EC-010 · Oxidation and Reduction Reactions in CorrosionJuly 11, 2026
EC TRACK · ELECTROCHEMISTRY & THE GALVANIC SERIES
Oxidation and Reduction Reactions in Corrosion
The matched pair behind every corroding pipe - oxidation and reduction, always running together, and the inversion that makes CP work.
Foundation~8 minutesPDH/CEC eligible
A right-of-way mower took the test station off at the grade — a brush tractor caught the post and sheared it, which out here happens more than anybody likes. You’re back to set a new one. The old anode’s lead wire lands in that station, so once the box is off you follow the lead down into the hole to see what’s on the end of it. Nobody on the crew knows how old this system is — you inherited it — so the call is to pull the anode and drop in a fresh one while the ground’s already open. A foot or so down, the lead runs into it: a magnesium anode that went in years ago as a 17-pound bar — better than two feet of bright metal, a few inches thick, packed in a cloth sack of backfill that ran better than forty pounds going into the ground. What you wrestle up now is a gnawed, gray-white lump, eaten down on every face, a fraction of what it started as.
A few feet over, the steel pipe that anode was wired to comes out of the trench clean. Same soil, same water. One piece of metal is most of the way gone and the other is barely touched. That’s not bad luck and it’s not a defect — that anode was put there to disappear so the pipe wouldn’t have to. Somebody set the job up so the magnesium would give up its metal and the steel would be spared. The question this module answers is the one underneath that: what does it actually mean for a piece of metal to “give up its metal,” what’s happening on the steel surface at the very same moment, and why are those two events locked together so tightly that you genuinely cannot have one without the other?
The answer is a matched pair of reactions — oxidation and reduction. One is the metal leaving. The other is what makes room for it to leave. They run at the same time, in lockstep, every time metal corrodes anywhere. Learn this pair cold and the rest of the external-corrosion track stops being a list of facts and starts being one machine you can see.
Why this module sits where it does
From the cell to the reactions inside it
The module before this one assembled the corrosion cell — the four parts every cell needs: an anode where metal leaves, a cathode where it doesn’t, an electrolyte carrying ions, and a metallic path carrying electrons. That gave you the machine. What it didn’t do was open up the two ends and show you the actual chemistry running at each one.
That’s this module. We zoom all the way in on the anode and the cathode and look at the reactions happening on those surfaces — the half that destroys metal and the half that consumes the electrons the destruction releases. Everything here lives at the level of “what reaction is running, and where.” We are not ranking which metals corrode first — that’s the next module. We are not turning electrons into pounds of lost steel with a formula — that’s the module after. We’re learning the two reactions themselves, because every one of those later topics is built on top of them.
The audience is working CP techs. We don’t need you to balance equations for a living. We need you to read a dig, a coupon, or a soil report and know which reactions are running and what that tells you. So that’s the depth: enough redox chemistry to read the field, and not one bit more.
Oxidation and reduction: the two words
Strip away everything else and corrosion chemistry comes down to where electrons go. Two definitions carry the whole module, and they are simpler than they sound:
Oxidation is the loss of electrons. A metal atom lets go of electrons and leaves the surface as a positively charged ion. This is the metal-loss side — the corrosion you can see and measure.
Reduction is the gain of electrons. Something in the environment picks up those electrons. No metal is lost here; this is the consuming side that keeps the whole thing running.
There’s a mnemonic that has outlived generations of chemistry classes because it works, and it’s worth burning in right now: OIL RIG. Oxidation Is Loss, Reduction Is Gain — of electrons. If you remember nothing else off this page, remember OIL RIG. It will keep the two straight every time.
OIL RIG. Oxidation Is Loss of electrons; Reduction Is Gain of electrons. Oxidation is the metal leaving. Reduction is what takes the electrons it leaves behind.
Oxidation does not mean oxygen
Here’s the trap in the word, and it catches good techs. “Oxidation” sounds like it has to involve oxygen. It does not. The name is a historical accident — back when the chemistry was first worked out, the electron-losing reactions people noticed most were metals combining with oxygen. The name stuck. But the definition has nothing to do with oxygen being present. Oxidation is the loss of electrons, full stop.
Why does that matter on the job? Because steel corrodes just fine in places with no oxygen at all — waterlogged clay, the bottom of a flooded trench, a sealed crevice under disbonded coating. In every one of those, iron is still losing electrons. The iron is still oxidizing. What changes when the oxygen is gone isn’t the metal-loss reaction — it’s the other half, the reduction, which simply switches to a different partner. The anode never cares what’s consuming its electrons. It only cares that something is.
