2.7.11 · D5Redox & Electrochemistry (Intro)
Question bank — Corrosion — electrochemical mechanism; cathodic protection, galvanization
Reminder of the sign convention used everywhere here: a more negative standard reduction potential means the metal gives up electrons more eagerly (is a stronger reducer, corrodes more readily). , and means the process is spontaneous (Gibbs Free Energy: ).
The figure below is the mental picture to hold while answering: a single water droplet on iron is a complete circuit — read it before the questions.
And the "potential ladder" — who sits above whom — settles almost every sacrificial-anode trap at a glance:
True or false — justify
TF1. Corrosion of iron can happen in perfectly dry air with plenty of oxygen.
False — the water film is the electrolyte that carries and ions between anode and cathode; with no liquid water there is no ion path, so the micro-cell cannot complete its circuit.
TF2. Rusting stops once a layer of rust covers the whole surface.
False — rust (, where is a variable amount of trapped water) is porous and flaky, so water and oxygen reach the fresh metal underneath and corrosion continues inward.
TF3. A more negative V means iron resists oxidation.
False — a more negative reduction potential means the reverse (oxidation) is favoured, so iron gives up electrons readily; that is exactly why it corrodes.
TF4. In cathodic protection the protected iron never changes into an electrode.
False — it very much becomes the cathode; the whole point is to force it into the reduction role so it accepts electrons instead of losing them.
TF5. Zinc protects iron only as long as the zinc physically covers every scratch.
False — even where the coating is scratched and bare iron is exposed, zinc still protects it electrochemically as a sacrificial anode because .
TF6. Copper would make a good sacrificial anode for an iron pipeline.
False — V is more positive than iron's V, so iron would corrode to protect the copper — the opposite of what we want.
TF7. Saltwater speeds up corrosion mainly because chloride ions oxidise the iron.
False — chloride is not the oxidiser. It accelerates corrosion two ways: it raises the electrolyte conductivity so ions shuttle faster between anode and cathode, and crucially it depassivates — chloride ions penetrate and break down protective oxide films, exposing fresh metal (a form of pitting attack).
TF8. Oxygen is oxidised during the cathodic step of iron corrosion.
False — oxygen is reduced (); it is the electron acceptor, which is the driving force pulling electrons from the iron.
TF9. A galvanised (zinc-coated) nail and a tin-coated nail behave identically when scratched.
False — zinc is more active than iron and sacrifices itself, but tin is less active than iron, so a scratched tin coating makes the exposed iron the anode and it corrodes faster.
TF10. Aluminium doesn't corrode because its reduction potential is very positive like gold's.
False — aluminium is actually very active ( strongly negative); it survives because its oxide is dense and adherent, sealing the surface — a phenomenon called Passivation.
Spot the error
SE1. " V, so corrosion is not spontaneous."
The formula is inverted: V, which is positive, so corrosion is spontaneous.
SE2. "The electrons travel from anode to cathode through the water film."
Electrons travel through the metal; it is the ions (, ) that move through the water film to complete the circuit.
SE3. "In the sacrificial method, magnesium is the cathode because it is being protected."
Magnesium is the anode (it is sacrificed); the iron structure is the protected cathode.
SE4. "Impressed-current protection uses a battery to push electrons out of the iron so it stops rusting."
The external supply pushes electrons into the iron, forcing it to be the cathode; pulling electrons out would only make it oxidise faster.
SE5. "For Mg protecting Fe, V."
The anode potential must be subtracted with its own sign: V; forgetting the double-negative flips the sign.
SE6. "Rust is , the direct product of meeting hydroxide."
is only an intermediate; it must be oxidised further to and dehydrate to before it is true rust.
SE7. "Since for corrosion is positive, the Nernst equation is unnecessary — corrosion always proceeds at the same rate."
tells you direction, not rate; the Nernst Equation gives the actual under real concentrations via , and kinetics/access to set the speed.
Why questions
WQ1. Why does the tip and bent region of an iron nail turn blue (Fe²⁺) rather than the smooth middle?
