Corrosion

Corrosion in Isolation

About once a quarter, a customer delivers a piece of pipe to my office and asks a fair question: “Why did this corrode?”

I take a deep breath before responding, because corrosion in isolation is nearly meaningless.

A piece of failed metal, removed from its environment, stripped of its operating history, and divorced from its chemistry, cannot explain itself. Corrosion is not a property of metal alone. It is a process enabled, accelerated, and shaped by its surroundings.

So, I ask questions. Lots of questions.

• What type of system was it connected to?

• How was it positioned in that system?

• What other metals are present in that system?

• Have you seen any evidence of corrosion elsewhere in that system?

• Do you have a recent water analysis for that system?

Almost invariably, I am met with a blank stare. “Can’t you just tell me what caused it?”

I rarely say what I am thinking, but the honest answer is simple: the metal did what metals do. And staring at corrosion in isolation is merely an act of nihilism. Metal returning to its natural state is simply a law of nature. Metals corrode because corrosion is thermodynamically favored; it is inevitable.

But inevitability does not mean helplessness. While metal wants to corrode, it does not corrode without assistance. The environment must cooperate. And when we shift our focus from the artifact to the process, from the pipe to the system, corrosion becomes not only explainable, but controllable.

We must first understand why metal corrodes. Then we must clarify the mechanism of corrosion. Finally, we must explore the many ways in which that mechanism enables vicious forms of corrosion to occur.

This is the process of providing a suitable answer, and actionable solution, for when our customers ask, “where did my metal go?”

The Restless Nature of Metals

Metal does not exist in nature. At least not in the form we’re used to.

Iron appears familiar to us as pipes, plates, and vessels – but these forms are temporary. In its natural state, iron is bound to oxygen as stable oxides: hematite, magnetite, goethite. Low energy. Comfortable. Content.

Everything we call “metal” in an industrial system is the result of force.

Enormous amounts of energy are spent mining ore, stripping away oxygen, and forcing atoms into rigid, crystalline lattices. These are orderly, conductive, and useful, but they are not stable. They exist in opposition to entropy.

Corrosion, therefore, is not a defect in material. It is the cost of borrowed order.

The same shared electrons that give metals their strength and conductivity also make them vulnerable. Those electrons are mobile. They are exposed. And once a path exists for them to leave, they will.

In the preceding chapter, we introduced irreversible oxidation and the minimal conditions required for electron loss to occur. When those conditions align (an anodic site, a cathodic reaction, an ionic pathway, and an electrical connection) the corrosion cell is complete.

At that point, corrosion is no longer a possibility. It is a certainty. What remains uncertain is not whether corrosion will occur, but where, how fast, and how violently it will localize.

That uncertainty is explained by a battery.

The Corrosion Cell

The corrosion cell is nature’s entropic battery.

It begins with a thermodynamic imbalance.

In the previous chapter, we explored how oxidants exert their reduction potential until they find a surface willing to give up electrons. For now, we will limit our focus to the role of oxygen’s potential to snatch up electrons. O₂ exists as a stable covalent molecule, but its high electronegativity and favorable reduction thermodynamics make it an exceptionally strong electron sink.

Metals exist at a higher energy state than they’d prefer, refined into a thermodynamically unstable crystalline structure. Within the lattice, electrons are delocalized from their respective atoms, providing both the glue that holds the structure together and a pathway for electrical current to flow. The vulnerability arises at the surface, where individual metal atoms can lose electrons, become ions, and leave the lattice.

This is the first step in forming a corrosion cell: a mismatch of energy.

The difference between these states is voltage: the electromotive force that compels electrons to move. When this voltage is paired with a continuous path for electrons and ions, chemistry gives way to electricity and a closed circuit forms. 

Once this circuit exists, corrosion is no longer a chemical reaction happening to the metal. It is an electrical current flowing through it.

Like any battery, the corrosion cell requires four components:

The Anode: Lower Reduction Potential → Site of oxidation (corrosion)

The Cathode: Higher Reduction Potential → Site of reduction (electron consumption)

The Electrolyte: Allows dissolved ions to move, completing one half of the circuit

The Metal Path: Allows electrons to move, completing the other half of the circuit 

The Anode: Where Metal Leaves

At the anode, metal atoms lose electrons and enter solution as ions:

Fe → Fe²⁺ + 2e⁻

Cu → Cu⁺ + e⁻

(For copper, the first anodic step is commonly Cu → Cu⁺ + e⁻. Under more oxidizing conditions, Cu⁺ can be further oxidized to Cu²⁺.)

