Chemistry · Redox reactions and electrochemistry · Grade 9

Electrolysis and corrosion

Current turns table salt into chlorine and alkali — so why does an iron pipe rust?

The situation

At the chlor-alkali shop of Navoiyazot, electric current is used to obtain chlorine, hydrogen and caustic soda from a solution of table salt. The same flow of electrons works in a household water pipe too — only it does no good: the iron slowly turns into rust. If current is passed through copper chloride solution in a school laboratory, one electrode turns red and the other gives off a gas with a sharp smell. What happens at the electrodes, how does a chemical reaction produce current, and how can iron be protected from rust?

Topic

Electrolysis and corrosion. Navoiyazot chlor-alkali shop and a rusty water pipe: a flow of electrons gives chlorine, hydrogen and caustic soda from brine, and eats away iron

Key ideas

Electrolysis of solutions and melts
Galvanic cell
Protection against corrosion

Formulas

CuCl_{2} \overset{}{\to} Cu + Cl_{2} ↑

Electrolysis of copper(II) chloride solution: copper at the cathode, chlorine at the anode (lab work 7).

2 NaCl + 2 H_{2} O \overset{}{\to} 2 NaOH + H_{2} ↑ + Cl_{2} ↑

Table salt solution: hydrogen and caustic soda at the cathode, chlorine at the anode (chlor-alkali production).

m = \frac{M I t}{z F}, F = 96 485 C/mol

Faraday's laws: the mass of substance deposited is proportional to the charge ( $q = I t$ ) and to $M / z$ .

ℰ = E_{cathode}^{\circ} - E_{anode}^{\circ}

Emf of a galvanic cell: for Zn–Cu (the Daniell cell) $0.34 - (- 0.76) = 1.10$ V.

4 Fe + 3 O_{2} + 6 H_{2} O \overset{}{\to} 4 Fe (OH)_{3}

Rusting of iron: when water and oxygen are present together; rust is hydrated iron(III) oxide.

Theory

Electrolysis: processes at the cathode and anode

Electrolysis is the redox process at the electrodes when direct current passes through an electrolyte solution or melt. At the cathode, connected to the negative pole of the source, cations (or water) gain electrons and are reduced; at the positive anode, anions (or water) give up electrons and are oxidised. Cathode rule (in solution): metals up to aluminium in the activity series (Li … Al) are not deposited — water is reduced: $2 H_{2} O + 2 e^{-} \overset{}{\to} H_{2} ↑ + 2 OH^{-}$ ; those after hydrogen (Cu, Ag, Au) — the metal is deposited: $Cu^{2 +} + 2 e^{-} \overset{}{\to} Cu$ ; those in the middle (Zn … Pb) — the metal and often a little hydrogen. Anode rule (inert graphite anode): oxygen-free residues (Cl⁻, Br⁻, I⁻, S²⁻) are oxidised: $2 Cl^{-} - 2 e^{-} \overset{}{\to} Cl_{2} ↑$ ; oxygen-containing residues (SO₄²⁻, NO₃⁻) and F⁻ are not oxidised — water is oxidised and oxygen is released: $2 H_{2} O - 4 e^{-} \overset{}{\to} O_{2} ↑ + 4 H^{+}$ .

Lab work 7: CuCl₂ and KI solutions

The solution is poured into a U-tube, two graphite electrodes are lowered in and direct current is connected. CuCl₂: the cathode is coated with reddish copper, chlorine is released at the anode — moist iodine–starch paper turns blue: $CuCl_{2} \overset{}{\to} Cu + Cl_{2} ↑$ . KI (with phenolphthalein and starch added): the region around the cathode is pink — water was reduced and OH⁻ formed; the region around the anode is deep blue — iodine was released and combined with the starch: $2 KI + 2 H_{2} O \overset{}{\to} 2 KOH + H_{2} ↑ + I_{2}$ . In Na₂SO₄ solution only water is split: the volumes of hydrogen and oxygen are 2 : 1, and universal indicator is violet at the cathode and red at the anode. In copper sulfate solution, however, copper is deposited at the cathode, oxygen is released at the anode, and sulfuric acid accumulates in the solution.

Faraday's laws and problems

The amount of substance deposited at an electrode is proportional to the charge passed: $n = q / (z F)$ , $q = I t$ , $m = n M = M I t / (z F)$ (z is the number of electrons needed per formula unit). Example: if a current of 0.80 A passes through CuCl₂ solution for 20 minutes (1200 s): $q$ = 960 C, $n (e^{-})$ = 960/96 485 = 9.95 mmol, $n (Cu)$ = 4.97 mmol → 0.316 g of copper; at the anode the same number of moles of Cl₂ — 111 mL (at STP). The same charge deposits masses of different substances in proportion to $M / z$ . Detailed measurement and finding the charge of the electron are in the “Electrolysis, Faraday's laws” experiment.