You cannot have one without the other
This is the heart of the module, so we’ll land it hard. Oxidation and reduction are not two separate events that happen to occur near each other. They are two halves of a single process, welded together by the electrons that pass between them. Every electron set loose by oxidation has to be picked up by a reduction somewhere. If nothing is there to take the electrons, the metal can’t keep giving them up, and the corrosion stalls.
Think about what that means. The two reactions run at the same time, usually at different spots on the metal, and at matched rates. Speed up the reaction that consumes electrons and the metal-loss reaction speeds up to feed it. Choke off the consuming reaction and the metal loss slows down with it. Cut the electron supply entirely — we’ll see how — and the metal stops dissolving. The two are chained together. That chain is the single most useful idea you’ll carry out of this module, because every method of controlling corrosion is really a way of reaching in and breaking one end of it.
Oxidation and reduction always pair. The electrons released by the metal at the anode are the exact same electrons consumed at the cathode — same count, same instant. Block either half and both stop.
Why we write them as “half-reactions”
Because each reaction is only half the story, we write each one on its own and call it a half-reaction. The giveaway that it’s only a half is the loose electron sitting in the equation — the e−. A reaction with a free electron hanging off it can’t stand alone in the real world; that electron has to come from somewhere or go somewhere. So two half-reactions add together — one that releases electrons, one that consumes them — and the loose electrons cancel out, leaving a balanced whole reaction with none left over. That’s the bookkeeping. When you see e− in an equation for the rest of this track, read it as “this is one half of a pair, and its partner is close by.”
It’s the same chemistry as a battery — just spread out. Reduction-oxidation reactions — “redox” for short — are everywhere. They run in the battery in your truck and in a beaker in a lab. The only thing that makes corrosion special is geography: instead of the two halves being separated by design into a tidy package, they’re smeared across the same piece of pipe, inches or feet apart. Same physics. The metal just happens to be both the battery and the thing being destroyed.
The oxidation side: four reactions to know cold
Oxidation — the anodic half — always does the same thing no matter which metal it’s working on: an atom at the surface gives up electrons and floats off as a positive ion. The general shape of every anodic reaction is this:
General anodic reaction — oxidation M → Mn+ + ne− A metal atom (M) gives up n electrons, becomes an ion with charge n+, and leaves the surface. The number of electrons depends on the metal.
Four versions of that reaction cover almost everything you’ll meet in pipeline and CP work. Three of them are the metals we bury on purpose to corrode — sacrificial anodes like the magnesium lump from the hook. The first is the one we spend our careers trying to stop:
The four anodic reactions to know cold Fe → Fe2+ + 2e− (iron — carbon-steel pipe) Zn → Zn2+ + 2e− (zinc anode) Mg → Mg2+ + 2e− (magnesium anode) Al → Al3+ + 3e− (aluminum anode) Iron is the enemy reaction — it’s your pipe dissolving. The other three are friendly: they’re the metals we wire to a structure so they corrode instead of it. All four are the same move — metal to ion, electrons left behind.
Notice iron, zinc, and magnesium each give up two electrons, and aluminum gives up three. That count is fixed by the metal, and it matters later when we start tallying electrons. For now the takeaway is the pattern: every one of these is metal turning into a dissolved ion and leaving electrons behind in the structure. That’s oxidation, every time, and it’s the only thing that physically removes metal from your pipe.
The reduction side: what eats the electrons
Now the other half. The electrons the iron left behind travel through the steel to the cathode, and there they get consumed. Unlike the anode, the cathode is a bit of a chameleon — which reduction reaction runs depends entirely on what the environment hands it. Three reactions cover nearly all of it, and reading which one is running tells you a lot about the ground your pipe is sitting in.
Cathodic reaction 1 — oxygen reduction (neutral / alkaline, oxygen present) O2 + 2H2O + 4e− → 4OH− Dissolved oxygen and water soak up the electrons and make hydroxide ions. This is the everyday one for most buried pipe in most soils. It nudges the surface alkaline.
Cathodic reaction 2 — hydrogen evolution (acidic / oxygen-starved) 2H+ + 2e− → H2 Hydrogen ions take the electrons and bubble off as hydrogen gas. Expect this in low-pH ground or sealed-off, oxygen-dead spots.
Cathodic reaction 3 — metal-ion reduction (less common) Fe3+ + e− → Fe2+ A dissolved ion drops a charge instead. You’ll meet this far less often — certain scaled or chemistry-heavy environments — but it rounds out the set.
