Chemistry (Part 1 & Part 2)

1. Structure of Atom

Chemistry begins with understanding the atom, the smallest particle of an element that retains its chemical properties. According to modern atomic theory, an atom consists of three fundamental particles: protons, neutrons, and electrons.

Protons are positively charged particles located in the nucleus. The number of protons determines the atomic number (Z) of an element and identifies the element itself. Hydrogen has one proton, carbon has six, oxygen has eight, and sodium has eleven.

Neutrons are neutral particles also present in the nucleus. Together with protons, they contribute to the mass number (A). Atoms of the same element having different numbers of neutrons are called isotopes. For example, Carbon-12 and Carbon-14 are isotopes of carbon.

Electrons are negatively charged particles revolving around the nucleus in various energy levels or shells. Their mass is negligible compared to protons and neutrons.

Electronic configuration describes the arrangement of electrons in shells.

• Maximum electrons in a shell = 2n² (where n is the shell number)
• K shell = 2 electrons
• L shell = 8 electrons
• M shell = 18 electrons
• N shell = 32 electrons

Important Principles Governing Electron Filling

• Aufbau Principle: Electrons occupy lower energy orbitals first.
• Pauli Exclusion Principle: No two electrons in an atom can have the same four quantum numbers.
• Hund's Rule: Electrons occupy orbitals singly before pairing.

Orbitals are regions where the probability of finding an electron is maximum. Types of orbitals:

• s orbital – spherical
• p orbital – dumbbell
• d orbital – cloverleaf
• f orbital – complex shape

Valence electrons are the electrons present in the outermost shell and determine chemical properties. Atoms having complete outer shells are chemically stable.

Examples of noble gases: Helium = 2, Neon = 2,8, Argon = 2,8,8.

2. Periodic Table

The modern periodic table was developed by Henry Moseley based on increasing atomic number.

• 7 Periods
• 18 Groups
• Elements arranged according to atomic number

Classification

• Metals: Generally found on the left side. They are good conductors, malleable, ductile and lose electrons to form positive ions.
• Non-metals: Located on the right side. Poor conductors except graphite. Gain electrons to form negative ions.
• Metalloids: Show both metallic and non-metallic properties. Examples: Boron, Silicon, Germanium, Arsenic, Antimony, Tellurium.

Periodic Trends

• Atomic Radius: Increases down a group due to addition of electron shells. Decreases across a period because nuclear attraction increases.
• Ionization Energy: Energy required to remove an electron. Increases across a period. Decreases down a group.
• Electron Affinity: Tendency to gain electrons. Generally increases across a period.
• Electronegativity: Ability of an atom to attract bonding electrons. Fluorine has the highest electronegativity.
• Metallic Character: Increases down a group. Decreases across a period.
• Reactivity: Alkali metals are highly reactive because they lose electrons easily. Halogens are highly reactive because they gain electrons easily.

3. Chemical Bonding

Atoms combine to attain stable electronic configuration. Types of chemical bonds:

Ionic Bond

Formed by complete transfer of electrons. Usually between metals and non-metals.

• High melting point
• Conduct electricity in molten or aqueous state
• Hard and brittle
• Water soluble

Covalent Bond

Formed by sharing of electrons. Occurs between non-metals. Examples: H₂, O₂, N₂, CH₄.

• Low melting point
• Poor electrical conductor
• Generally insoluble in water

Coordinate Covalent Bond

Both shared electrons are donated by one atom. Examples: NH₄⁺, H₃O⁺, CO.

Metallic Bond

Positive metal ions are surrounded by a sea of free electrons. Responsible for: Electrical conductivity, Thermal conductivity, Malleability, Ductility, Lustre.

Hydrogen Bonding

A special intermolecular force responsible for: High boiling point of water, Ice floating on water, Protein structure, DNA stability.

4. States of Matter

Matter exists in three common states.

