Electrochemistry and Redox Reactions Study Guide

Q1. When properly balanced using the smallest whole-number coefficients, what is the coefficient of S in the redox equation: H2S + HNO3 → S + NO (acidic solution)?

Background

Topic: Redox Reactions and Balancing Equations

This question tests your ability to balance a redox equation in acidic solution, specifically determining the coefficient for sulfur (S).

Key Terms and Formulas:

• Redox reaction: A chemical reaction involving the transfer of electrons.

• Oxidation: Loss of electrons.

• Reduction: Gain of electrons.

• Balancing in acidic solution: Add H+ and H2O as needed.

Step-by-Step Guidance

1. Identify the oxidation and reduction half-reactions for H2S and HNO3.

2. Write the half-reactions and balance atoms other than O and H.

3. Balance O by adding H2O, and H by adding H+.

4. Balance the electrons in each half-reaction so they cancel when added together.

5. Combine the half-reactions and check the coefficients, especially for S.

Try solving on your own before revealing the answer!

Final Answer: 3

When balanced, the coefficient of S is 3. This comes from balancing the electrons and atoms in the redox process.

Q2. What is the balanced overall (net) cell reaction for the electrochemical cell: Pt(s) | H2(g) | H+(aq) || Ag+(aq) | Ag(s)?

Background

Topic: Electrochemical Cells and Cell Notation

This question tests your understanding of cell notation and how to write the net cell reaction from the given cell diagram.

Key Terms and Formulas:

• Cell notation: Describes the arrangement of electrodes and solutions.

• Anode: Site of oxidation.

• Cathode: Site of reduction.

• Net cell reaction: Sum of the oxidation and reduction half-reactions.

Step-by-Step Guidance

1. Identify which species is oxidized and which is reduced based on the cell notation.

2. Write the half-reactions for H2/H+ and Ag+/Ag.

3. Balance the electrons transferred in each half-reaction.

4. Add the half-reactions to get the net cell reaction, making sure all species are balanced.

Try solving on your own before revealing the answer!

Final Answer: H2(g) + 2Ag+(aq) → 2H+(aq) + 2Ag(s)

The net cell reaction shows hydrogen gas being oxidized and silver ions being reduced.

Q3. For an electrochemical cell with Fe electrode in 1.0 M FeCl2 and Sn electrode in 1.0 M Sn(NO3)2, which choice includes only true statements when the cell runs spontaneously?

Background

Topic: Electrochemical Cells, Anode/Cathode Identification

This question tests your ability to determine which electrode is the anode/cathode and whether it gains or loses mass during cell operation.

Key Terms and Formulas:

• Anode: Electrode where oxidation occurs (loses mass).

• Cathode: Electrode where reduction occurs (gains mass).

• Spontaneous cell: Electrons flow from anode to cathode.

Step-by-Step Guidance

1. Determine which metal is oxidized and which is reduced using standard reduction potentials.

2. Identify the anode (site of oxidation) and cathode (site of reduction).

3. Recall that the anode loses mass and the cathode gains mass.

4. Match these facts to the answer choices to find the one with only true statements.

Try solving on your own before revealing the answer!

Final Answer: The tin electrode gains mass and the tin electrode is the cathode.

Reduction occurs at the cathode, so tin gains mass as Sn2+ is reduced to Sn(s).

Q4. For the cell reaction Zn + HgO → ZnO + Hg, which is the half-reaction occurring at the anode?

Background

Topic: Electrochemical Cells, Anode Reaction Identification

This question tests your ability to identify the oxidation half-reaction at the anode.

Key Terms and Formulas:

• Anode: Site of oxidation.

• Half-reaction: Shows electron transfer for one species.

Step-by-Step Guidance

1. Write the overall cell reaction and identify which species is oxidized.

2. Write the oxidation half-reaction for the anode.

3. Balance the electrons and charges in the half-reaction.

Try solving on your own before revealing the answer!