Two of these you’ll see constantly; the third is a footnote. The split between the first two is the one to internalize: oxygen present and near-neutral → oxygen reduction; acidic or oxygen-dead → hydrogen evolution. The anode is doing the identical iron-dissolution reaction in both cases. It’s the cathode that tells you what kind of place you’re standing in.
The one you can actually see
Hydrogen evolution has a tell. When the cathodic reaction is making hydrogen gas, it makes bubbles. A tech who exposes bare steel in acidic groundwater and sees a fine fizz of tiny bubbles clinging to the surface is watching the cathodic half-reaction happen in real time — live reduction, right there in the bell hole. Oxygen reduction, by contrast, is invisible; it just quietly raises the pH and leaves a chalky scale where the conditions are right. Reading those two signatures — bubbles versus scale — is reading the cathode straight off the dig.
Why oxygen makes ground so corrosive
Here’s a piece that runs against intuition. People assume aerated, well-drained soil is friendly and that the wet, airless muck is the dangerous stuff. For the metal-loss rate, it’s often the other way around, and oxygen is the reason. Oxygen is a fast electron consumer. Where there’s plenty of it, the cathodic reaction runs hard, which means it pulls electrons away quickly, which means the anode can keep dumping metal quickly to keep up — remember, the two rates are chained. Take the fast consumer away and the cell bogs down. So oxygen rarely attacks your pipe directly. What it does is keep the cathode hungry, and a hungry cathode keeps the anode dissolving.
Oxygen is the fast electron consumer. It doesn’t eat your pipe — it keeps the cathode running fast, and a fast cathode drags the anode along with it. That’s why well-aerated ground is often more corrosive, not less.
The electrons have to balance
Because every electron released has to be caught, the books always balance: electrons out of the anode equal electrons into the cathode. That sounds obvious, but it has a hard consequence the moment the two half-reactions use different electron counts. Iron oxidation lets go of two electrons per atom. Oxygen reduction swallows four per oxygen molecule. Those numbers don’t match — so to make the electrons come out even, you need two iron atoms dissolving for every one oxygen molecule reduced. The chemistry forces the ratio.
Worked example — put the two halves together
Setup. Buried steel in moist, aerated, near-neutral soil. The anode runs iron dissolution; the cathode runs oxygen reduction. Combine the two half-reactions into one balanced overall reaction, and account for the electrons.
The two halves. Anode: Fe → Fe2+ + 2e− (releases 2 electrons). Cathode: O2 + 2H2O + 4e− → 4OH− (consumes 4 electrons).
Balance the electrons. The anode gives 2 and the cathode needs 4. Run the anode reaction twice so it releases 4: 2Fe → 2Fe2+ + 4e−. Now 4 released, 4 consumed — even.
Add them up. 2Fe + O2 + 2H2O → 2Fe2+ + 4OH−. The four electrons cancel out — gone from the overall reaction, because they never left the system. The ferrous ions and hydroxide ions then pair up into ferrous hydroxide, Fe(OH)2, the first solid corrosion product — the gunk that goes on to become rust.
The takeaway. Two iron atoms lost per oxygen molecule reduced. That fixed ratio — electrons released equal electrons consumed — is the bookkeeping behind every “how much metal did we lose?” calculation you’ll do later. Land the idea here; the arithmetic comes in a later module.
That last line is the boundary. Turning this electron count into actual pounds of steel and amps of current is its own module further up the track. Here we just plant the principle: the electrons balance, and that balance sets the ratio of everything else.
Same iron, different cathode
To prove the “anode doesn’t care, cathode adapts” idea, put the exact same steel in two different places and watch only the cathode change.
Worked example — one anode reaction, two cathodes
Setup. The same carbon-steel pipe in two spots: (A) damp, aerated, near-neutral soil, and (B) an acidic, oxygen-starved pocket — low pH, no air getting in. What runs at the anode and the cathode in each, and what would you see?
Spot A — aerated, neutral. Anode: Fe → Fe2+ + 2e−. Cathode: O2 + 2H2O + 4e− → 4OH−. Signs: no bubbling; hydroxide makes the surface alkaline, and where the iron ions meet oxygen and water you get the familiar rust crust and maybe a chalky scale on the protected metal.
Spot B — acidic, oxygen-dead. Anode: Fe → Fe2+ + 2e− — identical. Cathode: 2H+ + 2e− → H2. Signs: a fine fizz of hydrogen bubbles, and often clean, dark pits instead of heaped rust, because the iron ions tend to stay dissolved and wash away rather than pile up.
