Acid Base Titration with a pH meter Lab Report

Acid-Base Titration with a pH meter

NAME: ______________________________ Date: ______________________

Data sheet: Strong Acid – Strong Base Titration

Molarity of HCl ~0.105 M Molarity of NaOH ~0.0983 M

Volume of NaOH added at which phenolphthalein indicator turns pink __Between 3.25 and 9.00 mL

 Initial Burette Reading = 0.0 mL pH Volume of NaOH (mL) pH Volume of NaOH (mL) 1.60 0.00 2.85 26.00 1.80 2.00 2.95 26.50 1.90 3.00 3.30 27.00 2.00 4.00 10.00 27.50 2.10 6.00 11.20 28.00 2.15 8.00 12.00 28.50 2.20 10.00 12.30 29.00 2.30 12.00 12.40 30.00 2.35 14.00 12.45 31.00 2.40 18.00 12.50 32.00 2.45 20.00 12.55 34.00 2.50 21.00 12.60 36.00 2.55 22.00 2.60 22.50 2.60 23.00 2.70 23.50 2.75 24.50 2.80 25.5

Plot the data above in excel with pH values on the Y axis and the volume of NaOH added on the X axis

Analysis of the graph: Pay attention to the notes below.

Extend the line between the 2:00 mL and 26 mL points; and between 31 to 36 ml. Find the mid-point between these two lines. Through this point draw a line parallel to the x-axis up to the y-axis. This should give you the pH value at the equivalence point.

Data: Weak Acid – Strong Base Titration

Molarity of Acetic acid 0.105 M Molarity of NaOH 0.0950 M

Volume of NaOH added at which phenolphthalein indicator turns pink: Between 6.6 & 9.4 mL

 Initial Burette Reading = 0.0 mL pH Volume of NaOH (mL) pH Volume of NaOH (mL) 2.80 0.0 5.55 23.5 3.30 2.0 5.65 24.5 3.40 3.0 5.75 25.5 3.50 4.0 5.9 26.5 3.70 5.0 6.2 27.0 3.80 6.0 6.6 27.5 3.90 7.0 9.4 28.0 4.00 8.0 10.8 28.5 4.20 9.0 11.3 29.0 4.30 10.0 11.5 29.5 4.40 11.0 11.8 30.0 4.60 12.5 12.0 31.0 4.70 14.5 12.2 32.0 4.80 16.5 12.2 33.0 4.90 18.5 12.3 34.0 5.20 19.5 12.35 36.0 5.25 20.5 12.45 38.0 5.35 21.5 12.45 40.0 5.45 22.5

Plot the data above in excel where pH values are on Y axis & volume of NaOH added is on the X axis

Analysis of the graph:

Extend the line between the 2:00 mL and 26 mL points; and between 32 to 38 ml. Find the mid-point between these two lines. Through this point draw a line parallel to the x-axis to the y-axis. This should give you the pH value at the equivalence point.

Calculations: Part A: Titration of HCl with NaOH solution

Table 1: Predicted equivalence point calculations for a strong acid-strong base titration

 *PredictedEquivalence Predicted pH [H3O+] [OH–] Table 2: Calculations for titration of a strong acid with a strong base

Read the pH values from the graph at the required volume.

 Volume of NaOH added pH [H3O+] [OH–] *Before adding NaOH After adding 10.00 mL NaOH 0.50 mL Before the equivalence point At the equivalence point 0.50 mL past the equivalence point Volume after last addition36.00

*a. Determine the volume of base required by calculating the moles of acid and divide it by the molarity of the base

*b. Read pH from the graph for the required volume and calculate [H+] = 10-pH

Part B: Titration of Acetic Acid with NaOH solution

Table 3: Predicted equivalence point calculations for a weak acid-strong base titration

 *PredictedEquivalence Predicted pH [H3O+] [OH–] Table 4: Calculated for titration of a weak acid with a strong base

Read the pH values from the graph at the required volume.

 Volume of NaOH added pH [H3O+] [OH–] *Before adding NaOH After adding 10.00 mL NaOH 0.50 mL Before the equivalence point At the equivalence point 0.50 mL past the equivalence point Volume after last addition

*a. Determine the volume of base required by calculating the moles of acid and divide it by the molarity of the base.