This is where metal is lost.
This is where pits begin.

A neutral atom becomes a charged ion, is ejected from the lattice, and dissolves into the surrounding water. It generally begins at a minor defect in the surface of the metal, a microscopic imperfection that leaves the area especially vulnerable: grain boundaries, inclusions, or physical scratches.

Once corrosion starts, the metal does not “thin” evenly. It departs atom by atom, wherever conditions make departure easiest.

The Cathode: Where Electrons Go

Electrons released at the anode do not vanish. They must be consumed elsewhere.

In neutral to alkaline aerated water, that consumption commonly occurs through oxygen reduction:

O₂ + 2H₂O + 4e⁻ → 4OH⁻

This reaction generates hydroxide at the cathode, which means the pH directly at the metal surface can be much higher than the pH measured in the bulk water.

In acidic systems, hydrogen ions can also participate in the cathodic reaction:

2H⁺ + 2e⁻ → H₂

These reactions do not remove metal directly. They provide the electron demand that allows metal dissolution to continue at the anode. This is why oxygen, oxidizers, and localized cathodic reactions are so dangerous. By consuming electrons, they allow anodic metal loss to continue uninterrupted.

The Electrolyte: One Half of the Circuit

Water acts as an electrolyte by providing a pathway for ions to move.

Charged particles are formed in both the anodic and cathodic reactions. Water allows iron and copper ions to migrate away from the anode, and hydroxide to disperse from the cathode. By allowing these oppositely charged corrosion by-products to move freely, the corrosion current continues to flow.

Without an electrolyte to allow this migration, these charged particles would accumulate at the surface, creating electrical resistance. When the charges are not able to move freely and balance, corrosion comes to a halt. This is why corrosion of dry metals largely comes to a halt. 

The Metal Path: The Other Half of the Circuit

Metal allows the freed electrons to move.

When the metal surface is oxidized, the metal ions dissolve into water, while the electrons are transported through the metal to the cathode. The free electrons are consumed in the cathodic reaction, removing them from the metal.

This is why corrosion is not confined to the point of attack. So long as they are connected by an electrolyte, anodic and cathodic sites can be separated over a distance. The metal becomes a wire. The electrolyte distributes the charged ions. Oxygen provides the demand.

The Completed Circuit

When all four elements of the corrosion cell are present, the circuit is energized, and chemistry leaves a record.

At the anode, metal atoms enter solution as positively charged ions.
At the cathode, oxygen reduction generates hydroxide ions.
These charged particles diffuse through the electrolyte, but they don’t stay independent for long.

Near the metal surface of mild steel, ferrous ions (Fe²⁺) encounter hydroxide (OH⁻), and the two combine:

Fe²⁺ + 2OH⁻ → Fe(OH)₂

This reaction does not occur uniformly throughout the water. It occurs at the interface, where ions are being produced and consumed by the circuit itself. The result is ferrous hydroxide, a corrosion byproduct that forms in intimate contact with the metal surface.

As corrosion continues, more ferrous hydroxide accumulates at or near the surface. In the presence of oxygen, it is further oxidized into more stable iron oxides and oxyhydroxides, the familiar red-brown corrosion byproducts we call rust. These compounds adhere loosely or tightly depending on conditions, forming films, tubercles, or deposits that partially cover the metal.

The appearance of corrosion products marks more than material loss.
It marks a completed electrical loop.

Electrons have moved through the metal.
Ions have moved through the water.
Charge has been balanced.
Energy has been released.

Corrosion is not simply metal disappearing into water. It is metal reorganizing itself into a lower-energy form. Guided, constrained, and made visible by the completed circuit.

Passivation: Metal’s Brief Defense

Left to corrosion, a steel surface should not last long in open water.

And yet, many metals survive for years without catastrophic loss. They do so not because corrosion has stopped, but because it has slowed.

This slowing is the result of passivation.

It is not immunity from corrosion.
It is more like a ceasefire.

