Overview — Module 5: Electrochemistry

Electrochemistry deals with the conversion of electrical energy into chemical energy and vice versa. Passing current through solutions or molten salts drives chemical change (electrolysis). Dry cells, button cells, and lead-acid batteries reverse the process — chemical reactions produce electricity. This lesson covers redox as electron transfer, oxidation numbers, balancing redox equations, electrolytic conduction and molar conductivity, Kohlrausch's law, Faraday's laws, electrolytic and galvanic cells, electrode potentials, electrochemical series, Nernst equation, batteries and fuel cells, ΔG–emf relation, and corrosion.

Mastering these ideas links L12 ionic equilibria (ions that move) with L9 thermodynamics (ΔG = −nFE°) and explains everyday technology from car batteries to rust prevention.

Section 1: Redox and Oxidation Number (13.1–13.3)

Oxidation: loss of electrons. Reduction: gain of electrons. They occur together as redox reactions. The species that loses electrons is the reductant; the species that gains electrons is the oxidant. Na → Na⁺ + e⁻ (oxidation); Cl + e⁻ → Cl⁻ (reduction).

Oxidation number (ON) is the apparent charge when shared electrons are assigned to the more electronegative atom. Rules: elemental form = 0; monatomic ion = charge; O usually −2 (peroxides −1, superoxides −½); H is +1 with non-metals, −1 with metals; alkali metals +1; sum of ON in neutral compound = 0; in polyatomic ion = ion charge.

Examples: S in H₂SO₄ is +6; N in NO₃⁻ is +5; Cl in ClO₄⁻ is +7. Balancing redox uses oxidation-number method or ion-electron method (half-reactions in acidic or basic medium). Example 13.1 balances skeletal equations by writing oxidation and reduction half-reactions, equalizing electrons, then combining.

In the oxidation-number method you identify atoms that change ON, find the total increase and decrease in ON, and multiply half-reactions so the electron loss equals electron gain. The ion-electron method is preferred for aqueous ionic reactions: write half-reactions, balance atoms (H₂O/H⁺ in acid; H₂O/OH⁻ in base), balance charge with electrons, then add. Both methods must give the same balanced equation — a useful check in exams.

Potassium permanganate (KMnO₄) in acid is a classic oxidant: Mn goes from +7 to +2 (Mn²⁺). Dichromate Cr₂O₇²⁻ goes from +6 to +3. Recognizing these common changes speeds up balancing and helps when writing half-cell reactions for electrochemical series problems later in the chapter.

Section 2: Electrolytic Conduction (13.4–13.5)

Electrolytes conduct electricity in solution by movement of ions. Solutions obey Ohm's law: V = IR. Conductance L = 1/R (siemens, S). Conductivity κ (kappa) = 1/ρ; κ = L × (l/A) where l/A is the cell constant.

Strong electrolytes (KCl) are fully dissociated — Λ_m rises gradually on dilution as inter-ionic forces weaken. Weak electrolytes (CH₃COOH) show sharp increase in Λ_m on dilution because degree of ionization α increases. Conductivity rises with temperature (~2–3% per degree) as viscosity falls and ion mobility increases.

13.5.1 Kohlrausch's Law

At infinite dilution, each ion contributes a fixed amount to molar conductivity independent of other ions: Λ_m° = ν₊λ₊° + ν₋λ₋°. This allows calculation of Λ_m° for weak electrolytes from strong ones.

Example 13.2: Λ°(NaCl)=126, HCl=426, CH₃COONa=91 → Λ°(CH₃COOH) = 91+426−126 = 391 S cm² mol⁻¹.

Table 13.1 compares conductivities: pure water is a very poor conductor (6×10⁻⁸ S cm⁻¹), 0.1 M HCl is much better (3.5×10⁻²), metals are orders of magnitude higher. Glass is an insulator (~10⁻¹⁴). This ranking explains why conductivity measurements distinguish electrolytes from non-electrolytes and strong from weak acids — exactly as in L12 ionic equilibrium.

Factors affecting conductivity: (i) nature of electrolyte — weak vs strong, ion charge, ion mobility; (ii) temperature — higher T lowers viscosity, speeds ions, and increases α for weak electrolytes; (iii) concentration — dilution decreases κ but increases Λ_m. Graph of Λ_m vs √c is nearly linear for strong electrolytes (Debye–Hückel–Onsager type behaviour at dilute limit) and curved for weak electrolytes.

