18.5 Oxidation Numbers, Electrochemical Cells, Standard Potentials & Batteries

Oxidation Numbers, Electrochemical Cells, Standard Potentials & Batteries

In the Intermediate Chemistry Part I syllabus and the UET ECAT entry test, electrochemistry is a mathematically intensive, heavily weighted domain. ECAT electrochemistry questions assess oxidation state determinations, redox equation balancing via the ion-electron method, electrochemical cell thermodynamics, cell potential calculations from standard reduction series, Nernst equation concentration shifts, commercial battery discharge-charge chemistry, and electrochemical rust prevention.

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1. Oxidation Numbers & Half-Reaction Balancing

The oxidation number (oxidation state) of an atom is defined as the apparent electrical charge that an atom would bear in a molecule or polyatomic ion if all bonding electron pairs were completely assigned to the more electronegative constituent atom.

Systematic Rules for Assigning Oxidation Numbers

1. Free Elements: Any element in its uncombined native elemental state possesses an oxidation number of zero (e.g., $\text{Na}, \text{Fe}, \text{H}_2, \text{O}_2, \text{O}_3, \text{P}_4, \text{S}_8 = 0$).
2. Monoatomic Ions: The oxidation number is identically equal to the charge of the ion (e.g., $\text{Na}^+ = +1, \text{Mg}^{2+} = +2, \text{Al}^{3+} = +3, \text{Cl}^- = -1, \text{S}^{2-} = -2$).
3. Hydrogen: In the vast majority of chemical compounds, Hydrogen is $+1$ (e.g., $\text{HCl}, \text{H}_2\text{O}, \text{NH}_3$). However, in binary metal hydrides (compounds with active metals like $\text{NaH}, \text{CaH}_2, \text{LiAlH}_4$), Hydrogen is more electronegative than the metal and has an oxidation state of $-1$.
4. Oxygen: In almost all covalent and ionic compounds, Oxygen is $-2$. Crucial exceptions tested on the ECAT include:
   • Peroxides (e.g., $\text{H}_2\text{O}_2, \text{Na}_2\text{O}_2, \text{BaO}_2$), where the oxygen-oxygen single bond gives an oxidation state of $-1$.
   • Superoxides (e.g., $\text{KO}_2, \text{RbO}_2$), where each oxygen atom has an average state of $-\frac{1}{2}$.
   • Oxygen Difluoride ($\text{OF}_2$), where Fluorine is more electronegative than oxygen, forcing oxygen into a state of $+2$.
5. Fluorine: As the most electronegative element in the universe, Fluorine is $-1$ in all of its compounds without exception.
6. Algebraic Sum: The sum of oxidation numbers of all atoms in a neutral molecule must equal zero; in a polyatomic ion, the sum must equal the net charge of the ion.

Balancing Redox Equations via the Ion-Electron Method

The ion-electron method (half-reaction method) divides a complex redox equation into two separate half-reactions: oxidation (electron loss) and reduction (electron gain).

Step-by-Step Algorithm in Acidic Medium

1. Identify elements undergoing changes in oxidation number and split the skeletal equation into an oxidation half-reaction and a reduction half-reaction.
2. Balance all atoms other than Oxygen and Hydrogen in each half-reaction using stoichiometric coefficients.
3. Balance Oxygen atoms by adding water molecules ($\text{H}_2\text{O}$) to the oxygen-deficient side.
4. Balance Hydrogen atoms by adding hydrogen ions ($\text{H}^+$) to the hydrogen-deficient side.
5. Balance electrical charge on both sides by adding electrons ($e^-$) to the more positive side.
6. Multiply each balanced half-reaction by the smallest common integer factor so that the total electrons released in oxidation equal the total electrons gained in reduction.
7. Add the two half-reactions together, canceling identical terms ($\text{H}^+, \text{H}_2\text{O}, e^-$) appearing on opposing sides.

Step-by-Step Algorithm in Basic Medium

Follow steps 1 through 5 of the acidic procedure. Then:

1. For every $\text{H}^+$ ion present in the equation, add an equal number of hydroxide ions ($\text{OH}^-$) to BOTH sides of the equation.
2. Combine adjacent $\text{H}^+$ and $\text{OH}^-$ ions on the same side to form $\text{H}_2\text{O}$ ($\text{H}^+ + \text{OH}^- \rightarrow \text{H}_2\text{O}$).
3. Cancel excess $\text{H}_2\text{O}$ molecules present on opposing sides of the combined reaction.

