentration of H3O+ and OH- at a fairly constant level when small amounts of strong acid or base are added to the solution. This property gives these solutions a number of uses in the laboratory.

You should refer to your textbook for additional discussion of the topics of pH of salt solutions and buffer solutions. pH Scale: As an alternate way to express the concentrations of H3O+ and OH- in solution, the quantities pH and pOH are defined: pH = -log[H3O+] pOH = -log[OH-] (7) where the square brackets signify the molar concentration of that species. Based on this, and the fact that the product of [H3O+] and [OH-] equals Kw, the pH of a neutral solution is exactly 7.00. The further below this value, the more acidic the solution (higher [H3O+]), the farther above 7.00, the more basic (higher [OH-]). You should consult your textbook for further discussion of the calculation of pH from [H3O+] and vice versa, the relationship of pH and pOH, and for the calculation of [H3O+] or pH for a solution (including a buffer) using the Ka or Kb values for its solutes.

In CHEM 131L, you performed a titration in which the end point of the titration was signaled by a color change due to an added indicator. In this case, a slight excess of titrant was necessary to produce the change, so the point at which you stopped the titration (end point) was at a slightly higher volume that the true equivalence point of the titration. In Parts B and D of this experiment, you will be using an instrument to measure the electrical potential (voltage) of a circuit, so it is referred to as a potentiometric titration. In this experiment, you will not be measuring the potential directly, but it will be converted by the instrument to the pH of the solution, and this will be used to determine the equivalence point. You will be working on this experiment for two laboratory periods.

You will be performing a series of procedures which explore the concepts of acids, bases, salts, and buffers using the measure pH of a solution as an experimental tool. In Part A, you will investigate one difference in behavior between a strong acid and a weak acid. In Part B, you will perform a potentiometric titration of a strong acid with a strong base and use the pH of the solution to determine the equivalence point of the titration, allowing you to determine the concentration of the acid. In Part C, the acid-base behavior of a series of salts will be investigated.

In Part D, you will perform a potentiometric titration of a weak acid with a strong base to determine the value of Ka for an unknown weak acid. Part E is an investigation of the behavior of buffers.

Laboratory Safety

• Wear safety goggles when working in the laboratory.

• Solutions of hydrochloric acid can cause burns and give off irritating vapors.

• Solutions of sodium hydroxide can cause burns, especially to the eyes.

• All solutions used must be disposed of in the Acid-Base Waste container provided. Procedure Work with a partner on all parts of this experiment. General Procedure for Using the pH Meter

1. Standardize the pH meter as described in Appendix 7 using buffers of pH 4.00 and 7.00 as the “low” and “high” pH respectively. The pH meter should be standardized before use each lab period.

2. When switching the electrode from one solution to another, thoroughly rinse the pH electrode with DI water from a squirt bottle into a waste beaker.

3. Wipe the sides of the electrode and gently blot dry the tip with a clean, soft tissue (KimWipe) before measuring the pH of the next solution.

4. When not in use, the pH electrode should be in the special storage container provided. At no time should the pH electrode be held horizontally or turned upside-down.

Part A: Acid Dissociation Supplies:

pH meter, pH electrode, 10 glass vials, wax pencil, squirt bottle of DI water, KimWipes, pH 4.00 buffer, pH 7.00 buffer, 5 HCl solutions (1.0 x 10-4 M, 5.0 x 10-4 M, 1.0 x 10-3 M, 5.00 x 10-3 M and 1.0 x 10-2 M), 5 CH3COOH solutions (1.0 x 10-4 M, 5.0 x 10-4 M, 1.0 x 10-3 M, 5.00 x 10-3 M and 1.0 x 10-2 M)

1. Obtain 5 clean, dry glass vials and label them.

2. Fill each vial about half full with one of the standard hydrochloric acid (HCl) solutions provided with a different solution in each vial. Record the exact concentration of each solution in your lab notebook.

3. Standardize the pH meter as directed in Appendix 7.

4. Rinse the electrode thoroughly with DI water and insert it into the vial containing the HCl solution of lowest concentration. Swirl the solution gently.

5. When the “stable” icon appears, record the pH in your notebook.

6. Repeat with the other vials of HCl solution. Remember to rise the electrode thoroughly between solutions.

7. Discard the HCl solutions in the proper waste container.

8. Using 5 clean, dry glass vials, repeat the steps above using the five acetic acid (CH3COOH) solutions.

9. Discard the acetic acid solutions in the proper waste container.

10. Rinse all 10 glass vials thoroughly with DI water and return them to their proper location.

Calculations  Complete during lab:

1. From the measured pH values, calculate the molar concentration of H3O+ in each of the ten acid solutions, and record those in your lab notebook. Show one example calculation.

