Corrosion and its Prevention - Comprehensive Study Guide

Introduction to Corrosion and Metal Nature

• Metallic Properties: Metals are electropositive. They have a natural tendency to lose electrons to complete their octet and reach a stable state by forming compounds.

• Natural State: Due to their reactive nature, most metals exist in nature in a combined state as oxides, carbonates, sulphides, or chlorides.

• Definition of Corrosion:     * Any process of destruction and consequent loss of a solid metallic material through unwanted chemical or electrochemical attack by its environment, starting at its surface.     * It occurs when pure metal is exposed to environments containing moisture, dry gases, or liquids, causing the exposed surface to form compounds with environmental elements, leading to destruction.

• Common Examples of Corrosion:     * Iron (Fe): When exposed to atmospheric conditions, it forms a layer of reddish scale, causing the iron to become weak.     * Copper (Cu): When exposed to moist air containing $CO_2$, a greenish film of carbonate $CuCO_3 + Cu(OH)_2$ forms on the surface.

Types of Corrosion

• Corrosion is classified into two primary types based on the factors responsible:     1. Atmospheric Corrosion or Chemical Corrosion: Also known as dry corrosion (e.g., rusting of iron due to oxygen).     2. Electrochemical Corrosion or Immersed Corrosion: Also known as wet corrosion (e.g., rusting of iron in the presence of a solution).

• Sub-categories of Atmospheric Corrosion:     * Corrosion due to $O_2$.     * Corrosion due to other gases (e.g., $Cl_2$, $CO_2$, $SO_2$, $H_2S$, etc.).     * Liquid metal corrosion.     * High-temperature corrosion.

• Sub-categories of Electrochemical Corrosion:     * Galvanic cell corrosion.     * Concentration cell corrosion.

Atmospheric Corrosion (Dry Corrosion)

• Definition: Occurs when metal comes into direct contact with atmospheric gases such as oxygen, halogens, hydrogen sulphide, carbon dioxide, or sulphur dioxide without the presence of moisture.

• Oxidation Corrosion (Corrosion due to Oxygen): The metal surface reacts with oxygen gas in the air.

• Mechanism of Atmospheric Corrosion:     * Oxidation at Metal Surface: The metal loses its valence electrons and gets oxidized to form metal ions.         * $2M \rightarrow 2M^{n+} + 2ne^-$ (Oxidation / Loss of electron)     * Reduction of Oxygen: Oxygen gas accepts the valence electrons to form oxide ions.         * $2ne^- + \frac{n}{2}O_2 \rightarrow nO^{2-}$ (Reduction / Gain of electron)     * Formation of Metal Oxide Layer: Metal ions combine with oxide ions to form a layer on the surface.         * $2M^{n+} + nO^{2-} \rightarrow [M_2O_n]$ (Metal Oxide)

• Diffusion Process: A thin oxide layer forms at the meeting point of metal and oxygen. Further corrosion depends on the diffusion of ions. Usually, the outward diffusion of metal ions is fast, while the inward diffusion of oxide ions through the scale is slow.

Types of Metal Oxide Layers

• Unstable Layer: The oxide layer formed decomposes back into metal and oxygen immediately. In this case, oxidation corrosion is not possible.     * Examples: Silver ($Ag$), Gold ($Au$), and Platinum ($Pt$).

• Volatile Layer: The oxide layer evaporates as soon as it forms, leaving the underlying metal continuously exposed for further attack. This leads to rapid and continuous corrosion.     * Examples: Molybdenum oxide ($MoO_3$) and Tin chloride ($SnCl_4$).

• Stable Layers: These layers adhere tightly to the parent metal.     * Porous Layer: Occurs when the volume of the oxide formed is less than the volume of the metal consumed. The layer does not fully cover the surface, resulting in stress, cracks, and pores. This allows oxygen free access to the fresh metal surface below.         * Examples: Alkali and alkaline earth metals such as $Li$, $K$, $Na$, $Ca$, and $Mg$.     * Non-porous Layer: Occurs when the volume of the oxide film formed is more than the volume of the metal. This film adheres tightly, cutting off the diffusion of oxygen to the underlying metal and acting as a protective coating.         * Example: Aluminum ($Al$), where the protective $Al_2O_3$ film prevents further corrosion.

Electrochemical Corrosion (Wet Corrosion)

• Definition: Corrosion brought about through ionic reactions in the presence of a solution acting as a conducting medium. It typically occurs at a solid-liquid interface or when two dissimilar metals are in contact.

• Necessary Conditions:     * Formation of anodic and cathodic parts.     * Electrical contact between anode and cathode for electron movement.     * Presence of a conducting medium (electrolyte).

• Common Examples:     1. Corrosion at riveted joints.     2. Steel pipes connected to copper plumbing.     3. Steel screws in contact with brass in marine environments.     4. Fencing wires corroding under the joints.

Mechanism: Evolution of Hydrogen

• Environment: Occurs when metals are exposed to acidic environments (e.g., industrial waste or non-oxidizing acids like $HCl$).

