Metallic Corrosion

• Corrosion is an undesirable process that limits progress and incurs unlimited replacement costs for materials and equipment.
• Corrosion is the spontaneous destruction of metals through chemical, electrochemical, or biochemical interactions with their environment.
• The rusting of iron involves the formation of a reddish scale of oxide (Fe3O4).
• Copper corrosion results in a green film of basic carbonate.
• Consequences of corrosion include machinery failure, product contamination, equipment replacement, toxic releases, and health risks.

Causes of Corrosion

• Metals exist in nature in either a native or combined state.
• Native state metals are uncombined, non-reactive, and have good corrosion resistance (e.g., Au, Pt, Ag).
• Combined state metals are reactive and form stable compounds like ores and minerals (e.g., Fe2O3, ZnO, PbS, CaCO3).
• Metallic corrosion occurs because metals extracted from ores are in a higher energy state and decay to more stable compounds when exposed to the environment.
• Corrosion leads to a loss of useful properties like malleability, ductility, hardness, luster, and electrical conductivity.

Classification of Corrosion

• Based on the environment, corrosion is classified as either dry (chemical) or wet (electrochemical).
• Dry corrosion involves a direct chemical attack by atmospheric gases (e.g., oxygen, halogens).
• Silver corrosion is an example of dry corrosion where silver materials undergo direct chemical attack by atmospheric H2S gas.

Types of Dry or Chemical Corrosion

• Corrosion can occur due to oxygen (oxidation), hydrogen, or liquid metals.
• Oxidation corrosion results from the direct attack of oxygen on metal surfaces.
• Alkali and alkaline earth metals oxidize rapidly at low temperatures.
• Most metals, except Ag, Au, and Pt, oxidize at high temperatures.
• During oxidation, metal ions form at the surface, oxygen converts to oxide ions, and a metal oxide film forms: 2M + n/2 O2 → 2 Mn+ + nO2-.
• The nature of the oxide film determines the extent of further corrosion.

Types of Oxide Layers

• Stable layers are fine-grained, tightly adhere to the metal surface, and act as protective coatings (e.g., oxides of Al, Sn, Pb, Cu, Pt).
• Unstable oxide layers form on noble metals (Ag, Au, Pt), decomposing back into metal and oxygen, preventing oxidation corrosion.
• Volatile oxide layers volatilize as they form, continuously exposing fresh metal surfaces to corrosion (e.g., MoO3).
• Porous layers have pores or cracks that allow atmospheric oxygen to access the underlying metal, leading to continuous corrosion.
• The Pilling-Bedworth rule states that an oxide is protective if its volume is at least as great as the volume of metal from which it formed.
• Alkali and alkaline earth metals form porous oxides, while aluminum forms a tightly-adhering, non-porous layer.

Corrosion by Other Gases

• Hydrogen embrittlement is the loss of ductility due to the presence of hydrogen.
• Embrittlement occurs when iron liberates atomic hydrogen with hydrogen sulphide Fe + H₂S → FeS + 2H
• Upon collection in voids, the hydrogen forms high pressure gas, causing cracking and blistering.
• Decarburization occurs when steel is exposed to high-temperature hydrogen, reducing the carbon content and weakening the steel.
• Liquid metal corrosion involves the chemical action of flowing liquid metal at high temperatures: leading to dissolution or internal penetration.

Wet or Electrochemical Corrosion

• Electrochemical corrosion involves the formation of anodic and cathodic areas, a conducting medium, and electron flow.
• At anodic areas, oxidation occurs, destroying the metal.
• At cathodic areas, reduction occurs, and dissolved constituents accept electrons.

Hydrogen Evolution Type

• Metals above hydrogen in the electrochemical series dissolve in acidic solutions with the simultaneous evolution of hydrogen (Fe → Fe2+ + 2e-).
• Hydrogen ions are eliminated as hydrogen gas, leading to the overall reaction Fe + 2H+ → Fe2+ + H2.

Oxygen Absorption Type

• Rusting of iron in a neutral aqueous solution with electrolytes is a common example.
• Metal dissolves as ferrous ions at the anode and the liberated electrons are intercepted by dissolved oxygen at the cathode.
• Ferrous hydroxide is precipitated and can be further oxidized to ferric hydroxide (yellow rust) or magnetite (black anhydrous).

Differences Between Dry and Wet Corrosion

• Dry corrosion occurs in the absence of moisture, while wet corrosion requires moisture or an electrolyte.
• Dry corrosion involves direct chemical attack, while wet corrosion involves the formation of anodic and cathodic areas.
• Dry corrosion can affect homogeneous metal surfaces, while wet corrosion typically affects heterogeneous surfaces.
• In dry corrosion, products accumulate at the corrosion site, while in wet corrosion, products form elsewhere.
• Dry corrosion is self-controlled, while wet corrosion is continuous.

