Corrosion

date2026-07-24tags:chem:

Pourbaix diagrams can predict a metal or alloy's corrosion behavior. They need the chemical composition of the environment. Real-world environments change over time. Corrosion depends on the microenvironment at the metal surface. It does not depend on the bulk environment. Experiments on the nominal environment are inadequate. They ignore flow, local pH variations, deposits, and galvanic coupling. They cannot predict long-term corrosion performance.

Corrosion is a redox reaction. The metal oxidizes: $$ M \rightarrow M^{n+} + ne^- $$

Standard Electrode Potentials

An electrochemical cell has a potential difference between its two electrodes. Electrons flow from the more negative (or less positive) electrode to the more positive one. Conventional current flows in the opposite direction. A voltmeter measures the overall cell potential. It cannot measure individual electrode potentials. You need a reference electrode.

The standard hydrogen electrode (SHE) is the primary historical reference. It bubbles hydrogen gas over a platinum electrode in acidic solution. Its potential is zero at all temperatures. Some people call it the normal hydrogen electrode (NHE). The SHE requires unit activity for hydrogen ions. It requires unit fugacity for hydrogen gas. These conditions are hard to meet. The SHE is often impractical because it needs hydrogen gas.

The reversible potential (E) is the potential of an electrode measured against the SHE. The standard electrode potential (E°) is the reversible potential under standard conditions. Tables of standard electrode potentials compare electrodes to a reference like the SHE.

Nernst Equation

The Nernst equation relates the potential of a cell to ion concentrations. It uses the relationship between Gibbs free energy change (ΔG) and the reaction quotient (Q). For a reaction aA + bB → mM + nN (capital letters are reactants and products, lowercase letters are coefficients), the Nernst equation uses ΔG = -nFE and ΔG° = -nFE°.

$$ Q = \frac{{\left(a_M\right)}^m \cdot {\left(a_N\right)}^n}{{\left(a_A\right)}^a \cdot {\left(a_B\right)}^b} $$

$$ E = E^0 - \frac{RT}{nF} \ln Q_{reaction} $$

At 25°C (298.15 K), the equation simplifies: $$ E = E^0 - \frac{0.059}{n} \log_{10} Q_{reaction} $$

The electrode potential (E) is the difference between the half-cell and the SHE. You can apply the Nernst equation to two half-reactions. Then combine the results to calculate the overall cell potential: $$ E_{cell} = \left( E_{cathode}^0 - E_{anode}^0 \right) - \frac{0.059}{n} \log_{10} Q_{reaction} $$

Measuring E_corr

The corrosion potential $E_{corr}$ is key in corrosion studies. Measure it as the potential difference between the corroding metal and a reference electrode. Both the magnitude and the sign of the voltage matter. The reference electrode connects to the low point or instrumental ground. Some equipment manufacturers use the opposite convention.

Low corrosion rates produce low current. Low currents are hard to measure. Corrosion-resistant samples make this worse. Many factors are hard to control during corrosion measurement. Use a minimum of 3 replicates for corrosion measurements. A relative standard deviation of 15% is excellent. Electrochemical instruments are highly accurate.

Open circuit potential

The open circuit potential $E_{OC}$ is the voltage between the working electrode (the metal) and the reference electrode. Both electrodes sit in the electrolyte. The $E_{OC}$ is a mixed potential. It depends on electrochemical half-reactions, not on the potentiostat. It is the potential at which corrosion occurs under normal conditions. It is a critical parameter.

The $E_{OC}$ is the start for almost all electrochemical corrosion experiments. A stable $E_{OC}$ indicates steady-state corrosion. This stability is a prerequisite for experiments. It shows that perturbation experiments can begin. Achieving a stable $E_{OC}$ can take minutes or days. Do not start experiments before confirming stability.

No current flows at the open circuit potential. Applying a voltage positive to the EOC accelerates oxidation (corrosion). Applying a voltage negative to the EOC accelerates reduction of solution species.

E_{OC} vs E_{corr}

"EOC" and "Ecorr" are often interchangeable. "Ecorr" can refer to the zero-current point during an experiment.

The EOC value alone is not a strong predictor of corrosion behavior.

Mixed-Potential Theory

$E_{corr}$ is a mixed potential. Wagner and Traud formalized this in 1938. A corroding surface has at least two independent partial reactions. These reactions run at the same interface. The first is anodic dissolution of the metal. Others are cathodic reactions. They consume electrons from the metal (oxygen reduction, hydrogen evolution, etc.). Wagner-Traud theory treats these as superimposed but independent. Charge cannot accumulate. Total anodic current equals total cathodic current. $E_{corr}$ is the potential where this balance holds. It is not the equilibrium potential of any single reaction. No net reaction is at equilibrium.

This theory justifies treating $E_{OC}$/$E_{corr}$ as real. It comes from superposition of electrochemical processes. Each follows its own kinetics.

Tafel Kinetics and Evans Diagrams

Near $E_{corr}$, partial reactions follow Butler-Volmer kinetics. Far from equilibrium potential (beyond about 50 mV of overpotential) the Butler-Volmer expression becomes the Tafel limit. Overpotential grows linearly with the logarithm of current density. Plot $\log|i|$ vs $E$ for a straight line. This line is the Tafel slope for each partial reaction: $$ \eta = \beta \log_{10}\left(\frac{i}{i_0}\right) $$

The symbols $\beta_a$ and $\beta_c$ are anodic and cathodic Tafel slopes. Their unit is V/decade. The symbol $i_0$ is the exchange current density of that reaction.

