Electrochemistry Notes

Electrochemistry

• Electrochemistry studies electricity production from spontaneous chemical reactions and using electrical energy for non-spontaneous transformations.

• It's important for both theoretical and practical reasons.

• Electrochemical methods are used to produce:

  • Metals

  • Sodium hydroxide

  • Chlorine

  • Fluorine

  • Other chemicals

• Batteries and fuel cells convert chemical energy into electrical energy.

• Electrochemical reactions are energy efficient and less polluting, making them important for eco-friendly technologies.

• Sensory signal transmission and cell communication have electrochemical origins.

• Electrochemistry is interdisciplinary.

Objectives

• Chemical reactions can produce electrical energy, and electrical energy can drive non-spontaneous chemical reactions.

Electrochemical Cells

• Daniell Cell: Converts chemical energy from the redox reaction \Zn(s) + Cu^{2+}(aq) \rightleftharpoons Zn^{2+}(aq) + Cu(s)

  • Electrical potential of 1.1 V when ion concentrations are unity (1 mol dm⁻³).

• Galvanic (Voltaic) Cell: A device like the Daniell cell.

• Applying an opposing voltage \E_{ext}:

  • When \E_{ext} < 1.1 V: Reaction continues, electrons flow from Zn to Cu, current flows from Cu to Zn.

    • Zinc dissolves at the anode, and copper deposits at the cathode.

  • When \E_{ext} = 1.1 V: Reaction stops, no current flows.

  • When \E_{ext} > 1.1 V: Reaction reverses, functions as an electrolytic cell.

    • Electrons flow from Cu to Zn, current flows from Zn to Cu.

    • Zinc deposits at the zinc electrode, and copper dissolves at the copper electrode.

• Electrolytic Cell: Uses electrical energy to drive non-spontaneous chemical reactions.

Galvanic Cells

• Galvanic cells convert the chemical energy of spontaneous redox reactions into electrical energy.

• Gibbs energy is converted into electrical work.

• Example: Daniell cell with the redox reaction:
  \Zn(s) + Cu^{2+}(aq) \rightleftharpoons Zn^{2+}(aq) + Cu(s)

• Half-Reactions:

  • Reduction: \Cu^{2+} + 2e^- \rightarrow Cu(s)

  • Oxidation: \Zn(s) \rightarrow Zn^{2+} + 2e^-

• Half-cells (Redox Couples):

  • Reduction half-cell: Copper electrode

  • Oxidation half-cell: Zinc electrode

• A metallic wire connects the two half-cells externally with a voltmeter and switch.

• A salt bridge connects the electrolytes internally.

• Electrode Potential:

  • Metal ions from the solution deposit on the metal electrode, making it positively charged.

  • Metal atoms from the electrode go into solution as ions, leaving electrons behind and making the electrode negatively charged.

  • At equilibrium, charge separation creates a potential difference called electrode potential.

• Standard Electrode Potential:

  • Electrode potential when all species concentrations are unity.

  • IUPAC convention: Standard reduction potentials are called standard electrode potentials.

• Anode (oxidation) has a negative potential, and cathode (reduction) has a positive potential.

• Electrons flow from the negative electrode to the positive electrode when the switch is on.

• Current flow direction is opposite to electron flow.

Cell Potential

• Cell potential: Potential difference between two electrodes of a galvanic cell, measured in volts.

• Electromotive force (emf): Cell potential when no current is drawn.

• Convention: Anode on the left, cathode on the right.

• Galvanic cell representation:

  • Single vertical line: Metal and electrolyte solution interface.

  • Double vertical line: Salt bridge between two electrolytes.

• emf of the cell is positive:
  \E{cell} = E{right} - E_{left}

• Example:

  • Cell reaction: \Cu(s) + 2Ag^+(aq) \rightleftharpoons Cu^{2+}(aq) + 2Ag(s)

  • Half-cell reactions:

    • Cathode (reduction): \2Ag^+(aq) + 2e^- \rightarrow 2Ag(s)

    • Anode (oxidation): \Cu(s) \rightarrow Cu^{2+}(aq) + 2e^-

  • Cell representation: \Cu(s)|Cu^{2+}(aq)||Ag^+(aq)|Ag(s)

  • \E{cell} = E{Ag^+/Ag} - E_{Cu^{2+}/Cu}

• Individual half-cell potentials cannot be measured; only the difference is measurable.

