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

Two substances, C₂H₆ (ethane) and CH₃OH (methanol), are both gases at room temperature but exhibit very different physical properties.

Based on electronegativity differences, classify the bond between carbon and hydrogen atoms in C₂H₆ as nonpolar covalent, polar covalent, or ionic. Justify your answer.

Identify the strongest type of intermolecular force between molecules of:

i) C₂H₆

ii) CH₃ OH

Draw a Lewis diagram for a single CH₃OH molecule.

Draw a diagram showing an intermolecular interaction between two methanol (CH₃OH) molecules. Clearly indicate the interaction and partial charges on the relevant atoms.

Explain why the boiling point of CH₃OH is significantly higher than that of C₂H₆. Your answer should refer to the differences in intermolecular forces between the two substances.

Methanol is highly soluble in water, whereas ethane is nearly insoluble. Explain this difference in solubility in terms of molecular polarity and interactions with water.

Carbon tetrachloride, CCl₄ (g), can be synthesized according to the reaction represented below.

CS₂ (g) + 3Cl₂ (g)   → CCl₄ (g) + S₂Cl₂ (g)

A chemist runs the reaction at a constant temperature of 120°C in a rigid 25.0 L container.

Chlorine gas, Cl₂ (g), is initially present in the container at a pressure of 0.40 atm.

i) How many moles of Cl₂ (g) are in the container?

ii) How many grams of carbon disulfide, CS₂ (g), are needed to react completely with the Cl₂ (g) ?

At 30°C the reaction is thermodynamically favorable, but no reaction is observed to occur. However, at 120°C, the reaction occurs at an observable rate.

i) Explain how the higher temperature affects the collisions between the reactant molecules so that the reaction occurs at an observable rate at 120°

ii) The graph below shows a distribution for the collision energies of reactant molecules at 120°.

Draw a second curve on the graph that shows the distribution for the collision energies of reactant molecules at 30°C.

S₂Cl₂ is a product of the reaction .

i) In the box below, complete the Lewis electron-dot diagram for the S₂Cl₂  molecule by drawing in all of the electron pairs.

ii) What is the approximate value of the Cl-S-S bond angle in the S₂Cl₂molecule that you drew in part (c)(i) ? (If the two Cl-S-S bond angles are not equal, include both angles.)

CCl₄(g) can also be produced by reacting CHCl₃(g) with Cl₂(g) at 400°C , as represented by the equation below.

CHCl₃(g) + Cl₂(g)   →  CCl₄(g) + HCl(g)

At the completion of the reaction a chemist successfully separates the CCl₄(g) from the HCl(g) by cooling the mixture to 70°C, at which temperature the CCl₄(g) condenses while the HCl(g) remains in the gaseous state.

i) Identify all types of intermolecular forces present in HCl(l) .

ii) What can be inferred about the relative strengths of the intermolecular forces in CCl₄(l)  and  HCl(l) ? Justify your answer in terms of the information above.

A student analyzes a 2.00 g sample of a solid mixture known to contain only NaCl and KCl. The student dissolves the sample in water and adds an excess of AgNO₃ (aq) to precipitate all of the chloride ions as AgCl (s). After filtering and drying, the mass of the AgCl precipitate is 4.78 g.

Write the net ionic equation for the reaction between chloride ions and silver(I) ions in aqueous solution.

i) Using the mass of the AgCl precipitate, calculate the number of moles of Cl⁻ in the sample.

ii) Calculate the mass of Cl⁻ in the sample.

The student determines that the mixture contains 0.0222 mol of NaCl.

i) Calculate the mass of KCl in the sample.

ii) Calculate the number of moles of KCl in the sample.

The student forgets to fully dry the AgCl precipitate before weighing it.

Would this error cause the calculated mass percent of KCl in the mixture to be too high, too low, or unchanged? Justify your answer.

Suppose the mixture also contained a small amount of NaBr.

Explain how the presence of NaBr would affect the calculated mass percent of KCl in the mixture.

The precipitates AgCl (s) and AgBr (s) are both pale in color.

Explain how this could affect the student’s ability to detect the presence of NaBr in the sample.

A student investigates the enthalpy of solution, ΔH_soln, for two alkali metal halides, LiCl and NaCl. In addition to the salts, the student has access to a calorimeter, a balance with a precision of ±0.1 g, and a thermometer with a precision of ±0.1°C.

To measure ΔH_soln  for LiCl, the student adds 100.0 g of water initially at 15.0°C to a calorimeter and adds 10.0 g of LiCl(s) , stirring to dissolve. After the LiCl dissolves completely, the maximum temperature reached by the solution is 35.6°C.

i) Calculate the magnitude of the heat absorbed by the solution during the dissolution process , assuming that the specific heat capacity of the solution is 4.18 J/(g·°C). Include units with your answer.

ii) Determine the value of ΔH_soln for LiCl in kJ /mol_rxn .

To explain why  ΔH_soln  for NaCl  is different than that for LiCl, the student investigates factors that affect ΔH_soln  and finds that ionic radius and lattice enthalpy (which can be defined as the ΔH associated with the separation of a solid crystal into gaseous ions) contribute to the process . The student consults references and collects the data shown in the table below.

Write the complete electron configuration for the Na⁺ ion in the ground state.

Using principles of atomic structure, explain why the Na⁺ ion is larger than the Li⁺ ion .

Which salt, LiCl or NaCl, has the greater lattice enthalpy? Justify your answer.

