Volatic Cells
Related Examples and Practice Problems
Additional Worked Out Examples/ Practice
Identifying classification types: Differentiation between elements, compounds or mixtures and homogeneous and heterogenous mixtures
Separation techniques: Selected and explaining limitation of appropriate separation
Relating Properties to Composition: Predicting classification based on descriptive properties
Topic Summary & Highlights
and Help Videos
Core Concept
A voltaic cell is an electrochemical cell in which a spontaneous redox reaction generates an electric current.
Key Purpose: Converts chemical energy into electrical energy.
Example: Batteries, such as alkaline or lead-acid batteries, are practical applications of voltaic cells.
Practice Tips
A voltaic cell generates electrical energy from spontaneous redox reactions.
The anode is the site of oxidation, and the cathode is the site of reduction.
The standard cell potential ($E°^{\text{cell}}$) determines if the reaction is spontaneous ($E°^{\text{cell}}$ > 0).
Understanding voltaic cells is essential for studying batteries, corrosion, and energy
Topic Overview Podcast
Topic Related Resources
|
LABORATORY
|
DEMONSTRATIONS
|
ACTIVITIES
|
VIRTUAL SIMULATIONS
|
Components of a Voltaic Cell
Electrodes:
Anode:
Where oxidation occurs (loss of electrons\text{loss of electrons}loss of electrons).
Electrons flow away from the anode.
Cathode:
Where reduction occurs (gain of electrons\text{gain of electrons}gain of electrons).
Electrons flow toward the cathode.
Electrolyte:
Ionic solution that facilitates the flow of ions to balance charges during the reaction.
Salt Bridge:
Contains a salt solution (e.g., KNO3KNO_3KNO3) that allows ion exchange to maintain electrical neutrality in the half-cells.
External Circuit:
Allows the flow of electrons from the anode to the cathode.
Key Concepts
Redox Reactions:
The overall cell reaction is the combination of two half-reactions:
Oxidation: Occurs at the anode.
Reduction: Occurs at the cathode.
Electron Flow:
Electrons flow from the anode (oxidation) to the cathode (reduction) through the external circuit.
Ion Flow:
Anions (negative ions\text{negative ions}negative ions) migrate toward the anode.
Cations (positive ions\text{positive ions}positive ions) migrate toward the cathode.
Cell Potential (Ecell∘E^\circ_{\text{cell}}Ecell∘):
The voltage generated by the cell, calculated using the standard reduction potentials of the half-reactions.
Standard Cell Notation
Voltaic cells are represented using standard cell notation:
$\text{Anode} | \text{Anode Solution} (\text{M}) || \text{Cathode Solution} (\text{M}) | \text{Cathode}Anode∣Anode Solution(M)∣∣Cathode Solution(M)∣Cathode$
Example: \text{Zn (s)} | \text{Zn}^{2+} (1.0 \, \text{M}) || \text{Cu}^{2+} (1.0 \, \text{M}) | \text{Cu (s)}
Calculating Cell Potential
The standard cell potential (Ecell∘E^\circ_{\text{cell}}Ecell∘) is calculated using the standard reduction potentials (E∘E^\circE∘) of the half-reactions:
Ecell∘=Ecathode∘−Eanode∘E^\circ_{\text{cell}} = E^\circ_{\text{cathode}} - E^\circ_{\text{anode}}Ecell∘=Ecathode∘−Eanode∘
If Ecell∘>0E^\circ_{\text{cell}} > 0Ecell∘>0, the reaction is spontaneous.
Example of a Voltaic Cell
Zinc-Copper Cell:
Half-Reactions:
Anode (oxidation): Zn (s)→Zn2+(aq)+2e−\text{Zn (s)} \rightarrow \text{Zn}^{2+} (aq) + 2e^-Zn (s)→Zn2+(aq)+2e−, E∘=−0.76 VE^\circ = -0.76 \, \text{V}E∘=−0.76V
Cathode (reduction): Cu2+(aq)+2e−→Cu (s)\text{Cu}^{2+} (aq) + 2e^- \rightarrow \text{Cu (s)}Cu2+(aq)+2e−→Cu (s), E∘=+0.34 VE^\circ = +0.34 \, \text{V}E∘=+0.34V
Overall Reaction:
Zn (s)+Cu2+(aq)→Zn2+(aq)+Cu (s)\text{Zn (s)} + \text{Cu}^{2+} (aq) \rightarrow \text{Zn}^{2+} (aq) + \text{Cu (s)}Zn (s)+Cu2+(aq)→Zn2+(aq)+Cu (s)
Cell Potential:
Ecell∘=+0.34 V−(−0.76 V)=+1.10 VE^\circ_{\text{cell}} = +0.34 \, \text{V} - (-0.76 \, \text{V}) = +1.10 \, \text{V}Ecell∘=+0.34V−(−0.76V)=+1.10V