Electrolysis, Batteries & Fuel Cells Big picture: moving between chemical and electrical energy Electrolysis forces a chemical reaction using an external power supply (non-spontaneous). Batteries and fuel cells do the reverse: they let a spontaneous chemical reaction run and deliver electrical energy to your circuit. This chapter helps you quickly decide what gets deposited or liberated during electrolysis, calculate “how much” using Faraday’s laws, identify common batteries in daily life, make sense of the H2–O2 fuel cell, and explain rusting of iron as an electrochemical process. Electrolysis vs battery: external source drives electrolysis; a battery generates power from a spontaneous redox. Spot where energy enters or leaves the system. Real-life: Electrolysis enables electroplating, copper electrorefining, chlor-alkali (Cl2, NaOH, H2) and aluminium extraction. Batteries power phones, laptops, cars, medical pacemakers, and EVs. Fuel cells run spacecraft and zero-emission vehicles. remember Electrolytic cell basics In an electrolytic cell, an external DC source pushes electrons into one electrode (cathode) and pulls electrons from the other (anode), forcing a non-spontaneous redox reaction. Think of the power supply as a “pump” that drives electrons against the natural tendency. - Oxidation occurs at the anode; reduction occurs at the cathode (AnOx, RedCat — always true). - Polarity flips compared to a galvanic cell: in an electrolytic cell, anode is positive, cathode is negative; in a galvanic cell, anode is negative, cathode is positive. Electrolysis of water: O2 at the anode (+), H2 at the cathode (−). Gas volumes show a 1:2 ratio (O2:H2). Polarity reverses: in galvanic cells anode is negative and cathode positive; in electrolytic cells anode is positive and cathode negative. Polarity of anode and cathode is the same in all cells. Faraday’s 1st law (Z form) Mass (m) deposited is proportional to charge passed (Q = It). Z is the electrochemical equivalent. Faraday’s 1st law (equivalent mass form) E is equivalent mass (atomic or formula mass divided by valency), F is 96485 C mol -1 . Predicting products of electrolysis (molten vs aqueous; inert vs active) Key idea: At the cathode, the species easiest to reduce (highest reduction potential under the conditions) is chosen; at the anode, the species easiest to oxidize is chosen. In aqueous solutions, water can also be reduced or oxidized, often beating metal ions with very negative E° values (like Na+). Overpotential can tilt outcomes, especially for O2 evolution. - Molten salts: only the cation and anion are present; products are usually the metal at cathode and the nonmetal at anode. - Aqueous solutions: compare the ion’s E° with water’s reduction/oxidation; also consider concentration (e.g., concentrated brine favors Cl2 evolution at anode). Case Electrolysis outcomes you must know Electrolyte (electrodes) Cathode product Anode product Reason/Note Molten NaCl (graphite/Pt) Na(l) Cl2(g) Only Na+ and Cl− present; no water. Na+ reduced; Cl− oxidized. Aqueous NaCl (concentrated brine, inert electrodes) H2(g) Cl2(g) H2O reduction (E° −0.83 V) is easier than Na+ → Na (−2.71 V). Cl2 forms at anode in concentrated brine due to kinetics/overpotential; NaOH remains in solution. Aqueous NaCl (very dilute, inert electrodes) H2(g) O2(g) (often) At low [Cl−], water oxidation can dominate (O2) because Cl− is scarce and O2 overpotential varies with electrode. CuSO4(aq), inert electrodes (Pt/C) Cu(s) O2(g) Cu2+ reduced at cathode; water oxidized at anode: 2H2O → O2 + 4H+ + 4e−. CuSO4(aq), Cu electrodes (electrorefining) Cu(s) deposited (pure) Cu(s) dissolves