Understanding Electrochemical Corrosion

• Corrosion is the deterioration of a material due to chemical reactions with its environment.
• Electrochemical corrosion involves the flow of electrons and ions.
• It typically occurs with metals in aqueous environments.

Electrochemical Cell Formation

• An electrochemical cell consists of an anode, cathode, electrolyte, and an external circuit.
• The anode is where oxidation (corrosion) occurs, releasing electrons.
• The cathode is where reduction occurs, consuming electrons.
• The electrolyte is an ionically conductive solution.
• The external circuit provides a path for electron flow between the anode and cathode.

Half-Cell Reactions

• Oxidation at the anode:
  • Metal atoms lose electrons and become ions.
  • Example: Fe -> Fe2+ + 2e-
• Reduction at the cathode:
  • Electrons are consumed by a reducing agent.
  • Common examples:
    • Oxygen reduction in acidic solution: O2 + 4H+ + 4e- -> 2H2O
    • Oxygen reduction in neutral or alkaline solution: O2 + 2H2O + 4e- -> 4OH-
    • Hydrogen evolution in acidic solution: 2H+ + 2e- -> H2

Standard Electrode Potential

• The standard electrode potential (E°) is the potential of a half-cell under standard conditions (25°C, 1 M concentration, 1 atm pressure).
• It is a measure of the tendency of a species to be reduced.
• A more positive E° indicates a greater tendency for reduction.
• The standard hydrogen electrode (SHE) is the reference electrode with E° = 0 V.
• The electrochemical series ranks metals based on their standard electrode potentials.
• Metals with lower (more negative) E° are more easily oxidized (more prone to corrosion).

Nernst Equation

• The Nernst equation relates the electrode potential (E) to the standard electrode potential (E°) and the activities (or concentrations) of the species involved in the half-cell reaction.
• The Nernst equation is: E = E° - (RT/nF) * ln(Q)
  • R is the ideal gas constant (8.314 J/mol·K).
  • T is the temperature in Kelvin.
  • n is the number of moles of electrons transferred in the cell reaction.
  • F is the Faraday constant (96485 C/mol).
  • Q is the reaction quotient.
• For the half-cell reaction: Mn+ + ne- -> M, the Nernst equation simplifies to: E = E° + (RT/nF) * ln([Mn+])

Types of Corrosion

• Uniform Corrosion:
  • Corrosion occurs evenly over the entire surface.
• Galvanic Corrosion:
  • Occurs when two dissimilar metals are in contact in the presence of an electrolyte.
  • The more active metal (lower E°) corrodes preferentially (anode).
  • The less active metal (higher E°) is protected (cathode).
• Crevice Corrosion:
  • Occurs in narrow gaps or crevices where stagnant electrolyte conditions exist.
  • Differential aeration leads to oxygen depletion inside the crevice, making it anodic.
• Pitting Corrosion:
  • Localized corrosion that results in small pits or holes.
  • Often occurs on passive metals due to local breakdown of the passive film.
• Intergranular Corrosion:
  • Corrosion occurs along grain boundaries.
  • Often due to segregation of impurities at grain boundaries or formation of precipitates.
• Stress Corrosion Cracking (SCC):
  • Cracking of a metal under the combined action of tensile stress and a corrosive environment.
• Erosion Corrosion:
  • Corrosion enhanced by the relative motion of a corrosive fluid and a metal surface.
• Dealloying:
  • Selective removal of one or more components from an alloy.
  • Example: Dezincification of brass (removal of zinc).

Passivity

• Passivity is the phenomenon where a metal forms a thin, protective oxide film on its surface, which significantly reduces the corrosion rate.
• Metals like stainless steel, aluminum, and titanium exhibit passivity.
• The passive film can be disrupted by certain ions (e.g., chloride ions), leading to localized corrosion like pitting.

Factors Affecting Corrosion Rate

• Temperature: Higher temperatures generally increase corrosion rates.
• Electrolyte Concentration: Increased concentration can accelerate corrosion up to a limiting point.
• pH: Acidic or alkaline environments can accelerate corrosion depending on the metal.
• Oxygen Concentration: Oxygen is a common cathodic reactant.
• Flow Rate: Higher flow rates can increase corrosion due to erosion corrosion.
• Galvanic Coupling: Contact between dissimilar metals promotes galvanic corrosion.
• Presence of Inhibitors: Corrosion inhibitors can reduce corrosion rates.

Corrosion Prevention Methods

• Material Selection: Choosing corrosion-resistant materials.
• Protective Coatings: Applying coatings to isolate the metal from the environment.
  • Examples: Paint, galvanizing (zinc coating), electroplating.
• Cathodic Protection: Making the metal a cathode, thereby preventing oxidation.
  • Sacrificial Anode: Connecting a more active metal (e.g., zinc, magnesium) to the metal structure.
  • Impressed Current: Applying an external DC current to make the metal cathodic.
• Anodic Protection: Making the metal passive by applying an external anodic current (less common).
• Corrosion Inhibitors: Adding chemicals to the environment to reduce corrosion rates.
• Design Considerations: Avoiding crevices, sharp corners, and stagnant areas.

Corrosion Measurement Techniques

• Weight Loss Method: Measuring the weight loss of a sample over time.
• Electrochemical Techniques:
  • Potentiodynamic Polarization: Measuring the current as a function of potential.
  • Electrochemical Impedance Spectroscopy (EIS): Measuring the impedance of the metal-electrolyte interface.
• Visual Inspection: Detecting signs of corrosion visually.
• Non-Destructive Testing (NDT): Using techniques like ultrasonic testing or radiography to detect internal corrosion.