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2. Electrochemical Cells: Electrolytic vs. Galvanic Cells

Electrochemistry studies the interconversion of chemical energy and electrical energy across phase boundaries. All electrochemical cells partition into two fundamental classes:

Universal ECAT Rule: In ALL electrochemical cells without exception:

• AN OX: The Anode is ALWAYS the electrode where Oxidation occurs.
• RED CAT: The Cathode is ALWAYS the electrode where Reduction occurs.

The Daniell Cell Mechanics

A classic galvanic cell invented by John Frederic Daniell in 1836 consists of a Zinc plate immersed in $1.0\text{ M } \text{ZnSO}_4$ solution (anode half-cell) and a Copper plate immersed in $1.0\text{ M } \text{CuSO}_4$ solution (cathode half-cell), connected externally by a metallic conductor and internally by a salt bridge:

                  Voltmeter (1.10 V)
             [ Zn ] ───► e⁻ ───► [ Cu ]
             (Anode, -)          (Cathode, +)
                │                    │
        ┌───────┴────────┐   ┌───────┴────────┐
        │  ZnSO₄ (1.0 M) │   │  CuSO₄ (1.0 M) │
        └────────────────┘   └────────────────┘
                ▲                    ▲
                └──── Salt Bridge ───┘
                      (KCl / KNO₃)

• Anode Reaction (Oxidation): $\text{Zn}(s) \rightarrow \text{Zn}^{2+}(aq) + 2e^- \quad E^\circ_{\text{ox}} = +0.76\text{ V} \quad (E^\circ_{\text{red}} = -0.76\text{ V})$
• Cathode Reaction (Reduction): $\text{Cu}^{2+}(aq) + 2e^- \rightarrow \text{Cu}(s) \quad E^\circ_{\text{red}} = +0.34\text{ V}$
• Net Overall Cell Reaction: $\text{Zn}(s) + \text{Cu}^{2+}(aq) \rightarrow \text{Zn}^{2+}(aq) + \text{Cu}(s) \quad E^\circ_{\text{cell}} = +1.10\text{ V}$
• Standard Cell Notation: $\text{Zn}(s) \mid \text{Zn}^{2+}(aq, 1\text{ M}) \parallel \text{Cu}^{2+}(aq, 1\text{ M}) \mid \text{Cu}(s)$

Functions of the Salt Bridge

The salt bridge is an inverted U-tube packed with a solidified gel of agar-agar containing a concentrated solution of an inert electrolyte such as $\text{KCl}, \text{KNO}_3$, or $\text{NH}_4\text{NO}_3$. The salt bridge serves three vital functions:

1. Completes the Electrical Circuit: It provides an internal ionic pathway connecting the two half-cells, allowing current to flow continuously without mechanical mixing of solutions.
2. Maintains Electrical Neutrality: As zinc dissolves at the anode, excess positive $\text{Zn}^{2+}$ cations accumulate in the anode compartment. To neutralize this charge, anions ($\text{Cl}^-$ or $\text{NO}_3^-$) migrate from the salt bridge into the anode beaker. At the cathode, $\text{Cu}^{2+}$ cations deposit as copper metal, leaving unpartnered $\text{SO}_4^{2-}$ anions; cations ($\text{K}^+$) migrate from the bridge into the cathode beaker. Without this charge neutralization, opposing polarization would halt current within microseconds.
3. Eliminates Liquid Junction Potential: When two dissimilar electrolyte solutions touch directly, ions diffuse across the boundary at unequal velocities, creating an undesirable electrostatic contact potential known as the liquid junction potential. In the salt bridge, $\text{K}^+$ and $\text{Cl}^-$ (or $\text{NO}_3^-$) possess virtually identical ionic mobilities, ensuring that diffusion rates match and canceling junction potentials.

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3. Standard Hydrogen Electrode (SHE) & Standard Potentials ($E^\circ$)

The absolute single electrode potential of an isolated half-cell cannot be measured directly because an oxidation or reduction half-reaction cannot occur in isolation. To measure relative potentials, IUPAC adopted the Standard Hydrogen Electrode (SHE) as the universal reference electrode.