*b. Predict the initial pH by calculating , and compare it with that from the graph.

*c. Read pH from the graph for the required volume and calculate [H+] = 10-pH

Acid-Base Titration with a pH meter

Name: _________________________ Date: _______________

Post-Lab Questions

1. Compare the predicted equivalence point values for pH, [H3O+], and [OH] to the experimentally determined equivalence point values of pH, [H3O+], and [OH] for the titration of HCl with NaOH, how do they compare? Why might they not be equal?
2. Compare the predicted equivalence point values for pH, [H3O+], and [OH] to the experimentally determined equivalence point values of pH, [H3O+], and [OH] for the titration of acetic acid with NaOH, how to they compare? Why might they not be equal?
3. Compare the pH of the solution at the indicator endpoint vs the equivalence point. Are they the same values? Explain why they are the same or why they are different. What is the pKa of phenolphthalein and methyl orange?
4. Calculate the pH and pOH at the equivalence point for 25.00 mL of 0.250 M CH3COOH, when it is titrated with 25.00 mL of 0.250 M NaOH. Ka for CH3COOH is 1.8 x 10-5 M. Write the ICE tables for calculating the pH and pOH.

Name: __________________________ Date: ____________________

Pre-Lab Questions

1. Describe the safety precautions needed when using HCl and sodium hydroxide solutions.
2. A student titrated 20.0 mL of 0.413 M HCl with 0.321 M NaOH
1. Determine the volume of NaOH needed at equivalence point; what should the pH be at the equivalence point?
 Volume NaOH added pH 0.00 0.39 2.00 0.46 4.00 0.54 6.00 0.62 8.00 0.7 10.00 0.78 12.00 0.87 14.00 0.96 16.00 1.07 18.00 1.19 20.00 1.35 22.00 1.56 24.00 1.93 24.50 2.09 25.00 2.35 25.50 3.06 26.00 11.4 26.50 11.8 27.00 12 28.00 12.2 29.00 12.3
1. The student collected the following data:
1. Plot the data above in excel with pH values on the Y axis and the volume of NaOH added on the X axis
2. Graphically determine the equivalence point and indicate the pH on the graph.

Show the following calculations on a separate sheet of paper:

1. Calculate the [H3O+] Before adding NaOH
2. Calculate the [H3O+] After adding 10.00 mL NaOH
3. Calculate the [H3O+] 2 mL Before the equivalence point
4. Calculate the [H3O+] At the equivalence point
5. Calculate the [H3O+] 2 mL past the equivalence point
6. Calculate the [H3O+] Volume after last addition

Acid-Base Titration with a pH meter

Introduction

In this experiment we will be conducting acid-base titrations of strong and weak acids with a strong base. We will monitor the progress of an acid-base titration by measuring the pH of the acid solution as a function of the volume of the base added. A plot between the pH and volume of the base added is called a titration curve. The titration can be monitored by using indicators and measuring pH of the solution. The equivalence/end point of an acid-base titration is determined visually as a color change of an indicator. Common indicators are phenolphthalein and methyl orange.

We will be using two acids: HCl and CH3COOH. HCl is a strong acid and CH3COOH is a weak acid. The nature of the two titrations curves is different due to the nature of the acids. HCl is a strong acid and dissociates completely into hydronium ions (H3O+) and chloride ions as given below. CH3COOH is a weak acid and ionizes to about 5% in aqueous solutions.

HCl (aq) + H2O (l) H3O+ (aq) + Cl (aq) (1)

CH3COOH (aq) + H2O (l) H3O+ (aq) + CH3COO (aq) (2)

Sodium hydroxide is a strong base and dissociates completely in water to form sodium and hydroxide ions as shown below.

NaOH (aq) Na+ (aq) + OH(aq) (3)

The acid-base titration of hydrochloric acid and acetic acid with sodium hydroxide is described by the following molecular, complete ionic and net ionic equations 3, 4 & 5 respectively.

