AP Chemistry 9.8 Galvanic (Voltaic) Cells Study Notes - New Syllabus Effective fall 2024
AP Chemistry 9.8 Galvanic (Voltaic) Cells Study Notes- New syllabus
AP Chemistry 9.8 Galvanic (Voltaic) Cells Study Notes – AP Chemistry – per latest AP Chemistry Syllabus.
LEARNING OBJECTIVE
Explain the relationship between the physical components of an electrochemical cell and the overall operational principles of the cell.
Key Concepts:
- Components and Functioning of an Electrochemical Cell
Galvanic vs. Electrolytic Cells
- Oxidation and Reduction in Electrochemical Cells
Components and Functioning of an Electrochemical Cell
An electrochemical cell is a system that converts chemical energy into electrical energy (or vice versa) through oxidation–reduction (redox) reactions. Each part of the cell plays a specific and essential role in ensuring electron flow, ion balance, and completion of the redox process.
Electrochemical cells operate under two main types:
- Galvanic (Voltaic) Cells: spontaneous chemical reactions that generate electrical energy.
- Electrolytic Cells: nonspontaneous reactions driven by an external energy source.
Key Components of an Electrochemical Cell
| Component | Description / Function | Macroscopic Role |
|---|---|---|
| Anode | Electrode where oxidation occurs (loss of electrons). | Electrons are released into the external circuit. |
| Cathode | Electrode where reduction occurs (gain of electrons). | Electrons are accepted from the circuit. |
| Half-cells | Each half-cell contains an electrode and an electrolyte solution of its ions. | Maintains ionic environment for oxidation or reduction. |
| Salt bridge | Connects two half-cells, allowing ion flow to balance charge buildup. | Completes the circuit electrically while preventing direct mixing. |
| External circuit | Wire or conductor that carries electrons between electrodes. | Allows current flow from anode → cathode. |
| Voltmeter / Ammeter | Measures cell potential (voltage) or current. | Indicates electrical output or input of the cell. |
Macroscopic vs. Particulate Descriptions
- Macroscopic level: Observes measurable changes such as voltage, current, electrode mass, and gas evolution.
- Particulate level: Involves electron transfer between ions and electrodes, ion migration through the salt bridge, and redox processes at surfaces.
General Representation of a Galvanic Cell
- Left side (anode): oxidation → \( \mathrm{Zn(s) \rightarrow Zn^{2+}(aq) + 2e^-} \)
- Right side (cathode): reduction → \( \mathrm{Cu^{2+}(aq) + 2e^- \rightarrow Cu(s)} \)
- Electrons flow: from Zn (anode) → Cu (cathode) through external wire.
- Cations flow from salt bridge to cathode; anions to anode.
Example:
In a galvanic cell made from Zn(s) and Cu(s) electrodes, identify what happens at each electrode and explain the role of the salt bridge.
▶️ Answer / Explanation
Step 1: Determine the oxidation and reduction half-reactions:
- Anode: \( \mathrm{Zn(s) \rightarrow Zn^{2+}(aq) + 2e^-} \) (oxidation)
- Cathode: \( \mathrm{Cu^{2+}(aq) + 2e^- \rightarrow Cu(s)} \) (reduction)
Step 2: Direction of electron flow:
Electrons flow from Zn (anode) → Cu (cathode).
Step 3: Role of the salt bridge:
Allows ions to move between half-cells to maintain electrical neutrality:
- Anions (e.g., Cl⁻) move toward the anode compartment to balance \( \mathrm{Zn^{2+}} \) formation.
- Cations (e.g., Na⁺) move toward the cathode compartment to replace \( \mathrm{Cu^{2+}} \) ions reduced at the electrode.
Step 4: Observations:
- Zn electrode loses mass as it dissolves into solution.
- Cu electrode gains mass as solid Cu is deposited.