Passivation occurs naturally with most metals, resulting from the corrosion process itself. As the oxidation reaction at the anode releases metal into solution, those ions combine with anions in solution to form thin surface films at the metal-water interface.

These films do not eliminate anodic or cathodic reactions, but they increase resistance by restricting electron transfer, ion migration, and oxidant access at the metal surface. They make it harder for the circuit to be maintained.

Mild Steel

Carbon steel forms a protective oxide layer composed primarily of iron oxides. Depending on the water, it can form loose and porous, offering little resistance, or, under the right conditions, dense, adherent, and genuinely protective. High chloride and sulfate push it toward the porous, soluble end. Nitrite and phosphate push it toward the dense, protective end.

The porosity is what matters. A porous film leaves the surface open to oxygen and ions, and the more porous it is, the more corrosion concentrates into localized, severe metal loss rather than spreading thin and slow. Mild steel survives not because its film is inherently robust, but because conditions often stay just favorable enough to keep that film on the protective side of the line.

Copper

Copper behaves differently.

As copper ions react with oxygen and carbonate species, they form a dense, adherent patina. Unlike rust, this patina is often protective. It limits further electron transfer and shields the underlying metal.

This is why copper pipes darken instead of dissolving, and why the Statue of Liberty is enshrined in green armor, rather than dissolving.

Stainless Steel

Stainless steel relies on precision.

Chromium in the alloy forms an exceptionally thin oxide film, only atoms thick, but remarkably tenacious. When damaged, it reforms almost instantly, as a protective chromium oxide layer. This self-healing film is what allows stainless steel to survive environments that would rapidly destroy carbon steel. But it does not mean that stainless steel is always the best choice.

The self-healing capacity can be severely limited by two constraints: low oxygen (anaerobic) environments interfere with repassivation, and chloride-rich environments make localized attack more likely to occur.

Galvanized Steel

Galvanized steel avoids reliance on iron’s weak defenses entirely.

Instead, it coats steel with zinc, a less noble metal that willingly corrodes first. Zinc sacrifices itself, supplying electrons and satisfying cathodic demand before iron ever needs to participate.

As long as zinc remains, the steel beneath is spared. When the zinc is gone, the truce ends. 

Breaking the Ceasefire 

In most situations, passivation is nothing more than a fragile truce.

Passivation is effective when it is able to shield the vulnerable anodic sites from the corrosive forces of water. Any condition that destabilizes the passive film reveals these localized weaknesses.

Given that the conditions for corrosion still exist, corrosion will return to these sites with a vengeance. Oxidation will concentrate where resistance is lowest, and voltage is highest.

This is why corrosion rarely looks fair: it feeds on itself.

It’s also why we need to know what exactly is breaking the ceasefire of passivation.

Film Saboteurs

Film saboteurs break the truce.

They are not the primary cause of corrosion, but they determine where corrosion concentrates and how fast it becomes destructive. By disrupting passive films, they expose reactive metal beneath and allow dissolved metal complexes to be carried into the bulk water, lost for good.

This is how uniform corrosion becomes pitting.
And pitting is how systems fail quickly.

Even in otherwise stable systems, a small number of saboteurs can dominate corrosion behavior. Four deserve our utmost attention.

Mechanical Disruption: Flow and Erosion

As water velocity increases, shear forces at the metal surface rise. Turbulence, caused by abrupt changes in direction, diameter, or pressure, amplifies those forces and scours away protective films.

This is why corrosion appears first at elbows, tees, pump impellers, control valves, and heat exchanger inlets. This effect is especially pronounced in copper, where protective films can be swept away by high velocity. The film is stripped away, and anodic sites are exposed faster than they can re-passivate.

We rarely control system geometry or fluid dynamics. But understanding where mechanical disruption occurs explains why corrosion localizes, and where it will reappear after repair.

Chemical Dissolution: Acids

Passive films are stable only within a narrow pH window.

As pH drops, hydrogen ions protonate the metal oxide layer, converting protective films back into soluble species. A film that took weeks or months to form can dissolve in minutes under acidic conditions.

This is why transient acid exposure is so dangerous. The damage persists long after pH is restored. Acids also increase corrosion potential by accelerating anodic reactions, compounding the loss.