Final Answer: Zn → Zn2+ + 2e-

Zinc is oxidized at the anode, releasing electrons.

Q5. Calculate E°cell for the reaction: 2Fe2+(aq) + Cd2+(aq) → 2Fe3+(aq) + Cd(s)

Background

Topic: Standard Cell Potentials

This question tests your ability to calculate the standard cell potential (E°cell) using standard reduction potentials.

Key Terms and Formulas:

• Standard reduction potential ($E^$): The voltage for a reduction half-reaction under standard conditions.

• Cell potential formula: $E^0_{cell} = E^0_{cathode} - E^0_{anode}$

Step-by-Step Guidance

1. Identify the cathode and anode reactions from the overall equation.

2. Look up the standard reduction potentials for Fe3+/Fe2+ and Cd2+/Cd.

3. Assign the correct values to cathode and anode.

4. Use the formula $E^0_{cell} = E^0_{cathode} - E^0_{anode}$ to set up the calculation.

Try solving on your own before revealing the answer!

Final Answer: -1.17 V

The negative value indicates the reaction is not spontaneous under standard conditions.

Q6. According to the cell diagram Sn | Sn2+ || NO3- (acid soln), NO(g) | Pt, which chemical species undergoes reduction?

Background

Topic: Cell Diagrams and Redox Identification

This question tests your ability to interpret cell diagrams and identify which species is reduced.

Key Terms and Formulas:

• Reduction: Gain of electrons.

• Cell diagram: Left side is anode (oxidation), right side is cathode (reduction).

Step-by-Step Guidance

1. Identify the anode and cathode sides of the cell diagram.

2. Determine which species is present at the cathode.

3. Recall that reduction occurs at the cathode.

Try solving on your own before revealing the answer!

Final Answer: NO3-

Nitrate ion is reduced at the cathode in this cell setup.

Q7. Which one of the following reactions will occur spontaneously at standard-state conditions and 25°C?

Background

Topic: Spontaneity of Redox Reactions

This question tests your ability to predict spontaneity using standard reduction potentials.

Key Terms and Formulas:

• Spontaneous reaction: Positive cell potential ($E^0_{cell} > 0$).

• Standard reduction potential table: Used to compare oxidizing/reducing strength.

Step-by-Step Guidance

1. Write the half-reactions for each option and look up their standard reduction potentials.

2. Calculate $E^0_{cell}$ for each reaction.

3. Identify which reaction has a positive cell potential.

Try solving on your own before revealing the answer!

Final Answer: Mg2+ + Ca → Mg + Ca2+

This reaction is spontaneous because calcium is a stronger reducing agent than magnesium.

Q8. Using a table of standard electrode potentials, which statement is completely true?

Background

Topic: Standard Electrode Potentials and Redox Predictions

This question tests your ability to use standard electrode potentials to predict which species can oxidize or reduce others.

Key Terms and Formulas:

• Oxidizing agent: Causes oxidation, is reduced.

• Reducing agent: Causes reduction, is oxidized.

• Standard reduction potential table: Used to compare strengths.

Step-by-Step Guidance

1. For each statement, check the standard reduction potentials of the species involved.

2. Determine if the reactions described are spontaneous based on the potentials.

3. Eliminate statements with any false claims.

Try solving on your own before revealing the answer!

Final Answer: H+ can oxidize Fe, and Ni can reduce Br2.

Both parts of this statement are true according to standard electrode potentials.

Q9. For the electrochemical cell Ni(s) | Ni2+(1 M) || H+(1 M) | H2(1 atm) | Pt(s), which change will cause a decrease in cell voltage?

Background

Topic: Factors Affecting Cell Voltage

This question tests your understanding of how changes in concentration, pressure, and electrode mass affect cell voltage.

Key Terms and Formulas:

• Nernst equation: $E_{cell} = E^0_{cell} - \frac{0.0592}{n} \log Q$

• Q: Reaction quotient, depends on concentrations and pressures.