Construction of SHE

• A piece of platinum foil coated with finely divided, colloidal platinum black (to maximize surface area and catalytically accelerate hydrogen gas adsorption).
• Immersed in an aqueous solution of a strong monoprotic acid having a hydronium ion activity of unit molarity ($[\text{H}^+] = 1.0\text{ M}$ or $1.0\text{ M } \text{HCl}$).
• Pure hydrogen gas ($\text{H}_2$) is bubbled continuously over the platinum foil at a standard pressure of $1.0\text{ atmosphere}$ ($101.3\text{ kPa}$) and standard temperature of $298.15\text{ K}$ ($25^\circ\text{C}$).

$\text{Half-reaction: } 2\text{H}^+(aq, 1\text{ M}) + 2e^- \rightleftharpoons \text{H}_2(g, 1\text{ atm})$

Universal Reference Standard: By international convention, the standard electrode potential of the Standard Hydrogen Electrode is assigned a value of identically zero volts at all temperatures: $E^\circ_{\text{SHE}} = 0.000\text{ V}$

To determine the standard reduction potential of any unknown metal electrode (e.g., $\text{Zn}^{2+}/\text{Zn}$), the metal half-cell is coupled with the SHE in a galvanic circuit under standard thermodynamic conditions ($25^\circ\text{C}, 1\text{ atm}, 1\text{ M}$). If the metal electrode acts as anode (spontaneous oxidation), electrons flow toward SHE; the measured cell potential equals the negative reduction potential of that metal ($E^\circ_{\text{red}} = -0.76\text{ V}$ for Zinc).

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4. The Electrochemical Series & Its Engineering Applications

The Electrochemical Series (Standard Reduction Potential Series) is an arrangement of chemical elements and redox half-reactions in order of their increasing standard reduction potentials ($E^\circ$) relative to the SHE at $298\text{ K}$:

Top of Series (Most Negative E°):
Li⁺ + e⁻  ⇄  Li        E° = -3.05 V   [Strongest Reducing Agent, Most Electropositive]
K⁺ + e⁻   ⇄  K         E° = -2.93 V
Ca²⁺ + 2e⁻ ⇄  Ca        E° = -2.87 V
Na⁺ + e⁻  ⇄  Na        E° = -2.71 V
Mg²⁺ + 2e⁻ ⇄  Mg        E° = -2.37 V
Al³⁺ + 3e⁻ ⇄  Al        E° = -1.66 V
Zn²⁺ + 2e⁻ ⇄  Zn        E° = -0.76 V
Fe²⁺ + 2e⁻ ⇄  Fe        E° = -0.44 V
─────────────────────────────────────────────────────────────────────────────
Reference Benchmark:
2H⁺ + 2e⁻ ⇄  H₂(g)     E° =  0.00 V   [SHE Standard]
─────────────────────────────────────────────────────────────────────────────
Bottom of Series (Most Positive E°):
Cu²⁺ + 2e⁻ ⇄  Cu        E° = +0.34 V
I₂(s) + 2e⁻ ⇄  2I⁻       E° = +0.54 V
Fe³⁺ + e⁻  ⇄  Fe²⁺      E° = +0.77 V
Ag⁺ + e⁻  ⇄  Ag        E° = +0.80 V
Br₂(l) + 2e⁻ ⇄ 2Br⁻     E° = +1.07 V
Cl₂(g) + 2e⁻ ⇄ 2Cl⁻     E° = +1.36 V
F₂(g) + 2e⁻ ⇄  2F⁻      E° = +2.87 V   [Strongest Oxidizing Agent, Highest Electronegativity]

Four Major Applications Tested on ECAT

1. Calculating Standard Cell Electromotive Force ($E^\circ_{\text{cell}}$)

For any galvanic cell where both potentials are expressed as standard reduction potentials:

$E^\circ_{\text{cell}} = E^\circ_{\text{cathode}} - E^\circ_{\text{anode}}$

$\text{For Daniell Cell: } E^\circ_{\text{cell}} = E^\circ(\text{Cu}^{2+}/\text{Cu}) - E^\circ(\text{Zn}^{2+}/\text{Zn}) = (+0.34\text{ V}) - (-0.76\text{ V}) = +1.10\text{ V}$

2. Predicting Reaction Spontaneity & Gibbs Free Energy ($\Delta G^\circ$)

The standard cell potential relates directly to standard Gibbs free energy change via the fundamental thermodynamic relation:

$\Delta G^\circ = -n F E^\circ_{\text{cell}}$

where $n$ is the moles of electrons transferred per mole of reaction, and $F$ is Faraday's constant ($F = 96,485\text{ C/mol} \approx 96,500\text{ C/mol}$).

