Galvanic Cells — Revision Notes

⚡ 30-Second Revision

Galvanic Cell: — Converts chemical energy (spontaneous redox) to electrical energy.
Anode: — Oxidation, negative electrode, electrons flow from here.
Cathode: — Reduction, positive electrode, electrons flow to here.
Salt Bridge: — Maintains electrical neutrality by ion migration (not electron flow).
Electron Flow: — Anode $\rightarrow$ External Circuit $\rightarrow$ Cathode.
Standard Cell Potential: — $E^\circ_{cell} = E^\circ_{cathode} - E^\circ_{anode}$ (both reduction potentials).
Nernst Equation (at $298\,\text{K}$): — $E_{cell} = E^\circ_{cell} - \frac{0.0592}{n} \log Q$.
Reaction Quotient ($Q$): — Products/Reactants (exclude solids/liquids).
Gibbs Free Energy: — $\Delta G = -nFE_{cell}$.
Spontaneity: — $E_{cell} > 0$, $\Delta G < 0$, $K_{eq} > 1$.

2-Minute Revision

Galvanic cells are devices that generate electricity from spontaneous redox reactions. Remember the core components: the anode (negative electrode, site of oxidation, electron source), the cathode (positive electrode, site of reduction, electron acceptor), the external circuit (for electron flow), and the salt bridge (for ion migration to maintain charge neutrality).

Electrons always flow from the anode to the cathode through the external wire. The cell potential ($E_{cell}$) is the driving force. Under standard conditions, it's $E^\circ_{cell}$, calculated as $E^\circ_{cathode} - E^\circ_{anode}$ (using standard reduction potentials).

For non-standard conditions, the Nernst equation, $E_{cell} = E^\circ_{cell} - \frac{0.0592}{n} \log Q$ (at $298\,\text{K}$), is used, where $Q$ is the reaction quotient and $n$ is the number of electrons transferred.

The spontaneity of the cell is linked to Gibbs free energy: $\Delta G = -nFE_{cell}$. For a spontaneous cell, $E_{cell}$ is positive, and $\Delta G$ is negative, indicating that the reaction favors product formation and can do electrical work.

5-Minute Revision

A galvanic cell is an electrochemical system that converts chemical energy into electrical energy via a spontaneous redox reaction. It comprises two half-cells: the anode, where oxidation occurs (loss of electrons), and the cathode, where reduction occurs (gain of electrons).

In a galvanic cell, the anode is the negative electrode, and the cathode is the positive electrode. Electrons flow from the anode to the cathode through an external wire. A salt bridge connects the two half-cells, allowing the migration of ions to maintain electrical neutrality and complete the circuit, preventing charge buildup that would otherwise stop the reaction.

Cell representation follows the IUPAC convention: Anode | Anode Electrolyte || Cathode Electrolyte | Cathode.

**Standard Electrode Potentials ($E^\circ$):** These are measured relative to the Standard Hydrogen Electrode (SHE, $0.00\,\text{V}$). The standard cell potential ($E^\circ_{cell}$) is calculated as $E^\circ_{cell} = E^\circ_{cathode} - E^\circ_{anode}$, where both are standard reduction potentials. A positive $E^\circ_{cell}$ indicates a spontaneous reaction.

Nernst Equation: This equation is crucial for calculating cell potential ($E_{cell}$) under non-standard conditions (i.e., when concentrations are not $1\,\text{M}$ or partial pressures are not $1\,\text{atm}$).

At $298\,\text{K}$, it is:

$$E_{cell} = E^\circ_{cell} - \frac{0.0592}{n} \log Q$$

where $n$ is the number of electrons transferred, and $Q$ is the reaction quotient (products/reactants, excluding pure solids/liquids).

Remember that increasing reactant concentration or decreasing product concentration generally increases $E_{cell}$.

**Gibbs Free Energy ($\Delta G$):** The thermodynamic spontaneity of a galvanic cell is quantified by $\Delta G = -nFE_{cell}$. For a spontaneous reaction, $E_{cell} > 0$ and $\Delta G < 0$. At equilibrium, $E_{cell} = 0$, and $\Delta G = 0$. The relationship between $E^\circ_{cell}$ and the equilibrium constant ($K_{eq}$) is $E^\circ_{cell} = \frac{0.0592}{n} \log K_{eq}$ at $298\,\text{K}$. A spontaneous reaction has $K_{eq} > 1$.