HCl (aq) + NaOH (aq) NaCl (aq) + H2O (l) (4)

H+ (aq) + Cl(aq) + Na+ (aq)+ OH (aq) Cl (aq) + Na+ (aq) + H2O (l) (5)

H+ (aq) + OH (aq) H2O (l) (6)

For both the titrations, the net ionic equation is given by equation 6.

The acid-base titrations can be monitored by determining changes in the [H3O+]. Typically a pH indicator is added to the analyte solution and pH recorded as the base is added to the acid. The indicator color change is near but not exactly at the equivalence point. In this experiment the endpoint of the titration is indicated by phenolphthalein. At the equivalence point of the titration, moles of titrant OH are equal to the moles of H3O+. In this experiment, we will monitor the changes in [H3O+] by measuring the pH of the solution.

The pH of a solution is equal to the negative logarithm of hydrogen ion concentration and is represented by equation 7 below.

pH =-log [H3O+] (7)

Figures 1a. and 1b. below, show titration curves for strong acid with a strong base and a weak acid with a strong base respectively.

For the first titration, we will use HCl and NaOH and for the second CH3COOH and NaOH. In a strong acid-strong base (SA-SB) titration, the initial pH of the solution is low as only the completely dissociated acid is present. The pH increases gradually until the titration is close to the equivalence point, where a dramatic change occurs. This behavior is due to the fact that early in the titration there is a relatively large amount of H+ in solution and the addition of OH produces a small change in pH. However, near the equivalence point, a very small amount of H+ is present and addition of a very small amount of OH produces a large change in pH. After the equivalence point, as more base is added, the pH does not change and reaches a steady state.

At the equivalence point, moles acid = moles base, and the solution contains only water and the salt (NaCl) from the cation of the base and the anion of the acid. Figure 1: Titration curve

The titration curve for the acetic acid and sodium hydroxide titration is distinctly different from SA-SB titration. The initial pH of the acid is low, but higher than that of the SA-SB titration. As the base is added, pH rises steadily and has a steeper slope than the SA-SB titration. Just before the equivalence point, the pH rises sharply, however, the height of the rise is less than that of the SA-SB titration. At the end of the titration, the pH is high (similar to the SA_SB titration) as only NaOH is present in solution. All the H3O+ ions are neutralized.

We will compare the predicted equivalence points to the experimentally determined equivalence point for the two titrations.

To calculate a titration curve for a strong acid with a strong base, calculations must be performed at several titration points: initial, pre-equivalence, equivalence and post-equivalence points.

SA and SB are completely dissociated and the calculations to obtain the pH curve for titrations are quite straightforward. In the pre-equivalence stage, we can calculate the concentration of the acid from the starting concentration and the volumetric data (amount of titrant added). At the equivalence point the H3O+ and OH concentrations are equal and the concentrations can be derived from the ion-product constant of water, Kw. Post-equivalence, the excess OHconcentration can be calculated from the volumetric data.

Calculation of the equivalence point for a WA-SB titration involves the reaction of the conjugate base of the weak acid with water.

A useful expression for pOH and pH comes from the ion-product constant, Kw for water.

Kw = [H3O+] [OH] = 1.0 x 10-14 (8)

Taking the logarithms of Kw, [H3O+] & [OH] results in the pKw, pH and pOH values.

-log Kw = – log [H3O+] – log [OH]

pKw = – log (1.0 x 10-14) = 14.00 = pH + pOH (9)

A sample calculation for a titration curve for a titration of 50 mL of 0.10 M HCl with 0.10 M NaOH is shown below.

• The initial pH of the acid is

pH = – log [H3O+] = -log (0.10) = 1.00

• After addition of 10 mL 0.1 M NaOH, the total solution volume increases to 60 mL and the concentration of the acid is decreased.

Initial # of moles of HCl = 0.10 M * 0.050 L = 0.005 mol

Moles of NaOH added = 0.10 M* 0.010 L= 0.001 mol

Total volume of the solution = 50.0 + 10.0 Ml = 60 ml

[HCl] = 0.005 – 0.001 mol = 0.004 mol = 6.67 x 10-2 M

0.060 L 0.060 L

pH = – log [H3O+] = -log (6.67 x 10-2) = 1.18

• Other pre-equivalence points can be determined in a similar way. At the equivalence point, the H3O+and OHconcentrations are equal and we can use Kw for water to calculate the pH.