• Spontaneous Process: $\Delta G^\circ < 0 \iff E^\circ_{\text{cell}} > 0$ (positive emf indicates forward spontaneity).
• Non-Spontaneous Process: $\Delta G^\circ > 0 \iff E^\circ_{\text{cell}} < 0$ (reaction proceeds spontaneously in reverse).
• Equilibrium State: $\Delta G^\circ = 0 \iff E^\circ_{\text{cell}} = 0$.

3. Predicting Displacement Reactions in Aqueous Solution

A metal situated higher in the electrochemical series (more negative reduction potential) is a stronger reducing agent; it will spontaneously displace any metal situated lower in the series (more positive reduction potential) from its aqueous salt solution:

$\text{Zn}(s) + \text{CuSO}_4(aq) \rightarrow \text{ZnSO}_4(aq) + \text{Cu}(s) \quad (\text{Spontaneous: } E^\circ_{\text{cell}} = +1.10\text{ V})$ $\text{Cu}(s) + \text{ZnSO}_4(aq) \rightarrow \text{No Reaction} \quad (\text{Non-spontaneous: } E^\circ_{\text{cell}} = -1.10\text{ V})$

4. Predicting Hydrogen Gas Evolution with Mineral Acids

Any metal with a negative standard reduction potential ($E^\circ < 0.00\text{ V}$, placed above hydrogen in the series: $\text{Li, Na, Mg, Al, Zn, Fe}$) can spontaneously reduce aqueous $\text{H}^+$ ions to liberate hydrogen gas from dilute non-oxidizing acids ($\text{HCl}, \text{H}_2\text{SO}_4$):

$\text{Zn}(s) + 2\text{H}^+(aq) \rightarrow \text{Zn}^{2+}(aq) + \text{H}_2(g) \quad E^\circ_{\text{cell}} = 0.00 - (-0.76) = +0.76\text{ V} \quad (\text{Spontaneous})$

Metals with positive reduction potentials ($E^\circ > 0.00\text{ V}$, placed below hydrogen: $\text{Cu, Ag, Au}$) cannot reduce $\text{H}^+$ ions and cannot liberate $\text{H}_2$ gas from non-oxidizing mineral acids.

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5. The Nernst Equation: Concentration Dependence of Cell Potential

Standard potentials ($E^\circ$) apply exclusively when all dissolved solutes are at unit activity ($1.0\text{ M}$) and gases are at $1.0\text{ atm}$ at $298\text{ K}$. In real cells, concentrations vary continuously. The German physical chemist Walther Nernst derived the fundamental relationship between cell potential ($E$) and the reaction quotient ($Q$):

$\Delta G = \Delta G^\circ + RT \ln Q$

Substituting $\Delta G = -nFE$ and $\Delta G^\circ = -nFE^\circ$:

$-nFE = -nFE^\circ + RT \ln Q$

$E = E^\circ - \frac{RT}{nF} \ln Q = E^\circ - \frac{2.303 RT}{nF} \log_{10} Q$

At standard room temperature $T = 298.15\text{ K}$ ($25^\circ\text{C}$), substituting $R = 8.314\text{ J}\cdot\text{mol}^{-1}\cdot\text{K}^{-1}$ and $F = 96,485\text{ C}\cdot\text{mol}^{-1}$ evaluates the numerical pre-factor to $0.0592\text{ V}$:

$E = E^\circ - \frac{0.0592}{n}\log_{10} Q$

Application to the Daniell Cell

For $\text{Zn}(s) + \text{Cu}^{2+}(aq) \rightleftharpoons \text{Zn}^{2+}(aq) + \text{Cu}(s)$ where $n=2$ and pure solids carry unit activity:

$Q = \frac{[\text{Zn}^{2+}]}{[\text{Cu}^{2+}]}$

$E_{\text{cell}} = 1.10 - \frac{0.0592}{2}\log_{10}\left(\frac{[\text{Zn}^{2+}]}{[\text{Cu}^{2+}]}\right)$