Below is a representation of a portion of a crystal of LiCl. Identify the ions in the representation by writing the appropriate formulas ( Li⁺ or Cl⁻ ) in the boxes below.

The lattice enthalpy of LiCl is positive, indicating that it takes energy to break the ions apart in LiCl. However, the dissolution of LiCl in water is an exothermic process. Identify all particle-particle interactions that contribute significantly to the dissolution process being exothermic. For each interaction , include the particles that interact and the specific type of intermolecular force between those particles.

At 700 K, carbon monoxide gas reacts reversibly with steam in a closed container according to the following balanced equation:

CO (g) + H₂O (g) ⇌ CO₂ (g) + H₂ (g)

An experiment is conducted where 0.400 mol of each species is placed in a 2.00 L container. The system reaches equilibrium at constant temperature.

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

Determine the initial concentrations (in mol L^-1) of all four species in the container.

At equilibrium, the concentration of H₂O (g) is found to be 0.160 mol L⁻¹. Calculate the equilibrium concentrations of the other three species.

Using your answers from parts (a) and (c), calculate the value of the equilibrium constant, K_c​, for the reaction.

A second reaction mixture is prepared at the same temperature with the following concentrations:

Determine whether the system will shift to the left, right, or remain unchanged. Justify your answer using a calculation.

The reaction vessel is opened briefly, allowing some H₂ gas to escape before being resealed. Explain how this change affects the concentrations of all four species after the system re-establishes equilibrium.

A student is given 50.0 mL of a solution of Na₂CO₃(aq) of unknown concentration. To determine the concentration of the solution, the student mixes the solution with excess 1.0 M Ca(NO₃)₂(aq) , causing a precipitate to form. The balanced equation for the reaction is shown below.

Na₂CO₃ (aq) + Ca(NO₃)₂ (aq) → 2NaNO₃ (aq) + CaCO₃ (s)

Write the net ionic equation for the reaction that occurs when the solutions of Na₂CO₃ and Ca(NO₃)₂ are mixed.

The diagram below is incomplete. Draw in the species needed to accurately represent the major ionic species remaining in the solution after the reaction has been completed.

The student filters and dries the precipitate of CaCO₃ (molar mass 100.1 g/mol) and records the data in the table below.

Determine the number of moles of Na₂CO₃ in the original 50.0 mL of solution.

The student realizes that the precipitate was not completely dried and claims that as a result, the calculated Na₂CO₃ molarity is too low. Do you agree with the student’s claim? Justify your answer.

After the precipitate forms and is filtered, the liquid that passed through the filter is tested to see if it can conduct electricity. What would be observed? Justify your answer.

The student decides to determine the molarity of the same Na₂CO₃ solution using a second method. When Na₂CO₃ is dissolved in water, CO₃²⁻(aq) hydrolyzes to form HCO₃⁻  (aq), as shown by the following equation.

 CO₃²⁻(aq) + H₂O(l)  → HCO₃⁻(aq) + OH⁻ (aq)          

 K_b = = 2.1 × 10⁻⁴

The student decides to first determine [OH⁻] in the solution, then use that result to calculate the initial concentration of CO₃²⁻(aq).

i) Identify a laboratory method (not titration) that the student could use to collect data to determine [OH⁻] in the solution.

ii) Explain how the student could use the measured value in part (f)(i) to calculate the initial concentration of CO₃²⁻(aq). (Do not do any numerical calculations.)

In the original Na₂CO₃ solution at equilibrium, is the concentration of HCO₃⁻(aq) greater than, less than, or equal to the concentration of CO₃²⁻(aq) ? Justify your answer.

The student needs to make a CO₃²⁻ / HCO₃⁻ buffer. Is the Na₂CO₃ solution suitable for making a buffer with a pH of 6 ? Explain why or why not.

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.

To determine the molar mass of an unknown metal, M, a student reacts iodine with an excess of the metal to form the water-soluble compound Ml₂ , as represented by the equation below.

M + I₂  →  MI₂

The reaction proceeds until all of the I₂  is consumed . The Ml₂(aq) solution is quantitatively collected and heated to remove the water , and the product is dried and weighed to constant mass. The experimental steps are represented below , followed by a data table.

Given that the metal M is in excess, calculate the number of moles of I₂ that reacted.

Calculate the molar mass of the unknown metal M .

The student hypothesizes that the compound formed in the synthesis reaction is ionic.

Propose an experimental test the student could perform that could be used to support the hypothesis. Explain how the results of the test would support the hypothesis if the substance was ionic.

The student hypothesizes that Br₂  will react with metal M more vigorously than I₂  did because  Br₂  is a liquid at room temperature.

Explain why I₂  is a solid at room temperature whereas Br₂  is a liquid. Your explanation should clearly reference the types and relative strengths of the intermolecular forces present in each substance.

While cleaning up after the experiment, the student wishes to dispose of the unused solid I₂  in a responsible manner. The student decides to convert the solid  I₂  to I⁻ (aq) anion. The student has access to three solutions, H₂O₂ (aq), Na₂S₂O₃ (aq) ,  and  Na₂S₄O₆ (aq) ,  and the standard reduction table shown below.

Which solution should the student add to l₂ (s) to reduce it to I⁻ (aq)? Circle your answer below. Justify your answer, including a calculation of E°  for the overall reaction.

H₂O₂ (aq) Na₂S₂O₃ (aq) Na₂S₄O₆ (aq) 

Write the balanced net-ionic equation for the reaction between  I₂  and the solution you selected in part (e).