What are buffers what are the two components of a

Question

Lab: Preparation and Properties of Buffers

PRE-LAB QUESTIONS:

1. What are buffers? What are the two components of a buffer?

2. Derive the Henderson-Hasselbalch equation from the Ka expression.

3. Which of the following pairs of chemicals will produce a buffer?

a. HCl and HC₂H₃O₂ c. H₂CO₃/Na₂CO₃

b. NaF and HF d. NH₃ and NaOH

4. Write balanced net ionic equations that show how a buffer containing NaHCO₃ and Na₂CO₃ maintains its pH when a small amount of HCl or NaOH is added.

5. Calculate the pH of a buffer made from mixing 10.0 mL of 0.100 M NaC₂H₃O₂ and 10.0 mL of 0.100 M HC₂H₃O₂.

a. Calculate the pH of the buffer when 5.00 mL of a 0.0100 M NaOH is added.

b. Calculate the pH of the buffer when 5.00 mL of a 0.0100 M HCl is added.

EXPERIMENTAL PROCEDURE:

Part 1: Preparation of a Buffer from a Weak Acid/Base and Its Conjugate

1. Obtain your buffer assignment (pH and total molarity) from your professor.

2. Select an appropriate pair from the list below to prepare 50.00 mL of your assigned buffer. Consult your text for K_a or K_b values.

a. NaC₂H₃O₂·3H₂O and HC₂H₃O₂ (aq., 1.00 M)
b. NaHCO₃ (s) and Na2CO3 (s)
c. Na₂HPO₄(s) and Na₃PO₄·12 H₂O (s)
d. NaH₂PO₄(s) and Na₂HPO₄(s)
e. NH₄Cl(s) and NH₃ (aq, 1.00 M)

3. Calculate the mass or volume of acid and base components needed to obtain the pH and molarity assigned (see example calculation for Method 1 in the Introduction). Note: Some salts contain waters of hydration in the formula so be sure to include them in the calculation for molar mass.) Check your answers with your professor.

4. Write a detailed procedure on how you would prepare 50.00 mL of the desired buffer using the chemicals above. Obtain instructor's approval.

5. Check out a 50 mL volumetric flask and prepare the buffer according to your procedure.

6. Calibrate the pH electrode using the pH 4/7 or pH 7/10 buffer solutions, depending on whether your buffer falls in the acidic or basic range. Follow proper calibration procedure as directed by your professor; guidelines for calibration are provided with the pH meters.

7. Pour the prepared buffer into a clean, dry 150 mL beaker and measure and record the pH of the buffer. The measured pH value of your buffer should be identical (or very close) to the assigned value. Save the prepared buffer for Part 2.

8. Rinse the volumetric flask with tap water and do a final rinse with DI water.

Part 2: Properties of Buffered and Non-Buffered Systems

1. Check out a 10 mL graduated pipet and pipet pump from your professor.

2. Obtain and label (#1-4) 4 clean and dry beakers (50 or 100 mL).

3. Add the following liquids/solutions into beakers using a 50-mL graduated cylinder.

a. Beaker 1-20.0 mL of 0.1M NaCl
b. Beaker 2-20.0 mL of 0.1 M NaCl
c. Beaker 3-20.0 mL of your prepared buffer from Part I.
d. Beaker 4-20.0 mL of your prepared buffer from Part I.

4. Measure and record the pH of the contents in each beaker. Be sure to rinse the electrode with deionized water in between uses to avoid contamination. DO NOT blot the electrode with paper towels or Kimwipes.

5. Effect of acid: Add 5.00 mL of 0.0100 M HCl solution (be sure to condition the pipet first) into beakers 1 and 3. Mix the contents well and measure and record the pH.

6. Effect of base: Add 5.00 mL of 0.0100 M NaOH solution (be sure to condition the pipet first) into beakers 2 and 4. Mix the contents well and measure and record the pH.

7. Dispose of all solutions into the chemical waste container.

Part 3: Preparation of a Buffer Using One Component (Optional)

1. Write a detailed procedure (using Method 2) for the preparation of 50.00 mL of the same assigned buffer (same pH and molarity) using just ONE of the chemicals from the pair you selected in step 2 and either 1.00 M NaOH OR 1.00 M HCl to generate the conjugate species in situ. Once your calculations and procedure are approved by your professor, check out the 50-mL volumetric flask and prepare the buffer accordingly.