Step-by-Step Guidance

1. Analyze how each change affects the reaction quotient Q.

2. Recall that decreasing the concentration of a reactant or increasing product concentration lowers cell voltage.

3. Use the Nernst equation to predict the effect of each change.

Try solving on your own before revealing the answer!

Final Answer: Decrease the concentration of Ni2+ ion.

Lowering Ni2+ concentration increases Q, decreasing cell voltage.

Q10. Magnesium reacts with dilute hydrochloric acid according to which equation?

Background

Topic: Metal-Acid Reactions

This question tests your ability to write the correct balanced equation for magnesium reacting with hydrochloric acid.

Key Terms and Formulas:

• Single displacement reaction: Metal replaces hydrogen in acid.

• Balanced equation: Reactants and products with correct stoichiometry.

Step-by-Step Guidance

1. Write the reactants: Mg and HCl.

2. Predict the products: MgCl2 and H2 gas.

3. Balance the equation for atoms and charges.

Try solving on your own before revealing the answer!

Final Answer: Mg + 2HCl → H2 + MgCl2

Magnesium displaces hydrogen, forming magnesium chloride and hydrogen gas.

Q11. How many coulombs (C) of electrical charge must pass through an electrolytic cell to reduce 0.44 mol Ca2+ ion to calcium metal?

Background

Topic: Electrolysis and Faraday's Law

This question tests your ability to calculate the total charge needed for a given amount of reduction using Faraday's law.

Key Terms and Formulas:

• Faraday's law: $Q = n 00 F$

• $n$: Number of moles of electrons transferred.

• $F$: Faraday constant ($96,485$ C/mol e-).

Step-by-Step Guidance

1. Write the reduction half-reaction: $Ca^{2+} + 2e^- ightarrow Ca$.

2. Calculate the total moles of electrons needed: $0.44$ mol Ca2+ × $2$ mol e- per mol Ca2+.

3. Multiply the moles of electrons by the Faraday constant to get total charge.

Try solving on your own before revealing the answer!

Final Answer: 85,000 C

Two moles of electrons per mole of Ca2+ are required, so multiply accordingly.

Q12. How long will it take to produce 78.0 g of Al metal by the reduction of Al3+ in an electrolytic cell with a current of 2.00 A?

Background

Topic: Electrolysis, Charge, and Time Calculations

This question tests your ability to relate mass, charge, and current in an electrolytic cell.

Key Terms and Formulas:

• Faraday's law: $Q = n 00 F$

• Current: $I = \frac{Q}{t}$

• Moles of Al: $n = \frac{mass}{M_{Al}}$

• Reduction half-reaction: $Al^{3+} + 3e^- ightarrow Al$

Step-by-Step Guidance

1. Calculate moles of Al produced from the given mass and molar mass.

2. Determine total moles of electrons needed (3 per Al atom).

3. Calculate total charge required using Faraday's law.

4. Use $I = \frac{Q}{t}$ to solve for time $t$.

Try solving on your own before revealing the answer!

Final Answer: 116 h

Calculate the charge needed, then divide by the current to find the time.

Q13. The reduction of 1.00 mole of Cr3+ to Cr requires 9.65 × 104 C of electrical charge. True or False?

Background

Topic: Faraday's Law and Stoichiometry

This question tests your understanding of the relationship between moles of electrons and charge required for reduction.

Key Terms and Formulas:

• Reduction half-reaction: $Cr^{3+} + 3e^- ightarrow Cr$

• Faraday constant: $F = 96,485$ C/mol e-

• Total charge: $Q = n 00 F$

Step-by-Step Guidance

1. Determine the number of moles of electrons needed for 1 mole of Cr3+ reduction.

2. Multiply the moles of electrons by the Faraday constant to find the total charge.

Try solving on your own before revealing the answer!

Final Answer: False

Three moles of electrons are needed, so the charge is three times the Faraday constant.