Kw = [H3O+] [OH] = 1.0 x 10-14

[H3O+] = = 1.0 x 10-7

pH = 7.00

• At post-equivalence, the solution now contains an excess of OH, and we can switch to calculating [OH]. For example after addition of 60 mL 0.1 M NaOH,

Initial # of moles of HCl = 0.10 M * 0.050 L = 0.005 mol

Moles of NaOH added = 0.10 M* 0.060 L= 0.006 mol

Total volume of the solution = 50.0 + 60.0 Ml = 110 ml

[NaOH] = 0.006 – 0.005 mol = 0.001 mol = 9.09 x 10-3 M

0.110 L 0.110 L

pOH = – log [OH] = -log (9.09 x 10-3) = 2.04

pH = 14.00-2.04 = 11.96

The data from several calculations for different NaOH additions produces a sigmoidal titration curve shown above in figure 1

The characteristics of the SA-SB and WA-WB titrations are:

1. The pH at the equivalence point of a SA-SB titration is 7. The solution at the equivalence point of a strong acid is truly neutral only when a strong acid is titrated with a strong base and vice-versa.
2. The pH at the equivalence point of a WA-SB titration is not 7. The pH > 7 in a WA-WB titration due to presence of the conjugate base of the WA.

Phenolphthalein indicator solution is used in each of the titrations for visual assistance. The pH will begin to change rapidly as the equivalence point are approached. Thus, the equivalence point can be detected simply by monitoring the pH of the analyte solution. Also, the equivalence point of the titration can be determined from the titration curve. It is the volume of titrant where the slope of the titration curve is the greatest. You will be asked to compare the endpoints predicted by the indicators to the equivalence points determined from the titration curves. You will also compare the predicted pH for various amounts of titrant added in each titration to the experimentally determined pH values.

Supplies

• Vernier Labquest
• Lab pH Measurement Instruction Sheet
• pH electrode in buffer solution
• 30 mL beakers
• 50 mL beakers
• 250 mL beaker (2)
• 25.00 mL volumetric pipet
• Pipet bulb
• Magnetic stir plate
• Magnetic stir bar
• 50 mL burette
• Ring stand
• Utility clamp
• burette clamp
• Deionized water squirt bottle
• Kimwipes
• ~0.10 M sodium hydroxide (NaOH)
• Phenolphthalein solution
• ~0.10 M hydrochloric acid (HCl)
• ~0.10 M acetic acid solution (HC2H3O2)
• pH 4.00 buffer
• pH 7.00 buffer
• pH 10.00 buffer

Procedure:

In this experiment, you will be using pH electrodes connected to the LabQuest interface. pH electrodes have a thin glass bulb at the tip. They break easily and are costly to replace. Be careful not to shove the electrode into the bottom of a beaker or drop the electrode. Please use extreme care when using this equipment.

Calibrating the pH Electrode

1. Make sure the pH electrode is plugged into the interface.
2. Calibrate the pH electrode using the instructions provided to you in a separate handout.
3. The calibration standards for the pH electrode will be a pH = 4.00 (pink) buffer solution, a pH = 7.00 (yellow) buffer solution, and a pH = 10.00 (blue) buffer solution. Use about 30 mL of each in 50 mL beakers.
4. Make sure the electrode is immersed in the solution and allow for a few seconds equilibration.
5. Rinse the electrode with distilled water in between solutions.

