Unit 3 | Engineering Chemistry Notes | AKTU Notes

UNIT 3: Electrochemistry, Batteries, Corrosion and Chemistry of Engineering Materials

PART A: ELECTROCHEMISTRY AND BATTERIES

3.1 Basic Concepts of Electrochemistry

Electrochemistry deals with the relationship between electrical energy and chemical reactions.

Oxidation-Reduction (Redox) Reactions:

• Oxidation: Loss of electrons. The substance that loses electrons is oxidized and is called the reducing agent.
• Reduction: Gain of electrons. The substance that gains electrons is reduced and is called the oxidizing agent.
• Both always occur together — one cannot happen without the other.

Electrochemical Cell:

• A device that converts chemical energy to electrical energy (or vice versa) through redox reactions.
• Galvanic (Voltaic) Cell: Converts chemical energy → electrical energy spontaneously. (Batteries are galvanic cells).
• Electrolytic Cell: Uses electrical energy to drive a non-spontaneous chemical reaction (electroplating, electrolysis of water).

Components of a Galvanic Cell:

• Anode (−): Where oxidation occurs. Metal loses electrons. Negative electrode in galvanic cell.
• Cathode (+): Where reduction occurs. Metal gains electrons. Positive electrode in galvanic cell.
• Electrolyte: Ionic solution that conducts ions between electrodes.
• Salt bridge: Tube filled with inert electrolyte (KCl, KNO₃) that maintains electrical neutrality by allowing ion flow between two half-cells.
• External wire: Electrons flow from anode to cathode through the external circuit.

Standard Electrode Potential (E°):

• The tendency of an electrode to gain or lose electrons is measured by its electrode potential.
• Standard hydrogen electrode (SHE) is assigned E° = 0 V (reference).
• More positive E° → stronger oxidizing agent (easily reduced).
• More negative E° → stronger reducing agent (easily oxidized).

EMF of a Cell:

• E°cell = E°cathode − E°anode
• Positive E°cell → spontaneous reaction → can produce electrical energy.

Nernst Equation:

• Relates electrode potential to concentration: E = E° − (RT/nF) ln Q
• At 25°C: E = E° − (0.0592/n) log Q
• Where R = gas constant, T = temperature, n = number of electrons, F = Faraday's constant (96500 C/mol), Q = reaction quotient.

3.2 Batteries — Classification

A battery is a device that converts stored chemical energy directly into electrical energy through electrochemical reactions.

Classification of Batteries:

• Primary Cells (Disposable Batteries): Cannot be recharged. Once the chemical reactants are consumed, the battery is discarded. Example: Dry Cell (Leclanché cell), alkaline battery.
• Secondary Cells (Rechargeable Batteries): Can be recharged by passing electric current in reverse direction, restoring the original chemical reactants. Example: Lead-acid battery, Li-ion battery, Ni-Cd battery.
• Fuel Cells: Convert chemical energy of a fuel (like hydrogen) directly to electricity. Reactants are continuously fed — does not run out. Example: H₂-O₂ fuel cell.

3.3 Primary Cell — Dry Cell (Leclanché Cell)

Construction:

• Anode (−): Zinc container (the outer casing of the battery is the anode).
• Cathode (+): Carbon (graphite) rod in the center.
• Electrolyte: Paste of ammonium chloride (NH₄Cl) and zinc chloride (ZnCl₂) — this is why it is called a "dry" cell (no liquid).
• Depolarizer: MnO₂ surrounds the carbon rod — it prevents buildup of H₂ gas (which would stop cell action).

Reactions:

• At Anode (oxidation): Zn → Zn²⁺ + 2e⁻
• At Cathode (reduction): 2MnO₂ + 2NH₄⁺ + 2e⁻ → Mn₂O₃ + 2NH₃ + H₂O
• Cell voltage: ~1.5 V

Limitations:

• Cannot be recharged.
• Short shelf life — zinc corrodes even when not in use.
• Voltage drops over time.

Applications:

• Torches, clocks, remote controls, radios, toys.

3.4 Secondary Cell — Lead Acid Battery

The lead acid battery is the most widely used rechargeable battery — found in automobiles, UPS systems, inverters.