• Electrochemical Series: Metals located above hydrogen in the electrochemical series tend to dissolve in acidic solutions with the simultaneous evolution of $H_2$ gas.

• Configuration: Usually involves a large anodic area and a small cathodic area.

• Example: A steel tank (anode) containing acidic waste with small copper scraps (cathode).     * Anode Reaction: $Fe \rightarrow Fe^{2+} + 2e^-$ (Oxidation)     * Cathode Reaction: $2H^+ + 2e^- \rightarrow H_2$ (Reduction)     * Net Reaction: $Fe + 2H^+ \rightarrow Fe^{2+} + H_2$

Mechanism: Absorption of Oxygen

• Environment: Rusting of iron in neutral aqueous solutions (e.g., $NaCl$) in the presence of atmospheric oxygen.

• Configuration: Small anodic area and large cathodic area. Cracks in the iron oxide film expose small metal areas that act as anodes.

• Process:     * Anode (Small Area): $Fe \rightarrow Fe^{2+} + 2e^-$ (Oxidation)     * Cathode (Large Area): Electrons flow to the cathodic area and react with dissolved oxygen and water.         * $O_2 + 2H_2O + 4e^- \rightarrow 4OH^-$ (Reduction)     * Formation of Precipitate: $Fe^{2+}$ and $OH^-$ ions meet in solution to form ferrous hydroxide.         * $Fe^{2+} + 2OH^- \rightarrow Fe(OH)_2 \text{ (Precipitate)}$

• Oxidation of the Product:     * Yellow Rust: If oxygen is abundant, ferrous hydroxide is oxidized to ferric hydroxide ($Fe(OH)_3$), known as yellow rust with the formula $Fe_2O_3 \times 2H_2O$.         * $4Fe(OH)_2 + O_2 + 2H_2O \rightarrow 4Fe(OH)_3$     * Black Magnetite: If oxygen supply is limited, the product may be black anhydrous magnetite ($Fe_3O_4$). This precipitate can fill cracks and prevent further corrosion.

Factors Affecting Corrosion Rate

Nature of the Metal

1. Nature of Oxide Film: Non-porous films ($Al_2O_3$) stop corrosion; porous films (iron rust) allow continuous corrosion; volatile films ($MoO_3$) lead to very fast corrosion.

2. Position in Galvanic Series: More active metals (top of series) act as anodes and corrode. The rate increases with the potential difference between two metals in contact.

3. Purity of Metal: Impurities increase the rate of corrosion. Increasing metal purity improves corrosion resistance.

4. Relative Area of Anode and Cathode: Corrosion is more rapid when the anode is small and the cathode is large. The rate is proportional to:     * $\text{Rate of corrosion} \times \frac{\text{Area of Cathode}}{\text{Area of Anode}}$

5. Solubility of Corrosion Product: Insoluble products act as barriers; soluble products allow fast corrosion.

6. Hydrogen Overvoltage: Extra voltage needed to evolve hydrogen. High overvoltage diminishes corrosion; impurities with low hydrogen overvoltage increase corrosion.

7. Physical State: Coarser grains (larger grain size) lead to lower corrosion rates. Hardening processes can improve grain size.

Nature of the Environment

1. Temperature: Rate of corrosion increases with temperature for both chemical and electrochemical processes.     * $\text{Rate of corrosion} \times \text{Temperature}$

2. Humidity: Higher moisture acts as a conducting medium, setting up galvanic cells. Electrochemical corrosion is faster than atmospheric corrosion.     * $\text{Rate of corrosion} \times \text{Humidity}$

3. Effect of pH: Acidic media are more corrosive than alkaline or neutral media (Acidic > Basic > Neutral).

4. Gaseous Impurities: Corrosive gases like $H_2S$, $CO_2$, $Cl_2$, $SO_2$, and acid fumes ($HCl$, $H_2SO_4$) increase corrosion rates.

5. Suspended Particles:     * Chemically active particles ($NaCl$, $Na_2SO_4$) absorb moisture and act as electrolytes.     * Chemically inactive particles (charcoal) absorb moisture and deleterious gases like $SO_2$.

6. Conductance of Medium: High conductance (sea water, mineralized soil) leads to rapid electrochemical corrosion compared to dry soil or clay.

7. Differential Aeration: Areas with less oxygen concentration act as anodes and corrode, while well-aerated parts are protected.

Corrosion Prevention and Design

• Material Selection:     * Noble metals are resistant but uneconomical.     * Alloying improves strength and resistance (e.g., Stainless Steel with $Cr$ forms a resistant oxide film; Cupro Nickel: $70\% Cu + 30\% Ni$ with $0.2\% Fe$ for condenser tubes).     * Active metals should be insulated from cathodic metals.

• Proper Designing:     * Avoid sharp bends and stresses.     * Avoid screws, nuts, and bolts where possible; prefer proper welding.     * Surfaces of joining parts should be smooth.