Stress concentrations and impurities at bends and tips lower the activation energy for oxidation, so those spots become the anodes where is released.
WQ2. Why does the smooth middle of the nail turn pink (OH⁻) instead?
Those regions have the best oxygen access, so oxygen reduction () happens there, making them cathodes; the released turns phenolphthalein pink.
WQ3. Why is oxygen reduction, not hydrogen evolution, the usual cathodic reaction in neutral rusting, and where does its working potential come from?
Because dissolved is plentiful while is tiny at neutral pH. Its standard value is V ( M , atm ); under real air ( atm) and pH 7 the Nernst equation, , actually raises it to about V because M is far below standard — so oxygen reduction is even more favourable than V suggests.
WQ4. Why can a small magnesium block protect a large steel ship hull?
The huge cell potential ( V) makes Mg oxidation vastly more favourable than Fe oxidation, so electrons flood into the hull, holding every iron atom as a cathode regardless of hull size.
WQ5. Why must sacrificial anodes be replaced periodically?
They corrode on purpose — the anode metal is consumed as it dissolves (), so once used up it can no longer supply protective electrons.
WQ6. Why does a difference in oxygen concentration alone (a "differential aeration cell") drive corrosion even on a single pure iron surface?
Regions starved of oxygen become anodes and oxygen-rich regions become cathodes; the potential difference set up by unequal acts like a concentration cell, driving iron dissolution at the poorly aerated spot.
WQ7. Why does lowering the pH accelerate iron corrosion?
More boosts the acidic cathodic reduction and also dissolves any protective oxide film, exposing fresh metal to attack.
WQ8. Why is a positive equivalent to saying ?
Because ; with and Faraday's constant C/mol both positive, a positive forces negative, i.e. spontaneous (Faraday's Laws, Gibbs Free Energy).
Edge cases
EC1. What happens to a galvanised iron sheet after the last of its zinc coating has been consumed?
With no active zinc left to sacrifice, the bare iron loses cathodic protection and begins to rust normally.
EC2. Two identical iron plates: one wholly submerged, one half in air and half in water. Which corrodes faster and where?
The half-in/half-out plate corrodes faster, concentrated near the waterline, because that zone gets both electrolyte contact and abundant dissolved oxygen for the cathodic reaction.
EC3. Gold has V; what does this say about its corrosion?
Its strongly positive reduction potential means removing electrons from gold is thermodynamically unfavourable, so it essentially does not corrode — the limiting case of a "noble" metal.
EC4. If dissolved oxygen were completely removed from neutral water, would iron still corrode significantly?
Corrosion nearly stops, because the dominant cathodic reaction (oxygen reduction) has no reactant; deaeration is in fact a real corrosion-control strategy for boiler water.
EC5. In the limit where the protected metal and the sacrificial metal have equal , does protection still occur?
No — with there is no driving force to force electrons onto the protected metal, so neither is preferentially protected; the anode must be genuinely more active.
EC6. How does raising the temperature change corrosion rate, and why?
Rate rises sharply because reaction rates follow Arrhenius behaviour () — a rough rule of thumb is that corrosion roughly doubles for every rise; higher speeds electron-transfer and ion diffusion (though very high can also lower solubility, partly offsetting it).
EC7. Is a "greener" alternative to zinc sacrificial anodes conceptually possible?
Yes in principle — any sufficiently active, less toxic metal (or impressed-current systems powered by renewables) reduces heavy-metal release, connecting corrosion control to Green Chemistry goals.
Recall Quick self-test
Corrosion needs which three things? ::: An anode (oxidation site), a cathode (reduction site), and an electrolyte to carry ions. The rule for picking a sacrificial anode metal? ::: Its reduction potential must be more negative than the protected metal's, so it oxidises preferentially. Why is rust worse than aluminium's oxide layer? ::: Rust is porous and flaky and doesn't seal the surface, whereas passivates and blocks further attack. What does the in mean? ::: A variable, non-fixed amount of trapped water — rust's water content changes with humidity.