Section 3: Electrolytic Cells and Faraday's Laws (13.6–13.7)

Electrolytic cells use electrical energy to drive non-spontaneous reactions. Anode is positive (oxidation); cathode is negative (reduction). Molten NaCl: Na⁺ + e⁻ → Na at cathode; 2Cl⁻ → Cl₂ + 2e⁻ at anode.

Faraday I: mass liberated ∝ quantity of electricity (w = zQ = zIt). Faraday II: when same charge passes different electrolytes in series, masses liberated ∝ chemical equivalent masses. Example 13.4: 32.4 g Ag deposited → copper deposited = 9.53 g (Eq Cu/Eq Ag ratio).

Products of electrolysis: Prefer process with easiest discharge. Aqueous NaCl: H₂ at cathode (water easier than Na⁺), Cl₂ at anode (O₂ overpotential). Aqueous CuSO₄ on Pt: Cu deposited, O₂ evolved. On Cu electrodes: refining — Cu dissolves at anode and deposits at cathode.

Discharge (decomposition) potential is the minimum voltage needed for an electrode process. Overpotential is the extra voltage often needed for gas evolution — H₂ overpotential is ~0 on Pt but ~1.5 V on mercury. That is why industrial NaOH manufacture uses a mercury cathode: Na⁺ is reduced instead of water, forming sodium amalgam, which then reacts with water to give NaOH and H₂. Exam questions love comparing molten vs aqueous NaCl and Pt vs Cu electrodes for CuSO₄.

Equivalent mass = molar mass / number of electrons lost or gained. For Ag⁺ + e⁻ → Ag, Eq = 108. For Cu²⁺ + 2e⁻ → Cu, Eq = 63.5/2 = 31.75. Faraday's second law then becomes a simple proportion when cells are in series — same Q, so m_A/m_B = Eq_A/Eq_B.

Section 4: Galvanic Cells and Electrode Potential (13.8–13.11)

Galvanic (voltaic) cells convert spontaneous redox into electricity. Classic Daniell cell: Zn rod in ZnSO₄ and Cu rod in CuSO₄, connected by a wire and salt bridge.

Anode (oxidation): Zn → Zn²⁺ + 2e⁻. Cathode (reduction): Cu²⁺ + 2e⁻ → Cu. Electrons flow external circuit from Zn to Cu. Salt bridge completes ionic circuit and prevents charge build-up without mixing solutions.

Cell notation: Zn | Zn²⁺(aq) || Cu²⁺(aq) | Cu · Double line = salt bridge · Left = anode · Right = cathode.

Standard electrode potential E° is measured against the Standard Hydrogen Electrode (SHE), assigned E° = 0. Electrodes with E° < 0 are better reductants than H₂; E° > 0 better oxidants.

Electrochemical series applications: (1) Predict redox — metal reduces ions of metals above it in series (more positive E°). (2) Calculate E°_cell. (3) Feasibility: positive E°_cell required. (4) Metals above H₂ liberate H₂ from acids. Example: Cu²⁺ + 2Ag → Cu + 2Ag⁺ has E° = −0.46 V — not feasible; reverse is spontaneous.

Cell emf is the potential difference measured in an open circuit (no current); under load with current flowing it is called potential difference and is lower due to internal resistance. Standard cell emf E°_cell requires both half-cells in standard states (1 M ions, 1 bar gases, pure solids). Measurement of unknown E° uses a galvanic cell with SHE: if the study electrode is positive, E° is positive; if negative (as for Zn), E° is negative.

Key E° values to remember: Zn²⁺/Zn −0.76 V, Fe²⁺/Fe −0.44 V, SHE 0, Cu²⁺/Cu +0.34 V, Ag⁺/Ag +0.80 V, Cl₂/Cl⁻ +1.36 V, F₂/F⁻ +2.87 V. Lithium (−3.05 V) is the strongest metallic reductant; fluorine is the strongest oxidant. Zn displaces Fe²⁺ (Example 13.5); Mg|Mg²⁺||Ag⁺|Ag gives E°_cell = 3.165 V.

Section 5: Nernst Equation (13.12)

Electrode potential depends on ion concentration. For Mⁿ⁺ + ne⁻ → M:

For cell reaction aA + bB → xX + yY: E_cell = E°_cell − (0.0591/n) log([X]^x[Y]^y/[A]^a[B]^b). Example 13.8: Ni|Ni²⁺(0.001)||Ag⁺(0.1)|Ag — first find E°_cell = 1.05 V, then apply Nernst with n=2.