• Solid: Definite shape, definite volume, strong intermolecular force, least compressible.
• Liquid: Definite volume, no definite shape, moderate intermolecular force, flows easily.
• Gas: No definite shape, no definite volume, highly compressible, weak intermolecular force, diffuses rapidly.

Changes of State

• Melting: Solid → Liquid
• Freezing: Liquid → Solid
• Vaporization: Liquid → Gas
• Condensation: Gas → Liquid
• Sublimation: Solid → Gas directly (Examples: Camphor, Naphthalene, Dry ice)

Evaporation is a surface phenomenon. Boiling occurs throughout the liquid.

Factors Affecting Evaporation

• Temperature
• Surface area
• Humidity
• Wind speed

5. Gas Laws

• Boyle's Law: At constant temperature, pressure is inversely proportional to volume. P₁V₁ = P₂V₂
• Charles' Law: At constant pressure, volume is directly proportional to absolute temperature. V₁/T₁ = V₂/T₂
• Gay-Lussac Law: Pressure is directly proportional to temperature. P₁/T₁ = P₂/T₂
• Combined Gas Law: P₁V₁/T₁ = P₂V₂/T₂

Ideal Gas Equation

Where: P = Pressure, V = Volume, n = Number of moles, R = Universal gas constant, T = Absolute temperature.

Real gases deviate from ideal behaviour at high pressure and low temperature.

6. Mole Concept

The mole is the SI unit representing the amount of substance. One mole contains 6.022 × 10²³ particles, called the Avogadro Number.

Molar Mass: Mass of one mole of a substance.

• 1 mole Carbon = 12 g
• 1 mole Oxygen = 32 g
• 1 mole Water = 18 g

Applications

• Chemical calculations
• Stoichiometry
• Gas calculations
• Preparation of solutions

7. Stoichiometry

Stoichiometry deals with quantitative relationships between reactants and products. Balanced equations are essential.

• Law of Conservation of Mass: Mass of reactants equals mass of products.
• Limiting reagent: Reactant consumed first.
• Excess reagent: Remaining reactant after reaction.
• Percentage yield: (Actual Yield / Theoretical Yield) × 100

Important Laws

• Law of Constant Proportions
• Law of Multiple Proportions
• Gay-Lussac's Law of Combining Volumes
• Avogadro's Law

8. Solutions

A solution is a homogeneous mixture.

• Solute: Substance dissolved.
• Solvent: Medium in which solute dissolves.

Types

• Solid in liquid
• Liquid in liquid
• Gas in liquid

Concentration Terms

• Molarity (M)
• Molality (m)
• Normality (N)
• Mole Fraction
• Percentage composition

Factors Affecting Solubility

• Temperature
• Pressure (for gases)
• Nature of solute
• Nature of solvent

• Saturated solution: Cannot dissolve more solute.
• Unsaturated solution: Can dissolve additional solute.
• Supersaturated solution: Contains more dissolved solute than saturation limit.

9. Acids, Bases and Salts

An acid is a substance that produces H⁺ ions in aqueous solution. A base produces OH⁻ ions, while an alkali is a water-soluble base. A salt is formed by the reaction between an acid and a base (neutralization).

Theories of Acids and Bases

• Arrhenius Theory: Acid → Produces H⁺ ions in water. Base → Produces OH⁻ ions in water.
• Bronsted-Lowry Theory: Acid → Proton donor. Base → Proton acceptor.
• Lewis Theory: Acid → Electron pair acceptor. Base → Electron pair donor.

Examples: HCl, H₂SO₄, HNO₃ (Acids); NaOH, KOH, Ca(OH)₂ (Bases).

Properties of Acids

• Sour taste
• Turn blue litmus red
• React with metals to produce hydrogen gas
• React with carbonates to produce carbon dioxide
• Conduct electricity in aqueous solution

Properties of Bases

• Bitter taste
• Slippery feel
• Turn red litmus blue
• Neutralize acids
• Conduct electricity in solution

pH Scale

• pH < 7 → Acidic
• pH = 7 → Neutral
• pH > 7 → Basic

Strong acids: HCl, HNO₃, H₂SO₄. Weak acids: CH₃COOH, H₂CO₃. Strong bases: NaOH, KOH. Weak bases: NH₄OH.

