OPOpenPithMake curiosity visibleChemistry lab / Cells and electrolysis

About this lab

Seven topics, one thread: where the electrons come from and where they go.

In a galvanic cell a spontaneous redox reaction turns chemical energy into electrical energy: the more active electrode loses electrons, and they flow through the wire to the other. In an electrolytic cell a power supply pushes electrons into the cathode and pulls them from the anode, driving a reaction that would not go by itself. In both, the solution and the salt bridge carry only ions and the wire carries only electrons; the amount of product follows the electrons transferred: n = Q / (zF).

Every reading, every bubble and every colour comes from one model, not a pre-drawn animation. Change one condition and all the data are recalculated together.

What the seven topics show

Seven common ideas, tested with data

01 Electrolysis of water: “There should be as much hydrogen as oxygen”

At 0.40 A for 8 minutes, 192 C passes: 24.3 mL of hydrogen at the cathode and 12.2 mL of oxygen at the anode (25 °C, 101 kPa). Each O₂ needs 4 electrons and each H₂ only 2, so the volume ratio is 2 : 1. With pure water, even at the 12 V limit the current is only 0.051 mA.

02 Copper–zinc cell: “Bubbles on the copper mean the copper is reacting”

With the zinc 3 cm deep the emf is 0.76 V and the current 99 mA; in 5 minutes 11.1 mg of zinc dissolves and 3.8 mL of hydrogen forms on the copper. The bubbles are H⁺ from the solution gaining electrons on the copper; the copper itself is unchanged. Magnesium gives 2.37 V and 309 mA; iron gives only 57 mA.

03 Salt bridge: “Electrons flow through the salt bridge”

With both beakers at 1 mol/L the emf is 1.100 V and the current 6.11 mA. Dilute the Zn²⁺ to 0.01 mol/L and the emf rises to 1.159 V. Only K⁺ and Cl⁻ move in the bridge; pull it out and the current is zero at once. Practical cells follow the same idea: alkaline 1.5 V, lead–acid 2.04 V, and a hydrogen–oxygen fuel cell 1.23 V in theory.

04 Electrolysis: “Whatever ions are in solution come out”

At 0.30 A for 10 minutes, 180 C passes. CuCl₂ solution gives 59.3 mg of copper at the cathode and 22.8 mL of chlorine at the anode. NaCl solution gives 22.8 mL of hydrogen at the cathode: Na⁺ does not discharge. Na₂SO₄ solution gives 22.8 mL of hydrogen and 11.4 mL of oxygen: really the electrolysis of water. The products depend on both the discharge order and the electrode material.

05 Chlor-alkali: “The membrane just keeps the two gases apart”

At 2 A for 20 minutes with a cation-exchange membrane: 304 mL of chlorine and 24.9 mmol of NaOH. Without the membrane OH⁻ reaches the anode region: only 122 mL of chlorine is collected, 7.5 mmol of NaClO forms, and the NaOH yield drops to 40 %.

06 Plating and refining: “Refining, like plating, leaves the solution unchanged”

At 0.50 A for 20 minutes, 600 C passes: the key gains 197.6 mg of copper, about 22.1 µm thick, and the anode loses the same. In refining, the crude anode loses 198.4 mg, more than the cathode gains: zinc and iron enter the solution, silver and gold fall as 0.99 mg of anode slime, and c(Cu²⁺) falls. The same charge would deposit 670.8 mg of silver.

07 Corrosion: “Rusting is only about water”

Over six simulated hours: in dilute hydrochloric acid, hydrogen-evolution corrosion pushes the U-tube level on the tube side down 68.5 mm; in salt water, oxygen-absorption corrosion pulls it up 23.0 mm, and the iron loses 1.25 mg. With a zinc block the iron loses nothing and the zinc loses 2.20 mg; with an applied current the iron stops dissolving too.

Model and sources

Model and sources

  • Faraday’s law: Q = It, n(e⁻) = Q / F, F = 96485 C/mol; product n = Q / (zF), where z is the electrons transferred per particle.
  • Gas volume: V = nRT / p at 25 °C and 101.325 kPa, Vm = 24.46 L/mol.
  • emf: E = E°(positive) − E°(negative); concentration correction by the Nernst equation E = E° − (RT / zF)·ln Q.
  • Circuit: I = E / (R_ext + R_int); in the single-liquid cell R_int = L / (κA) changes with immersion depth.
  • Discharge order: at the cathode Ag⁺ > Cu²⁺ > H⁺ > Pb²⁺ > Fe²⁺ > Zn²⁺ > H₂O > Al³⁺ > Mg²⁺ > Na⁺ > K⁺; at the anode first check whether the electrode is active; on an inert electrode S²⁻ > I⁻ > Br⁻ > Cl⁻ > OH⁻ > oxoanions.
  • Corrosion manometer: (n₀ + Δn)RT = (p₀ + ρg·2x)(V₀ + a·x), level difference Δh = 2x.

Standard electrode potentials are the usual textbook values (Zn −0.76 V, Fe −0.44 V, Mg −2.37 V, Cu 0.34 V, O₂ 1.23 V). The following are teaching assumptions, used only to compare sizes: a decomposition voltage of 1.83 V and a cell resistance of 12 Ω (pure water 2×10⁵ Ω); a 5 Ω external circuit, 20 S/m acid conductivity and no polarisation in the single-liquid cell; the direct reaction rates of Mg, Zn and Fe with acid; an 80 Ω bridge and 100 Ω load; the internal resistance of the practical cells; 60 % of the NaOH mixing with chlorine when there is no membrane; the crude-copper composition; a 10 cm² key; corrosion currents of 0.3 mA and 0.2 mA and the zinc and applied-current factors; chlorine dissolving in water is ignored. Colours and ion beads are for legibility; their numbers scale with the model’s amounts.

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