10 Point Questions (College Board AP® Chemistry): Exam Questions

A researcher is studying the properties of silver (Ag), nickel (Ni), and silver bromide (AgBr), which differ in bonding type and electrical behavior. These substances are used in photography, electronics, and electrochemical processes.

Identify the type of solid formed by AgBr and justify your answer in terms of the particles and forces involved.

Predict whether AgBr or NaBr has a higher lattice energy. Justify your answer.

AgBr is used in photographic film because it undergoes decomposition when exposed to light.

Write a balanced equation for the photochemical decomposition of AgBr.

Explain why metallic silver conducts electricity, but solid AgBr does not.

A nickel solution is often used in electroplating.

Write the half-equation for the reduction of Ni²⁺ to Ni(s), and
identify what type of bonding is present in the plated nickel layer.

Predict which substance, Ag or AgBr, is more malleable and justify your answer based on the bonding and structure of the solid.

Silver bromide (AgBr) is a sparingly soluble salt. The K_sp of AgBr is 5.0 × 10⁻¹³ at 25 °C.

Calculate the molar solubility of AgBr in pure water at this temperature. Show all working.

A solution is prepared by mixing AgNO₃ and NaBr, each at an initial concentration of 1.00 × 10^-4 M.

Predict whether a precipitate of AgBr will form.
Justify your answer by calculating the ion product Q and comparing it to K_sp.

Ethene, C₂H₄ (g) (molar mass 28.1 g/mol), may be prepared by the dehydration of ethanol, C₂H₅OH (g) (molar mass 46.1 g/mol), using a solid catalyst. A setup for the lab synthesis is shown in the diagram below.

The equation and relevant data for the dehydration reaction is given below.

C₂H₅OH (g) C₂H₄ (g) + H₂O (g)

H^o₂₉₈ = 45.5 kJ/mol_rxn

S^o₂₉₈ = 126 J/(K.mol_rxn)

A student added a 0.200 g sample of C₂H₅OH (l) to a test tube using the setup shown above. The student heated the test tube gently with a Bunsen burner until all of the C₂H₅OH (l) evaporated and gas generation stopped. When the reaction stopped, the volume of collected gas was 0.0854 L at 0.822 atm and 305 K. (The vapor pressure of water at 305 K is 35.7 torr.)

Calculate the number of moles of C₂H₄ (g)

i) that are actually produced in the experiment and measured in the gas collection tube and

ii) that would be produced if the dehydration reaction went to completion.

Calculate the percent yield of C₂H₄(g) in the experiment.

Because the dehydration reaction is not observed to occur at 298 K, the student claims that the reaction has an equilibrium constant less than 1.00 at 298 K.

Does the thermodynamic data for the reaction support the student’s claim? Justify your answer, including a calculation of ΔG°₂₉₈ for the reaction. 

The Lewis electron-dot diagram for C₂H₄ is shown below in the box on the left. In the box on the right, complete the Lewis electron-dot diagram for C₂H₅OH by drawing in all of the electron pairs.

What is the approximate value of the C–O–H bond angle in the ethanol molecule?

During the dehydration experiment, C₂H₄(g) and unreacted C₂H₅OH (g) passed through the tube into the water. The C₂H₄ was quantitatively collected as a gas, but the unreacted C₂H₅OH was not.

Explain this observation in terms of the intermolecular forces between water and each of the two gases.

In aqueous solution, hydroxylamine (NH₂OH) reacts with hydrogen fluoride (HF):

NH₂OH (aq) + HF (aq) ⇌ NH₃OH⁺ (aq) + F⁻(aq)

Identify the Brønsted–Lowry acid and base in the forward reaction. Justify your answer in terms of proton transfer.

Identify the conjugate acid–base pairs and explain how they are related.

Predict which is the stronger conjugate base. Justify your answer.

CH₃COCH₂COCH₃ + NH₃ → CH₃COCH⁻COCH₃ + NH₄⁺

Identify the Brønsted–Lowry acid and base in the reaction above. Explain your answer.

Phosphine (PH₃) is considered a weaker base than ammonia (NH₃).