Final Answer: The salt bridge maintains charge balance, while the anode undergoes oxidation and the cathode reduction. The Zn–Cu cell generates a spontaneous voltage of approximately 1.10 V.
Galvanic vs. Electrolytic Cells
Electrochemical cells are classified into two main types based on whether the overall redox reaction is thermodynamically favored or requires an external source of energy to proceed:
- Galvanic (Voltaic) Cell: uses a spontaneous redox reaction (\( \mathrm{\Delta G^\circ < 0} \)) to produce electrical energy.
- Electrolytic Cell: uses electrical energy from an external source to drive a nonspontaneous reaction (\( \mathrm{\Delta G^\circ > 0} \)).
Key Equation for Spontaneity
\( \mathrm{\Delta G^\circ = -nFE^\circ_{cell}} \)
- \( \mathrm{n} \): moles of electrons transferred
- \( \mathrm{F = 96,485\ C/mol\ e^-} \): Faraday’s constant
- \( \mathrm{E^\circ_{cell}} \): standard cell potential (V)
When \( \mathrm{E^\circ_{cell} > 0} \), the reaction is spontaneous (galvanic). When \( \mathrm{E^\circ_{cell} < 0} \), the reaction is nonspontaneous (electrolytic).
Comparison Between Galvanic and Electrolytic Cells
| Feature | Galvanic (Voltaic) Cell | Electrolytic Cell |
|---|---|---|
| Nature of Reaction | Spontaneous (\( \mathrm{\Delta G^\circ < 0} \)) | Nonspontaneous (\( \mathrm{\Delta G^\circ > 0} \)) |
| Energy Conversion | Chemical → Electrical | Electrical → Chemical |
| Source of Energy | Generated by reaction itself | Supplied from an external power source |
| Anode Reaction | Oxidation | Oxidation |
| Cathode Reaction | Reduction | Reduction |
| Electron Flow | Anode → Cathode (spontaneous) | Forced Anode → Cathode (by power source) |
| Sign of Electrodes | Anode (−), Cathode (+) | Anode (+), Cathode (−) |
| Examples | Daniell cell, batteries | Electrolysis of water, electroplating |
Visual Representation of Flow and Reactions
- In a galvanic cell: electrons flow spontaneously from anode → cathode through a wire; ions move through a salt bridge to maintain neutrality.
- In an electrolytic cell: an external voltage source drives electrons in the reverse direction (nonspontaneous redox).
Example :
Write the half-reactions, direction of electron flow, and electrode signs for a Daniell cell composed of Zn and Cu electrodes.
▶️ Answer / Explanation
Step 1: Half-reactions:
- Anode (oxidation): \( \mathrm{Zn(s) \rightarrow Zn^{2+}(aq) + 2e^-} \)
- Cathode (reduction): \( \mathrm{Cu^{2+}(aq) + 2e^- \rightarrow Cu(s)} \)
Step 2: Electron flow:
Electrons flow from Zn (anode) → Cu (cathode).
Step 3: Electrode signs:
- Anode = negative (source of electrons)
- Cathode = positive (receives electrons)
Step 4: Cell representation:
\( \mathrm{Zn(s) | Zn^{2+}(aq) || Cu^{2+}(aq) | Cu(s)} \)
Final Answer: The Daniell cell operates spontaneously with \( \mathrm{E^\circ_{cell} = +1.10\ V} \); oxidation at the anode and reduction at the cathode generate electrical current.
Example :
Describe what happens at each electrode during the electrolysis of molten \( \mathrm{NaCl(l)} \), including electron flow and products formed.
▶️ Answer / Explanation
Step 1: External DC power source forces nonspontaneous redox:
- Cathode (reduction): \( \mathrm{Na^+ + e^- \rightarrow Na(l)} \)
- Anode (oxidation): \( \mathrm{2Cl^- \rightarrow Cl_2(g) + 2e^-} \)
Step 2: Direction of flow:
Electrons move from the anode (+) to the cathode (−) through the external circuit (opposite of galvanic flow).
