Corrosion Notes FE Sem 1 Applied Chemistry | Mumbai University

Corrosion is the slow destruction of metals by chemical or electrochemical reaction with their surroundings, and it costs industry enormous sums every year. This module explains the two basic mechanisms, dry (chemical) and wet (electrochemical) corrosion in acidic and neutral media. It then examines the factors that decide how fast a metal corrodes and the practical methods used to control corrosion, from design and material choice to cathodic protection, inhibitors and metallic coatings. It ends with corrosion problems in electronic devices, which matter for every engineering branch.

Summary

Corrosion is the gradual destruction or deterioration of a metal by unwanted chemical or electrochemical attack of its environment, starting at the surface. Most metals occur in nature as oxides, sulphides or carbonates; during extraction they are raised to a higher energy state, and corrosion is the natural tendency to return to the stable combined state, which is why it is called reverse metallurgy. Rusting of iron (hydrated ferric oxide, Fe2O3.xH2O) and green film on copper are common examples.

Dry or atmospheric (chemical) corrosion is the direct attack of dry gases such as O2, Cl2, H2S and SO2 on a metal surface in the absence of a liquid electrolyte. In oxidation corrosion the metal loses electrons (2M gives 2Mn+ + 2ne-) and oxygen gains them to form an oxide film. The nature of the film decides further attack: a stable, compact, non-porous film (Al, Cr, Sn) protects the metal; an unstable film (Ag, Au, Pt) decomposes back to metal; a volatile film (MoO3) evaporates and exposes fresh metal; and a porous film allows continued attack. The Pilling-Bedworth rule states that if the volume of the oxide is less than the volume of metal consumed (ratio below 1, e.g. Li, Na, K, Mg) the film is porous and non-protective, while a ratio of 1 or more (e.g. Al, Cr) gives a protective film.

Wet or electrochemical corrosion occurs when a conducting liquid (moisture or electrolyte) is in contact with the metal, or when two dissimilar metals or areas are in contact through it. Separate anodic and cathodic areas form, and current flows between them. Oxidation always occurs at the anode (Fe gives Fe2+ + 2e-), so the anode corrodes. In an acidic medium free of oxygen, the cathodic reaction is hydrogen evolution: 2H+ + 2e- gives H2. In a neutral or slightly alkaline aerated medium, the cathodic reaction is oxygen absorption: 1/2 O2 + H2O + 2e- gives 2OH-. The Fe2+ and OH- ions meet to form Fe(OH)2, which is oxidised further by dissolved oxygen to Fe(OH)3 and rust, Fe2O3.xH2O, which deposits near the cathode.

The rate of corrosion depends on the nature of the metal and of the environment. (i) Position in the galvanic series: the farther apart two metals are, the faster the more active (anodic) one corrodes. (ii) Purity: impurities form tiny galvanic cells, so pure metals corrode less. (iii) Nature of the corrosion product: a stable, non-porous, insoluble film slows attack, while a porous, soluble or volatile one speeds it. (iv) Temperature: rate generally rises with temperature. (v) pH: acidic media (low pH) corrode most metals faster; iron corrodes slowly above about pH 10. (vi) Concentration of the medium: higher concentration of corrosive ions such as Cl- generally increases the rate, though some metals are passivated by concentrated acids (Al and Fe in conc. HNO3). (vii) Moisture: corrosion increases sharply above the critical humidity. (viii) Relative areas: a small anode with a large cathode corrodes intensely. (ix) Overvoltage: a high hydrogen overvoltage on the metal slows hydrogen-evolution corrosion.

Corrosion is controlled first by choosing the right metal: noble or pure metals, or corrosion-resistant alloys such as stainless steel, for the given environment. Proper design avoids contact of dissimilar metals (or uses insulating washers), keeps the anodic metal area large when contact is unavoidable, prefers welding to riveting, avoids crevices, sharp bends and stagnant pockets, and allows free drainage and easy cleaning and painting.

