Galvanic Cells & Electrode Potential (NTCH07/03) Why this matters Batteries power phones, remotes, and even EVs. Inside, chemistry pushes electrons through a wire. That is a galvanic (voltaic) cell: a spontaneous redox reaction creates electrical energy. NEET loves this chapter because one sign mistake flips the answer. Master anode/cathode roles, cell notation, E° values, and the link to ΔG° and K. Core terms An electrochemical cell in which a spontaneous redox reaction produces electrical energy (no external power). Galvanic (voltaic) cell Electrode where oxidation occurs. In a galvanic cell it is the negative terminal. Anode Electrode where reduction occurs. In a galvanic cell it is the positive terminal. Cathode Salt bridge A U-tube (often gel) containing an inert electrolyte (e.g., KCl, KNO3) that allows ions to flow to maintain charge neutrality, completing the internal circuit. Electrode potential Potential developed at an electrode due to a redox couple. By convention, tabulated as reduction potentials vs SHE. Electrode potential measured under standard conditions: all solutes 1 M, gases 1 atm, pure solids/liquids, 298 K, relative to SHE set as 0.00 V. Standard electrode potential (E°) Standard hydrogen electrode (SHE) Reference half-cell: Pt(s, platinized) | H2(g, 1 atm) | H+(aq, 1 M), E° = 0.00 V by convention. Electromotive force — the cell potential. Positive for a spontaneous galvanic cell in the forward direction. EMF (Ecell) Compact representation: Anode | Anode solution ‖ Cathode solution | Cathode. Single bar = phase boundary, double bar = salt bridge. Cell notation Charge per mole of electrons, F = 96485 C mol⁻¹. Faraday constant (F) Galvanic vs electrolytic: the big picture In a galvanic cell, chemistry drives electricity: oxidation at anode sends electrons through the external wire to the cathode, where reduction happens. In an electrolytic cell, an external power source pushes a non-spontaneous reaction. Polarity flips between these two, which often causes confusion — so keep the contrast handy. Side-by-side contrast: galvanic (spontaneous, anode negative, cathode positive) vs electrolytic (non-spontaneous, anode positive, cathode negative). Track electron flow in the external circuit and ion flow through the electrolyte/salt bridge. The anode is always negative and the cathode always positive. In a galvanic cell, anode is negative and cathode is positive. In an electrolytic cell (with external DC supply), the anode is positive and the cathode is negative — the polarities are reversed. The Daniell cell: your model galvanic cell Classic setup: zinc electrode dipped in 1 M ZnSO4 solution and copper electrode dipped in 1 M CuSO4 solution, connected by a salt bridge (KCl/KNO3) and an external wire with a voltmeter. Oxidation happens at zinc (anode), reduction at copper (cathode). Part What it is Role/sign Daniell cell components and roles Component Anode (left) Zn(s) in 1 M ZnSO4(aq) Oxidation; negative terminal; source of electrons Cathode (right) Cu(s) in 1 M CuSO4(aq) Reduction; positive terminal; accepts electrons Salt bridge KCl/KNO3 (gel or U-tube) Maintains electrical neutrality: anions to anode, cations to cathode External circuit Wire + voltmeter/load Electrons flow from anode (Zn) to cathode (Cu) Standard Daniell cell: label the anode (Zn), cathode (Cu), salt bridge, electron flow (Zn → Cu externally), and ion flow (Cl⁻ to anode, K⁺ to cathode). Overall Daniell redox: Zn is oxidized, Cu²⁺ is reduced. Electrons released at anode travel through the wire to the cathode. Half-reactions in Daniell cell Overall reaction: Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s). If standard concentrations (1 M), this is a standard cell. Its standard EMF comes from tabulated E° values. Cell