Net: copper transfers from anode (impure) to cathode (pure); anode slime collects Ag/Au. Chlor-alkali: Concentrated brine electrolysis yields Cl2 (anode), H2 (cathode), and NaOH(aq). Industrial chlor-alkali membrane cell schematic. Show anode chamber with concentrated brine, Cl2 bubbles at Ti/RuO2 anode; cation-exchange membrane; cathode chamber with Na+ migrating, H2 evolution at steel cathode, NaOH formation. Clear flow arrows and labels. Clean vector style. Chlor-alkali membrane cell: Cl2 at anode, H2 at cathode, NaOH collected in catholyte; brine feed and ion-selective membrane shown. 2026-05-26T17:04:41.191Z gpt-image-2 In any aqueous electrolysis, water must be the species reduced or oxidized. Not always. Competing ions can discharge preferentially based on standard potentials and overpotential. Example: concentrated brine gives Cl2 at anode rather than O2. Electrorefining and electroplating (copper as the model) Electrorefining of copper uses impure Cu slabs as anodes and thin pure Cu sheets as cathodes in CuSO4/H2SO4. At the anode, Cu dissolves: Cu → Cu2+ + 2e−. At the cathode, Cu2+ + 2e− → Cu is deposited as high-purity copper. Precious metals (Ag, Au) do not dissolve readily and settle as anode slime, later recovered. Electroplating is similar: the object to be plated is the cathode; a solution containing the metal ion to coat is the electrolyte; a soluble anode of that metal helps maintain ion concentration. gpt-image-2 Diagram of a Cu electrorefining tank with impure copper anodes, pure copper cathode sheets, CuSO4/H2SO4 electrolyte, electron flow through external circuit, and labeled anode slime tray. Vector style, clear oxidation/reduction labels. Copper electrorefining cell: impure Cu anode dissolves; pure Cu grows on cathode; anode slime collects Ag/Au. 2026-05-26T17:04:41.221Z Faraday’s laws — how much gets deposited or liberated? Faraday’s 1st law: mass of a substance deposited or liberated at an electrode is directly proportional to the charge passed. Faraday’s 2nd law: for the same charge passed through different electrolytes, masses deposited are proportional to their equivalent masses (E = molar mass/valency). Q is charge (C), n is moles of electrons, F is Faraday constant (96485 C mol -1 ). Charge–moles of electrons link The Faraday constant One Faraday equals charge carried by one mole of electrons. Ag+ (1) 107.87 107.87 Cu2+ (2) 31.78 31.78 Al3+ (3) 9.00 9.00 Substance Electrolyte (ion, valency) Equivalent mass E (g eq -1 ) Mass deposited by 1 F (g) Quick results for 1 F of charge Worked numericals (NEET-style) Mass of Cu from CuSO4(aq): Current 2.0 A for 30.0 min. Q = It = 2.0 × 1800 = 3600 C. E(Cu2+) = 63.55/2 = 31.78 g eq−1. m = (EIt)/F = 31.78 × 3600 / 96485 ≈ 1.19 g of Cu. Volume of H2 at STP from 9650 C in acid solution: n(e−) = Q/F ≈ 0.100 mol. 2 e− → 1 H2, so n(H2) = 0.050 mol. Volume ≈ 0.050 × 22.4 L = 1.12 L (older STP); ≈ 0.050 × 22.7 L = 1.14 L (IUPAC). Exams often use 22.4 L mol−1. Same charge through AgNO3 and CuSO4: mAg/mCu = EAg/ECu = 107.87 / 31.78 ≈ 3.39. If 2.00 g Cu deposits, Ag would be ≈ 6.78 g. Traps: Don’t forget valency in E = M/valency; include n in Q = nF for gas volumes; check whether water competes in aqueous electrolysis; and watch 22.4 vs 22.7 L mol -1 at STP. neet-alert Batteries — primary (non-rechargeable) vs secondary (rechargeable) Primary cells are discarded after use (chemistry not reversible easily). Secondary cells can be recharged by applying reverse current. Know the chemistry and EMF ranges of common cells used in daily life and exams. Type Anode (discharge) Cathode (discharge) Electrolyte EMF (per cell) Typical uses Common batteries