T1 revision material.

Page 1: Atomic Structure Overview

  • Atoms: Basic unit of matter; consists of protons, neutrons, and electrons.

  • Subatomic Particles:

    • Protons: Positively charged; found in the nucleus.

    • Neutrons: Neutral particles; located in the nucleus.

    • Electrons: Negatively charged; orbit nucleus in energy levels (shells).

  • Atomic Number: Number of protons in an atom, defining the element (e.g., carbon = 6).

  • Mass Number: Total number of protons and neutrons (e.g., carbon with 6 protons and 6 neutrons = 12).

  • Isotopes: Same element, same protons, different neutrons (e.g., carbon-12, carbon-14).

  • Electron Configuration: Distribution of electrons in orbitals; e.g., Oxygen (atomic number 8): 1s² 2s² 2p.

  • Valence Electrons: Outermost electrons that determine chemical properties/reactivity (e.g., sodium (Na) has 1 valence electron).

  • Ions: Charged particles from losing or gaining electrons.

    • Cations: Positive ions (e.g., Na⁺).

    • Anions: Negative ions (e.g., Cl⁻).

Page 2: Key Concepts and Electron Configuration Overview

  • Nuclear Forces: Hold protons/neutrons together in the nucleus.

  • Electromagnetic Forces: Keep electrons in orbit around the nucleus.

  • Periodic Table: Arranges elements by increasing atomic number and groups similar properties.

Electron Configuration Overview

  • Energy Levels:

    • 1st shell: 2 electrons

    • 2nd shell: 8 electrons

    • 3rd shell: 18 electrons

    • 4th shell: 32 electrons.

  • Subshells:

    • s: 2 electrons

    • p: 6 electrons

    • d: 10 electrons

    • f: 14 electrons.

Page 3: Principles and Writing Electron Configurations

  • Aufbau Principle: Fill lowest energy orbitals first.

  • Pauli Exclusion Principle: No two electrons can have the same set of four quantum numbers (max 2 electrons per orbital with opposite spins).

  • Hund’s Rule: Fill degenerate orbitals singly before pairing.

Steps to Write Electron Configurations:

  1. Identify Atomic Number (number of electrons).

  2. Fill Orbitals (following Aufbau principle).

Examples:

  • Hydrogen (H): 1s¹

  • Carbon (C): 1s² 2s² 2p²

  • Oxygen (O): 1s² 2s² 2p⁶

  • Sodium (Na): 1s² 2s² 2p⁶ 3s¹

  • Chlorine (Cl): 1s² 2s² 2p⁶ 3s² 3p⁵

  • Argon (Ar): 1s² 2s² 2p⁶ 3s² 3p⁶.

Page 4: Orbital Notation and Visualization

  • Orbital Notation: Visual representation of electron configuration using arrows for electrons/spins.

Orbitals:

  • s: 2 electrons (1 orientation)

  • p: 6 electrons (3 orientations)

  • d: 10 electrons (5 orientations)

  • f: 14 electrons (7 orientations).

Example of Orbital Notation:

  • Oxygen (O):

    • Electron configuration: 1s² 2s² 2p.

    • 1s: ↑↓

    • 2s: ↑↓

    • 2p: ↑↑.

  • Carbon (C):

    • Electron configuration: 1s² 2s² 2p².

  • Sodium (Na):

    • Electron configuration: 1s² 2s² 2p 3s¹.

Page 5: Periodic Trends Overview

  • Periodic Trends: Predictable patterns in properties across the periodic table.

Key Trends:

  1. Atomic Radius:

  • Definition: Distance from nucleus to outermost shell.

  • Trend: Increases down a group, decreases across a period.

  1. Ionization Energy:

  • Definition: Energy needed to remove an electron.

  • Trend: Decreases down a group, increases across a period.

  1. Electronegativity:

  • Definition: Atom's tendency to attract electrons in a bond.

  • Trend: Decreases down a group, increases across a period.

Visualizing Trends:

  • Groups: Vertical columns (similar properties).

