Learning Objectives:
- To test the acidic and basic properties of ionic compounds.
- To create a buffer solution and calculate its pH.
- To determine the acid dissociation constant of acetic acid using the Henderson-Hasselbach equation.
- To test a buffer’s capacity by reacting it with a strong acids and a strong base.
Discussion:
Buffers are solutions that resist changes in pH when small quantities of an acid or base are added to them. These solutions contain a weak acid and its conjugate base or a weak base and its conjugate acid. Because these mixtures contain both the acidic and basic species, they can easily neutralize small quantities of either H+ or OH–. Buffers are also resistant to dilution effects, maintaining the same pH even when water is added to them.
Part A: Acidic and Basic Properties of Ionic Compouds
The first part of the experiment involves using a pH meter to determine the pH of several salt solutions.
A pH meter consists of a meter and two electrodes. The measurement of pH requires two electrodes, a sensing electrode that is sensitive to H3O+ concentration and a reference electrode. This is because the pH meter is really a voltmeter that measures the electrical potential of a solution. The reference electrode is an electrode that develops a known potential that is essentially independent of the contents of the solution into which it is placed. The glass electrode is sensitive to the H3O+ concentration of the solution into which it is placed.
Part B: Buffers
A buffer system will be prepared from solutions of conjugate acid – base pairs. The pH of the buffer system will be determined using the pH meter. The Ka of the weak acid used in the buffer will be calculated using the pH data. The effect of adding an acid and a base to the system will be monitored.
For a weak acid,
Reaction: HA (aq) + H2O ↔ H3O+ (aq) + A– (aq)
And Ka = [H3O+] [A–]
[HA]
Knowing the experimentally determined pH of the solution the [H3O+] can be calculated from pH = -log[H3O+]
The Ka of the buffer system can be calculated from the Henderson-Hasselbalch equation
pH = pka + log [A–]
[HA]
How to Calibrate a Vernier pH meter:
You will be using two buffer solutions to calibrate the pH meter. One solution has a pH of 4 and the other has a pH of 7. You will also need a small beaker with deionized water to rinse the pH probe between each reading.
- Connect the pH probe to the LabQuest.
- Gently unscrew the cap and remove the pH probe from its protective solution. Dip the pH probe in a beaker with deionized water and gently wipe the probe with a Kim wipe.
- Next, click on “Sensors” at the upper left hand corner of the LabQuest screen.
- Click on “Calibrate” and select the correct channel where the probe is connected: CH1, CH2, or CH3 etc.
- Click “Calibrate Now” and dip the pH meter into the first buffer solution. Enter the correct pH value (7 or 4) in the window and select “Keep”
- Rinse the probe in the deionized water and wipe with a Kim wipe or paper towel. Then, dip the pH probe into the second buffer solution and enter the correct pH value (7 or 4) in the window and select “Keep”
- Click “Ok” and your pH probe is now calibrated.
Part A Procedures:
- Label five 50mL beakers with the following chemical formulas: NaCl, Na2CO3, NaC2H3O2, NH4Cl, and NaHSO4.
- Measure out 20mL of each of the 0.1M solutions of NaCl, Na2CO3, NaC2H3O2, NH4Cl and NaHSO4 into the beakers.
- Measure the pH of each solution and record it in the data table. Rinse the pH meter between each reading to avoid contaminating the solutions. Indicate which ion hydrolyzes in water and write the net ionic equation in the data table. If pH of the solution is neutral, then none of the ions hydrolyze in water and there is no net ionic equation.
- DO NOT THROW AWAY THE SOLUTIONS IN THE SINK! Return them to the original container.
Part B Procedures:
- Weigh about 3.5g of sodium acetate, record the exact mass on your data table and add it to a 150mL beaker.
- Using a 10mL graduated cylinder measure 8.8mL of 3.0M acetic acid and add it to the beaker containing the sodium acetate.
- Using a graduated cylinder measure 55.2mL of distilled water and add it to the solution in the beaker. Stir the solution until all the sodium acetate is dissolved.
- Calibrate the pH meter using the 4.00pH buffer solution. Click on the “one point calibration” option in the calibration menu. Measure the pH of your buffer solution. Determine the value of Ka for acetic acid using the pH of your buffer and the Henderson Hasselbach equation. Use the buffer solution you prepared for the next step.
- Pour half (32mL) of the buffer solution into another 150mL beaker. Label the two beakers as 1 and 2. Pipette out 1.0mL of 6.0M HCl into beaker 1, mix and then measure the pH and record it. Remember to rinse the electrode between measurements. Pipette out 1.0mL of 6.0M NaOH into beaker 2, mix and measure the pH and record it.
Calculations:
- Calculate the Ka for acetic acid using the measured pH of your buffer solution. Compare the calculated value of Ka with the standard value of Ka for acetic acid, which is 1.8×10-5.
- Calculate the pH of the original buffer you prepared using the known value of Ka for acetic acid.
- Calculate the pH of the buffer after the additions of HCl and NaOH.
- Compare the calculated values with the measured values.
CHM 112 Name: ________________________
Date: ______________ Lab Partner: ______________________
Data Sheet pH of Salt solutions and Properties of Buffers
A. pH of Salt Solutions
| Solution | pH measured | Ionization of Salt | Ion Hydrolyzed |
| NaCl | |||
| Na2CO3 | |||
| NaC2H3O2 | |||
| NH4Cl | |||
| NaHSO4 |
Net Ionic Equation
B. Study of Buffers
1. Mass of Sodium Acetate (NaC2H3O2) _____________________
2. pH of Original Buffer _____________________
3. pH of Buffer + HCl _____________________
4. pH of Buffer + NaOH _____________________
- Calculate Ka for acetic acid using measured pH of buffer solution and the Henderson Hasselbach equation.
6. Calculated pH of Original Buffer _____________________
(Show calculations below and use theoretical value Ka = 1.8×10-5)
7. Calculated pH of Buffer + HCl _____________________
(Show calculations below and use theoretical value Ka = 1.8×10-5)
- Calculated pH of Buffer + NaOH _____________________
(Show calculations below and use theoretical value Ka = 1.8×10-5)
- How do your measured pH values compare with calculated pH values in 6, 7 and 8? Give percent error.
- How does your calculated value of Ka compare with the standard value of Ka for acetic acid? Discuss why your value may be larger or smaller than the standard value.
Post Lab Questions
- Identify and discuss the implications of two sources of error you found during this lab.
- During the course of the experiment, a student prepared their buffer solution using the wrong bottle of acetic acid. Instead of using the 3.0M acetic acid, they grabbed 3.0M hydrochloric acid. How would this impact the pH of the buffer, and how the effectiveness of the buffer to resist change in pH?
- Determine the pH of a buffer created by adding these
solutions together:
- 50 mL 0.50M acetic acid and 50 mL 0.50M sodium acetate
- 20 mL 0.10M carbonic acid and 10 mL 0.10M sodium bicarbonate
- 20 mL 0.50M sodium acetate and 10 mL 0.50 hydrochloric acid
- 20 mL 0.20M NaHSO4 and 10 mL 0.20M NaOH