Part A Titration of HCl with NaOH:

1. Using a clean, dry 100 mL beaker, obtain about 50 mL of 0.100 M HCl solution. Record the exact concentration on your datasheet.
2. Condition a 25.00 mL volumetric pipet with the HCl solution.
3. Pipet 25.00 mL of the HCl solution into a 250 mL beaker, and add 3 drops of phenolphthalein indicator. Figure 2 Apparatus

1. Obtain some 0.100 M NaOH solution, condition the 50.00 mL burette with NaOH solution.
2. Fill the burette with 0.100 M NaOH and carefully clamp it with the burette clamp to the ring stand. Record the initial volume on the data sheet.
3. Carefully slide the stir bar into the 250 mL beaker while tilted to avoid splashing or damage to the beaker. Position the stir plate under the 250 mL beaker and begin stirring slowly.
4. Position the burette so that the tip of the burette is just inside the beaker
5. Carefully position the pH electrode in the 250 mL breaker until about 1/2 inch of the tip is in the solution. Clamp the electrode to the ring stand with a utility clamp. Be sure that the stir bar will not strike the pH electrode. If necessary, add distilled water from a graduated cylinder. See Figure 2 for the complete apparatus.
6. Measure the pH of the HCl solution and record the pH on your data sheet. Begin adding the sodium hydroxide solution to the HCl solution in 2 mL increments from the burette, stir about 15 seconds. Then read the exact volume on the burette, Record the new volume and pH after each addition. Record the data on the data sheet.
7. Remember to read the burette to the nearest 0.01 mL. Reading a burette to this accuracy is tricky; the last significant figure is expected to be an estimate.
8. When the pH begins to change more rapidly (or when you are within 2 mL of the predicted equivalence point), the increments of titrant should be changed to 0.5 mL. Note the pH range over which the indicator changes color on your Data Sheet.
9. Record the volume and pH the indicator turns pink
10. Return to 2 mL increments of titrant as the changes in pH decrease beyond the equivalence point. Do not stop the titration until the solution has reached a pH of approximately 11.5.
11. When you are finished with your titration, carefully remove the pH electrode from the solution, rinse it off with distilled water and place it in the storage buffer until you are ready to use it in Part B
12. Rinse the titration mixture down the drain with plenty of water.

Part B Titration of a weak acid with NaOH:

1. Condition a 25.00 mL volumetric pipet with the acetic acid solution.
2. Pipet 25.00 mL of the acetic acid solution into a 250 mL beaker, and 3 drops of phenolphthalein indicator.
3. Refill the burette with NaOH and carefully clamp it with the burette clamp to the ring stand. Record the initial volume on the data sheet.
4. Carefully slide the stir bar into the 250 mL beaker while tilted to avoid splashing or damage to the beaker. Position the stir plate under the 250 mL beaker and begin stirring slowly.
5. Position the burette so that the tip of the burette is just inside the beaker
6. Carefully position the pH electrode in the 250 mL breaker until about 1/2 inch of the tip is in the solution. Clamp the electrode to the ring stand with the clamp provided. Be sure that the stir bar will not strike the pH electrode. If necessary, add distilled water from a graduated cylinder.
7. Measure the pH of the acetic acid solution. Begin adding the sodium hydroxide solution to the acetic acid solution in 2 mL increments from the burette, stir about 15 seconds. Then read the exact volume on the burette, Record the new volume and pH after each addition.
8. Remember to read the burette to the nearest 0.01 mL. Reading a burette to this accuracy is tricky; the last significant figure is expected to be an estimate.
9. When the pH begins to change more rapidly (or when you are within 2 mL of the predicted equivalence point), the increments of titrant should be changed to 0.5 mL. Note the volume and pH the indicator turns pink on the datasheet.
10. Return to 2 mL increments of titrant as the changes in pH decrease beyond the equivalence point. Do not stop the titration until the solution has reached a pH of approximately 11.5
11. When you are finished with your titration, carefully remove the pH electrode from the solution, rinse it off and place it in storage buffer.

Waste Disposal and cleanup

• Calibration buffers and titration solutions can be rinsed down the drain with lots of water
• Empty the sodium hydroxide from the burette into a waste beaker, neutralize with acetic acid and rinse down the drain with lots of water.
• Rinse the burette with tap water 3 times, then 2 times with distilled water. Clamp the burette upside-down in the burette clamp to dry
• Wash all used glassware with soap and water, rinse and return to the appropriate location.
• Refill the distilled water bottle at your station.