Types of Electrochemical Corrosion

• Electrochemical corrosion is classified into galvanic and differential aeration.
• Galvanic corrosion occurs when two dissimilar metals are electrically connected and exposed to an electrolyte, with the more active metal corroding.
• Steel screws in brass marine hardware are examples.
• In acidic environments, hydrogen evolution drives corrosion; in neutral or alkaline, oxygen absorption occurs.
• The corrosion occurs at the anode metal while the cathode is protected.

Concentration Cell Corrosion

• Concentration cell corrosion stems from electrochemical attacks on metal surfaces exposed to varying electrolyte concentrations or aeration levels.
• It occurs when one part of the metal experiences a different air concentration than another.
• Poorly oxygenated sections become anodic.
• Examples include waterline corrosion and corrosion of partially buried metal in soil.
• The metal part above solution is more aerated (cathodic), the part in the solution is less aerated (anodic).

Anode/Cathode Reactions

• M → M2+ + 2e- happens at the anode (less aerated).
• 1/2 O2 + H2O + 2e¯ → 2OH- happens at the cathode (more aerated).

Examples of Concentration Cell Corrosion

• Pitting corrosion happens
• Crevice corrosion happens
• Pipeline corrosion happens
• Wire fence corrosion happens

Pitting

• Pitting is a localized attack, resulting in holes, with metal around being relatively untouched.
• This corrosion involves a differential aeration or concentration cell setup.
• Metal covered by water, dust, sand, etc. acts as an aeration or concentration cell.
• The area with less oxygen (covered by salt solution) acts as the anode.

Crevice Corrosion

• Crevice corrosion happens when a narrow opening is in contact with liquid.
• A crack between metallic and non-metallic material is an example, here the crevice becomes anodic and undergoes corrosion.
• Oxygen supply to the crevice is less, the exposed area acts as the cathode.
• Bolts, nuts, rivets, and joints are examples.

Pipeline Corrosion

• Pipeline corrosion is when buried pipelines or cables from one type of soil (clay less aerated) to another soil (sand more aerated) corrode.
• This happens due to differential aeration.
• Corrosion in the wire fence is when the areas where the wires cross (anodic) are less aerated than the rest of the fence (cathodic).

Factors Influencing Corrosion

• There are two factors that influence the rate of corrosion: nature of the metal, and nature of the corroding environment.
• Understanding these factors and how they affect corrosion rate is essential.

Purity of Metal & Corrosion Rate

• Impurities in metal cause electrochemical cells to form, leading to corrosion of anodic parts
• Higher purity reduces corrosion, i.e. aluminum's corrosion rate in hydrochloric acid increases with more impurities.

Over Voltage & Corrosion

• Over voltage in a corrosive environment is inversely proportional to corrosion rate e.g the corrosion rate of zinc increases when the hydrogen over voltage is reduced by adding copper sulphate.

Relative Areas & Corrosion Rate

• With two dissimilar metals, the corrosion of the anodic part happens at a rate directly proportional to the ratio of the anodic and cathodic areas.
• If the anodic area is small, corrosion can be faster due to higher current density.

Galvanic Series Position

• Positioning in the galvanic series and passive character impacts corrosion.
• Also the solubility and volatility of corrosion products impacts corrosion.

Temperature's Role

• The rate of corrosion is directly proportional to temperature, as higher temperatures increase ion diffusion rates.

Humidity Levels

• Rate will be higher when humidity is highest
• Moisture acts as a solvent for setting up the corrosion cell

Impurities In Air

• Atmosphere in industrial areas increase acidity and electrical conductivity, increasing rate of atmospheric corrosion.

Airborne Particles

• If they are chemically active, they result in enchanced corrosion
• If they are inactive, they absorb sulphur and moisture, so they slowly enhance corrosion rate.

Impact of pH

• Acidic solution are corrosive relative to alkaline ones.
• Zn suffers minimum corrosion at pH 11.

Corrosion Control Methods

• To minimize corrosion control: modify the metal, or modify the environment.
• The first choice of metal is noble metals.

Choice of Metals & Alloys

• Noble metals such as gold and platinum are corrosion-resistant but too expensive for general use.
• The use of purest possible metal is next choice, but high chemical purity is difficult to achieve.
• The next best choice is corrosion-resistant alloys.
• Stainless steel with chromium forms a protective oxide film to prevent steel from further corrosion.
• Nimonic alloys are resistant to not gases.

Proper Designing Techniques

• Use simple designs.
• Do not use complicated shapes.
• Insulate dissimilar metals to avoid galvanic corrosion.

Surface Area Requirements

• Increase anodic area, make cathodic area as small as possible to avoid corrosion when dissimlar metals are used in contact.

Crevices

• Avoid these
• Bolt and rivet joints, use welding instead

Washers

• Replace metal washers with plastic ones

Pipelines

• Use smooth bends to control corrosion in these.
• Use annealing minimize stress corrosion.

