Comprehensive Notes on Corrosion Science, Metal Finishing, and E-Waste Management

Introduction to Corrosion Science and Electrochemical Theory

Corrosion is defined as an insidious process involving the destruction of metals and alloys through chemical and electrochemical changes caused by the surrounding environment. The primary initiating factors for this process are atmospheric air and water. Since metals are proficient conductors, they undergo surface-level electrochemical changes that eventually lead to the destruction of the entire metal structure, thereby shortening its functional life. The specific rate at which corrosion proceeds is determined by both the nature of the metal itself and the characteristics of the environment to which it is exposed. Conceptually, corrosion is viewed as the reverse process of metal extraction, represented by the relationship where metal extraction converts metal oxides into metals, and corrosion converts metals back into metal oxides ($M \rightleftharpoons MO$).

According to the electrochemical theory of corrosion, the process occurs when iron is exposed to a wet environment, leading to the formation of many minute galvanic cells consisting of distinct anodic and cathodic areas. Corrosion acts as an oxidation reaction specifically occurring at the anodic site. These galvanic cells arise due to surface heterogeneity from three primary conditions: first, when a metal is exposed to varying concentrations of air or oxygen, where the region in contact with higher oxygen levels acts as the cathode and the region with lower oxygen concentration acts as the anode; second, when two dissimilar metals are in contact within a conducting medium, such as copper-iron contact where copper serves as the cathode and iron as the anode; and third, when a metal is subjected to mechanical stress, causing the strained area to act as the anode. At the anodic area, iron atoms are converted into ferrous ions with the liberation of electrons, expressed by the reaction $Fe(s) \rightarrow Fe^{2+}(aq) + 2e^{-}$. These released electrons migrate toward the cathodic area, forming what is known as the corrosion current, denoted as $i_{\text{Corr}}$.

Cathodic Reactions and the Formation of Rust

Cathodic reactions are significantly more complex than anodic reactions because they involve constituents of the surrounding medium and vary based on environmental conditions. The most common cathodic reactions are the liberation of hydrogen and the absorption of oxygen. Hydrogen liberation occurs in the absence of oxygen. In acidic solutions, hydrogen ions are reduced to gas via $2H^{+}(aq) + 2e^{-} \rightarrow H_{2}(g)$. In neutral or alkaline solutions without oxygen, the reaction follows a two-step process: $2H_{2}O(l) \rightleftharpoons 2H^{+}(aq) + 2OH^{-}(aq)$ followed by the reduction of ions to hydrogen gas, resulting in the net reaction $2H_{2}O(l) + 2e^{-} \rightarrow H_{2}(g) + 2OH^{-}(aq)$.

Oxygen absorption occurs when oxygen is present in the environment. In acidic solutions, oxygen reacts with hydrogen ions to form water: $4H^{+}(aq) + O_{2}(g) + 4e^{-} \rightarrow 2H_{2}O(l)$. In neutral or alkaline environments with oxygen, hydroxide ions are produced as follows: $2H_{2}O(l) + O_{2}(g) + 4e^{-} \rightarrow 4OH^{-}(aq)$. The ferrous ions from the anode and the hydroxide ions from the cathode diffuse toward each other. Because ferrous ions are smaller, they diffuse more rapidly, leading to the formation of insoluble ferrous hydroxide, $2Fe^{2+}(aq) + 4OH^{-}(aq) \rightarrow 2Fe(OH)_{2}(s)$, near the cathodic area. In an oxidizing environment, this is further oxidized to hydrated ferric oxide, commonly known as yellow rust: $4Fe(OH)_{2}(s) + O_{2}(g) + 2H_{2}O(l) \rightarrow 2[Fe_{2}O_{3} \cdot 3H_{2}O](s)$. If oxygen is limited, the product is magnetic black rust ($Fe_{3}O_{4}$), which is a mix of ferrous and ferric oxides: $3Fe(OH)_{2}(s) + \frac{1}{2}O_{2}(g) \rightarrow [Fe_{3}O_{4} \cdot 3H_{2}O](s)$.