An Evans diagram overlays anodic and cathodic $\log|i|$-$E$ lines. They cross at equal anodic and cathodic current. This intersection is $E_{corr}$. The current density at the intersection is $i_{corr}$ (corrosion current density). Evans diagrams are visually useful. They show whether $E_{corr}$ shifted due to metal dissolution or due to oxidizer availability. A small change in one partial reaction's kinetics moves $E_{corr}$ a lot if the other reaction's Tafel slope is shallow.

Extrapolating the linear (Tafel) region of a polarization curve back to $E_{corr}$ estimates $i_{corr}$. This method is Tafel extrapolation. The small-perturbation linear polarization resistance method and the Stern-Geary equation are in the LPR notes. This page covers only the kinetic picture.

Thermodynamics vs. kinetics

Standard electrode potentials, the Nernst equation, and Pourbaix diagrams are thermodynamic. They describe which reaction is favorable. They show direction and equilibrium extent. They do not show rate. A metal can be thermodynamically unstable. It can corrode extremely slowly if kinetics are sluggish. Passivation is the everyday example. A thermodynamically favored oxide forms a kinetic barrier. It suppresses further reaction. The practical corrosion rate drops by orders of magnitude. Conversely, a weakly favored thermodynamic reaction can proceed at an appreciable rate if the kinetic barrier is low. Tafel/Evans analysis and $i_{corr}$ show rate. Standard potentials and Pourbaix diagrams show tendency. Treating either as a substitute for the other is the most common mistake in corrosion data analysis.

Pourbaix Diagrams (Potential-pH)

A Pourbaix diagram (named after Marcel Pourbaix) maps thermodynamically stable phases. It covers one metal and water. It uses fixed temperature and species activities. Stable phases are bare metal (immunity), dissolved ion (corrosion), or solid oxide/hydroxide (passivation). The diagram plots electrode potential (y-axis) against pH (x-axis). It uses the Nernst equation and solubility data. It is a purely thermodynamic tool. Read the caveat above. Do not treat a stable-oxide region as a guarantee of low corrosion rate.

Two limitations matter in practice:

Bulk vs microenvironment. A Pourbaix diagram uses specified bulk conditions. Common values: $10^{-6}$ M dissolved species, room temperature, 1 atm. Inside a pit or crevice, local chemistry differs from bulk. Metal-ion hydrolysis acidifies the occluded solution. Chloride migrates in to balance charge. Oxygen depletes from the surrounding cathodic reaction. The bulk diagram can show "passive" at nominal pH and potential. The pit-tip environment may sit in the "corrosion" region of a local diagram. This positive feedback drives pit and crevice propagation.
No kinetics, no rate. The diagram does not show transition speed. It does not show how protective a "stable" oxide is against ion transport. It does not show whether the system reaches equilibrium. It ignores mass-transport limitations. It assumes species activity is at the bulk value.
**Temperature.** A default Pourbaix diagram is at 25°C. Phase boundaries shift with temperature. The Nernst slope ($2.303RT/F$) changes. Equilibrium constants change. A passive region at 25°C can shrink or disappear at 70°C. Stable solid phases can also change. Hydrated oxides at room temperature may dehydrate to a different oxide. This different oxide has different protectiveness. Treat the 25°C diagram as a starting hypothesis. Experiment confirms or refutes it.

Passivity and the passive film

A reactive metal corrodes at a negligible rate is passive. The Pourbaix diagram favors corrosion. Passivity is a kinetic phenomenon. A thin oxide or hydroxide film (typically 1-10 nm) forms on the surface. It acts as a diffusion barrier. It slows metal-ion transport outward and oxidizer transport inward. It reduces rates by orders of magnitude.

Stainless steel is the classic example. The bulk alloy is mostly iron. Iron is thermodynamically unstable in aerated water. Yet stainless steel corrodes far slower than thermodynamics predicts. A Cr-rich oxide film (mixed Fe-Cr spinel or Cr₂O₃) is the barrier. The film is not static. It dissolves slowly at the film/solution interface. It reforms at the metal/film interface. The passive state is a steady-state condition. It is not equilibrium. Film thickness and composition respond to potential and solution chemistry.

Critical breakdown

Passive films fail when local chemistry exceeds their stability range:

Chloride ions (and other halides, Br⁻, I⁻) cause breakdown. Chloride adsorbs at the oxide surface. It penetrates film defects. It forms metal-chloride complexes at the metal/film interface. These complexes undermine the film locally. The result is a metastable pit. It either repassivates or grows into a stable pit. The pitting potential $E_{pit}$ is the potential below which pitting does not initiate. It is a key parameter alongside $E_{corr}$.
Local acidification inside a pit or crevice lowers the pH. Metal-ion hydrolysis causes this. The acidified, chloride-rich environment prevents repassivation. The pit grows. The surrounding surface remains passive.
Mechanical damage removes the film. Abrasion, cavitation, or erosion expose bare metal. Repassivation kinetics determine the result. Fast repassivation causes a transient. The OCP page describes the OCP spike-and-recovery pattern. Slow repassivation starts sustained local corrosion.

Film analysis methods

XPS (X-ray photoelectron spectroscopy) and AES (Auger electron spectroscopy) are separate from the DC electrochemical techniques on this page. They give elemental composition and chemical state of the film. They show composition as a function of depth. They combine with sputter depth profiling. Mott-Schottky analysis probes electronic properties of the film. It measures film capacitance versus potential. It determines whether the film acts as an n-type or p-type semiconductor. It measures donor and acceptor density. This density correlates with film protectiveness.