Question

A common shorthand way to represent a voltaic cell is anode | anode solution || cathode solution |cathode A double vertical line represents a salt bridge or a porous barrier. A single vertical line represents a change in phase, such AS from solid to solution. (a) Write the half-reactions and overall cell reaction represented by Fe \(\left|\mathrm{Fe}^{2+}\right|\left|\mathrm{Ag}^{+}\right| \mathrm{Ag}_{\mathrm{g}}\) sketch the cell. (b) Write the half-reactions and overall cell reaction represented by \(\mathrm{Zn}\left|\mathrm{Zn}^{2+}\right|\left|\mathrm{H}^{+}\right| \mathrm{H}_{3}\) s sketch the cell. (c) Using the notation just described, represent a cell based on the following reaction: $$ \begin{aligned} \mathrm{ClO}_{3}^{-}(a q)+3 \mathrm{Cu}(s)+6 \mathrm{H}^{+}(a q) & \longrightarrow \mathrm{Cl}^{-}(a q)+3 \mathrm{Cu}^{2+}(a q)+3 \mathrm{H}_{2} \mathrm{O}(l) \end{aligned} $$ Pt is used as an inert electrode in contact with the \(\mathrm{ClO}_{3}^{-}\)and Cl. Sketch the cell.

Short answer

The shorthand notations for the voltaic cells are as follows: a) \(Fe | Fe^{2+} || Ag^+ | Ag\) b) \(Zn | Zn^{2+} || H^{+} | H_{2}\) c) \(Pt | ClO_3^-, Cl^-, H^+ || Cu^{2+} | Cu\)

Step-by-step solution

1. Identify Anode and Cathode

In this case, Fe is the anode (where oxidation occurs) and Ag is the cathode (where reduction occurs).

2. Write Half-Reactions

Anode half-reaction (oxidation): \(Fe \rightarrow Fe^{2+} + 2 e^-\) Cathode half-reaction (reduction): \(Ag^{+} + e^- \rightarrow Ag\)

3. Write Overall Cell Reaction

To get the overall cell reaction, balance the electrons and add the half-reactions: \(Fe + 2Ag^+ \rightarrow Fe^{2+} + 2Ag\)

4. Represent Voltaic Cell Using Shorthand Notation

The shorthand notation for this cell is: \[Fe | Fe^{2+} || Ag^+ | Ag\] #Part B: Zn/Zn²⁺ and H⁺/H₂ System#

1. Identify Anode and Cathode

In this case, Zn is the anode (oxidation) and H⁺/H₂ is the cathode (reduction).

2. Write Half-Reactions

Anode half-reaction (oxidation): \(Zn \rightarrow Zn^{2+} + 2 e^-\) Cathode half-reaction (reduction): \(2H^{+} + 2 e^- \rightarrow H_{2}\)

3. Write Overall Cell Reaction

Balance the electrons and add the half-reactions: \(Zn + 2H^{+} \rightarrow Zn^{2+} + H_{2}\)

4. Represent Voltaic Cell Using Shorthand Notation

The shorthand notation for this cell is: \[Zn | Zn^{2+} || H^{+} | H_{2}\] #Part C: ClO₃⁻/Cl⁻ and Cu/Cu²⁺ System#

1. Identify Anode and Cathode

In this case, ClO₃⁻ is the anode (oxidation) and Cu is the cathode (reduction).

2. Write Half-Reactions

Anode half-reaction (oxidation): \(2ClO_3^- + 12 H^+ + 10 e^- \rightarrow 2Cl^- + 6 H_2O\) Cathode half-reaction (reduction): \(3 Cu^{2+} + 6 e^- \rightarrow 3 Cu\)

3. Write Overall Cell Reaction

Balance the electrons and add the half-reactions: \(2 ClO_3^- + 3 Cu + 12 H^+ \rightarrow 2Cl^- + 3 Cu^{2+} + 6 H_2O\)

4. Represent Voltaic Cell Using Shorthand Notation

The shorthand notation for this cell is: \[Pt | ClO_3^-, Cl^-, H^+ || Cu^{2+} | Cu\] The sketch of the cell should show the two compartments with the platinum electrode in contact with the ClO₃⁻ and Cl⁻ in the anode half-cell and the copper electrode in contact with the Cu²⁺ in the cathode half-cell. A salt bridge or porous barrier should be shown connecting the two compartments.

Key concepts

Half-reactions

In electrochemistry, half-reactions are crucial for understanding how voltaic cells work. A half-reaction is a representation of either the oxidation or reduction process separately in a redox reaction. Each half-reaction shows how electrons are either gained or lost. In the voltaic cell, the sum of these half-reactions gives the overall cell reaction.
The anode is where oxidation occurs, and electrons are released. For example, in the provided exercise, the half-reaction for iron is:

• Anode reaction: \( Fe \rightarrow Fe^{2+} + 2e^- \)

Here, iron (Fe) loses electrons to form iron ions (\( Fe^{2+} \)).
The cathode is where reduction occurs, and electrons are gained. It involves a half-reaction such as:

• Cathode reaction: \( Ag^{+} + e^- \rightarrow Ag \)

This indicates that silver ions (\( Ag^{+} \)) gain electrons to form metallic silver (Ag).
By balancing the electrons lost and gained in these half-reactions, we can derive the overall reaction occurring in the cell.

Electrochemistry

Electrochemistry is the branch of chemistry that deals with the interplay between electrical energy and chemical reactions. Voltaic, or galvanic, cells are devices that convert chemical energy into electrical energy through redox reactions. These cells consist of two half-cells connected by a salt bridge or a porous partition.
A salt bridge serves to maintain electrical neutrality within the internal circuit, preventing the solutions in each half-cell from becoming excessively charged. Each half-cell is prepared with specific electrodes immersed in electrolyte solutions. In a voltaic cell:

• The anode is the electrode where oxidation occurs, and it has a negative charge, as it releases electrons.
• The cathode is the electrode where reduction occurs, and it has a positive charge, as it gains electrons.

The potential difference between the electrodes, known as the cell potential or electromotive force (EMF), drives the flow of electrons from the anode to the cathode through an external circuit, producing electricity. Understanding these basic principles of electrochemistry helps explain how devices like batteries and fuel cells work.

Redox reactions

Redox, or reduction-oxidation reactions, are chemical processes in which electrons are transferred between two substances. These reactions occur in each half-cell of a voltaic cell, allowing it to generate electrical energy.
In a redox reaction:

• Oxidation is defined as the loss of electrons by a substance. For instance, in the half-reaction: \( Zn \rightarrow Zn^{2+} + 2e^- \), zinc loses electrons, undergoing oxidation.
• Reduction involves the gain of electrons by a substance. For example, in the half-reaction: \( 2H^{+} + 2e^- \rightarrow H_2 \), hydrogen ions gain electrons, resulting in reduction.

Combining the oxidation and reduction reactions allows us to write the net ionic equation for the electrochemical cell. This demonstrates how the flow of electrons between the anode and cathode generates electrical energy.
Understanding the redox chemistry in voltaic cells is essential for designing and optimizing electrochemical devices used in a wide range of technological applications, from energy storage to chemical synthesis.