• Standard Hydrogen Electrode (SHE):

  • Assigned a zero potential at all temperatures.

  • Representation: \Pt(s)|H_2(g)|H^+(aq)

  • Reaction: $H^+ (aq) + e^- \rightleftharpoons \frac{1}{2} H_2(g)$

  • Platinum electrode coated with platinum black in acidic solution with hydrogen gas bubbled through it.

  • Hydrogen gas pressure is one bar, and hydrogen ion concentration is one molar.

Measurement of Electrode Potential

• At 298 K, the emf of the cell (standard hydrogen electrode || second half-cell) gives the reduction potential of the other half-cell.

• If the concentrations of oxidized and reduced forms are unity, the cell potential equals the standard electrode potential:\ \E^o = E^oR - E^oL

  • Since \E^oL for SHE is zero: \E^o = E^oR

• Example 1:

  • Cell: \Pt(s) | H_2(g, 1 bar) | H^+ (aq, 1 M) || Cu^{2+} (aq, 1 M) | Cu

  • emf = 0.34 V, which is the standard electrode potential for the reaction:
    \Cu^{2+} (aq, 1M) + 2 e^- \rightarrow Cu(s)

• Example 2:

  • Cell: \Pt(s) | H_2(g, 1 bar) | H^+ (aq, 1 M) || Zn^{2+} (aq, 1M) | Zn

  • emf = -0.76 V, corresponding to the standard electrode potential of:
    \Zn^{2+} (aq, 1 M) + 2e^- \rightarrow Zn(s)

• Positive \E^o: Reduced form is more stable than hydrogen gas (e.g., $Cu^{2+}$ is more easily reduced than $H^+$, so Cu doesn't dissolve in HCl)

• Negative \E^o: Hydrogen gas is more stable than the reduced form (e.g., $H^+$ can oxidize Zn, or Zn can reduce $H^+$).

• Daniell Cell Half-Reactions:

  • Left electrode (anode): \Zn(s) \rightarrow Zn^{2+} (aq, 1 M) + 2 e^-

  • Right electrode (cathode): \Cu^{2+} (aq, 1 M) + 2 e^- \rightarrow Cu(s)

  • Overall reaction: \Zn(s) + Cu^{2+} (aq) \rightarrow Zn^{2+} (aq) + Cu(s)

  • emf of the cell: \E^o{cell} = E^oR - E^o_L = 0.34V - (-0.76V) = 1.10 V

• Inert Electrodes:

  • Metals like platinum or gold provide a surface for oxidation or reduction reactions but do not participate directly.

  • Hydrogen electrode: \Pt(s)|H2(g)| H^+(aq), with reaction: \H^+ (aq)+ e^- \rightarrow \frac{1}{2} H2(g)

  • Bromine electrode: \Pt(s)|Br2(aq)| Br^-(aq), with reaction: $\frac{1}{2} Br</em>2(aq) + e^- \rightarrow Br^-(aq)$

Standard Electrode Potentials

• If \E^o > 0: the reduced form is more stable than hydrogen gas.

• If \E^o < 0: hydrogen gas is more stable than the reduced form.

• Fluorine (F2) has the highest \E^o, making it the strongest oxidizing agent and fluoride ion (F⁻) the weakest reducing agent.

• Lithium has the lowest electrode potential, making lithium ion the weakest oxidizing agent and lithium metal the most powerful reducing agent in aqueous solution.

• As you go down the table, the oxidizing power of the species on the left decreases, and the reducing power of the species on the right increases.

• Electrochemical cells are used to determine pH, solubility product, equilibrium constant, thermodynamic properties, and for potentiometric titrations.

Nernst Equation

• Nernst Equation for Electrode Reaction:
  \Mn^+(aq) + ne^- \rightleftharpoons M(s)

  • Electrode potential at any concentration:
    \E{M^{n+}/M} = E^o{M^{n+}/M} - \frac{RT}{nF} \ln \frac{[M]}{[M^{n+}]}

  • Since the concentration of solid M is unity:
    \E{M^{n+}/M} = E^o{M^{n+}/M} - \frac{RT}{nF} \ln \frac{1}{[M^{n+}]}

  • R = gas constant (8.314 JK⁻¹ mol⁻¹), F = Faraday constant (96487 C mol⁻¹), T = temperature in Kelvin, and [Mn+] is the concentration of the species Mn+.