2. Create Tables of acid concentration (from reagent bottle), measured pH, and calculated [H3O+]. Prepare one table for hydrochloric acid and a separate table for acetic acid. Include a descriptive caption for each table.

Complete outside of lab:

3. Prepare a graph of concentration of H3O+ (M) versus reagent concentration (concentration of HCl or CH3COOH on the reagent bottles). Think carefully regarding which data set is plotted on which axis. Enter the data in order of increasing concentration of acid.

Plot both acids on the same graph. See the Freezing Point of a Solution experiment for instructions on how to plot two data sets in the same graph. Insert a trendline for each acid. Include a descriptive caption including the trendline and R2 value for each acid.

Part B: Potentiometric Titration of a Strong Acid with a Strong Base

Supplies:

pH meter, pH electrode, 3-250 mL beakers, 2 watch glass, wax pencil, magnetic stir bar, stir plate, 50 mL buret, ring stand, buret clamp, 10 mL volumetric pipet, pipet pump (or bulb), transparent one foot rulers, squirt bottle of DI water, KimWipes, standard 0.10 M NaOH solution, HCl solutions of unknown concentration, buffers of pH 4.00 and 7.00.

1. Obtain about 70 mL of the standard 0.10 M NaOH solution in a clean, dry beaker and cover with a watch glass. Label the beaker. Record the exact concentration of the solution from the reagent bottle.

2. Obtain a 50 mL buret and prepare it for titration by rinsing, and filling it with the standard 0.10 M NaOH solution as described in Appendix 5. Be sure to remove any air bubbles from the tip.

3. Obtain about 25 mL of the unknown HCl solution in a second clean, dry beaker and cover it with a watch glass. Label the beaker.

4. Obtain a 10 mL volumetric pipet and rinse it 2-3 times with a small portion of the HCl solution (less than 5 mL).

5. Carefully pipet 10.00 mL of the unknown HCl solution into a clean 250 mL beaker.

6. Add 90 mL of DI water and the magnetic stir bar to the beak with HCl. 7. Place the beaker on the stir plate.

8. Position the stir plate, ring stand, buret and pH meter so that the pH electrode can be placed into the solution in the beaker and the stir bar can spin without hitting the electrode tip.

9. Adjust the stir plate setting so that the magnetic stir bar is spinning at a moderate rate.

10. Position the buret tip in the beaker, but not touching the beaker walls.

11. Read and record the initial buret reading and pH.

12. Add 2 mL increments of NaOH from the buret into the beaker, allow the pH reading to stabilize, then record the pH and the buret reading.

13. Repeat this procedure with additional 2 mL increments of NaOH until the pH approaches 2.8. After the addition of each increment of NaOH, allow the pH reading to stabilize, then record the pH and buret reading.

14. In the pH range of 2.8 to 10, add 1-2 drops (~0.10 mL) of NaOH, allow the pH reading to stabilize, then record the pH and buret reading.

15. After pH 10, return to adding 2 mL increments of NaOH until you have recorded five readings beyond a pH of 11.0.

Calculations – Complete outside of lab

1. Prepare a plot of pH versus cumulative volume of NaOH added.

Note that your initial buret reading must be subtracted from every reading before preparing the graph. Include a descriptive caption.

2. Determine which points are essential for determining the equivalence point (see Figure 1). Generally, you need 4-5 points before and 4-5 points after the portion of the graph that is rising sharply. Expand this part of the graph by adjusting the “high” and “low” values of the x-axis to obtain a graph like the one shown in Figure 2.

3. Be sure to use a sufficiently small increment on the y-axis so that you can read the pH scale precisely.

4. Print a copy of the expanded graph at full-page size in Landscape mode.

5. Using a ruler and sharp pencil, draw lines A, B and C onto the expanded graph as shown in Figure 3.

6. Measure the length of line C between its intersections with lines A and B. The equivalence point is the midpoint of this line. Read the volume of NaOH added at the equivalence point to ± 0.1 mL (or better) and record along with the pH at the equivalence point. Include a photo or scan of your plot with lines A, B and C and report the pH and NaOH volume at the equivalence point.

7. Using the volume of NaOH at the equivalence point and the molarity of NaOH, calculate the number of moles of HCl in the unknown. Show the calculation.