Indicators

• Natural: Litmus, Turmeric, China rose
• Synthetic: Phenolphthalein, Methyl orange, Methyl red

Neutralization

• Antacids relieve acidity
• Liming neutralizes acidic soil
• Baking soda treats insect bites

10. Redox Reactions

Redox reactions involve oxidation and reduction occurring simultaneously.

• Oxidation: Addition of oxygen, removal of hydrogen, loss of electrons.
• Reduction: Removal of oxygen, addition of hydrogen, gain of electrons.
• Oxidizing agent: Accepts electrons and causes oxidation.
• Reducing agent: Donates electrons and causes reduction.

Examples: Rusting, Combustion, Photosynthesis, Respiration, Electroplating.

11. Electrochemistry

Electrochemistry studies the relationship between electricity and chemical reactions.

Electrolytes

Substances conducting electricity in molten or aqueous state.

• Strong electrolytes: HCl, NaOH, NaCl
• Weak electrolytes: CH₃COOH, NH₄OH

Electrolysis

Chemical decomposition using electricity. Applications: Electroplating, Metal extraction, Purification of copper, Manufacture of NaOH, Cl₂ and H₂.

Galvanic Cell

Converts chemical energy into electrical energy. Components: Anode (Oxidation), Cathode (Reduction), Salt bridge, Electrolyte. Standard electrode potential determines cell voltage.

Applications: Batteries, Fuel cells, Corrosion studies.

12. Thermodynamics

Thermodynamics deals with heat and energy changes.

• System: Part under study.
• Surroundings: Everything outside the system.

Types of Systems

• Open: Exchange matter and energy.
• Closed: Exchange only energy.
• Isolated: Exchange neither matter nor energy.

First Law

Energy can neither be created nor destroyed. ΔU = q + w

Second Law

Entropy of the universe increases.

Enthalpy (ΔH)

Heat content at constant pressure.

• Exothermic: Heat released. Examples: Combustion, Neutralization.
• Endothermic: Heat absorbed. Examples: Photosynthesis, Thermal decomposition.

13. Chemical Kinetics

Studies the speed of chemical reactions.

Factors Affecting Reaction Rate

• Concentration
• Temperature
• Catalyst
• Surface area
• Pressure (gases)

• Catalyst: Changes reaction rate without being consumed.
• Positive catalyst: MnO₂, Pt, Ni.
• Negative catalyst: Retards reaction.
• Activation energy: Minimum energy required for reaction.
• Collision theory: Effective collisions produce products.

Applications: Industrial synthesis, Food preservation, Drug manufacture.

14. Surface Chemistry

Deals with surface phenomena.

Adsorption

Accumulation of molecules on a surface. Types: Physical adsorption, Chemical adsorption.

Applications: Gas masks, Water purification, Catalysis, Chromatography.

Colloids

Particle size: 1–1000 nm. Types: Sol, Gel, Foam, Emulsion. Examples: Milk, Fog, Smoke, Butter.

• Tyndall Effect: Scattering of light by colloids.
• Brownian Movement: Random motion of particles.

15. Organic Chemistry

Organic chemistry deals with carbon compounds except CO, CO₂, carbonates, carbides etc.

Characteristics

• Covalent bonding
• Catenation
• Isomerism
• Homologous series

Functional Groups

• Alcohol –OH
• Aldehyde –CHO
• Ketone >C=O
• Carboxylic acid –COOH
• Amine –NH₂
• Ether –O–
• Ester –COO–

Isomerism

• Structural isomerism
• Geometrical isomerism
• Optical isomerism

16. Hydrocarbons

Compounds containing carbon and hydrogen only.

