Electrochemical Cells and Applications

Lessons

Senior 6 · Chemistry

Electrochemical Cells and Applications

5 lessons

Lesson 4 of 5

Corrosion and Its Prevention

25 min

A galvanic cell nobody wanted

Look up at an old corrugated iron roof and you may see orange-brown patches where the metal is crumbling. Iron left in damp air slowly turns into rust, weakening roofs, bridges, water pipes, vehicles and machinery, and costing money to repair and replace.

Rusting is not a simple reaction between iron and oxygen. It is an electrochemical cell, with an anode and a cathode on the same piece of metal and the water film as the electrolyte. Seeing it that way explains why salt and acid make it worse, and why a coat of zinc protects iron even where it is scratched.

The underside of an old corrugated iron roof: the grey zinc coating is still intact in places, while orange-brown patches show where the iron beneath has rusted. Photo: Bill Nicholls, CC BY-SA 2.0, via Wikimedia Commons

Definition

Corrosion: the gradual destruction of a metal by chemical or electrochemical reaction with its surroundings. Example: the rusting of iron; the tarnishing of silver.

Rust: hydrated iron(III) oxide, $\mathrm{Fe_2O_3 \cdot xH_2O}$ , the product of the corrosion of iron and steel in the presence of both water and oxygen.

Sacrificial protection: protecting a metal by connecting it to a more reactive metal (more negative $E^{\circ}$ ), which is oxidised instead. Example: zinc blocks bolted to a steel ship's hull.

Galvanising: coating iron or steel with a layer of zinc, which acts both as a barrier and as a sacrificial metal.

Cathodic protection: making the protected metal the cathode of a cell, either with a sacrificial anode or with an external power supply (impressed current).

Rusting as an electrochemical cell

Figure 4.1 shows a drop of water resting on iron. Oxygen dissolves most easily near the edge of the drop, where the water layer is thin.

Figure 4.1: The water-drop model of rusting. The anode is where oxygen is scarce; the cathode is where it is plentiful.

Anode (centre, little oxygen): $\mathrm{Fe(s) \rightarrow Fe^{2+}(aq) + 2e^{-}}$ , $E^{\circ} = -0.44\,\mathrm{V}$ for the reverse reaction. A pit forms here.
Cathode (edge, plenty of oxygen): $\mathrm{O_2(g) + 2H_2O(l) + 4e^{-} \rightarrow 4OH^{-}(aq)}$ , $E^{\circ} = +0.40\,\mathrm{V}$ .
Electrons travel through the iron itself from anode to cathode; ions move through the water.

$E^{\circ}_{\text{cell}} = (+0.40) - (-0.44) = +0.84\,\mathrm{V}$

The ions then meet and react further:

$\mathrm{Fe^{2+}(aq) + 2OH^{-}(aq) \rightarrow Fe(OH)_2(s)}$

$\mathrm{4Fe(OH)_2(s) + O_2(g) + 2H_2O(l) \rightarrow 4Fe(OH)_3(s)}$

Iron(III) hydroxide loses water to give rust, $\mathrm{Fe_2O_3 \cdot xH_2O}$ . Rust is porous and flakes off, so, unlike the oxide layer on aluminium, it does not protect the metal underneath.

What speeds rusting up

Factor	Effect and reason
both water and oxygen present	essential: water is the electrolyte and a reactant; oxygen is reduced at the cathode
dissolved salt (near the sea, salted roads)	faster: more ions make the water a better electrolyte
acid (acid rain, industrial air)	faster: in acid the cathode reaction becomes $\mathrm{O_2 + 4H^{+} + 4e^{-} \rightarrow 2H_2O}$ , $E^{\circ}_{\text{cell}} = 1.23 - (-0.44) = +1.67\,\mathrm{V}$
contact with a less reactive metal (e.g. copper)	faster: iron becomes the anode of a two-metal cell
higher temperature	faster, as for most reactions

Preventing corrosion

Method	How it works	Typical use
paint, oil, grease, plastic coating	barrier: keeps out water and oxygen	bridges, gates, bicycle chains, tools
tin plating	barrier only; tin is unreactive and non-toxic	food cans
galvanising (zinc coating)	barrier plus sacrificial protection	roofing sheets, buckets, wire fences
sacrificial anodes (Mg or Zn blocks)	the block is oxidised instead of the steel, and is replaced when used up	ships' hulls, buried pipelines, water heaters
impressed current	a power supply makes the steel the cathode	pipelines, steel in concrete
alloying	stainless steel contains chromium, which forms a thin, self-healing oxide layer	sinks, cutlery, surgical tools

Worked example

Problem. A galvanised (zinc-coated) bucket and a tin-plated can are both scratched through to the iron. Use $E^{\circ}$ values to explain why the bucket does not rust at the scratch, while the can rusts faster than uncoated iron.