Construction:

• Anode (−): Lead (Pb) plates.
• Cathode (+): Lead dioxide (PbO₂) plates.
• Electrolyte: Dilute sulfuric acid (H₂SO₄) solution (about 38% concentration).
• A typical car battery has 6 cells connected in series → 6 × 2V = 12V total.

Discharge Reactions (when battery is providing current):

• At Anode: Pb + SO₄²⁻ → PbSO₄ + 2e⁻ (oxidation)
• At Cathode: PbO₂ + SO₄²⁻ + 4H⁺ + 2e⁻ → PbSO₄ + 2H₂O (reduction)
• Both electrodes become coated with PbSO₄ during discharge. H₂SO₄ is consumed → electrolyte becomes dilute.
• Cell voltage: ~2 V per cell.

Charging Reactions (when external electricity is applied):

• At Anode: PbSO₄ + 2H₂O → PbO₂ + SO₄²⁻ + 4H⁺ + 2e⁻
• At Cathode: PbSO₄ + 2e⁻ → Pb + SO₄²⁻
• PbSO₄ is converted back to Pb and PbO₂. H₂SO₄ is regenerated → electrolyte becomes concentrated again.

Applications:

• Car batteries (starter batteries for engines).
• UPS (Uninterruptible Power Supply).
• Inverters for home power backup.
• Electric forklifts, golf carts.

PART B: CORROSION

3.5 Introduction to Corrosion

Corrosion is the gradual destruction of a metal by chemical or electrochemical reaction with its surrounding environment.

• The most common example: rusting of iron — iron reacts with oxygen and water to form iron oxide (rust).
• Corrosion costs industries billions of dollars every year in repairs and replacements.
• Corrosion is essentially the reverse of metal extraction — metals tend to return to their natural oxide/ore form.

3.6 Types of Corrosion

1. Chemical (Dry) Corrosion:

• Occurs by direct chemical reaction of metal with gases in the environment — without any electrolyte (water).
• Usually occurs at high temperatures.
• Examples: Oxidation (reaction with O₂), sulfidation (reaction with H₂S), carburization.
• Example: 2Fe + O₂ → 2FeO (at high temperature)

2. Electrochemical (Wet) Corrosion:

• Occurs in the presence of an electrolyte (moisture, water).
• Works like a galvanic cell — areas of the metal act as anode and cathode.
• Most common type of corrosion (rusting of iron).
• Two types based on mechanism: Hydrogen evolution type and oxygen absorption type.

Types of Electrochemical Corrosion:

• Galvanic Corrosion: When two different metals are in electrical contact in the presence of electrolyte. The more active metal (lower in electrochemical series) acts as anode and corrodes. Example: Iron and copper in contact — iron corrodes.
• Concentration Cell Corrosion: Corrosion due to difference in concentration of electrolyte or dissolved oxygen at different parts of the metal surface.
• Pitting Corrosion: Localized corrosion forming small pits or holes in the metal surface.
• Crevice Corrosion: Corrosion in confined spaces (crevices) where stagnant solution accumulates.
• Stress Corrosion Cracking: Combined effect of mechanical stress and corrosive environment causing cracks.
• Intergranular Corrosion: Corrosion along grain boundaries of the metal.

3.7 Causes of Corrosion

• Presence of moisture/water: Acts as electrolyte — necessary for electrochemical corrosion.
• Presence of oxygen: Oxidizes the metal surface.
• Presence of CO₂ and acids: Form acidic solutions (carbonic acid) that accelerate corrosion.
• Presence of salts: Increase conductivity of electrolyte — accelerate corrosion. (Seawater causes faster corrosion than fresh water).
• Dissimilar metals in contact: Galvanic corrosion — the more active metal corrodes.
• Stress in metal: Stressed regions act as anodes and corrode faster.
• Surface defects: Scratches, cracks expose fresh metal surface.
• High temperature: Accelerates chemical corrosion.

3.8 Corrosion Prevention and Control

1. Proper Material Selection:

• Use corrosion-resistant materials: stainless steel, titanium, copper, gold, platinum.
• Avoid using dissimilar metals in contact.