  • Periods: Horizontal rows (progressive property change).

Summary of Trends:

  • Atomic Radius: Increases down, does not increase across.

  • Ionization Energy: Does not increase down, increases across.

  • Electronegativity: Does not increase down, increases across.

Page 6: Ionic and Covalent Bonding

Ionic Bonding:

  • Definition: Atoms transfer electrons, forming cations/anions attracted by electrostatic forces.

  • Properties: High melting points, dissolves in water, conducts electricity when dissolved/melted.

  • Example: Sodium (Na) → Na⁺, Chlorine (Cl) → Cl⁻.

Covalent Bonding:

  • Definition: Atoms share electrons to fill outer shells.

  • Example: Water (H₂O) forms through shared electrons.

  • Properties: Low melting points, poor conductors, often liquids/gases.

Lewis Structures:

  • Shows connections between atoms in a molecule via shared electrons.

  • Example: Water (H₂O) represented as H-O-H.

Page 7: Polar and Non-Polar Molecules

Polar Molecule:

  • Definition: Molecules with a positive and negative end due to electronegativity differences.

  • Example: Water (H₂O).

Non-Polar Molecule:

  • Definition: Electrons are shared equally, no charge separation.

  • Example: O₂.

Identifying Bond Polarity:

  • Polar Bonds: Different electronegativities (e.g., H-Cl).

  • Non-Polar Bonds: Similar electronegativities (e.g., Cl-Cl).

Identifying Molecular Polarity:

  • Asymmetric shape or bonds → polar (e.g., H₂O).

  • Symmetric shape → non-polar (e.g., CO₂).

Page 8: Key Points

  1. Ionic Bonds:

  • Formed through electron transfer.

  • Charged ions, high melting, conduct electricity in solution.

  1. Covalent Bonds:

  • Formed through electron sharing.

  • No charged particles, low melting points, often gases/liquids.

  1. Polar Molecules:

  • Uneven charge distribution (e.g., H₂O).

  1. Non-Polar Molecules:

  • Even charge distribution (e.g., O₂).

Page 9: Revision Worksheet

Ionic Bonding:

  1. Define ionic bonding and the role of electrons.

  2. Formation of ions with examples.

  3. Properties of ionic compounds.

  4. Writing ionic formulas.

  5. Concept of lattice structure.

Covalent Bonding:

  1. Define covalent bonding and contrast with ionic bonding.

  2. Electron sharing with examples.

  3. Properties of covalent compounds.

  4. Drawing Lewis structures for given molecules.

  5. Define types of bonds (single, double, triple).

Polar and Non-Polar Molecules:

  1. Define polar vs non-polar molecules.

  2. Explain electronegativity and its effect on polarity with examples.

  3. Identify bond polarity of given bonds.

  4. Determine molecular polarity of examples.

  5. Draw dipole moment for a polar molecule.

Page 10: Answer Key (for teacher)

Section A: Ionic Bonding:

  • Definition: Bond formed via electron transfer, creating ions.

  • Ion Formation: Na⁺, Cl⁻, Mg²⁺, O²⁻.

  • Properties: High melting, soluble in water, conduct electricity in solution.

  • Ionic Formulas: NaCl, MgO, KBr.

  • Lattice Structure: Repeating patterns of ions contributing to strength/high melting points.

Section B: Covalent Bonding:

  • Definition: Sharing of electrons.

  • Electron Sharing: H₂O → oxygen shares with hydrogen.

  • Properties: Low melting/boiling, poor conductors, exist as gases/liquids.

  • Lewis Structures: H₂O=H-O-H, CO₂=O=C=O, CH₄=H-C-H.

  • Types of Bonds: Single (H-H), Double (O=O), Triple (N≡N).

Section C: Polar and Non-Polar Molecules:

  • Polar vs Non-Polar: Uneven vs equal charge distribution.

  • Electronegativity: Differences cause partial charges in polar bonds.

  • Bond Polarity: H-Cl = polar, N-N = non-polar, O-H = polar, Cl-Cl = non-polar.