The Film Breaker: Chloride

Chloride ions are small, mobile, and relentless.

They penetrate oxide films and concentrate at the metal surface, where they form soluble metal–chloride complexes. These complexes carry metal away from the surface, preventing repassivation.

Chloride does not remove the entire film.
It creates weak points. And then exploits them.

The result is not uniform thinning, but deep, localized pitting.

The Double Agent: Sulfate

Sulfate ions are less aggressive on their own, but more dangerous in combination.

Chemically, they can destabilize passive films through metal-complex formation. Biologically, sulfate can serve as an electron acceptor for sulfate-reducing bacteria. Under low-oxygen deposits and biofilms, that metabolism can become part of the corrosion circuit itself. Sulfide may appear as a byproduct, but in cooling-water MIC the deeper concern is the electron demand created by the microbial reaction.

Sulfate is therefore a double agent: chemically disruptive, and biologically enabling under the right conditions. Systems with stagnant zones, deposits, or low-flow regions provide ideal territory for both mechanisms to operate.

Quantifying the Sabotage

Film saboteurs weaken the defenses.

Water treaters like indices because they promise simplicity. They take a complex water analysis and compress it into one number. That can be useful for scale, where carbonate chemistry often dominates the question. Corrosion is less cooperative.

Corrosion is not controlled by bulk water chemistry alone. It depends on metallurgy, oxygen availability, pH, temperature, velocity, deposits, biofilm, inhibitor concentration, surface condition, galvanic coupling, and the localized chemistry that develops at the metal surface. The water in the pit is not the same as the water in the sample bottle.

Chloride and sulfate still matter. They interfere with protective films, increase conductivity, participate in localized chemistry, and become more troublesome as they concentrate in recirculating systems. But their effect cannot be reliably converted into one universal “corrosivity number.”

The Larson-Skold Index is sometimes used to show the relationship between chloride, sulfate, and alkalinity. That is its best use here: not as a predictor, but as a reminder. Chloride and sulfate are not innocent spectators. They make it harder for steel to maintain protective films.

In cycled cooling water, however, the index can be misleading. If chloride, sulfate, and alkalinity all concentrate together, the ratio may remain nearly unchanged even though the water has become more concentrated, more conductive, and often more aggressive. This is why corrosion control cannot rely on source-water indices alone.

The better question is, “What conditions are helping the corrosion cell organize itself?”

Saboteurs do more than restart corrosion. By stripping resistance, concentrating oxidants, or accelerating electron transfer, they increase the voltage driving it. When this happens, the corrosion cell does not merely persist. It intensifies.

Supercharging the Corrosion Cell

The corrosion cell is governed by the same rule as any electrical circuit.

Ohm’s Law:

Current = $\frac{\mathbf{Voltage}}{\mathbf{Resistance}}$

When the passive film fails, both terms move in the wrong direction. Resistance drops because the protective barrier is gone. Voltage increases because the newly active site develops an electrochemical potential difference relative to the surrounding passive surface.

Current density increases. Metal loss accelerates. Damage localizes.

Once this imbalance forms, the corrosion cell begins to reinforce itself. The environment at the anode diverges from the bulk water. As dissolved metal ions accumulate and hydrolyze, acidity rises. Chloride migrates inward to preserve charge balance. Oxygen remains more available outside the pit than inside it, strengthening the differential-aeration cell. Repassivation becomes increasingly unlikely.

Oxygen Differential Corrosion

Once a pit forms, the environment inside becomes starved of oxygen, while the surrounding metal remains well-aerated. This sets up an oxygen differential cell: high oxygen at the rim (the cathode), low oxygen at the pit base (the anode).

The Result: A powerful electrochemical potential forms. The large cathode (surface) drives intense dissolution at the tiny anode (pit). Chloride ions migrate into the pit to balance the charge, creating an acidic, salt-rich soup that drills through metal at terrifying speeds.

Under-Deposit Corrosion

Deposition in water systems can come from corrosion by-products, scale, biofouling, and suspended solids. They’re all lumped in by the rather encompassing term “deposition,”  which indicates that the surface of the metal is covered by a foreign material. This creates a localized oxygen-depleted region, which sets up the conditions for oxygen differential corrosion.