8. Use the moles of HCl and the volume of HCl used in the titration to calculate the molarity of the HCl solution. Show the calculation. Figure 1. Plot of pH versus volume of NaOH added during titration with HCl. Figure 2. Expanded area to determine the equivalence point.

Figure 3. Expanded area to determine equivalence point with lines A, B and C.

Part C:  pH of Salt Solutions Supplies:

pH meter, pH electrode, 3 glass vials, wax pencil, squirt bottle of DI water, KimWipes, 0.10 M KNO3, 0.10 M Na2HPO4, 0.10 M NH4Br, buffers of pH 4.00 and 7.00

1. Obtain 3 clean, dry glass vials and label them.

2. Fill each vial about half full with one of the three salt solutions. 3. Measure and record the pH of each of the solutions as outlined in Part A. Remember to rinse and dry the pH electrode between solutions.

Calculations – Complete outside of lab Complete during lab:

1. From the pH, calculate the molar concentration of H3O+ in each of the three solutions.

2. Create a Table that includes salt formula, measured pH, and calculated [H3O+]. Include a descriptive caption.

Complete outside of lab:

3. For each of the three solutions measured, indicate if the cation or anion was a pH active species. For any pH active ion, indicate if it would be classified as monoprotic or polyprotic. Write the balanced equation showing the interaction of the ion with water.

Part D: Potentiometric Determination of Ka for a Weak Acid Supplies:

pH meter, pH electrode, 3-250 mL beakers, 2 watch glass, 50 mL beaker, wax pencil, magnetic stir bar, stir plate, 50 mL buret, ring stand, buret clamp, 10 mL volumetric pipet, pipet pump (or bulb), transparent one foot rulers, squirt bottle of DI water, KimWipes, standard 0.10 M NaOH solution, unknown solid weak acid, 95% ethanol, buffers of pH 4.00 and 7.00.

1. Obtain about 100 mL of the standard 0.10 M NaOH solution in a clean, dry beaker and cover with a watch glass. Label the beaker. Record the exact molarity as shown on the reagent bottle.

2. Obtain a 50 mL buret and prepare it for titration by rinsing, filling, etc. with the standard 0.10 M NaOH solution as described in Appendix 5. Remember to remove any air bubbles from the buret tip.

3. Standardize the pH meter as directed in Appendix 7.

4. Place about 1.0 g of the unknown solid into a clean dry beaker.

5. Transfer 0.4-0.5 g of the unknown solid acid to a clean beaker (no larger than 250 mL). Record the exact mass of the unknown solid acid.

6. Add 20 mL of 95% ethanol to the solid acid. Add the magnetic stir bar to the beaker and place the beaker on the stir plate. Stir the mixture until the acid completely dissolves.

7. Add 50 mL of DI water to the acid solution and mix well.

8. Position the stir plate, ring stand, buret and pH meter so that the pH electrode can be placed into the acid solution and the stir bar can spin without hitting the pH electrode tip.

9. Adjust the stir plate so that the stir bar is spinning at a moderate rate.

10. Place the tip of the buret into the beaker, but not touching its walls.

11. Read and record the initial pH and buret reading.

12. Add NaOH in 2 mL increments, allow the pH to stabilize, then record the pH and buret reading.

13. Repeat this procedure with additional 2 mL increments of NaOH until the pH approaches 5.5.

14. For the pH range of 5.5-11.0, add 1-2 drops (~0.10 mL) increments of NaOH, allow the pH to stabilize, then record the pH and buret reading after each addition.

15. Beyond pH 11.0, add 1.0 mL increments of NaOH until pH 12. Calculations – Complete outside of lab The pH of a weak acid solution (HA in equation 3) can be calculated using the Henderson-Hasselbalch Equation: pH = pKa + log([A-]/[HA]) where pKa = -log(Ka). At the equivalence point, all of the weak acid is converted to its conjugate base, and halfway to the equivalence point (midpoint), exactly half of the weak acid would have been converted to its conjugate base. At this point, [A-] = [HA] and the ratio of [A-]/[HA] is equal to 1. Since log(1) = 0, the Henderson-Hasselbalch equation simplifies to pH = pKa at the midpoint. Since pKa = -log(Ka), the acid dissociation constant (Ka) of the unknown weak acid can be determined by taking the antilogarithm (inverse logarithm) of the negative of pKa. That is, Ka = 10-pKa.

1. Prepare a plot of pH versus cumulative volume of NaOH added. Use all of the data to get an overview of the results and note that the initial buret reading must be subtracted from every reading before preparing the graph. Include a descriptive caption.