Example: For a Daniell cell ($Zn | Zn^{2+} || Cu^{2+} | Cu$), $E^\circ_{Zn^{2+}/Zn} = -0.76\,\text{V}$ and $E^\circ_{Cu^{2+}/Cu} = +0.34\,\text{V}$. Anode: Zn (oxidized), Cathode: Cu (reduced). $E^\circ_{cell} = (+0.34\,\text{V}) - (-0.76\,\text{V}) = +1.10\,\text{V}$. If $[Zn^{2+}] = 0.01\,\text{M}$ and $[Cu^{2+}] = 0.1\,\text{M}$, then $Q = \frac{[Zn^{2+}]}{[Cu^{2+}]} = \frac{0.01}{0.1} = 0.1$. $E_{cell} = 1.10 - \frac{0.0592}{2} \log(0.1) = 1.10 - 0.0296 \times (-1) = 1.1296\,\text{V}$.

Prelims Revision Notes

Galvanic Cells: NEET Quick Recall

1. Definition & Function:

Converts chemical energy from spontaneous redox reactions into electrical energy.
Also known as Voltaic cells.

2. Components & Roles:

Anode: — Site of oxidation (loss of electrons). It is the negative electrode in a galvanic cell. Electrons flow from the anode.
Cathode: — Site of reduction (gain of electrons). It is the positive electrode in a galvanic cell. Electrons flow to the cathode.
External Circuit: — Metallic wire connecting anode and cathode, allowing electron flow.
Salt Bridge: — U-tube containing inert electrolyte (e.g., $KCl$, $KNO_3$).

* Function 1: Completes the internal circuit by allowing ion migration. * Function 2: Maintains electrical neutrality in half-cells (anions to anode, cations to cathode). * Function 3: Prevents mixing of electrolytes.

3. Electron & Ion Flow:

Electrons: — Anode $\rightarrow$ External Wire $\rightarrow$ Cathode.
Anions (from salt bridge): — Towards Anode compartment.
Cations (from salt bridge): — Towards Cathode compartment.

4. Cell Representation (IUPAC):

Anode | Anode Electrolyte || Cathode Electrolyte | Cathode
Single line ($|$) = phase boundary; Double line ($||$) = salt bridge.

5. Standard Electrode Potentials ($E^\circ$):

Measured relative to SHE ($E^\circ_{SHE} = 0.00\,\text{V}$). Values are typically standard reduction potentials.
Anode: — Has more negative (or less positive) $E^\circ_{red}$.
Cathode: — Has more positive $E^\circ_{red}$.

6. Standard Cell Potential ($E^\circ_{cell}$):

$E^\circ_{cell} = E^\circ_{cathode} - E^\circ_{anode}$ (using standard reduction potentials for both).
For a spontaneous reaction (galvanic cell), $E^\circ_{cell}$ must be positive.

7. Nernst Equation (for non-standard conditions):

General form: $E_{cell} = E^\circ_{cell} - \frac{RT}{nF} \ln Q$
At $298\,\text{K}$ (most common in NEET): $E_{cell} = E^\circ_{cell} - \frac{0.0592}{n} \log Q$
$n$ — Number of electrons transferred in the balanced reaction.
$Q$ (Reaction Quotient): — $\frac{[Products]^{coeff}}{[Reactants]^{coeff}}$ (exclude pure solids/liquids).

8. Relationship with Gibbs Free Energy ($\Delta G$):

$\Delta G = -nFE_{cell}$
$\Delta G^\circ = -nFE^\circ_{cell}$
For spontaneity: $\Delta G$ must be negative.

9. Relationship with Equilibrium Constant ($K_{eq}$):

At equilibrium, $E_{cell} = 0$ and $Q = K_{eq}$.
$E^\circ_{cell} = \frac{RT}{nF} \ln K_{eq}$ (or $\frac{0.0592}{n} \log K_{eq}$ at $298\,\text{K}$).
For spontaneity: $K_{eq}$ must be greater than 1.

10. Key Takeaways for NEET:

Always identify anode/cathode correctly from $E^\circ_{red}$ values.
Balance electrons to find 'n' for Nernst and $\Delta G$ calculations.
Pay attention to signs in Nernst equation and $\Delta G$ calculations.
Understand the role of the salt bridge thoroughly.

Vyyuha Quick Recall

AN OX RED CAT

ANode: OXidation
REDuction: CAThode

And for polarity in galvanic cells:

ANode is NEGative (A-N)
CAThode is POSitive (C-P)