🎯 Learning Objectives:
• Understand different types of chemical reactions and their characteristics
• Identify reactants, products, and conditions in chemical equations
• Classify reactions based on electron transfer and structural changes
• Apply knowledge of activity series in displacement reactions
• Understand the role of catalysts and energy changes in reactions
• Connect chemical reactions to real-world applications and processes
Chemical Reaction: A process in which one or more substances (reactants) are transformed into different substances (products) with new chemical properties, involving the breaking and forming of chemical bonds.
Reactants: The starting substances that undergo chemical change in a reaction. Written on the left side of a chemical equation.
Products: The new substances formed as a result of chemical reaction. Written on the right side of a chemical equation.
⚡ CHEMICAL REACTION FUNDAMENTALS ⚡
The Heart of Chemistry - Transformation of Matter
REACTANTS → PRODUCTS
Energy changes, bond breaking & forming, new properties emerge!
🧠 Understanding Chemical Reactions:
Think of chemical reactions like cooking recipes. Just as ingredients (reactants) combine under specific conditions (temperature, mixing) to create a new dish (products), chemical substances combine under specific conditions to form entirely new compounds with different properties.
Characteristics of Chemical Reactions
1. Observable Changes
Physical Evidence:
• Color change
• Gas evolution (bubbling)
• Precipitate formation
• Temperature change
• Light or sound production
• Odor change
Examples:
• Iron rusting → red-brown color
• Acid + Base → heat evolution
• AgNO₃ + NaCl → white precipitate
2. Energy Changes
Energy Transformations:
• Exothermic: Heat released
• Endothermic: Heat absorbed
• Activation energy required
• Bond breaking (energy input)
• Bond formation (energy release)
Examples:
• Combustion → exothermic
• Photosynthesis → endothermic
• Respiration → exothermic
3. Conservation Laws
What is Conserved:
• Mass (atoms neither created nor destroyed)
• Energy (converted between forms)
• Charge (electrical neutrality maintained)
• Atomic number of each element
• Total number of atoms
Lavoisier's Law: "In a chemical reaction, mass is neither created nor destroyed, only rearranged."
Energy Profile of Chemical Reactions
Activation Energy
↓
Reactants
→
Products
↓
Energy Released (Exothermic)
Activation energy is the minimum energy required to start a reaction
Types of Chemical Reactions
🔄 CLASSIFICATION OF CHEMICAL REACTIONS 🔄
Based on structural changes and electron transfer
COMBINATION
DECOMPOSITION
DISPLACEMENT
DOUBLE DISPLACEMENT
Each type follows specific patterns and principles!
1. Combination Reactions
Definition: Two or more reactants combine to form a single product.
General Form: A + B → AB
Examples:
• 2H₂ + O₂ → 2H₂O (Water formation)
• N₂ + 3H₂ → 2NH₃ (Ammonia synthesis)
• CaO + H₂O → Ca(OH)₂ (Slaked lime)
• C + O₂ → CO₂ (Carbon dioxide)
Also called: Synthesis reactions or Addition reactions
2. Decomposition Reactions
Definition: A single compound breaks down into two or more simpler substances.
General Form: AB → A + B
Examples:
• 2H₂O → 2H₂ + O₂ (Electrolysis)
• CaCO₃ → CaO + CO₂ (Thermal decomposition)
• 2KClO₃ → 2KCl + 3O₂ (Oxygen preparation)
• 2HgO → 2Hg + O₂ (Mercury oxide heating)
Types: Thermal, electrolytic, and photolytic decomposition
3. Displacement Reactions
Definition: A more reactive element displaces a less reactive element from its compound.
General Form: A + BC → AC + B
Examples:
• Zn + CuSO₄ → ZnSO₄ + Cu (Zinc displaces copper)
• Fe + CuSO₄ → FeSO₄ + Cu (Iron displaces copper)
• Mg + 2HCl → MgCl₂ + H₂ (Hydrogen gas evolution)
• 2Al + 3CuSO₄ → Al₂(SO₄)₃ + 3Cu
Key Factor: Follows the reactivity series of metals
4. Double Displacement
Definition: Two compounds exchange their ions to form two new compounds.
