Corrosion Kinetics

Introduction

Thermodynamics tells us that iron in moist air can rust, but not whether a bridge will last ten years or two hundred. Rate is a kinetic question, and it is governed by the same electrode kinetics you met with the Butler–Volmer and Tafel equations. The key idea is that on a corroding surface two different reactions run at once, and the surface adopts the single potential at which their rates match.

Core explanation

No net current. A freely corroding piece of metal is not connected to any external circuit. Every electron released by metal oxidation must be consumed by a cathodic reaction on the same piece. At steady state:

I anodic (metal oxidation) = I cathodic (e.g. oxygen reduction)

Mixed potential. Each half-reaction has its own equilibrium potential, but the metal can have only one potential. That potential, the corrosion potential E corr, lies between the two equilibrium values. At E corr the metal oxidation is driven anodically (positive overpotential) and the cathodic reaction is driven cathodically (negative overpotential), each by just enough to make their currents equal. This theory of mixed potentials was set out by Wagner and Traud in 1938.

Evans diagrams. Plot log i on one axis against E on the other. The anodic Tafel line for metal dissolution rises from its equilibrium potential; the cathodic Tafel line falls from the more positive equilibrium potential of the oxidant. Where the lines cross gives E corr and the corrosion current density i corr. This picture shows immediately how kinetics control the rate:

- a larger exchange current density for the cathodic reaction (a better catalyst surface) moves the crossing to higher current; - a steeper Tafel slope (more sluggish kinetics) lowers i corr; - a larger gap between equilibrium potentials generally increases i corr.

Diffusion control. In neutral aerated water, dissolved oxygen is only about 0.25 mmol/dm³ at room temperature. The cathodic line therefore hits a limiting current set by oxygen diffusion. The corrosion rate then equals this limiting current and depends on stirring, flow and temperature rather than on the metal. This is why steel corrodes faster in flowing aerated water.

Measuring the rate. An electrochemical instrument can polarise a sample away from E corr and record current. Extrapolating the linear Tafel regions back to E corr gives i corr. Alternatively, near E corr the curve is almost linear, and the polarisation resistance R p = ΔE ÷ Δi is inversely proportional to i corr (the Stern–Geary relation).

From current to thickness loss. By Faraday's law, a corrosion current density i gives a mass loss per unit area per second of iM ÷ (nF). Dividing by density ρ gives the penetration rate:

rate = iM ÷ (nFρ)

Galvanic coupling. When two metals touch in an electrolyte, they share one mixed potential. The less noble metal becomes more anodic and corrodes faster, while the more noble metal is partly protected. A large cathode area connected to a small anode area is especially damaging, because the entire cathodic current must be supplied by the small anode.

Formulae

At E corr: i a = i c = i corr; penetration rate = i corr M ÷ (nFρ); Stern–Geary: i corr = B ÷ R p, with B = b a b c ÷ [2.303(b a + b c)].

Step-by-step reasoning

To construct an Evans diagram:

1. Mark the equilibrium potentials of the metal and of the oxidant. 2. Draw the anodic Tafel line from the metal's exchange current density. 3. Draw the cathodic Tafel line from the oxidant's exchange current density, adding a limiting-current plateau if transport is slow. 4. Read E corr and i corr at the intersection. 5. Convert i corr into a penetration rate.

Visual explanation

Imagine two straight lines on a log-current chart forming an X. The rising line is iron dissolving; the falling line is oxygen reduction. If the falling line bends to a vertical plateau before reaching the rising line, the crossing moves to the plateau: the rate is fixed by oxygen supply.

Real-world analogy

A corroding metal is like a shop with one cashier (the anode) and one delivery driver (the cathode). Goods cannot leave faster than money comes in; the shop settles at a single speed where the two are balanced, and the slower partner sets the pace.

Real-world example

Steel screws holding copper sheeting on an old ship's hull corroded rapidly: the small steel anode was coupled to a very large copper cathode. The reverse arrangement, copper screws in large steel sheets, causes much less damage because the anodic current is spread over a large area.

Why?

Why does zinc corrode slowly in pure acid but faster when touching platinum? Hydrogen evolution on zinc has a very small exchange current density. Platinum catalyses hydrogen evolution strongly, so connecting it shifts the cathodic line to higher currents and increases the corrosion current of zinc.

Common misconception

"The metal with the most negative electrode potential always corrodes fastest." Rate depends on kinetics: exchange current densities, Tafel slopes, oxygen supply and protective films. Aluminium and titanium are very reactive thermodynamically yet corrode extremely slowly.

Worked example

Question: Iron (M = 55.8 g/mol, n = 2, ρ = 7.87 g/cm³) corrodes with i corr = 10 μA/cm². Estimate the penetration rate in mm per year.

Reasoning: Rate = iM ÷ (nFρ) = (10 × 10⁻⁶ × 55.8) ÷ (2 × 96 485 × 7.87) ≈ 3.67 × 10⁻¹⁰ cm/s. One year is about 3.15 × 10⁷ s, giving 1.16 × 10⁻² cm per year.

Answer: About 0.12 mm per year.

Quick check

1. At the corrosion potential of a freely corroding metal, how do the anodic and cathodic currents compare? Answer: They are equal in magnitude, so there is no net external current.

Exam focus

Draw and interpret Evans diagrams, explain how exchange current density and oxygen transport shift i corr, and convert current density to penetration rate with Faraday's law. Explain galvanic corrosion and the importance of the anode-to-cathode area ratio.

Advanced insight

Real surfaces are heterogeneous: grain boundaries, inclusions and scratches create local anodes and cathodes, and local chemistry can differ sharply from the bulk. In a crevice or pit, metal hydrolysis acidifies the trapped solution and chloride migrates in, so the local corrosion current can be orders of magnitude higher than the average measured value.

Summary

A corroding metal adopts a mixed potential at which anodic metal dissolution and cathodic reduction proceed at equal rates. Evans diagrams show how exchange current densities, Tafel slopes and mass-transport limits set the corrosion current. Faraday's law converts that current into a penetration rate, and galvanic coupling with unfavourable area ratios can accelerate attack dramatically.

Practice questions

1. Why must the total anodic current equal the total cathodic current on an isolated corroding metal? Answer: Charge cannot accumulate on the metal, so every electron released by oxidation must be consumed by reduction. 2. How would increasing the flow rate of aerated water affect the corrosion of steel under oxygen diffusion control? Answer: It thins the diffusion layer, raising the oxygen limiting current and therefore the corrosion rate. 3. A metal shows R p = 2000 Ω cm² with B = 0.026 V. Calculate i corr. Answer: i corr = 0.026 ÷ 2000 = 1.3 × 10⁻⁵ A/cm², or 13 μA/cm². 4. Explain why a small aluminium rivet in a large copper panel is a poor design. Answer: Aluminium becomes the anode for the whole copper cathode area, so a large cathodic current is concentrated on the small rivet, which corrodes rapidly.