• Effect of Increasing Reactant Concentration ($[\text{Cu}^{2+}] \uparrow$): Decreases the ratio $\frac{[\text{Zn}^{2+}]}{[\text{Cu}^{2+}]}$, making $\log Q$ negative, thereby increasing cell emf ($E_{\text{cell}} > 1.10\text{ V}$).
• Effect of Increasing Product Concentration ($[\text{Zn}^{2+}] \uparrow$): Increases $Q$, depressing cell emf ($E_{\text{cell}} < 1.10\text{ V}$).
• At Dynamic Chemical Equilibrium ($E_{\text{cell}} = 0$): The battery is fully dead. The reaction quotient equals the equilibrium constant ($Q = K_{\text{eq}}$): $0 = E^\circ - \frac{0.0592}{n}\log_{10} K_{\text{eq}} \implies \log_{10} K_{\text{eq}} = \frac{n E^\circ}{0.0592}$ For the Daniell cell at $25^\circ\text{C}$: $\log K_{\text{eq}} = \frac{2(1.10)}{0.0592} = 37.16 \implies K_{\text{eq}} \approx 1.5 \times 10^{37}$, demonstrating that the reaction proceeds virtually to absolute completion.

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6. Commercial Cells & Batteries

1. Dry Cell (Leclanché Cell)

An irreversible primary galvanic cell generating approximately $1.5\text{ V}$:

• Anode Container: Zinc metal cylinder: $\text{Zn}(s) \rightarrow \text{Zn}^{2+}(aq) + 2e^-$
• Cathode: Central graphite rod surrounded by a moist paste of manganese dioxide ($\text{MnO}_2$), carbon black, ammonium chloride ($\text{NH}_4\text{Cl}$), and zinc chloride ($\text{ZnCl}_2$): $2\text{MnO}_2(s) + 2\text{NH}_4^+(aq) + 2e^- \rightarrow \text{Mn}_2\text{O}_3(s) + 2\text{NH}_3(aq) + \text{H}_2\text{O}(l)$
• Ammonia gas is prevented from rupturing the cell by reacting with $\text{Zn}^{2+}$ to form the soluble complex $[\text{Zn}(\text{NH}_3)_4]^{2+}$.

2. Lead Storage Battery (Automobile Lead-Acid Accumulator)

A reversible secondary (rechargeable) battery consisting of six $2\text{ V}$ cells connected in series to deliver a terminal potential of $12\text{ V}$:

• Cell Construction:
  • Anode Plates: Spongy metallic Lead ($\text{Pb}$).
  • Cathode Plates: Lead grids packed with Lead Dioxide ($\text{PbO}_2$).
  • Electrolyte: Aqueous sulfuric acid ($\text{H}_2\text{SO}_4$) with an initial specific gravity of $1.28\text{ g/cm}^3$ (approx. $38\%\text{ w/w}$). Plates are interleaved and separated by porous insulating sheets.
• Discharge Chemistry (Delivering Electrical Power): $\text{Anode: } \text{Pb}(s) + \text{SO}_4^{2-}(aq) \rightarrow \text{PbSO}_4(s) + 2e^- \quad (E^\circ_{\text{ox}} = +0.36\text{ V})$ $\text{Cathode: } \text{PbO}_2(s) + 4\text{H}^+(aq) + \text{SO}_4^{2-}(aq) + 2e^- \rightarrow \text{PbSO}_4(s) + 2\text{H}_2\text{O}(l) \quad (E^\circ_{\text{red}} = +1.69\text{ V})$ $\text{Net Discharge Reaction: } \text{Pb}(s) + \text{PbO}_2(s) + 2\text{H}_2\text{SO}_4(aq) \rightarrow 2\text{PbSO}_4(s) + 2\text{H}_2\text{O}(l) \quad (E^\circ_{\text{cell}} \approx 2.05\text{ V})$ Diagnostic Shift: During discharge, sulfuric acid is consumed and water is produced. The specific gravity of the electrolyte drops from $1.28\text{ g/cm}^3$ down to below $1.15\text{ g/cm}^3$, indicating the battery is depleted. Furthermore, both electrodes become coated with insoluble solid lead sulfate ($\text{PbSO}_4$).
• Recharge Chemistry (External DC Power Reversal): When connected to an external direct-current charging source whose voltage exceeds the battery's emf, electrons are forced in reverse, driving the non-spontaneous reformation of original reactants: $2\text{PbSO}_4(s) + 2\text{H}_2\text{O}(l) \xrightarrow{\text{recharge}} \text{Pb}(s) + \text{PbO}_2(s) + 2\text{H}_2\text{SO}_4(aq)$ This restores the original electrodes and replenishes sulfuric acid, returning the density to $1.28\text{ g/cm}^3$.

