7.3 Pourbaix Diagrams and Corrosion Chemistry

Pourbaix Diagrams

Understanding Pourbaix Diagrams

A Pourbaix diagram (also called a potential-pH diagram) maps out the thermodynamic stability of different metal species in aqueous solution. It answers a straightforward question: given a particular electrode potential and pH, what's the most stable form of the metal?

• The vertical axis represents electrode potential ($E_h$), measured in volts vs. the standard hydrogen electrode (SHE).
• The horizontal axis shows pH, typically ranging from 0 to 14.
• Each region on the diagram corresponds to a predominance area where one species (metal, metal ion, oxide, hydroxide, etc.) is thermodynamically favored.

These diagrams exist for specific metal-water systems. The iron-water and copper-water Pourbaix diagrams are the ones you'll encounter most often in this course.

Key Components and Features

The boundaries between predominance areas represent equilibrium conditions between adjacent species. The type of line tells you what variables are involved in that equilibrium:

• Horizontal lines represent reactions that depend only on potential (pure electron-transfer reactions, no $H^+$ or $OH^-$ involved).
• Vertical lines represent reactions that depend only on pH (acid-base equilibria with no electron transfer).
• Diagonal lines represent reactions involving both electrons and protons. These are the most common boundaries you'll see.

Two dashed lines appear on every Pourbaix diagram regardless of the metal:

• The upper dashed line (a) marks the equilibrium for $O_2$ evolution: $2H_2O \rightarrow O_2 + 4H^+ + 4e^-$. Above this line, water is thermodynamically unstable with respect to oxygen evolution.
• The lower dashed line (b) marks the equilibrium for $H_2$ evolution: $2H^+ + 2e^- \rightarrow H_2$. Below this line, water is thermodynamically unstable with respect to hydrogen evolution.

The region between these two lines is where liquid water is stable. Most practical corrosion scenarios fall within this window.

Three Stability Regions

On a typical Pourbaix diagram for a metal like iron, you can identify three broad types of regions:

1. Corrosion regions — The dissolved metal ion (e.g., $Fe^{2+}$, $Fe^{3+}$) is the stable species. The metal actively dissolves.
2. Passivation regions — A solid oxide or hydroxide (e.g., $Fe_2O_3$, $Fe(OH)_2$) is stable on the surface. This layer can protect the metal from further attack.
3. Immunity regions — The metal itself (e.g., $Fe$) is the thermodynamically stable species. No corrosion occurs because oxidation is not favored.

Applications and Limitations

Pourbaix diagrams are widely used to:

• Predict whether a metal will corrode, passivate, or remain immune under given conditions
• Guide the selection of corrosion prevention strategies (e.g., shifting conditions into the immunity or passivation region)
• Interpret behavior in geochemistry and environmental aqueous systems

However, they have real limitations you should keep in mind:

• They are purely thermodynamic. A diagram might predict corrosion, but the reaction could be extremely slow in practice. Kinetics are not captured.
• They assume pure substances in contact with pure water. Effects of complexing agents, alloying elements, or mixed electrolytes are not accounted for.
• They don't address mixed potential theory or the actual corrosion potential a metal adopts in a real solution.

Corrosion and Passivation

Corrosion Mechanisms and Types

Corrosion is the degradation of a metal through electrochemical reactions with its environment. At the atomic level, metal atoms at the surface lose electrons (oxidation) and enter solution as ions, while a corresponding reduction reaction occurs elsewhere.

The main types of corrosion you should know:

• Uniform corrosion — Metal dissolves evenly across the entire exposed surface. It's the most common type and the easiest to predict and manage.
• Pitting corrosion — Highly localized attack that produces small holes or pits. Dangerous because it can penetrate deeply while the surrounding surface looks fine.
• Crevice corrosion — Occurs in confined spaces (under gaskets, bolt heads, overlapping joints) where stagnant solution develops a different chemistry from the bulk.
• Galvanic corrosion — Results from electrical contact between two dissimilar metals in an electrolyte. The more active metal (more negative $E^0$) corrodes preferentially, acting as the anode.
• Stress corrosion cracking — A combination of tensile mechanical stress and a corrosive environment causes crack propagation that neither factor would cause alone.