Step 3: Observation:
- Liquid sodium forms at the cathode.
- Chlorine gas evolves at the anode.
Final Answer: Electrolysis of molten NaCl is an electrolytic process driven by external electrical energy, resulting in the decomposition of NaCl into \( \mathrm{Na(l)} \) and \( \mathrm{Cl_2(g)} \).
Oxidation and Reduction in Electrochemical Cells
In all electrochemical cells — whether galvanic (voltaic) or electrolytic — the fundamental rule remains the same:
Oxidation occurs at the anode, and reduction occurs at the cathode.
This principle applies universally, even though the signs of the electrodes differ between galvanic and electrolytic cells.
Key Terms and Equations
- Oxidation: loss of electrons → occurs at the anode
- Reduction: gain of electrons → occurs at the cathode
- Mnemonic: “OIL RIG” — Oxidation Is Loss, Reduction Is Gain (of electrons)
General representation for a redox reaction:
\( \mathrm{Ox + e^- \leftrightharpoons Red} \)
Electrode Behavior in Different Cell Types
| Property | Galvanic (Voltaic) Cell | Electrolytic Cell |
|---|---|---|
| Reaction Type | Spontaneous (\( \mathrm{\Delta G^\circ < 0} \)) | Nonspontaneous (\( \mathrm{\Delta G^\circ > 0} \)) |
| Anode process | Oxidation | Oxidation |
| Cathode process | Reduction | Reduction |
| Electron flow | From anode → cathode (spontaneous) | From anode → cathode (forced by external source) |
| Anode sign | Negative (source of electrons) | Positive (connected to positive terminal of power supply) |
| Cathode sign | Positive (electron sink) | Negative (connected to negative terminal) |
Mnemonic Summary
- Oxidation → Anode (Always!)
- Reduction → Cathode (Always!)
- In Galvanic Cell: Anode (−), Cathode (+)
- In Electrolytic Cell: Anode (+), Cathode (−)
Example :
In the galvanic cell \( \mathrm{Zn(s) | Zn^{2+}(aq) || Cu^{2+}(aq) | Cu(s)} \), determine where oxidation and reduction occur, and identify the anode and cathode.
▶️ Answer / Explanation
Step 1: Write half-reactions:
- \( \mathrm{Zn(s) \rightarrow Zn^{2+}(aq) + 2e^-} \) — oxidation
- \( \mathrm{Cu^{2+}(aq) + 2e^- \rightarrow Cu(s)} \) — reduction
Step 2: Identify electrodes:
- Anode → Zn(s) electrode (oxidation site)
- Cathode → Cu(s) electrode (reduction site)
Step 3: Determine flow:
Electrons flow from Zn (anode, −) → Cu (cathode, +)
Final Answer: Oxidation occurs at the Zn anode; reduction occurs at the Cu cathode. This is a galvanic cell that generates electricity spontaneously (\( \mathrm{E^\circ_{cell} = +1.10\ V} \)).
Example :
During the electrolysis of molten \( \mathrm{NaCl(l)} \), identify the anode, cathode, and reactions at each electrode.
▶️ Answer / Explanation
Step 1: Write half-reactions:
- Anode (oxidation): \( \mathrm{2Cl^-(l) \rightarrow Cl_2(g) + 2e^-} \)
- Cathode (reduction): \( \mathrm{Na^+(l) + e^- \rightarrow Na(l)} \)
Step 2: Assign electrodes:
- Anode = positive (connected to external power supply’s + terminal)
- Cathode = negative (connected to − terminal)
Step 3: Direction of flow:
Electrons move from anode → cathode through the external circuit (forced by power source).
Final Answer: In the electrolysis of molten NaCl, oxidation occurs at the anode (+) and reduction at the cathode (−). The universal rule — oxidation at anode, reduction at cathode — holds true.