Solution. $E^{\circ}(\mathrm{Zn^{2+}/Zn}) = -0.76\,\mathrm{V}$ , $E^{\circ}(\mathrm{Fe^{2+}/Fe}) = -0.44\,\mathrm{V}$ , $E^{\circ}(\mathrm{Sn^{2+}/Sn}) = -0.13\,\mathrm{V}$ .

Zinc and iron: zinc has the more negative potential, so in the cell formed at the scratch zinc is the anode and is oxidised: $\mathrm{Zn \rightarrow Zn^{2+} + 2e^{-}}$ . The iron is the cathode, where oxygen is reduced; the iron itself is not oxidised. ( $E^{\circ}_{\text{cell}}$ for zinc reducing $\mathrm{Fe^{2+}}$ is $-0.44 - (-0.76) = +0.32\,\mathrm{V}$ , so any $\mathrm{Fe^{2+}}$ formed would be reduced back.)

Tin and iron: iron now has the more negative potential, so iron is the anode and is oxidised, while the large area of tin acts as the cathode. The cell drives the corrosion of the exposed iron faster than if no tin were present.

Answer: zinc is sacrificed to protect iron; tin protects only as long as the coating is unbroken.

Worked example

Problem. A zinc sacrificial anode on a ship supplies an average protective current of 0.50 A. What mass of zinc is used up in one year (365.25 days)? ( $F = 96\,485\,\mathrm{C\,mol^{-1}}$ , $M(\mathrm{Zn}) = 65.38\,\mathrm{g\,mol^{-1}}$ )

Solution.

$Q = It = 0.50 \times (365.25 \times 24 \times 3600) = 1.58 \times 10^{7}\,\mathrm{C}$

$n(e^{-}) = \frac{1.58 \times 10^{7}}{96\,485} = 164\,\mathrm{mol}, \qquad n(\mathrm{Zn}) = \frac{164}{2} = 81.8\,\mathrm{mol} \quad (\mathrm{Zn \rightarrow Zn^{2+} + 2e^{-}})$

$m = 81.8 \times 65.38 = 5.35 \times 10^{3}\,\mathrm{g}$

Answer: about 5.3 kg of zinc per year, which is why sacrificial anodes are inspected and replaced regularly.

Warning

Common mistakes

Saying iron rusts in dry air or in water with no dissolved oxygen. Both water and oxygen are needed.
Calling aluminium "unreactive" because it does not corrode much. It is very reactive ( $E^{\circ} = -1.66\,\mathrm{V}$ ) but is protected by a thin, tough layer of aluminium oxide; rust gives iron no such protection.
Confusing galvanising with electroplating in general. The point of zinc is that it is more reactive than iron; a coating of a less reactive metal (tin, copper, nickel) protects only while unbroken.
Thinking sacrificial anodes last for ever. They are consumed and must be replaced.

Tip

To decide what a coating does when scratched, compare $E^{\circ}$ with iron's −0.44 V: more negative (Mg, Zn, Al) = sacrificial, iron is protected; more positive (Sn, Cu, Ni) = barrier only, iron corrodes faster once exposed.

FuseSchool: Corrosion. Watch for the conditions needed for iron to rust and the ways of preventing it, including sacrificial protection.