The Galvanic Series and Dissimilar Metal Corrosion

The galvanic series is an arrangement of metals and alloys based on their relative corrosion resistance in a specific environment. Metals at the top are more anodic or active, while those at the bottom are noble. The sequence includes Magnesium and its alloys at the most active end, followed by Zinc, Galvanized steel, Aluminium, Cadmium, Steel/Iron, Cast Iron, Chromium steel, active Stainless steel, Lead, Tin, Nickel, Brasses, Copper, Silicon bronze, Silver solder, passive Nickel, Titanium, passive Stainless steel, and finally the noble metals Silver, Graphite, Gold, and Platinum. When alloys are grouped closely together in this series, their potential difference is minimal, and there is little danger of galvanic corrosion when they are in contact.

Differential metal corrosion, also known as galvanic corrosion, occurs when two dissimilar metals are in contact across a conducting medium. A potential difference creates a galvanic current. The more active metal (lower electrode potential) becomes the anode and undergoes corrosion, while the metal with the higher potential becomes the cathode and remains unattacked. For example, if iron is coupled with copper, iron acts as the anode ($Fe(s) \rightarrow Fe^{2+}(aq) + 2e^{-}$) and undergoes corrosion while copper acts as the cathode site for peroxide or hydrogen reactions and remains protected. Conversely, if iron is coupled with zinc, zinc becomes the anode and corrodes because it is more active than iron. Practical examples include steel screws in copper sheets, steel pipes connected to copper plumbing, and lead-antimony solder used on copper wires.

Differential Aeration, Waterline, and Pitting Corrosion

Differential aeration corrosion occurs when a single metal surface is exposed to varying concentrations of air or oxygen. This difference initiates galvanic cells where the part of the metal exposed to lower oxygen concentrations acts as the anode ($M(s) \rightarrow M^{n+}(aq) + ne^{-}$) and undergoes corrosion. The part exposed to higher oxygen levels acts as the cathode ($O_{2}(g) + 2H_{2}O(l) + 4e^{-} \rightarrow 4OH^{-}(aq)$). Common examples include segments of a nail inside a wall, window rods inside a frame, paper pins within paper, metals under dust or scale, and pipelines partially buried in soil or submerged in water.

A specific instance of this is waterline corrosion, often seen in steel water tanks or ships. The metal below the waterline is exposed to less oxygen (dissolved only) and becomes the anode, whereas the metal above the waterline is exposed to the atmosphere and becomes the cathode. This leads to a distinct brown line of rust deposition just below the waterline. The rate is often determined by the amount of meniscus or "creep" of the water. Another destructive form is pitting corrosion, which happens when small particles like sand, dust, or water drops settle on a metal surface. The area beneath the deposit is less aerated and acts as a small anode, while the surrounding exposed metal acts as a large cathode. This configuration accelerates corrosion, leading to the formation of holes, pits, or perforations. These pits are often difficult to detect because they are small and covered by corrosion products.

Corrosion Control through Metallic and Conversion Coatings

Since corrosion is a spontaneous reaction, control is more realistic than complete prevention. Protective metallic coatings can be classified as anodic or cathodic. Anodic coatings involve covering a base metal with a more active metal, such as zinc, magnesium, or aluminium on iron. In these cases, even if the coating is ruptured, the base metal is protected because it remains cathodic relative to the coating. Galvanization is the process of coating steel with zinc through hot dipping. The process involves washing the steel with organic solvents to remove grease, pickling with dilute sulfuric acid to remove rust, treating it with a mixture of aqueous zinc chloride and ammonium chloride, and finally dipping it in molten zinc at $430 - 470\,^{\circ}\text{C}$. Excess zinc is removed by hot rollers. Galvanized items like roofing sheets and buckets are common, though they are unsuitable for food storage because zinc dissolves in dilute acids to form toxic compounds.

Surface conversion coatings involve converting the metal surface into a protective compound through chemical or electrochemical reactions. This coating becomes an integral part of the metal. Anodizing is a common method for non-ferrous metals like Aluminium. The aluminium is made the anode in an electrolytic bath of acids (chromic, sulfuric, phosphoric, or boric). Direct current forms a porous anodic oxide film. To increase corrosion resistance, the film is "sealed" by treatment with boiling water or steam, which converts alumina ($Al_{2}O_{3}$) into its monohydrate ($Al_{2}O_{3} \cdot H_{2}O$). This monohydrate occupies more volume, effectively plugging the pores. Anodized aluminium is used in soap boxes, window frames, and tiffin carriers, and can be produced in various colors.