• Nernst Equation for Daniell Cell:

  • For Cathode:
    \E{Cu^{2+}/Cu} = E^o{Cu^{2+}/Cu} - \frac{RT}{2F} \ln \frac{1}{[Cu^{2+}(aq)]}

  • For Anode:
    \E{Zn^{2+}/Zn} = E^o{Zn^{2+}/Zn} - \frac{RT}{2F} \ln \frac{1}{[Zn^{2+}(aq)]}

• Cell Potential:
  \E{cell} = E{Cu^{2+}/Cu} - E{Zn^{2+}/Zn} \E{cell} = E^o_{cell} - \frac{RT}{2F} \ln \frac{[Zn^{2+}]}{[Cu^{2+}]}

  • \E_{cell} depends on $[Cu^{2+}]$ and $[Zn^{2+}]$; increases with $\uparrow [Cu^{2+}]$ and $\downarrow [Zn^{2+}]$.

• At T = 298 K:
  \E{cell} = E^o{cell} - \frac{0.0592}{2} \log \frac{[Zn^{2+}]}{[Cu^{2+}]}

• General Electrochemical Reaction:

  • $aA + bB + ne^- \rightleftharpoons cC + dD$

  • Nernst equation:
    \E{cell} = E^o{cell} - \frac{RT}{nF} \ln Q \E{cell} = E^o{cell} - \frac{RT}{nF} \ln \frac{[C]^c[D]^d}{[A]^a[B]^b}

Equilibrium Constant from Nernst Equation

• At equilibrium, \E_{cell} = 0

• $E^o<em>{cell} = \frac{2.303RT}{nF} \log K</em>C$

• For Daniell cell at equilibrium: \E{cell} = 0 = E^o{cell} - \frac{2.303RT}{2F} \log \frac{[Zn^{2+}]}{[Cu^{2+}]}

  • $E^o<em>{cell} = \frac{2.303RT}{2F} \log K</em>C$

  • At T = 298K, $E^o<em>{cell} = \frac{0.0592}{2} \log K</em>C$

  • If \E^o{cell} = 1.1V $\log K</em>C = \frac{(1.1V \times 2)}{0.059 V} = 37.288$

  • $K_C = 2 \times 10^{37}$

• General Relationship:
  \E^o{cell} = \frac{2.303RT}{nF} \log KC

• Equilibrium constants can be calculated from Eo values.

Electrochemical Cell and Gibbs Energy of the Reaction

• Electrical work done in one second is electrical potential multiplied by total charge passed.

• Reversible work done by a galvanic cell equals the decrease in Gibbs energy: $\Delta<em>rG = -nFE</em>{cell}$

  • E(cell) is intensive, but ΔrG is extensive.

• For standard conditions:
  $\Delta<em>rG^o = -nFE^o</em>{cell}$

• Measuring \E^o{cell} gives $\Delta</em>rG^o$, from which equilibrium constant K can be found:
  $\Delta_rG^o = -RT \ln K$

Conductance of Electrolytic Solutions

• Electrical Resistance (R):

  • Measured in ohms (Ω), which equals $\frac{kg \cdot m^2}{S^3 \cdot A^2}$

  • Measured using a Wheatstone bridge.

  • $R \propto \frac{l}{A}$, where l is length and A is the area of cross-section.

  • $R = \rho \frac{l}{A}$, where $\rho$ (rho) is resistivity.

• Resistivity (Specific Resistance) $\rho$:

  • SI units: ohm metre (Ω m) or ohm centimetre (Ω cm).

  • 1 Ω m = 100 Ω cm, 1 Ω cm = 0.01 Ω m

  • Resistivity is the resistance when the substance is one meter long and has a cross-section area of one square meter.

• Conductance (G):

  • $G = \frac{1}{R} = \kappa \frac{A}{l}$

  • SI unit: siemens (S), equal to ohm⁻¹ (Ω⁻¹ or mho).

• Conductivity (Specific Conductance) $\kappa$ (kappa):

  • $\kappa = \frac{1}{\rho}$

  • SI units: S m⁻¹, often expressed in S cm⁻¹.

  • Conductivity is the conductance when the substance is 1 m long with a cross-section area of 1 m².

  • 1 S cm⁻¹ = 100 S m⁻¹.

Metallic and Electrolytic Conductance

• Metallic (Electronic) Conductance:

  • Electrical conductance through metals due to electron movement.

  • Depends on:

    • Nature and structure of the metal.