Water Storage Design

• Good water storage design must allow to drain/clean easily.

Cathodic Protection

• It involves making a metallic structure cathodic to prevent corrosion.
• It works by preventing anodic areas.
• Two types: Sacrificial anodic protection ; Impressed current cathodic protection.

Sacrificial Anodic Protection

• This is where, the protected structure is connected it with a more active metal.
• All corrosion concentrates on the active metal (sacrificial anode), protecting the parent stricture.
• Examples: magnesium, zinc, aluminium
• Applications include underground cables and ship marine protection.

Impressed

• Opposes the corrosion current to turn the corroding metal from anode to cathode.
• High silica Iron material is used for anode.
• Its used to protect large structures for long periods, like water coolers, water tanks and marine piers.

Sacrificial Anode vs Impressed Current

• While sacrificial doesnt need external power, impressed does
• Sacrificial investment costs are low, impressed, high
• Sacrificial requires periodic replacement, where impressed stabilizes from the start
• Sacrifice does not consider soil defects, whereas impresse accounts for it
• In the end, sacrifice is economical, impressed is well suited for large structures.

Environmental Corrosion Control

• You can prevent corrosion by modifying the environment.
• Deaeration: Reduce corrosion by removing dissolved oxygen from water via heating.
• Remove dissolved carbon dioxide from water.

Using Inhibitors

• Use an organic or inorganic substance that reduces corrosion rate
• Anodic ihibitors are chemical perservatives
• Cathodic inhibitors are adsorption perservatives
• Vapor phase inhibotrs are volatile corrosion inhibitors

Anodic Inhibitors

• Anodic inhibitors form a sparingly soluble compound, retarding corrosion by forming a passive film or barrier.
• Used to repair the oxide film over metal, or pitting corrosion.

Cathodic Inhibitors

• For acidic environments, use amines, mercaptans, and thiourea to slow Hydrogen diffusion through the cathode

Neutral Solution Usage

• For neutral sloutions, prevent oxygen from corroding metal.
• Eliminate dissolved oxygen by adding reducing agents like Na2SO3.

Vapor Phase Inhibitors

• Organic inhibitors are readily vaporized
• Used to prevent corrosion

Anodic Protection

• Controls corrosion via potentistat electrochemical corrosion
• Works when passive film is produced onto the metal

Protective Coatings

• They act as Barriers
• They afford special properties
• They are classified as organic/inorganic types.
• Inorganic are either Metallic or Non-metallic
• Metallic: Hot Dipping, Metal Cladding, Cementation and electroplating

Paints

• Made out of viscous opaque mechanical dispersion mixture of dye in a vehicle
• A good paint must have: high hiding power, uniform fillm, film wont crack when drying.
• It should be washable and atmospheric condition resistant.

Other Paint Requirements

• Possess high material capacity
• Not fade
• Dry quickly

Constituents

• Made out of pigment, vehicle, thinner, driers, fillers, plasticizers, etc.
• pigment imparts color
• Vehicles are liquid susbstances made out of film forming material.

Thinners for Paints

• You must use violatile substanes to thin the quality of paint.
• Examples: Dipentine, turpentine, toluol, xylol

Driers for Paints

• Accelerate process of drying, are catalysts and examples include zinc, lead, cobalt.

Fillers/Extenders

• Inert materials that improve paint
• Fill voids and reduce painr material costs
• Add to durability
• Increase arrangeent of pigment paritcles

Placticizers

• Paint ingredients enhance elasticity and minimze cracking, eg. triphehyl phosphate
• There are anti-skinning agents which prevents thickening and hardening.

Metallic Coatings

• Corrosion can be controlled when you prevent with galvanisation, electroplating, etc
• Hot Dipping requires putting a coating of Zinc - 419, Tin, aluminum-on iron.

Galvanisation

• Zinc coats iron/steel to prevent rusting.

Metal Cladding

• Dense Material bonded firmly to bse material
• Silver and nickel materials fall under this
• Tinning: A coating of thing over ferrous material passes steel sheets in sulphuric acid then rollers covered in oil.

Electroplating

• Deposits coasting material via electrolysis
• Objectives: decoration better appearance, chemical attack, electro refining.
• Additives such as glues gels and acids are added to electrolytic baths along with brightness
• Use an opptimum temperature, current density and lower metal ion consiturations
• To make it work
• The surface to be plated is first degreased by using organic solvents or alkali, followed by acid treatment.
• Examples include Gold(electrical) and ornomental jewelry (thin coat).

Electroless Plating

• Electroless plating deposits without using electrical energy
• Nickel is pretreated with 50% sulpuric acid and plated at various heats.

Explore the causes, types, and prevention of metallic corrosion, including electrochemical and dry corrosion. Learn about factors influencing corrosion, such as exposure to oxygen, hydrogen sulfide, and the properties of noble metals.