for NEET Battery Dry cell (Leclanché, primary) Zn → Zn2+ + 2e− (complexed as [Zn(NH3)4]2+) 2MnO2 + 2NH4+ + 2e− → Mn2O3 + 2NH3 + H2O NH4Cl/ZnCl2 paste ≈ 1.5 V Torches, remotes Alkaline MnO2 (primary) Zn + 2OH− → ZnO + H2O + 2e− MnO2 + H2O + e− → MnOOH + OH− (approx.) KOH ≈ 1.5 V High-drain gadgets Mercury cell (primary) Zn(Hg) + 2OH− → ZnO + H2O + 2e− HgO + H2O + 2e− → Hg + 2OH− KOH ≈ 1.35 V (flat) Watches, instruments Lead–acid (secondary) Pb + SO4 2− → PbSO4 + 2e− PbO2 + SO4 2− + 4H+ + 2e− → PbSO4 + 2H2O H2SO4(aq) ≈ 2.0 V Car starter batteries (≈12 V: 6 cells) Ni–Cd (secondary) Cd + 2OH− → Cd(OH)2 + 2e− NiO(OH) + H2O + e− → Ni(OH)2 + OH− KOH ≈ 1.2 V Cordless tools Li-ion (secondary) C6Li → C6 + Li+ + e− (graphite anode) Li1−xCoO2 + xLi+ + xe− → LiCoO2 Li+ salt in organic solvent ≈ 3.6–3.7 V Phones, laptops, EV packs gpt-image-2 Cross-sectional diagram of a 6-cell lead–acid car battery. Show alternating Pb and PbO2 plates with separators, sulfuric acid electrolyte, and arrows indicating discharge (Pb to PbSO4, PbO2 to PbSO4) and reverse on charging. Clean vector style. Lead–acid battery cross-section: Pb (anode) and PbO2 (cathode) plates in H2SO4, separators, terminals, and discharge/charge directions. 2026-05-26T17:04:41.854Z Remember nominal EMFs: Dry/alkaline ~1.5 V, Mercury ~1.35 V, Lead–acid ~2.0 V per cell, Ni–Cd ~1.2 V, Li-ion ~3.6–3.7 V. neet-alert Fuel cells — continuous fuel supply, high efficiency A fuel cell converts chemical energy of a continuously supplied fuel and oxidant directly into electricity. For the hydrogen–oxygen fuel cell, either alkaline or PEM operation is used. Alkaline fuel cell (AFC): - Anode: H2 + 2OH− → 2H2O + 2e− - Cathode: O2 + 2H2O + 4e− → 4OH− Overall: 2H2 + O2 → 2H2O; E°cell ≈ 1.23 V at 298 K. It runs quietly, with high efficiency and water as the only product (clean). PEM cells run similarly but with proton conduction in an acidic membrane. Standard EMF of H2–O2 fuel cell At 298 K for liquid water formation. Hydrogen–oxygen fuel cell: H2 oxidized at anode, O2 reduced at cathode; water forms; continuous fuel feed. gpt-image-2 2026-05-26T17:04:43.035Z Diagram of a PEM fuel cell stack layer: H2 flow field to Pt anode, proton-exchange membrane, O2/air flow field to Pt cathode. Show electron path in external circuit, proton path through membrane, and net water generation. Clean 2D vector style. PEM fuel cell schematic: H2 fed to anode (H+ to membrane, e− to circuit), O2 to cathode, water out; labels for anode/cathode catalysts and proton-exchange membrane. Corrosion of iron — an electrochemical process On a moist iron surface, tiny anodic and cathodic regions form. At the anode: Fe → Fe2+ + 2e−. Electrons flow to a nearby cathodic site where dissolved oxygen is reduced: O2 + 2H2O + 4e− → 4OH−. The Fe2+ and OH− form Fe(OH)2, which further oxidizes to hydrated iron(III) oxide, Fe2O3·xH2O (rust). Both O2 and H2O are needed, explaining why dry or oxygen-free conditions prevent rusting. gpt-image-2 Annotated surface diagram showing an anodic pit (Fe → Fe2+ + 2e−), electron flow through metal to a cathodic area (O2 + 2H2O + 4e− → 4OH−), diffusion of ions, and formation of Fe(OH)2 → Fe2O3·xH2O. Use arrows and labels, clean vector style. Rusting map on an iron surface: anodic pits where Fe dissolves, nearby cathodic regions reducing O2, and rust deposit forming. 