3. Alkaline Fuel Cells (Bacon $\text{H}_2-\text{O}_2$ Cell)

Fuel cells are electrochemical galvanic devices that convert the chemical energy of a continuously supplied fuel and oxidizer directly into electricity without intermediate thermal combustion:

• Reagents & Electrolyte: Hydrogen gas (fuel) and oxygen gas (oxidizer) are bubbled through porous carbon electrodes impregnated with platinum/palladium catalysts into hot concentrated aqueous potassium hydroxide ($\text{KOH}$) at $200^\circ\text{C}$: $\text{Anode: } 2\text{H}_2(g) + 4\text{OH}^-(aq) \rightarrow 4\text{H}_2\text{O}(l) + 4e^-$ $\text{Cathode: } \text{O}_2(g) + 2\text{H}_2\text{O}(l) + 4e^- \rightarrow 4\text{OH}^-(aq)$ $\text{Overall Cell Reaction: } 2\text{H}_2(g) + \text{O}_2(g) \rightarrow 2\text{H}_2\text{O}(l) \quad E^\circ_{\text{cell}} = 1.23\text{ V}$
• Key Advantages: High thermodynamic conversion efficiency ($>70\%$ vs $35-40\%$ for conventional thermal power plants), continuous operation with zero pollution, and the byproduct is pure potable water used by astronauts during Apollo space flights.

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7. Corrosion and Rust Prevention Mechanisms

Corrosion is the destructive deterioration of a metal through spontaneous electrochemical oxidation induced by environmental agents (oxygen, moisture, electrolytes).

Electrochemical Mechanism of Rusting of Iron

Rusting requires both oxygen gas and liquid water simultaneously:

1. Anodic Site Formation: Strained regions, impurities, or crystal dislocations on the iron surface act as anodes where iron oxidizes: $\text{Fe}(s) \rightarrow \text{Fe}^{2+}(aq) + 2e^- \quad E^\circ = -0.44\text{ V}$
2. Cathodic Site Formation: The liberated electrons migrate through the conductive iron metal to an adjacent surface region with high dissolved oxygen and moisture concentration, which serves as the cathode: $\text{O}_2(g) + 4\text{H}^+(aq) + 4e^- \rightarrow 2\text{H}_2\text{O}(l) \quad E^\circ = +1.23\text{ V}$ Acidic gases ($\text{CO}_2, \text{SO}_2$) dissolve in water droplets to form $\text{H}^+$ ions, accelerating the reduction.
3. Rust Precipitation: The $\text{Fe}^{2+}$ ions diffuse through the water droplet and are further oxidized by atmospheric oxygen to hydrated iron(III) oxide (rust): $4\text{Fe}^{2+}(aq) + \text{O}_2(g) + 4\text{H}_2\text{O}(l) + 2x\text{H}_2\text{O} \rightarrow 2\text{Fe}_2\text{O}_3 \cdot x\text{H}_2\text{O}(s) + 8\text{H}^+(aq)$

Corrosion Prevention Technologies

1. Galvanizing (Zinc Coating): Iron is coated with a thin layer of metallic zinc ($E^\circ_{\text{Zn}^{2+}/\text{Zn}} = -0.76\text{ V}$). Because zinc's reduction potential is more negative than iron's ($E^\circ_{\text{Fe}^{2+}/\text{Fe}} = -0.44\text{ V}$), zinc is more easily oxidized. Even if the galvanized zinc layer is scratched or chipped, an electrochemical couple forms where zinc acts as a sacrificial anode and corrodes preferentially, while iron acts as the cathode and remains protected.
2. Tin Plating (Tinning): Iron is coated with tin ($E^\circ_{\text{Sn}^{2+}/\text{Sn}} = -0.14\text{ V}$). Because tin is non-toxic and resists organic acids, it is widely used in food cans. However, tin has a less negative reduction potential than iron. Tin protects iron strictly as an impermeable mechanical barrier. If the tin coating is scratched, an electrochemical cell forms in which iron has the more negative reduction potential: iron oxidizes rapidly as the anode, corroding much faster than if the tin were absent.
3. Cathodic Protection (Sacrificial Anodes): Underground steel fuel pipelines, storage tanks, and ship hulls are electrically connected by insulated copper cables to blocks of more electropositive sacrificial metals, such as Magnesium ($E^\circ = -2.37\text{ V}$) or Zinc ($E^\circ = -0.76\text{ V}$) buried in damp soil. The active metal oxidizes sacrificially, driving electrons into the iron structure and keeping it permanently cathodic and uncorroded.

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8. Electrochemical Parameters & Battery Summary Table