Electrolysis of melts and industry

There is no water in a melt, so even the most reactive metals are deposited. Sodium: a melt of NaCl and CaCl₂ (≈ 600 °C; pure NaCl melts at 801 °C), a graphite anode, a steel cathode and an iron mesh between them — $2 NaCl \overset{}{\to} 2 Na + Cl_{2} ↑$ . Aluminium (the Hall–Héroult process, 1886): Al₂O₃ dissolved in molten cryolite (≈ 950–970 °C) is electrolysed; the carbon anodes burn to CO₂: $2 Al_{2} O_{3} + 3 C \overset{}{\to} 4 Al + 3 CO_{2} ↑$ ; about 13–14 kW·h per 1 kg of Al. Chlor-alkali: saturated brine in membrane cells (earlier diaphragm cells and mercury cathodes): chlorine at the anode, hydrogen and NaOH at the cathode; the membrane lets only Na⁺ through. The complex that started up at Navoiyazot in December 2019 is designed to produce ≈ 75 thousand t of caustic soda and ≈ 100 thousand t of PVC a year. Electrolysis is also used to refine copper, to produce zinc and to coat objects (electroplating).

Galvanic cells and corrosion

A more reactive metal gives electrons to a salt of a less reactive metal; if we put them in separate vessels and join them with a wire and a salt bridge, the electrons travel through the wire — this is a galvanic cell. The reactive metal is the anode (it is oxidised, “−”), the less reactive one the cathode (“+”). Emf: $ℰ = E °_{c} - E °_{a}$ ; the Daniell cell Zn|ZnSO₄||CuSO₄|Cu — 1.10 V. If the concentration changes, the potential shifts according to the Nernst equation: $E = E ° + \frac{0.059}{z} \lg c$ . Accumulators are also galvanic cells, but they can be recharged: $Pb + PbO_{2} + 2 H_{2} SO_{4} \overset{}{\to} 2 PbSO_{4} + 2 H_{2} O$ (≈ 2 V). Corrosion is the destruction of metals by their environment: chemical (dry gases, high temperature) and electrochemical (in an electrolyte solution — microgalvanic couples). Iron rusts when water and oxygen are present together: at the anode $Fe - 2 e^{-} \overset{}{\to} Fe^{2 +}$ , at the cathode $O_{2} + 2 H_{2} O + 4 e^{-} \overset{}{\to} 4 OH^{-}$ ; salt speeds up rusting, and so does contact with a less reactive metal (Cu).

Protection, laboratory methods, safety and the model

Protection: paint, enamel, oil; metal coatings — zinc (galvanised iron stays protected even when scratched: zinc is more reactive) and tin (if tin plate is scratched, the iron rusts faster); chromium, nickel; stainless alloys; sacrificial anodes (Mg or Zn blocks attached to a pipe or a ship's hull — they are eaten away themselves); cathodic protection from an external source; inhibitors. Ferroxyl indicator (K₃[Fe(CN)₆] + phenolphthalein, in agar): the anode areas where Fe²⁺ comes out turn blue, the cathode areas where OH⁻ forms turn pink. Corrosion rate: gravimetric — $v = Δ m / (S t)$ , g/(m²·day); electrochemical — the corrosion current: for iron 1 µA/cm² ≈ 0.25 g/(m²·day) ≈ 0.012 mm/year. Safety: chlorine is toxic — carry out electrolysis in a fume cupboard and with a small current; do not taste solutions of KI, CuCl₂ or CuSO₄, and wash your skin; melts and industrial electrolysis only at a plant. Model: the solutions are ideal and the gases are released completely (the dissolving of Cl₂ in water is not taken into account); the coloured zones around the electrodes are drawn using one-dimensional diffusion; the current efficiencies and voltages of the industrial cells are typical values; the rusting rates are representative; time is sped up (60 to 21 600 times, depending on the mode).

Did you know?

The polyvinyl chloride and caustic soda complex at Navoiyazot started up in December 2019: it is designed to produce about 75 thousand t of caustic soda and 100 thousand t of PVC a year — chlorine and alkali are obtained by electrolysis of table salt solution.
According to the IMPACT study published by NACE International in 2016, corrosion costs the world about 2.5 trillion US dollars a year — ≈ 3.4 % of world GDP (2013 data).

Inside the experiment

The experiment starts from a real-life situation: you predict the result, change the controls, run the experiment and draw a conclusion. Inside are missions, theory, a quiz and a guided tour in which the camera moves and explains the key points.

More experiments in this section

Redox reactions — Grade 8
Galvanic cells and corrosion — Year 1