Cathodic protection makes the whole structure the cathode. In sacrificial anode protection, a more active metal (Mg, Zn or Al) is connected to the structure (buried pipelines, ship hulls, water tanks) and corrodes instead of it, being replaced periodically. In impressed current protection, a DC source drives current from an insoluble anode (graphite, high-silicon iron, scrap iron) buried in a backfill (coke, gypsum, bentonite) to the structure, so the structure becomes the cathode. Corrosion inhibitors are substances added in small amounts to the medium: anodic inhibitors (chromates, phosphates, molybdates, tungstates) form an insoluble film on anodic areas, and cathodic inhibitors either raise hydrogen overvoltage or form films in acid (amines, thiourea, heavy-metal salts of As and Sb) or remove dissolved oxygen in neutral media (Na2SO3, hydrazine) or precipitate hydroxide films on cathodes (Mg, Zn and Ni salts).

Metallic coatings are anodic or cathodic. An anodic coating such as zinc on iron (galvanizing) protects even when scratched, because zinc corrodes sacrificially. A cathodic coating such as tin on iron (tinning) protects only while intact; at a scratch, iron becomes a small anode facing a large tin cathode and corrodes rapidly. In galvanizing, steel is pickled in dilute H2SO4 at 60 to 90 C, washed, dried, fluxed with NH4Cl and dipped in molten zinc at 425 to 450 C, and excess zinc is removed with rollers. Galvanized iron is not used for food containers because zinc forms toxic compounds with food acids; tinned iron is used instead. Other methods are tinning by hot dipping, electroplating, metal cladding, metal spraying and cementation (sherardizing).

Corrosion in electronic devices is caused by moisture films, ionic contaminants (chlorides, flux residues, fingerprints), polluting gases (SO2, H2S, NOx) and the contact of dissimilar metals in connectors, solder joints and PCBs. Under applied voltage, metal ions dissolve at the anode and redeposit as dendrites (electrochemical migration, typical of Ag, Cu and Sn), causing leakage currents and short circuits. Other forms include galvanic attack between gold plating and the underlying Ni or Cu at pores, fretting corrosion of connector contacts, and creep corrosion of copper sulphide films. Control measures are conformal coatings and encapsulation, hermetic sealing, humidity control, thorough cleaning of flux residues, noble-metal (gold) plating of contacts and use of vapour-phase corrosion inhibitors in packaging.

Mind map

Mechanism of wet corrosion of iron in a neutral aerated medium (oxygen absorption)

Important definitions

Formulas

Key points to remember

• Anode always corrodes; cathode is protected.
• Corrosion is called reverse metallurgy because the metal returns to its stable combined state.
• Acidic, oxygen-free medium: hydrogen evolution; neutral aerated medium: oxygen absorption.
• Rust is deposited near the cathode, while pitting occurs at the anode.
• Small anode with large cathode gives very fast corrosion (e.g. steel rivets in copper plates).
• Alkali and alkaline earth metals form porous oxides (PB ratio below 1); Al and Cr form protective oxides.
• Pure zinc corrodes slowly in dilute H2SO4 because of high hydrogen overvoltage; impure zinc corrodes fast.
• Galvanizing is anodic protection of the coating type; tinning is cathodic.
• Galvanized containers are not used for food; tinned containers are used.
• Sacrificial anodes: Mg, Zn, Al. Impressed current anodes: graphite, scrap iron, high-silicon iron in backfill.
• Anodic inhibitors: chromates, phosphates, molybdates, tungstates. Cathodic: amines, thiourea, Na2SO3, hydrazine, Mg/Zn/Ni salts.
• Electronics corrode mainly through moisture plus ionic contamination plus voltage, causing dendrites and short circuits.

Solved examples

Example 1. Calculate the Pilling-Bedworth ratio for aluminium and predict whether its oxide film is protective. (Al2O3: M = 101.96 g/mol, density 3.95 g/cm3; Al: M = 26.98 g/mol, density 2.70 g/cm3)

Step 1: PB = (M_oxide x rho_metal)/(n x M_metal x rho_oxide), n = 2 Al per Al2O3. Step 2: Numerator = 101.96 x 2.70 = 275.29. Denominator = 2 x 26.98 x 3.95 = 213.14. Step 3: PB = 275.29/213.14 = 1.29. Since PB > 1, the oxide occupies more volume than the metal consumed, forms a continuous non-porous film, and is protective. This is why aluminium resists atmospheric corrosion.