Cooling towers are prime victims. Dirt, sludge, and biofilm act just like a pit. The metal under the deposit is starved of oxygen (Anodic), while the clean metal nearby is oxygen-rich (Cathodic).

The Result: Beneath these innocuous-looking deposits, corrosion accelerates in the shadows. The process is quiet, persistent, and perfectly designed to escape detection until it’s too late.

Microbiologically Influenced Corrosion

Microbiologically influenced corrosion (MIC) is a form of corrosion that has been reorganized by life. Biofilms create deposits, oxygen gradients, stagnant microenvironments, and localized chemistry that differs sharply from the bulk water. Under those deposits, small anodic regions can form next to larger cathodic regions, producing intense localized attack.

Sulfate-reducing bacteria are especially important because they use sulfate as an electron acceptor under low-oxygen conditions. In doing so, they can accelerate corrosion by consuming electrons that ultimately originate from the metal.

Sulfide may be present as a byproduct, and in some environments it can contribute to corrosion or form iron sulfide deposits. But the deeper problem is not the smell of sulfide. The deeper problem is that microbial metabolism can become part of the corrosion circuit itself.

The result: a living electrochemical cell. Biofilm lowers resistance, oxygen gradients increase voltage, and bacterial metabolism helps consume electrons. This is why MIC is so localized, aggressive, and difficult to diagnose from bulk water chemistry alone.

Galvanic Corrosion

Galvanic corrosion occurs without subtlety.

It happens when two dissimilar metals are placed in electrical contact and immersed in an electrolyte. Metals differ in their willingness to oxidize, which is referred to as their nobility. The nobility of different metals can be compared by their relative positions in the galvanic series.

Galvanic corrosion increases voltage by electrically coupling two metals with different nobility. The more noble metal exerts an electric potential, leading to corrosion of the less noble metal. The further apart the metals are on the Galvanic Series, the higher the voltage and the faster the anode dissolves.

This is why we use dielectric unions: insulating barriers that break the electrical path and stop the battery. Mixed metallurgies in systems sometimes makes this challenging.

The Result: Galvanic corrosion always targets the less noble metal. It will proceed until there is no more anodic surface available, which often results in rapid, catastrophic failure of the anodic component. The severity of attack depends not just on how far apart the metals sit on the galvanic series, but on the ratio of their surface areas: small anode/ big cathode being far more dangerous than the reverse.

From Supercharging to Suppression

Up to this point, we have watched corrosion become efficient.

Passive films fail. Resistance drops. Voltage concentrates. Current focuses into smaller and smaller regions of metal until damage accelerates and localizes. Whether through oxygen differentials, deposits, biology, or metallurgy, the story has been the same: corrosion thrives when its circuit is allowed to organize itself efficiently.

Protection begins by doing the opposite.

Corrosion Protection: Interrupting the Circuit

Corrosion cannot be eliminated.

But it can be disrupted.

Once corrosion is understood as an electrical circuit, protection becomes a matter of strategy. By interfering with any one part of the circuit, we can slow the inevitable march of entropy to a crawl.

• Anode

• Cathode

• Electrolyte

• Metallic path

For many industrial systems, two of these are fixed. Water must remain. Metal must remain. That leaves two levers we can realistically pull: the anodic and cathodic reactions.

Every corrosion control strategy works by targeting these parts of the circuit. Some raise resistance. Some remove reactants. Some prevent the circuit from assembling at all.

Cathodic Control: Increasing Resistance

At the cathode, electrons are consumed. If that consumption can be slowed, the entire circuit weakens.

In most water systems, cathodic reactions generate hydroxide ions, raising pH at the metal surface:

O₂ + 2H₂O + 4e⁻ → 4OH⁻

This means the pH at the cathode can be much higher than the pH measured in the bulk water. That localized high pH is critical.

It shifts carbonate equilibria and creates conditions favorable for calcium carbonate precipitation directly at the surface. When properly controlled, this precipitation forms a whisper-thin, adherent, self-renewing film that interferes with oxygen access and electron transfer.

This film is not permanent. It is fugitive.

It dissolves back into the bulk water when conditions change, and reforms when conditions return. Its strength lies not in durability, but in equilibrium.