General Form: AB + CD → AD + CB
Examples:
• AgNO₃ + NaCl → AgCl + NaNO₃ (Precipitation)
• BaCl₂ + Na₂SO₄ → BaSO₄ + 2NaCl (White precipitate)
• HCl + NaOH → NaCl + H₂O (Neutralization)
• Pb(NO₃)₂ + 2KI → PbI₂ + 2KNO₃ (Yellow precipitate)
Also called: Metathesis or Exchange reactions
Displacement Reaction Example: Zn + CuSO₄ → ZnSO₄ + Cu
Zn
+
Cu
SO₄
→
Zn
SO₄
+
Cu
Zinc (more reactive) displaces copper from copper sulfate
Oxidation and Reduction
Oxidation: A process involving loss of electrons, gain of oxygen, or loss of hydrogen. The substance that gets oxidized is called a reducing agent.
Reduction: A process involving gain of electrons, loss of oxygen, or gain of hydrogen. The substance that gets reduced is called an oxidizing agent.
🔄 REDOX REACTIONS 🔄
Oxidation and Reduction occur simultaneously
Remember: OIL RIG (Oxidation Is Loss, Reduction Is Gain of electrons)
Oxidation Examples
Electron Loss:
• Mg → Mg²⁺ + 2e⁻ (Magnesium oxidized)
• Fe²⁺ → Fe³⁺ + e⁻ (Iron oxidized)
Oxygen Gain:
• 2Mg + O₂ → 2MgO (Magnesium gains oxygen)
• C + O₂ → CO₂ (Carbon gains oxygen)
Hydrogen Loss:
• H₂S + Cl₂ → 2HCl + S (H₂S loses hydrogen)
Reduction Examples
Electron Gain:
• Cu²⁺ + 2e⁻ → Cu (Copper reduced)
• Cl₂ + 2e⁻ → 2Cl⁻ (Chlorine reduced)
Oxygen Loss:
• CuO + H₂ → Cu + H₂O (CuO loses oxygen)
• Fe₂O₃ + 3CO → 2Fe + 3CO₂ (Iron oxide reduced)
Hydrogen Gain:
• Cl₂ + H₂ → 2HCl (Chlorine gains hydrogen)
🔍 Identifying Redox Reactions
1Assign oxidation numbers: Determine oxidation states of all atoms
2Compare oxidation numbers: Look for changes in oxidation states
3Identify oxidation: Element with increased oxidation number
4Identify reduction: Element with decreased oxidation number
5Name agents: Reducing agent (gets oxidized), Oxidizing agent (gets reduced)
Example: Zn + CuSO₄ → ZnSO₄ + Cu
Zn⁰ → Zn²⁺ (Oxidation, Zn is reducing agent)
Cu²⁺ → Cu⁰ (Reduction, CuSO₄ is oxidizing agent)
Reactivity Series and Activity of Metals
Reactivity Series: An arrangement of metals in order of decreasing reactivity, used to predict the outcome of displacement reactions.
⚡ REACTIVITY SERIES OF METALS ⚡
From Most Reactive to Least Reactive
More reactive metals displace less reactive ones!