2. Examine the graph and determine which points are essential to determine the equivalence point (see Figure 1). Generally, you need 4-5 points before and 4-5 points after the portion of the graph that is rising sharply. Expand this area to be able to read the NaOH volume at the equivalence point to ± 0.10 mL by adjusting the High and Low of the x-axis to obtain a graph similar to Figure 2.

3. Print the expanded graph at full-page size in Landscape mode.

4. Using a ruler and sharp pencil, draw lines A, B and C on the expanded graph (Figure 3).

5. Measure the length of line C between its intersections with lines A and B. The equivalence point is the midpoint of this line. Read the volume of NaOH added at the equivalence point to ± 0.1 mL (or better) and record it along with the pH at the equivalence point. Include a photo or scan of your plot with lines A, B and C and report the pH and NaOH volume at the equivalence point.

6. Divide the volume of NaOH at the equivalence point by two to obtain the volume of NaOH added. For example, if the volume NaOH added at the equivalence point is 27.2 mL, the volume of NaOH at the midpoint is 13.6 mL.

7. Using the same data as before, expand the midpoint area of the graph and expand the pH scale so that it can be read to ± 0.01 pH unit (Figures 4 and 5). You should have 5-7 points to establish a straight line. Print this graph.

8. Use a ruler to draw the best fit line through the points on your expanded graph.

9. Locate the midpoint on the graph. Record the corresponding pH. In Figure 5, the volume of NaOH added at the midpoint was 13.6 mL and the pH was determined to be 3.36. Report the pH at the midpoint.

10. Calculate the pKa and Ka of the weak acid from your data. In Figure 5, the pH at the midpoint was 3.36. This means that the pKa of the acid would be 3.36 and Ka would be 10-3.36 = 4.4 x 10-4. Review the rules for keeping track of significant figures in logarithm calculations.

11. Look at the Ka tables in the textbook and identify the unknown acid. Figure 4. Locate the midpoint region of the titration curve.

Figure 5. Identify pH at the midpoint.

Part E: Buffer Action

You will probably find it easier to understand the procedures below if you keep in mind the meaning of “buffer action” and the effect of the presence of a buffer in a solution. In this part of the experiment, you are examining the effect on the pH as you add NaOH to a buffer solution and compare it with the effect of adding NaOH to a non-buffer solution. Do you expect the two solutions to react the same?

Before you begin, predict how you expect the pH to change in the two cases.

Supplies: pH meter, pH electrode, stir plate, magnetic stir bar, 50 mL buret, 20 mL volumetric pipet, 10 mL volumetric pipet, pipet pump (or bulb), 4 small beakers, 25 mL pipet, squeeze bottle of DI water, wax pencil, 0.1 M KNO3, 0.100 M CH3COOH, 0.100 M NaOH, buffers of pH 4.00 and 7.00 Protocol

1. Into separate small, clean beakers, obtain about 30 mL of 0.100 acetic acid and about 30 mL of 0.100 M sodium hydroxide. Label the beakers. Record the exact concentration of each solution from the reagent bottles.

2. Use the volumetric pipet to dispense 20.00 mL of 0.100 M acetic acid into a separate small, clean beaker.

3. Add 10.00 mL (using volumetric pipet) of DI water to the acetic acid.

4. Add 10.00 mL (using volumetric pipet) of 0.100 M sodium hydroxide.

5. Mix the buffer solution thoroughly and label the beaker.

6. Set up a buret containing 0.10 M NaOH.

7. Use the pipet to measure 25.0 mL of the buffer solution into a small clean, dry beaker.

8. Place the pH electrode and stir bar into the beaker. Place the beaker on the stir plate and stir on medium speed.

9. Place the pH electrode into the beaker and measure the initial pH. Record the pH and initial buret reading.

10. While mixing continuously, add the NaOH dropwise from the buret until the pH changes by 0.20-0.40 pH units. Allow the pH to stabilize, record the pH and buret reading.

11. Repeat this process for a second 0.20-0.40 pH unit change, then a third, etc., until a total change of about 2 pH units from the initial pH is observed. In other words, perform 6-8 additions of NaOH.

12. Repeat the steps above using 25.0 mL of the 0.1 M KNO3 solution in place of the buffer solution. Remember to record the new initial buret reading.

13. Add the NaOH dropwise and mix well after each drop. Be patient! Read the pH after each drop, but only record the pH when it has changed by 0.20-0.40 pH units. It is possible that the first drop will give a larger change than this. If this occurs, record the pH and buret reading after the first drop of NaOH.