Waters with sufficient calcium and alkalinity can support this defense naturally. Waters that lack these building blocks cannot. This is why ultra-soft, low-alkalinity waters are often corrosive: the cathode has nothing to hide behind.

The Langelier Saturation Index (LSI) and Ryznar Stability Index (RSI) can be used to evaluate whether a given water chemistry has the potential to support this fugitive calcium-carbonate layer. These indices describe the thermodynamic tendency of water to precipitate or dissolve calcium carbonate under bulk conditions.

They do not predict corrosion.

Local conditions at the metal surface (oxygen availability, flow, deposits, aggressive anions, and microbiology) ultimately determine how fast corrosion occurs. LSI and RSI simply tell us whether calcium carbonate has the opportunity to participate in raising cathodic resistance at all.

In practice, we don’t rely on calcium carbonate alone. Calcium phosphate, zinc compounds, and calcium-phosphonate complexes are also highly dependent on local pH, and become less soluble near the cathode and form thin precipitated films at the surface.

None of these films make the metal surfaces immune. But they can increase resistance. By making it harder for oxygen to reach the surface and harder for electrons to be consumed, they slow the cathodic half of the corrosion cell.

Phosphonates help refine this process when properly controlled. At low concentrations, they interfere with uncontrolled crystal growth, prevent bulk deposition, and help favor thinner, more dispersed surface films rather than heavy scale. But more is not always better. Overfeeding phosphonates can interfere with calcium carbonate film formation and can complex soluble metal ions, which may increase corrosion rather than suppress it.

Phosphates can operate on both sides of the cell. At the cathode, they can participate in precipitated calcium-phosphate films formed by the high-pH surface environment. At the anode, orthophosphate can also interact with iron to promote a denser protective film.

The goal is a chemically enhanced, self-renewing boundary layer: thin enough to avoid heat-transfer penalties, dynamic enough to reform if disrupted, and robust enough to suppress oxygen-driven cathodic corrosion.

Cathodic control does not stop corrosion.

It raises resistance, which makes the circuit harder to run.

Anodic Protection: Locking the Lattice

At the anode, metal atoms leave the lattice.

Fe → Fe²⁺ + 2e⁻

Anodic control aims to prevent that departure.

Anodic inhibitors work by strengthening or enforcing passive films that lock metal atoms in place. When effective, they make electron loss energetically unfavorable and suppress metal dissolution at its source.

Nitrite is the classic example. It is itself an oxidizer, and that is precisely how it works: it drives the oxidation of the metal surface toward a stable, protective ferric oxide film rather than toward dissolution. By forcing the formation of that film, it suppresses anodic dissolution in closed systems, provided concentration and pH are maintained. When maintained within proper limits, nitrite provides excellent corrosion protection with minimal scaling concerns.

But anodic protection carries inherent danger.

Partial coverage concentrates corrosion. If inhibitor levels fall too low, unprotected sites become intense anodes surrounded by protected cathodic surfaces. The result is rapid, localized attack.

Anodic control works best where chemistry is stable, monitoring is reliable, and the system can be kept within narrow limits.

Removing the Oxidizer: Starving the Circuit

Corrosion requires an electron acceptor.

In most industrial water systems, that acceptor is oxygen.

In systems where air ingress and re-aeration can be controlled, such as boiler feedwater systems, oxygen can be removed almost entirely. This is achieved mechanically through deaeration, or chemically through oxygen scavengers such as sulfite, bisulfite, or organic reducing agents. Even trace oxygen concentrations can sustain corrosion, which is why both mechanical and chemical methods are commonly used together.

Though often discussed as boiler-specific practices, oxygen control is fundamentally electrochemical: fewer oxidizers mean fewer cathodic reactions. This does not protect the metal directly, but it deprives the cathode of its preferred reaction. Without oxygen, the circuit weakens regardless of surface chemistry.

Eliminating oxygen is one of the most effective ways to suppress the circuit entirely, but sustained oxygen removal is only practical where re-aeration can be controlled. In an open cooling tower, oxygen is continuously restored by contact with air.

Film-Forming Inhibitors: Preventing the Circuit from Assembling

Some inhibitors do not slow reactions. They prevent contact with the surface.

Azoles are used specifically to control copper corrosion. These organic molecules chemically bond to the copper surface, forming hydrophobic films that block water access and suppress both anodic and cathodic reactions. 