2026-05-26T17:04:43.255Z Method How it protects Examples/Notes Preventing corrosion — methods and principles Technique Galvanization (Zn coating) Physical barrier; Zn acts as sacrificial anode if scratched Iron pipes, roofing sheets Cathodic protection (sacrificial anode) Attach a more active metal so it oxidizes instead of iron Zn/Mg blocks on ship hulls, pipelines Tinning, painting, polymer coatings Barrier to O2/H2O Food cans (Sn), paint layers Alloying/passivation Stable oxide film resists attack Stainless steel forms Cr2O3 film Choose sacrificial anode more reactive than Fe (e.g., Mg, Zn). Replace it periodically as it corrodes. tip Industrial highlights to connect Chlor-alkali (brine electrolysis): Cl2(g), H2(g), NaOH(aq) — membrane cells dominate. Aluminium extraction (Hall–Héroult): Electrolysis of molten Al2O3 dissolved in cryolite (Na3AlF6); carbon anodes consumed. Copper electrorefining: High-purity Cu and recovery of precious metals from anode slime. Electroplating: Au, Ag, Ni, Cr for protection and appearance. Batteries for transport and electronics: lead–acid car starters; Li-ion for phones and EVs. Fuel-cell vehicles: H2–O2 fuel cells produce electricity on-board with water as exhaust. 2026-05-26T17:04:43.267Z Hall–Héroult cell: molten Al2O3 in cryolite bath, carbon anodes (CO2 off-gas), Al(l) collected at cathode (carbon-lined). Cutaway of an aluminium electrolysis pot: carbon cathode lining, molten cryolite with dissolved Al2O3, carbon anodes dipping from top releasing CO2, molten Al pool at bottom with tap. Labels and arrows; vector technical style. gpt-image-2 Electrolytic vs galvanic: side-by-side Galvanic cell: spontaneous reaction generates electrical energy (chemical → electrical). Electrolytic cell: electrical energy is consumed to drive a non-spontaneous reaction (electrical → chemical). Both follow AnOx and RedCat; the sign of electrodes differs. Daniell (galvanic) cell: chemical energy → electrical energy; anode(−): Zn | Zn2+, cathode(+): Cu2+ | Cu, salt bridge completes the circuit. Used for galvanic cells and to estimate fuel-cell EMF under standard conditions. Cell EMF relation (recall) External EMF drives non-spontaneous redox; product depends on species and conditions. Decarboxylative dimerization of carboxylate anions at the anode (organics link to electrochemistry). Electrolysis is simply the reverse of a galvanic cell in every respect. Structurally, yes (energy direction reverses), but remember: in electrolytic cells anode is positive and cathode negative, opposite to galvanic cells. Reactions are also constrained by overpotentials and medium. AnOx, RedCat always; for polarity: Electrolytic has Positive Anode (EP-A), Galvanic has Negative Anode (GN-A). Electrolytic cell Cell where an external EMF drives a non-spontaneous redox; anode is positive, cathode is negative. Electrolysis Chemical decomposition by passing electric current through an electrolyte. Faraday’s laws Quantitative laws relating mass deposited to charge (1st law) and to equivalent mass (2nd law). Charge per mole of electrons: 96485 C mol−1. Faraday constant (F) Primary cell Non-rechargeable battery; chemistry not reversed easily (e.g., dry cell). Rechargeable battery; chemistry is reversible by applying external current (e.g., lead–acid, Li-ion). Secondary cell Fuel cell Device generating electricity from continuous supply of fuel and oxidant (e.g., H2–O2). Polymer membrane that conducts H+ in PEM fuel cells. PEM (Proton-exchange membrane) Electrochemical deterioration of metals; for iron, rust forms as Fe2O3·xH2O. Corrosion A more active metal attached to protect iron by preferentially oxidizing (e.g., Zn, Mg). Sacrificial anode Galvanization Coating iron with zinc to prevent corrosion. Purifying a metal (e.g., Cu) by electrolytic dissolution at anode and deposition at cathode. Electrorefining Chlor-alkali process Industrial electrolysis of brine to produce Cl2, H2, and NaOH. Key terms at a glance