📊 Metal Activity Series
Most Reactive (Top)
K
Na
Ca
Mg
Al
Zn
Fe
Pb
Sn
Hydrogen (Reference)
Cu
Hg
Ag
Au
Pt
Least Reactive (Bottom)
🧠 Memory Aid for Reactivity Series:
"Please Send Charlie's Monkeys And Zebras In Lead Suits, Climbing Mighty Silver And Gold Peaks"
K, Na, Ca, Mg, Al, Zn, Fe, Pb, Sn, H, Cu, Hg, Ag, Au, Pt
Applications of Reactivity Series
Predictions:
• Displacement reactions
• Hydrogen gas evolution
• Metal extraction methods
• Corrosion tendencies
• Reactivity with acids/water
Rules:
• Metals above H₂ displace hydrogen from acids
• Higher metal displaces lower metal from solutions
• More reactive metals are harder to extract
Displacement Predictions
Will React:
• Zn + CuSO₄ → ZnSO₄ + Cu ✓
• Fe + ZnSO₄ → FeSO₄ + Zn ✗
• Mg + HCl → MgCl₂ + H₂ ✓
• Cu + HCl → No reaction ✗
Remember: Only metals above hydrogen can displace hydrogen from acids
Conditions Affecting Chemical Reactions
🌡️ FACTORS AFFECTING REACTION RATES 🌡️
External conditions influence how fast reactions occur
Optimizing conditions is crucial in industry!
1. Temperature
Effect: Higher temperature increases reaction rate by providing more kinetic energy to molecules.
Examples:
• Cooking food faster at higher temperatures
• Industrial ammonia synthesis at 400-500°C
• Decomposition reactions need heating
• Refrigeration slows food spoilage reactions
Rule: Generally, 10°C increase doubles reaction rate
2. Concentration
Effect: Higher concentration of reactants increases reaction rate due to more frequent collisions.
Examples:
• Concentrated acids react faster than dilute acids
• Pure oxygen burns substances faster than air
• Industrial processes use concentrated reactants
• Breathing pure oxygen vs. air
3. Surface Area
Effect: Larger surface area (smaller particles) increases reaction rate by exposing more reactive sites.
Examples:
• Powdered iron rusts faster than iron bars
• Sugar cubes vs. powdered sugar in tea
• Catalytic converters use fine particles
• Antacid tablets vs. powder
4. Pressure (for gases)
Effect: Higher pressure increases concentration of gaseous reactants, thus increasing reaction rate.
Examples:
• Haber process for ammonia (200-300 atm)
• Contact process for sulfuric acid
• Combustion in pressurized engines
• Deep-sea diving and gas solubility
Catalysts
Catalyst: A substance that increases the rate of a chemical reaction without being consumed in the process. It provides an alternative pathway with lower activation energy.
⚗️ CATALYSTS - REACTION ACCELERATORS ⚗️
Speed up reactions without being used up
Essential for industrial processes and biological systems!
Types of Catalysts
Homogeneous Catalyst: Same phase as reactants
• Example: H₂SO₄ in esterification
Heterogeneous Catalyst: Different phase from reactants
• Example: Ni in hydrogenation
Enzyme: Biological catalyst
• Example: Amylase in digestion
Industrial Catalysts
Important Examples:
• Iron: Haber process (NH₃ production)
• Vanadium(V) oxide: Contact process (H₂SO₄)
• Nickel: Hydrogenation of oils
• Platinum: Catalytic converters
• Manganese dioxide: Decomposition of H₂O₂
🔬 How Catalysts Work
Step 1: Adsorption - Reactants bind to catalyst surface
Step 2: Activation - Bonds weakened, activation energy lowered
Step 3: Reaction - Products formed on catalyst surface
Step 4: Desorption - Products leave, catalyst regenerated
Energy Profile with Catalyst
Without Catalyst:
High Activation Energy
With Catalyst:
Lower Activation Energy
Same starting and ending energy, but easier pathway!
Real-World Applications
🌍 CHEMICAL REACTIONS IN DAILY LIFE 🌍
Chemistry shapes our world and everyday experiences
From digestion to technology, reactions are everywhere!