Filming Amines or Film-Forming Amines operate similarly in boilers and condensate systems. They adsorb onto the metal surface, orienting their polar groups toward the metal and their hydrophobic portions away from it. This hydrophobic barrier protects against oxygen, carbonic acid, and iron transport.

Choosing the Lever

None of these corrosion control strategies are competing philosophies.

They are tools.

Some systems favor dynamic, self-renewing defenses. Others require rigid, enforced passivity. Most real systems rely on a balance, nudging both reactions just enough to keep the circuit from organizing itself efficiently.

The mistake is not choosing the “wrong” inhibitor.
The mistake is forgetting which part of the circuit you are touching.

Corrosion control is not about chemistry in isolation.

It is about where electrons are allowed to move, and where they are not.

Every inhibitor, every adjustment, every material choice ultimately answers the same question:

How hard are we making it for the circuit to run?

The Two Ends of Thermodynamics

Corrosion describes what happens when metal yields to water.

Metal exists in a restless, high-energy state. Oxidation offers a path downhill, and water provides the conditions for that path to open. Once the circuit is complete, electron transfer is rewarded, the metal begins to dissolve, and corrosion becomes not just possible, but thermodynamically favored. It is entropy at work.

But this is only one half of the story.

Thermodynamics does not only tear structures apart. Under the right conditions, it builds them.

In corrosion, atoms leave a solid lattice and the structure is destroyed.
In scale, dissolved ions leave solution and assemble into a new solid.

Both processes are driven by the same search for lower-energy arrangements. One dismantles order. The other creates it.

That is the tension at the heart of water treatment.

The conditions that suppress corrosion often encourage scale. Waters capable of forming protective films may also be poised to deposit stone. Waters that resist scale may leave metal exposed and vulnerable. The water treater lives between those outcomes, managing one without inviting the other.

Up to this point, we have followed electrons as they escape metal.

In the next chapter, we will follow ions as they decide whether to remain dissolved or assemble into stone.

Engineering Notes: Corrosion

“If you are reading this straight through, you can skip this section and lose nothing essential to the story. These notes are for the operators, engineers, and technicians who need to do the math.”

Why This Matters in the Field

Corrosion is electron flow. If you can account for the circuit, you can:

• predict where corrosion will localize (cells, gradients, deposits)

• choose control strategies intelligently (anodic vs cathodic vs oxygen removal)

• interpret coupons/probes quantitatively (MPY)

• anticipate risk from water chemistry (indices + chloride/sulfate ratios)

Most corrosion failures are not “mysteries.” They are batteries operating in disguise.

Core Tools & Constants

The Corrosion Cell

A functioning corrosion cell requires all four components:

1. Anode (metal dissolves): Fe → Fe²⁺ + 2e⁻

2. Cathode (electrons consumed): O₂ + 2H₂O + 4e⁻ → 4OH⁻

3. Electrolyte Path: ionic conduction through water

4. Metallic Path: electron conduction through metal

Ohm’s Law (Circuit Logic)

The corrosion cell is governed by the same rule as any electrical circuit.

Current = $\frac{Voltage}{Resistance}$

The rate of metal loss is directly proportional to the corrosion current flowing from the anode to the cathode. Increased voltage drives higher current; increased resistance reduces current.

• Voltage increases with strong oxidizers, oxygen gradients, galvanic couples

• Resistance increases with passivation, protective films, coatings

Engineering Qualification: Ohm’s Law provides the circuit logic for corrosion: greater electrochemical driving force and lower resistance permit greater corrosion current. Actual corrosion current is also limited by electrode kinetics, polarization, and mass transfer. Therefore, $I = V/R$is a conceptual model here, not a stand-alone equation for calculating corrosion rate.

Oxygen Differentials

Oxygen does not have to be “high” for corrosion to be severe. It only has to be uneven.

A deposit, crevice, gasket, or biofilm creates:

• Low Oxygen Environments (Under deposit, inside crevice): Anode

• High Oxygen Environments (Bulk water): Cathode

Small anode + big cathode = high current density at the anode → pitting corrosion.

Measuring corrosion

Corrosion is measured as thickness lost.

