How Corrosion Develops in Industrial Water Systems

The electrochemical picture

Corrosion in water is an electrochemical reaction. At the anode, metal gives up electrons and passes into solution as an ion. At the cathode, those electrons are consumed — in near-neutral aerated water, usually by dissolved oxygen. The water itself completes the circuit.

That is why removing any one of the four requirements stops it. No anode, no cathode, no electrolyte or no electron path means no corrosion. Practical control works by interfering with the anodic or cathodic reaction rather than by removing the water.

What accelerates it

• Dissolved oxygen feeds the cathodic reaction. It is the single most important driver in most open cooling and untreated feedwater systems.
• Chloride breaks down passive oxide films locally, which is why it is associated with pitting rather than general thinning.
• Low pH removes the protective film that would otherwise form and slow the reaction.
• Temperature speeds up reaction rate, as it does with most chemistry.
• Flow velocity matters at both extremes: too low allows deposits and stagnation, too high causes erosion-corrosion.

Why localised attack is the real problem

General corrosion removes metal evenly and slowly. It is measurable, predictable and can be designed around with a corrosion allowance. It is not usually what takes equipment out of service.

Localised corrosion concentrates the entire anodic reaction into a very small area. The metal loss per unit area is enormous compared with general attack, and the overall thickness measurement barely moves — so a wall thickness survey can look reassuring right up until a tube perforates.

• Pitting — chloride breaks the passive film at a point, and the pit chemistry becomes self-sustaining once established.
• Under-deposit corrosion — a deposit creates an oxygen-depleted zone beneath it, and the metal under the deposit becomes anodic to the metal around it.
• Crevice corrosion — the same differential aeration mechanism at a gasket face, a joint or a tube-to-tubesheet gap.
• Galvanic corrosion — where two different metals share an electrolyte, the less noble one corrodes preferentially.
• Microbiologically influenced corrosion — biofilm creates its own local chemistry, and some organisms produce corrosive metabolic products directly.

How inhibitors work — and how they fail

Anodic inhibitors form a protective film at anodic sites. They are effective, but they carry a specific risk: at insufficient dose they can protect most of the surface while leaving small unprotected areas, concentrating the attack there. An underdosed anodic inhibitor can be worse than none.

Cathodic inhibitors reduce the cathodic reaction rate and are more forgiving of underdosing, though generally slower to establish. Film-forming inhibitors coat the metal surface as a physical barrier.

All of them share one requirement: the film has to be established and then maintained. Interrupted dosing does not simply pause protection — it can leave a system in a worse condition than a system that never had a programme, because a partially broken film concentrates the attack.

Measuring what is actually happening

Corrosion coupons remain the standard method: pre-weighed metal specimens of the relevant alloys exposed in a representative flow for a defined period, then cleaned and re-weighed. They give an average rate over the exposure, and — importantly — the visual condition shows whether the attack was general or localised.

Online corrosion-rate sensors give a live reading, which is useful for seeing the response to a change rather than waiting for the next coupon interval. Iron and copper levels in the water indicate what is being lost from the system, and rising values point to where.