Acid-Base Titrations (College Board AP® Chemistry): Exam Questions

Potassium sorbate, KC₆H₇O₂ (molar mass 150 g/mol) is commonly added to diet soft drinks as a preservative. A stock solution of KC₆H₇O₂(aq) of known concentration must be prepared. A student titrates 45.00 mL of the stock solution with 1.25 M HCl(aq) using both an indicator and a pH meter. The value of K_a for sorbic acid, HC₆H₇O₂, is 1.7 × 10⁻⁵.

Write the net-ionic equation for the reaction between KC₆H₇O₂(aq) and HCl(aq).

A total of 29.95 mL of 1.25 M HCl(aq) is required to reach the equivalence point. Calculate [KC₆H₇O₂] in the stock solution.

The pH at the equivalence point of the titration is measured to be 2.54. Which of the following indicators would be the best choice for determining the end point of the titration? Justify your answer.

Calculate the pH at the half-equivalence

The initial pH and the equivalence point are plotted on the graph below. Accurately sketch the titration curve on the graph below. Mark the position of the half-equivalence point on the curve with an X.

The pH of the soft drink is 3.37 after the addition of the KC₆H₇O₂(aq). Which species, HC₆H₇O₂ or C₆H₇O₂⁻, has a higher concentration in the soft drink? Justify your answer.

Answer the following questions relating to Fe and its ions, Fe²⁺ and Fe³⁺.

Write the ground-state electron configuration of the Fe²⁺ ion.

The radii of the ions are given in the table above. Using principles of atomic structure, explain why the radius of the Fe²⁺ ion is larger than the radius of the Fe³⁺ ion.

Fe³⁺ ions interact more strongly with water molecules in aqueous solution than Fe²⁺ ions do.

Give one reason for this stronger interaction, and justify your answer using Coulomb’s law.

A student obtains a solution that contains an unknown concentration of Fe²⁺ (aq). To determine the concentration of Fe²⁺ (aq) in the solution, the student titrates a sample of the solution with MnO₄^– (aq), which converts Fe²⁺ (aq) to Fe³⁺ (aq) , as represented by the following equation.

5 Fe²⁺ (aq) + MnO₄^– (aq) + 8 H⁺ (aq) →  5 Fe³⁺ (aq) + Mn²⁺ (aq) + 4H₂O(l)

Write the balanced equation for the half-reaction for the oxidation of Fe²⁺ (aq) to Fe³⁺ (aq) .

The student titrates a 10.0 mL sample of the Fe²⁺ (aq) solution . Calculate the value of [Fe²⁺] in the solution if it takes 17.48 mL of added 0.0350 M KMnO₄ (aq) to reach the equivalence point of the titration.

To deliver the 10.0 mL sample of the Fe²⁺ (aq) solution in part (e), the student has the choice of using one of the pieces of glassware listed below.

Explain why the 25 mL volumetric flask would be a poor choice to use for delivering the required volume of the Fe²⁺ (aq) solution.

In a separate experiment, the student is given a sample of powdered Fe (s) that contains an inert impurity. The student uses a procedure to oxidize the Fe(s) in the sample to Fe₂O₃ (s) . The student collects the following data during the experiment.

Calculate the number of moles of Fe in the Fe₂O₃(s) produced.

Calculate the percent by mass of Fe in the original sample of powdered Fe (s) with the inert impurity.

If the oxidation of the Fe (s) in the original sample was incomplete so that some of the 7 .531 g of product was FeO (s) instead of Fe₂O₃, would the calculated mass percent of Fe (s) in the original sample be higher, lower, or the same as the actual mass percent of Fe (s)? Justify your answer.

A student performs a titration of 50.0 mL of 0.100 M acetic acid (CH₃COOH) with 0.100 M NaOH. The titration curve is recorded using a pH meter.

The K_a of acetic acid is 1.76 x 10^-5.

Write the net ionic equation for the reaction.

i) Describe what occurs at the equivalence point during the titration of acetic acid with sodium hydroxide.

i) Explain why the pH at the equivalence point is greater than 7.

Identify the pH at the half-equivalence point and explain how it relates to the pK_a of acetic acid.

An aqueous solution of acetic acid (CH₃COOH) is titrated with 0.150 M sodium hydroxide (NaOH).

Label the half-equivalence point on the graph by marking it with an "X."

The titration reaches the equivalence point when 25.0 mL of acetic acid has reacted with an equal number of moles of NaOH.

The K_a of acetic acid is 1.75 x 10⁻⁵

(i) Explain why the pH at the half-equivalence point is numerically equal to the pK_a of acetic acid.

(ii) Calculate the pH at the half-equivalence point.

After completing the titration, the student wants to verify their experimental results by calculating the volume of NaOH required to reach the first equivalence point.

Determine the volume of 0.150 M NaOH needed to fully neutralize 50.0 mL of 0.100 M acetic acid.