Example 2. Calculate the Pilling-Bedworth ratio for magnesium (MgO: M = 40.30 g/mol, density 3.58 g/cm3; Mg: M = 24.31 g/mol, density 1.74 g/cm3) and comment.

Step 1: n = 1. PB = (40.30 x 1.74)/(1 x 24.31 x 3.58). Step 2: Numerator = 70.12, denominator = 87.03. Step 3: PB = 0.81. Since PB < 1, the oxide cannot cover the metal fully; the film is porous and non-protective, so oxidation of Mg continues.

Example 3. Calculate the Pilling-Bedworth ratio for copper forming Cu2O (Cu2O: M = 143.09 g/mol, density 6.0 g/cm3; Cu: M = 63.55 g/mol, density 8.96 g/cm3).

Step 1: n = 2. PB = (143.09 x 8.96)/(2 x 63.55 x 6.0). Step 2: Numerator = 1282.1; denominator = 762.6. Step 3: PB = 1.68, greater than 1, so the Cu2O film is protective at ordinary temperatures.

Example 4. A copper plate is fastened with steel bolts in sea water. Predict which metal corrodes and how fast. What change in design would you suggest?

Step 1: In the galvanic series steel (iron) is more active than copper, so steel becomes the anode and copper the cathode. Step 2: The bolts present a small anodic area while the plate is a large cathodic area; the cathodic reaction (O2 reduction) consumes electrons over a large area, so the current density at the small anode is very high and the bolts corrode rapidly. Step 3: Remedy: use copper or bronze bolts (or bolts of a metal close to copper in the series), or insulate the bolts with non-conducting washers and sleeves; if steel must be used, the anodic part should have the larger area.

Example 5. A galvanized iron sheet and a tinned iron sheet are both scratched, exposing iron. Explain what happens in each case.

Galvanized sheet: zinc is above iron in the galvanic series, so at the scratch zinc acts as anode and iron as cathode. Zinc corrodes sacrificially and iron is protected until the zinc near the scratch is used up. Tinned sheet: tin is more noble than iron, so the exposed iron becomes the anode and the large tin surface the cathode. The small anode with large cathode causes rapid local corrosion (pitting) of iron at the scratch. Hence anodic coatings protect even when broken, while cathodic coatings protect only when continuous.

Example 6. Pure zinc reacts very slowly with dilute H2SO4, but the reaction becomes fast when a few drops of CuSO4 solution are added. Explain.

Step 1: Pure zinc has a high hydrogen overvoltage, so the cathodic reaction 2H+ + 2e- gives H2 is slow on its surface, which slows corrosion. Step 2: Zn displaces Cu from CuSO4 and copper deposits on the zinc. Step 3: Copper has a lower hydrogen overvoltage and acts as an efficient cathode, forming many Zn-Cu micro galvanic cells. Step 4: Hydrogen is now evolved easily on copper, and zinc (anode) dissolves rapidly.

Important questions and answers

Q1. Define corrosion. Why is it called reverse metallurgy? (3 marks)

Corrosion is the gradual destruction of a metal by chemical or electrochemical reaction with its environment. In metallurgy, stable ores (oxides, sulphides) are reduced to metals by supplying energy; the metals are thus in a higher energy state and tend to return to the stable combined state by reacting with the environment. Since corrosion reverses extraction, it is called reverse metallurgy.

