Chemistry Paper 1 Lecture Notes

Atoms and Subatomic Particles

• Atoms: Smallest part of an element, make up all matter.
• Subatomic Particles: Protons, neutrons, and electrons.
  • Protons: Positive charge, mass of 1.
  • Neutrons: Neutral charge, mass of 1.
  • Electrons: Negative charge, negligible mass.
• Nucleus: Center of the atom containing protons and neutrons.
• Atom Size: Nucleus ~10,000 times smaller than the atom.
• Atomic Neutrality: Equal number of protons and electrons.
• Chemical Symbols: Represent elements with mass number (neutrons + protons) and atomic number (protons).

Elements and Isotopes

• Element: Substance made of one type of atom.
• Periodic Table: Contains 100+ elements.
• Isotopes: Same number of protons, different neutrons.
• Relative Atomic Mass: Accounts for isotope abundance.

Mixtures and Compounds

• Compounds: Chemically combined substances, difficult to separate.
• Mixtures: Physically combined substances, easy to separate.

Separation Methods

1. Filtration: Separates insoluble solids from liquids (e.g., sand and water).
2. Evaporation and Crystallization: Separate soluble solids from liquids.
3. Simple Distillation: Separates liquids with different boiling points.
4. Fractional Distillation: Separates more than two liquids.
5. Chromatography: Separates and analyzes components in mixtures.

History of Atomic Models

• 1803 Dalton: Atoms as solid spheres.
• 1897 JJ Thompson: Discovery of electrons; "plum pudding model."
• Rutherford's Model: Nucleus discovered; electrons around it.
• Bohr's Model: Electrons in shells/energy levels.
• Chadwick: Discovery of neutrons.

Electron Configuration and Periodic Table

• Electron Shells: Determine chemical properties.
• Periodic Table: Organized by atomic number and chemical properties.
• Groups: Elements with similar outer electron configurations.
• Periods: Represent number of electron shells.

Reactivity Trends

• Metals: More reactive down a group (easier to lose electrons).
• Non-metals: Less reactive down a group (harder to gain electrons).

Alkaline Metals (Group 1)

• Reactivity Increases: Down the group.
• Reactions:
  • With water: Forms metal hydroxides and hydrogen.
  • With chlorine: Forms metal chlorides.
  • With oxygen: Forms metal oxides.

Halogens (Group 7)

• Exist as Diatomic Molecules: E.g., F2, Cl2.
• Reactivity Decreases: Down the group.

Noble Gases (Group 0)

• Full Outer Shells: Inert, non-reactive.
• Exist as Monatomic Gases.

Ions and Ionic Bonding

• Ions: Atoms with charge (by losing/gaining electrons).
• Ionic Compounds: Metal + Non-metal, high melting/boiling points.

Covalent Bonding

• Covalent Bonds: Sharing of electrons between non-metal atoms.
• Simple Molecular Substances: Low melting/boiling points.

Metallic Bonding

• Metallic Bonding: Metals with delocalized electrons.
• Properties: High melting points, good conductors, malleable.

States of Matter

• Solid, Liquid, Gas: Different particle arrangements and energy levels.
• State Changes: Involve energy changes and particle movement.

The Mole and Chemical Calculations

• The Mole: Measurement unit, Avogadro's constant.
• Conservation of Mass: Total mass of reactants equals products.
• Calculating Moles: Moles = Mass / Molar Mass.

Reactivity Series and Metal Extraction

• Reactivity Series: Predicts reactivity with water and acids.
• Metal Extraction: Via reduction or electrolysis.

Redox Reactions

• Oxidation/Reduction: Involves electron transfer.

Acids, Bases, and pH

• pH Scale: Measures acidity/alkalinity.
• Neutralization Reactions: Acid + Base = Salt + Water.

Electrolysis

• Electrolysis Process: Separates ionic compounds using electricity.
• Applications: Metal extraction.

Energy Changes in Reactions

• Exothermic/Endothermic: Energy released to/absorbed from surroundings.

These notes summarize key concepts in chemistry, including atomic structure, the periodic table, bonding types, chemical reactions, and more, as outlined in the lecture.