Biological Processes
Cellular Respiration:
C₆H₁₂O₆ + 6O₂ → 6CO₂ + 6H₂O + ATP
(Glucose oxidation for energy)
Photosynthesis:
6CO₂ + 6H₂O + light → C₆H₁₂O₆ + 6O₂
(Plant food production)
Digestion:
Enzymes catalyze breakdown of food molecules
Industrial Applications
Steel Production:
Fe₂O₃ + 3CO → 2Fe + 3CO₂
(Iron extraction from ore)
Cement Making:
CaCO₃ → CaO + CO₂
(Limestone decomposition)
Fertilizer Production:
N₂ + 3H₂ → 2NH₃
(Ammonia synthesis)
Environmental Chemistry
Corrosion:
4Fe + 3O₂ + 6H₂O → 4Fe(OH)₃
(Iron rusting)
Acid Rain Formation:
SO₂ + H₂O → H₂SO₃
(Environmental pollution)
Ozone Depletion:
Cl + O₃ → ClO + O₂
(Stratospheric reactions)
Household Chemistry
Cooking:
• Browning reactions (Maillard)
• Fermentation in bread
• Caramelization of sugars
Cleaning:
• Soap action (saponification)
• Bleaching reactions
• Antacid neutralization
Food Preservation:
• Oxidation prevention
• Pickling (acid reactions)
• Refrigeration slowing reactions
⚠️ Common Misconceptions:
• Misconception: "Catalysts are consumed in reactions"
Truth: Catalysts are regenerated and can be used repeatedly
• Misconception: "All reactions need high temperature"
Truth: Many biological reactions occur at body temperature
• Misconception: "Chemical equations show what happens step by step"
Truth: Equations show overall change, not necessarily the mechanism
🎯 Practice Examples
Classify these reactions:
1.
2KClO₃ → 2KCl + 3O₂ → Decomposition
2.
Mg + 2HCl → MgCl₂ + H₂ → Displacement
3.
NaOH + HCl → NaCl + H₂O → Double displacement
4.
C + O₂ → CO₂ → Combination
5.
Zn + CuSO₄ → ZnSO₄ + Cu → Displacement + Redox
Section B: Short Answer Questions
Q1. Define chemical reaction and list four observable signs of chemical reactions.
Answer: Chemical reaction is a process where reactants transform into products with new chemical properties through bond breaking and forming. Observable signs: (1) Color change, (2) Gas evolution, (3) Precipitate formation, (4) Temperature change, (5) Light/sound production, (6) Odor change.
Q2. Classify the following reactions: (a) 2H₂ + O₂ → 2H₂O (b) CaCO₃ → CaO + CO₂ (c) Zn + H₂SO₄ → ZnSO₄ + H₂
Answer: (a) Combination reaction - two reactants form one product, (b) Decomposition reaction - one compound breaks into two products, (c) Displacement reaction - zinc displaces hydrogen from sulfuric acid.
Q3. What is meant by oxidation and reduction? Give one example each.
Answer: Oxidation: Loss of electrons, gain of oxygen, or loss of hydrogen. Example: Mg → Mg²⁺ + 2e⁻. Reduction: Gain of electrons, loss of oxygen, or gain of hydrogen. Example: Cu²⁺ + 2e⁻ → Cu. Both occur simultaneously in redox reactions.
Q4. Explain the reactivity series and its applications.
Answer: Reactivity series arranges metals in decreasing order of reactivity (K, Na, Ca, Mg, Al, Zn, Fe, Pb, H, Cu, Ag, Au). Applications: Predicting displacement reactions, hydrogen evolution from acids, metal extraction methods, and corrosion tendencies.
Q5. What is a catalyst? How does it affect chemical reactions?
Answer: Catalyst is a substance that increases reaction rate without being consumed. It lowers activation energy by providing alternative pathway. Effects: Speeds up reactions, remains unchanged after reaction, enables reactions at lower temperatures, improves efficiency in industrial processes.
Q6. List four factors that affect the rate of chemical reactions with examples.
Answer: (1) Temperature: Cooking faster at high heat, (2) Concentration: Concentrated acids react faster, (3) Surface area: Powdered sugar dissolves faster than cubes, (4) Pressure: High pressure in Haber process, (5) Catalyst: MnO₂ speeds H₂O₂ decomposition.
Q7. Differentiate between combination and decomposition reactions with examples.
Answer: Combination: Two or more reactants form one product (A + B → AB). Example: 2H₂ + O₂ → 2H₂O. Decomposition: One compound breaks into two or more products (AB → A + B). Example: 2H₂O → 2H₂ + O₂.