Q2. Explain dry corrosion. Discuss the nature of oxide films and the Pilling-Bedworth rule. (6 marks)

Dry corrosion is the direct attack of dry gases on a metal. In oxidation corrosion, metal atoms lose electrons at the surface (2M gives 2Mn+ + 2ne-) and oxygen gains them (n/2 O2 + 2ne- gives nO2-), forming metal oxide. Further attack depends on the film: stable (Al, Cr, Cu): compact, adherent, protective; unstable (Ag, Au, Pt): decomposes back to metal and oxygen, so no net corrosion; volatile (MoO3): evaporates, exposing fresh metal, so rapid corrosion; porous: cracks and pores allow O2 to reach the metal. Pilling-Bedworth rule: if volume of oxide is less than volume of metal (ratio below 1, alkali and alkaline earth metals) the film is porous and non-protective; if the ratio is 1 or more (Al, Cr) the film is non-porous and protective.

Q3. Explain the mechanism of wet corrosion by (i) hydrogen evolution and (ii) oxygen absorption. (6 marks)

(i) Hydrogen evolution occurs in acidic media without oxygen, e.g. iron in contact with acid industrial waste. Anode: Fe gives Fe2+ + 2e-. Cathode: 2H+ + 2e- gives H2. Overall Fe + 2H+ gives Fe2+ + H2. Here anodes are large and cathodes small, and all metals above hydrogen in the electrochemical series can corrode this way. (ii) Oxygen absorption occurs in neutral or slightly alkaline aerated media, e.g. iron in water or moist air. Anode: Fe gives Fe2+ + 2e-. Cathode: 1/2 O2 + H2O + 2e- gives 2OH-. Fe2+ + 2OH- gives Fe(OH)2; with excess O2, 4Fe(OH)2 + O2 + 2H2O gives 4Fe(OH)3, i.e. rust Fe2O3.xH2O. Anodic areas are small and pitting results; rust deposits near the cathode. Diagrams of the anodic and cathodic areas should be drawn.

Q4. Discuss any four factors affecting the rate of corrosion. (4 marks)

Galvanic series position: greater separation of two metals in contact gives faster corrosion of the anodic one. Purity: impurities set up local galvanic cells, so impure metals corrode faster. Relative areas: a small anode and large cathode corrode the anode intensely. pH: lower pH increases corrosion; iron corrodes little above pH 10. (Other acceptable factors: nature of corrosion product, temperature, concentration, moisture, overvoltage.)

Q5. Explain the effect of overvoltage and relative anodic-cathodic areas on corrosion. (4 marks)

Overvoltage: in hydrogen-evolution corrosion, H2 must be liberated at the cathode. A high hydrogen overvoltage at the cathode surface slows this step and hence slows corrosion; pure Zn in dil. H2SO4 corrodes slowly for this reason, while adding Cu (low overvoltage) speeds it up. Areas: the anodic and cathodic currents are equal, so when the anode is small and the cathode is large, the current density at the anode is very high and corrosion is rapid and localised; a large anode with a small cathode corrodes slowly and uniformly.

Q6. What is cathodic protection? Explain sacrificial anode and impressed current methods. (6 marks)

Cathodic protection is the method of protecting a metal by forcing it to act as the cathode, so that it does not corrode. Sacrificial anode: the structure is connected by a wire to a more active metal such as Mg, Zn or Al, which acts as anode and corrodes, while the structure becomes the cathode. The anodes are replaced when consumed. Used for buried pipelines, ship hulls, water tanks and marine structures. Impressed current: a direct current from an external source (battery or rectifier) is applied so that the structure is the cathode; an insoluble anode such as graphite, high-silicon iron or scrap iron is buried in a backfill of coke, gypsum or bentonite to improve conductivity. Used for long pipelines, underground tanks, transmission line towers. Diagrams of both arrangements are expected.

Q7. What are corrosion inhibitors? Explain anodic and cathodic inhibitors with examples. (5 marks)

Inhibitors are substances added in small amount to a corrosive medium to decrease corrosion. Anodic inhibitors such as chromates, phosphates, tungstates and molybdates react with metal ions formed at the anode to give a sparingly soluble, adherent film that blocks the anodic reaction; if added in insufficient amount they can cause pitting. Cathodic inhibitors act on the cathodic reaction: in acidic solution, organic inhibitors such as amines, thiourea and mercaptans adsorb on the metal, and salts of heavy metals like As and Sb deposit and raise hydrogen overvoltage; in neutral solution, oxygen scavengers such as Na2SO3 and hydrazine remove dissolved O2, and salts of Mg, Zn or Ni precipitate insoluble hydroxide films on cathodic areas.