The takeaway. The metal-loss reaction never changed — iron oxidizes the same way in both. Only the cathode swapped partners to match the environment. When you can name which cathodic reaction is running from what you see, you’re reading the ground through the chemistry.
The trick behind every CP system
Now the payoff — the reason this module is the gateway to the whole protection half of your training. We said the iron only keeps oxidizing because it needs to release electrons, and a reduction reaction nearby keeps taking them. So ask the obvious question: what if the electrons came from somewhere else?
That’s the entire idea of cathodic protection. We supply electrons to the structure from an outside source — from a sacrificial anode wired to it, like the magnesium that came up with the spoil, or from a powered system pushing current through the ground. Once electrons are arriving from outside, the iron has no reason to give up its own. The cathodic reaction still runs on the pipe surface, fed by the outside electrons, but the iron-dissolution half goes quiet. The metal stops leaving. We’ve turned the whole structure into a cathode.
Cathodic protection works by feeding a structure electrons from outside, so its own iron never has to give any up. Kill the reason for the oxidation half and the metal stops dissolving. That’s it — that’s why CP works.
Look back at the hook with that in hand. The magnesium anode wasn’t unlucky — it was the outside source. It oxidized on purpose, pumping its electrons into the steel, so the steel could sit there as a cathode and keep its metal. The anode disappeared so the pipe didn’t have to. The whole job, every reading at every test station, comes back to one question: is this structure still getting enough electrons from outside that its own iron stays put?
Why corrosion wants to happen at all. It helps to know the metal is fighting uphill the whole time it’s a pipe. Iron in nature is an ore — an oxide. We spend a lot of energy at the mill turning that ore into clean steel. The metal “wants” to go back to its ore state, and corrosion is simply that slide back downhill, acting itself out. Rust is just iron returning to what it was before we refined it. We don’t stop that wish; with CP, we just keep paying its electron bill from the outside so it never has to cash in your pipe.
The road ahead — and why this pair is the hinge
EC-010 is module 3 of 7 in the Electrochemistry & the Galvanic Series series. You now have the two reactions that drive every corrosion cell: oxidation, the metal leaving, and reduction, what takes the electrons it leaves behind — always paired, always balanced. The rest of the set builds straight up from here. The next module ranks the metals, so you can predict which one in a pair will be the one to oxidize. After that comes the arithmetic that turns this electron balance into real pounds of metal and amps of current. Later still, how the rates actually behave once the cell gets running.
Hold onto the chain. Oxidation and reduction are locked together by their electrons, and every way we fight corrosion is a way of reaching in and breaking that chain — shrinking the metal that’s exposed, cutting the path the electrons travel, or feeding the structure electrons from outside so its own iron never leaves. Control the electrons and you control the corrosion. That’s the whole game, and you just learned the move it’s played with.
Key takeaways
Oxidation is loss of electrons; reduction is gain. OIL RIG. Oxidation is the metal leaving (the anodic half); reduction is what consumes the electrons it releases (the cathodic half).
Oxidation does not require oxygen. Iron oxidizes even in oxygen-free ground — the name is historical. What changes without oxygen is the cathodic reaction, not the metal loss.
You cannot have one without the other. The two half-reactions run at the same time, at matched rates; electrons released equal electrons consumed. Block either half and both stop.
Four anodic reactions to know: Fe, Zn, Mg (two electrons each) and Al (three) — all metal-to-ion. Iron is your pipe corroding; the other three are sacrificial anodes corroding on purpose.
Three cathodic reactions: oxygen reduction (O2 + 2H2O + 4e− → 4OH−) in neutral/aerated ground; hydrogen evolution (2H+ + 2e− → H2) in acidic/oxygen-dead spots — the one that bubbles; and metal-ion reduction, the rare one.
The electrons balance. Iron gives 2, oxygen takes 4 — so 2 iron atoms dissolve per oxygen molecule reduced. That ratio is the basis of the mass-loss math in a later module.
Oxygen is the fast electron consumer — it keeps the cathode running hard, which keeps the anode dissolving. That’s why aerated soil is often more corrosive, not less.
CP is the inversion. Feed a structure electrons from outside and its iron has no reason to oxidize — the metal stops leaving. Every CP system is this one move.
References & further reading
The Electrochemical Reaction Leading to Corrosion & Reactions at the Anode Surface — Field Notes from RCS article on the anodic and cathodic half-reactions; included with this module as a PDF supplement.