2. Pour the buffer into a 100 mL beaker. Measure and record the pH of the solution.

3. Obtain two clean and dry 50-mL beakers (beakers 1 and 2). Using a 50-mL graduated cylinder, pour 20.0 mL of the buffer into each of the beakers.

4. Calculate the volume of 0.100 M HCl required to decrease the pH of the buffer solution by 0.15 pH units in beaker 1.

5. Calculate the volume of 0.100 M NaOH required to increase the pH of the buffer solution by 0.10 pH units in beaker 2.

6. Once your calculations and procedure are approved b your professor, use a conditioned pipet and add the calculated volume of acid (or base) into the beakers. Measure and record the new pH.

7. Dispose of the solutions into the waste containers. Returned the clean volumetric flask to your professor.

DATA ANALYSIS:

Part 1: Preparation of a Buffer from a Weak Acid/Base and Its Conjugate

1. Complete the table below. If the actual amounts of the weak acid/base and its conjugate used during the preparation differ from the calculated amounts, compute the new target pH.

	Target pH	Measured pH	% error
Buffer

2. What are some likely sources of error for the preparation of the buffer?

Part 2: Properties of Buffered and Non-buffered Systems

1. Using your measured pH values, calculate the concentration of H⁺ (M) in beakers 1 and 3, before and after the addition of HCl. Organize your answers in a table as shown below

	Initial [H+] (M)	[H+] (M) after addition of HCl	By what factor did the [H+] (M) increase?
Beaker 1 (0.1 M NaCl)
Beaker 3 (Buffer)

2. Using your measured pH values, calculate the concentration of H⁺ (M) in beakers 2 and 4, before and after the addition of NaOH. Organize your answers in a table as shown below.

	Initial [H+] (M)	[H+] (M) after addition of HCl	By what factor did the [H+] (M) increase?
Beaker 2 (0.1 M NaCl)
Beaker 4 (Buffer)

3. Which system (buffered or non-buffered) shows a smaller increase/decrease in the concentration of H⁺? Explain.

4. For each beaker, calculate the expected (theoretical) pH when 5.00 mL of 0.0100 M HCl or 5.00 mL of 0.0100 M NaOH is added to 20.0 mL of buffered and non-buffered systems. Use the actual volume of acid/base added in the calculation. Organize your answer in a table as shown below.

	Actual volume of acid or base added (mL)	Measured pH	Calculated (Theoretical) pH	% Error
0.1 M NaCl + acid (beaker 1)
0.1 M NaCl + base (beaker 2)
Buffer + acid (beaker 3)
Buffer + base (beaker 4)

Part 3: Preparation of a Buffer Using One Component (Optional)

1. Complete the table below.

	Target pH	Measured pH	% error
Buffer
Buffer + acid (beaker 1)
Buffer + base (beaker 2)

2. What are some likely sources of error for the preparation of the buffer and the addition of acid/base?

POST-LAB QUESTIONS:

1. What is the effect on the pH of the solutions in beaker 1 and 3 if the pipet was rinsed with DI H2O but not conditioned with the HCl solution?

2. Which of the following pairs of solutions will produce a buffer solution? Explain your reasoning.

a. 10.0 mL of 0.10 M HCl + 10.0 mL of 0.10 M NaOH

b. 10.0 mL of 0.10 M HCl + 5.0 mL of 0.10 M NaOH

c. 10.0 mL of 0.10 M HF + 10.0 mL of 0.10 M NaOH

d. 10.0 mL of 0.10 M HF + 5.0 mL of 0.10 M NaOH

3. A student prepared 50.0 mL of "Buffer Q" using 0.50 moles of HA and 0.50 moles of A- while another student prepared 50.0 mL of "Buffer S" using 0.25 moles of HA and 0.25 moles of A-.

a. Do the two buffer solutions have the same or different pH? Explain.

b. If 1.00 mL of 0.010 M NaOH were added to the two buffer solutions, would the pH of the two solutions increase or decrease? Explain.

c. Which solution (Buffer Q or Buffer S) would show a smaller change in the pH for question (b)? Explain.