Q8. Distinguish between anodic and cathodic coatings with examples. (4 marks)

Anodic coating: coating metal is more active (anodic) than base metal, e.g. Zn on Fe (galvanizing). It protects even if broken because the coating corrodes sacrificially. Cathodic coating: coating metal is more noble, e.g. Sn on Fe (tinning). It protects only if continuous; at a break the base metal corrodes rapidly due to small anode and large cathode. Galvanized iron is not used for food, tinned iron is.

Q9. Describe the process of galvanizing with a diagram. (5 marks)

Galvanizing is coating iron or steel with zinc by hot dipping. Steps: the article is cleaned with dilute H2SO4 at 60-90 C for 15-20 min (pickling) to remove scale and rust; washed with water and dried; dipped in molten zinc at 425-450 C; the surface is kept covered with NH4Cl flux to prevent oxidation of zinc; the article is taken out and passed through rollers to remove excess zinc and give a uniform coat; it is then cooled. Uses: roofing sheets, wires, pipes, buckets, nails. The diagram shows the pickling bath, washing, drier, molten zinc bath with flux and rollers.

Q10. Write a note on corrosion in electronic devices. (4 marks)

Electronic components contain closely spaced conductors of different metals (Cu, Sn, Ag, Au, Ni) at different potentials. Moisture films with ionic contaminants (chloride, flux residues) and pollutant gases (SO2, H2S) act as electrolytes. Effects: galvanic corrosion at connectors and at pores in gold plating; electrochemical migration in which Ag, Cu or Sn dissolves at the anode and grows dendrites towards the cathode, causing leakage and short circuits; fretting corrosion of contacts; corrosion of PCB tracks and solder joints. Prevention: conformal coating and encapsulation, hermetic sealing, humidity control, proper cleaning of flux residues, gold plating of contacts and vapour-phase inhibitors in packaging.

Q11. How does proper designing help control corrosion? (4 marks)

Avoid contact of dissimilar metals; if unavoidable, keep them close in the galvanic series and insulate the joint. Keep the anodic area large and the cathodic area small. Prefer welding to riveting and butt joints to lap joints to avoid crevices. Avoid sharp corners, bends and stagnant areas where moisture and dirt collect; provide drainage. Design for easy inspection, cleaning, painting and replacement of parts likely to corrode.

Exam tips

• Always write both anode and cathode reactions and label anodic and cathodic areas in diagrams.
• State the medium: acidic (H2 evolution) vs neutral aerated (O2 absorption) before writing reactions.
• For factors, write the factor name as a heading, then one line of reason and one example.
• Cathodic protection answers need two neat diagrams; marks are lost without them.
• For coating questions, use the scratched galvanized vs tinned sheet comparison; examiners expect it.
• Corrosion in electronic devices is new in the NEP syllabus; prepare a 4-mark note.

FAQs

Where does rust form in oxygen absorption corrosion?

Iron dissolves at the anode, but Fe2+ diffuses towards the cathode where OH- ions are formed, so rust deposits near the cathode while pitting occurs at the anode.

Why is zinc used for galvanizing rather than tin?

Zinc is more active than iron, so it protects iron sacrificially even when the coating is damaged; tin protects only while unbroken.

Why are gold and platinum not corroded by oxygen?

Their oxides are unstable and decompose back to metal and oxygen, so no permanent oxide film or corrosion results.

Is galvanic series the same as electrochemical series?

No. The electrochemical series ranks pure metals by standard electrode potentials, whereas the galvanic series is an experimental ranking of metals and alloys in an actual environment such as sea water, and is more useful for corrosion prediction.

These notes follow the University of Mumbai first-year engineering syllabus (NEP 2020, Applied Chemistry) and were prepared with AI help and checked.