notation: quick and clean Write anode (oxidation) on the left, cathode (reduction) on the right. Single vertical bar | = phase boundary; double vertical bar ‖ = salt bridge. Species in the same phase are separated by commas; inert conductors (Pt) are written when needed. Example Daniell cell: Zn(s) | Zn2+(1 M) ‖ Cu2+(1 M) | Cu(s). Electrons are NOT written in cell notation. Rules you must memorize Standard electrode potential and SHE We cannot measure the absolute potential of a single electrode, only differences. So all half-cells are compared against the standard hydrogen electrode (SHE): Pt(s, platinized) | H2(g, 1 atm) | H+(aq, 1 M), assigned E° = 0.00 V. By convention, tables list all standard potentials as reduction potentials. 2026-05-26T17:04:38.446Z Diagram of the standard hydrogen electrode: beaker with 1 M acid, platinized platinum electrode connected to an external circuit, H2 gas inlet at 1 atm bubbling over Pt surface. Labels: Pt (platinized), H2(1 atm), [H+] = 1 M, 298 K, E° = 0.00 V. Clean 2D vector, arrows in red. SHE construction: platinized Pt electrode dipping in 1 M H+; H2 gas bubbled at 1 atm; temperature 298 K; potential defined as 0.00 V. gpt-image-2 Combine two reduction potentials Use reduction potentials for both half-cells; subtract anode (written as reduction) from cathode. For Daniell: E°(Cu2+/Cu) = +0.34 V, E°(Zn2+/Zn) = −0.76 V. So E°cell = 0.34 − (−0.76) = +1.10 V (spontaneous). Electrochemical series: who oxidizes/reduces whom? Order half-cells by E° (reduction). More positive E° = stronger oxidizing agent (as oxidant on left gets reduced). More negative E° = stronger reducing agent (as metal on right tends to oxidize). Predict displacement reactions and feasibility quickly from this ladder. F2(g) + 2e− → 2F−(aq) +2.87 MnO4−(aq) + 8H+(aq) + 5e− → Mn2+(aq) + 4H2O(l) +1.51 Au3+(aq) + 3e− → Au(s) +1.50 Cr2O7 2− (aq) + 14H+(aq) + 6e− → 2Cr3+(aq) + 7H2O(l) +1.33 O2(g) + 4H+(aq) + 4e− → 2H2O(l) +1.23 Ag+(aq) + e− → Ag(s) +0.80 Cu2+(aq) + 2e− → Cu(s) +0.34 2H+(aq) + 2e− → H2(g) (SHE) 0.00 Sn2+(aq) + 2e− → Sn(s) −0.14 Fe2+(aq) + 2e− → Fe(s) −0.44 Zn2+(aq) + 2e− → Zn(s) −0.76 Al3+(aq) + 3e− → Al(s) −1.66 Mg2+(aq) + 2e− → Mg(s) −2.37 Na+(aq) + e− → Na(s) −2.71 Li+(aq) + e− → Li(s) −3.05 Half-reaction (reduction direction) E° (V) Selected standard reduction potentials at 298 K (vs SHE) 2026-05-26T17:04:38.386Z Vertical ladder chart with half-reactions placed by E°. Top: F2/2F− (+2.87 V), then MnO4−/Mn2+, Au3+/Au, O2/H2O, Ag+/Ag, Cu2+/Cu, SHE at 0.00; below: Sn2+/Sn, Fe2+/Fe, Zn2+/Zn, Al3+/Al, Mg2+/Mg, Na+/Na, Li+/Li. Clean vector, E° labeled, arrows showing oxidizing/reducing strength. Electrochemical series ladder: top (most positive E°) are strong oxidants like F2; bottom (most negative E°) are strong reductants like Li(s). gpt-image-2 Spontaneity, ΔG°, and equilibrium K A positive E°cell means the forward reaction is spontaneous under standard conditions. The energy link is direct: electrical work equals decrease in Gibbs energy. From ΔG° you can jump to the equilibrium constant K — very testable. General (non-standard) relation between Gibbs energy change and cell potential. At 298 K, relates standard EMF to equilibrium constant. From G ° = -nFE ° = -RT K. Daniell example (n = 2, E°cell = +1.10 V): ΔG° = −nFE° = −2 × 96485 × 1.10 ≈ −2.12 × 10 5 J mol⁻¹ (≈ −212 kJ mol⁻¹). K = 10 (nE°/0.0592) ≈ 10 (2.2/0.0592) ≈ 10 37.2 — enormously product-favored. neet-alert Always use reduction potentials to compute E°cell. Do not change signs when multiplying half-reactions to balance electrons — E° values are intensive, they do not scale with stoichiometry. Preview: concentration matters (Nernst form) Under non-standard conditions, the cell potential depends on reaction quotient Q. You will derive