Electrochemical Cells (College Board AP® Chemistry): Exam Questions

Metal-air cells are a relatively new type of portable energy source consisting of a metal anode, an alkaline electrolyte paste that contains water, and a porous cathode membrane that lets in oxygen from the air. A schematic of the cell is shown below.

Reduction potentials for the cathode and three possible metal anodes are given in the table below.

Early forms of metal-air cells used zinc as the Zinc oxide is produced as the cell operates according to the overall equation below.

2Zn (s) + O₂ (g) → 2ZnO (s)

 i) Using the data in the table above, calculate the cell potential for the zinc-air cell.

 ii) The electrolyte paste contains OH⁻. On the diagram of the cell above, draw an arrow to indicate the direction of migration of OH⁻ ions through the electrolyte as the cell operates.

A fresh zinc-air cell is weighed on an analytical balance before being placed in a hearing aid for

i) As the cell operates, does the mass of the cell increase, decrease, or remain the same?

ii) Justify your answer to part (b)(i) in terms of the equation for the overall cell

The zinc-air cell is taken to the top of a mountain where the air pressure is lower

i) Will the cell potential be higher, lower, or the same as the cell potential at the lower elevation?

ii) Justify your answer to part (c)(i) based on the equation for the overall cell reaction and the information.

Metal-air cells need to be lightweight for many In order to transfer more electrons with a smaller mass, Na and Ca are investigated as potential anodes.

A 1.0 g anode of which of these metals would transfer more electrons, assuming that the anode is totally consumed during the lifetime of a cell? Justify your answer with calculations.

The only common oxide of zinc has the formula ZnO.

i) Write the electron configuration for a Zn atom in the ground state.

ii) From which sublevel are electrons removed when a Zn atom in the ground state is oxidized?

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).

A student wants to determine the concentration of H₂O₂ in a solution of H₂O₂ (aq). The student can use one of two titrants, either dichromate ion, Cr₂O₇²⁻ (aq), or cobalt(II) ion,  Co²⁺ (aq). The balanced  chemical equations for the two titration reactions are shown below.

Dichromate as titrant:

Cr₂O₇²⁻ (aq) + 3H₂O₂ (aq) + 8H⁺ (aq) → 2Cr³⁺ (aq) + 3O₂ (g) + 7H₂O (l)

Cobalt(II) as titrant:

2Co²⁺ (aq) + H₂O₂ (aq) + 2H⁺ (aq) → 2Co³⁺ (aq) + 2H₂O (l)

The half-reactions and the E° values for the systems related to the titrations above are given in the following table.

Use the information in the table to calculate the following.

i) E° for the reaction between Cr₂O₇²⁻(aq) and H₂O₂ (aq) at 298 K

ii) E° for the reaction between Co²⁺ (aq) and H₂O₂ (aq) at 298 K

Based on the calculated values of E° , the student must choose the titrant for which the titration reaction is thermodynamically favorable at 298 K.

i) Which titrant should the student choose? Explain your reasoning.

ii) Calculate the value of ΔG° in kJ/mol_rxn , for the reaction between the chosen titrant and H₂O₂ (aq).

A student sets up a galvanic cell at 298 K that has an electrode of Ag (s) immersed in a 1.0 M solution of Ag⁺ (aq) and an electrode of Cr (s) immersed in a 1.0 M solution of Cr³⁺ (aq), as shown in the diagram above.

The student measures the voltage of the cell shown above and discovers that it is zero. Identify the missing component of the cell, and explain its importance for obtaining a nonzero voltage.

The student adds the missing component to the cell and measures E°_cell  to be +1.54 V. As the cell operates, Ag⁺ ions are reduced. Use this information and the information in the table above to do the following.

i) Calculate the value of E° for the half-reaction Cr³⁺ (aq) + 3e⁻ → Cr (s).

ii) Write the balanced net-ionic equation for the overall reaction that occurs as the cell operates.

iii) Calculate the value of ΔG° for the overall cell reaction in J/mol_rxn.

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.

Answer the following questions relating to the element aluminum, Al.

Write the complete ground-state electron configuration of an Al atom.

Based on principles of atomic structure, explain why the radius of the Al atom is larger than the radius of the Al³⁺ ion.

A student plans to combine solid aluminum with an aqueous solution of silver ions. The student determines the mass of solid AgNO₃ needed to prepare the solution with a specific concentration.

In the following table, briefly list the steps necessary to prepare 200.0 mL of an aqueous solution of AgNO₃ using only equipment selected from the choices given. Assume that all appropriate safety measures are already in place. Not all equipment or lines in the table may be needed.

• Solid AgNO₃

• Weighing paper and scoop

• Distilled water

• 200.00 mL volumetric flask

• Balance

• 50.0 mL graduated cylinder

• 250 mL beakers

• Pipet 

After preparing the solution, the student places some of the solution into a beaker and adds a sample of aluminum. The reaction represented by the following equation occurs.

Al (s) + 3Ag⁺ (aq) → Al³⁺ (aq) + 3Ag (s)   

The following diagram gives an incomplete particulate representation of the reaction. The beaker on the left represents the system before the mixture reacts. Complete the drawing on the right to represent the system after the reaction has occurred. Be sure to include 1) the correct type and number of particles based on the number shown on the left and 2 ) the relative spacing to depict the appropriate phases.

The student finds the standard reduction potentials given in the table, which are related to the reaction that occurs.

 Using the standard reduction potentials, calculate the value of E ° for the reaction.

Based on the value of E °, would the standard free energy change of the reaction under standard conditions, ∆G°, be positive, negative, or zero? Justify your answer.

Once the reaction appears to stop progressing, would the change in free energy, ∆G, be positive, negative, or zero? Justify your answer.

Molten MgCl₂  can be decomposed into its elements if a sufficient voltage is applied using inert electrodes. The products of the reaction are liquid Mg (at the cathode) and Cl₂  gas (at the anode). A simplified representation of the cell is shown above. The reduction half-reactions related to the overall reaction in the cell are given in the table.

Draw an arrow on the diagram to show the direction of electron flow through the external circuit as the cell operates.

Would an applied voltage of 2.0 V be sufficient for the reaction to occur? Support your claim with a calculation as part of your answer.

If the current in the cell is kept at a constant 5.00 amps, calculate how many seconds it takes to produce 2.00 g of Mg (l) at the cathode.

A galvanic cell is constructed using a silver electrode in a 1.00 M AgNO₃ solution and a zinc electrode in a 1.00 M Zn(NO₃)₂ solution. The two half-cells are connected by a salt bridge, and a voltmeter measures the cell potential.

Write the half-equations for the reactions at the anode and the cathode.

The standard reduction potentials for the relevant half-reactions are given in the table below:

Calculate the standard cell potential (E_cell^∘).

Explain the role of the salt bridge and how it maintains electrical neutrality in the cell.

Predict how the cell potential will change if the concentration of Ag⁺ is decreased to 0.010 M while keeping [Zn²⁺] constant at 1.00 M. Justify your answer using the Nernst equation.

A metal object is plated with copper using an electrolytic cell. A 2.00 A current is passed through a solution of CuSO₄ for 45.0 minutes. Write the half-equation for the reaction at the cathode.

Using your answer from part (a), calculate the mass of copper deposited on the object.

The efficiency of the electroplating process is reduced due to side reactions. Write a chemical equation for one possible side reaction and explain how it impacts on the mass of copper deposited.

If the current is doubled to 4.00 A, explain how the time required to deposit the same mass of copper would change. Support your answer with a calculation.