2. Protective Coatings:

• Metallic coatings: Galvanizing (coating iron with zinc), tinning (coating with tin), chromium plating, nickel plating.
• Non-metallic coatings: Paint, varnish, enamel, plastic coatings. Act as physical barrier between metal and environment.
• Chemical conversion coatings: Phosphating, chromating, anodizing — convert surface to a protective compound.

3. Cathodic Protection:

• Sacrificial anode method: A more active metal (like zinc or magnesium) is connected to the metal to be protected. The active metal acts as anode and corrodes sacrificially, protecting the main metal (cathode). Example: Zinc blocks on ship hulls, Mg anodes for underground pipelines.
• Impressed current method: An external DC current is applied to make the metal a cathode — prevents oxidation. Used for large structures like oil tanks, ship hulls.

4. Corrosion Inhibitors:

• Chemicals added to the environment (usually liquid) that slow down or stop corrosion.
• Types: Anodic inhibitors (form protective film on anode), cathodic inhibitors (slow down cathode reaction), mixed inhibitors.
• Examples: Chromates, phosphates, benzotriazole, amines.

5. Alloying:

• Adding alloying elements to improve corrosion resistance.
• Example: Stainless steel = iron + 10-30% chromium. Chromium forms a thin, stable Cr₂O₃ layer on surface (passivation) that protects iron underneath.

3.9 Corrosion in Specific Industries

1. Power Generation Industry:

• Boilers suffer from oxygen pitting, caustic embrittlement, and steam corrosion.
• Steam turbine blades corrode due to high-temperature steam and dissolved salts.
• Prevention: Deaeration of boiler water, water treatment, use of stainless steel components.

2. Chemical Processing Industry:

• Equipment exposed to acids, alkalis, solvents, and high temperatures.
• Stress corrosion cracking of stainless steel in chloride solutions.
• Prevention: Use of Hastelloy, titanium, FRP (fiber-reinforced plastics) for highly corrosive environments.

3. Oil and Gas Industry:

• Pipelines corrode due to H₂S (hydrogen sulfide — causes hydrogen embrittlement), CO₂ (forms carbonic acid), saline water.
• Internal corrosion: by produced water and dissolved gases.
• External corrosion: soil corrosion for buried pipelines.
• Prevention: Chemical inhibitors injected into pipelines, cathodic protection, corrosion-resistant alloys.

4. Pulp and Paper Industry:

• Digesters (used in paper making) exposed to hot alkali (white liquor) and acidic solutions.
• Bleaching chemicals (chlorine, chlorine dioxide) are highly corrosive.
• Prevention: Use of stainless steel, titanium, and rubber linings.

PART C: CHEMISTRY OF ENGINEERING MATERIALS

3.10 Cement — Constituents

Cement is a binding material used in construction to bind sand, gravel, and bricks together to form concrete and mortar.

Raw Materials for Cement:

• Limestone (CaCO₃) — provides CaO (calcia).
• Clay (silicates and aluminates of sodium/potassium) — provides SiO₂, Al₂O₃, Fe₂O₃.
• Gypsum (CaSO₄·2H₂O) — added after clinker formation to control setting time.

Composition of Portland Cement:

Oxide	Percentage (%)	Role
CaO (Lime)	60–67%	Main component, provides strength
SiO₂ (Silica)	17–25%	Provides strength and durability
Al₂O₃ (Alumina)	3–8%	Reduces melting point, accelerates setting
Fe₂O₃ (Iron oxide)	0.5–6%	Gives gray color, flux in kiln
MgO (Magnesia)	<5%	Small amount needed; excess causes expansion
SO₃ (Sulfate)	1–3%	Controls setting time (from gypsum)

Main Compounds in Cement Clinker:

• C₃S (Tricalcium silicate, Alite) — 3CaO·SiO₂: Most important compound. Responsible for early strength gain (hardens within first week). About 45–65%.
• C₂S (Dicalcium silicate, Belite) — 2CaO·SiO₂: Responsible for long-term strength (hardens over weeks to months). About 15–30%.
• C₃A (Tricalcium aluminate) — 3CaO·Al₂O₃: Sets very quickly — causes flash setting. Gypsum is added to control this. About 5–10%.
• C₄AF (Tetracalcium aluminoferrite) — 4CaO·Al₂O₃·Fe₂O₃: Gives cement its gray color. Low contribution to strength. About 5–15%.