and use this fully in the next subtopic, but keep the forms handy for quick checks. Identify anode/cathode fast Decision steps Write both half-cells as reductions with their E° (from the series). The half-cell with more positive E° acts as cathode (reduction). The other becomes anode (oxidation). Compute E°cell = E°cathode − E°anode. If > 0, the written cell is spontaneous in that direction. Write the overall reaction by reversing the anode half-reaction and adding electrons to cancel. AnOx, RedCat: Anode = Oxidation; Reduction = Cathode. In galvanic cells: Electron flow is A → C; Conventional current is C → A (externally). No. Only ions move through the salt bridge to keep solutions electrically neutral. Electrons travel through the external metallic circuit. Electrons flow through the salt bridge. Ion traffic in a galvanic cell: anions migrate into the anode half-cell, cations into the cathode half-cell through the salt bridge. Close-up schematic of salt bridge between Zn2+ and Cu2+ solutions. Show Cl− moving to anode, K+ to cathode, while electrons move in external wire from Zn to Cu. Clear labels; vector style; arrows in red. 2026-05-26T17:04:40.164Z gpt-image-2 Industrial relevance and real-world links remember Galvanic cells underlie all batteries, from alkaline AA cells to lithium-ion packs in phones and EVs. Electrolytic cells drive electroplating (gold, chrome) and metal purification (Cu), and large-scale production (e.g., Al). Electroplating is an electrolytic process (requires a DC power source). Recognize the polarity flip versus a galvanic cell. Corrosion control uses sacrificial anodes: attach Zn or Mg to an iron pipeline. The more negative E° metal oxidizes first, protecting iron from rusting. Pipeline cross-section with attached Zn block labeled 'sacrificial anode'. Show galvanic coupling: electrons from Zn to Fe, oxygen/water at Fe surface reduced, Zn dissolves as Zn2+. Include E° values: Zn2+/Zn (−0.76 V) vs Fe2+/Fe (−0.44 V). Clean vector diagram. Sacrificial-anode protection: Mg/Zn bolted to an iron pipeline corrodes preferentially, keeping Fe in its metallic state. 2026-05-26T17:04:40.214Z gpt-image-2 Worked predictions: displacement and feasibility Quick checks using E° Will Zn(s) displace Cu2+ from CuSO4? Yes. E°cell = 0.34 − (−0.76) = +1.10 V → spontaneous. Reaction: Zn + Cu2+ → Zn2+ + Cu. Can Ag(s) displace H+ from acid to make H2? Compare Ag+/Ag (+0.80) vs H+/H2 (0.00). E°cell = 0.00 − 0.80 = −0.80 V for Ag + H+ → Ag+ + H2 (non-spontaneous). So no. Will Fe(s) displace Cu2+? E°cell = 0.34 − (−0.44) = +0.78 V → yes. Expect copper deposition and iron(II) formation. tip When choosing a salt bridge electrolyte, pick ions that do not react with either half-cell (KCl/KNO3 are common). Precipitation or complexation can ruin the cell. Conventional current flows from cathode to anode in the external circuit for a galvanic cell (opposite to electron flow). In circuit diagrams, arrows typically show conventional current. neet-alert Electrons flow from anode to cathode externally. Conventional current direction is opposite to electron flow: from cathode to anode in a galvanic cell. Current always flows from anode to cathode externally. Fuel cell preview (application of galvanic concept) Preview only; detailed treatment later. Galvanic cell where H2 is oxidized and O2 is reduced to H2O with continuous reactant feed. gpt-image-2 Simple schematic of H2–O2 fuel cell: separate compartments with gas feeds, electrodes, and ion-conducting medium showing oxidation of H2 and reduction of O2. Fuel cell diagram: left anode fed with H2, right cathode fed with O2/air, electrolyte separating. Show half-reactions and electron flow through load. Vector style, labels only, no heavy math. 2026-05-26T17:04:39.026Z