NAME: __________________________________ Date: _________________________

Partner: _______________________________

Data sheet: Strong Acid – Strong Base Titration

Molarity of HCl ~0.10 M Molarity of NaOH ~0.105 M

Volume of NaOH added at which phenolphthalein indicator turns pink ______

Plot the data above in excel with pH values on the Y axis and the volume of NaOH added on the X axis

Molarity of Acetic acid Molarity of NaOH

Volume of NaOH added at which phenolphthalein indicator turns pink _________

Plot the data above in excel where pH values on the Y axis and the volume of NaOH added is on the X axis

Calculations: Part A: Titration of HCl with NaOH solution

Table 1: Predicted equivalence point calculations for a strong acid-strong base titration

 PredictedEquivalenceVolume NaOH (use the [HCl] & [NaOH] and volume of acid pH [H3O+] [OH–]

Table 2: Calculations for titration of a strong acid with a strong base

 Volume of NaOH added pH [H3O+] [OH–] Before adding NaOH After adding 10.00 mL NaOH 0.50 mL Before the equivalence point At the equivalence point 0.50 mL past the equivalence point Volume after last addition

Part B: Titration of Acetic Acid with NaOH solution

Table 3: Predicted equivalence point calculations for a weak acid-strong base titration

 PredictedEquivalenceVolume NaOH (use the [HCl] & [NaOH] and volume of acid pH [H3O+] [OH–]

Table 4: Calculated for titration of a weak acid with a strong base

 Volume of NaOH added pH [H3O+] [OH–] Before adding NaOH After adding 10.00 mL NaOH 0.50 mL Before the equivalence point At the equivalence point 0.50 mL past the equivalence point Volume after last addition

Acid-Base Titration with a pH meter

Name: _______________________________ Partner: _____________________________

Post-Lab Questions

1. Compare the predicted equivalence point values for pH, [H3O+], and [OH] to the experimentally determined equivalence point values of pH, [H3O+], and [OH] for the titration of HCl with NaOH, how do they compare? Why might they not be equal?
2. Compare the predicted equivalence point values for pH, [H3O+], and [OH] to the experimentally determined equivalence point values of pH, [H3O+], and [OH] for the titration of acetic acid with NaOH, how to they compare? Why might they not be equal?
3. Compare the pH of the solution at the indicator endpoint vs the equivalence point. Are they the same values? Explain why they are the same or why they are different. What is the pKa of phenolphthalein and methyl orange?
4. Calculate the pH and pOH at the equivalence point for 25.00 mL of 0.250 M CH3COOH when it is titrated with 25.00 mL of 0.250 M NaOH. Ka for CH3COOH is 1.8 x 10-5 M. Write the ICE tables for calculating the pH and pOH.

Note: weak acid, its conjugate base reacts with water to produce OH-

Name: _______________________________ Partner: _____________________________

Pre-Lab Questions

1. Describe the safety precautions needed when using HCl and sodium hydroxide solutions.
2. A student titrated 20.0 mL of 0.413 M HCl with 0.321 M NaOH
1. Determine the volume of NaOH needed at equivalence point; what should the pH be at the equivalence point?
 Volume NaOH added pH 0.00 0.39 2.00 0.46 4.00 0.54 6.00 0.62 8.00 0.7 10.00 0.78 12.00 0.87 14.00 0.96 16.00 1.07 18.00 1.19 20.00 1.35 22.00 1.56 24.00 1.93 24.50 2.09 25.00 2.35 25.50 3.06 26.00 11.4 26.50 11.8 27.00 12 28.00 12.2 29.00 12.3
1. The student collected the following data:
1. Plot the data above in excel with pH values on the Y axis and the volume of NaOH added on the X axis
2. Graphically determine the equivalence point and indicate the pH on the graph.

Show the following calculations on a separate sheet of paper:

1. Calculate the [H3O+] Before adding NaOH
2. Calculate the [H3O+] After adding 10.00 mL NaOH
3. Calculate the [H3O+] 2 mL Before the equivalence point
4. Calculate the [H3O+] At the equivalence point
5. Calculate the [H3O+] 2 mL past the equivalence point
6. Calculate the [H3O+] Volume after last addition