Nernst equation shows how concentration cells and non-standard conditions shift emf — critical for batteries as they discharge and for corrosion under varying oxygen/pH conditions.

At 298 K the factor 2.303RT/F simplifies to 0.0591 V. For n = 2, each 10-fold change in Q shifts E by about 0.03 V. Only aqueous ions and gases (partial pressures) appear in Q — pure solids and pure liquids are taken as unit activity. Writing half-reactions first is the safest way to find n: the number of electrons cancelled when anode and cathode half-reactions are combined.

As a battery discharges, ion concentrations change, Q increases, and E_cell falls toward zero. When E_cell = 0 the cell is dead (equilibrium). Recharging a secondary cell drives the reverse electrolysis, restoring original concentrations and reactants — Faraday's laws then tell you how much charge is needed to restore a given mass of active material.

Section 6: Batteries, Fuel Cells, ΔG and Corrosion (13.13–13.15)

Primary cells (dry cell / Leclanché): Zn anode, graphite cathode, NH₄Cl–MnO₂ paste; ~1.5 V; not rechargeable. Secondary cells (lead storage): Pb anode, PbO₂ cathode, H₂SO₄ electrolyte; discharge forms PbSO₄; recharge reverses reaction. Fuel cells: continuous fuel feed; H₂–O₂ cell with KOH electrolyte produces water (Apollo program, E ≈ 0.9 V).

Gibbs energy: ΔG° = −nFE°. Negative ΔG° means spontaneous. Example 13.9: Daniell cell n=2, E°=1.1 V → ΔG° = −212.3 kJ.

Corrosion: deterioration of metals by environment. Iron rusts when one spot acts as anode (Fe → Fe²⁺) and another as cathode (O₂ reduction). Fe²⁺ oxidizes further to Fe₂O₃·xH₂O (rust). Accelerated by CO₂, SO₂ (acidic); needs moisture and oxygen. Protection: paint, grease, phosphate coatings, tinning, galvanizing (Zn sacrificial anode), impressed current.

Cathodic protection makes the entire metal surface the cathode of an electrochemical cell so oxidation of the structural metal is suppressed. Zinc coating on iron (galvanizing) works because Zn has more negative E° — zinc corrodes preferentially and protects iron even if the coating is scratched. Tin plating protects only while the coating is intact; if scratched, iron becomes the anode and rusts faster. Completely homogeneous iron rusts less — local anodic/cathodic sites need surface differences or impurities.

Fuel cells combine continuous fuel supply (like a power plant) with high efficiency of electrochemical conversion. The H₂–O₂ cell overall reaction is simply water formation; porous carbon electrodes with Pt/Pd catalyst and aqueous KOH electrolyte give ~0.9 V. Dry cells cannot be recharged because Zn is consumed irreversibly; lead-acid batteries reverse cleanly because PbSO₄ can be converted back to Pb and PbO₂ on charging.

Exam Connections and Chapter Summary

Key numerical skills: oxidation numbers; Λ_m and Kohlrausch; Faraday mass calculations; E°_cell from series; Nernst at non-standard concentrations; ΔG° = −nFE°. Conceptual: electrolytic vs galvanic polarity; products of electrolysis with overpotential; salt bridge role; ECS applications; battery types; corrosion mechanism.

Intext questions test: conduction factors, Faraday laws, cell notation, SHE, series feasibility, Nernst setup. Terminal exercises combine multi-step problems — e.g. calculate E_cell then ΔG, or mass deposited then series cell partner.

Electrochemistry unifies redox chemistry with electrical work: ions move (conductivity), charge transfers mass (Faraday), spontaneous redox drives emf (galvanic), concentration shifts potential (Nernst), and thermodynamics links energy (ΔG). From industrial electrolysis of NaCl to preventing ship hull corrosion, these principles power modern technology.

Module 5 path: chemical equilibrium → ionic equilibrium (L12) → electrochemistry (L13). Conductivity and Faraday connect to lab quantitative analysis; Nernst and ΔG connect to chemical thermodynamics; batteries and corrosion connect to materials science and engineering. When you see a redox problem, ask: which half is oxidized? What is n? Is E° positive? Does concentration change E via Nernst? That decision tree covers almost every NIOS numerical and conceptual question on this chapter.