Q8. What are double displacement reactions? Give two examples.
Answer: Double displacement reactions involve exchange of ions between two compounds (AB + CD → AD + CB). Examples: (1) AgNO₃ + NaCl → AgCl + NaNO₃ (precipitation), (2) HCl + NaOH → NaCl + H₂O (neutralization).
Q9. Explain why iron displaces copper from copper sulfate solution.
Answer: Iron is more reactive than copper in the reactivity series. According to displacement rules, more reactive metals can displace less reactive metals from their salt solutions. Fe + CuSO₄ → FeSO₄ + Cu. Iron loses electrons (oxidized) while copper gains electrons (reduced).
Q10. Give three examples of chemical reactions occurring in daily life.
Answer: (1) Respiration: C₆H₁₂O₆ + 6O₂ → 6CO₂ + 6H₂O + energy (cellular energy production), (2) Photosynthesis: 6CO₂ + 6H₂O + light → C₆H₁₂O₆ + 6O₂ (plant food making), (3) Rusting: 4Fe + 3O₂ + 6H₂O → 4Fe(OH)₃ (iron corrosion).
Section C: Long Answer Questions
Q1. Describe the different types of chemical reactions with definitions, general equations, and detailed examples.
Solution: Four main types exist. (1) Combination (A + B → AB): Two or more reactants form single product. Examples: 2H₂ + O₂ → 2H₂O (water formation), N₂ + 3H₂ → 2NH₃ (ammonia synthesis), CaO + H₂O → Ca(OH)₂ (slaked lime). (2) Decomposition (AB → A + B): Single compound breaks into simpler substances. Examples: 2H₂O → 2H₂ + O₂ (electrolysis), CaCO₃ → CaO + CO₂ (thermal decomposition), 2KClO₃ → 2KCl + 3O₂ (oxygen preparation). (3) Displacement (A + BC → AC + B): More reactive element displaces less reactive. Examples: Zn + CuSO₄ → ZnSO₄ + Cu, Mg + 2HCl → MgCl₂ + H₂. (4) Double Displacement (AB + CD → AD + CB): Ion exchange between compounds. Examples: AgNO₃ + NaCl → AgCl + NaNO₃ (precipitation), HCl + NaOH → NaCl + H₂O (neutralization). Each type follows specific patterns and has distinct applications in industry and nature.
Q2. Explain oxidation and reduction in detail with examples, including electron transfer and oxidizing/reducing agents.
Solution: Oxidation: Process involving electron loss, oxygen gain, or hydrogen loss. Examples: Mg → Mg²⁺ + 2e⁻ (electron loss), 2Mg + O₂ → 2MgO (oxygen gain), H₂S + Cl₂ → 2HCl + S (hydrogen loss). Reduction: Process involving electron gain, oxygen loss, or hydrogen gain. Examples: Cu²⁺ + 2e⁻ → Cu (electron gain), CuO + H₂ → Cu + H₂O (oxygen loss), Cl₂ + H₂ → 2HCl (hydrogen gain). Redox Reactions: Oxidation and reduction occur simultaneously. In Zn + CuSO₄ → ZnSO₄ + Cu: Zn loses electrons (oxidized), Cu²⁺ gains electrons (reduced). Agents: Reducing agent gets oxidized (Zn), oxidizing agent gets reduced (CuSO₄). Memory aid: OIL RIG (Oxidation Is Loss, Reduction Is Gain of electrons). Applications: Metal extraction, electroplating, batteries, corrosion, biological processes like respiration and photosynthesis.
Q3. Discuss the reactivity series of metals in detail, including its arrangement, applications, and prediction of chemical reactions.
