Buffers (College Board AP® Chemistry): Exam Questions

At high temperatures, N₂ (g) and O₂ (g) can react to produce nitrogen monoxide, NO (g), as represented  by the equation below.

  N₂ (g) + O₂ (g)  2NO (g)

Write the expression for the equilibrium constant, K_P , for the forward reaction.

A student injects N₂ (g) and O₂ (g) into a previously evacuated, rigid vessel and raises the temperature of the vessel to 2000° At this temperature the initial partial pressures of N₂ (g) and O₂ (g) are 6.01 atm and 1.61 atm, respectively. The system is allowed to reach equilibrium. The partial pressure of NO (g)  at equilibrium is 0.122 atm. Calculate the value of K_P.

Nitrogen monoxide, NO (g), can undergo further reactions to produce acids such as HNO₂, a weak acid with a K_a of 4.0 x 10⁻⁴ and a pK_a of 3.40.

A student is asked to make a buffer solution with a pH of 3.40 by using 0.100 M HNO₂ (aq) and 0.100 M NaOH (aq).

i) Explain why the addition of 0.100 M NaOH (aq)  to 0.100 M HNO₂ (aq) can result in the formation of a buffer solution. Include the net ionic equation for the reaction that occurs when the student adds the NaOH (aq)  to the HNO₂ (aq).

ii) Determine the volume, in mL, of 0.100 M NaOH (aq)  the student should add to 100. mL of 0.100 M HNO₂ (aq) to make a buffer solution with a pH of 3.40. Justify your answer.

A second student makes a buffer by dissolving 0.100 mol of NaNO₂ (s) in 100. mL of 1.00 M HNO₂ (aq). Which is more resistant to changes in pH when a strong acid or a strong base is added, the buffer made by the second student or the buffer made by the first student in part (c)? Justify your answer.

A new buffer is made using HNO₂(aq) as one of the A particulate representation of a small representative portion of the buffer solution is shown below. (Cations and water molecules are not shown.)

Is the pH of the buffer represented in the diagram greater than, less than, or equal to 3.40 ? Justify your answer.

Two buffer solutions are prepared:

• Buffer 1: 1.00 M CH₃COOH + 1.00 M CH₃COO^-

• Buffer 2: 0.10 M CH₃COOH + 0.10 M CH₃COO^-

A student adds 0.05 moles of HCl to both 1.0 L buffer solutions.

Define buffer capacity.

Predict which buffer in part (a) will experience a smaller pH change upon HCl addition and justify your answer.

Explain why the ratio of [CH₃COO^-] to [CH₃COOH] remains nearly constant in Buffer 1 in part (a).

Use Le Châtelier’s principle to explain how the addition of HCl affects the equilibrium in buffer 2.

A buffer solution is prepared by dissolving 0.25 moles of acetic acid (CH₃COOH) and 0.15 moles of sodium acetate (CH₃COONa) in 1.00 L of water. The pK_a of acetic acid is 4.76.

Use the Henderson-Hasselbalch equation to calculate the pH of the buffer solution.

The student accidentally adds 0.05 moles of NaOH to the buffer. Predict whether the pH will increase, decrease, or remain unchanged. Justify your answer

The student decides to dilute the buffer solution by adding 500.0 mL of distilled water. Explain why this dilution affects the buffer’s capacity but not its pH.

The solubility of calcium fluoride (CaF₂) in water is governed by the following equilibrium:

CaF₂ (s) ⇌ Ca²⁺ (aq) + 2F⁻ (aq)

A student investigates how pH affects the solubility of CaF₂.

Write the expression for the solubility product constant (K_sp) for CaF₂.

Explain how the solubility of CaF₂ would change if HCl is added to the solution. Justify your answer using equilibrium principles

The student then adds NaF instead of HCl. Predict and explain how the solubility of CaF₂ would be affected.

A marine chemist is investigating how the H₂CO₃/HCO₃^- buffer system regulates pH in seawater.

To model this system, the chemist prepares a buffer by mixing:

• 75.0 mL of 0.200 M H₂CO₃

• 25.0 mL of 0.400 M NaHCO₃

At 25 ^oC, the value of Ka₁ for carbonic acid (H₂CO₃) is 4.3 × 10^-7.

Calculate the pH of the resulting buffer solution.

Explain how the H₂CO₃/HCO₃^- buffer system resists a change in pH when a small amount of HCl (aq) is added

A 5.0 mL sample of 0.100 M HCl is added to the buffer solution.

Predict whether the pH will increase, decrease, or remain approximately the same. Justify your answer.

A second buffer is prepared using:

• 75.0 mL of 0.0200 M H₂CO₃

• 25.0 mL of 0.0400 M NaHCO₃

This new buffer has the same initial pH as the original buffer.

Explain which buffer would better resist a change in pH. Justify your answer in terms of buffer capacity.

Predict what happens to pH if all H₂CO₃ is neutralized by NaOH. Justify your answer.

A student investigates the solubility of lead(II) oxalate (PbC₂O₄), which has the following solubility equilibrium:

PbC₂O₄ (s) ⇌ Pb²⁺ (aq) + C₂O₄²⁻(aq)

The student prepares two test tubes:

• Test Tube 1 contains PbC₂O₄ in pure water.

• Test Tube 2 contains PbC₂O₄ in a solution of HCl (pH 2.0).

The image shows the difference in PbC₂O₄ solubility between the two solutions.

Write the solubility product (K_sp) expression for PbC₂O₄.

Explain how the solubility of PbC₂O₄ changes in the acidic solution. Justify your answer using Le Châtelier’s Principle.

The student wants to confirm that the increased solubility of PbC₂O₄ in acidic conditions is due to the removal of oxalate ions (C₂O₄^2-) rather than any direct reaction between HCl and PbC₂O₄.

Describe an additional experiment the student could perform to test this hypothesis.

A 500.0 mL buffer solution is prepared with 0.200 M sodium bicarbonate (NaHCO₃) and 0.300 M sodium carbonate (Na₂CO₃). The pK_a​ of HCO₃^- is 4.68 x 10^-11.

Write the net ionic equation for the reaction that occurs when a strong acid (HCl) is added to this buffer system.

Calculate the initial pH of the buffer solution using the Henderson-Hasselbalch equation. Show your working.

The student adds a large excess of HCl to the buffer solution. Predict what will happen to the pH and explain why the buffer is no longer effective.

A buffer solution contains 0.300 M NH₃ (ammonia) and 0.200 M NH₄Cl (ammonium chloride). The K_b​ of ammonia is 1.8 x 10⁻⁵.

Write the equation for the equilibrium reaction that occurs in this buffer system.

Explain why the buffer in part (a) is more effective at resisting pH changes when small amounts of strong acid (HCl) are added, compared to when strong base (NaOH) is added.

Predict what happens to the buffer’s ability to maintain pH if a large excess of HCl is added, and justify your answer.