Interactive Lecture
Understanding Material Degradation and Protection Strategies
by Dr Arnaldo Delli Carri
Corrosion is the gradual degradation of materials (usually metals) through electrochemical or chemical reactions with their environment. It is the process of materials returning to their thermodynamically stable oxide or compound state. Corrosion costs the global economy an estimated $2.5 trillion annually (3-4% of global GDP).
Left unmanaged, corrosion threatens safety, drains maintenance budgets, wastes raw materials, and constrains how a component can be designed:
Most corrosion is electrochemical in nature, requiring:
Oxidation at the Anode (metal dissolution):
\(\text{M} \rightarrow \text{M}^{n+} + ne^-\)
Example for iron: \(\text{Fe} \rightarrow \text{Fe}^{2+} + 2e^-\)
Reduction at the Cathode (consumes electrons):
In acidic solution: \(2\text{H}^+ + 2e^- \rightarrow \text{H}_2\) (hydrogen evolution)
In neutral/basic solution: \(\text{O}_2 + 2\text{H}_2\text{O} + 4e^- \rightarrow 4\text{OH}^-\) (oxygen reduction)
In aerated acidic solution: \(\text{O}_2 + 4\text{H}^+ + 4e^- \rightarrow 2\text{H}_2\text{O}\)
The standard electrode potential (E°) indicates the tendency of a metal to corrode:
More negative E° = More active (anodic) = Corrodes more easily
More positive E° = More noble (cathodic) = Resists corrosion
Selected values (V vs. SHE):
The practical corrosion behaviour in real environments (not standard conditions):
Noble (Cathodic) end:
Active (Anodic) end
Pourbaix diagrams show regions of corrosion, immunity, and passivation as functions of pH and electrode potential.
Three key regions:
Faraday's Law relates corrosion current to mass loss:
\(\text{Corrosion Rate} = \frac{I_{corr} \cdot M}{nF\rho A}\)
Where:
Written with the corrosion current density \(i_{corr}\) (lower-case, in mA/cm²), the area cancels and the penetration rate becomes:
Penetration rate (mm/year): \(\text{CR (mm/y)} = \frac{3.27 \times i_{corr}\,(mA/cm²) \times M}{n \times \rho}\)
Explore different corrosion mechanisms with a cross-section view showing electrochemical processes. Adjust pH and temperature to see how environmental conditions affect the corrosion rate via Faraday's law and Arrhenius kinetics.
A virtual electrochemistry experiment. Two metals are connected in an electrolyte — watch ions dissolve from the anode, electrons flow through the wire, and hydroxide form at the cathode. Use the area ratio slider to see why a small anode with a large cathode is dangerous.
Click or drag on the Pourbaix diagram to explore how pH and electrode potential determine whether iron will corrode, form a passive film, or remain immune. The dashed lines show the stability limits of water.
Mechanism: Corrosion occurs evenly over the entire exposed surface.
Characteristics: Most common, predictable, easiest to design for.
Examples: Rusting of steel, tarnishing of silver, atmospheric corrosion.
Prevention: Coatings, cathodic protection, material selection.
Mechanism: Two different metals in electrical contact in an electrolyte. The more active metal (anode) corrodes preferentially.
Characteristics: Corrosion rate depends on potential difference and anode/cathode area ratio.
Examples: Steel screw in copper plate, aluminium rivets in steel, brass fittings on steel pipes.
Prevention: Use similar metals, insulate dissimilar metals, use sacrificial anodes.
Mechanism: Localised attack forming small holes or cavities. Protective film breaks down at specific sites.
Characteristics: Extremely dangerous - small surface damage but deep penetration. Difficult to predict.
Examples: Stainless steel in chloride environments, aluminium in seawater.
Prevention: Alloy modification (Mo in stainless), avoid stagnant conditions, cathodic protection.
Mechanism: Occurs in shielded areas (gaps, under gaskets, deposits) where oxygen is depleted, creating local anodes.
Characteristics: Similar to pitting but in confined spaces. Autocatalytic process.
Examples: Under bolt heads, washers, gaskets, lap joints, deposits.