Cathodic Protection and Impressed Current Methods

Cathodic protection aims to eliminate corrosion by forcing the metal to act as a cathode, reversing the electron flow. In the sacrificial anode method, the metal structure is connected to a more active metal like magnesium or zinc. The active metal corrodes preferentially, "sacrificing" itself. Examples include magnesium blocks on buried oil tanks, zinc wires on buried pipelines, and magnesium bars on ship hulls. This method is cost-effective and requires no external power.

The Impressed Current Cathodic Protection (ICCP) method uses an external DC power source (rectifier) and inert anodes like graphite, high-silicon iron, or MMO-coated titanium. The power source pushes current from the inert anode through the medium (soil or water) to the structure. This makes the entire metal surface a cathode where corrosion cannot occur. ICCP allows for adjustable protection levels and is better suited for very large structures compared to the sacrificial method. The components include the DC source, inert anodes, the structure (cathode), and reference electrodes for monitoring.

Corrosion Penetration Rate (CPR)

Corrosion Penetration Rate (CPR) is defined as the speed at which a metal deteriorates in a specific environment, measured as the thickness loss per year. The rate depends on environmental conditions and the type of metal. It is calculated using the formula $CPR = \frac{KW}{DAT}$, where $W$ is the total weight loss in $mg$, $T$ is the time in hours, $A$ is the exposed surface area, and $D$ is the metal density in $g/cm^{3}$. The constant $K$ varies based on units: for results in mils per year ($mpy$), $K = 534$ and area is in square inches; for millimeters per year ($mmpy$), $K = 87.5$ and area is in square centimeters. CPR is vital for predicting component lifespan, assessing safety risks, selecting materials, optimizing maintenance schedules, and minimizing economic losses from material failures.

Metal Finishing: Electroplating and Electroless Plating

Metal finishing involves modifying a metal's surface properties by depositing a layer of another metal, polymer, or oxide film. It is used to provide a shiny appearance, increase corrosion and wear resistance, improve electrical conductivity, and enhance solderability. Electroplating is the electrochemical process of coating a base metal with a thin layer of another metal using electrical energy. In contrast, electroless plating is a method of depositing a metal or alloy over a substrate via a controlled chemical reaction with a reducing agent, requiring no external electricity. The substrate surface must be catalytic. Catalytic metals like nickel and steel require no preparation, whereas non-catalytic metals like copper require activation by dipping in palladium chloride. Non-conductors like glass or plastic are activated using stannous chloride followed by palladium chloride.

Key differences between the two methods involve the driving force (current vs. autocatalytic redox reaction) and application. Electroplating requires a separate anode ($M \rightarrow M^{n+} + ne^{-}$) and is limited to conductors, typically producing thickness between $1 - 100\,\mu\text{m}$. Electroless plating uses the catalytic surface of the substrate as the anode while chemical reagents provide electrons for reduction. It is applicable to conductors, semiconductors, and non-conductors, offering superior "throwing power" (uniformity of coating on complex shapes) compared to electroplating.

Chromium and Copper Plating Applications

Chromium plating is used for decorative and hard coatings. Both use a bath of $250\,g/dm^{3}$ chromic acid and $2.5\,g/dm^{3}$ sulfuric acid at $45 - 55\,^{\circ}\text{C}$. Decorative chromium is thin ($0.25 - 0.75\,\mu\text{m}$) with a current density of $145 - 430\,A/ft^{2}$, while hard chromium is thicker ($2.5 - 300\,\mu\text{m}$) with higher current density ($290 - 580\,A/ft^{2}$). In the bath, $Cr^{6+}$ is reduced to $Cr^{3+}$, which is then reduced to metal ($Cr^{3+} + 3e^{-} \rightarrow Cr$). Lead or lead-alloy anodes are used because chromium anodes passivate and cause burnt deposits. Applications include automotive finishes, surgical tools (decorative), and cutting tools, piston rings, and marine engine crankshafts (hard).