14. When the pH has changed by 2 or more units, add NaOH in larger increments, recording as before. There is no need to have all increments be the same size. The increment size may be increased when the pH is no longer changing rapidly. Approximately the same total number of increments and approximately the same total volume of NaOH should be added to the KNO3 solution as were added to the buffer solution.

15. Stop when the total volume of NaOH added is about the same as the total volume of NaOH added to the buffer solution. You should have performed 6-8 additions of NaOH at this point.

Calculations – Complete outside of lab

1. Calculate the number of moles of sodium hydroxide and moles of acetic acid used in preparation of the buffer solution.

2. Write the balance equation for the reaction of sodium hydroxide with acetic acid.

3. Assuming the reaction goes to completion, calculate the moles of acetate ions formed and the moles of excess acetic acid left over in the reaction.

4. Calculate the concentrations of acetate ion and acetic acid in this mixture. Assume the solution volumes are additive.

5. Calculate the expected pH of this acetic acid-acetate solution using the Henderson-Hasselbalch equation. For acetic acid, Ka is 1.8 x 10-5.

6. From the initial and subsequent buret readings, determine the cumulative volume of NaOH added for each of the pH values recorded for the buffer solution and then for the KNO3 solution. You will need to subtract the initial buret reading from every reading.

7. Prepare a graph of the observed pH versus cumulative volume of NaOH added by entering data for both the buffer solution and the KNO3 solution on the same spreadsheet, with one set of data immediately followed by the second set. The graph should now show the points for both data sets.

Adjust the x-axis minimum (Low for the x-axis) so that the first points (volume NaOH added = 0.0 mL for each solution) are visible on your graph. Use a smooth curve for this graph to show the trend, but do NOT insert a trendline.

Make certain you distinguish the two sets of data from each other and include a descriptive caption for your figure. Report Include the following information at the top of the first page of your report: Name: Lab Section Number: Experiment Title: Due Date:

General Instructions

The entire report should be word processed (typed), as handwritten portions will NOT be graded. The report should be submitted to Blackboard before the due date. Organize the report as shown below. All data in the report must match the data recorded by you in your laboratory notebook.

The first report will cover Parts A-C and the second report will cover Parts D and E. Data and Calculations See each part above for instructions on calculations and graphs.

Part A:

plot of [H3O+] vs [acid] for both hydrochloric acid and acetic acid on same plot

Part B:

plot of pH vs volume NaOH added, work to determine pH and NaOH volume at equivalence point from plot, calculate moles HCl in unknown, calculate [HCl]

Part C:

anion or cation active species in each solution, balance equations for reaction of each active ion with water

Part D:

plot of pH vs volume NaOH added, work to determine pH and NaOH volume at equivalence point, calculate volume NaOH at midpoint, work to determine pH at midpoint, calculate weak acid pKa and Ka, and identify unknown weak acid

Part E:

calculate moles sodium hydroxide and acetic acid using to prepare buffer solution, write balanced equation for reaction of acetic acid with sodium hydroxide, calculate moles acetate formed and moles acetic acid leftover when reaction goes to completion, calculate [acetate] and [acetic acid], calculate buffer pH using H-H, plot of pH vs vol NaOH added to both buffer solution and potassium nitrate solution on same plot Questions All answers should be in complete sentences and in your own words. Parts A, B and C

Questions:

1. Compare the slopes of the acids in the graph from Part A. Clearly explain what chemical property relates to the difference in slopes between the two acids.

2. Write the balanced chemical equation for the titration performed in Part B.

3. Was the pH at the equivalence point in Part B acidic, basic or neutral? Explain why the pH makes sense based on the chemical equation for Part B.

4. In Part C, the pH of a 0.1 M KNO3 solution was measured. How would the pH of a 0.1 M NaNO3 solution compare with that of a 0.1 M KNO3 solution?

Parts D and E

Questions:

1. What was the pH at the equivalence point in Part D? Was it acidic, basic or neutral?

2. What chemical reaction is taking place at the equivalence point of the titration in Part D that gives rise to the pH at the equivalence point? Use the generic formula HA to represent the unknown weak acid and write a chemical equation to help explain your answer.

3. Would you expect the equivalence point for the titration of a weak base (such as ammonia) with a strong acid (such as HCl) to be acidic, basic or neutral? Write a chemical equation to help explain your answer.

4. For Part E, compare the pH changes for the buffer solution with the KNO3 solution upon addition of NaOH.

Did the two solutions have large or small changes in pH with additions of NaOH (compare slopes on plot)? Do the results support the definition of a buffer? Explain.

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