Electrochemical Protection Methods

• Principle: Force the metal to be protected to behave as a cathode.

1. Sacrificial Anodic Protection:     * Metal structure is connected via insulated wire to a more active/anodic metal (e.g., $Zn$, $Mg$, $Al$).     * The active metal (Sacrificial Anode) corrodes and is replaced when consumed.     * Applications: Buried pipelines, underground cables, water tanks, ship hulls.

2. Impressed Current Cathodic Protection:     * D.C. current (from a battery or rectifier) is applied in the opposite direction to nullify corrosion current.     * The negative terminal is connected to the structure; the positive terminal to an insoluble anode (graphite rod, stainless steel) buried in the soil.     * Applications: Oil/water pipes, transmission lines, open water-box coolers.

Inorganic and Metal Coatings

• Inorganic Coatings:     1. Chemical Oxide Coating: Base metal is treated with oxidizing solution/gas to thicken the oxide film (e.g., $Al$, $Zn$, $Ni$, $Cr$).     2. Phosphate Coating: Reaction with phosphoric acid and metal phosphates ($Fe, Mn, Zn$). Formulated by dipping, spraying, or brushing.     3. Chromate Coating: Produced using potassium chromate acid; amorphous and non-porous. Used for $Zn$, $Cd$, $Al$, $Mg$.

• Metal Coatings:     * With More Active Metal: Base metal is cathode, coating is anode (e.g., Galvanization).     * With Less Active Metal / Noble Metal: Coating has higher resistance (e.g., Tinning).

Hot Dipping: Galvanizing and Tinning

• Overview: Dipping metal into molten metal covered by flux. Used for low melting metals like $Zn$ ($419^\text{o}C$) and $Sn$ ($232^\text{o}C$).

• Galvanizing (Zinc Coating):     1. Cleaning: $dil. H_2SO_4$ for $15\text{--}20$ minutes at $60\text{--}90^\text{o}C$.     2. Flux: Ammonium chloride ($NH_4Cl$) to prevent oxidation.     3. Dipping: Molten zinc at $425\text{--}430^\text{o}C$.     4. Finishing: Rolled for thickness, annealed at $650^\text{o}C$, cooled slowly.     5. Use: Roofing, fencing, pipes. Not for food (zinc forms poisonous compounds with organic acids).

• Tinning (Tin Coating):     1. Cleaning: $dil. H_2SO_4$.     2. Flux: Zinc chloride ($ZnCl_2$).     3. Dipping: Molten tin, then vegetable oil (palm oil) bath to prevent oxidation.     4. Use: Storing food (ghee, pickles), cooking utensils, refrigeration equipment.

Comparison: Galvanizing vs. Tinning

Specialized Coating Techniques

• Electroplating: Applying metal coating using electric current.     * Article = Cathode; Pure coating metal = Anode.     * Uses: Decoration (Gold/Silver), protection, machinery repair.

• Metal Cladding: Sandwiching base metal between sheets of coating metal and rolling at high temperature.     * Alclad: Duralumin sandwiched between pure Aluminum (Aircraft industry).     * Steel/Aluminum: Rolled at $400^\text{o}C$, annealed at $550^\text{o}C$ to form $FeAl_3$.

• Metal Spraying: Molten metal sprayed via gun using oxy-hydrogen flame and compressed air.     * Used for fabricated parts and non-metallic bases (wood, plastic).

Nernst Theory and Equation

• Definition: Relates cell potential to standard potential, temperature ($298\,K$), and reaction quotient ($Q$).

• Expression for Single Electrode Potential:     * $E_{cell} = E^o - \frac{RT}{nF} \text{ln} Q$     * For reduction reaction ($M^{n+} + ne^- \rightarrow nM$):     * $E_{red} = E^o_{Mn+/M} - \frac{2.303RT}{nF} \text{log} \frac{1}{[M^{n+}]}$

• Constants:     * $R = 8.314\,J/K \cdot mole$     * $F = 96487 \text{ (approx. 96500) } coulomb/mole$

• Simplification at $25^\text{o}C$ ($298\,K$):     * $E = E^o - \frac{0.0592}{n} \text{log}_{10} Q$

Derivation of Nernst Equation

1. Work Done ($W_{red}$): The movement of electrons corresponds to $W_{red} = nFE_{red}$.

2. Gibbs Free Energy ($\Delta G$): Change in free energy equals maximum useful work.     * $\Delta G = -nFE_{red}$ and $\Delta G^o = -nFE^o_{red}$.

3. Vant Hoff Isotherm: For a reversible reaction:     * $\Delta G = \Delta G^o + RT \text{ln} K$     * Where $K = \frac{[Product]}{[Reactant]} = \frac{[M]^n}{[M^{n+}]}$.

4. Substitution:     * $-nFE_{red} = -nFE^o_{red} + RT \text{ln} \frac{[M]}{[M^{n+}]}$     * Dividing by $-nF$ and converting to base-10 log yields the Nernst expression.