Prevention: Design to avoid crevices, use sealants, maintain clean surfaces.
Mechanism: Combined action of tensile stress and corrosive environment causes brittle cracking.
Characteristics: Sudden catastrophic failure with little warning. Highly specific to material-environment combinations.
Examples: Brass in ammonia, stainless steel in chlorides, high-strength aluminium in seawater.
Prevention: Stress relief, material selection, control environment, shot peening.
The most effective corrosion control strategy combines material selection, protective coatings, design and regular inspection. Prevention is always cheaper than failure.
| Concept | Equation | Description |
|---|---|---|
| Anodic Reaction | \(\text{M} \rightarrow \text{M}^{n+} + ne^-\) | Metal oxidation (dissolution) |
| Cathodic Reaction (acidic) | \(2\text{H}^+ + 2e^- \rightarrow \text{H}_2\) | Hydrogen evolution |
| Cathodic Reaction (aerated) | \(\text{O}_2 + 2\text{H}_2\text{O} + 4e^- \rightarrow 4\text{OH}^-\) | Oxygen reduction |
| Galvanic Potential | \(\Delta E = E_{cathode} - E_{anode}\) | Driving force for galvanic corrosion |
| Corrosion Rate | \(\text{CR} = \frac{3.27 \times i_{corr} \times M}{n \times \rho}\) | Penetration rate (mm/year), \(i_{corr}\) in mA/cm² |
A steel bolt (E° = -0.44 V) is used to fasten a copper plate (E° = +0.34 V) in a marine environment. Which metal will corrode, and what is the driving potential for this galvanic couple?
Show Solution
Given: \(E°_{Fe} = -0.44\) V, \(E°_{Cu} = +0.34\) V
Determining the anode and cathode:
The metal with more negative potential becomes the anode (corrodes).
Steel (Fe) is more negative → Steel is the anode
Copper (Cu) is more positive → Copper is the cathode
Driving potential:
\(\Delta E = E_{cathode} - E_{anode}\)
\(\Delta E = 0.34 - (-0.44) = 0.78\) V
The steel bolt corrodes preferentially while the copper plate is protected. A driving potential of 0.78 V is large and will cause significant galvanic corrosion.
Engineering solution: Use stainless steel bolts, or insulate the bolt from the copper plate.
Answer: Steel (Fe) is the anode and corrodes; driving potential ΔE = 0.78 V.
A corrosion current density of 0.1 mA/cm² is measured on an iron sample (M = 55.85 g/mol, ρ = 7.87 g/cm³, n = 2). Calculate the corrosion rate in mm/year.
Show Solution
Given: \(i_{corr} = 0.1\) mA/cm²
\(M = 55.85\) g/mol, \(\rho = 7.87\) g/cm³, \(n = 2\)
Using the corrosion rate equation:
\(\text{CR (mm/y)} = \displaystyle \frac{3.27 \times i_{corr} (mA/cm²) \times M}{n \times \rho}\)
\(\text{CR} = \displaystyle \frac{3.27 \times 0.1 \times 55.85}{2 \times 7.87}\)
\(\text{CR} = \displaystyle \frac{18.26}{15.74} = 1.16\) mm/year
This is a very high corrosion rate. For a 5 mm thick component, complete penetration would occur in about 4 years.
Answer: CR = 1.16 mm/year.
A steel pipeline requires cathodic protection. Zinc anodes (E° = -0.76 V) are to be used. The pipeline surface area is 100 m². The required protection current density is 0.02 A/m². If each zinc anode can deliver 0.5 A for 5 years before replacement, how many anodes are needed?
Show Solution
Given:
Step 1: Calculate total protection current required:
\(I_{total} = i \times A = 0.02 \times 100 = 2.0\) A
Step 2: Calculate number of anodes needed:
\(N = \displaystyle \frac{I_{total}}{I_{anode}} = \frac{2.0}{0.5} = 4\) anodes
Four zinc anodes supply the calculated protection current for the pipeline.
Engineering considerations:
Answer: N = 4 anodes (install 5–6 with a safety margin).