Oxidation Does Not Mean Oxygen — Field Notes from RCS companion on where the words “oxidation” and “reduction” came from, and the trap in the name; included with this module as a PDF supplement.
Electrochemical Basics — Field Notes from RCS companion on oxidation, reduction, and the chemistry terms behind them.
Corrosion Basics: An Introduction — foundational text on the electrochemistry of corrosion and the anodic and cathodic reactions.
Cathodic Protection Training Materials — industry credential materials covering corrosion reactions at the CP-technician level.
AUCSC Short Course Materials — pipeline corrosion technician training curriculum.
Peabody’s Control of Pipeline Corrosion — long-standing field
Listen — narrated walkthrough
Oxidation and Reduction Reactions in Corrosion
Same scope as the read — oxidation and reduction defined, OIL RIG, why the two halves can’t run apart, the four anodic and three cathodic reactions, the electron balance, and the inversion that makes cathodic protection work — walked through with the dug-up magnesium anode as the anchor.
Narrated by Mike Roberts · ~20 min
Listen on the drive in or while waiting for the coating to cure. Come back for the deck or the worked problems whenever you want.
Once you’ve worked through the audio or the deck, head to the Apply lesson for three field problems — matching the electrons on a magnesium anode, reading two CP setups through the reaction pair, and the anode that quit — and then the quiz to lock it in. EC-011 is next; the Electrochemistry & the Galvanic Series certificate posts to your profile when you complete the full set (EC-008 through EC-014).
Apply — three field-grounded problems
Match the pair, find the pair, lose the pair
Three short problems, all built from the matched pair you just learned — and all kept at the level you’d actually use on the job. In the first you match up the electrons in a sacrificial-anode reaction. In the second you find the anode and the cathode in the two kinds of cathodic protection. In the third you reason through what happens when protection quits on you. Read each setup, think it through, then click each step to compare with our working.
How to use this lesson. Read the setup. Think before you click. Each step reveals what we’d say.
Problem 1 · Match the electrons
The reaction that’s eating the anode
Setup. Picture the magnesium anode from the start of this module, doing its job in damp, aerated soil — slowly giving itself up to protect the line. Two half-reactions are running. The magnesium oxidizes: Mg → Mg²⁺ + 2e⁻. On the steel it’s protecting, oxygen reduction takes the electrons: O₂ + 2H₂O + 4e⁻ → 4OH⁻.
Put the two halves together into one balanced reaction — the reaction that’s actually consuming the anode. The one rule: every electron the magnesium gives up has to be used by the oxygen side.
Step 1 — match the electrons
Count what each side handles. Magnesium gives up 2 electrons per atom. Oxygen reduction takes 4 at a time. Those don’t match — so it takes two magnesium atoms to feed one oxygen reaction.
Now it’s 4 electrons given and 4 electrons taken. Even.
Step 2 — add the halves together
Put them together and the electrons cancel — they never actually left the system, they just went from the magnesium to the steel.
2Mg + O₂ + 2H₂O → 2Mg(OH)₂
The magnesium ions and the hydroxide ions pair up into magnesium hydroxide. That’s the magnesium turning into corrosion product — the anode wearing away, exactly what we buried it to do.
Step 3 — name the two halves
Tag each half with OIL RIG. The magnesium side — losing electrons — is the oxidation. The oxygen side — gaining them — is the reduction. Two halves, locked together, two magnesiums for every oxygen.
It’s the same move we made with iron and oxygen back in the Read. Different metal, same idea: match the electrons, add the halves, and the ratio falls out on its own.
Match the electrons, then add the halves. Magnesium gives 2 and oxygen takes 4, so it’s two magnesiums per oxygen: 2Mg + O₂ + 2H₂O → 2Mg(OH)₂. The electrons set the ratio — you don’t pick it.
Problem 2 · Find the anode, find the cathode
Two ways to protect the same line
Setup. Same buried steel line, protected two different ways — the two kinds of cathodic protection you’ll run into. You don’t need to write any chemistry for this one. Just figure out, in each case, which piece of metal is corroding (the anode) and which is being protected (the cathode).
SETUP A · GALVANIC
Hardwaremagnesium anode
Wired tothe steel line
Powernone – natural
SETUP B · IMPRESSED CURRENT
Hardwarerectifier + groundbed
Wired tothe steel line
Powerrectifier (AC in)
For each setup: which piece is the anode (the one that corrodes), and which is the cathode (the protected steel)? And in one sentence — why is the steel protected?