Solution: Reactivity Series: Arrangement of metals in decreasing order of reactivity: K, Na, Ca, Mg, Al, Zn, Fe, Pb, Sn, H, Cu, Hg, Ag, Au, Pt. Basis: Tendency to lose electrons and form positive ions. Applications: (1) Displacement Reactions: Metals above displace those below from salt solutions. Zn + CuSO₄ → ZnSO₄ + Cu (works), Cu + ZnSO₄ → No reaction. (2) Hydrogen Evolution: Metals above hydrogen displace H₂ from acids. Mg + 2HCl → MgCl₂ + H₂ (works), Cu + HCl → No reaction. (3) Metal Extraction: More reactive metals need stronger reducing agents/electrolysis. K, Na extracted by electrolysis; Fe, Zn by carbon reduction; Ag, Au found free. (4) Corrosion: More reactive metals corrode faster. (5) Industrial Applications: Predicting feasibility of metallurgical processes, galvanization (Zn protects Fe), metal purification. Memory aid: "Please Send Charlie's Monkeys And Zebras In Lead Suits, Climbing Mighty Silver And Gold Peaks."
Q4. Explain the factors affecting chemical reaction rates with detailed examples and practical applications.
Solution: Five main factors affect reaction rates: (1) Temperature: Higher temperature increases molecular kinetic energy, more frequent and energetic collisions. Examples: Food cooks faster at high heat, enzymes work optimally at body temperature, industrial reactions need controlled heating. Rule: 10°C increase approximately doubles rate. (2) Concentration: Higher concentration means more particles per unit volume, increased collision frequency. Examples: Concentrated HCl reacts faster with metals than dilute HCl, pure oxygen burns substances faster than air, industrial processes use concentrated reactants for efficiency. (3) Surface Area: Larger surface area exposes more reactive sites. Examples: Powdered iron rusts faster than iron bars, sugar powder dissolves faster than cubes, catalytic converters use fine particles, antacid powder works faster than tablets. (4) Pressure (gases): Higher pressure increases gas concentration. Examples: Haber process (200-300 atm for NH₃), contact process for H₂SO₄, combustion engines use compression. (5) Catalyst: Lowers activation energy, provides alternative pathway. Examples: Enzymes in digestion, platinum in catalytic converters, iron in Haber process, MnO₂ in H₂O₂ decomposition. Applications: Food preservation (refrigeration), industrial optimization, pharmaceutical manufacturing, environmental control.
Q5. Describe catalysts in detail, including types, mechanism of action, and industrial applications.
Solution: Catalysts: Substances that increase reaction rate without being consumed, providing alternative pathway with lower activation energy. Types: (1) Homogeneous: Same phase as reactants. Example: H₂SO₄ in esterification reactions, acid-base catalysis in solution. (2) Heterogeneous: Different phase from reactants. Example: Ni in hydrogenation of oils, Pt in catalytic converters, solid catalysts with gaseous reactants. (3) Enzymes: Biological catalysts with high specificity. Examples: Amylase (starch digestion), catalase (H₂O₂ decomposition), pepsin (protein digestion). Mechanism: (1) Adsorption of reactants on catalyst surface, (2) Activation and bond weakening, (3) Product formation at active sites, (4) Desorption of products, (5) Catalyst regeneration. Industrial Applications: Iron in Haber process (NH₃ production), V₂O₅ in contact process (H₂SO₄), Ni in margarine production, Pt in petroleum refining, zeolites in petrochemicals, enzymes in biotechnology. Characteristics: Unchanged after reaction, small quantities needed, specific for particular reactions, can be poisoned by impurities, provide faster route to equilibrium without changing position.
Q6. Analyze chemical reactions in biological systems and their importance to life processes.