    • Number of valence electrons per atom.

    • Temperature (decreases with increasing temperature).

  • Composition of the metallic conductor remains unchanged.

• Electrolytic (Ionic) Conductance:

  • Conductance of electricity by ions in solution.

  • Depends on:

    • Nature of the electrolyte added.

    • Size of ions and their solvation.

    • Nature of the solvent and its viscosity.

    • Concentration of the electrolyte.

    • Temperature (increases with increasing temperature).

  • Prolonged direct current can change the solution's composition due to electrochemical reactions.

Measurement of Conductivity of Ionic Solutions

• Using Alternating Current (AC):

  • Avoids changes in the solution's composition.

• Conductivity Cell:

  • Special vessel to contain the solution.

  • Two platinum electrodes coated with platinum black.

  • Electrodes have area A and are separated by distance l.

• Resistance of the solution column:

  • $R = \rho \frac{l}{A} = \frac{1}{\kappa} \frac{l}{A}$

• Cell Constant (G*):

  • $G* = \frac{l}{A}$, with dimension length⁻¹.

  • Determined by measuring the resistance of a solution with known conductivity (e.g., KCl).

  • $G* = \frac{l}{A} = R \kappa$

• Wheatstone Bridge Setup:

  • Two resistances R3 and R4, a variable resistance R1, and the conductivity cell (R2).

  • Oscillator O (AC power source).

  • Detector P (headphone or electronic device).

  • Bridge is balanced when no current passes through the detector.

  • Unknown resistance: $R<em>2 = \frac{R</em>1 R<em>4}{R</em>3}$

• Conductivity Meters:

  • Inexpensive devices that directly read conductance or resistance.

• Conductivity Calculation:

  • $\kappa = \frac{G*}{R}$

Molar Conductivity

• Molar conductivity ($\Lambda<em>m$): conductivity of the solution divided by the concentration. $\Lambda</em>m = \frac{\kappa}{c}$

  • If κ is in S m⁻¹ and c is in mol m⁻³, then Λm is in S m² mol⁻¹.

  • If κ is in S cm⁻¹ and c is in mol cm⁻³, then Λm is in S cm² mol⁻¹.

  • 1 mol m⁻³ = 1000 (L/m³) × molarity (mol/L).

  • $\Lambda_m (S cm^2 mol^{-1}) = \frac{\kappa (S cm^{-1}) \times 1000 (cm^3/L)}{molarity (mol/L)}$

• Units relationship:

  • 1 S m² mol⁻¹ = 10⁴ S cm² mol⁻¹

  • 1 S cm² mol⁻¹ = 10⁻⁴ S m² mol⁻¹

Variation of Conductivity and Molar Conductivity with Concentration

• Conductivity:

  • Decreases with decreased concentration.

  • Fewer ions per unit volume carry current on dilution.

  • $\kappa = G \frac{l}{A}$

• Molar Conductivity:

  • Increases with decreased concentration.

  • The total volume, V, of solution containing one mole of electrolyte increases.

  • $\Lambda_m = \kappa V$

• Limiting Molar Conductivity ($\Lambda_m^o$):

  • Molar conductivity at zero concentration.

• Different behavior for strong and weak electrolytes.

• Strong electrolytes undergo a small increase in molar conductivity with dilution.

• Empirical relationship:
  $\Lambda<em>m = \Lambda</em>m^o - A \sqrt{c}$

  • Plotting Λm against c^(1/2) gives a straight line with intercept Λm^o and slope -A.

  • Constant A depends on the type of electrolyte (e.g., 1-1, 2-1, 2-2).

• Weak electrolytes exhibit a large increase in molar conductivity with dilution, especially at low concentrations.

• Limiting molar conductivity ($\Lambda_m^o$) cannot be obtained by extrapolation due to steep increase.

• Kohlrausch Law is used to determine $\Lambda_m^o$

• Degree of dissociation ($\alpha$) can be approximated by:
  $\alpha = \frac{\Lambda<em>m}{\Lambda</em>m^o}$

  • For weak electrolytes like acetic acid, the dissociation constant (Ka) is:
    $K<em>a = \frac{c\alpha^2}{1 - \alpha} = \frac{c(\frac{\Lambda</em>m}{\Lambda<em>m^o})^2}{1 - \frac{\Lambda</em>m}{\Lambda_m^o}}$

• Kohlrausch Law of Independent Migration of Ions:

• The limiting molar conductivity of an electrolyte is the sum of the individual contributions of the ions:
  \Lambdam^o (AB) = \lambda^oA^+ + \lambda^o_B^-

  • For an electrolyte giving v+ cations and v- anions:
    $\Lambda<em>m^o = v</em>+\lambda^o<em>+ + v</em>-\lambda^o<em>-$ Where $\lambda^o</em>+$ and $\lambda^o_-$ are the limiting molar conductivities of the cation and anion respectively.