Step 1 — the one idea behind both
Cathodic protection does one thing, two ways: it makes your steel the cathode. A cathode only runs reduction — electrons coming in — so the steel’s iron has no reason to give up its own. The oxidation gets shoved off the pipe and onto something else. Find that “something else” and you’ve found the anode.
Step 2 — Setup A: the magnesium anode
Anode: the magnesium. It’s more active than steel, so it corrodes on its own and feeds its electrons into the line.
Cathode: the steel — protected.
Why protected: the magnesium is handing the steel electrons for free, so the steel’s iron never has to give up any. The anode dissolves so the pipe doesn’t. (That’s the lump from the start of the module.)
Step 3 — Setup B: the rectifier
Anode: the groundbed out in the field, pushed by the rectifier.
Cathode: the steel — protected, same as before.
Why protected: there’s no sacrificial metal here — the rectifier takes wall power and pushes electrons onto the steel from outside. Same result: the steel runs reduction, and the corroding happens out at the groundbed instead of on your pipe.
Step 4 — the common thread
Galvanic or impressed current, the steel ends up doing the same thing: only reduction. One setup gets the electrons from a metal that’s happy to corrode; the other gets them from a rectifier. Either way, your structure is the cathode, and a cathode doesn’t lose metal.
Both kinds of cathodic protection do one thing: make the steel the cathode and move the corrosion somewhere you chose. Find what’s corroding and what’s protected, and any CP setup reads straight off the pair.
Problem 3 · When protection quits
The anode that quit
Setup. That buried line is protected by its magnesium anode, and for a couple of years the readings look good — the steel isn’t losing metal. Then a contractor’s mini-excavator, working an unrelated job nearby, catches the anode’s lead wire and snaps it clean underground. Nobody sees it happen. The line looks exactly the same from the surface.
No math here — just reason it through. (1) While the anode was connected, which metal was corroding, and why wasn’t the steel losing any? (2) The instant that lead wire snaps, what changes for the steel? (3) If you dig the line up a year later, what would you expect to find?
Step 1 — while the anode was hooked up
The magnesium was the one corroding. It oxidized and fed its electrons down the lead wire and into the steel. The steel sat there as the cathode, taking those electrons in.
Why didn’t the steel lose metal? Because its iron only oxidizes when it needs to let go of electrons — and the magnesium was supplying them from outside. With its electrons coming for free, the steel had no reason to give up any of its own. Protected.
Step 2 — the moment the wire snaps
The outside electron supply is gone. The steel is on its own again. With nothing feeding it electrons, the steel’s own iron goes right back to doing the oxidizing — Fe → Fe²⁺ + 2e⁻ starts running on the pipe wall again.
Nothing looks different at the surface. But from that moment, the pipe is corroding — quietly, with no anode covering for it anymore.
Step 3 — a year later in the ditch
You’d expect fresh corrosion on the steel — new pitting and metal loss that wasn’t there before — because the pipe has been an unprotected, bare anode for a year. And the magnesium would be frozen right where it was the day the wire broke; it stopped wearing away the moment the circuit opened, because it had nothing to feed anymore.
If anyone had taken a reading, the line would have flipped from protected to not — the day of the snap, not the day you dug it up.
Cathodic protection is just an outside electron supply. Cut it — a snapped lead, a dead rectifier — and the steel goes straight back to oxidizing. The matched pair never went anywhere; you were just feeding the cathode from outside. That’s why a broken anode lead is an emergency, not a paperwork problem.
What this Apply lesson was after. Three uses of one idea, all at field level. You matched the electrons in the reaction eating a magnesium anode. You found the anode and the cathode in the two kinds of cathodic protection and saw they do the same thing — leave the steel running only reduction. And you reasoned through what happens when the protection quits: cut the outside electrons, and the steel is right back to oxidizing.
The Quiz at the end of this module checks the core ideas with a mix of recall and recognition questions. After that, the Electrochemistry set moves on to ranking the metals — which one in a pair becomes the anode, and by how much. You’ve been calling the anode all module; next you’ll learn how to predict it.
Time limit: 0
Quiz Summary
0 of 10 Questions completed
Questions:
Information
You have already completed the quiz before. Hence you can not start it again.
Foundation tier · EC TRACK · ELECTROCHEMISTRY & THE GALVANIC SERIES
One module done. Keep going — you'll earn the certificate when you finish this section, and the Foundation medal when you complete every section in the tier.