Solution: Cellular Respiration: C₆H₁₂O₆ + 6O₂ → 6CO₂ + 6H₂O + ATP. Glucose oxidation releases energy for cellular activities. Occurs in mitochondria through glycolysis, Krebs cycle, electron transport. Essential for metabolism, movement, growth, reproduction. Photosynthesis: 6CO₂ + 6H₂O + light energy → C₆H₁₂O₆ + 6O₂. Plants convert light energy to chemical energy, producing food and oxygen. Occurs in chloroplasts using chlorophyll catalyst. Foundation of food chains, oxygen production for atmosphere. Digestion: Enzyme-catalyzed breakdown of complex molecules. Amylase converts starch to sugars, pepsin breaks proteins, lipase digests fats. Enables nutrient absorption and utilization. Protein Synthesis: Amino acids combine to form proteins through peptide bonds. DNA → RNA → Proteins. Essential for growth, repair, enzyme production. Detoxification: Liver enzymes break down harmful substances through oxidation, conjugation reactions. Blood Chemistry: Hemoglobin-oxygen binding (reversible), pH buffering systems, clotting cascades. Importance: Energy production, nutrient processing, waste removal, growth and repair, homeostasis maintenance, adaptation to environment. All controlled by enzyme catalysts with high specificity and efficiency.
Q7. Explain industrial chemical reactions and their applications in manufacturing and technology.
Solution: Ammonia Production (Haber Process): N₂ + 3H₂ ⇌ 2NH₃. Uses iron catalyst, 400-500°C, 200-300 atm pressure. Produces fertilizers feeding billions, explosives, plastics. Reversible reaction requiring optimization of conditions. Sulfuric Acid (Contact Process): S + O₂ → SO₂, then 2SO₂ + O₂ → 2SO₃ (V₂O₅ catalyst), finally SO₃ + H₂SO₄ → H₂S₂O₇, then H₂S₂O₇ + H₂O → 2H₂SO₄. Used in fertilizers, petroleum refining, steel processing, batteries. Steel Production: Fe₂O₃ + 3CO → 2Fe + 3CO₂ in blast furnaces. Carbon monoxide reduces iron ore at high temperatures. Essential for construction, automobiles, machinery. Petroleum Refining: Cracking long hydrocarbons into shorter ones, reforming for octane improvement, polymerization for plastics. Catalysts include zeolites, platinum, aluminum compounds. Pharmaceutical Manufacturing: Synthesis of drugs through controlled reactions, crystallization for purification, enzyme catalysis for specificity. Electrochemical Industries: Aluminum production by electrolysis, chlorine-alkali process for NaOH and Cl₂, electroplating for corrosion protection. Environmental Applications: Catalytic converters reduce vehicle emissions, scrubbing systems remove SO₂ from power plants, water treatment uses oxidation/reduction reactions. Polymer Industry: Addition and condensation polymerization create plastics, rubbers, fibers that revolutionized modern life.
Q8. Discuss environmental chemical reactions and their impact on ecosystems and human health.
Solution: Atmospheric Chemistry: (1) Ozone Layer: O₂ + UV → 2O, then O + O₂ → O₃ (ozone formation). CFCs destroy ozone: Cl + O₃ → ClO + O₂, leading to UV radiation increase, skin cancer, ecosystem damage. (2) Acid Rain: SO₂ + H₂O → H₂SO₃, NO₂ + H₂O → HNO₃. Causes forest damage, building corrosion, aquatic ecosystem acidification. Greenhouse Effect: CO₂, CH₄, N₂O trap heat through molecular vibrations. Fossil fuel combustion: C₈H₁₈ + 12.5O₂ → 8CO₂ + 9H₂O causes climate change, sea level rise, weather pattern disruption. Water Pollution: Heavy metal contamination through industrial discharge, eutrophication from excess nutrients causing algal blooms and oxygen depletion, pesticide persistence in food chains. Soil Chemistry: Nutrient cycling through decomposition reactions, pH changes affecting plant growth, contamination by industrial chemicals and heavy metals. Corrosion: 4Fe + 3O₂ + 6H₂O → 4Fe(OH)₃ (rusting) costs billions annually in infrastructure damage. Bioremediation: Using bacterial reactions to break down pollutants, phytoremediation with plants to absorb contaminants. Solutions: Catalytic converters, scrubbing technologies, green chemistry principles, renewable energy adoption, sustainable industrial processes. Impact: Health effects from air/water pollution, ecosystem disruption, economic costs, need for environmental regulations and clean technologies.