In a gas-phase experiment, PH₃ reacts with hydrogen bromide:

PH₃​ (g) + HBr (g) → PH₄⁺ ​(g) + Br⁻ (g)

Explain how this reaction supports the role of PH₃ as a Brønsted–Lowry base, and justify why PH₃ is less basic than NH₃ based on structure and electronegativity.

A student performs an experiment to determine the value of the enthalpy change, ΔH°_rxn, for the oxidation-reduction reaction represented by the balanced equation below.

Na₂S₂O₃(aq) + 4 NaOCl(aq) + 2 NaOH(aq) → 2 Na₂SO₄(aq) + 4 NaCl(aq) + H₂O(l)

Determine the oxidation number of Cl in NaOCl .

Calculate the number of grams of Na₂S₂O₃ needed to prepare 100.00 mL of 0.500 M Na₂S₂O₃(aq).

In the experiment, the student uses the solutions shown in the table below.

 Using the balanced equation for the oxidation-reduction reaction and the information in the table above, determine which reactant is the limiting reactant. Justify your answer.

The solutions, all originally at 20.0°C, are combined in an insulated calorimeter. The temperature of the reaction mixture is monitored, as shown in the graph below.

According to the graph, what is the temperature change of the reaction mixture?

The mass of the reaction mixture inside the calorimeter is 15.21 g.

i) Calculate the magnitude of the heat energy, in joules, that is released during the reaction. Assume that the specific heat of the reaction mixture is 3.94 J/(g·°C) and that the heat absorbed by the calorimeter is negligible.

ii) Using the balanced equation for the oxidation-reduction reaction and your answer to part (c), calculate the value of the enthalpy change of the reaction, ΔH°_rxn, in kJ/mol_rxn. Include the appropriate algebraic sign with your answer.

The student repeats the experiment, but this time doubling the volume of each of the reactants, as shown in the table below.

The magnitude of the enthalpy change, ΔH°_rxn, in kJ/mol_rxn, calculated from the results of the second experiment is the same as the result calculated in part (e)(ii) . Explain this result.

Write the balanced net ionic equation for the given reaction.

Chlorine gas reacts reversibly with iodine monochloride gas in a sealed container at constant pressure:

Cl₂ (g) + ICl (g) ⇌ ICl₃ (g)

This reaction is studied at two different temperatures. The equilibrium concentrations of each species are recorded for each trial in identical 2.00 L containers:

Write the expression for the equilibrium constant, K_c​, for this reaction.

Calculate the value of K_c​ at 400 K.

Calculate the value of K_c​ at 500 K.

Determine whether the forward reaction is endothermic or exothermic. Justify your answer using the values of K_c​.

Each of the following particle diagrams shows a 2.00 L container at equilibrium.

Which diagram best represents the system at 400 K? Justify your answer based on the value of K_c​.

A new mixture is prepared at 400 K with the following concentrations:

[Cl₂] = 0.10 M, [ICl] = 0.10 M, [ICl₃] = 1.20 M

Determine whether the system will shift toward the reactants, the products, or remain unchanged. Justify your answer with a calculation.

Propanoic acid, CH₃CH₂COOH, is a carboxylic acid that reacts with water according to the equation below.

CH₃CH₂COOH(aq) + H₂O(l)   CH₃CH₂COO⁻(aq) + H₃O⁺(aq)

At 25°C the pH of a 50.0 mL sample of 0.20 M CH₃CH₂COOH is 2.79.

Identify a Brønsted-Lowry conjugate acid-base pair in the reaction. Clearly label which is the acid and which is the base.

 Determine the value of K_a for propanoic acid at 25°C. 

For each of the following statements, determine whether the statement is true or false. In each case, explain the reasoning that supports your answer.

i) The pH of a solution prepared by mixing the 50.0 mL sample of 0.20 M CH₃CH₂COOH with a 50.0 mL sample of 0.20 M NaOH is 7.00 .

ii) If the pH of a hydrochloric acid solution is the same as the pH of a propanoic acid solution, then the molar concentration of the hydrochloric acid solution must be less than the molar concentration of the propanoic acid solution.