• Applications of Kohlrausch Law:

  • Calculating °m for electrolytes.

  • Determining dissociation constants of weak electrolytes.

Electrolytic Cells and Electrolysis

• Electrolytic Cell: External voltage drives a chemical reaction.

• Electrolysis: Using an electrolytic cell to carry out chemical reactions.

• Simple Electrolytic Cell: Two copper strips in copper sulfate solution.

  • At Cathode: \Cu^{2+}(aq) + 2e^- \to Cu(s)

    • Copper deposits on the cathode.

  • At Anode: \Cu(s) \to Cu^{2+}(aq) + 2e^-

    • Copper dissolves (oxidizes) at the anode.

• Applications:

  • Industrial purification of copper.

  • Production of metals like Na, Mg, and Al by electrochemical reduction.

Quantitative Aspects of Electrolysis

• Faraday's Laws of Electrolysis:

  • First Law: The amount of chemical reaction is proportional to the quantity of electricity passed through the electrolyte.

  • Second Law: Amounts of different substances liberated by the same quantity of electricity are proportional to their equivalent weights.

• Quantity of Electricity (Q):

  • $Q = It$

    • Q is in coulombs (C), I is in amperes (A), and t is in seconds (s).

• Stoichiometry and Charge:

  • For the reaction: $Ag^+(aq) + e^- \to Ag(s)$

    • 1 mole of electrons is required for 1 mole of silver ions.

• Faraday (F):

  • The charge on one mole of electrons.

  • $F = N_A \times e = 6.02 \times 10^{23} mol^{-1} \times 1.6021 \times 10^{-19} C = 96487 C mol^{-1} \approx 96500 C mol^{-1}$

• For reactions like \Mg^{2+}(l) + 2e^- \to Mg(s), 2F are needed.

• For reactions like \Al^{3+}(l) + 3e^- \to Al(s), 3F are needed.

Products of Electrolysis

• Products depend on:

  • Nature of the material being electrolyzed.

  • Type of electrodes (inert or reactive).

• Inert electrodes (Pt, Au) do not participate in the reaction.

• Reactive electrodes participate in the electrode reaction.

• Electrolysis of Molten NaCl:

  • Products: Sodium metal and Cl2 gas.

  • Cathode: $Na^+ + e^- \to Na$

  • Anode: $Cl^- \to \frac{1}{2}Cl_2 + e^-$

• Electrolysis of Aqueous NaCl:

  • Products: NaOH, Cl2, and H2.

  • Cathode:

    • $H<em>2O(l) + e^- \rightleftharpoons \frac{1}{2}H</em>2(g) + OH^-$

  • Anode: $Cl^- \to \frac{1}{2}Cl_2 + e^-$

  • Net reaction:
    $NaCl(aq) + H<em>2O(l) \to Na^+(aq) + OH^-(aq) + \frac{1}{2}H</em>2(g) + \frac{1}{2}Cl_2(g)$

• Electrolysis of Sulfuric Acid:

  • Anode:

    • Dilute: $2H<em>2O(l) \rightleftharpoons O</em>2(g) + 4H^+(aq) + 4e^-$$</p></li><li><p>Concentrated:$2SO4^{2-}(aq) \rightleftharpoons S2O_8^{2-}(aq) + 2e^- $</p></li></ul></li></ul></li></ul><h4 id="4b814f82-4c7d-4616-88f8-f039816cdecc" data-toc-id="4b814f82-4c7d-4616-88f8-f039816cdecc" collapsed="false" seolevelmigrated="true">Batteries</h4><ul><li><p>Galvanic cells converting chemical energy into electrical energy.</p></li><li><p>Practical batteries are light, compact, and maintain voltage during use.</p></li></ul><h5 id="d6850692-c265-4c95-b089-12934896fdc1" data-toc-id="d6850692-c265-4c95-b089-12934896fdc1" collapsed="false" seolevelmigrated="true">Primary Batteries</h5><ul><li><p>Reactions occur only once; cannot be reused.</p></li><li><p><strong>Dry Cell (Leclanche Cell):</strong></p><ul><li><p>Anode: Zinc container.</p></li><li><p>Cathode: Carbon rod surrounded by MnO2 and carbon.</p></li><li><p>Electrolyte: Moist paste of NH4Cl and ZnCl2.</p></li><li><p>Anode:$Zn(s) \to Zn^{2+} + 2e^- $</p></li><li><p>Cathode:$MnO2 + NH4^+ + e^- \to MnO(OH) + NH_3$</p></li><li><p>Cell potential: ~1.5 V.</p></li></ul></li><li><p><strong>Mercury Cell:</strong></p><ul><li><p>Anode: Zinc-mercury amalgam.</p></li><li><p>Cathode: HgO and carbon paste.</p></li><li><p>Electrolyte: KOH and ZnO paste.</p></li><li><p>Anode:$Zn(Hg) + 2OH^- \to ZnO(s) + H_2O + 2e^- $</p></li><li><p>Cathode:$HgO + H_2O + 2e^- \to Hg(l) + 2OH^- $</p></li><li><p>Overall:$Zn(Hg) + HgO(s) \to ZnO(s) + Hg(l)$</p></li><li><p>Cell potential: ~1.35 V, constant during its life.</p></li></ul></li></ul><h5 id="0ee6d11d-f611-4845-bd48-46ea7ed832b6" data-toc-id="0ee6d11d-f611-4845-bd48-46ea7ed832b6" collapsed="false" seolevelmigrated="true">Secondary Batteries</h5><ul><li><p>Rechargeable by passing current in the opposite direction.</p></li><li><p><strong>Lead Storage Battery:</strong></p><ul><li><p>Anode: Lead.</p></li><li><p>Cathode: Lead grid packed with PbO2.</p></li><li><p>Electrolyte: 38$Pb(s) + SO4^{2-}(aq) \to PbSO4(s) + 2e^- $</p></li><li><p>Cathode:<br>$PbO2(s) + SO4^{2-}(aq) + 4H^+(aq) + 2e^- \rightleftharpoons PbSO4(s) + 2H2O(l) $</p></li><li><p>Overall:<br>$Pb(s) + PbO2(s) + 2H2SO4(aq) \rightleftharpoons 2PbSO4(s) + 2H_2O(l) $</p></li></ul></li><li><p><strong>Nickel-Cadmium Cell:</strong></p><ul><li><p>Longer life than lead storage cell but more expensive.</p></li><li><p>Overall:$Cd(s) + 2Ni(OH)3(s) \rightleftharpoons CdO(s) + 2Ni(OH)2(s) + H_2O(l) $</p></li></ul></li></ul><h4 id="c6f8e3d3-9909-481d-a51d-2d24460d5e7c" data-toc-id="c6f8e3d3-9909-481d-a51d-2d24460d5e7c" collapsed="false" seolevelmigrated="true">Fuel Cells</h4><ul><li><p>Convert fuel combustion energy directly into electrical energy.</p></li><li><p>Reactants are continuously fed, and products are removed.</p></li><li><p>Hydrogen-Oxygen Fuel Cell:</p><ul><li><p>Hydrogen and oxygen bubbled through porous carbon electrodes into NaOH solution.</p></li><li><p>Catalysts: Finely divided Pt or Pd.</p></li><li><p>Cathode:$O2(g) + 2H2O(l) + 4e^- \to 4OH^-(aq) $</p></li></ul></li><li><p>Anode:$2H2(g) + 4OH^-(aq) \to 4H2O(l) + 4e^- $</p><ul><li><p>Overall:$2H2(g) + O2(g) \to 2H_2O(l)$</p></li><li><p>Efficiency: ~70$2 Fe (s) \to 2 Fe^{2+} + 4 e^- $</p></li><li><p>Cathode:$O2(g) + 4 H^+(aq) + 4 e^- \to 2 H2O (l)$</p><ul><li><p>Ferrous ions oxidized to ferric ions, forming hydrated ferric oxide (rust):$Fe2O3. x H_2O$$

• Prevention:

  • Prevent surface contact with the atmosphere (paint, chemicals).

  • Covering the surface with other metals (Sn, Zn).

  • Sacrificial electrode (Mg, Zn) that corrodes instead of the object.