Alkanes

Single bonds. General formula: CₙH₂ₙ₊₂. Example: Methane, Ethane, Propane.

Alkenes

Double bond. Formula: CₙH₂ₙ. Example: Ethene.

Alkynes

Triple bond. Formula: CₙH₂ₙ₋₂. Example: Ethyne.

Aromatic Hydrocarbons

Contain benzene ring. Example: Benzene, Toluene, Naphthalene.

Important Reactions

• Combustion
• Halogenation
• Hydrogenation
• Polymerization

17. Polymers

Large molecules formed from repeating monomers.

Natural Polymers

Cellulose, Starch, Protein, Rubber, DNA.

Synthetic Polymers

Polyethylene, PVC, Nylon, Teflon, Bakelite, Polystyrene.

Types

• Addition polymerization
• Condensation polymerization

Applications: Packaging, Pipes, Textiles, Electrical insulation, Automobile parts.

18. Biomolecules

Carbohydrates

Energy-giving compounds.

• Monosaccharides: Glucose, Fructose
• Disaccharides: Sucrose, Lactose
• Polysaccharides: Starch, Cellulose, Glycogen

Proteins

Made of amino acids. Functions: Growth, Repair, Enzymes, Hormones, Antibodies.

Lipids

Store energy. Functions: Insulation, Cell membrane formation.

Nucleic Acids

• DNA: Stores genetic information.
• RNA: Protein synthesis.

19. Environmental Chemistry

Studies chemical changes affecting the environment.

Air Pollution

Major pollutants: CO, SO₂, NOₓ, Particulate matter, Ozone.

Effects: Acid rain, Smog, Global warming, Respiratory diseases.

Water Pollution

Sources: Industrial waste, Domestic sewage, Agricultural runoff.

Greenhouse Gases

CO₂, CH₄, N₂O, Water vapour, CFCs.

Ozone Layer

Protects Earth from UV radiation. Ozone depletion caused mainly by CFCs.

Acid Rain

Caused by SO₂ and NOₓ. Damages: Crops, Buildings, Lakes, Forests.

20. Industrial Chemistry

• Haber Process: N₂ + 3H₂ ⇌ 2NH₃. Uses: Fertilizers, Nitric acid manufacture.
• Contact Process: Manufacture of sulphuric acid. Catalyst: V₂O₅.
• Ostwald Process: Manufacture of nitric acid. Catalyst: Platinum-rhodium.
• Solvay Process: Manufacture of sodium carbonate.

21. Analytical Chemistry

Qualitative analysis identifies substances. Quantitative analysis determines amount.

Common Techniques

• Titration
• Gravimetric analysis
• Volumetric analysis
• Chromatography
• Spectroscopy

Applications: Pharmaceuticals, Food testing, Environmental monitoring, Forensic science.

22. Nuclear Chemistry

Radioactivity is the spontaneous emission of radiation.

Types

• Alpha (α)
• Beta (β)
• Gamma (γ)

Applications

• Cancer treatment
• Food preservation
• Power generation
• Carbon dating

Hazards

• Genetic mutation
• Radiation sickness
• Cancer

23. Corrosion

Gradual destruction of metals due to environmental reactions. Example: Rusting of iron.

Prevention

• Painting
• Galvanization
• Electroplating
• Alloy formation
• Cathodic protection

24. Important Named Laws

• Law of Conservation of Mass – Lavoisier
• Law of Constant Proportions – Proust
• Law of Multiple Proportions – Dalton
• Avogadro's Law
• Boyle's Law
• Charles' Law
• Gay-Lussac's Law
• Hess's Law
• Faraday's Laws of Electrolysis
• Raoult's Law

25. TNPSC CTS One-Liners

Study Tips

1. Learn the elements, symbols, and atomic numbers.
2. Understand acid-base, redox, and electrochemistry concepts.
3. Memorize important chemical formulas, reactions, and processes.
4. Revise the one-liners regularly for quick scoring marks.
5. Review previous year question papers to understand the pattern.