A student is given the task of determining the concentration of a propanoic acid solution of unknown concentration. A 0.173 M NaOH solution is available to use as the titrant. The student uses a 25.00 mL volumetric pipet to deliver the propanoic acid solution to a clean, dry flask. After adding an appropriate indicator to the flask, the student titrates the solution with the 0.173 M NaOH , reaching the end point after 20.52 mL of the base solution has been added.

Calculate the molarity of the propanoic acid solution.

The student is asked to redesign the experiment to determine the concentration of a butanoic acid solution instead of a propanoic acid solution. For butanoic acid the value of pK_a is 4.83. The student claims that a different indicator will be required to determine the equivalence point of the titration accurately. Based on your response to part (b), do you agree with the student’s claim? Justify your answer.

The compound urea, H₂NCONH₂ , is widely used in chemical fertilizers. The complete Lewis electron-dot diagram for the urea molecule is shown below.

Identify the hybridization of the valence orbitals of the carbon atom in the urea molecule.

Urea has a high solubility in water, due in part to its ability to form hydrogen bonds. A urea molecule and four water molecules are represented in the box below. Draw ONE dashed line ( ---- ) to indicate a possible location of a hydrogen bond between a water molecule and the urea molecule.

The dissolution of urea is represented by the equation above. A student determines that 5.39 grams of H₂NCONH₂ (molar mass 60.06 g/mol) can dissolve in water to make 5.00 mL of a saturated solution at 20°C.

Calculate the concentration of urea, in mol/L, in the saturated solution at 20.°C.

The student also determines that the concentration of urea in a saturated solution at 25 ^oC is 19.8 M. Based on this information, is the dissolution of urea endothermic or exothermic? Justify your answer in terms of Le Chatelier’s principle.

The equipment shown above is provided so that the student can determine the value of the molar heat of solution for urea. Knowing that the specific heat of the solution is 4.18 J/(g⋅°C), list the specific measurements that are required to be made during the experiment.

The entropy change for the dissolution of urea, ΔS°_soln, is 70.1 J/(mol⋅K) at 25°C. Using the information in the table above, calculate the absolute molar entropy, S°, of aqueous urea.

Using particle-level reasoning, explain why ΔS°_soln is positive for the dissolution of urea in water.

The student claims that ΔS° for the process contributes to the thermodynamic favorability of the dissolution of urea at 25°C. Use the thermodynamic information above to support the student’s claim.

A student is given a standard galvanic cell, represented below, that has a Cu electrode and a Sn electrode.

As current flows through the cell, the student determines that the Cu electrode increases in mass and the Sn electrode decreases in mass.

Identify the electrode at which oxidation is occurring. Explain your reasoning based on the student’s observations.

As the mass of the Sn electrode decreases, where does the mass go?

In the expanded view of the center portion of the salt bridge shown in the diagram below, draw and label a particle view of what occurs in the salt bridge as the cell begins to operate. Omit solvent molecules and use arrows to show the movement of particles.

A nonstandard cell is made by replacing the 1.0 M solutions of Cu(NO₃)₂ and Sn(NO₃)₂ in the standard cell with 0.50 M solutions of Cu(NO₃)₂ and Sn(NO₃)₂ . The volumes of solutions in the nonstandard cell are identical to those in the standard cell.

i) Is the cell potential of the nonstandard cell greater than, less than, or equal to the cell potential of the standard cell? Justify your answer.

 ii) Both the standard and nonstandard cells can be used to power an electronic device. Would the nonstandard cell power the device for the same time, a longer time, or a shorter time as compared with the standard cell? Justify your answer.

In another experiment, the student places a new Sn electrode into a fresh solution of 1.0 M Cu(NO₃)₂ .

 i) Using information from the table above, write a net-ionic equation for the reaction between the Sn electrode and the Cu(NO₃)₂ solution that would be thermodynamically favorable. Justify that the reaction is thermodynamically favorable.

 ii) Calculate the value of ΔG° for the reaction. Include units with your answer.