Volume of HCl (cm³)

A student performs a titration using a 0.10 M ammonia, NH₃ (aq), and a 0.10 M hydrochloric acid, HCl (aq).

Write the equation for the overall reaction that is occurring.

Mark on the curve the point at which the pOH is equal to pK_b of the weak base and deduce the pK_b of the acid.

A student repeats the titration using 0.100 M acetic acid (CH₃COOH), which has a pK_a value of 4.76. The resulting titration curve is shown below.

Explain why accurately determining the equivalence point for this titration is more challenging compared to a strong acid-strong base titration.

A student reacts 0.300 g of methyl salicylate (C₈H₈O₃) with a stoichiometric amount of a strong base. This product is then acidified to produce salicylic acid crystals (HC₇H₅O₃).

For every 1 mole of C₈H₈O₃ (molar mass 152.15 g/mol) reactant used, 1 mole of salicylic acid crystals (HC₇H₅O₃, molar mass 138.12 g/mol) is produced. Calculate the maximum mass, in grams, of HC₇H₅O₃ that could be produced in this reaction.

As part of the experimental procedure to purify the HC₇H₅O₃ crystals after the reaction is complete, the crystals are filtered from the reaction mixture, rinsed with distilled water, and dried. Some physical properties of HC₇H₅O₃ are given in the following table.

The student’s experiment results in an 87% yield of dry HC₇H₅O₃. The student suggests that some of the HC₇H₅O₃ crystals dissolved in the distilled water during the rinsing step. Is the student’s claim consistent with the calculated percent yield value? Justify your answer.

Given the physical properties in the table, calculate the quantity of heat that must be absorbed to increase the temperature of a 0.105 g sample of dry HC₇H₅O₃ (molar mass 138.12 g / mol) crystals from 25° C to the melting point of 159° C and melt the crystals completely.

The structures and melting points for methyl salicylate and salicylic acid are shown.

The same three types of intermolecular forces (London dispersion forces, dipole-dipole interactions, and hydrogen bonding) exist among molecules of each substance. Explain why the melting point of salicylic acid is higher than that of methyl salicylate.

The student titrates 20.0  mL of 0.0100 M  HC₇H₅O₃(aq) with 0.0200 M  NaOH, using a probe to monitor the pH of the solution. The data are plotted producing the following titration curve.

Using the information in the graph, estimate the pK_a  of HC₇H₅O₃.                                 

When the pH of the titration mixture is 4.00, is there a higher concentration of the weak acid, HC₇H₅O₃, or its conjugate base, C₇H₅O₃⁻, in the flask? Justify your answer.

The student researches benzoic acid (HC₇H₅O₂) and finds that it has similar properties to salicylic acid (HC₇H₅O₃). The K_a  for benzoic acid is 6.3 × 10⁻⁵. Calculate the value of pK_a  for benzoic acid.

The student performs a second titration, this time titrating 20.0 mL of a 0.0100 M benzoic acid solution with 0.0200 M NaOH. Sketch the curve that would result from this titration of benzoic acid on the following graph, which already shows the original curve from the titration of 20.0 mL of 0.0100 M salicylic acid. The initial pH of the benzoic acid solution is 3.11.

A 50.0 mL sample of 0.100 M phosphoric acid (H₃PO₄) is titrated with 0.100 M NaOH. The first, second, and third equivalence points occur at pH values of approximately 4.7, 9.8, and 12.0, respectively. The student must select an appropriate indicator to detect the first equivalence point and is given the indicators listed in the table.

Write the balanced chemical equation for the reaction occurring at the first equivalence point.

Identify the most appropriate indicator for detecting the first equivalence point and justify your choice.

Explain why bromothymol blue would not be a suitable indicator for detecting any of the equivalence points in this titration.

Explain why the solution acts as a buffer at the first half-equivalence point of the titration.

Sketch a titration curve for the reaction of 50.0 mL of 0.100 M HNO₃ (aq) with 25.0 mL of 0.100 M NH₃ (aq). Label the following points:

• The buffer region

• The equivalence point

NH₃​ (aq) + HNO₃ ​(aq) → NH₄⁺ ​(aq) + NO₃⁻ ​(aq)

The reaction between ammonia and nitric acid is represented by the equation above.

Using the equation to explain why the pH at the equivalence point is less than 7.

The student must select an appropriate indicator for this titration. The table shows the pK_a values and pH transition ranges for three indicators.

Based on the data, identify the most suitable indicator for this titration and justify your choice.

A 50.0 mL sample of 0.150 M benzoic acid (C₆H₅COOH) is titrated with 0.100 M NaOH. The pK_a of C₆H₅COOH is 4.20.

Write the balanced chemical equation for the reaction that occurs during the titration.

Calculate the volume of NaOH required to reach equivalence point.

Calculate the pH at equivalence point.

Assume that only benzoate ions (C₆H₅COO⁻) are present at equivalence, and the K_b of C₆H₅COO⁻ can be determined using the relationship K_w = K_a × K_b.