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Practice

Write the half-equations at the anode and the cathode when iron rusts in neutral water, and calculate $E^{\circ}_{\text{cell}}$ .
Explain why rust forms in a ring around a drop of water on iron, while the iron under the centre of the drop is eaten away.
Explain why cars rust faster where salt is spread on roads, and why iron rusts faster in acidic water.
Magnesium blocks are attached to buried steel pipelines. Explain how they work, and why they must be replaced from time to time.
A 1.0 kg magnesium anode supplies an average current of 0.20 A. How many days will it last? ( $M(\mathrm{Mg}) = 24.31\,\mathrm{g\,mol^{-1}}$ )
A plumber joins a copper pipe directly to a steel pipe. Predict which pipe corrodes faster near the joint, and explain with $E^{\circ}$ values.
A householder in a rainy area must choose between galvanised iron roofing sheets and painted (ungalvanised) iron sheets. Evaluate the two choices in terms of corrosion.

Show answersHide answers

Anode: $\mathrm{Fe \rightarrow Fe^{2+} + 2e^{-}}$ . Cathode: $\mathrm{O_2 + 2H_2O + 4e^{-} \rightarrow 4OH^{-}}$ . $E^{\circ}_{\text{cell}} = 0.40 - (-0.44) = +0.84\,\mathrm{V}$ .
Oxygen dissolves most easily near the edge, so the edge is the cathode; the centre, short of oxygen, is the anode, where iron dissolves as $\mathrm{Fe^{2+}}$ . The $\mathrm{Fe^{2+}}$ and $\mathrm{OH^{-}}$ ions meet in between, where the hydroxide precipitates and is oxidised to rust.
Dissolved salt adds ions, so the water is a better electrolyte and the corrosion cell passes more current. In acid, oxygen is reduced as $\mathrm{O_2 + 4H^{+} + 4e^{-} \rightarrow 2H_2O}$ ( $E^{\circ} = +1.23\,\mathrm{V}$ ), so $E^{\circ}_{\text{cell}}$ rises to +1.67 V.
Magnesium ( $E^{\circ} = -2.37\,\mathrm{V}$ ) is more reactive than iron, so it becomes the anode and is oxidised, while the steel is the cathode and is protected. The magnesium is used up as it corrodes, so it must be replaced.
$n(\mathrm{Mg}) = \frac{1000}{24.31} = 41.1\,\mathrm{mol}$ ; $n(e^{-}) = 82.3\,\mathrm{mol}$ ; $Q = 82.3 \times 96\,485 = 7.94 \times 10^{6}\,\mathrm{C}$ ; $t = \frac{7.94 \times 10^{6}}{0.20} = 3.97 \times 10^{7}\,\mathrm{s} = 459\,\mathrm{days}$ , about 15 months.
The steel pipe. Iron (−0.44 V) is more negative than copper (+0.34 V), so where the two touch in water, iron is the anode and corrodes faster ( $E^{\circ}_{\text{cell}} = +0.78\,\mathrm{V}$ ) while copper is protected.
Galvanised sheets are better: the zinc is a barrier, and where it is scratched or worn it still protects the iron sacrificially until the zinc is used up. Paint is only a barrier; once it chips, the exposed iron rusts, and rust can spread under the paint. Painting galvanised sheets as well gives the longest life. Cost, appearance and the life of the coating also affect the choice.

Sources and credits

Standard electrode potentials at 25 °C: Standard electrode potential (data page), Wikipedia contributors, Wikipedia, licensed under CC BY-SA 4.0; values compiled there from P. Vanýsek, "Electrochemical Series", CRC Handbook of Chemistry and Physics (CRC Press), and other cited sources; adapted by Ganzaa.
Rusting reactions and prevention methods: Rust and Galvanization, Wikipedia contributors, Wikipedia, licensed under CC BY-SA 4.0, adapted by Ganzaa.
Images: Rust on the roof - geograph.org.uk - 1422581.jpg by Bill Nicholls, licensed under CC BY-SA 2.0; Electrode protecting a screw.jpg by Rémi Kaupp, licensed under CC BY-SA 3.0; both via Wikimedia Commons.
Video: Corrosion, FuseSchool – Global Education, embedded from YouTube.
The text, water-drop diagram, tables, worked examples and Practice items were written and drawn by Ganzaa.
This lesson adapts CC BY-SA material and is shared under CC BY-SA 4.0.

1A galvanic cell nobody wanted

2Rusting as an electrochemical cell

3What speeds rusting up

4Preventing corrosion

A galvanic cell nobody wanted

Rusting as an electrochemical cell

What speeds rusting up

Preventing corrosion