Chemistry

🧬 Structure of the Atom & Atomic Theory Timeline

Early Atomic Models:

  • Democritus (400 BCE): First idea of "atomos" – indivisible particles.

Dalton (1808):

  • Matter is made of tiny, indivisible atoms.

  • Atoms combine in fixed ratios to form compounds.

J.J. Thomson (1904) – Plum Pudding Model:

  • Discovered the electron.

  • Atom is a positively charged sphere with negative electrons scattered inside.

Ernest Rutherford (1911) – Nuclear Model:

  • Gold foil experiment → discovered the nucleus.

  • Atom is mostly empty space with a dense, positive nucleus.

Niels Bohr (1913) – Bohr Model:

  • Electrons orbit the nucleus in energy levels/shells.

  • First 20 elements follow 2-8-8 electron arrangement.

Schrödinger (1926) – Quantum Mechanical Model:

  • Electrons in orbitals (clouds), not fixed paths.

James Chadwick (1932):

  • Discovered neutrons in the nucleus.

🔢 Atomic Notation & Subatomic Particles

Standard Atomic Notation:

Mass Number (A) | | Element Symbol (X) | Atomic Number (Z)

Subatomic Particles:

Particle

Charge

Location

Mass

Proton

+1

Nucleus

1 amu

Neutron

0

Nucleus

1 amu

Electron

-1

Electron cloud

~0 amu

Key Notes:

  • Atomic Number = # of protons = # of electrons (in neutral atoms)

  • Mass Number = protons + neutrons

  • Neutrons = Mass Number − Atomic Number

⚛ Bohr-Rutherford Diagrams & Lewis Structures

Bohr-Rutherford (First 20 elements):

  • 2 electrons in first shell

  • 8 in second

  • 8 in third (for the first 20 elements)

  • Protons and neutrons go in the nucleus.

Lewis Structures:

  • Only show valence electrons (outer shell electrons).

  • Dots placed around the element symbol.

🧭 Periodic Table Basics

Metals:

  • Left side of table.

  • Shiny, malleable, good conductors.

  • Lose electrons to become stable.

Non-Metals:

  • Right side.

  • Dull, brittle, poor conductors.

  • Gain electrons to become stable.

Metalloids:

  • Found in staircase pattern (B, Si, As, etc.).

  • Properties of both metals and non-metals.

Families (Groups):

Group

Key Features

Alkali Metals

Group 1 (except H), 1 valence e⁻, very reactive, reacts violently with water.

Alkaline Earth

Group 2, 2 valence e⁻, reactive, not found alone in nature.

Halogens

Group 17, 7 valence e⁻, very reactive non-metals.

Noble Gases

Group 18, full outer shell, unreactive gases.

Transition Metals

Groups 3–12, heavy, conductive, often found naturally.

Rare Earth Metals

Lanthanides + Actinides (bottom rows), very heavy, some radioactive.

📈 Trends in the Periodic Table

Across a Period (Left → Right):

  • Valence electrons increase.

  • Atoms become less metallic.

  • Atomic size decreases, reactivity of metals decreases.

Down a Group:

  • Valence shell number increases.

  • Atomic size increases.

  • Reactivity of metals increases, but non-metals decrease.

🧪 Types of Matter

Term

Definition

Atom

Smallest unit of an element.

Element

Pure substance made of one type of atom.

Molecule

Two or more atoms bonded together (can be same or different elements).

Compound

Molecule with two or more different elements.

Pure Substance

Made of only one type of particle (atom or molecule).

Mixture

Made of more than one type of particle, not chemically bonded.

🧾 Naming & Writing Formulas

Binary Molecular Compounds (Non-metal + Non-metal):

  • Use prefixes (mono-, di-, tri-, tetra-...)

  • Second element ends in “-ide”
    e.g. CO₂ = carbon dioxide

Binary Ionic Compounds (Metal + Non-metal):

  • Metal comes first (use roman numeral if it’s a transition metal).

  • Non-metal ends in “-ide”
    e.g. NaCl = sodium chloride

Formulas:

  • Use valence to balance charges
    e.g. Mg²⁺ and Cl⁻ → MgCl₂

🔥 Properties & Flame Tests

Physical Properties:

  • Colour, state, melting point, boiling point, density, solubility.

Chemical Properties:

  • Reactivity, flammability, pH, combustibility.

Flame Colour:

  • Different elements give off different colours when heated.

  • Reason: Valence electrons get “excited” → jump to higher energy → fall back → release energy as light.

⚡ Ionic Compounds & Electron Transfer

Electron Transfer:

  • Metals lose electrons → become positive ions (cations).

  • Non-metals gain electrons → become negative ions (anions).

Ionic Bond:

  • Formed between metal and non-metal by transferring electrons.

  • Results in a stable ionic compound.