Chemistry Lecture Notes

Section A: Principles of Chemistry

1) States of Matter

1.1) Definition and Properties

• Matter: Anything with mass occupying space.
• States: Solid, liquid, gaseous, determined by physical properties (e.g., color, odour).
• Particulokinetic Theory:
  • Matter comprises particles in constant motion.
  • Balance between forces of repulsion and attraction keeps particles together.
• Evidence: Crystals, diffusion, Brownian motion, osmosis.

1.2) Properties of Different States

• Solids: Definite shape, constant volume, not compressible.
• Liquids: Variable shape, constant volume, slightly compressible.
• Gases: Variable shape and volume, easily compressible.

1.3) Changes of State

• Melting: Solid to liquid.
• Evaporation: Liquid particles escape as vapor.
• Boiling: Liquid to vapor at boiling point.
• Condensation: Gas to liquid.
• Freezing: Liquid to solid.
• Sublimation: Solid directly to gas.

2) Mixtures and Separations

2.1) Definitions

• Pure Substance: Single element/compound with constant properties.
• Mixture: Two or more substances not chemically combined.
  • Homogeneous: Same composition throughout.
  • Heterogeneous: Non-uniform composition.

2.2) Types of Mixtures

• Solutions: Solute and solvent thoroughly mixed, do not separate.
• Suspensions: Components separate over time, visible particles.
• Colloids: Intermediate particle size, do not separate on standing.

2.4) Solubility

• Solubility depends on solute and solvent nature, temperature, pressure.
• Saturated Solution: Maximum solute dissolved.

2.5) Separation Techniques

• Filtration, Sublimation, Solvent Use: To separate components based on properties.
• Distillation: To separate based on boiling points.
• Chromatography: To separate based on movement through a medium.

3) Atomic Structure

3.1) Basic Concepts

• Atoms: Smallest part showing element properties.
• Electrons move in shells around nucleus.
• Valence electrons participate in bonding.

3.2) Subatomic Particles

• Protons: Positive charge, relative mass = 1.
• Neutrons: No charge, relative mass = 1.
• Electrons: Negative charge, relative mass ≈ 0.

3.6) Isotopes

• Same element, different neutron numbers.
• Used in medicine and energy generation.

4) Periodic Table and Periodicity

4.1) General Trends

• Arrangement based on atomic number.
• Periods: Rows with increasing electron shells.
• Groups: Columns with similar properties.

4.4) Period 3 Elements

• Trends in atomic radius, electronegativity, melting/boiling points.

5) Structure and Bonding

5.1) Ionic and Covalent Bonds

• Ionic Bonds: Transfer of electrons, forming ions.
• Covalent Bonds: Sharing electrons between non-metal atoms.

5.4) Metallic Bonding

• Positive ions in a 'sea' of electrons.

5.6) Properties of Compounds

• Ionic Compounds: High melting points, conduct electricity when molten.
• Covalent Compounds: Lower melting points, don't conduct electricity.

6) Mole Concept

6.1) Definitions

• Mole: Amount of substance with Avogadro's number of particles.
• Calculations involving moles, mass, and number of particles.

7) Acids, Bases, and Salts

7.1) Properties of Acids and Bases

• Acids produce H+ ions in solution.
• Bases accept H+ ions.
• pH scale indicates acidity/alkalinity.

7.4) Reactions of Acids

• With metals, bases, carbonates.

8) Oxidation-Reduction Reactions

8.1) Bleaching and Browning

• Bleaching agents as oxidizing agents.
• Browning of fruits due to enzymatic reaction.

8.3) Oxidation Numbers

• Rules for assigning oxidation states.

9) Electrochemistry

9.1) Conductivity

• Metallic Conduction: Via free electrons.
• Electrolytic Conduction: Via ions in molten/aqueous solutions.

9.9) Applications

• Electroplating, electrorefining, and anodizing.

10) Rates of Reaction

10.1) Factors Affecting Rates

• Concentration, pressure, temperature, catalysts, surface area.

11) Energetics

11.1) Exothermic and Endothermic

• Exothermic: Releases heat, ΔH negative.
• Endothermic: Absorbs heat, ΔH positive.