The standard industry unit is Mils per Year (MPY). 1 mil = 1/1,000 of an inch.

MPY = $\frac{534\ \times \ W}{(D\ \times \ A\ \times \ T)}$

Where:

• 534 = Unit Conversion Factor

• W = Weight loss (mg)

• D = Density of the metal (g/cm³)

• A = Surface Area of the coupon (in²)

• T = Time exposed (hours)

Note: The recommended time frame for coupon exposure is typically 90 days. This allows time for passivation to occur, providing a representative sample of long-term corrosion. Shorter timeframes may significantly overestimate corrosion activity.

Weight-loss coupons report average metal loss over the exposed area. A low or moderate MPY does not rule out severe localized pitting.

Interpretation:

Acceptable corrosion rates differ for open and closed systems because their operating environments are fundamentally different (specifically oxygen exposure, water chemistry stability, and contaminant levels).

General Corrosion Rates for Open Recirculating Cooling Water Systems

General Corrosion Rates for Closed Loop Systems

On Corrosion Indices

Operators often ask which index predicts corrosivity. The honest answer is that none reliably do.

The Larson-Skold Index is sometimes cited for cooling water. It should be used with caution, if at all. It is an empirical relationship derived from Great Lakes municipal water, expressing the ratio between: “Aggressive Anions” (chloride and sulfate) and “Protective Inhibitors” (bicarbonate and carbonate):

The calculation requires milliequivalents per liter (meq/L) to be used.

Conversions (mg/L to meq/L):

• Chloride**:** meq/L = mg/L ÷ 35.5 

• Sulfate**:** meq/L = mg/L ÷ 48 

• Alkalinity (as CaCO₃): meq/L = mg/L as CaCO₃ ÷ 50

Interpretation for uncycled municipal or makeup waters:

• < 0.8: Lower corrosion risk. The passive film is likely stable.

• 0.8 – 1.2: Potential corrosion risk. Interference of film formation likely.

• > 1.2: High corrosion risk. Localized corrosion and pitting are likely to occur.

Limitation:

Because it is a ratio, it is blind to cycling. Concentrate a water through six cycles and the index does not move, even though the chloride and sulfate driving localized attack have multiplied. For makeup or municipal water it offers a rough indication. For cycled cooling water it does not apply, and treating its output as a corrosion prediction is a mistake.

The reason no index succeeds is structural. General corrosion can be estimated. Localized corrosion (pitting, crevice attack, under-deposit cells) cannot. Localized attack is governed by conditions at the metal surface that a bulk water analysis never captures: local oxygen gradients, occluded crevice chemistry, deposit coverage, microbiological activity, and surface flow. These are the factors that actually determine where and how fast metal is lost, and they are invisible to any formula fed by a routine water panel.

Use the ratio, if you use it at all, for one purpose: to show a skeptic that chloride and sulfate are not bystanders.

Galvanic Corrosion

For galvanic corrosion to happen, three things must be present simultaneously:

1. Dissimilar Metals: Two metals with different "nobility" (electrical potential).

2. Metal-to-Metal Contact: The metals must be touching (or connected by a wire/bolt) so electrons can flow.

3. Electrolyte: Both metals must be exposed to a conductive liquid.

When these conditions occur, the "Less Noble" metal becomes the anode, loses electrons, and corrodes rapidly. The "More Noble" metal becomes the cathode, receives electrons, and is protected.

The Galvanic Hierarchy

The simplified list below shows the relative nobility of different metals. Exact ranking depends on environment (chlorides, temperature, passivity), but the direction is reliable.

Most Noble (Protected Cathode)

• Titanium

• Stainless Steel (Passive) – see note

• Copper / Brass

• Lead

• Cast Iron

• Carbon Steel

• Aluminum

• Zinc (Galvanizing)

• Magnesium

Least Noble (Sacrificial Anode)

Note: Stainless steel appears near the top of this list only while its passive film is intact. In low-oxygen or chloride-rich conditions (crevices, under deposits, stagnant zones) the film cannot maintain itself and the steel goes "active," dropping to a position near carbon steel. Stainless steel is a particularly important example because its position changes dramatically when its passive film is lost. The same general warning applies to other passive metals: the galvanic series describes behavior in a specified environment, not an immutable property of the material.

\