3.11 Manufacturing of Cement

Two main processes:

1. Dry Process (Modern, more efficient):

• Raw materials (limestone + clay) are crushed and dried.
• Ground into fine powder → mixed → fed into kiln as dry powder.
• Less energy consumption than wet process.
• Used in most modern cement plants.

2. Wet Process (Older):

• Raw materials are mixed with water to form a slurry.
• Slurry is fed into the kiln.
• More energy consumption (to evaporate water).

Steps in Cement Manufacturing:

• Step 1 — Crushing: Limestone and clay are crushed into small pieces.
• Step 2 — Mixing: Crushed materials are mixed in correct proportions.
• Step 3 — Burning in Rotary Kiln: The mixture is heated in a long rotary kiln at different temperature zones: drying zone (100–400°C), calcination zone (600–900°C, CaCO₃ → CaO + CO₂), clinkering zone (1300–1500°C, all compounds react to form clinker).
• Step 4 — Cooling: Hot clinker is cooled rapidly (quenching).
• Step 5 — Grinding: Clinker + Gypsum (3–5%) are ground finely → Cement powder.
• Step 6 — Packing and dispatch.

3.12 Hardening and Setting of Cement

When water is added to cement, two processes occur:

1. Setting:

• The process of cement becoming stiff and losing its plasticity.
• Initial set: Cement starts to stiffen — usually within 30 minutes (gypsum delays C₃A reaction to allow working time).
• Final set: Cement becomes completely rigid — usually within 10 hours.

2. Hardening:

• After setting, cement gains strength over time — this is hardening.
• Hardening continues for weeks, months, and even years.
• C₃S gives early strength (first 28 days), C₂S gives later strength (months to years).

Chemistry of Hardening (Hydration reactions):

• 2C₃S + 6H₂O → C₃S₂H₃ (CSH gel) + 3Ca(OH)₂
• 2C₂S + 4H₂O → C₃S₂H₃ (CSH gel) + Ca(OH)₂
• C₃S₂H₃ = Calcium Silicate Hydrate (CSH) gel — this is the main product responsible for the strength of cement.
• The CSH gel fills pores and binds everything together — giving mechanical strength.

3.13 Deterioration of Cement

Cement and concrete can deteriorate due to various factors:

• Sulfate attack: Sulfates (from soil, groundwater) react with Ca(OH)₂ and C₃A to form expansive compounds (ettringite) → cracking.
• Carbonation: CO₂ from air reacts with Ca(OH)₂ → CaCO₃. Reduces pH → corrodes steel reinforcement.
• Chloride attack: Chlorides from seawater or de-icing salts penetrate concrete → corrode steel reinforcement (rusting causes expansion → cracking).
• Alkali-silica reaction (ASR): Alkalis in cement react with reactive silica in aggregates → gel that expands with water → cracking.
• Freeze-thaw damage: Water in pores freezes → expands → cracks the concrete.
• Fire damage: High temperature breaks down CSH gel → loss of strength.

3.14 Plaster of Paris (POP)

Plaster of Paris is calcium sulphate hemihydrate — CaSO₄·½H₂O (or 2CaSO₄·H₂O).

Preparation:

• Gypsum (CaSO₄·2H₂O) is heated to 120–130°C:
• CaSO₄·2H₂O → CaSO₄·½H₂O + 3/2 H₂O
• If heated above 200°C → anhydrous CaSO₄ (dead burnt plaster) — does not set with water.

Setting of Plaster of Paris:

• When mixed with water, POP sets (hardens) within a few minutes:
• CaSO₄·½H₂O + 3/2 H₂O → CaSO₄·2H₂O (gypsum)
• It expands slightly on setting — this is useful as it fills molds completely.

Applications of Plaster of Paris:

• Medical: Casts for fractured bones (orthopedic plaster).
• Dentistry: Making dental molds and models.
• Construction: Plasterwork on walls and ceilings, decorative plaster moldings.
• Making statues, toys, chalk, and art models.
• Fire-resistant coating (gypsum boards).
• Mold making in ceramics and metal casting.