AP Chemistry Progress Check

A student measures the pH of a 0.0100M buffer solution made with HClO and NaClO, as shown above. The pKa of HClO is 7.40 at 25°C. Based on this information, which of the following best compares the relative concentrations of ClO⁻ and HClO in the buffer solution?

• [ClO⁻]=[HClO]

• [ClO⁻]>[HClO]

• [ClO⁻]<[HClO]

• It is not possible to compare the concentrations without knowing the pKb of ClO⁻.

HCNO is a weak acid with a pKa value of 3.5. The graph above shows the results of a titration of an aqueous solution of HCNO with 0.100MNaOH. Based on the results, the concentration of CNO−is greater than the concentration of HCNO at which of the following pH values?

• 2.0

• 3.0

• 3.5

• 4.0

Tags

NGSS.HS-PS1-2

The weak acid ionization equilibrium for C₂H₃COOH is represented by the equation above. A student measures the pH of C₂H₃COOH(aq) using a probe and a pH meter in the experimental setup shown. Based on the information given, which of the following is true?

• [C₂H₃COOH]>[C₂H₃COO⁻] since the pKa of the weak acid is less than pKw.

• [C₂H₃COOH]>[C₂H₃COO⁻] since the pKa of the weak acid is greater than the pH of the solution.

• [C₂H₃COOH]<[C₂H₃COO⁻] since the pKb for C₂H₃COO⁻ is less than pKw

• [C₂H₃COOH]<[C₂H₃COO⁻] since the Kb for C₂H₃COO⁻ is less than the Ka for C₂H₃COOH.

The acid equilibrium for CH₃CH₂COOH and the base equilibrium for CH₃CH₂COO⁻ are represented above. One liter of a buffer solution with pH=4.85 is made by mixing 0.100M CH₃CH₂COOH and 0.100 M NaCH₃CH₂COO. If 10.0mL of 0.500MNaOH is added to the buffer, which of the following is most likely the resulting pH, and why?

• The pH will be much less than 4.85, because the buffer will respond to the addition of NaOH by producing more CH₃CH₂COOH.

• The pH will be much greater than 4.85, because there is a large increase in the concentration of the weak base, CH₃CH₂COO⁻.

• The pH will be slightly less than 4.85, because the addition of NaOH reduces the autoionization of H₂O.

• The pH will be slightly greater than 4.85, because some CH₃CH₂COOH will react with the added bases, resulting in a slight decrease in [H₃O⁺].

Tags

NGSS.HS-PS1-5

A buffer solution that is 0.100M in both HCOOH and HCOOK has a pH=3.75. A student says that if a very small amount of 0.100MHCl is added to the buffer, the pH will decrease by a very small amount. Which of the following best supports the student’s claim?

• HCOO⁻ will accept a proton from HCl to produce more HCOOH and H₂O.

• HCOOH will accept a proton from HCl to produce more HCOO⁻ and H₂O.

• HCOO⁻ will donate a proton to HCl to produce more HCOOH and H₂O.

• HCOOH will donate a proton to HCl to produce more HCOO⁻ and H₂O.

The equilibrium representing the acid dissociation of CH₃COOH is shown above. A buffer solution is prepared by adding 0.10mol of NaOH(s) to 1.00L of 0.30M CH₃COOH. Assuming the change in volume is negligible, which of the following expressions will give the pH of the resulting buffer 25°C?

To prepare a buffer solution for an experiment, a student measured out 53.49g of NH₄Cl(s) (molar mass 53.49g/mol) and added it to 1.0MNH₃(aq). However, in the process of adding the NH3(aq), the student spilled some of the NH₄Cl(s) onto the bench top. As a result, only about 50.g of NH₄Cl(s) was actually added to the 1.0MNH₃(aq). Which of the following best describes how the buffer capacity of the solution is affected as a result of the spill?

• The solution has a greater buffer capacity for the addition of base than for acid, because [NH₃]<[NH₄⁺].

• The solution has a greater buffer capacity for the addition of base than for acid, because [NH₃]>[NH₄⁺].

• The solution has a greater buffer capacity for the addition of acid than for base, because [NH₃]<[NH₄⁺].

• The solution has a greater buffer capacity for the addition of acid than for base, because [NH₃]>[NH₄⁺].

The acid ionization equilibrium for HC₃H₅O₂ is represented by the equation above. A mixture of 1.00L of 0.100M HC₃H₅O₂ and 0.500L of 0.100M NaOH will produce a buffer solution with a pH=4.87. If the NaOH solution was mislabeled and was 1.00M instead of 0.100M, which of the following would be true?

• The pH of the resulting solution would still be 4.87 because buffer solutions regulate changes in pH regardless of how much NaOH is added.

• The pH of the resulting solution would be somewhat higher than 4.87 because adding NaOH that is 10 times more concentrated makes the concentration of the conjugate base 10 times larger.

• The pH of the resulting solution would be much higher than 4.87 because the weak acid would be completely neutralized by the larger amount of NaOH added.

• The pH of the resulting solution would still be 4.87 because the solution contains a large volume of the weak acid to react with the NaOH and maintain the same pH.

The acid ionization equilibrium for HNO₂ is represented by the equation above. A 250.0mL buffer solution is prepared by mixing 125.0mL of 0.20M HNO₂ and 125.0mL of 0.1M NaOH. To test the buffer capacity, the pH is measured and recorded in the table for four samples of the buffer and one sample of a mixture of the buffer and HCl. Which of the following best helps explain why the pH of sample 4 is lower than the pH of the other samples containing only buffer solution?

• Prior to measuring the pH of sample 4, some water evaporated, resulting in an increase in the concentration of HNO₂ in the sample.

• After measuring the pH of the more acidic sample 3, the pH probe was not rinsed and wiped, resulting in the neutralization of a very small amount of the conjugate base in sample 4.

• Prior to measuring the pH of sample 4, the room temperature dropped suddenly, resulting in an increase in the pKa for HNO₂.

• The volume used to measure the pH of sample 4 was less than 25.0mL, resulting in a decrease in the concentration of NO₂⁻.

Buffer solutions containing Na₂CO₃ and NaHCO₃ range in pH from 10.0 to 11.0. The chemical equation below represents the equilibrium between CO₃²⁻and H₂O, and the table lists the composition of four different buffer solutions at 25°C.

Which of the following chemical equilibrium equations best shows what happens in the buffer solutions to minimize the change in pH when a small amount of a strong base is added?

• H₃O⁺(aq) + OH⁻(aq) ⇄ 2 H₂O(l)

• HCO₃⁻(aq)+ OH⁻(aq) ⇄ CO₃²⁻(aq) + H₂O(l)

• CO₃²⁻(aq) + H₃O⁺(aq) ⇄ HCO₃⁻(aq) +H₂O(l)

• CO₃²⁻(aq) + H₂O(l) ⇄ HCO₃⁻(aq)+ OH⁻(aq)