Passivation and Corrosion Prevention

Passivation is the formation of a thin, adherent oxide or hydroxide layer on a metal surface that acts as a barrier between the metal and its environment. This layer dramatically slows further corrosion.

Some metals passivate spontaneously. Aluminum, for example, forms a thin $Al_2O_3$ layer almost instantly in air, which is why aluminum doesn't visibly corrode under normal conditions despite being thermodynamically quite reactive. Chromium behaves similarly, and its presence in stainless steel (at least ~10.5% Cr by mass) is what gives stainless steel its corrosion resistance.

Passivation can also be induced through surface treatments (chemical passivation baths, anodizing) or by alloying with elements that promote stable oxide formation.

Two other key protection concepts:

• Immunity — Shifting the metal's potential into the immunity region of the Pourbaix diagram, where the metal itself is the stable species. This is the basis of cathodic protection.
• Cathodic protection — An external intervention that forces the metal to act as the cathode in an electrochemical cell, preventing its oxidation. More detail on this below.

Factors Influencing Corrosion and Passivation

• pH — Affects the stability of passive films directly. Many oxide films dissolve in strongly acidic or strongly basic solutions, which is visible on Pourbaix diagrams as the passivation region narrowing at extreme pH values.
• Dissolved oxygen — Often serves as the cathodic reactant ($O_2 + 2H_2O + 4e^- \rightarrow 4OH^-$), so higher $O_2$ concentration generally accelerates corrosion.
• Chloride ions — Particularly aggressive toward passive films. $Cl^-$ ions can penetrate or undermine oxide layers, initiating pitting. This is why seawater and de-icing salts are so corrosive.
• Temperature — Higher temperatures increase reaction rates and can destabilize passive films.
• Mechanical stress — Can crack passive films, exposing fresh metal to the corrosive environment.
• Metallurgical factors — Grain boundaries, inclusions, and compositional inhomogeneities create local electrochemical cells that accelerate corrosion.

Electrochemistry Fundamentals

The Nernst Equation

The Nernst equation connects the reduction potential of an electrochemical reaction to the concentrations (or activities) of the species involved. It's the quantitative backbone of Pourbaix diagrams.

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

where:

• $E$ = cell potential under actual conditions (V)
• $E^0$ = standard electrode potential (V)
• $R$ = universal gas constant (8.314 J mol$^{-1}$ K$^{-1}$)
• $T$ = temperature (K)
• $n$ = number of electrons transferred
• $F$ = Faraday constant (96,485 C mol$^{-1}$)
• $Q$ = reaction quotient

At 25°C, this simplifies to the form you'll use most often:

$E = E^0 - \frac{0.0592}{n} \log Q$

Connection to Pourbaix diagrams: Every boundary line on a Pourbaix diagram is derived from the Nernst equation applied to a specific equilibrium. For reactions involving $H^+$, the $\log Q$ term introduces pH dependence, which is why those boundary lines are diagonal. For reactions with no $H^+$ involvement, the line is horizontal (potential-dependent only).

Cathodic Protection Techniques

Cathodic protection works by shifting a metal's potential into the immunity region on its Pourbaix diagram. There are two main approaches:

1. Sacrificial anode method — A more active metal (more negative $E^0$) is electrically connected to the metal you want to protect. The sacrificial anode corrodes instead. Zinc and magnesium are common choices for protecting steel because they sit lower in the galvanic series.

2. Impressed current cathodic protection (ICCP) — An external DC power source forces current onto the protected metal, making it the cathode. An inert or semi-inert anode (e.g., platinized titanium, graphite) completes the circuit. This method is used when sacrificial anodes alone can't supply enough current.

Common applications include underground pipelines, ship hulls, storage tanks, and reinforced concrete structures. ICCP systems require monitoring and control to maintain the correct protection potential without overprotecting (which can cause hydrogen embrittlement or coating damage).