Electroless copper plating is vital for Printed Circuit Boards (PCBs). The bath contains copper sulfate ($12\,g/dm^{3}$) as the metal source, formaldehyde ($8\,g/dm^{3}$) as the reducing agent, and sodium hydroxide and Rochelle salt as buffers to maintain pH $11$. Disodium EDTA ($20\,g/dm^{3}$) is used as a complexing agent, and the process runs at $25\,^{\circ}\text{C}$. The net reaction is $Cu^{2+} + 2HCHO + 4OH^{-} \rightarrow Cu + 2HCOO^{-} + 2H_{2}O + H_{2}$. In PCB manufacturing, holes are drilled through plastic bases like glass fiber reinforced plastic (GRP). These holes are activated and then electrolessly plated to create electrical connections between the two sides of the board. Unwanted copper is etched away to leave the circuit pattern.

E-Waste Sources, Composition, and Environmental Impact

Electronic Waste (E-Waste) refers to discarded electrical or electronic devices and components that are not intended for reuse or recycling. Sources include large household appliances (fridges, washing machines), small appliances (toasters, irons), IT equipment (PCs, laptops, phones), lighting, tools, and toys. E-waste is caused by rapid technological advancement, changes in fashion, and reached end-of-life. It contains valuable metals like gold, platinum, and silver; useful metals like copper and aluminum; and hazardous substances like mercury, radioactive isotopes, PCBs, and dioxins. A single mobile phone can contain over $40$ elements. Plastics like HIPS, ABS, and Polycarbonate, along with CRT glass ($SiO_{2}$, $CaO$, $Na$), are also part of its composition.

E-waste poses severe threats to the environment. Informal recycling through open-air burning releases toxic dioxins, furans, and hydrocarbons, contributing to smog and acid rain. Heavy metals like lead and mercury vaporize or turn into fine dust (PM10 and PM2.5), causing respiratory diseases. E-waste incineration also releases greenhouse gases like $CO_{2}$ and CFCs. In soil, heavy metals leach from landfills, altering pH, killing essential microbes, and reducing fertility. These toxins enter the food chain via bioaccumulation. Water systems are contaminated by leachate and industrial discharge from informal acid baths, leading to aquatic toxicity, mass mortality of river ecosystems, and hormonal disruptions in wildlife. Persistent organic pollutants (POPs) can travel through ocean currents to affect marine life thousands of miles away.

Human Health Risks and Bioleaching Gold Extraction

Exposure to e-waste toxins occurs via inhalation, skin contact, or contaminated food and water. Vulnerable groups include children, who face irreversible neurological damage and lower IQ, and pregnant women, as toxins like lead cross the placenta causing stillbirths or low birth weight. Health issues include neurotoxicity, respiratory damage (asthma, bronchitis), and organ failure (kidney and liver damage). Chronic cadmium exposure causes bone demineralization (Itai-itai disease). Genetic impacts include DNA strand breakage and endocrine disruption. Carcinogenic risks include lung, skin, and bladder cancers linked to beryllium, arsenic, and chromium.

Bioleaching is a "green" method for extracting gold from e-waste using microorganisms. The process begins with shredded e-waste. First, copper is removed using acidophilic bacteria like Acidithiobacillus ferrooxidans and thiooxidans: $4Fe^{2+} + O_{2} + 4H^{+} \rightarrow 4Fe^{3+} + 2H_{2}O$ and $CuFeS_{2} + 4Fe^{3+} \rightarrow Cu^{2+} + 5Fe^{2+} + 2S^{0}$. Next, gold leaching (cyanogenesis) occurs using Pseudomonas or Chromobacterium. These bacteria metabolize glycine ($C_{2}H_{5}NO_{2}$) to produce hydrogen cyanide ($HCN$). Under alkaline conditions (pH $7.3 - 8.6$), gold dissolves to form a soluble complex: $2Au + 4CN^{-} + \frac{1}{2}O_{2} + H_{2}O \rightarrow 2[Au(CN)_{2}]^{-} + 2OH^{-}$. Gold is then recovered via precipitation using agents like sodium metabisulfite ($2Au^{+} + 3SO_{2} + 6H_{2}O \rightarrow 2Au(s) + 3H_{2}SO_{4}$) or adsorption on activated carbon. While eco-friendly and cost-effective, bioleaching is slower and currently yields lower purity than traditional chemical methods.